A significant proportion of your second year is made up of core Chemistry modules and practical work which is common to all our Chemistry courses. The year covers more advanced concepts in chemistry through a blend of lectures, tutorials and workshops. You'll also choose one option module, allowing you to focus on a specific area in detail. Second Year Chemistry 2012.? 100 hours practical required in 2nd year. Organic & Medicinal Chemistry. Organic. Vidmantas Bieli?nas (2nd year BSc student, Vilnius University. Year One Year Two Year Three Year Four. Bsc 2nd Year Organic Chemistry Notes Created Date.
Chemistry |
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Chemistry is the scientific discipline involved with elements and compounds composed of atoms, molecules and ions: their composition, structure, properties, behavior and the changes they undergo during a reaction with other substances.[1][2]
In the scope of its subject, chemistry occupies an intermediate position between physics and biology.[3] It is sometimes called the central science because it provides a foundation for understanding both basic and applied scientific disciplines at a fundamental level.[4] For example, chemistry explains aspects of plant chemistry (botany), the formation of igneous rocks (geology), how atmospheric ozone is formed and how environmental pollutants are degraded (ecology), the properties of the soil on the moon (astrophysics), how medications work (pharmacology), and how to collect DNA evidence at a crime scene (forensics).
Chemistry addresses topics such as how atoms and molecules interact via chemical bonds to form new chemical compounds. There are four types of chemical bonds: covalent bonds, in which compounds share one or more electron(s); ionic bonds, in which a compound donates one or more electrons to another compound to produce ions (cations and anions); hydrogen bonds; and Van der Waals force bonds.
- 2Modern principles
- 2.1Matter
- 3History
- 4Practice
Etymology
The word chemistry comes from alchemy, which referred to an earlier set of practices that encompassed elements of chemistry, metallurgy, philosophy, astrology, astronomy, mysticism and medicine. It is often seen as linked to the quest to turn lead or another common starting material into gold,[5] though in ancient times the study encompassed many of the questions of modern chemistry being defined as the study of the composition of waters, movement, growth, embodying, disembodying, drawing the spirits from bodies and bonding the spirits within bodies by the early 4th century Greek-Egyptian alchemist Zosimos.[6] An alchemist was called a 'chemist' in popular speech, and later the suffix '-ry' was added to this to describe the art of the chemist as 'chemistry'.
The modern word alchemy in turn is derived from the Arabic word al-kīmīā (الكیمیاء). In origin, the term is borrowed from the Greek χημία or χημεία.[7][8] This may have Egyptian origins since al-kīmīā is derived from the Greek χημία, which is in turn derived from the word Kemet, which is the ancient name of Egypt in the Egyptian language.[7] Alternately, al-kīmīā may derive from χημεία, meaning 'cast together'.[9]
Modern principles
Laboratory, Institute of Biochemistry, University of Cologne in Germany.
The current model of atomic structure is the quantum mechanical model.[10] Traditional chemistry starts with the study of elementary particles, atoms, molecules,[11]substances, metals, crystals and other aggregates of matter. This matter can be studied in solid, liquid, or gas states, in isolation or in combination. The interactions, reactions and transformations that are studied in chemistry are usually the result of interactions between atoms, leading to rearrangements of the chemical bonds which hold atoms together. Such behaviors are studied in a chemistry laboratory.
The chemistry laboratory stereotypically uses various forms of laboratory glassware. However glassware is not central to chemistry, and a great deal of experimental (as well as applied/industrial) chemistry is done without it.
Solutions of substances in reagent bottles, including ammonium hydroxide and nitric acid, illuminated in different colors
A chemical reaction is a transformation of some substances into one or more different substances.[12] The basis of such a chemical transformation is the rearrangement of electrons in the chemical bonds between atoms. It can be symbolically depicted through a chemical equation, which usually involves atoms as subjects. The number of atoms on the left and the right in the equation for a chemical transformation is equal. (When the number of atoms on either side is unequal, the transformation is referred to as a nuclear reaction or radioactive decay.) The type of chemical reactions a substance may undergo and the energy changes that may accompany it are constrained by certain basic rules, known as chemical laws.
Energy and entropy considerations are invariably important in almost all chemical studies. Chemical substances are classified in terms of their structure, phase, as well as their chemical compositions. They can be analyzed using the tools of chemical analysis, e.g. spectroscopy and chromatography. Scientists engaged in chemical research are known as chemists.[13] Most chemists specialize in one or more sub-disciplines. Several concepts are essential for the study of chemistry; some of them are:[14]
Matter
In chemistry, matter is defined as anything that has rest mass and volume (it takes up space) and is made up of particles. The particles that make up matter have rest mass as well – not all particles have rest mass, such as the photon. Matter can be a pure chemical substance or a mixture of substances.[15]
Atom
A diagram of an atom based on the Bohr model
The atom is the basic unit of chemistry. It consists of a dense core called the atomic nucleus surrounded by a space occupied by an electron cloud. The nucleus is made up of positively charged protons and uncharged neutrons (together called nucleons), while the electron cloud consists of negatively charged electrons which orbit the nucleus. In a neutral atom, the negatively charged electrons balance out the positive charge of the protons. The nucleus is dense; the mass of a nucleon is approximately 1,836 times that of an electron, yet the radius of an atom is about 10,000 times that of its nucleus.[16][17]
The atom is also the smallest entity that can be envisaged to retain the chemical properties of the element, such as electronegativity, ionization potential, preferred oxidation state(s), coordination number, and preferred types of bonds to form (e.g., metallic, ionic, covalent).
Element
Standard form of the periodic table of chemical elements. The colors represent different categories of elements
A chemical element is a pure substance which is composed of a single type of atom, characterized by its particular number of protons in the nuclei of its atoms, known as the atomic number and represented by the symbol Z. The mass number is the sum of the number of protons and neutrons in a nucleus. Although all the nuclei of all atoms belonging to one element will have the same atomic number, they may not necessarily have the same mass number; atoms of an element which have different mass numbers are known as isotopes. For example, all atoms with 6 protons in their nuclei are atoms of the chemical element carbon, but atoms of carbon may have mass numbers of 12 or 13.[17]
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The standard presentation of the chemical elements is in the periodic table, which orders elements by atomic number. The periodic table is arranged in groups, or columns, and periods, or rows. The periodic table is useful in identifying periodic trends.[18]
Compound
Carbon dioxide (CO2), an example of a chemical compound
A compound is a pure chemical substance composed of more than one element. The properties of a compound bear little similarity to those of its elements.[19] The standard nomenclature of compounds is set by the International Union of Pure and Applied Chemistry (IUPAC). Organic compounds are named according to the organic nomenclature system.[20] The names for inorganic compounds are created according to the inorganic nomenclature system. When a compound has more than one component, then they are divided into two classes, the electropositive and the electronegative components.[21] In addition the Chemical Abstracts Service has devised a method to index chemical substances. In this scheme each chemical substance is identifiable by a number known as its CAS registry number.
Molecule
A ball-and-stick representation of the caffeine molecule (C8H10N4O2).
A molecule is the smallest indivisible portion of a pure chemical substance that has its unique set of chemical properties, that is, its potential to undergo a certain set of chemical reactions with other substances. However, this definition only works well for substances that are composed of molecules, which is not true of many substances (see below). Molecules are typically a set of atoms bound together by covalent bonds, such that the structure is electrically neutral and all valence electrons are paired with other electrons either in bonds or in lone pairs.
Thus, molecules exist as electrically neutral units, unlike ions. When this rule is broken, giving the 'molecule' a charge, the result is sometimes named a molecular ion or a polyatomic ion. However, the discrete and separate nature of the molecular concept usually requires that molecular ions be present only in well-separated form, such as a directed beam in a vacuum in a mass spectrometer. Charged polyatomic collections residing in solids (for example, common sulfate or nitrate ions) are generally not considered 'molecules' in chemistry. Some molecules contain one or more unpaired electrons, creating radicals. Most radicals are comparatively reactive, but some, such as nitric oxide (NO) can be stable.
A 2-D skeletal model of a benzene molecule (C6H6)
The 'inert' or noble gas elements (helium, neon, argon, krypton, xenon and radon) are composed of lone atoms as their smallest discrete unit, but the other isolated chemical elements consist of either molecules or networks of atoms bonded to each other in some way. Identifiable molecules compose familiar substances such as water, air, and many organic compounds like alcohol, sugar, gasoline, and the various pharmaceuticals.
However, not all substances or chemical compounds consist of discrete molecules, and indeed most of the solid substances that make up the solid crust, mantle, and core of the Earth are chemical compounds without molecules. These other types of substances, such as ionic compounds and network solids, are organized in such a way as to lack the existence of identifiable molecules per se. Instead, these substances are discussed in terms of formula units or unit cells as the smallest repeating structure within the substance. Examples of such substances are mineral salts (such as table salt), solids like carbon and diamond, metals, and familiar silica and silicate minerals such as quartz and granite.
One of the main characteristics of a molecule is its geometry often called its structure. While the structure of diatomic, triatomic or tetra-atomic molecules may be trivial, (linear, angular pyramidal etc.) the structure of polyatomic molecules, that are constituted of more than six atoms (of several elements) can be crucial for its chemical nature.
Substance and mixture
Examples of pure chemical substances. From left to right: the elements tin (Sn) and sulfur (S), diamond (an allotrope of carbon), sucrose (pure sugar), and sodium chloride (salt) and sodium bicarbonate (baking soda), which are both ionic compounds. |
A chemical substance is a kind of matter with a definite composition and set of properties.[22] A collection of substances is called a mixture. Examples of mixtures are air and alloys.[23]
Mole and amount of substance
The mole is a unit of measurement that denotes an amount of substance (also called chemical amount). The mole is defined as the number of atoms found in exactly 0.012 kilogram (or 12 grams) of carbon-12, where the carbon-12 atoms are unbound, at rest and in their ground state.[24] The number of entities per mole is known as the Avogadro constant, and is determined empirically to be approximately 6.022×1023 mol−1.[25]Molar concentration is the amount of a particular substance per volume of solution, and is commonly reported in mol/dm3.[26]
Phase
Diagram showing relationships among the phases and the terms used to describe phase changes.
In addition to the specific chemical properties that distinguish different chemical classifications, chemicals can exist in several phases. For the most part, the chemical classifications are independent of these bulk phase classifications; however, some more exotic phases are incompatible with certain chemical properties. A phase is a set of states of a chemical system that have similar bulk structural properties, over a range of conditions, such as pressure or temperature.
Physical properties, such as density and refractive index tend to fall within values characteristic of the phase. The phase of matter is defined by the phase transition, which is when energy put into or taken out of the system goes into rearranging the structure of the system, instead of changing the bulk conditions.
Sometimes the distinction between phases can be continuous instead of having a discrete boundary' in this case the matter is considered to be in a supercritical state. When three states meet based on the conditions, it is known as a triple point and since this is invariant, it is a convenient way to define a set of conditions.
The most familiar examples of phases are solids, liquids, and gases. Many substances exhibit multiple solid phases. For example, there are three phases of solid iron (alpha, gamma, and delta) that vary based on temperature and pressure. A principal difference between solid phases is the crystal structure, or arrangement, of the atoms. Another phase commonly encountered in the study of chemistry is the aqueous phase, which is the state of substances dissolved in aqueous solution (that is, in water).
Less familiar phases include plasmas, Bose–Einstein condensates and fermionic condensates and the paramagnetic and ferromagnetic phases of magnetic materials. While most familiar phases deal with three-dimensional systems, it is also possible to define analogs in two-dimensional systems, which has received attention for its relevance to systems in biology.
Bonding
An animation of the process of ionic bonding between sodium (Na) and chlorine (Cl) to form sodium chloride, or common table salt. Ionic bonding involves one atom taking valence electrons from another (as opposed to sharing, which occurs in covalent bonding)
Atoms sticking together in molecules or crystals are said to be bonded with one another. A chemical bond may be visualized as the multipole balance between the positive charges in the nuclei and the negative charges oscillating about them.[27] More than simple attraction and repulsion, the energies and distributions characterize the availability of an electron to bond to another atom.
A chemical bond can be a covalent bond, an ionic bond, a hydrogen bond or just because of Van der Waals force. Each of these kinds of bonds is ascribed to some potential. These potentials create the interactions which hold atoms together in molecules or crystals. In many simple compounds, valence bond theory, the Valence Shell Electron Pair Repulsion model (VSEPR), and the concept of oxidation number can be used to explain molecular structure and composition.
An ionic bond is formed when a metal loses one or more of its electrons, becoming a positively charged cation, and the electrons are then gained by the non-metal atom, becoming a negatively charged anion. The two oppositely charged ions attract one another, and the ionic bond is the electrostatic force of attraction between them. For example, sodium (Na), a metal, loses one electron to become an Na+ cation while chlorine (Cl), a non-metal, gains this electron to become Cl−. The ions are held together due to electrostatic attraction, and that compound sodium chloride (NaCl), or common table salt, is formed.
In the methane molecule (CH4), the carbon atom shares a pair of valence electrons with each of the four hydrogen atoms. Thus, the octet rule is satisfied for C-atom (it has eight electrons in its valence shell) and the duet rule is satisfied for the H-atoms (they have two electrons in their valence shells).
In a covalent bond, one or more pairs of valence electrons are shared by two atoms: the resulting electrically neutral group of bonded atoms is termed a molecule. Atoms will share valence electrons in such a way as to create a noble gas electron configuration (eight electrons in their outermost shell) for each atom. Atoms that tend to combine in such a way that they each have eight electrons in their valence shell are said to follow the octet rule. However, some elements like hydrogen and lithium need only two electrons in their outermost shell to attain this stable configuration; these atoms are said to follow the duet rule, and in this way they are reaching the electron configuration of the noble gas helium, which has two electrons in its outer shell.
Similarly, theories from classical physics can be used to predict many ionic structures. With more complicated compounds, such as metal complexes, valence bond theory is less applicable and alternative approaches, such as the molecular orbital theory, are generally used. See diagram on electronic orbitals.
Energy
In the context of chemistry, energy is an attribute of a substance as a consequence of its atomic, molecular or aggregate structure. Since a chemical transformation is accompanied by a change in one or more of these kinds of structures, it is invariably accompanied by an increase or decrease of energy of the substances involved. Some energy is transferred between the surroundings and the reactants of the reaction in the form of heat or light; thus the products of a reaction may have more or less energy than the reactants.
A reaction is said to be exergonic if the final state is lower on the energy scale than the initial state; in the case of endergonic reactions the situation is the reverse. A reaction is said to be exothermic if the reaction releases heat to the surroundings; in the case of endothermic reactions, the reaction absorbs heat from the surroundings.
Chemical reactions are invariably not possible unless the reactants surmount an energy barrier known as the activation energy. The speed of a chemical reaction (at given temperature T) is related to the activation energy E, by the Boltzmann's population factor – that is the probability of a molecule to have energy greater than or equal to E at the given temperature T. This exponential dependence of a reaction rate on temperature is known as the Arrhenius equation.The activation energy necessary for a chemical reaction to occur can be in the form of heat, light, electricity or mechanical force in the form of ultrasound.[28]
A related concept free energy, which also incorporates entropy considerations, is a very useful means for predicting the feasibility of a reaction and determining the state of equilibrium of a chemical reaction, in chemical thermodynamics. A reaction is feasible only if the total change in the Gibbs free energy is negative, ; if it is equal to zero the chemical reaction is said to be at equilibrium.
There exist only limited possible states of energy for electrons, atoms and molecules. These are determined by the rules of quantum mechanics, which require quantization of energy of a bound system. The atoms/molecules in a higher energy state are said to be excited. The molecules/atoms of substance in an excited energy state are often much more reactive; that is, more amenable to chemical reactions.
The phase of a substance is invariably determined by its energy and the energy of its surroundings. When the intermolecular forces of a substance are such that the energy of the surroundings is not sufficient to overcome them, it occurs in a more ordered phase like liquid or solid as is the case with water (H2O); a liquid at room temperature because its molecules are bound by hydrogen bonds.[29] Whereas hydrogen sulfide (H2S) is a gas at room temperature and standard pressure, as its molecules are bound by weaker dipole-dipole interactions.
The transfer of energy from one chemical substance to another depends on the size of energy quanta emitted from one substance. However, heat energy is often transferred more easily from almost any substance to another because the phonons responsible for vibrational and rotational energy levels in a substance have much less energy than photons invoked for the electronic energy transfer. Thus, because vibrational and rotational energy levels are more closely spaced than electronic energy levels, heat is more easily transferred between substances relative to light or other forms of electronic energy. For example, ultraviolet electromagnetic radiation is not transferred with as much efficacy from one substance to another as thermal or electrical energy.
The existence of characteristic energy levels for different chemical substances is useful for their identification by the analysis of spectral lines. Different kinds of spectra are often used in chemical spectroscopy, e.g. IR, microwave, NMR, ESR, etc. Spectroscopy is also used to identify the composition of remote objects – like stars and distant galaxies – by analyzing their radiation spectra.
Emission spectrum of iron
The term chemical energy is often used to indicate the potential of a chemical substance to undergo a transformation through a chemical reaction or to transform other chemical substances.
Reaction
During chemical reactions, bonds between atoms break and form, resulting in different substances with different properties. In a blast furnace, iron oxide, a compound, reacts with carbon monoxide to form iron, one of the chemical elements, and carbon dioxide.
When a chemical substance is transformed as a result of its interaction with another substance or with energy, a chemical reaction is said to have occurred. A chemical reaction is therefore a concept related to the 'reaction' of a substance when it comes in close contact with another, whether as a mixture or a solution; exposure to some form of energy, or both. It results in some energy exchange between the constituents of the reaction as well as with the system environment, which may be designed vessels—often laboratory glassware.
Chemical reactions can result in the formation or dissociation of molecules, that is, molecules breaking apart to form two or more molecules or rearrangement of atoms within or across molecules. Chemical reactions usually involve the making or breaking of chemical bonds. Oxidation, reduction, dissociation, acid-base neutralization and molecular rearrangement are some of the commonly used kinds of chemical reactions.
A chemical reaction can be symbolically depicted through a chemical equation. While in a non-nuclear chemical reaction the number and kind of atoms on both sides of the equation are equal, for a nuclear reaction this holds true only for the nuclear particles viz. protons and neutrons.[30]
The sequence of steps in which the reorganization of chemical bonds may be taking place in the course of a chemical reaction is called its mechanism. A chemical reaction can be envisioned to take place in a number of steps, each of which may have a different speed. Many reaction intermediates with variable stability can thus be envisaged during the course of a reaction. Reaction mechanisms are proposed to explain the kinetics and the relative product mix of a reaction. Many physical chemists specialize in exploring and proposing the mechanisms of various chemical reactions. Several empirical rules, like the Woodward–Hoffmann rules often come in handy while proposing a mechanism for a chemical reaction.
According to the IUPAC gold book, a chemical reaction is 'a process that results in the interconversion of chemical species.'[31] Accordingly, a chemical reaction may be an elementary reaction or a stepwise reaction. An additional caveat is made, in that this definition includes cases where the interconversion of conformers is experimentally observable. Such detectable chemical reactions normally involve sets of molecular entities as indicated by this definition, but it is often conceptually convenient to use the term also for changes involving single molecular entities (i.e. 'microscopic chemical events').
Ions and salts
The crystal lattice structure of potassium chloride (KCl), a salt which is formed due to the attraction of K+ cations and Cl− anions. Note how the overall charge of the ionic compound is zero.
An ion is a charged species, an atom or a molecule, that has lost or gained one or more electrons. When an atom loses an electron and thus has more protons than electrons, the atom is a positively charged ion or cation. When an atom gains an electron and thus has more electrons than protons, the atom is a negatively charged ion or anion. Cations and anions can form a crystalline lattice of neutral salts, such as the Na+ and Cl− ions forming sodium chloride, or NaCl. Examples of polyatomic ions that do not split up during acid-base reactions are hydroxide (OH−) and phosphate (PO43−).
Plasma is composed of gaseous matter that has been completely ionized, usually through high temperature.
Acidity and basicity
When hydrogen bromide (HBr), pictured, is dissolved in water, it forms the strong acid hydrobromic acid
A substance can often be classified as an acid or a base. There are several different theories which explain acid-base behavior. The simplest is Arrhenius theory, which states that acid is a substance that produces hydronium ions when it is dissolved in water, and a base is one that produces hydroxide ions when dissolved in water. According to Brønsted–Lowry acid-base theory, acids are substances that donate a positive hydrogenion to another substance in a chemical reaction; by extension, a base is the substance which receives that hydrogen ion.
A third common theory is Lewis acid-base theory, which is based on the formation of new chemical bonds. Lewis theory explains that an acid is a substance which is capable of accepting a pair of electrons from another substance during the process of bond formation, while a base is a substance which can provide a pair of electrons to form a new bond. According to this theory, the crucial things being exchanged are charges.[32] There are several other ways in which a substance may be classified as an acid or a base, as is evident in the history of this concept.[33]
Acid strength is commonly measured by two methods. One measurement, based on the Arrhenius definition of acidity, is pH, which is a measurement of the hydronium ion concentration in a solution, as expressed on a negative logarithmic scale. Thus, solutions that have a low pH have a high hydronium ion concentration and can be said to be more acidic. The other measurement, based on the Brønsted–Lowry definition, is the acid dissociation constant (Ka), which measures the relative ability of a substance to act as an acid under the Brønsted–Lowry definition of an acid. That is, substances with a higher Ka are more likely to donate hydrogen ions in chemical reactions than those with lower Ka values.
Redox
Redox (reduction-oxidation) reactions include all chemical reactions in which atoms have their oxidation state changed by either gaining electrons (reduction) or losing electrons (oxidation). Substances that have the ability to oxidize other substances are said to be oxidative and are known as oxidizing agents, oxidants or oxidizers. An oxidant removes electrons from another substance. Similarly, substances that have the ability to reduce other substances are said to be reductive and are known as reducing agents, reductants, or reducers.
A reductant transfers electrons to another substance and is thus oxidized itself. And because it 'donates' electrons it is also called an electron donor. Oxidation and reduction properly refer to a change in oxidation number—the actual transfer of electrons may never occur. Thus, oxidation is better defined as an increase in oxidation number, and reduction as a decrease in oxidation number.
Equilibrium
Although the concept of equilibrium is widely used across sciences, in the context of chemistry, it arises whenever a number of different states of the chemical composition are possible, as for example, in a mixture of several chemical compounds that can react with one another, or when a substance can be present in more than one kind of phase.
A system of chemical substances at equilibrium, even though having an unchanging composition, is most often not static; molecules of the substances continue to react with one another thus giving rise to a dynamic equilibrium. Thus the concept describes the state in which the parameters such as chemical composition remain unchanged over time.
Chemical laws
Chemical reactions are governed by certain laws, which have become fundamental concepts in chemistry. Some of them are:
- Boyle's law (1662, relating pressure and volume)
- Charles's law (1787, relating volume and temperature)
- Gay-Lussac's law (1809, relating pressure and temperature)
- Law of conservation of energy leads to the important concepts of equilibrium, thermodynamics, and kinetics.
- Law of conservation of mass continues to be conserved in isolated systems, even in modern physics. However, special relativity shows that due to mass–energy equivalence, whenever non-material 'energy' (heat, light, kinetic energy) is removed from a non-isolated system, some mass will be lost with it. High energy losses result in loss of weighable amounts of mass, an important topic in nuclear chemistry.
- Law of definite composition, although in many systems (notably biomacromolecules and minerals) the ratios tend to require large numbers, and are frequently represented as a fraction.
History
The history of chemistry spans a period from very old times to the present. Since several millennia BC, civilizations were using technologies that would eventually form the basis of the various branches of chemistry. Examples include extracting metals from ores, making pottery and glazes, fermenting beer and wine, extracting chemicals from plants for medicine and perfume, rendering fat into soap, making glass, and making alloys like bronze. Chemistry was preceded by its protoscience, alchemy, which is an intuitive but non-scientific approach to understanding the constituents of matter and their interactions. It was unsuccessful in explaining the nature of matter and its transformations, but, by performing experiments and recording the results, alchemists set the stage for modern chemistry. Chemistry as a body of knowledge distinct from alchemy began to emerge when a clear differentiation was made between them by Robert Boyle in his work The Sceptical Chymist (1661). While both alchemy and chemistry are concerned with matter and its transformations, the crucial difference was given by the scientific method that chemists employed in their work. Chemistry is considered to have become an established science with the work of Antoine Lavoisier, who developed a law of conservation of mass that demanded careful measurement and quantitative observations of chemical phenomena. The history of chemistry is intertwined with the history of thermodynamics, especially through the work of Willard Gibbs.[34]
Of definition
The definition of chemistry has changed over time, as new discoveries and theories add to the functionality of the science. The term 'chymistry', in the view of noted scientist Robert Boyle in 1661, meant the subject of the material principles of mixed bodies.[35] In 1663, the chemist Christopher Glaser described 'chymistry' as a scientific art, by which one learns to dissolve bodies, and draw from them the different substances on their composition, and how to unite them again, and exalt them to a higher perfection.[36]
The 1730 definition of the word 'chemistry', as used by Georg Ernst Stahl, meant the art of resolving mixed, compound, or aggregate bodies into their principles; and of composing such bodies from those principles.[37] In 1837, Jean-Baptiste Dumas considered the word 'chemistry' to refer to the science concerned with the laws and effects of molecular forces.[38] This definition further evolved until, in 1947, it came to mean the science of substances: their structure, their properties, and the reactions that change them into other substances – a characterization accepted by Linus Pauling.[39] More recently, in 1998, Professor Raymond Chang broadened the definition of 'chemistry' to mean the study of matter and the changes it undergoes.[40]
Of discipline
Democritus' atomist philosophy was later adopted by Epicurus (341–270 BCE).
Early civilizations, such as the Egyptians[41]Babylonians, Indians[42] amassed practical knowledge concerning the arts of metallurgy, pottery and dyes, but didn't develop a systematic theory.
A basic chemical hypothesis first emerged in Classical Greece with the theory of four elements as propounded definitively by Aristotle stating that fire, air, earth and water were the fundamental elements from which everything is formed as a combination. Greekatomism dates back to 440 BC, arising in works by philosophers such as Democritus and Epicurus. In 50 BCE, the Roman philosopher Lucretius expanded upon the theory in his book De rerum natura (On The Nature of Things).[43][44] Unlike modern concepts of science, Greek atomism was purely philosophical in nature, with little concern for empirical observations and no concern for chemical experiments.[45]
An early form of the idea of conservation of mass is the notion that 'Nothing comes from nothing' in Ancient Greek philosophy, which can be found in Empedocles (approx. 4th century BC): 'For it is impossible for anything to come to be from what is not, and it cannot be brought about or heard of that what is should be utterly destroyed.'[46] and Epicurus (3rd century BC), who, describing the nature of the Universe, wrote that 'the totality of things was always such as it is now, and always will be'.[47]
In the Hellenistic world the art of alchemy first proliferated, mingling magic and occultism into the study of natural substances with the ultimate goal of transmuting elements into gold and discovering the elixir of eternal life.[48] Work, particularly the development of distillation, continued in the early Byzantine period with the most famous practitioner being the 4th century Greek-Egyptian Zosimos of Panopolis.[49] Alchemy continued to be developed and practised throughout the Arab world after the Muslim conquests,[50] and from there, and from the Byzantine remnants,[51] diffused into medieval and Renaissance Europe through Latin translations.
Jābir ibn Hayyān (Geber), a Perso-Arab alchemist whose experimental research laid the foundations of chemistry.
The development of the modern scientific method was slow and arduous, but an early scientific method for chemistry began emerging among early Muslim chemists, beginning with the 9th century Perso-Arab chemist Jābir ibn Hayyān (known as 'Geber' in Europe), who is sometimes referred to as 'the father of chemistry'.[52][53][54][55] He introduced a systematic and experimental approach to scientific research based in the laboratory, in contrast to the ancient Greek and Egyptian alchemists whose works were largely allegorical and often unintelligible.[56] He also introduced the alembic (al-anbiq) of Persian encyclopedist Ibn al-Awwam to Europe, chemically analyzed many chemical substances, composed lapidaries, distinguished between alkalis and acids, and manufactured hundreds of drugs.[57] His books strongly influenced the medieval European alchemists[58] and justified their search for the philosopher's stone.[59][60]In the Middle Ages, Jabir's treatises on alchemy were translated into Latin and became standard texts for European alchemists. These include the Kitab al-Kimya (titled Book of the Composition of Alchemy in Europe), translated by Robert of Chester (1144); and the Kitab al-Sab'een (Book of Seventy) by Gerard of Cremona (before 1187). Later influential Muslim philosophers, such as Abū al-Rayhān al-Bīrūnī,[61]Avicenna[62] and Al-Kindi disputed the theories of alchemy, particularly the theory of the transmutation of metals; and al-Tusi described a version of the conservation of mass, noting that a body of matter is able to change but is not able to disappear.[63]
Under the influence of the new empirical methods propounded by Sir Francis Bacon and others, a group of chemists at Oxford, Robert Boyle, Robert Hooke and John Mayow began to reshape the old alchemical traditions into a scientific discipline. Boyle in particular is regarded as the founding father of chemistry due to his most important work, the classic chemistry text The Sceptical Chymist where the differentiation is made between the claims of alchemy and the empirical scientific discoveries of the new chemistry.[64] He formulated Boyle's law, rejected the classical 'four elements' and proposed a mechanistic alternative of atoms and chemical reactions that could be subject to rigorous experiment.[65]
Antoine-Laurent de Lavoisier is considered the 'Father of Modern Chemistry'.[66]
The theory of phlogiston (a substance at the root of all combustion) was propounded by the German Georg Ernst Stahl in the early 18th century and was only overturned by the end of the century by the French chemist Antoine Lavoisier, the chemical analogue of Newton in physics; who did more than any other to establish the new science on proper theoretical footing, by elucidating the principle of conservation of mass and developing a new system of chemical nomenclature used to this day.[67]
Before his work, though, many important discoveries had been made, specifically relating to the nature of 'air' which was discovered to be composed of many different gases. The Scottish chemist Joseph Black (the first experimental chemist) and the Dutchman J.B. van Helmont discovered carbon dioxide, or what Black called 'fixed air' in 1754; Henry Cavendish discovered hydrogen and elucidated its properties and Joseph Priestley and, independently, Carl Wilhelm Scheele isolated pure oxygen.
In his periodic table, Dmitri Mendeleev predicted the existence of 7 new elements,[68] and placed all 60 elements known at the time in their correct places.[69]
English scientist John Dalton proposed the modern theory of atoms; that all substances are composed of indivisible 'atoms' of matter and that different atoms have varying atomic weights.
The development of the electrochemical theory of chemical combinations occurred in the early 19th century as the result of the work of two scientists in particular, J.J. Berzelius and Humphry Davy, made possible by the prior invention of the voltaic pile by Alessandro Volta. Davy discovered nine new elements including the alkali metals by extracting them from their oxides with electric current.[70]
British William Prout first proposed ordering all the elements by their atomic weight as all atoms had a weight that was an exact multiple of the atomic weight of hydrogen. J.A.R. Newlands devised an early table of elements, which was then developed into the modern periodic table of elements[71] in the 1860s by Dmitri Mendeleev and independently by several other scientists including Julius Lothar Meyer.[72][73] The inert gases, later called the noble gases were discovered by William Ramsay in collaboration with Lord Rayleigh at the end of the century, thereby filling in the basic structure of the table.
Top: Expected results: alpha particles passing through the plum pudding model of the atom undisturbed.
Bottom: Observed results: a small portion of the particles were deflected, indicating a small, concentrated charge.
Bottom: Observed results: a small portion of the particles were deflected, indicating a small, concentrated charge.
At the turn of the twentieth century the theoretical underpinnings of chemistry were finally understood due to a series of remarkable discoveries that succeeded in probing and discovering the very nature of the internal structure of atoms. In 1897, J.J. Thomson of Cambridge University discovered the electron and soon after the French scientist Becquerel as well as the couple Pierre and Marie Curie investigated the phenomenon of radioactivity. In a series of pioneering scattering experiments Ernest Rutherford at the University of Manchester discovered the internal structure of the atom and the existence of the proton, classified and explained the different types of radioactivity and successfully transmuted the first element by bombarding nitrogen with alpha particles.
His work on atomic structure was improved on by his students, the Danish physicist Niels Bohr and Henry Moseley. The electronic theory of chemical bonds and molecular orbitals was developed by the American scientists Linus Pauling and Gilbert N. Lewis.
The year 2011 was declared by the United Nations as the International Year of Chemistry.[74] It was an initiative of the International Union of Pure and Applied Chemistry, and of the United Nations Educational, Scientific, and Cultural Organization and involves chemical societies, academics, and institutions worldwide and relied on individual initiatives to organize local and regional activities.
Organic chemistry was developed by Justus von Liebig and others, following Friedrich Wöhler's synthesis of urea which proved that living organisms were, in theory, reducible to chemistry.[75] Other crucial 19th century advances were; an understanding of valence bonding (Edward Frankland in 1852) and the application of thermodynamics to chemistry (J. W. Gibbs and Svante Arrhenius in the 1870s).
Practice
Subdisciplines
Chemistry is typically divided into several major sub-disciplines. There are also several main cross-disciplinary and more specialized fields of chemistry.[76]
- Analytical chemistry is the analysis of material samples to gain an understanding of their chemical composition and structure. Analytical chemistry incorporates standardized experimental methods in chemistry. These methods may be used in all subdisciplines of chemistry, excluding purely theoretical chemistry.
- Biochemistry is the study of the chemicals, chemical reactions and chemical interactions that take place in living organisms. Biochemistry and organic chemistry are closely related, as in medicinal chemistry or neurochemistry. Biochemistry is also associated with molecular biology and genetics.
- Inorganic chemistry is the study of the properties and reactions of inorganic compounds. The distinction between organic and inorganic disciplines is not absolute and there is much overlap, most importantly in the sub-discipline of organometallic chemistry.
- Materials chemistry is the preparation, characterization, and understanding of substances with a useful function. The field is a new breadth of study in graduate programs, and it integrates elements from all classical areas of chemistry with a focus on fundamental issues that are unique to materials. Primary systems of study include the chemistry of condensed phases (solids, liquids, polymers) and interfaces between different phases.
- Neurochemistry is the study of neurochemicals; including transmitters, peptides, proteins, lipids, sugars, and nucleic acids; their interactions, and the roles they play in forming, maintaining, and modifying the nervous system.
- Nuclear chemistry is the study of how subatomic particles come together and make nuclei. Modern Transmutation is a large component of nuclear chemistry, and the table of nuclides is an important result and tool for this field.
- Organic chemistry is the study of the structure, properties, composition, mechanisms, and reactions of organic compounds. An organic compound is defined as any compound based on a carbon skeleton.
- Physical chemistry is the study of the physical and fundamental basis of chemical systems and processes. In particular, the energetics and dynamics of such systems and processes are of interest to physical chemists. Important areas of study include chemical thermodynamics, chemical kinetics, electrochemistry, statistical mechanics, spectroscopy, and more recently, astrochemistry.[77] Physical chemistry has large overlap with molecular physics. Physical chemistry involves the use of infinitesimal calculus in deriving equations. It is usually associated with quantum chemistry and theoretical chemistry. Physical chemistry is a distinct discipline from chemical physics, but again, there is very strong overlap.
- Theoretical chemistry is the study of chemistry via fundamental theoretical reasoning (usually within mathematics or physics). In particular the application of quantum mechanics to chemistry is called quantum chemistry. Since the end of the Second World War, the development of computers has allowed a systematic development of computational chemistry, which is the art of developing and applying computer programs for solving chemical problems. Theoretical chemistry has large overlap with (theoretical and experimental) condensed matter physics and molecular physics.
Other disciplines within chemistry are traditionally grouped by the type of matter being studied or the kind of study. These include inorganic chemistry, the study of inorganic matter; organic chemistry, the study of organic (carbon-based) matter; biochemistry, the study of substances found in biological organisms; physical chemistry, the study of chemical processes using physical concepts such as thermodynamics and quantum mechanics; and analytical chemistry, the analysis of material samples to gain an understanding of their chemical composition and structure. Many more specialized disciplines have emerged in recent years, e.g. neurochemistry the chemical study of the nervous system (see subdisciplines).
Other fields include agrochemistry, astrochemistry (and cosmochemistry), atmospheric chemistry, chemical engineering, chemical biology, chemo-informatics, electrochemistry, environmental chemistry, femtochemistry, flavor chemistry, flow chemistry, geochemistry, green chemistry, histochemistry, history of chemistry, hydrogenation chemistry, immunochemistry, marine chemistry, materials science, mathematical chemistry, mechanochemistry, medicinal chemistry, molecular biology, molecular mechanics, nanotechnology, natural product chemistry, oenology, organometallic chemistry, petrochemistry, pharmacology, photochemistry, physical organic chemistry, phytochemistry, polymer chemistry, radiochemistry, solid-state chemistry, sonochemistry, supramolecular chemistry, surface chemistry, synthetic chemistry, thermochemistry, and many others.
Industry
The chemical industry represents an important economic activity worldwide. The global top 50 chemical producers in 2013 had sales of US$980.5 billion with a profit margin of 10.3%.[78]
Professional societies
See also
References
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- ^'chemical bonding'. Britannica. Encyclopædia Britannica. Retrieved 1 November 2012.
- ^Matter: Atoms from Democritus to Dalton by Anthony Carpi, Ph.D.
- ^IUPAC Gold BookDefinition
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- ^Armstrong, James (2012). General, Organic, and Biochemistry: An Applied Approach. Brooks/Cole. p. 48. ISBN978-0-534-49349-3.
- ^Burrows et al. 2008, p. 13.
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- ^Burrows et al. 2008, p. 12.
- ^'IUPAC Nomenclature of Organic Chemistry'. Acdlabs.com. Retrieved 2011-06-12.
- ^Connelly, Neil G.; Damhus, Ture; Hartshorn, Richard M.; Hutton, Alan T. (2005). Nomenclature of Inorganic Chemistry IUPAC Recommendations 2005. RSCPublishing. pp. 5–12. ISBN978-0-85404-438-2.
- ^Hill, J.W.; Petrucci, R.H.; McCreary, T.W.; Perry, S.S. (2005). General Chemistry (4th ed.). Upper Saddle River, New Jersey: Pearson Prentice Hall. p. 37.
- ^M.M. Avedesian; Hugh Baker. Magnesium and Magnesium Alloys. ASM International. p. 59.
- ^'Official SI Unit definitions'. Bipm.org. Retrieved 2011-06-12.
- ^Burrows et al. 2008, p. 16.
- ^Atkins & de Paula 2009, p. 9.
- ^Visionlearning. 'Chemical Bonding by Anthony Carpi, Ph'. visionlearning. Retrieved 2011-06-12.
- ^Reilly, Michael. (2007). Mechanical force induces chemical reaction, NewScientist.com news service, Reilly
- ^Changing States of Matter – Chemforkids.com
- ^Chemical Reaction Equation – IUPAC Goldbook
- ^Gold BookChemical Reaction IUPAC Goldbook
- ^'The Lewis Acid-Base Concept'. Apsidium. May 19, 2003. Archived from the original on 2008-05-27. Retrieved 2010-07-31.[unreliable source?]
- ^'History of Acidity'. Bbc.co.uk. 2004-05-27. Retrieved 2011-06-12.
- ^Selected Classic Papers from the History of Chemistry
- ^Boyle, Robert (1661). The Sceptical Chymist. New York: Dover Publications, Inc. (reprint). ISBN978-0-486-42825-3.
- ^Glaser, Christopher (1663). Traite de la chymie. Paris. as found in: Kim, Mi Gyung (2003). Affinity, That Elusive Dream – A Genealogy of the Chemical Revolution. The MIT Press. ISBN978-0-262-11273-4.
- ^Stahl, George, E. (1730). Philosophical Principles of Universal Chemistry. London.
- ^Dumas, J.B. (1837). 'Affinite' (lecture notes), vii, p 4. 'Statique chimique', Paris: Académie des Sciences
- ^Pauling, Linus (1947). General Chemistry. Dover Publications, Inc. ISBN978-0-486-65622-9.
- ^Chang, Raymond (1998). Chemistry, 6th Ed. New York: McGraw Hill. ISBN978-0-07-115221-1.
- ^First chemists, February 13, 1999, New Scientist
- ^Barnes, Ruth. Textiles in Indian Ocean Societies. Routledge. p. 1.
- ^Lucretius (50 BCE). 'de Rerum Natura (On the Nature of Things)'. The Internet Classics Archive. Massachusetts Institute of Technology. Retrieved 9 January 2007.Check date values in:
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(help) - ^Simpson, David (29 June 2005). 'Lucretius (c. 99–55 BCE)'. The Internet History of Philosophy. Retrieved 2007-01-09.
- ^Strodach, George K. (2012). The Art of Happiness. New York: Penguin Classics. pp. 7–8. ISBN978-0-14-310721-7.
- ^Fr. 12; see pp.291–2 of Kirk, G. S.; J. E. Raven; Malcolm Schofield (1983). The Presocratic Philosophers (2 ed.). Cambridge: Cambridge University Press. ISBN978-0-521-27455-5.
- ^Long, A. A.; D. N. Sedley (1987). 'Epicureanism: The principals of conservation'. The Hellenistic Philosophers. Vol 1: Translations of the principal sources with philosophical commentary. Cambridge: Cambridge University Press. pp. 25–26. ISBN978-0-521-27556-9.
- ^'International Year of Chemistry – The History of Chemistry'. G.I.T. Laboratory Journal Europe. Feb 25, 2011. Retrieved March 12, 2013.
- ^Bryan H. Bunch & Alexander Hellemans (2004). The History of Science and Technology. Houghton Mifflin Harcourt. p. 88. ISBN978-0-618-22123-3.
- ^Morris Kline (1985) Mathematics for the nonmathematician. Courier Dover Publications. p. 284. ISBN0-486-24823-2
- ^Marcelin Berthelot, Collection des anciens alchimistes grecs (3 vol., Paris, 1887–1888, p. 161); F. Sherwood Taylor, 'The Origins of Greek Alchemy,' Ambix 1 (1937), 40.
- ^Derewenda, Zygmunt S.; Derewenda, ZS (2007). 'On wine, chirality and crystallography'. Acta Crystallographica Section A. 64 (Pt 1): 246–258 [247]. Bibcode:2008AcCrA..64..246D. doi:10.1107/S0108767307054293. PMID18156689.
- ^John Warren (2005). 'War and the Cultural Heritage of Iraq: a sadly mismanaged affair', Third World Quarterly, Volume 26, Issue 4 & 5, pp. 815–830.
- ^Dr. A. Zahoor (1997), Jâbir ibn Hayyân (Geber)
- ^Paul Vallely, How Islamic inventors changed the world, The Independent, 10 March 2006
- ^Kraus, Paul, Jâbir ibn Hayyân, Contribution à l'histoire des idées scientifiques dans l'Islam. I. Le corpus des écrits jâbiriens. II. Jâbir et la science grecque,. Cairo (1942–1943). Repr. By Fuat Sezgin, (Natural Sciences in Islam. 67–68), Frankfurt. 2002:To form an idea of the historical place of Jabir's alchemy and to tackle the problem of its sources, it is advisable to compare it with what remains to us of the alchemical literature in the Greek language. One knows in which miserable state this literature reached us. Collected by Byzantine scientists from the tenth century, the corpus of the Greek alchemists is a cluster of incoherent fragments, going back to all the times since the third century until the end of the Middle Ages.'The efforts of Berthelot and Ruelle to put a little order in this mass of literature led only to poor results, and the later researchers, among them in particular Mrs. Hammer-Jensen, Tannery, Lagercrantz, von Lippmann, Reitzenstein, Ruska, Bidez, Festugiere and others, could make clear only few points of detail…The study of the Greek alchemists is not very encouraging. An even surface examination of the Greek texts shows that a very small part only was organized according to true experiments of laboratory: even the supposedly technical writings, in the state where we find them today, are unintelligible nonsense which refuses any interpretation.It is different with Jabir's alchemy. The relatively clear description of the processes and the alchemical apparatuses, the methodical classification of the substances, mark an experimental spirit which is extremely far away from the weird and odd esotericism of the Greek texts. The theory on which Jabir supports his operations is one of clearness and of an impressive unity. More than with the other Arab authors, one notes with him a balance between theoretical teaching and practical teaching, between the `ilm and the `amal. In vain one would seek in the Greek texts a work as systematic as that which is presented for example in the Book of Seventy.'(cf.Ahmad Y Hassan. 'A Critical Reassessment of the Geber Problem: Part Three'. Archived from the original on 2008-11-20. Retrieved 2008-08-09.)
- ^Will Durant (1980). The Age of Faith (The Story of Civilization, Volume 4), p. 162-186. Simon & Schuster. ISBN0-671-01200-2.
- ^Cite error: The named reference
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was invoked but never defined (see the help page). - ^Ragai, Jehane (1992). 'The Philosopher's Stone: Alchemy and Chemistry'. Journal of Comparative Poetics. 12 (Metaphor and Allegory in the Middle Ages): 58–77. doi:10.2307/521636. JSTOR521636.
- ^Holmyard, E. J. (1924). 'Maslama al-Majriti and the Rutbatu'l-Hakim'. Isis. 6 (3): 293–305. doi:10.1086/358238.
- ^Marmura, Michael E.; Nasr, Seyyed Hossein (1965). 'An Introduction to Islamic Cosmological Doctrines. Conceptions of Nature and Methods Used for Its Study by the Ikhwan Al-Safa'an, Al-Biruni, and Ibn Sina by Seyyed Hossein Nasr'. Speculum. 40 (4): 744–746. doi:10.2307/2851429. JSTOR2851429.
- ^Robert Briffault (1938). The Making of Humanity, pp. 196–197.
- ^Alakbarov, Farid (2001). 'A 13th-Century Darwin? Tusi's Views on Evolution'. Azerbaijan International. 9: 2.
- ^'Robert Boyle, Founder of Modern Chemistry' Harry Sootin (2011)
- ^'History – Robert Boyle (1627–1691)'. BBC. Retrieved 2011-06-12.
- ^Eagle, Cassandra T.; Jennifer Sloan (1998). 'Marie Anne Paulze Lavoisier: The Mother of Modern Chemistry'. The Chemical Educator. 3 (5): 1–18. doi:10.1007/s00897980249a.
- ^Mi Gyung Kim (2003). Affinity, that Elusive Dream: A Genealogy of the Chemical Revolution. MIT Press. p. 440. ISBN978-0-262-11273-4.
- ^Chemistry 412 course notes. 'A Brief History of the Development of Periodic Table'. Western Oregon University. Retrieved July 20, 2015.
- ^Note: '...it is surely true that had Mendeleev never lived modern chemists would be using a Periodic Table' and 'Dmitri Mendeleev'. Royal Society of Chemistry. Retrieved July 18, 2015.
- ^Davy, Humphry (1808). 'On some new Phenomena of Chemical Changes produced by Electricity, particularly the Decomposition of the fixed Alkalies, and the Exhibition of the new Substances, which constitute their Bases'. Philosophical Transactions of the Royal Society. 98: 1–45. doi:10.1098/rstl.1808.0001.
- ^Winter, Mark. 'WebElements: the periodic table on the web'. The University of Sheffield. Archived from the original on January 4, 2014. Retrieved January 27, 2014.
- ^'Julius Lothar Meyer and Dmitri Ivanovich Mendeleev'. Science History Institute. June 2016. Retrieved March 20, 2018.
- ^'What makes these family likenesses among the elements? In the 1860s everyone was scratching their heads about that, and several scientists moved towards rather similar answers. The man who solved the problem most triumphantly was a young Russian called Dmitri Ivanovich Mendeleev, who visited the salt mine at Wieliczka in 1859.' Bronowski, Jacob (1973). The Ascent of Man. Little, Brown and Company. p. 322. ISBN978-0-316-10930-7.
- ^'Chemistry'. Chemistry2011.org. Retrieved 2012-03-10.
- ^Ihde, Aaron John (1984). The Development of Modern Chemistry. Courier Dover Publications. p. 164. ISBN978-0-486-64235-2.
- ^W.G. Laidlaw; D.E. Ryan And Gary Horlick; H.C. Clark, Josef Takats, And Martin Cowie; R.U. Lemieux (1986-12-10). 'Chemistry Subdisciplines'. The Canadian Encyclopedia. Retrieved 2011-06-12.CS1 maint: Multiple names: authors list (link)
- ^Herbst, Eric (May 12, 2005). 'Chemistry of Star-Forming Regions'. Journal of Physical Chemistry A. 109 (18): 4017–4029. Bibcode:2005JPCA..109.4017H. doi:10.1021/jp050461c. PMID16833724.
- ^Tullo, Alexander H. (28 July 2014). 'C&EN's Global Top 50 Chemical Firms For 2014'. Chemical & Engineering News. American Chemical Society. Retrieved 22 August 2014.
Bibliography
- Atkins, Peter; de Paula, Julio (2009) [1992]. Elements of Physical Chemistry (5th ed.). New York: Oxford University Press. ISBN978-0-19-922672-6.
- Burrows, Andrew; Holman, John; Parsons, Andrew; Pilling, Gwen; Price, Gareth (2009). Chemistry3. Italy: Oxford University Press. ISBN978-0-19-927789-6.
- Housecroft, Catherine E.; Sharpe, Alan G. (2008) [2001]. Inorganic Chemistry (3rd ed.). Harlow, Essex: Pearson Education. ISBN978-0-13-175553-6.
Further reading
- Popular reading
- Atkins, P.W. Galileo's Finger (Oxford University Press) ISBN0-19-860941-8
- Atkins, P.W. Atkins' Molecules (Cambridge University Press) ISBN0-521-82397-8
- Kean, Sam. The Disappearing Spoon – and other true tales from the Periodic Table (Black Swan) London, 2010 ISBN978-0-552-77750-6
- Levi, PrimoThe Periodic Table (Penguin Books) [1975] translated from the Italian by Raymond Rosenthal (1984) ISBN978-0-14-139944-7
- Stwertka, A. A Guide to the Elements (Oxford University Press) ISBN0-19-515027-9
- 'Dictionary of the History of Ideas'. Archived from the original on March 10, 2008.
- 'Chemistry' . Encyclopædia Britannica. 6 (11th ed.). 1911. pp. 33–76.
- Introductory undergraduate text books
- Atkins, P.W., Overton, T., Rourke, J., Weller, M. and Armstrong, F. Shriver and Atkins inorganic chemistry (4th edition) 2006 (Oxford University Press) ISBN0-19-926463-5
- Chang, Raymond. Chemistry 6th ed. Boston: James M. Smith, 1998. ISBN0-07-115221-0.
- Clayden, Jonathan; Greeves, Nick; Warren, Stuart; Wothers, Peter (2001). Organic Chemistry (1st ed.). Oxford University Press. ISBN978-0-19-850346-0.
- Voet and Voet Biochemistry (Wiley) ISBN0-471-58651-X
- Advanced undergraduate-level or graduate text books
- Atkins, P.W. Physical Chemistry (Oxford University Press) ISBN0-19-879285-9
- Atkins, P.W. et al. Molecular Quantum Mechanics (Oxford University Press)
- McWeeny, R. Coulson's Valence (Oxford Science Publications) ISBN0-19-855144-4
- Pauling, L. The Nature of the chemical bond (Cornell University Press) ISBN0-8014-0333-2
- Pauling, L., and Wilson, E.B. Introduction to Quantum Mechanics with Applications to Chemistry (Dover Publications) ISBN0-486-64871-0
- Smart and Moore Solid State Chemistry: An Introduction (Chapman and Hall) ISBN0-412-40040-5
- Stephenson, G. Mathematical Methods for Science Students (Longman) ISBN0-582-44416-0
External links
- General Chemistry principles, patterns and applications.
Retrieved from 'https://en.wikipedia.org/w/index.php?title=Chemistry&oldid=898107136'
B. Sc. CBZ with Biotech 2nd Year
(Chemistry Practical)
1. Qualitative Analysis of Organic Compounds (2-33)
2. Inorganic Quantitative Analysis (33-45)
3. Inorganic Synthesis (46-54)
Qualitative Analysis of Organic Compounds.
The analysis and identification of unknown organic compounds constitutes a very important aspect of experimental organic chemistry. There is no definite set procedure that can be generally applied to organic qualitative analysis. Various books have different approaches, but a systematic approach based on the scheme given below will give good results.
Practical Notes
Before outlining the general scheme, one or two points of practical importance should be noted.
(a) Quantities of substance for tests. For most tests about 0.1 g solid or 0.1 - 0.2 mL (2 - 3 drops) of liquid material (NOT MORE) should be used.
(b) Reagents likely to be met within organic analysis are on the reagent shelves. Students are advised to develop a general knowledge of the physical characteristics of common organic compounds. If in doubt about the expected result of a test between a certain compound and a reagent, carry out a trial test with a known compound and compare with the unknown.
(c) Quantities of substance derivatives. Students have wasted much time and material in the past by taking too large a quantity of substance for preparation of a derivative. In general, 0.5 - 1 g (or 0.5 - 1 mL) of substance gives the most satisfactory results.
General Scheme of Analysis
A. Preliminary Tests
(a) Note physical characteristics - solid, liquid, colour and odour.
(b) Ignition test (heat small amount on metal spatula) to determine whether the compound is aliphatic or aromatic (i.e. luminous flame - aliphatic; sooty flame - aromatic).
B. Physical Constants
Determine the boiling point or melting point. Distillation is recommended in the case of liquids (see Appendix 3). It serves the dual purpose of determining the b.p., as well as purification of the liquid for subsequent tests.
C. Analysis for elements present (Lassaigne's Test)
In organic compounds the elements commonly occurring along with carbon and hydrogen, are oxygen, nitrogen, sulphur, chlorine, bromine and iodine. The detection of these elements depends upon converting them to water-soluble ionic compounds and the application of specific tests.
Theory
It is a general test for the detection of halogens, nitrogen & sulphur in an organic compound. These elements are bonded covalently in the organic compounds. In order to detect them, these have to be converted into their ionic forms. This is done by fusing the organic compound with sodium metal. The ionic compounds formed during the fusion are extracted in aqueous solution, and can be detected by simple chemical tests. The extract is called sodium fusion extract or Lassaigne's extract.
PROCEDURE
Place a piece of clean sodium metal, about the size of a pea into a fusion tube. Add a little of the compound (50 mg or 2 - 3 drops).* Heat the tube gently at first, allowing any distillate formed to drop back onto the molten sodium. When charring begins, heat the bottom of the tube to dull redness for about three minutes and finally plunge the tube, while still hot, into a clean dish containing cold distilled water (6 mL) and cover immediately with a clean wire gauze.**
Test for Nitrogen
The carbon and nitrogen present in the organic compound on fusion with sodium metal give sodium cyanide (NaCN) soluble in water. This is converted in to sodium ferrocyanide by the addition of sufficient quantity of ferrous sulphate .Ferric ions generated during the process reacts with ferrocyanide to form blue precipitate of ferric ferrocyanide.
Test for chlorine
Chlorine present in the organic compound forms sodium chloride on fusion with sodium metal. Sodium chloride, extracted with water, can be easily identified by adding silver nitrate solution after acidifying with dil. Nitric acid.
Test for sulphur
If sulphur is present in the organic compound, sodium fusion will convert it into sodium sulphide. Sulphide ions are readily identified by sodium nirtoprusside.
The 'fusion' filtrate which should be clear and colourless, is used for the SPECIFIC TESTS DESCRIBED BELOW:
1. To a portion (2 mL) of the 'fusion' filtrate add 0.2 g of powdered ferrous sulphate crystals. Boil the mixture for a half a minute, cool and acidify by adding dilute sulphuric acid dropwise. Formation of a bluish-green precipitate (Prussian blue) or a blue solution indicates that the original substance contains nitrogen. If no precipitate appears, allow to stand for 15 minutes, filter and inspect filter paper.
2. SULPHUR (SULPHIDE)
To the cold 'fusion' filtrate (1 mL) add a few drops of cold, freshly prepared, dilute solution of sodium nitroprusside. The latter may be prepared by adding a small crystal of the solid to 2 mL of water. Production of a rich purple colour indicates that the original substance contains sulphur. This test is very sensitive. Only strong positive results are significant.
3. HALOGENS (HALIDES)
Acidify a portion (1 mL) of the 'fusion' filtrate with 2N nitric acid, and if nitrogen and/or sulphur are present, boil for 1 - 2 minutes.* Cool and add aqueous silver nitrate (1 mL), compare with a blank. Formation of a heavy, white or yellow precipitate of silver halide indicates halogen. If a positive result is obtained: acidify the remaining portion of the 'fusion' filtrate with dilute sulphuric acid, boil and cool. Add carbon tetrachloride (1 mL) and a few drops of freshly prepared chlorine water. Shake the mixture.
(a) If the carbon tetrachloride layer remains colourless - indicates chlorine.
(b) If the carbon tetrachloride layer is brown - indicates bromine.
(c) If the carbon tetrachloride layer is violet - indicates iodine.
*If nitrogen and/or sulphur are also present, the addition of silver nitrate to the acidified 'fusion' solution will precipitate silver cyanide and/or silver sulphide in addition to the silver halides. The removal of hydrogen cyanide and/or hydrogen sulphide is effected by boiling the 'fusion' solution. GROUP CLASSIFICATION TESTS
D. Solubility tests
The solubility of the unknown in the following reagents provides very useful information. In general, about 3 mL of the solvent is used with 0.1 g or 0.2 mL (2 - 3 drops) of the substance. The class of compound may be indicated from the following table:
SOLUBILITY TABLE
REAGENT AND TEST CLASS GROUP OF COMPOUNDS
Soluble in cold or hot water. (If the unknown is soluble do NOT perform solubility tests below) Neutral, acidic or basic. (Test with litmus or universal indicator paper) Lower members of series. Neutral, e.g. alcohols; Acidic, e.g. acids, phenols; Basic, e.g. amines
Soluble in dil. HCl Basic Most amines (except III amines with only aromatic groups
Soluble in dil. NaOH Acidic Most acids, most phenols.
Soluble in NaHCO3 Strongly acidic Most carboxylic acids.
Insoluble in water, acid and alkali Neutral Hydrocarbons, nitrohydro-carbons, alkyl or aryl halides, esters and ethers. Higher molecular weight alcohols, aldehydes and ketones
E. Group Classification Tests
From the previous tests it is often possible to deduce the functional groups present in the unknown compound. Consult i.r. spectra when available.
Individual tests are then performed to identify and confirm the functional groups present.
NOTE:
1. Students are strongly advised against carrying out unnecessary tests, since not only are they a waste of time but also increase the possibility of error. Thus it is pointless to first test for alcohol or ketone in a basic compound containing nitrogen! Instead tests for amines, etc. should be done on such a compound.
2. A systematic approach cannot be overemphasised in group classification tests to avoid confusion and error.
Once the functional group has been identified, reference is made to tables in a book on organic analysis, for assessing possibilities and for the preparation of suitable solid derivatives.
It should be noted that whilst two substances with the same functional group may sometimes have very similar b.p. or m.p., solid derivatives canusually be chosen from the literature, with m.p. differences of about 10 (or more), which distinguish between the two possibilities.
Example:
COMPOUND B.P. DERIVATIVES (M.P.)
2,4-DNPH SEMICARBAZONE
Diethyl ketone 102 156 139
Methyl n-propyl ketone 102 144 112
G. Preparation of derivatives
The final characterisation of the unknown is made by the preparation of suitable solid derivatives. The derivative should be carefully selected and its m.p. should preferably be between 90 - 150 for ease of crystallisation and m.p. determination.
Preparation of one derivative should be attempted. The derivative should be purified by recrystallisation, dried and the m.p. determined. Derivatives should be submitted correctly labelled for assessment together with the record.
Recording of Results
The results should be recorded in a systematic manner. Results should be recorded in the practical book at the time (not written up afterwards).
A record should be made of every test carried out, no matter whether a NEGATIVE RESULT HAS BEEN OBTAINED.
Tests for unsaturation
1. Cold dilute potassium permanganate solution.
2. Solution of bromine in carbon tetrachloride.
Tests for compounds containing nitrogen
1. Amines
(a) Nitrous acid.
(b) Confirmatory tests.
2. Compounds which give amines or ammonia on acid or alkaline hydrolysis:
Amides, substituted amides, anilides, nitriles.
3. Compounds which give amines on reduction:
Nitro, nitroso, azo, hydrazo, nitriles.
Tests for compounds containing C, H and possibly oxygen
1. Carboxylic acids
Na2CO3 or NaHCO3 solution liberate carbon dioxide.
2. Phenols
(a) Sodium hydroxide solution (soluble). Insoluble in and no CO2 from NaHCO3 (except when electron attracting groups present, e.g. 2,4-dinitrophenol).
(b) Ferric chloride solution.
(c) Bromine water.
3. Aldehydes and Ketones
(a) 2,4-dinitrophenylhydrazine (as Brady's reagent) for C=O.
(b) Iodoform test for CH3CO-.
4. Aldehydes only (reducing properties)
(a) Fehling's solution.
(b) Tollen's reagent (ammoniacal AgNO3 solution).
(c) Jones reagent.
5. Alcohols
(a) Lucas' reagent to distinguish I, II and III alcohols.
(b) Jones reagent.
(c) Metallic sodium (use dry liquid and dry tube).
6. Sugars
(a) Molisch's test.
7. Esters
(a) Hydroxamic acid test.
(b) Hydrolysis.
________________________________________
Write up of the identification of an unknown organic compound
Compound containing C, H ,N, Halogens, S
Physical characteristics ...................... (solid, liquid, gas, colour, odour, etc.)
Ignition test .............................. (aromatic or aliphatic)
Physical constant ........................ (boiling point or melting point)
Solubility tests (in tabular form)
Group classification tests (in tabular form)
Test Observation Inference
From the above tests and observations the given compound is probably a
.........................(acid, phenol, aldehyde, etc.)
Consultation of literature (Possibilities) M.P. of derivative
(a)
(b)
(c)
Preparation of derivative (method of preparation)
Observed m.p. of derivative
Lit. m.p. of derivative
Result
Compound No. ........................ is ............................
(give formula)
TESTS FOR FUNCTIONAL GROUPS
I. UNSATURATED COMPOUNDS
Two common types of unsaturated compounds are alkenes and alkynes characterised by the carbon-carbon double and triple bond, respectively, as the functional group. The two common qualitative tests for unsaturation are the reactions of the compounds with (a) bromine in carbon tetrachloride and (b) potassium permanganate.
(a) 2% Bromine in carbon tetrachloride
Dissolve 0.2 g (or 0.2 mL) of the compound in 2 mL of carbon tetrachloride or another suitable solvent and add the solution dropwise to 2 ml of 2% bromine solution in carbon tetrachloride and shake. e. g.
Rapid disappearance of the bromine colour to give a colourless solution is a positive test for unsaturation.
NOTE: The reagent is potentially dangerous. Keep it off your skin and clothes; protect your eyes and nose.
(b) 2% Aqueous potassium permanganate
Dissolve 0.2 g (or 0.2 mL) of the substance in 2 mL of water (acetone may also be used as solvent). Add the potassium permanganate solution dropwise and observe the result.
e.g.
For a blank determination, count the number of drops added to 2 mL of acetone before the colour persists. A significant difference in the number of drops required in the two cases is a positive test for unsaturation.
II. COMPOUNDS CONTAINING NITROGEN
1. Amines
(a) Reaction with nitrous acid Dissolve the amine (0.5 mL) in concentrated acid (2.0 mL) and water (3 mL) and cool the solution to 0 - 5 in an ice-bath for 5 minutes. Add a cold solution (ice-bath) of sodium nitrite (0.5 g) in water (2.0 mL) from a dropper, with swirling of the test tube, still keeping the mixture in the ice-bath.
AMINE REACTION
Aliphatic N2 evolved.
RNH2 + HNO2 -> ROH + N2 + H2O
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Aromatic Diazonium salt is formed.
ArNH2 + HNO2 -> ArN=N+
Add the cold diazonium solution and with swirling
to a cold solution of 2-naphthol (0.2 g) in 5% NaOH
solution (2 mL). An orange-red azo dye is formed.
__________________________________________________________________
aliphatic and Yellow oily nitrosamines are generally formed.
aromatic R2NH + HNO2 -> R2N-NO
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(b) Reaction with benzenesulphonyl chloride
Benzenesulphonyl chloride reacts with primary and secondary but not with tertiary amines to yield substituted sulphonamides.
e.g. (a) C6H5SO2Cl + H-NHR + NaOH -> C6H5SO2NHR + NaCl + H2O
(b) C6H5SO2Cl + H-NR2 + NaOH -> C6H5SO2NR2 + NaCl + H2O
The substituted sulphonamide formed from a primary amine dissolves in the alkali medium whilst that produced from a secondary amine is insoluble in alkali.
Place 0.5 mL (or 0.5 g) of the compound, 15 - 10 mL of 5% NaOH and 1 mL of benzenesulphonyl chloride in a test tube, stopper the tube and shake until the odour of the sulphonyl chloride has disappeared. The solution must be kept alkaline (if no reaction has occurred, the substance is probably a tertiary amine).
If a precipitate appears in the alkaline solution, dilute with about 10 mL of water and shake; if the precipitate does not dissolve, a secondary amine is indicated.
If there is no precipitate, acidify it cautiously to congo red with concentrated hydrochloric acid (added dropwise): a precipitate is indicative of a primary amine.
2. Amides R-CO-NH2
Simple primary amides can be decomposed by boiling with alkali and thereby evolving ammonia.
e.g. CH3-CO-NH2 + NaOH -> CH3-CO2- Na+ + NH3 ¬
Boil 0.5 g of the compound with 5 mL of 10% sodium hydroxide solution and observe whether ammonia is evolved.
III. COMPOUNDS CONTAINING C, H AND POSSIBLY OXYGEN
1. Carboxylic acids - test with 5% aq. NaHCO3
R-CO2H + NaHCO3 -> R-CO2- Na+ + CO2 ¬ + H2O
Sodium hydrogen carbonate reacts with carboxylic acids to give the sodium salt of the acid and liberates carbon dioxide. If the acid is insoluble in water and the reaction is sluggish dissolve the acid in methanol and add carefully to a saturated sodium hydrogen carbonate solution, when a vigorous effervescence will be observed.
2. Phenols [Soluble in NaOH and produce no CO2 from NaHCO3]
(a) Bromine water
Phenols are generally highly reactive towards electrophilic reagents and are readily brominated by bromine water. e.g.
Dissolve or suspend about 0.05 g of the compound in 2 mL of dilute hydrochloric acid and add bromine water dropwise until the bromine colour remains. A white precipitate of the bromophenol may form. Solid bromophenol derivatives can be used for the confirmation of the structure of a phenol (cf the preparation of derivatives).
(b) Ferric chloride test
Most phenols react with iron (III) chloride to form coloured complexes. The colours vary - red, purple, blue or green - depending on various factors, e.g. the phenolic compound used, the solvent, concentration. Since some phenols do not give colours, a negative test must not be taken as significant without supporting information.
Dissolve 0.05 g of the compound in 2 mL water (or a mixture of water and ethanol if the compound is not water-soluble) and add an aqueous solution of ferric chloride dropwise. Observe any colour changes which may occur.
3. Aldehydes and ketones
(a) 2,4-Dinitrophenylhydrazine (as Brady's reagent)
A test for the carbonyl group (C=O) in aldehydes and ketones. 2,4- Dinitrophenylhydrazine gives sparingly soluble yellow or red 2,4-dinitrophenylhydrazones with aldehydes and ketones.
Add 3 mL of the reagent to 2 drops of the compound in a test tube and shake. If no precipitate forms immediately, warm and allow to stand for 5 - 10 minutes. A crystalline precipitate indicates the presence of a carbonyl compound.
The bench reagent is very dilute and is intended for qualitative tests only and should not be used in the preparation of a derivative for identification purposes.
(b) Iodoform test for CH3CO-
Dissolve 0.1 g (or 5 drops) of the compound in 2 mL of water; if it is insoluble in water add sufficient dioxan to produce a homogeneous solution. Add 2 mL of 5% NaOH solution and then introduce the potassium iodide - iodine reagent dropwise with shaking until a definite dark colour of iodine persists. Allow to stand for 2 - 3 minutes; if no iodoform separates at room temperature, warm the test tube in a beaker of water at 60 . Add a few more drops of the iodine reagent if the faint iodine colour disappears. Continue the addition of the reagent until a dark colour is not discharged after 2 minutes heating at 60 . Remove the excess of iodine by the addition of a few drops of dilute sodium hydroxide solution with shaking, dilute with an equal volume of water, and allow to stand for 10 minutes. The test is positive if a yellow precipitate of iodoform is deposited. Filter off the yellow precipitate, dry upon pads of filter paper and determine the m.p. Iodoform melts at 120 (it can be recrystallised from methanol- water).
The reaction is given by acetaldehyde and simple methyl ketones. Alcohols containing the CH3CHROH group will be oxidised under the reaction conditions and also give a positive test.
4. Aldehydes only (reducing properties).
(a) Fehling's solution
Aldehydes reduce Fehling's solution to yellow or red copper (I) oxide.
Preparation of the reagent: Mix equal volumes of Fehling's solution solution I (aqueous alkaline potassium tartrate) and Fehling's solution II (copper sulphate solution).
Add 2 drops (or 0.05 g) of the compound and 2 - 3 drops of the reagent and heat on a boiling water bath for 3 - 4 minutes.
The test is positive for aliphatic aldehydes, but is often indecisive for aromatic aldehydes where Jones' Reagent is often useful (see 5).
(b) Tollen's reagent (Ammonical silver nitrate solution)
Aldehydes are readily oxidised to carboxylic acids and will reduce Tollen's reagent to produce a silver mirror on the inside of a clean test tube.
FIRST clean up a test tube with a little hot nitric acid (fume cupboard) and rinse with distilled water.
Preparation of the reagent: To 1 mL of silver nitrate solution add a few drops of sodium hydroxide. Then add dilute ammonium hydroxide dropwise until the precipitate just dissolves.
Add 2 - 3 drops of the compound in methanol to 2 - 3 mL of Tollen's solution contained in a very clean test tube. If no reaction takes place in the cold, warm gently in a water bath.
CAUTION: After the test, pour the contents of the test tube into the sink and wash the test tube with dilute nitric acid. Any silver fulminate present, which is highly explosive when dry, will be destroyed.
(c) Jones Reagent (See section under alcohols).
5. Alcohols
The tests for the hydroxyl group not only detect the presence of the group, but may also indicate whether it is primary, secondary or tertiary.
(a) Jones Reagent (CrO3-H2SO4 in H2O)
This reagent distinguishes primary and secondary alcohols from tertiary alcohols; the test is based on the much greater resistance to oxidation of tertiary alcohols compared to the other two types. Aldehydes also give a positive test.
Place 1 mL of acetone in a test tube and dissolve one drop of a liquid or ca 10 mg of a solid alcohol or aldehyde in it. Add one drop of the reagent to the acetone solution and shake the tube to mix the contents. Primary and secondary alcohols react within two seconds as indicated by the disappearance of the orange colour of the reagent and the formation of a green or blue-green precipitate or emulsion.
Tertiary alcohols do not react even after 3 minutes.
(I) RCH2OH -> RCHO -> RCO2H
(II) R2CHOH -> R2C=O
(III) R3COH -> no visible reaction.
(b) Lucas' Reagent [ZnCl2 - conc. HCl]
This reagent converts alcohols into the corresponding alkyl chlorides. Zinc chloride (a Lewis acid) increases the reactivity of alcohols towards acid. The test depends on the rate of reaction of primary, secondary, and tertiary alcohols with the reagent at room temperature.
(I) RCH2OH -> no reaction at room temperature.
(II) R2CHOH -> R2CHCl + H2O (1 hour or maybe longer)
(III) R3COH -> R3CCl + H2O (immediately)
To 1 mL of the alcohol in a small test tube add 6 mL of Lucas' reagent at room temperature. Close the tube with a cork, shake and allow to stand.
(i) Primary alcohols - the aqueous phase remains clear (except allyl alcohol - droplets after 7 minutes).
(ii) Secondary alcohols - very slow reaction (~ 1 hour or maybe longer) when droplets of alkyl chloride may be seen.
(iii) Tertiary alcohols - very fast reaction and droplets of the alkyl chloride formed almost immediately.
6. Sugars, Carbohydrates
Molisch's Test
This is a general test for carbohydrates. Dissolve 20 - 30 mg of the compound in 2 mL water and add 0.5 mL of the reagent (a 20% solution of 2-naphthol in ethanol). Pour 2 mL of concentrated sulphuric acid from a dropper carefully down the side of the tube so that the acid forms a layer beneath the aqueous solution without mixing with it. A red colouration, changing to dark purple forms at the interface. Carry out a second test on a blank solution.
7. Esters
Hydroxamic acid test
R-CO-OR' + H2N-OH -> R-CO-NH-OH + R'-OH
Esters react with hydroxylamine in the presence of sodium hydroxide to form the sodium salt of the corresponding hydroxamic acid. On acidification and addition of ferric chloride the magenta-coloured iron (III) complex of the hydroxamic acid is formed.
It is always advisable to ensure that an unknown compound does not give a colour with iron (III) chloride before carrying out the hydroxamic acid test.
Procedure for hydroxamic acid test
(a) Ferric chloride test
Dissolve a drop or a few small crystals of the compound in 1 mL of 95% ethanol (rectified spirit) and add 1 mL of M hydrochloric acid. Note the colour produced when 1 drop of 5% iron (III) chloride is added to the solution. If a pronounced violet, blue, red or orange colour is produced, the hydroxamic acid test described below is NOT APPLICABLE.
(b) Hydroxamic acid test
Mix 1 drop or several small crystals (ca 0.05 g) of the compound with 1 mL of 0.5 M hydroxylamine hydrochloride in 95% ethanol and add 0.2 mL of 6 M aqueous sodium hydroxide. Heat the mixture to boiling and after the solution has cooled slightly add 2 mL of M hydrochloric acid. If the solution is cloudy, add 2 mL of 95% ethanol. Observe the colour produced when 1 drop of 5% iron (III) chloride solution is added. If the resulting colour does not persist, continue to add the reagent dropwise until the observed colour pervades the entire solution. Usually only 1 drop of the iron (III) chloride solution is necessary. Compare the colour with that produced in test (a). A positive test will be a distinct burgundy or magenta colour as compared with the yellow colour observed when the original compound is tested with iron (III) chloride solution in the presence of acid. It is often advisable to conduct in parallel the test with, say, ethyl acetate, to ensure that the conditions for this test are correct.
THE PREPARATION OF DERIVATIVES OF ORGANIC COMPOUNDS
The preliminary examination and group classification tests indicate the particular class (functional group) to which an unknown organic compound may belong. Further characterisation and identification depends on the selection and preparation of a suitable solid derivative and accurate determination of its melting point (best, between 90 - 150 ).
The following table lists some of the classes of organic compounds and a selection of derivatives that may be prepared to characterise them. Check with the tables of melting points in Vogel which derivatives are most suitable for the characterisation of your particular compound.
CLASS OF COMPOUND DERIVATIVES
1. ALCOHOLS 3,5-dinitrobenzoate
2. PHENOLS benzoate, acetate, bromo-derivative
3.ALDEHYDES AND KETONES semicarbazone, 2,4-dinitrophenyl-hydrazone, oxime
4. ACIDS anilide, amide, p-toluidide.
5. AMINES benzoyl, acetyl and sulphonamide derivatives
METHODS FOR THE PREPARATION OF DERIVATIVES
ALCOHOLS
(i) 3,5-Dinitrobenzoates
3,5-Dinitrobenzoyl chloride is usually partially hydrolysed and should be prepared in the pure state by heating gently a mixture of 3,5-dinitrobenzoic acid (1 g) and phosphorus pentachloride (1.5 g) in a dry test tube, until it liquifies (5 min).* The liquid is poured on a dry watch glass and allowed to solidify. The phosphoryl chlorides are removed by pressing the solid with a spatula on a wad of filter paper. The residual acid chloride is suitable for immediate use in the preparation of the derivatives.
*Work under fume hood. Fumes are irritating to the eyes and nose.
The 3,5-dinitrobenzoyl chloride is mixed with the alcohol (0.5 - 1 mL) in a loosely corked dry test tube and heated on a steam bath for about 10 min. Secondary and tertiary alcohols require up to 30 min. On cooling add 10 mL sodium hydrogen carbonate solution, stir until the ester crystallises out, and filter at the pump. Wash with a little carbonate solution, water and suck dry. Recrystallise from the minimum hot ethanol or light petroleum. Cool slowly to avoid the formation of oily droplets of your ester.
PHENOLS
(i) Benzoates (Schötten-Baumann method).
To the phenol (0.5 g) is added 5% sodium hydroxide (10 mL) in a well-corked boiling tube or a small conical flask. Benzoyl chloride (2 mL) is added in small quantities at a time, and the mixture shaken vigorously with occasional cooling under the tap or in ice-water. After 15 min the solid benzoate separates out: the solution should be alkaline at the end of the reaction; if not alkaline, or if the product is oily, add a solid pellet of sodium hydroxide and shake again. Collect the benzoate, wash thoroughly with cold water, and recrystallise from alcohol or light petroleum.
(ii) Acetates
Acetates of many simple phenols are liquids; however, this is a suitable derivative for polyhydric and substituted phenols. The phenol (0.5 g) is dissolved in 10% sodium hydroxide solution and an equal quantity of crushed ice is added, followed by acetic anhydride (2 mL). The mixture is vigorously shaken in a stoppered test tube until the acetate separates. The product is filtered and recrystallised from alcohol.
(iii) Bromo derivatives
The phenol (0.3 g) is suspended in dilute hydrochloric (10 mL) and bromine water added dropwise until no more decolourisation occurs. The bromo derivative which precipitates out is filtered off and recrystallised from alcohol.
ALDEHYDES AND KETONES
(i) Semicarbazones
Dissolve semicarbazide hydrochloride (1 g) and sodium acetate (1.5 g) in water (8 - 10 mL), add the aldehyde or ketone (0.3 mL) and shake. Shake the mixture for a few minutes and then cool in ice-water. Filter off the crystals, wash with a little cold water and recrystallise from methanol or ethanol.
(ii) 2,4-Dinitrophenylhydrazones
Suspend 0.25 g of 2,4-dinitrophenylhydrazine in 5 mL of methanol and add 0.5 mL of concentrated sulphuric acid cautiously. Filter the warm solution and add a solution of 0.2 g of the carbonyl compound in 1 mL of methanol. Recrystallise the derivative from methanol, ethanol or ethyl acetate.
(iii) Oximes
Hydroxylamine hydrochloride (0.5 g) is dissolved in water (2 mL). 10% sodium hydroxide (2 mL) and the carbonyl compound (0.2 - 0.3 g) dissolved in alcohol (1 - 2 mL) are added, the mixture warmed on a steam bath for 10 min and then cooled in ice. Crystallisation is induced by scratching the sides of the test tube with a glass rod. The oximes may be crystallised from alcohol.
ACIDS
(i) Amides, anilides and p-toluidides
The acid (0.5 g) is refluxed with thionyl chloride (2 - 3 mL) in a fume cupboard for about 30 mins.* It is advisable to place a plug of cotton wool in the top of the reflux condenser to exclude moisture. The condenser is removed and the excess of thionyl chloride is distilled off (b.p. 78 ). The acid chloride thus produced is treated with concentrated ammonia solution (5 mL) or aniline (0.5 - 1 mL) or p-toluidine (0.5 - 1 g), when the solid derivative separates out. It is collected and recrystallised from alcohol adding decolourising charcoal if found necessary.
*Alternately use PCl5 to form the acid chloride.
AMINES
(i) Acetyl derivatives (acetamides)
Reflux gently in a small dry flask under a dry condenser the amine (1 g) with acetic anhydride (3 mL) for 15 min. Cool the reaction mixture and pour into 20 mL cold water. Boil to decompose the excess acetic anhydride. Cool and filter by suction the insoluble derivative. Recrystallise from ethanol.
(ii) Benzoyl derivatives (benzamides)
Suspend 1 g of the amine in 20 mL of 5% aqueous sodium hydroxide in a well-corked flask, and add 2 mL benzoyl chloride (fume hood!), about 0.5 mL at a time, with constant shaking. Shake vigorously for 5 - 10 min until the odour of the benzoyl chloride has disappeared. Ensure that the mixture remains alkaline. Filter off the solid derivative, wash with a little cold water and recrystallise from ethanol.
(iii) Benzenesulphonamides
To 1 g of the amine in 20 mL of 5% sodium hydroxide solution in a well-corked flask add 1 mL benzenesulphonyl chloride (fume hood!). Shake the mixture until the odour of the sulphonyl chloride disappears. Check that the solution is alkaline. Acidify if necessary to obtain the precipitated derivative. Concentrated hydrochloric acid added dropwise should be used. Filter the product, wash with a little cold water and suck dry. Recrystallise from ethanol.
Inorganic Quantitative Analysis
Gravimetric Estimation of Ba2+, Fe2+, Zn2+ and Cu2+
All Gravimetric analyses rely on some final determination of weight as a means of quantifying an analyte. Since weight can be measured with greater accuracy than almost any other fundamental property, gravimetric analysis is potentially one of the most accurate classes of analytical methods available. These methods are among the oldest of analytical techniques, and they may be lengthy and tedious. Samples may have to be extensively treated to remove interfering substances. As a result, only a very few gravimetric methods are currently used in environmental analysis.
There are four fundamental types of gravimetric analysis: physical gravimetry, thermogravimetry, precipitative gravimetric analysis, and electrodeposition. These differ in the preparation of the sample before weighing of the analyte. Physical gravimetry is the most common type used in environmental engineering. It involves the physical separation and classification of matter in environmental samples based on volatility and particle size (e.g., total suspended solids). With thermogravimetry, samples are heated and changes in sample mass are recorded. Volatile solids analysis is an important example of this type of gravimetric analysis. As the name implies, precipitative gravimetry relies on the chemical precipitation of an analyte. Its most important application in the environmental field is with the analysis of sulfite. Electrodeposition involves the electrochemical reduction of metal ions at a cathode, and simultaneous deposition of the ions on the cathode.
Common Procedures in Gravimetric Analysis
a. Drying to a Constant Weight
All solids have a certain affinity for water, and may absorb moisture from the laboratory air. Reagents that readily pick up water are termed hygroscopic. Those that absorb so much water that they will dissolve in it and form a concentrated solution are called deliquescent (e.g., sodium hydroxide, trichloroacetic acid). These types of substances will continually increase in weight while exposed to the air. For this reason, many types of laboratory procedures require that a sample be dried to a constant, reproducible weight (i.e., absorbed moisture removed to some standard, low level). This is especially important for the gravimetric methods. Generally, the sample is dried in a 103 C to 110 C oven for about 1 hour and allowed to cool to room temperature in a desiccator. It is then weighed, and heated again for about 30 minutes. The sample is cooled and weighed a second time. The procedure is repeated until successive weighings agree to within 0.3 mg.
b. Description and Use of the Analytical Balance
The analytical balance is the most accurate and precise instrument in an environmental laboratory. Objects of up to 100 grams may be weighed to 6 significant figures. Volumetric glassware is accurate to no more than 4 significant figures, and the accuracy of complex analytical methods rarely justifies more than 2 significant figures. Analytical balances are generally used for gravimetric analyses, and for the preparation of standard solutions.
Summary of Gravimetric Methods for Environmental Analysis
Some gravimetric methods are in generally using for the analysis of waters and wastewaters.
Type Analyte Pretreatment
Physical Total Solids Evaporation
Suspended Solids Filtration
Dissolved Solids Filtration + Evaporation
Oil & Grease extraction with C2Cl3F3 + distillation of solvent
Surfactants extraction into ethylacetate + evaporation
Thermal Volatile Solids Evaporation + 550`C for 15 min
Volatile Suspended Solids Filtration + 550`C for 15 min
Precipitative Mg with Diammonium hydrogen phosphate and final pyrolysis
Na with zinc uranyl acetate
Silica precipitation/ ignition/ volatilization (with HF)
SO4 with Ba
PHYSICAL GRAVIMETRY
1. Total, Dissolved and Suspended Solids
a. Definitions
Total solids (TS) is generally defined as all matter in a water or wastewater sample that is not water. Because solids are not a specific chemical compound, but rather a diverse collection of dissolved and particulate matter, their concentration cannot be determined in an unambiguous way. Instead, they must be defined by the procedure used to estimated their concentration. Total solids may be differentiated according to size into total dissolved solids (TDS) and total suspended solids (TSS). Once again, this is an operational distinction, whereby all solids passing through filter paper of a certain pore size (e.g., 1.5 microns, Whatman #934AH) are called dissolved, and those retained are termed suspended.
b. Significance to Environmental Engineering
Most of the impurities in potable waters are in the dissolved state, principally as inorganic salts. Thus, the parameters, 'total solids' and especially 'total dissolved solids' are of primary importance here. Waters containing high concentrations of inorganic salts are not suitable as sources of drinking water, because such materials are often difficult to remove during treatment. Finished drinking waters containing more than 1000 mg/L TDS are generally considered unacceptable. Waters of this type may also be unsuitable for agricultural purposes due to the harmful effects of high ionic concentrations on plants. In most natural waters, the TDS (total dissolved solids) concentration correlates well with total hardness (i.e., [Ca] + [Mg]). This is useful in assessing the corrosivity of a water and the need for softening.
The total suspended solids (TSS) content of natural waters is of interest for the purpose of assessing particle bed load and transport. High concentrations of suspended matter may be detrimental to aquatic life. In theory, TSS could be used for assessing particle removals during water treatment. However, nearly always the concentration of colloidal particles in water is measured as turbidity since this latter technique is faster and more precise.
Some Typical Solids Concentrations
Source Concentration (mg/L)
Low Avg High
NATURAL WATERS
Fresh TDS 20 120 1,000
Brines TDS 5,000 300,000
DOMESTIC WASTEWATER
Raw TDS 350 600 900
VDS 165 285 600
TSS 100 200 350
VSS 75 135 215
Secondary Effluent TSS 10 30 60
Activated Sludge Mixed
Liquor (conventional) TSS 1,500 3,000
Activated Sludge Mixed
Liquor (extended aeration) TSS 3,000 6,000
Primary Sludge TSS 20,000 70,000
Secondary Sludge TSS 5,000 12,000
STORM WATER TSS 5 300 3,000
Procedures
Total Solids (Total Residue). Total solids is determined by the final weight of a dried sample (minus tare) divided by the original sample volume. Evaporating dishes of platinum, vycor or porcelain may be used. Platinum is preferred, because it is more inert than the other two, and can be heated to a constant weight more easily. However, platinum is very expensive, so porcelain is often used. Porcelain is difficult to bring to a constant weight, and its use should be avoided. Space permitting, evaporating dishes should be stored in a desiccator so as to avoid the collection of dust and absorption of moisture while not in use. The precision of this method has been estimated to be 4 mg or 5%. However, settled wastewater may give better precision, on the order of 1 mg.
1. Preheat a 100 mL evaporating dish at 550 50 C for 1 hour, cool in a in a drying oven or in the open air (protected from dust) for 15-20 minutes, bring to room temperature in a desiccator, and weigh. Repeat until a constant weight is achieved.
2. Measure 75 mL of sample or a volume sufficient to yield 200 mg TS, whichever is less. Add this to the preweighed dish and evaporate to dryness in a drying oven set at 98 C. Alternatively, a steam bath may be used.
3. Dry for an additional hour at 103-105 C.
4. Cool in a desiccator and weigh.
Dissolved Solids (Filtrable Residue). Dissolved solids may be determined directly by analysis of the filtered sample for total solids, or indirectly by determining the suspended solids and subtracting this value from the total solids. When using the direct method, final drying may be conducted at one of two temperatures.
1. Analyze the filtrate in accordance with the total solids procedure.
2. Final drying (1 hour period) may be conducted at either 103-105 C or 180 2 C.
Suspended Solids (Non-filtrable Residue). Suspended solids is measured directly by drying and weighing the solids retained during filtration.
1. Dry this filter at 103-105 C for 1 hour, and cool in a desiccator.
2. Weigh the filter, then pass a water sample of sufficient volume to yield 50-200 mg suspended solids through it. Smaller volumes will result in reduced accuracy.
3. Dry for at least one hour at 103-105 C.
4. Cool in a desiccator and weigh.
C. THERMOGRAVIMETRY AND COMBUSTION ANALYSIS
Thermogravimetry and combustion analysis involve the heating of a sample to 500 C or more with the oxidation and/or volatilization of some of the sample constituents. Either the change of sample weight is determined (thermogravimetry), or the combustion gases are trapped and weighed (combustion analysis). With thermogravimetric methods, it is especially important to return the sample to room temperature before weighing. Otherwise the differences in temperature will create convection currents around the balance pan, which will severely disrupt method accuracy. A steady increase in apparent weight while the sample is on the pan indicates a problem of this type. Large vessels and samples will require longer cooling times to dissipate their excess heat.
Volatile Solids and Fixed Solids
Fixed solids are those that remain as residue after ignition at 550 C for 15 minutes. The weight of material lost is called the volatile solids. Thus the total operational definition for volatile solids would be: all matter lost upon ignition at 550 C for 15 minutes, but not lost upon drying at 103-105 C for 1 hour. The portion lost upon ignition is generally assumed to be equivalent to the organic fraction. The portion remaining is considered the inorganic fraction. For waters of moderate to high hardness, most of this is calcium carbonate which decomposes only at temperatures exceeding 800 C. When igniting a filter with suspended matter, one must be especially careful of the temperature; above 600 C glass fiber filters begin to melt and can loose a significant amount of weight in 15 minutes.
Combining the fractionations resulting from ignition and filtration, one arrives at a total of 9 separate categories: total solids (TS), fixed solids, volatile solids, total dissolved solids (TDS), fixed dissolved solids, volatile dissolved solids, total suspended solids (TSS), fixed suspended solids, and volatile suspended solids (VSS). In practice, only four of these (TS, TDS, TSS, and VSS) are commonly used. When comparing fixed solids with inorganic content, one would expect positive bias from incomplete oxidation of organic matter, and negative bias from decomposition of certain inorganics. Ammonium salts may be lost during low temperature drying or upon ignition. Most others are stable under the conditions used for volatile solids determination with the exception of magnesium carbonate. Volatile solids may be effected by these as well as loss of recalcitrant waters of crystallation (positive bias), and previous losses of organic matter to volatilization during low-temperature drying (negative bias). A modest interlaboratory study found an average standard deviation of 11 mg/L on a sample of 170 mg/L volatile solids.
MgCO3 ------------> MgO + CO2 ¬
Ammonium compounds (often present in sludge in the form of ammonium bicarbonate) may be lost during low temperature drying and therefore should not introduce a bias in volatile solids
NH4HCO3 ---------> NH3 ¬ + H2O ¬ + CO2 ¬
1. Dry and weigh a vessel containing the solids to be analyzed. For volatile and fixed suspended solids analysis, the filter (with residue) prepared for suspended solids analysis and dried to a constant weight may be used. For volatile and fixed total (or dissolved) solids, the evaporating dish (with residue) prepared for total (or dissolved) solids analysis dried to a constant weight should be used.
2. Ignite the sample and vessel in a preheated muffle furnace set at 550 50 C for 15-20 min (water & wastewater) or 1 hour (sludge, sediment & soil).
3. Cool for 15 minutes in the open air in an area protected from dust.
4. Place vessel in a desiccator for final cooling to room temperature and weigh. Due to the approximate nature of this test samples are not generally re-heated and dried to a constant weight.
D. PRECIPITATIVE GRAVIMETRIC ANALYSIS
Precipitative gravimetric analysis requires that the substance to be weighed be readily removed by filtration. In order for a non-filtrable precipitate to form, it must be supersaturated with respect to its solubility product constant. However, if it is too far above the saturation limit, crystal nucleation may occur at a rate faster than crystal growth (the addition of molecules to a crystal nucleus, eventually forming a non-filtrable crystal). When this occurs, numerous tiny micro-crystals are formed rather than a few large ones. In the extreme case, micro-crystals may behave as colloids and pass through a fibrous filter. To avoid this, precipitating solutions may be heated. Because the solubility of most salts increases with increasing temperature, this treatment will lower the relative degree of supersaturation and slow the rate of nucleation. Also, one might add the precipitant slowly with rapid mixing to avoid the occurrence of locally high concentrations.
It is very important that the precipitate be pure and have the correct stoichiometry. Coprecipitation occurs when an unwanted ion or molecule becomes trapped in the precipitate. This may be due to inclusion or occlusion. Inclusion is the term used of a single subsitution in the crystal lattice by an ion of similar size. Occlusion refers to the physical trapping of a large pocket of impurities within the crystal. One technique for minimizing these problems is to remove the mother liquor, re-dissolve the precipitate and then re-precipitate. The second time the mother liquor will contain fewer unwanted ions capable of coprecipitation.
Sulfate Determination
The method of choice for sulfate in waters and wastewaters is the precipitative gravimetric procedure using barium. If Ba(+II) is added in excess under acidic conditions, BaSO4 is precipitated quantitatively. The reaction is allowed to continue for 2 hours or more at 80-90oC. This is to encourage the formation of BaSO4 crystals (non-filtrable) from the initially formed colloidal precipitate (partially filtrable). The precipitate is washed, and then dried at 800`C for 1 hour. Low pH is needed to avoid the precipitation of BaCO3 and Ba3(PO4)
Ba+2 + SO4-2 = BaSO4(ppt.)
Chloride Determination
Chloride may be determined by precipitation with silver. Interfering ions likely to form insoluble silver salts are the other halogens (bromide, iodide), cyanide, and reduced sulfur species (sulfite, sulfide, and thiosulfate). Fortunately, the reduced sulfur compounds can be pre-oxidized with hydrogen peroxide, and the others are rarely present at high concentrations. Although AgCl can be determined gravimetrically, the recommended procedure for water and wastewater is to use a volumetric procedure with chromate as an indicator.
3. Inorganic Synthesis
(i) Synthesis of Cuprous Chloride (work in a fume hood)
(a) Cu2+ + Cu + 2 Cl- → 2 CuCl
Material: CuSO4.5H2O, Cu, NaCl, HCl, Na2SO3, CH3COOH
Procedure: Prepare a solution of 10 g of powdered CuSO4.5H2O and 15 ml of
concentrated HCl in a 250-ml round-bottom flask, add 4 g of NaCl and heat to boiling.
Cover the flask with a little funnel. Add copper to hot solution in small portions. The
green colour of the solution will turn to yellow. Filter off the remaining Cu. Pour the
solution to one litre of cold water with 2 g of Na2SO3.
Wash the precipitated cuprous chloride 2 – 3 times with a solution of 1.5 g of
Na2SO3, 6 ml of HCl and 300 ml of H2O by decantation. Filter out the precipitate and wash it with concentrated acetic acid. Dry it in a drying oven at 100 oC.
( b) 2 Cu2+ + SO2 + 2 Cl- + 2 H2O → 2 CuCl + SO42- + 4 H+
Material: CuSO4.5H2O, NaCl, HCl, SO2, CH3COOH
Procedure: Add 5 g of NaCl to the warm solution (70 oC) of 10 g CuSO4.5H2O and
bubble SO2 through the mixture. CuCl precipitates. Filter out the precipitate and wash
it with a solution of SO2 in water and then with concentrated acetic acid.
CuCl – white crystals, insoluble in water, on air turns to green alkali copper chloride
(ii) Preparation of chrome alum - KCr(SO4)2.12H2O
a) K2Cr2O7 + 3 SO2 + H2SO4 → KCr(SO4)2 + H2O
Material: K2Cr2O7, SO2, H2SO4
Procedure: Blow the SO2 gas through a gas washing bottle with K2Cr2O7 solution acidified with H2SO4. Do not permit the temperature to rise above 60 oC. Above this temperature complexes of chromium (III) sulphate are formed. These complexes contain sulphate in a non-ionisable form and are difficult to crystallise.
Bubble the unreacted gas through a washing bottle with 10 % NaOH solution. Test for the end of the reaction – to the sample of reduced solution in a test tube add small amount of Na2CO3 crystals and heat the mixture just below the boiling point. Let the precipitate settle, the solution over the precipitate has to be colourless. Set the solution aside to crystallise after the end of reduction.When the crystallization is complete, filter off the crystals and wash them with a small amount of water.
Transfer the product to a dry filter paper and let them dry in air.
b) K2Cr2O7 + 3 C2H5OH + 4 H2SO4 → KCr(SO4)2 + 3 CH3CHO + 7 H2O
Material: K2Cr2O7, C2H5OH, H2SO4
Procedure: Dissolve crushed K2Cr2O7 in diluted H2SO4 (1:3) and add, in small portions with stirring, calculated volume (+ 10 % excess) of C2H5OH. Do not permit the temperature to rise above 60°C. Continue like in procedure a).
KCr(SO4)2.12H2O – dark violet crystals, crystallize in regular octahedra, soluble in
water.
(ii) Preparation of potassium tris(oxalate)ferrate(III) trihydrate
Fe(OH)3 + 3 KHC2O4 → K3[Fe(C2O4)3] + 3 H2O
Material: FeSO4.7H2O or (NH4)2Fe(SO4)2.4H2O, K2C2O4, H2C2O4, HNO3, ethanol
Procedure: Dissolve 35 g of FeSO4.7H2O in 100 ml of warm water and add slowly
diluted HNO3 (1:1) to oxidize Fe2+. Add NH3(aq) to the solution until the precipitation
of Fe(OH)3 is completed. Let the precipitate settle and decant the liquid. Filter out the
precipitate and wash it with hot water. Prepare a hot solution of 44 g of KHC2O4
(calculate it as a mixture of K2C2O4 and H2C2O4) in 100 ml H2O. Add precipitate of
Fe(OH)3 in small portions to this solution. Filter the resulting solution and evaporate it
on a steam bath to crystallization. Filter out and wash the crystals on the Buchner
funnel with ethanol/water 1:1 and finally with acetone. Transfer the product to a dry
filter paper and let it dry in air.
K3[Fe(C2O4)3].3H2O – green crystals, photosensitive and decomposes due to
influence of light:
2 K3[Fe(C2O4)3] → K2[Fe(C2O4)2] + K2C2O4 + 2 CO2
(iv) Preparation of iron alum (Ferrous ammoninium sulphate)
2 FeSO4 + H2O2 +H2SO4 → Fe2(SO4)3
Fe2(SO4)3 + (NH4)2SO4 → 2 NH4Fe(SO4)2
Material: FeSO4.7H2O, H2O2, H2SO4, (NH4)2SO4
Procedure: Dissolve FeSO4.7H2O in water to a solution with w(FeSO4) = 0.15. Filter
this solution if necessary and carefully add concentrated H2SO4 in 10 % excess to the
stoichiometry. Then slowly add H2O2 (double amount compared to stoichiometry),
while stirring the mixture continuously. Heat the mixture to the boiling. Make sure that
Fe2+ was oxidized to Fe3+ by a reaction of sample with K3[Fe(CN)6]. Add further H2O2
if necessary.
Evaporate the solution on a steam bath to half the volume and add warm saturated
(by 60 oC) solution of calculated amount of (NH4)2SO4. Let the solution crystallize.
Put the crystals on a dry filter paper and let them dry in air.
NH4Fe(SO4)2.12H2O – colourless or light violet crystals, turn brown on air.
(v) Preparation of lead carbonate
Pb(CH3COO)2 + (NH4)2CO3 → PbCO3 + 2 CH3COONH4
Pb(NO3)2 + 2 NaHCO3 → PbCO3 + 2 NaNO3 + CO2 + H2O
Material: Pb(CH3COO)2 or Pb(NO3)2, (NH4)2CO3 or NaHCO3
Procedure: Add saturated (NH4)2CO3 or NaHCO3 solution (10 % excess to stoichiometry) to saturated Pb2+ salt solution while continuously stirring. Decant the precipitated product with water, filter it and dry at laboratory temperature.
PbCO3 – white powder, insoluble in water, easy soluble in acids and hydroxides,
decomposes by on heating. If Na2CO3 is used for precipitation, the alkali carbonate
2PbCO3.Pb(OH)2 creates.
(vi) Preparation of lead dioxide
a) Pb(NO3)2 + 2 NaOH → Pb(OH)2 + 2 NaNO3
Pb(OH)2 + 2 NaOH → Na2[Pb(OH)4]
Na2[Pb(OH)4] + CaOCl2 → PbO2 + CaCl2 + 2 NaOH + H2O
Material: Pb(NO3)2 or Pb(CH3COO)2, NaOH, CaOCl2 or NaClO, HNO3
Procedure: Dissolve 0.1 mol of Pb2+ salt in 300 ml of hot water and cool the solution.
If a small amount of substance remains undissolved, it will not affect the result. Add a solution of 20 g NaOH in 180 ml of water. First white precipitate of Pb(OH)2 appears, which dissolves in excess of NaOH to Na2[Pb(OH)4].
Mix 40 g of CaOCl2 with 50 ml of water and the necessary amount of Na2CO3 for
reaction with Ca2+. Add 200 ml of water to the mixture and filter it.
Add the filtrate to the boiling solution of Na2[Pb(OH)4] until the test for presence of
Pb2+ in the final solution is negative.
Test for presence of Pb2+: add one drop of Na2S solution to one drop of the
reaction mixture on filter paper. Black precipitate (PbS) indicates presence of Pb2+.
Pour the mixture into 500 ml of water and decant it. Add 100 ml of diluted HNO3
(1:3) to the precipitate, stir and decant it until pH is neutral. Filter off the lead dioxide,
wash it with boiling water and dry.
b) Pb(NO3)2 + 2NaOH → Pb(OH)2 + 2NaNO3
Pb(OH)2 + 2NaOH + Cl2 → PbO2 + 2NaCl + 2H2O
Material: Pb(NO3)2 or Pb(CH3COO)2, NaOH, Cl2
Procedure: Dissolve 10.0 g of Pb(NO3)2 in 80 ml of water acidified with a drop of
concentrated nitric acid and add solution of 2.4 g of sodium hydroxide dissolved in
50 ml of water slightly with constant stirring. Heat the precipitated lead(II) hydroxide
suspension to 70 - 80 oC and bubble chlorine through it at the same time (do not heat
it over 80 oC or PbO will be obtained!). Decant the precipitated brown-black product
with diluted nitric acid (1:3) and then with water. Dry it at 100 oC.
PbO2 – brown powder, insoluble in water, decomposes on heating to Pb3O4 or PbO and oxygen. Good oxidizer.
Preparation of Potash alum K2(SO4).Al(SO4)3•12H2O)
It is white crystalline solid, soluble in water, used for the purification of water, leather industry paper industry and as fire extinguisher.
Melting point is 92oC
Potash alum is commonly known as 'PHITKARI'
Potash alum is prepared by mixing equi-molecular masses of potassium sulphate and aluminum sulphate in water followed by evaporation
K2SO4 + Al2(SO4)3 + 24H2O-- K2SO4.Al2(SO4)3.24H2O
(vii) Preparation of copper ammine sulphate - perform this experiment in a fume hood
CuSO4 + 4 NH3(aq) + H2O → [Cu(NH3)4]SO4.H2O
Material: CuSO4.5H2O, NH3(aq), C2H5OH
Procedure: Place 5 g of finely powdered copper sulphate, CuSO4.5H2O, in a small
beaker, pour upon it 7.5 ml of concentrated ammonia and 3 ml of water. Shake it for
about 1 minute and then heat it gently until all the solid dissolves. Add about 10 ml of
ethanol to the solution, let it stand for about one hour and filter off the crystals. Wash
them with a mixture of 5 ml of concentrated ammonia and 5 ml of ethanol. Dry them
on air in the hood.
[Cu(NH3)4]SO4.H2O – dark blue crystals, soluble in water (18 g in 100 ml of water at
21.5 oC), stable on air.
(viii) Preparation of cuprous chloride – work in a fume hood
a) Cu2+ + Cu + 2 Cl- → 2 CuCl
Material: CuSO4.5H2O, Cu, NaCl, HCl, Na2SO3, CH3COOH
Procedure: Prepare a solution of 10 g of powdered CuSO4.5H2O and 15 ml of concentrated HCl in a 250-ml round-bottom flask, add 4 g of NaCl and heat to boiling.
Cover the flask with a little funnel. Add copper to hot solution in small portions. The green colour of the solution will turn to yellow. Filter off the remaining Cu. Pour the solution to one litre of cold water with 2 g of Na2SO3.
Wash the precipitated cuprous chloride 2 – 3 times with a solution of 1.5 g of Na2SO3, 6 ml of HCl and 300 ml of H2O by decantation. Filter out the precipitate and wash it with concentrated acetic acid. Dry it in a drying oven at 100 oC.
b) 2 Cu2+ + SO2 + 2 Cl- + 2 H2O → 2 CuCl + SO4
2- + 4 H+
Material: CuSO4.5H2O, NaCl, HCl, SO2, CH3COOH
Procedure: Add 5 g of NaCl to the warm solution (70 oC) of 10 g CuSO4.5H2O and bubble SO2 through the mixture. CuCl precipitates. Filter out the precipitate and wash
it with a solution of SO2 in water and then with concentrated acetic acid.
CuCl – white crystals, insoluble in water, on air turns to green alkali copper chloride
(Chemistry Practical)
1. Qualitative Analysis of Organic Compounds (2-33)
2. Inorganic Quantitative Analysis (33-45)
3. Inorganic Synthesis (46-54)
Qualitative Analysis of Organic Compounds.
The analysis and identification of unknown organic compounds constitutes a very important aspect of experimental organic chemistry. There is no definite set procedure that can be generally applied to organic qualitative analysis. Various books have different approaches, but a systematic approach based on the scheme given below will give good results.
Practical Notes
Before outlining the general scheme, one or two points of practical importance should be noted.
(a) Quantities of substance for tests. For most tests about 0.1 g solid or 0.1 - 0.2 mL (2 - 3 drops) of liquid material (NOT MORE) should be used.
(b) Reagents likely to be met within organic analysis are on the reagent shelves. Students are advised to develop a general knowledge of the physical characteristics of common organic compounds. If in doubt about the expected result of a test between a certain compound and a reagent, carry out a trial test with a known compound and compare with the unknown.
(c) Quantities of substance derivatives. Students have wasted much time and material in the past by taking too large a quantity of substance for preparation of a derivative. In general, 0.5 - 1 g (or 0.5 - 1 mL) of substance gives the most satisfactory results.
General Scheme of Analysis
A. Preliminary Tests
(a) Note physical characteristics - solid, liquid, colour and odour.
(b) Ignition test (heat small amount on metal spatula) to determine whether the compound is aliphatic or aromatic (i.e. luminous flame - aliphatic; sooty flame - aromatic).
B. Physical Constants
Determine the boiling point or melting point. Distillation is recommended in the case of liquids (see Appendix 3). It serves the dual purpose of determining the b.p., as well as purification of the liquid for subsequent tests.
C. Analysis for elements present (Lassaigne's Test)
In organic compounds the elements commonly occurring along with carbon and hydrogen, are oxygen, nitrogen, sulphur, chlorine, bromine and iodine. The detection of these elements depends upon converting them to water-soluble ionic compounds and the application of specific tests.
Theory
It is a general test for the detection of halogens, nitrogen & sulphur in an organic compound. These elements are bonded covalently in the organic compounds. In order to detect them, these have to be converted into their ionic forms. This is done by fusing the organic compound with sodium metal. The ionic compounds formed during the fusion are extracted in aqueous solution, and can be detected by simple chemical tests. The extract is called sodium fusion extract or Lassaigne's extract.
PROCEDURE
Place a piece of clean sodium metal, about the size of a pea into a fusion tube. Add a little of the compound (50 mg or 2 - 3 drops).* Heat the tube gently at first, allowing any distillate formed to drop back onto the molten sodium. When charring begins, heat the bottom of the tube to dull redness for about three minutes and finally plunge the tube, while still hot, into a clean dish containing cold distilled water (6 mL) and cover immediately with a clean wire gauze.**
Test for Nitrogen
The carbon and nitrogen present in the organic compound on fusion with sodium metal give sodium cyanide (NaCN) soluble in water. This is converted in to sodium ferrocyanide by the addition of sufficient quantity of ferrous sulphate .Ferric ions generated during the process reacts with ferrocyanide to form blue precipitate of ferric ferrocyanide.
Test for chlorine
Chlorine present in the organic compound forms sodium chloride on fusion with sodium metal. Sodium chloride, extracted with water, can be easily identified by adding silver nitrate solution after acidifying with dil. Nitric acid.
Test for sulphur
If sulphur is present in the organic compound, sodium fusion will convert it into sodium sulphide. Sulphide ions are readily identified by sodium nirtoprusside.
The 'fusion' filtrate which should be clear and colourless, is used for the SPECIFIC TESTS DESCRIBED BELOW:
1. To a portion (2 mL) of the 'fusion' filtrate add 0.2 g of powdered ferrous sulphate crystals. Boil the mixture for a half a minute, cool and acidify by adding dilute sulphuric acid dropwise. Formation of a bluish-green precipitate (Prussian blue) or a blue solution indicates that the original substance contains nitrogen. If no precipitate appears, allow to stand for 15 minutes, filter and inspect filter paper.
2. SULPHUR (SULPHIDE)
To the cold 'fusion' filtrate (1 mL) add a few drops of cold, freshly prepared, dilute solution of sodium nitroprusside. The latter may be prepared by adding a small crystal of the solid to 2 mL of water. Production of a rich purple colour indicates that the original substance contains sulphur. This test is very sensitive. Only strong positive results are significant.
3. HALOGENS (HALIDES)
Acidify a portion (1 mL) of the 'fusion' filtrate with 2N nitric acid, and if nitrogen and/or sulphur are present, boil for 1 - 2 minutes.* Cool and add aqueous silver nitrate (1 mL), compare with a blank. Formation of a heavy, white or yellow precipitate of silver halide indicates halogen. If a positive result is obtained: acidify the remaining portion of the 'fusion' filtrate with dilute sulphuric acid, boil and cool. Add carbon tetrachloride (1 mL) and a few drops of freshly prepared chlorine water. Shake the mixture.
(a) If the carbon tetrachloride layer remains colourless - indicates chlorine.
(b) If the carbon tetrachloride layer is brown - indicates bromine.
(c) If the carbon tetrachloride layer is violet - indicates iodine.
*If nitrogen and/or sulphur are also present, the addition of silver nitrate to the acidified 'fusion' solution will precipitate silver cyanide and/or silver sulphide in addition to the silver halides. The removal of hydrogen cyanide and/or hydrogen sulphide is effected by boiling the 'fusion' solution. GROUP CLASSIFICATION TESTS
D. Solubility tests
The solubility of the unknown in the following reagents provides very useful information. In general, about 3 mL of the solvent is used with 0.1 g or 0.2 mL (2 - 3 drops) of the substance. The class of compound may be indicated from the following table:
SOLUBILITY TABLE
REAGENT AND TEST CLASS GROUP OF COMPOUNDS
Soluble in cold or hot water. (If the unknown is soluble do NOT perform solubility tests below) Neutral, acidic or basic. (Test with litmus or universal indicator paper) Lower members of series. Neutral, e.g. alcohols; Acidic, e.g. acids, phenols; Basic, e.g. amines
Soluble in dil. HCl Basic Most amines (except III amines with only aromatic groups
Soluble in dil. NaOH Acidic Most acids, most phenols.
Soluble in NaHCO3 Strongly acidic Most carboxylic acids.
Insoluble in water, acid and alkali Neutral Hydrocarbons, nitrohydro-carbons, alkyl or aryl halides, esters and ethers. Higher molecular weight alcohols, aldehydes and ketones
E. Group Classification Tests
From the previous tests it is often possible to deduce the functional groups present in the unknown compound. Consult i.r. spectra when available.
Individual tests are then performed to identify and confirm the functional groups present.
NOTE:
1. Students are strongly advised against carrying out unnecessary tests, since not only are they a waste of time but also increase the possibility of error. Thus it is pointless to first test for alcohol or ketone in a basic compound containing nitrogen! Instead tests for amines, etc. should be done on such a compound.
2. A systematic approach cannot be overemphasised in group classification tests to avoid confusion and error.
Once the functional group has been identified, reference is made to tables in a book on organic analysis, for assessing possibilities and for the preparation of suitable solid derivatives.
It should be noted that whilst two substances with the same functional group may sometimes have very similar b.p. or m.p., solid derivatives canusually be chosen from the literature, with m.p. differences of about 10 (or more), which distinguish between the two possibilities.
Example:
COMPOUND B.P. DERIVATIVES (M.P.)
2,4-DNPH SEMICARBAZONE
Diethyl ketone 102 156 139
Methyl n-propyl ketone 102 144 112
G. Preparation of derivatives
The final characterisation of the unknown is made by the preparation of suitable solid derivatives. The derivative should be carefully selected and its m.p. should preferably be between 90 - 150 for ease of crystallisation and m.p. determination.
Preparation of one derivative should be attempted. The derivative should be purified by recrystallisation, dried and the m.p. determined. Derivatives should be submitted correctly labelled for assessment together with the record.
Recording of Results
The results should be recorded in a systematic manner. Results should be recorded in the practical book at the time (not written up afterwards).
A record should be made of every test carried out, no matter whether a NEGATIVE RESULT HAS BEEN OBTAINED.
Tests for unsaturation
1. Cold dilute potassium permanganate solution.
2. Solution of bromine in carbon tetrachloride.
Tests for compounds containing nitrogen
1. Amines
(a) Nitrous acid.
(b) Confirmatory tests.
2. Compounds which give amines or ammonia on acid or alkaline hydrolysis:
Amides, substituted amides, anilides, nitriles.
3. Compounds which give amines on reduction:
Nitro, nitroso, azo, hydrazo, nitriles.
Tests for compounds containing C, H and possibly oxygen
1. Carboxylic acids
Na2CO3 or NaHCO3 solution liberate carbon dioxide.
2. Phenols
(a) Sodium hydroxide solution (soluble). Insoluble in and no CO2 from NaHCO3 (except when electron attracting groups present, e.g. 2,4-dinitrophenol).
(b) Ferric chloride solution.
(c) Bromine water.
3. Aldehydes and Ketones
(a) 2,4-dinitrophenylhydrazine (as Brady's reagent) for C=O.
(b) Iodoform test for CH3CO-.
4. Aldehydes only (reducing properties)
(a) Fehling's solution.
(b) Tollen's reagent (ammoniacal AgNO3 solution).
(c) Jones reagent.
5. Alcohols
(a) Lucas' reagent to distinguish I, II and III alcohols.
(b) Jones reagent.
(c) Metallic sodium (use dry liquid and dry tube).
6. Sugars
(a) Molisch's test.
7. Esters
(a) Hydroxamic acid test.
(b) Hydrolysis.
________________________________________
Write up of the identification of an unknown organic compound
Compound containing C, H ,N, Halogens, S
Physical characteristics ...................... (solid, liquid, gas, colour, odour, etc.)
Ignition test .............................. (aromatic or aliphatic)
Physical constant ........................ (boiling point or melting point)
Solubility tests (in tabular form)
Group classification tests (in tabular form)
Test Observation Inference
From the above tests and observations the given compound is probably a
.........................(acid, phenol, aldehyde, etc.)
Consultation of literature (Possibilities) M.P. of derivative
(a)
(b)
(c)
Preparation of derivative (method of preparation)
Observed m.p. of derivative
Lit. m.p. of derivative
Result
Compound No. ........................ is ............................
(give formula)
TESTS FOR FUNCTIONAL GROUPS
I. UNSATURATED COMPOUNDS
Two common types of unsaturated compounds are alkenes and alkynes characterised by the carbon-carbon double and triple bond, respectively, as the functional group. The two common qualitative tests for unsaturation are the reactions of the compounds with (a) bromine in carbon tetrachloride and (b) potassium permanganate.
(a) 2% Bromine in carbon tetrachloride
Dissolve 0.2 g (or 0.2 mL) of the compound in 2 mL of carbon tetrachloride or another suitable solvent and add the solution dropwise to 2 ml of 2% bromine solution in carbon tetrachloride and shake. e. g.
Rapid disappearance of the bromine colour to give a colourless solution is a positive test for unsaturation.
NOTE: The reagent is potentially dangerous. Keep it off your skin and clothes; protect your eyes and nose.
(b) 2% Aqueous potassium permanganate
Dissolve 0.2 g (or 0.2 mL) of the substance in 2 mL of water (acetone may also be used as solvent). Add the potassium permanganate solution dropwise and observe the result.
e.g.
For a blank determination, count the number of drops added to 2 mL of acetone before the colour persists. A significant difference in the number of drops required in the two cases is a positive test for unsaturation.
II. COMPOUNDS CONTAINING NITROGEN
1. Amines
(a) Reaction with nitrous acid Dissolve the amine (0.5 mL) in concentrated acid (2.0 mL) and water (3 mL) and cool the solution to 0 - 5 in an ice-bath for 5 minutes. Add a cold solution (ice-bath) of sodium nitrite (0.5 g) in water (2.0 mL) from a dropper, with swirling of the test tube, still keeping the mixture in the ice-bath.
AMINE REACTION
Aliphatic N2 evolved.
RNH2 + HNO2 -> ROH + N2 + H2O
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Aromatic Diazonium salt is formed.
ArNH2 + HNO2 -> ArN=N+
Add the cold diazonium solution and with swirling
to a cold solution of 2-naphthol (0.2 g) in 5% NaOH
solution (2 mL). An orange-red azo dye is formed.
__________________________________________________________________
aliphatic and Yellow oily nitrosamines are generally formed.
aromatic R2NH + HNO2 -> R2N-NO
__________________________________________________________________
(b) Reaction with benzenesulphonyl chloride
Benzenesulphonyl chloride reacts with primary and secondary but not with tertiary amines to yield substituted sulphonamides.
e.g. (a) C6H5SO2Cl + H-NHR + NaOH -> C6H5SO2NHR + NaCl + H2O
(b) C6H5SO2Cl + H-NR2 + NaOH -> C6H5SO2NR2 + NaCl + H2O
The substituted sulphonamide formed from a primary amine dissolves in the alkali medium whilst that produced from a secondary amine is insoluble in alkali.
Place 0.5 mL (or 0.5 g) of the compound, 15 - 10 mL of 5% NaOH and 1 mL of benzenesulphonyl chloride in a test tube, stopper the tube and shake until the odour of the sulphonyl chloride has disappeared. The solution must be kept alkaline (if no reaction has occurred, the substance is probably a tertiary amine).
If a precipitate appears in the alkaline solution, dilute with about 10 mL of water and shake; if the precipitate does not dissolve, a secondary amine is indicated.
If there is no precipitate, acidify it cautiously to congo red with concentrated hydrochloric acid (added dropwise): a precipitate is indicative of a primary amine.
2. Amides R-CO-NH2
Simple primary amides can be decomposed by boiling with alkali and thereby evolving ammonia.
e.g. CH3-CO-NH2 + NaOH -> CH3-CO2- Na+ + NH3 ¬
Boil 0.5 g of the compound with 5 mL of 10% sodium hydroxide solution and observe whether ammonia is evolved.
III. COMPOUNDS CONTAINING C, H AND POSSIBLY OXYGEN
1. Carboxylic acids - test with 5% aq. NaHCO3
R-CO2H + NaHCO3 -> R-CO2- Na+ + CO2 ¬ + H2O
Sodium hydrogen carbonate reacts with carboxylic acids to give the sodium salt of the acid and liberates carbon dioxide. If the acid is insoluble in water and the reaction is sluggish dissolve the acid in methanol and add carefully to a saturated sodium hydrogen carbonate solution, when a vigorous effervescence will be observed.
2. Phenols [Soluble in NaOH and produce no CO2 from NaHCO3]
(a) Bromine water
Phenols are generally highly reactive towards electrophilic reagents and are readily brominated by bromine water. e.g.
Dissolve or suspend about 0.05 g of the compound in 2 mL of dilute hydrochloric acid and add bromine water dropwise until the bromine colour remains. A white precipitate of the bromophenol may form. Solid bromophenol derivatives can be used for the confirmation of the structure of a phenol (cf the preparation of derivatives).
(b) Ferric chloride test
Most phenols react with iron (III) chloride to form coloured complexes. The colours vary - red, purple, blue or green - depending on various factors, e.g. the phenolic compound used, the solvent, concentration. Since some phenols do not give colours, a negative test must not be taken as significant without supporting information.
Dissolve 0.05 g of the compound in 2 mL water (or a mixture of water and ethanol if the compound is not water-soluble) and add an aqueous solution of ferric chloride dropwise. Observe any colour changes which may occur.
3. Aldehydes and ketones
(a) 2,4-Dinitrophenylhydrazine (as Brady's reagent)
A test for the carbonyl group (C=O) in aldehydes and ketones. 2,4- Dinitrophenylhydrazine gives sparingly soluble yellow or red 2,4-dinitrophenylhydrazones with aldehydes and ketones.
Add 3 mL of the reagent to 2 drops of the compound in a test tube and shake. If no precipitate forms immediately, warm and allow to stand for 5 - 10 minutes. A crystalline precipitate indicates the presence of a carbonyl compound.
The bench reagent is very dilute and is intended for qualitative tests only and should not be used in the preparation of a derivative for identification purposes.
(b) Iodoform test for CH3CO-
Dissolve 0.1 g (or 5 drops) of the compound in 2 mL of water; if it is insoluble in water add sufficient dioxan to produce a homogeneous solution. Add 2 mL of 5% NaOH solution and then introduce the potassium iodide - iodine reagent dropwise with shaking until a definite dark colour of iodine persists. Allow to stand for 2 - 3 minutes; if no iodoform separates at room temperature, warm the test tube in a beaker of water at 60 . Add a few more drops of the iodine reagent if the faint iodine colour disappears. Continue the addition of the reagent until a dark colour is not discharged after 2 minutes heating at 60 . Remove the excess of iodine by the addition of a few drops of dilute sodium hydroxide solution with shaking, dilute with an equal volume of water, and allow to stand for 10 minutes. The test is positive if a yellow precipitate of iodoform is deposited. Filter off the yellow precipitate, dry upon pads of filter paper and determine the m.p. Iodoform melts at 120 (it can be recrystallised from methanol- water).
The reaction is given by acetaldehyde and simple methyl ketones. Alcohols containing the CH3CHROH group will be oxidised under the reaction conditions and also give a positive test.
4. Aldehydes only (reducing properties).
(a) Fehling's solution
Aldehydes reduce Fehling's solution to yellow or red copper (I) oxide.
Preparation of the reagent: Mix equal volumes of Fehling's solution solution I (aqueous alkaline potassium tartrate) and Fehling's solution II (copper sulphate solution).
Add 2 drops (or 0.05 g) of the compound and 2 - 3 drops of the reagent and heat on a boiling water bath for 3 - 4 minutes.
The test is positive for aliphatic aldehydes, but is often indecisive for aromatic aldehydes where Jones' Reagent is often useful (see 5).
(b) Tollen's reagent (Ammonical silver nitrate solution)
Aldehydes are readily oxidised to carboxylic acids and will reduce Tollen's reagent to produce a silver mirror on the inside of a clean test tube.
FIRST clean up a test tube with a little hot nitric acid (fume cupboard) and rinse with distilled water.
Preparation of the reagent: To 1 mL of silver nitrate solution add a few drops of sodium hydroxide. Then add dilute ammonium hydroxide dropwise until the precipitate just dissolves.
Add 2 - 3 drops of the compound in methanol to 2 - 3 mL of Tollen's solution contained in a very clean test tube. If no reaction takes place in the cold, warm gently in a water bath.
CAUTION: After the test, pour the contents of the test tube into the sink and wash the test tube with dilute nitric acid. Any silver fulminate present, which is highly explosive when dry, will be destroyed.
(c) Jones Reagent (See section under alcohols).
5. Alcohols
The tests for the hydroxyl group not only detect the presence of the group, but may also indicate whether it is primary, secondary or tertiary.
(a) Jones Reagent (CrO3-H2SO4 in H2O)
This reagent distinguishes primary and secondary alcohols from tertiary alcohols; the test is based on the much greater resistance to oxidation of tertiary alcohols compared to the other two types. Aldehydes also give a positive test.
Place 1 mL of acetone in a test tube and dissolve one drop of a liquid or ca 10 mg of a solid alcohol or aldehyde in it. Add one drop of the reagent to the acetone solution and shake the tube to mix the contents. Primary and secondary alcohols react within two seconds as indicated by the disappearance of the orange colour of the reagent and the formation of a green or blue-green precipitate or emulsion.
Tertiary alcohols do not react even after 3 minutes.
(I) RCH2OH -> RCHO -> RCO2H
(II) R2CHOH -> R2C=O
(III) R3COH -> no visible reaction.
(b) Lucas' Reagent [ZnCl2 - conc. HCl]
This reagent converts alcohols into the corresponding alkyl chlorides. Zinc chloride (a Lewis acid) increases the reactivity of alcohols towards acid. The test depends on the rate of reaction of primary, secondary, and tertiary alcohols with the reagent at room temperature.
(I) RCH2OH -> no reaction at room temperature.
(II) R2CHOH -> R2CHCl + H2O (1 hour or maybe longer)
(III) R3COH -> R3CCl + H2O (immediately)
To 1 mL of the alcohol in a small test tube add 6 mL of Lucas' reagent at room temperature. Close the tube with a cork, shake and allow to stand.
(i) Primary alcohols - the aqueous phase remains clear (except allyl alcohol - droplets after 7 minutes).
(ii) Secondary alcohols - very slow reaction (~ 1 hour or maybe longer) when droplets of alkyl chloride may be seen.
(iii) Tertiary alcohols - very fast reaction and droplets of the alkyl chloride formed almost immediately.
6. Sugars, Carbohydrates
Molisch's Test
This is a general test for carbohydrates. Dissolve 20 - 30 mg of the compound in 2 mL water and add 0.5 mL of the reagent (a 20% solution of 2-naphthol in ethanol). Pour 2 mL of concentrated sulphuric acid from a dropper carefully down the side of the tube so that the acid forms a layer beneath the aqueous solution without mixing with it. A red colouration, changing to dark purple forms at the interface. Carry out a second test on a blank solution.
7. Esters
Hydroxamic acid test
R-CO-OR' + H2N-OH -> R-CO-NH-OH + R'-OH
Esters react with hydroxylamine in the presence of sodium hydroxide to form the sodium salt of the corresponding hydroxamic acid. On acidification and addition of ferric chloride the magenta-coloured iron (III) complex of the hydroxamic acid is formed.
It is always advisable to ensure that an unknown compound does not give a colour with iron (III) chloride before carrying out the hydroxamic acid test.
Procedure for hydroxamic acid test
(a) Ferric chloride test
Dissolve a drop or a few small crystals of the compound in 1 mL of 95% ethanol (rectified spirit) and add 1 mL of M hydrochloric acid. Note the colour produced when 1 drop of 5% iron (III) chloride is added to the solution. If a pronounced violet, blue, red or orange colour is produced, the hydroxamic acid test described below is NOT APPLICABLE.
(b) Hydroxamic acid test
Mix 1 drop or several small crystals (ca 0.05 g) of the compound with 1 mL of 0.5 M hydroxylamine hydrochloride in 95% ethanol and add 0.2 mL of 6 M aqueous sodium hydroxide. Heat the mixture to boiling and after the solution has cooled slightly add 2 mL of M hydrochloric acid. If the solution is cloudy, add 2 mL of 95% ethanol. Observe the colour produced when 1 drop of 5% iron (III) chloride solution is added. If the resulting colour does not persist, continue to add the reagent dropwise until the observed colour pervades the entire solution. Usually only 1 drop of the iron (III) chloride solution is necessary. Compare the colour with that produced in test (a). A positive test will be a distinct burgundy or magenta colour as compared with the yellow colour observed when the original compound is tested with iron (III) chloride solution in the presence of acid. It is often advisable to conduct in parallel the test with, say, ethyl acetate, to ensure that the conditions for this test are correct.
THE PREPARATION OF DERIVATIVES OF ORGANIC COMPOUNDS
The preliminary examination and group classification tests indicate the particular class (functional group) to which an unknown organic compound may belong. Further characterisation and identification depends on the selection and preparation of a suitable solid derivative and accurate determination of its melting point (best, between 90 - 150 ).
The following table lists some of the classes of organic compounds and a selection of derivatives that may be prepared to characterise them. Check with the tables of melting points in Vogel which derivatives are most suitable for the characterisation of your particular compound.
CLASS OF COMPOUND DERIVATIVES
1. ALCOHOLS 3,5-dinitrobenzoate
2. PHENOLS benzoate, acetate, bromo-derivative
3.ALDEHYDES AND KETONES semicarbazone, 2,4-dinitrophenyl-hydrazone, oxime
4. ACIDS anilide, amide, p-toluidide.
5. AMINES benzoyl, acetyl and sulphonamide derivatives
METHODS FOR THE PREPARATION OF DERIVATIVES
ALCOHOLS
(i) 3,5-Dinitrobenzoates
3,5-Dinitrobenzoyl chloride is usually partially hydrolysed and should be prepared in the pure state by heating gently a mixture of 3,5-dinitrobenzoic acid (1 g) and phosphorus pentachloride (1.5 g) in a dry test tube, until it liquifies (5 min).* The liquid is poured on a dry watch glass and allowed to solidify. The phosphoryl chlorides are removed by pressing the solid with a spatula on a wad of filter paper. The residual acid chloride is suitable for immediate use in the preparation of the derivatives.
*Work under fume hood. Fumes are irritating to the eyes and nose.
The 3,5-dinitrobenzoyl chloride is mixed with the alcohol (0.5 - 1 mL) in a loosely corked dry test tube and heated on a steam bath for about 10 min. Secondary and tertiary alcohols require up to 30 min. On cooling add 10 mL sodium hydrogen carbonate solution, stir until the ester crystallises out, and filter at the pump. Wash with a little carbonate solution, water and suck dry. Recrystallise from the minimum hot ethanol or light petroleum. Cool slowly to avoid the formation of oily droplets of your ester.
PHENOLS
(i) Benzoates (Schötten-Baumann method).
To the phenol (0.5 g) is added 5% sodium hydroxide (10 mL) in a well-corked boiling tube or a small conical flask. Benzoyl chloride (2 mL) is added in small quantities at a time, and the mixture shaken vigorously with occasional cooling under the tap or in ice-water. After 15 min the solid benzoate separates out: the solution should be alkaline at the end of the reaction; if not alkaline, or if the product is oily, add a solid pellet of sodium hydroxide and shake again. Collect the benzoate, wash thoroughly with cold water, and recrystallise from alcohol or light petroleum.
(ii) Acetates
Acetates of many simple phenols are liquids; however, this is a suitable derivative for polyhydric and substituted phenols. The phenol (0.5 g) is dissolved in 10% sodium hydroxide solution and an equal quantity of crushed ice is added, followed by acetic anhydride (2 mL). The mixture is vigorously shaken in a stoppered test tube until the acetate separates. The product is filtered and recrystallised from alcohol.
(iii) Bromo derivatives
The phenol (0.3 g) is suspended in dilute hydrochloric (10 mL) and bromine water added dropwise until no more decolourisation occurs. The bromo derivative which precipitates out is filtered off and recrystallised from alcohol.
ALDEHYDES AND KETONES
(i) Semicarbazones
Dissolve semicarbazide hydrochloride (1 g) and sodium acetate (1.5 g) in water (8 - 10 mL), add the aldehyde or ketone (0.3 mL) and shake. Shake the mixture for a few minutes and then cool in ice-water. Filter off the crystals, wash with a little cold water and recrystallise from methanol or ethanol.
(ii) 2,4-Dinitrophenylhydrazones
Suspend 0.25 g of 2,4-dinitrophenylhydrazine in 5 mL of methanol and add 0.5 mL of concentrated sulphuric acid cautiously. Filter the warm solution and add a solution of 0.2 g of the carbonyl compound in 1 mL of methanol. Recrystallise the derivative from methanol, ethanol or ethyl acetate.
(iii) Oximes
Hydroxylamine hydrochloride (0.5 g) is dissolved in water (2 mL). 10% sodium hydroxide (2 mL) and the carbonyl compound (0.2 - 0.3 g) dissolved in alcohol (1 - 2 mL) are added, the mixture warmed on a steam bath for 10 min and then cooled in ice. Crystallisation is induced by scratching the sides of the test tube with a glass rod. The oximes may be crystallised from alcohol.
ACIDS
(i) Amides, anilides and p-toluidides
The acid (0.5 g) is refluxed with thionyl chloride (2 - 3 mL) in a fume cupboard for about 30 mins.* It is advisable to place a plug of cotton wool in the top of the reflux condenser to exclude moisture. The condenser is removed and the excess of thionyl chloride is distilled off (b.p. 78 ). The acid chloride thus produced is treated with concentrated ammonia solution (5 mL) or aniline (0.5 - 1 mL) or p-toluidine (0.5 - 1 g), when the solid derivative separates out. It is collected and recrystallised from alcohol adding decolourising charcoal if found necessary.
*Alternately use PCl5 to form the acid chloride.
AMINES
(i) Acetyl derivatives (acetamides)
Reflux gently in a small dry flask under a dry condenser the amine (1 g) with acetic anhydride (3 mL) for 15 min. Cool the reaction mixture and pour into 20 mL cold water. Boil to decompose the excess acetic anhydride. Cool and filter by suction the insoluble derivative. Recrystallise from ethanol.
(ii) Benzoyl derivatives (benzamides)
Suspend 1 g of the amine in 20 mL of 5% aqueous sodium hydroxide in a well-corked flask, and add 2 mL benzoyl chloride (fume hood!), about 0.5 mL at a time, with constant shaking. Shake vigorously for 5 - 10 min until the odour of the benzoyl chloride has disappeared. Ensure that the mixture remains alkaline. Filter off the solid derivative, wash with a little cold water and recrystallise from ethanol.
(iii) Benzenesulphonamides
To 1 g of the amine in 20 mL of 5% sodium hydroxide solution in a well-corked flask add 1 mL benzenesulphonyl chloride (fume hood!). Shake the mixture until the odour of the sulphonyl chloride disappears. Check that the solution is alkaline. Acidify if necessary to obtain the precipitated derivative. Concentrated hydrochloric acid added dropwise should be used. Filter the product, wash with a little cold water and suck dry. Recrystallise from ethanol.
Inorganic Quantitative Analysis
Gravimetric Estimation of Ba2+, Fe2+, Zn2+ and Cu2+
All Gravimetric analyses rely on some final determination of weight as a means of quantifying an analyte. Since weight can be measured with greater accuracy than almost any other fundamental property, gravimetric analysis is potentially one of the most accurate classes of analytical methods available. These methods are among the oldest of analytical techniques, and they may be lengthy and tedious. Samples may have to be extensively treated to remove interfering substances. As a result, only a very few gravimetric methods are currently used in environmental analysis.
There are four fundamental types of gravimetric analysis: physical gravimetry, thermogravimetry, precipitative gravimetric analysis, and electrodeposition. These differ in the preparation of the sample before weighing of the analyte. Physical gravimetry is the most common type used in environmental engineering. It involves the physical separation and classification of matter in environmental samples based on volatility and particle size (e.g., total suspended solids). With thermogravimetry, samples are heated and changes in sample mass are recorded. Volatile solids analysis is an important example of this type of gravimetric analysis. As the name implies, precipitative gravimetry relies on the chemical precipitation of an analyte. Its most important application in the environmental field is with the analysis of sulfite. Electrodeposition involves the electrochemical reduction of metal ions at a cathode, and simultaneous deposition of the ions on the cathode.
Common Procedures in Gravimetric Analysis
a. Drying to a Constant Weight
All solids have a certain affinity for water, and may absorb moisture from the laboratory air. Reagents that readily pick up water are termed hygroscopic. Those that absorb so much water that they will dissolve in it and form a concentrated solution are called deliquescent (e.g., sodium hydroxide, trichloroacetic acid). These types of substances will continually increase in weight while exposed to the air. For this reason, many types of laboratory procedures require that a sample be dried to a constant, reproducible weight (i.e., absorbed moisture removed to some standard, low level). This is especially important for the gravimetric methods. Generally, the sample is dried in a 103 C to 110 C oven for about 1 hour and allowed to cool to room temperature in a desiccator. It is then weighed, and heated again for about 30 minutes. The sample is cooled and weighed a second time. The procedure is repeated until successive weighings agree to within 0.3 mg.
b. Description and Use of the Analytical Balance
The analytical balance is the most accurate and precise instrument in an environmental laboratory. Objects of up to 100 grams may be weighed to 6 significant figures. Volumetric glassware is accurate to no more than 4 significant figures, and the accuracy of complex analytical methods rarely justifies more than 2 significant figures. Analytical balances are generally used for gravimetric analyses, and for the preparation of standard solutions.
Summary of Gravimetric Methods for Environmental Analysis
Some gravimetric methods are in generally using for the analysis of waters and wastewaters.
Type Analyte Pretreatment
Physical Total Solids Evaporation
Suspended Solids Filtration
Dissolved Solids Filtration + Evaporation
Oil & Grease extraction with C2Cl3F3 + distillation of solvent
Surfactants extraction into ethylacetate + evaporation
Thermal Volatile Solids Evaporation + 550`C for 15 min
Volatile Suspended Solids Filtration + 550`C for 15 min
Precipitative Mg with Diammonium hydrogen phosphate and final pyrolysis
Na with zinc uranyl acetate
Silica precipitation/ ignition/ volatilization (with HF)
SO4 with Ba
PHYSICAL GRAVIMETRY
1. Total, Dissolved and Suspended Solids
a. Definitions
Total solids (TS) is generally defined as all matter in a water or wastewater sample that is not water. Because solids are not a specific chemical compound, but rather a diverse collection of dissolved and particulate matter, their concentration cannot be determined in an unambiguous way. Instead, they must be defined by the procedure used to estimated their concentration. Total solids may be differentiated according to size into total dissolved solids (TDS) and total suspended solids (TSS). Once again, this is an operational distinction, whereby all solids passing through filter paper of a certain pore size (e.g., 1.5 microns, Whatman #934AH) are called dissolved, and those retained are termed suspended.
b. Significance to Environmental Engineering
Most of the impurities in potable waters are in the dissolved state, principally as inorganic salts. Thus, the parameters, 'total solids' and especially 'total dissolved solids' are of primary importance here. Waters containing high concentrations of inorganic salts are not suitable as sources of drinking water, because such materials are often difficult to remove during treatment. Finished drinking waters containing more than 1000 mg/L TDS are generally considered unacceptable. Waters of this type may also be unsuitable for agricultural purposes due to the harmful effects of high ionic concentrations on plants. In most natural waters, the TDS (total dissolved solids) concentration correlates well with total hardness (i.e., [Ca] + [Mg]). This is useful in assessing the corrosivity of a water and the need for softening.
The total suspended solids (TSS) content of natural waters is of interest for the purpose of assessing particle bed load and transport. High concentrations of suspended matter may be detrimental to aquatic life. In theory, TSS could be used for assessing particle removals during water treatment. However, nearly always the concentration of colloidal particles in water is measured as turbidity since this latter technique is faster and more precise.
Some Typical Solids Concentrations
Source Concentration (mg/L)
Low Avg High
NATURAL WATERS
Fresh TDS 20 120 1,000
Brines TDS 5,000 300,000
DOMESTIC WASTEWATER
Raw TDS 350 600 900
VDS 165 285 600
TSS 100 200 350
VSS 75 135 215
Secondary Effluent TSS 10 30 60
Activated Sludge Mixed
Liquor (conventional) TSS 1,500 3,000
Activated Sludge Mixed
Liquor (extended aeration) TSS 3,000 6,000
Primary Sludge TSS 20,000 70,000
Secondary Sludge TSS 5,000 12,000
STORM WATER TSS 5 300 3,000
Procedures
Total Solids (Total Residue). Total solids is determined by the final weight of a dried sample (minus tare) divided by the original sample volume. Evaporating dishes of platinum, vycor or porcelain may be used. Platinum is preferred, because it is more inert than the other two, and can be heated to a constant weight more easily. However, platinum is very expensive, so porcelain is often used. Porcelain is difficult to bring to a constant weight, and its use should be avoided. Space permitting, evaporating dishes should be stored in a desiccator so as to avoid the collection of dust and absorption of moisture while not in use. The precision of this method has been estimated to be 4 mg or 5%. However, settled wastewater may give better precision, on the order of 1 mg.
1. Preheat a 100 mL evaporating dish at 550 50 C for 1 hour, cool in a in a drying oven or in the open air (protected from dust) for 15-20 minutes, bring to room temperature in a desiccator, and weigh. Repeat until a constant weight is achieved.
2. Measure 75 mL of sample or a volume sufficient to yield 200 mg TS, whichever is less. Add this to the preweighed dish and evaporate to dryness in a drying oven set at 98 C. Alternatively, a steam bath may be used.
3. Dry for an additional hour at 103-105 C.
4. Cool in a desiccator and weigh.
Dissolved Solids (Filtrable Residue). Dissolved solids may be determined directly by analysis of the filtered sample for total solids, or indirectly by determining the suspended solids and subtracting this value from the total solids. When using the direct method, final drying may be conducted at one of two temperatures.
1. Analyze the filtrate in accordance with the total solids procedure.
2. Final drying (1 hour period) may be conducted at either 103-105 C or 180 2 C.
Suspended Solids (Non-filtrable Residue). Suspended solids is measured directly by drying and weighing the solids retained during filtration.
1. Dry this filter at 103-105 C for 1 hour, and cool in a desiccator.
2. Weigh the filter, then pass a water sample of sufficient volume to yield 50-200 mg suspended solids through it. Smaller volumes will result in reduced accuracy.
3. Dry for at least one hour at 103-105 C.
4. Cool in a desiccator and weigh.
C. THERMOGRAVIMETRY AND COMBUSTION ANALYSIS
Thermogravimetry and combustion analysis involve the heating of a sample to 500 C or more with the oxidation and/or volatilization of some of the sample constituents. Either the change of sample weight is determined (thermogravimetry), or the combustion gases are trapped and weighed (combustion analysis). With thermogravimetric methods, it is especially important to return the sample to room temperature before weighing. Otherwise the differences in temperature will create convection currents around the balance pan, which will severely disrupt method accuracy. A steady increase in apparent weight while the sample is on the pan indicates a problem of this type. Large vessels and samples will require longer cooling times to dissipate their excess heat.
Volatile Solids and Fixed Solids
Fixed solids are those that remain as residue after ignition at 550 C for 15 minutes. The weight of material lost is called the volatile solids. Thus the total operational definition for volatile solids would be: all matter lost upon ignition at 550 C for 15 minutes, but not lost upon drying at 103-105 C for 1 hour. The portion lost upon ignition is generally assumed to be equivalent to the organic fraction. The portion remaining is considered the inorganic fraction. For waters of moderate to high hardness, most of this is calcium carbonate which decomposes only at temperatures exceeding 800 C. When igniting a filter with suspended matter, one must be especially careful of the temperature; above 600 C glass fiber filters begin to melt and can loose a significant amount of weight in 15 minutes.
Combining the fractionations resulting from ignition and filtration, one arrives at a total of 9 separate categories: total solids (TS), fixed solids, volatile solids, total dissolved solids (TDS), fixed dissolved solids, volatile dissolved solids, total suspended solids (TSS), fixed suspended solids, and volatile suspended solids (VSS). In practice, only four of these (TS, TDS, TSS, and VSS) are commonly used. When comparing fixed solids with inorganic content, one would expect positive bias from incomplete oxidation of organic matter, and negative bias from decomposition of certain inorganics. Ammonium salts may be lost during low temperature drying or upon ignition. Most others are stable under the conditions used for volatile solids determination with the exception of magnesium carbonate. Volatile solids may be effected by these as well as loss of recalcitrant waters of crystallation (positive bias), and previous losses of organic matter to volatilization during low-temperature drying (negative bias). A modest interlaboratory study found an average standard deviation of 11 mg/L on a sample of 170 mg/L volatile solids.
MgCO3 ------------> MgO + CO2 ¬
Ammonium compounds (often present in sludge in the form of ammonium bicarbonate) may be lost during low temperature drying and therefore should not introduce a bias in volatile solids
NH4HCO3 ---------> NH3 ¬ + H2O ¬ + CO2 ¬
1. Dry and weigh a vessel containing the solids to be analyzed. For volatile and fixed suspended solids analysis, the filter (with residue) prepared for suspended solids analysis and dried to a constant weight may be used. For volatile and fixed total (or dissolved) solids, the evaporating dish (with residue) prepared for total (or dissolved) solids analysis dried to a constant weight should be used.
2. Ignite the sample and vessel in a preheated muffle furnace set at 550 50 C for 15-20 min (water & wastewater) or 1 hour (sludge, sediment & soil).
3. Cool for 15 minutes in the open air in an area protected from dust.
4. Place vessel in a desiccator for final cooling to room temperature and weigh. Due to the approximate nature of this test samples are not generally re-heated and dried to a constant weight.
D. PRECIPITATIVE GRAVIMETRIC ANALYSIS
Precipitative gravimetric analysis requires that the substance to be weighed be readily removed by filtration. In order for a non-filtrable precipitate to form, it must be supersaturated with respect to its solubility product constant. However, if it is too far above the saturation limit, crystal nucleation may occur at a rate faster than crystal growth (the addition of molecules to a crystal nucleus, eventually forming a non-filtrable crystal). When this occurs, numerous tiny micro-crystals are formed rather than a few large ones. In the extreme case, micro-crystals may behave as colloids and pass through a fibrous filter. To avoid this, precipitating solutions may be heated. Because the solubility of most salts increases with increasing temperature, this treatment will lower the relative degree of supersaturation and slow the rate of nucleation. Also, one might add the precipitant slowly with rapid mixing to avoid the occurrence of locally high concentrations.
It is very important that the precipitate be pure and have the correct stoichiometry. Coprecipitation occurs when an unwanted ion or molecule becomes trapped in the precipitate. This may be due to inclusion or occlusion. Inclusion is the term used of a single subsitution in the crystal lattice by an ion of similar size. Occlusion refers to the physical trapping of a large pocket of impurities within the crystal. One technique for minimizing these problems is to remove the mother liquor, re-dissolve the precipitate and then re-precipitate. The second time the mother liquor will contain fewer unwanted ions capable of coprecipitation.
Sulfate Determination
The method of choice for sulfate in waters and wastewaters is the precipitative gravimetric procedure using barium. If Ba(+II) is added in excess under acidic conditions, BaSO4 is precipitated quantitatively. The reaction is allowed to continue for 2 hours or more at 80-90oC. This is to encourage the formation of BaSO4 crystals (non-filtrable) from the initially formed colloidal precipitate (partially filtrable). The precipitate is washed, and then dried at 800`C for 1 hour. Low pH is needed to avoid the precipitation of BaCO3 and Ba3(PO4)
Ba+2 + SO4-2 = BaSO4(ppt.)
Chloride Determination
Chloride may be determined by precipitation with silver. Interfering ions likely to form insoluble silver salts are the other halogens (bromide, iodide), cyanide, and reduced sulfur species (sulfite, sulfide, and thiosulfate). Fortunately, the reduced sulfur compounds can be pre-oxidized with hydrogen peroxide, and the others are rarely present at high concentrations. Although AgCl can be determined gravimetrically, the recommended procedure for water and wastewater is to use a volumetric procedure with chromate as an indicator.
3. Inorganic Synthesis
(i) Synthesis of Cuprous Chloride (work in a fume hood)
(a) Cu2+ + Cu + 2 Cl- → 2 CuCl
Material: CuSO4.5H2O, Cu, NaCl, HCl, Na2SO3, CH3COOH
Procedure: Prepare a solution of 10 g of powdered CuSO4.5H2O and 15 ml of
concentrated HCl in a 250-ml round-bottom flask, add 4 g of NaCl and heat to boiling.
Cover the flask with a little funnel. Add copper to hot solution in small portions. The
green colour of the solution will turn to yellow. Filter off the remaining Cu. Pour the
solution to one litre of cold water with 2 g of Na2SO3.
Wash the precipitated cuprous chloride 2 – 3 times with a solution of 1.5 g of
Na2SO3, 6 ml of HCl and 300 ml of H2O by decantation. Filter out the precipitate and wash it with concentrated acetic acid. Dry it in a drying oven at 100 oC.
( b) 2 Cu2+ + SO2 + 2 Cl- + 2 H2O → 2 CuCl + SO42- + 4 H+
Material: CuSO4.5H2O, NaCl, HCl, SO2, CH3COOH
Procedure: Add 5 g of NaCl to the warm solution (70 oC) of 10 g CuSO4.5H2O and
bubble SO2 through the mixture. CuCl precipitates. Filter out the precipitate and wash
it with a solution of SO2 in water and then with concentrated acetic acid.
CuCl – white crystals, insoluble in water, on air turns to green alkali copper chloride
(ii) Preparation of chrome alum - KCr(SO4)2.12H2O
a) K2Cr2O7 + 3 SO2 + H2SO4 → KCr(SO4)2 + H2O
Material: K2Cr2O7, SO2, H2SO4
Procedure: Blow the SO2 gas through a gas washing bottle with K2Cr2O7 solution acidified with H2SO4. Do not permit the temperature to rise above 60 oC. Above this temperature complexes of chromium (III) sulphate are formed. These complexes contain sulphate in a non-ionisable form and are difficult to crystallise.
Bubble the unreacted gas through a washing bottle with 10 % NaOH solution. Test for the end of the reaction – to the sample of reduced solution in a test tube add small amount of Na2CO3 crystals and heat the mixture just below the boiling point. Let the precipitate settle, the solution over the precipitate has to be colourless. Set the solution aside to crystallise after the end of reduction.When the crystallization is complete, filter off the crystals and wash them with a small amount of water.
Transfer the product to a dry filter paper and let them dry in air.
b) K2Cr2O7 + 3 C2H5OH + 4 H2SO4 → KCr(SO4)2 + 3 CH3CHO + 7 H2O
Material: K2Cr2O7, C2H5OH, H2SO4
Procedure: Dissolve crushed K2Cr2O7 in diluted H2SO4 (1:3) and add, in small portions with stirring, calculated volume (+ 10 % excess) of C2H5OH. Do not permit the temperature to rise above 60°C. Continue like in procedure a).
KCr(SO4)2.12H2O – dark violet crystals, crystallize in regular octahedra, soluble in
water.
(ii) Preparation of potassium tris(oxalate)ferrate(III) trihydrate
Fe(OH)3 + 3 KHC2O4 → K3[Fe(C2O4)3] + 3 H2O
Material: FeSO4.7H2O or (NH4)2Fe(SO4)2.4H2O, K2C2O4, H2C2O4, HNO3, ethanol
Procedure: Dissolve 35 g of FeSO4.7H2O in 100 ml of warm water and add slowly
diluted HNO3 (1:1) to oxidize Fe2+. Add NH3(aq) to the solution until the precipitation
of Fe(OH)3 is completed. Let the precipitate settle and decant the liquid. Filter out the
precipitate and wash it with hot water. Prepare a hot solution of 44 g of KHC2O4
(calculate it as a mixture of K2C2O4 and H2C2O4) in 100 ml H2O. Add precipitate of
Fe(OH)3 in small portions to this solution. Filter the resulting solution and evaporate it
on a steam bath to crystallization. Filter out and wash the crystals on the Buchner
funnel with ethanol/water 1:1 and finally with acetone. Transfer the product to a dry
filter paper and let it dry in air.
K3[Fe(C2O4)3].3H2O – green crystals, photosensitive and decomposes due to
influence of light:
2 K3[Fe(C2O4)3] → K2[Fe(C2O4)2] + K2C2O4 + 2 CO2
(iv) Preparation of iron alum (Ferrous ammoninium sulphate)
2 FeSO4 + H2O2 +H2SO4 → Fe2(SO4)3
Fe2(SO4)3 + (NH4)2SO4 → 2 NH4Fe(SO4)2
Material: FeSO4.7H2O, H2O2, H2SO4, (NH4)2SO4
Procedure: Dissolve FeSO4.7H2O in water to a solution with w(FeSO4) = 0.15. Filter
this solution if necessary and carefully add concentrated H2SO4 in 10 % excess to the
stoichiometry. Then slowly add H2O2 (double amount compared to stoichiometry),
while stirring the mixture continuously. Heat the mixture to the boiling. Make sure that
Fe2+ was oxidized to Fe3+ by a reaction of sample with K3[Fe(CN)6]. Add further H2O2
if necessary.
Evaporate the solution on a steam bath to half the volume and add warm saturated
(by 60 oC) solution of calculated amount of (NH4)2SO4. Let the solution crystallize.
Put the crystals on a dry filter paper and let them dry in air.
NH4Fe(SO4)2.12H2O – colourless or light violet crystals, turn brown on air.
(v) Preparation of lead carbonate
Pb(CH3COO)2 + (NH4)2CO3 → PbCO3 + 2 CH3COONH4
Pb(NO3)2 + 2 NaHCO3 → PbCO3 + 2 NaNO3 + CO2 + H2O
Material: Pb(CH3COO)2 or Pb(NO3)2, (NH4)2CO3 or NaHCO3
Procedure: Add saturated (NH4)2CO3 or NaHCO3 solution (10 % excess to stoichiometry) to saturated Pb2+ salt solution while continuously stirring. Decant the precipitated product with water, filter it and dry at laboratory temperature.
PbCO3 – white powder, insoluble in water, easy soluble in acids and hydroxides,
decomposes by on heating. If Na2CO3 is used for precipitation, the alkali carbonate
2PbCO3.Pb(OH)2 creates.
(vi) Preparation of lead dioxide
a) Pb(NO3)2 + 2 NaOH → Pb(OH)2 + 2 NaNO3
Pb(OH)2 + 2 NaOH → Na2[Pb(OH)4]
Na2[Pb(OH)4] + CaOCl2 → PbO2 + CaCl2 + 2 NaOH + H2O
Material: Pb(NO3)2 or Pb(CH3COO)2, NaOH, CaOCl2 or NaClO, HNO3
Procedure: Dissolve 0.1 mol of Pb2+ salt in 300 ml of hot water and cool the solution.
If a small amount of substance remains undissolved, it will not affect the result. Add a solution of 20 g NaOH in 180 ml of water. First white precipitate of Pb(OH)2 appears, which dissolves in excess of NaOH to Na2[Pb(OH)4].
Mix 40 g of CaOCl2 with 50 ml of water and the necessary amount of Na2CO3 for
reaction with Ca2+. Add 200 ml of water to the mixture and filter it.
Add the filtrate to the boiling solution of Na2[Pb(OH)4] until the test for presence of
Pb2+ in the final solution is negative.
Test for presence of Pb2+: add one drop of Na2S solution to one drop of the
reaction mixture on filter paper. Black precipitate (PbS) indicates presence of Pb2+.
Pour the mixture into 500 ml of water and decant it. Add 100 ml of diluted HNO3
(1:3) to the precipitate, stir and decant it until pH is neutral. Filter off the lead dioxide,
wash it with boiling water and dry.
b) Pb(NO3)2 + 2NaOH → Pb(OH)2 + 2NaNO3
Pb(OH)2 + 2NaOH + Cl2 → PbO2 + 2NaCl + 2H2O
Material: Pb(NO3)2 or Pb(CH3COO)2, NaOH, Cl2
Procedure: Dissolve 10.0 g of Pb(NO3)2 in 80 ml of water acidified with a drop of
concentrated nitric acid and add solution of 2.4 g of sodium hydroxide dissolved in
50 ml of water slightly with constant stirring. Heat the precipitated lead(II) hydroxide
suspension to 70 - 80 oC and bubble chlorine through it at the same time (do not heat
it over 80 oC or PbO will be obtained!). Decant the precipitated brown-black product
with diluted nitric acid (1:3) and then with water. Dry it at 100 oC.
PbO2 – brown powder, insoluble in water, decomposes on heating to Pb3O4 or PbO and oxygen. Good oxidizer.
Preparation of Potash alum K2(SO4).Al(SO4)3•12H2O)
It is white crystalline solid, soluble in water, used for the purification of water, leather industry paper industry and as fire extinguisher.
Melting point is 92oC
Potash alum is commonly known as 'PHITKARI'
Potash alum is prepared by mixing equi-molecular masses of potassium sulphate and aluminum sulphate in water followed by evaporation
K2SO4 + Al2(SO4)3 + 24H2O-- K2SO4.Al2(SO4)3.24H2O
(vii) Preparation of copper ammine sulphate - perform this experiment in a fume hood
CuSO4 + 4 NH3(aq) + H2O → [Cu(NH3)4]SO4.H2O
Material: CuSO4.5H2O, NH3(aq), C2H5OH
Procedure: Place 5 g of finely powdered copper sulphate, CuSO4.5H2O, in a small
beaker, pour upon it 7.5 ml of concentrated ammonia and 3 ml of water. Shake it for
about 1 minute and then heat it gently until all the solid dissolves. Add about 10 ml of
ethanol to the solution, let it stand for about one hour and filter off the crystals. Wash
them with a mixture of 5 ml of concentrated ammonia and 5 ml of ethanol. Dry them
on air in the hood.
[Cu(NH3)4]SO4.H2O – dark blue crystals, soluble in water (18 g in 100 ml of water at
21.5 oC), stable on air.
(viii) Preparation of cuprous chloride – work in a fume hood
a) Cu2+ + Cu + 2 Cl- → 2 CuCl
Material: CuSO4.5H2O, Cu, NaCl, HCl, Na2SO3, CH3COOH
Procedure: Prepare a solution of 10 g of powdered CuSO4.5H2O and 15 ml of concentrated HCl in a 250-ml round-bottom flask, add 4 g of NaCl and heat to boiling.
Cover the flask with a little funnel. Add copper to hot solution in small portions. The green colour of the solution will turn to yellow. Filter off the remaining Cu. Pour the solution to one litre of cold water with 2 g of Na2SO3.
Wash the precipitated cuprous chloride 2 – 3 times with a solution of 1.5 g of Na2SO3, 6 ml of HCl and 300 ml of H2O by decantation. Filter out the precipitate and wash it with concentrated acetic acid. Dry it in a drying oven at 100 oC.
b) 2 Cu2+ + SO2 + 2 Cl- + 2 H2O → 2 CuCl + SO4
2- + 4 H+
Material: CuSO4.5H2O, NaCl, HCl, SO2, CH3COOH
Procedure: Add 5 g of NaCl to the warm solution (70 oC) of 10 g CuSO4.5H2O and bubble SO2 through the mixture. CuCl precipitates. Filter out the precipitate and wash
it with a solution of SO2 in water and then with concentrated acetic acid.
CuCl – white crystals, insoluble in water, on air turns to green alkali copper chloride
Publication Date: 06-13-2010
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