Chemistry - Nature Of Matter

NATURE OF MATTER Matter is anything that has mass and occupies space. Everything around us e.g. plants, animals, food, water, air including ourselves are matter. The mass of an object or a body is the quantity of matter it contains and it is the same everywhere. The S. I. unit for measuring mass is the mole but for practical purposes, it is measured in kilograms (kg) with a weighing balance. States of Matter The three main physical states of matter are solid, liquid and gas A. • • • • • • • • • B. • • • • • • • • Solid It has a definite shape It has a definite volume It has a fixed mass It is very dense Particles of solid are very closely packed and are held firmly together by forces of cohesion. Particles of solid possess less kinetic energy Particles of a solid only vibrate and rotate about fixed positions but cannot translate i.e. cannot move from one place to another. It is incompressible (it is difficult to compress. Examples of solids commonly used in nursing are tablets, capsules, caplets, injection powder, scissors, scalpels, forceps, sample bottles, needles, syringes, cotton wool, plaster, bandages, gloves, wheelchairs/walkers, beds, gauze pads/rolls, suppositories, implants etc. Liquid It has a definite volume. It has no definite shape. It assumes the shape of its container. It has a fixed mass. It is less dense than solid. Particles of liquid are slightly further apart than those of solid. Particles of liquid possess more kinetic energy than those of solid. Particles can vibrate, rotate and translate though they only move within restricted space because they are still under the influence of cohesive forces. It is incompressible i.e. it is difficult to compress. • Examples of liquids commonly used in nursing are water, methylated spirit, liquid hydrogen, syrups, injections, kerosene, normal saline, dextrose saline, dextrose water, oral rehydration solution (ORS), detail, liquid paraffin, eusol lotion, hydrogen peroxide etc. C. • • • • • • • • • Gas It has no definite shape It has no definite volume. It fills the entire space of its container. It has a fixed mass. It is the least dense among the three states of matter. Particles of gas are far apart and are loosely packed. Particles possess much more kinetic energy and move randomly at great speed Cohesive forces in gas are negligible. It is compressible. Examples oxygen, air, nitrogen, nitrous oxide, medical air, medical vacuum helium, heliox etc. Plasma • • • • • • It is a super heated state of matter It is a state of matter that exists when atoms are in excited state. This occurs when atoms are heated in such a way that electrons are ripped off. Plasma consists of charged particles with extremely high kinetic energy. Like gas, plasma has no definite volume and shape. Unlike gas, plasma possess electrical charges hence, it is electrically conductive Examples are stars(sun), lightning, aurorae, neon sign, fluorescent bulb, welding arcs, fireball of nuclear explosion, plasma TVs, Earth’s ionosphere, solar wind etc. Physical and Chemical Properties of Matter Properties are attributes by which an object or a thing can be identified. A matter possesses both physical and chemical properties. A. Physical Properties of Matter Physical properties of matter are properties that can be measured without matter changing to an entirely different substance. They are properties that can be detected with our senses and are associated with physical changes. Examples are melting point, boiling point, freezing point, density, hardness, texture, heat and electrical conductivity, solubility in other substances, colour, taste, odour, volatility, shape, volume, malleability etc. B. Chemical Properties of Matter Chemical properties of matter are properties that can be measured only when matter undergoes a change to become an entirely different kind of matter. They are properties associated with chemical changes or reactions. Examples are flammability, rusting/corrosion, toxicity, acidity, alkalinity, decomposition, reactivity etc. Assignment Outline the importance of the knowledge of properties of matter to nursing students. PHYSICAL AND CHEMICAL CHANGES • Physical Change Physical change is a change which is easily reversed and in which no new substances are formed. It is a temporary change which takes place without any changes in molecular composition. There is no change in the mass of substance undergoing the change. Physical change does not involves any great heat changes except latent heat changes which occur during changes of state. Examples of physical changes are: • • • • • • • Dissolution of common salt in water. Melting of candle wax. Freezing of liquids to solids. Vapourization of liquids to gases. Liquefaction of gases to liquids. Separation of mixtures by various separation techniques like evaporation, distillation, crystallization, filtration, sublimation etc. Chemical Change Chemical change is a change which is not easily reversed and in which new substances are formed. It is a permanent change which involves changes in molecular composition of substances. There is always a change in the mass of substance that undergoes the change. It involves a considerable amount of heat changes. Examples are: • • • Burning of wood Rusting of iron rods e.g. nails Dissolution of metals in acids • • • • Rotting of leaves Oxidation of alcohols to alkanals, alkanones and alkanoic acids. Fermentation of sugars Fading of colours etc. SEPARATION MIXTURES Mixtures are impure substances whose constituents are physically combined together. Hence, they can be separated by physical methods known as Separation Techniques. The basic principles behind the choice of separation methods for separating mixtures are the physical properties of the individual constituents that make up the mixtures. The table below shows different separation techniques/methods and the basic principles behind the choice of each of them. 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. Separation Techniques Sieving Magnetic separation Sublimation Decantation Filtration Centrifugation Evaporation Crystallization Fractional crystallization Precipitation Distillation Fractional distillation Separation funnel Chromatography What to separate (Basic Principle) Solids of different sizes Magnetic solids from nonmagnetic ones Solid which sublimes from other solids Insoluble solids from solution Soluble solid from solution Solvent from its solution Miscible liquids of different boiling points Immiscible liquids Solutes from a solution. 1. Sieving: It is used to separate solid particles of different sizes. It involves the use of a sieve with a mesh of a particular size. When the mixture is placed on the sieve, particles smaller than the mesh size pass through the sieve while the bigger particles remain on it. Sieving is used in mining industries involving diamond and gold as well as garri making industries in West Africa. 2. Magnetic Separation: This involves the use of a magnet to separate magnetic substances from nonmagnetic ones. It is used in steel industries and to remove magnetic impurities from tin ore. 3. Sublimation: This is used to separate a solid which sublimes (i.e. changes directly from solid to gaseous state without passing through the liquid state) when heated. Examples of substances that sublime on heating are iodine, ammonium chloride, benzoic acid, sulphur, iron(III)chloride, phosphorus (V) Chloe, camphor, naphthalene etc. 4. Decantation: This is used to separate a mixture of solid and liquid which separate into two distinct layers (upper liquid layer and lower solid layer) on standing. It involves the use of glass rod to carefully pour or decant the upper liquid layer into a second container. Decantation is a quick but inaccurate method of separation. 5. Filtration: It is used to separate insoluble solid particles from liquid. It involves the use of a porous material like filter paper which is selectively permeable. When the mixture is placed on filter paper in a funnel, the liquid drains through the filter paper while the solid particles remain on it. The liquid which passes through the filter paper is called filtrate while the solid left on the paper is called residue. Filtration is used in water purification plants and breweries to remove solid particles from liquids. 6. Centrifugation: Centrifugation is used in hospitals and medical laboratories to separate blood samples into blood cells and plasma. It involves the use of a centrifuge, a machine which can spin test-tubes containing suspensions at high speed. The spinning causes heavier solid particles to be thrown to the bottom of the test-tubes. The upper clear liquid layer can easily be decanted. 7. Evaporation: This is used to recover a solid solute from a solution while the solvent is usually sacrificed. The solid to be recovered by evaporation must withstand dry heating i.e. must not decompose on heating. Water-bath or sand-bath is used to bring about a steady rate of evaporation. Evaporation is used in salt-making industries where sea water is allowed to evaporate in trenches under the heat of sun. Note: Evaporation is not suitable for salts that easily decompose on heating. Solution —> solute + solvent. 8. Crystallization: It is used to separate solids which decompose easily on heating from their solutions. The salt solution is first heated to drive away some of the solvent until the solution becomes saturated. When the saturated solution is cooled, crystals of the salt begin to form. Crystal formation can be induced by: • • Adding few crystals of the same salt to serve as seeds; and Scratching the inside of the vessel containing the solution. The salt crystals obtained in this way are pure and contain water of crystallization e.g. i. ii. iii. Copper (II) tetraoxosulphate (VI)-pentahydrate, CuSO4. 5H2O. Copper (II) trioxonitrate (V)-trihydrate, Cu(NO3)2.. 3H2O. Iron (II) tetraoxosulphate (VI)-heptahydrate, FeSO4. 7H2O. Crystallization is widely used in drug and sugar production where purity of products is highly essential. 9. Fractional Crystallization: This is used to separate two or more solutes which are present in the same solution in roughly equal amounts. This method is based on the difference in the solubilities of the different solutes at different temperatures. During the cooling process, at a particular temperature, crystals of the relevant solute will come out of the solution leaving behind others which are still within their limits of solubility. 10. Precipitation: This method of separation is based on the difference in the solubility of a solid in two miscible liquids. For example, iron (II) tetraoxosulphate (VI), FeSO4, is soluble in water but insoluble in ethanol. If ethanol is added to the aqueous solution of FeSO4, the salt precipitates out and can be filtered from the solution. 11. Distillation: Distillation is used to recover a solvent from its solution. The solution is heated in a flask to vapourize the solvent. The vapour passes along a condenser which is cooled by the circulating water in its outer jacket. This recondenses the vapour into a liquid called distillate, which is then collected in a receiver. The solute and other impurities are left behind in the distillation flask. Distillation is used in water and gin distilleries. 12. Fractional Distillation: This method is based on the difference in boiling points of miscible liquids. It is used to separate a mixture of two or more miscible liquids into its components or fractions. The fractions distil over in ascending order of boiling points starting with the fraction with the lowest boiling point. In laboratory, the apparatus is the same as simple distillation except that a fractionating column is introduced between the distillation flask and the condenser. Fractional distillation is used to separate petroleum into its components. 13. Separating Funnel: This is used to separate a mixture of two immiscible liquids like petrol and water. The mixture is poured into the separating funnel and allowed to stand. On standing it separate into two distinct layer (lower denser layer and upper less dense layer). The lower denser layer can easily be tapped off leaving the upper less dense liquid layer in the funnel. 14. Chromatography: It involves the use of a solvent moving over a porous adsorbent medium to separate a mixture of solutes. It is used to separate coloured as well as colourless substances. The four main types of chromatography are: i. ii. iii. iv. Paper chromatography Gas chromatography Column chromatography Thin layer chromatography The commonest method is the ascending paper chromatography which makes use of a strip of chromatographic or filter paper. The solution containing the mixture of solutes is spotted onto the strip near one end. The paper is then suspended in a closed air-tight jar with the spotted end dipping into an appropriate solvent. As the system moves up the strip, the different solutes also move along the strip together with the solvent at different speeds and become separated. The paper is then removed and dried and if necessary, sprayed with the appropriate chemical reagents to locate the positions of the various solutes on the paper. There are two phases in chromatography namely: a) Stationary phase (e.g. the paper) on which the solutes are adsorbed; and b) Moving phase(e.g. the solvent) which dissolves the solutes and carries them. Assignment 1. Explain how you can obtain pure nitrogen and oxygen from air 2. State the importance of the following separation techniques in nursing i. ii. iii. iv. v. Centrifugation Filtration Distillation Crystallization Chromatography ELEMENTS, COMPOUNDS AND MIXTURES Element An element is a substance which cannot be split into simpler units by an ordinary chemical process. Elements are pure substances. They are the building units of other substances. Elements are grouped into three main classes namely: metals, semi-metals/metalloids and nonmetals. 1. Metals:- These are elements whose atoms ionize by electron loss to form positive ions. Examples are iron, sodium, aluminum, magnesium, potassium, gold, etc. Physical Properties of Metals • • • • • • • • • They are good conductors of heat and electricity They are solids at room temperature except mercury which is liquid. They have high melting and boiling points They have high relative density They are malleable i.e. they can be hammered into sheets They are ductile i.e. they can be drawn into thin wires. They are sonorous I.e. they give off notes when hit. They have great tensile strength They are lustrous i.e. they have shining appearances Chemical Properties of Metals • • Their atoms ionize by electron loss to form positive ions e.g. Na —> Na+ + 2eMg —> Mg2+ + 2eAl —> Al3+ + 3e-0 They are good reducing agents They form basic or amphoteric oxides with oxygen. • • • They have low electronegativities. They tend to form ionic compounds More electropositive metals displace hydrogen from acids. 2. Nonmetals:- These are elements whose atoms ionize by electron gain to form negative ions e.g. oxygen, chlorine, carbon, nitrogen, neon, xenon, sulphur, phosphorus, bromine, iodine etc. Physical Properties of Nonmetals • • • • • • They are poor or non conductors heat and electricity except carbon (graphite) They may be solids, liquids or gases at room temperature They have low melting and boiling points except carbon. They have low relative density They are not malleable, ductile and sonorous. They are non-lustrous. Chemical Properties of Nonmetals • • • • • Their atoms ionize by electron gain to form negative ions e.g. Cl + e- —> ClO + 2e- —> O2N + 3e- —> N3They have high electronegativities. They form acidic and neutral oxides with oxygen. They tend to form covalent compounds. Nonmetals cannot displace hydrogen from acids. 3. Semi-metals/metalloids:- These are elements which possess both metallic and nonmetallic properties. They are semi-conductors of heat and electricity. E.g. boron, silicon, germanium etc. Chemical Symbols of Elements Symbol of an element represents the atom of that element. The modern system of representing elements with symbols was suggested by Berzilius in 1814. The symbols of elements could be written in the following ways. 1. Elements with the first letter of their names as their symbols Atomic Number 1 5 6 7 8 9 15 16 23 53 92 Name of Element Hydrogen Boron Carbon Nitrogen Oxygen Fluorine Phosphorus Sulphur Vanadium Iodine Uranium Symbol H B C N O F P S V I U 2. Elements with the first and one other letter of their names as their symbols Atomic Number 2 3 4 Name of Element Helium Lithium Beryllium Symbol Li Be 10 12 13 14 17 18 20 21 22 24 25 30 32 33 34 35 36 37 38 48 55 Neon Magnesium Aluminum Silicon Chlorine Argon Calcium Scandium Titanium Chromium Manganese Zinc Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Cadmium Caesium Me Mg Al Si Cl At Ca Sc Ti Cr Mn Zn Ge As Se Br Kr Rb Sr Cd Cs 3. Elements that derived symbols from their Latin names Atomic Number 11 19 26 29 47 50 51 79 80 82 Name of Element Sodium Potassium Iron Copper Silver Tin Antimony Gold Mercury Lead Latin name Symbol Natrium Kalium Ferrum Cuprum Argentum Stannum Stibium Aurum Hydragyrum Plumbum Na K Fe Cu Ag Sn Sb Au Hg Pb Compounds ✓ A compound is a substance which contains two or more elements chemically combined together. Compounds are pure substances. They are always homogeneous. Components of compounds are present in fixed proportion by mass Compounds can be represented by chemical formulae. The properties of a compound is entirely different from those of its component elements. ✓ The two main types of compounds are organic and inorganic compounds. Organic compounds are compounds that contain carbon-hydrogen covalent bonds (C—H). They contain large number of atoms per molecule and are mainly found as natural products in plants and animals. Inorganic compounds are compounds that do not contain carbon-hydrogen covalent bonds. They contain few atoms per molecule. ✓ Examples of compounds, their component elements and chemical formulae are shown in the table below ✓ ✓ ✓ ✓ ✓ Compound Water Carbon (IV) oxide Sand Common salt Caustic soda Washing soda Ethanol Limestone Sugars (sucrose) Component elements Hydrogen, oxygen Carbon, oxygen Silicon, oxygen Sodium, chlorine Sodium, oxygen, hydrogen Sodium, carbon, oxygen, hydrogen Carbon, hydrogen, oxygen Calcium, carbon, oxygen Carbon, hydrogen, oxygen Chemical formula H 2O CO2 SiO2 NaCl NaOH Na2CO3.10H2O C2H5OH CaCO3 C12H22O11 3. Mixtures:✓ A mixture is a substance which contains two or more constituents which can be separated by physical methods. ✓ They are impure substances ✓ The constituents can be elements or compounds or both. ✓ The individual constituents can be added in any proportion by mass. ✓ They cannot be represented by chemical formulae. ✓ The properties of a mixture is the sum of those of the individual constituents that make up the mixture. ✓ Mixture can be homogeneous or heterogeneous. ✓ Examples of mixtures and their constituents are shown in the table below Mixture Air Soil Urine Coca-cola Blood Sea water Milk Crude oil Ampiclox Constituents Nitrogen, oxygen, carbon(IV) oxide, water vapour, dust Sand, clay, humus, air, water, mineral salts Water, urea, mineral salts Water, carbon(IV) oxide, sugar, coca-cola concentrate Water, proteins, sugars, vitamins, mineral salts, haemoglobin, blood cells, enzymes Water, mineral salts, bacteria, remains of organic matter Water, proteins, sugar, fats, mineral salts, vitamins Petrol, heavy oil, gas oil, kerosene, naphtha, gas bitumen Ampicillin, cloxacilin PARTICULATE NATURE OF MATTER The three main discrete particles of matter are atoms, molecules and ions. Atom An atom is the smallest particle of an element which takes part in chemical reaction. Or an atom is the smallest part of an element that can ever exist and still possess the chemical Properties of that element. Atom is the basic unit of simple substances or elements. Molecule A molecule is the smallest particle of a substance that can normally exist alone and still retain the chemical Properties of that substance be it an element or a compound. Molecules may be made up of atoms of the same element or different elements. The number of atoms in each molecule of an element is called the atomicity of that element. Molecules may be: i. ii. iii. Monatomic i.e. having one atom per molecule e.g. helium (He), sodium (Na), magnesium (Mg), potassium (K), neon(Ne) etc. Diatomic i.e. having two atoms per molecule e.g. oxygen (O2), chlorine (Cl2), hydrogen (H2), nitrogen (N2), bromine (Br2) etc. Polyatomic i.e. having more than two atoms per molecule e.g. ozone (O3), phosphorus (P4), sulphur (S8) Ion An ion is any atom or group of atoms which possess electrical charge. A group of atoms carrying electrical charge which enters into chemical reactions as a single unit is called radical. Ions may be positively charged or negatively charged. The positively charged ions are called cations e.g. Na+, Ca2+, K+, Fe2+, Mg2+, Al3+, NH4+, Zn2+, Mn2+ etc. while the negatively charged ions are called anions e.g. Cl-, O2-, Br-, N3-, OH-, CO32-, SO42-, SO32-, HSO4-, PO43-, ClO- etc. Constituents/Sub particles of the Atom The main sub particles of the atom are protons, neutrons and electrons Proton The proton is located inside the nucleus of the atom. It is positively charged It has a mass equal to 1 Neutron It located inside the nucleus of the atom It has no charge i.e. it is neutral It has a mass equal to that of the proton Electron It is located outside the nucleus of the atom. It is negatively charged. It has a negligible mass approximately 0.0005. Note: The atomic particles residing inside the nucleus are collectively known as nucleons e.g. protons and neutrons. Discoveries about the Atom 1. Enest Rutheford (1911). He discovered the planetary structure of the atom. According to him, the protons and neutrons are concentrated in the centre (nucleus) while the electrons revolve round the nucleus. He also proposed the nuclear theory of the atom which states that “the atom consists of a positive core called the nucleus, where most of the mass of the atom is contained, and electrons which move round the nucleus”. Rutherford was later awarded the Nobel Prize for his work on the atom. 2. Niels Bohr (1913). He suggested that the electrons in an atom revolve round the nucleus along certain imaginary circular lines or paths known as orbits or shells situated at various distances from the nucleus. He named the shells K, L, M, N, O, P, Q. Bohr use the Quantum Theory ideas developed by Plank and Einstein to account for the spectrum of hydrogen. He made the following assumptions: i. ii. iii. Rutherford’s model of the atom is basically correct. Each spectral line is caused by an electron. Electrons can exist only in circular orbits of definite quantum energy. iv. v. An electron emits energy in the form of radiation when it moves from a higher to a lower permitted orbit. This produces a line in atomic emission spectrum. Since the energies of the higher and lower orbits are fixed, the line will be of a particular energy and frequency. The difference in energy, ΔE, of the two orbits is related to the frequency of radiation given by Plank’s equation, ΔE = En2 – En1 where En1 is a state with lower energy level compared to En2. The investigation of the hydrogen atom led Bohr to postulate that the circular orbits of electrons were quantize i.e. the electrons were orbiting at certain specific energy levels within the atom. 3. J. J Thompson (1897). He discovered electrons and protons using cathode ray tube and discharge tube respectively. He successfully measured the charge to mass ratio 𝑐ℎ𝑎𝑟𝑔𝑒 ( 𝑚𝑎𝑠𝑠 ) of the electron to be 1.76 x 1011Ckg-1 which is called the specific charge of the electron. He also measured the charge to mass ratio of the proton to be 9.58x107Ckg-1 4. R. A. Millikan (1910). Millikan in his oil-drop experiment successfully measured the charge of the electron to be -1.6 x 10-19C. Using the above values , the mass of an electron is found to be 9.1 x10-31kg which is very small to make any difference to the mass of the atom. 5. Roentgen (1885). He discovered a penetrating radiation known as X-ray when high energy electrons in a discharge tube collide with the anode. 6. James Chadwick (1932): He discovered neutron when he bombarded a thin sheet of beryllium with alpha particles. He proved that the neutron has no charge (i.e. neutral) and has a mass equal to that of the proton. 7. Moseley(1914): He found that atomic number could be assigned to an element based on the number of X-rays the element emits when it is used as the anode of an X-ray tube. Further experiments enabled Moseley to conclude that the atomic number represents the number of protons in the nucleus of the atom of the element. 8. Geiger and Marsden: Geiger and Marsden in their gold-foil experiment developed an idea that an atom has a minute centre carrying a positive charge. Calculations based on their work indicated that the number of positive charge (protons) in the nucleus of an atom are approximately half its relative atomic mass. DALTON’S ATOMIC THEORY John Dalton in 1808 proposed the atomic theory which can be summarized as follows: 1. All elements are made up of small indivisible particles called atoms. 2. Atoms can neither be created nor destroyed. 3. Atoms of the same element are alike in every aspect and differ from atoms of other elements. 4. When atoms combine with other atoms, they do so in simple ratios. 5. All chemical changes result from the combination or separation of atoms. Modifications to Dalton’s Atomic Theory 1. All Elements are made up of small indivisible particles Called Atoms: This statement has been proven wrong by Rutherford’s discovery that atom is built up of three sub particles namely: protons, electrons and neutrons. Hence, atom is not indivisible solid piece. 2. Atoms can neither be Created nor Destroyed: This statement still hold good for ordinary chemical reactions but nuclear reactions (nuclear fission and nuclear fusion) show that atoms can be destroyed or created e.g. the nucleus of uranium235 can be broken into smaller units while a tremendous amount of energy is released. 3. Atoms of the same Element are alike in every Aspect and Differ from Atoms of other Elements: The discovery of isotopes makes this statement unacceptable e.g. chlorine has two isotopes, 3517Cl and 3717Cl which differ in their neutron number and relative atomic masses. 4. During Chemical Reaction, there is a Combination of Atoms in Small whole Numbers: This statement is generally true only for inorganic compounds which usually contain few atoms per molecule. Carbon, however, forms very large organic molecules such as starch, proteins, fats etc. which contain thousands of atoms. Modern Atomic Theory Modern Atomic Theory, grounded in quantum mechanics, describes the atom as a central, dense, positively charged nucleus surrounded by a "cloud" of negatively charged electrons existing in specific energy levels or orbitals. This model, a culmination of work by scientists like John Dalton, Ernest Rutherford, Niels Bohr, Erwin Schrödinger, and Werner Heisenberg, is the foundation of modern chemistry and physics. Key Principles of Modern Atomic Theory • All matter is composed of atoms: Atoms are the fundamental building blocks of all matter. • Atoms have a complex internal structure: Contrary to Dalton's original theory that atoms are indivisible, they are composed of subatomic particles: protons, neutrons, and electrons. • The Nucleus: Protons (positive charge) and neutrons (no charge) reside in a tiny, dense, central nucleus that contains almost all the atom's mass. • Electrons and Quantum Mechanics: Negatively charged electrons exist in regions of space around the nucleus called orbitals or electron clouds. Their exact position and momentum cannot be known simultaneously (Heisenberg Uncertainty Principle); rather, their location is described by probability functions (Schrödinger equation). • Discrete Energy Levels: Electrons occupy specific, discrete energy levels. They can move between these levels by absorbing or emitting energy in fixed packets (quanta), but they cannot exist in between these levels. • Atomic Number Determines Element Identity: The number of protons (atomic number) in the nucleus determines the identity of an element. All atoms of the same element have the same number of protons. • Isotopes: Atoms of the same element can have different numbers of neutrons. These variations are called isotopes and have different atomic masses. • Ions: Atoms can gain or lose electrons to form ions, which are electrically charged particles. • Atoms combine to form compounds: Atoms of different elements combine in small, whole-number ratios to form chemical compounds, which can be rearranged during chemical reactions. • Protons and Neutrons are further divisible: Protons and neutrons are themselves composed of even smaller particles called quarks. The four Quantum Numbers 1. The Principal Quantum Number (n): This determines the energy levels or shells of an atom. For K, L, M, N … shells, n has integral values of 1, 2, 3, 4 … respectively. i.e. for K shell, n = 1, for L shell, n = 2 and so on. Note: The maximum number of electrons in a shell is given by 2n2 and the total number of orbitals is given by n2. 2. The Subsidiary or Azimuthal Quantum Number (l): This shows the number of energy sublevels in each electron shell. The Subsidiary quantum number has integral values ranging from 0 (n-1). The electrons with Subsidiary quantum numbers 0, 1, 2 and 3 are usually referred to as the s, p, d, and f electrons respectively. 3. The Magnetic Quantum Number (m): This shows the number of orbitals in each energy sublevel. It has integral values ranging from -1 through 0 to 1. E.g. p subshell has three p orbitals. 4. The Spin Quantum Number (s): This describes the angular momentum of an electron. The spin quantum number has the values of -½ and +½. An orbital can only hold two electrons which spin on their axes in opposite directions. The first electron in an orbital will have a spin of +½ while the second will have a spin of -½. The Energy Sublevels Each energy level denoted by the principal quantum number, n, is divided into sublevels. The sublevels are assigned letters s, p, d and f. s stands for sharp, p for principle, p, d for diffuse and f for fine or fundamental. These are given in the table below: Name of Electron shell Value of n Value of l (n- Number of 1) value of l Number of sublevels Name of sublevel K 1 0 One One s L 2 0 and 1 Two Two s and p M 3 0, 1 and 2 Three Three s, p and d N 4 0, 1, 2 and 3 Four Four s, p, d and f Number of Orbitals in each Sublevel Value of l (n- Name of 1) sublevel Value of m Number of value of m Number of orbitals Name of orbital 0 s 0 One One s 1 p 1, 0, -1 Three Three p 2 d 2, 1, 0, -1, -2 Five Five d 3 f 3, 2, 1, 0, -1, -2, -3 Seven Seven f Principles or Rules Governing Electron Filling into Orbitals 1. Pauli’s Exclusion Principle: This principle states that “two electrons in the same orbital of an atom cannot have same values for all the four quantum numbers”. In other words, two electrons in an atom cannot behave in the same manner. 2. Hund’s Rule of Maximum Multiplicity: This rule states that “electrons occupy each orbital singly first before pairing takes place in a degenerate orbital”. 3. Aufbau’s Principle: This principle states that “in the building up of an atom, electrons enter into Orbitals in order of increasing energy”. This means that the electrons are fed into Orbitals starting at the lowest energy level before filling higher energy levels. Aufbau’s principle can be derived using the illustration below: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p i.e. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p etc. Arrangement of Electrons around the Nucleus The arrangement of electrons around the nucleus of an atom is known as electronic configuration. Niels Bohr (1913) suggested that the electrons in an atom revolve round the nucleus along certain imaginary circular lines or paths known as orbits or shells situated at various distances from the nucleus. He named the shells or orbits K, L, M, N, O, P, and Q with K being the nearest shell to the nucleus followed by L, M and so on. The maximum number of electrons in a shell numbered “n” is given by 2n2 where n is the shell number. The maximum number of electrons in the outermost shell of any atom is 8. Shell/orbit Value of n (shell number) Maximum number of electrons (2n²) K 1 2x1² = 2x1 = 2 L 2 2x2² = 2x4 = 8 M 3 2x3² = 2x9 = 18 N 4 2x4² = 2x16 = 32 O… 5 2x5² = 2x25 = 50… Electronic structure of oxygen, 2, 6. Electronic structure of sodium, 2, 8, 1. Element Symbol Atomic Electronic configuration number Hydrogen H 1 1s¹ Helium He 2 1s² Lithium Li 3 1s² 2s¹ Beryllium Be 4 1s² 2s² Boron B 5 1s² 2s² 2p¹ Carbon C 6 1s² 2s² 2p² Nitrogen N 7 1s² 2s² 2p³ Oxygen O 8 1s² 2s² 2p⁴ Fluorine F 9 1s² 2s² 2p⁵ Neon Ne 10 1s² 2s² 2p⁶ Sodium Na 11 1s² 2s² 2p⁶ 3s¹ Magnesium Mg 12 1s² 2s² 2p⁶ 3s² Aluminum Al 13 1s² 2s² 2p⁶ 3s² 3p¹ Silicon Si 14 1s² 2s² 2p⁶ 3s² 3p² Phosphorus P 15 1s² 2s² 2p⁶ 3s² 3p³ Sulphur S 16 1s² 2s² 2p⁶ 3s² 3p⁴ Chlorine Cl 17 1s² 2s² 2p⁶ 3s² 3p⁵ Argon Ar 18 1s² 2s² 2p⁶ 3s² 3p⁶ Potassium K 19 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ Calcium Ca 20 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² Scandium Sc 21 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹ Titanium Ti 22 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d² Vanadium V 23 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d³ Chromium Cr 24 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵ Manganese Mn 25 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁵ Iron Fe 26 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ Cobalt Co 27 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁷ Nickel Ni 28 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁸ Copper Cu 29 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰ Zinc Zn 30 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ Atomic Number and Mass Number The atomic number, Z, of an element is the number of protons in one atom of the element. The mass number, A, of an element is the sum of the protons and neutrons in one atom of the element. The atom of an element can be represented by writing it’s atomic number as subscript to the left and mass number as superscript to the left of its symbol. E.g.⁴⁰20Ca, ¹⁶8O, ³⁹19K, ²⁷13Al etc. Summary about the Atom • • • • • • The number of electrons in a neutral atom equals the number of protons i.e. e- = p. The atomic number, Z, equals the number of protons i.e. Z = p. The mass number, A, equals the sum of protons and neutrons i.e. A = p + n. For an ion, the number of protons is unaffected while the number of electrons either increases or decreases depending on the charge on the ion For cations (positively charged ions) the number of electrons equals the number of protons minus number of charges i.e. e- = p - charge. For anions (negatively charged ions) the number of electrons equals the number of protons plus the number of charges i.e. e- = p + charge. Examples 1. Write the representation of the atom of element X containing 15 protons, 15 electrons and 16 neutrons. Solution: ³¹15X 2. If the atom of element Z has 19 protons, 19 electrons and mass number of 39, find the number of neutrons in the nucleus of its atom. Solution: A=p+n n=A–p n = 39 – 19 n = 20 3. Find the number of protons, electrons and neutrons in each of the following (i) ¹⁶8O2- (ii) ²⁷13jAl3+ (iii) ⁴⁰20Ca2+ (iv) ³⁵17ClParticle Proton Electron Neutron (i) ¹⁶8O2- 8 8 + 2 = 10 16 – 8 = 8 (ii) ²⁷13jAl3+ 13 13 – 3 = 10 27 – 13 = 14 (iii) ⁴⁰20Ca2+ 20 20 – 2 = 18 40 – 20 = 20 (iv) ³⁵17Cl- 17 17 + 1 = 18 35 – 17 = 18 ISOTOPY Isotopy is a phenomenon whereby atoms of an element have different mass numbers but the same atomic number. The different atoms of an element having the same atomic number but different mass numbers are called isotopes. The difference in mass is due to the difference in the number of neutrons present in the nucleus of the atoms. Examples of elements exhibiting isotopy include • • • • Chlorine ³⁵17Cl and ³⁷17Cl Oxygen: ¹⁶8O, ¹⁷8O and ¹⁸8O. Hydrogen: ¹1H, ²1H and ³1H Carbon: ¹²6C, ¹³6C and ¹⁴6C etc. The relative atomic mass of an element exhibiting isotopy is the average mass of its various isotopes as they occur naturally in any quantity of the element e.g. The relative atomic mass of chlorine is 35.5 because a given quantity of chlorine always contains 75% of ³⁵17Cl and 25% of ³⁷17Cl i.e. 75/100 X 35) + (25/100 X 37) 26.25 + 9.25 35.5 Further example An element X has two isotopes ²⁰10X and ²²10X present in ratio1 : 3, calculate the relative atomic mass of X. Solution Sum of ratios = 1 + 3 = 4 :. The relative atomic mass of X = (¼ X 20) + (¾ X 22) 20/4 + 66/4) = 5 + 16.5 = 21.5 Note: The relative atomic mass of an element exhibiting isotopy is not a whole number because it’s is the weighted average of the masses of the various isotopic forms. Relative Atomic Mass The relative atomic mass, Ar, of an element is the number of times the average mass of one atom of that element is heavier than one-twelfth the mass of one atom of carbon-12. Relative atomic mass, Ar, is given by Ar = Average mas of 1 atom of the element 1 mass of 1 atomof carbon−12 12 Carbon-12 Scale Carbon-12 Scale is a standard reference for comparing the relative atomic masses of other elements. Or it is a mass reference scale in which the atom of the ¹²C isotope is taken to have twelve units of atomic mass. Relative Molecular Mass, Mr. The relative molecular mass, Mr, of an element or a compound is the number of times the average mass of one molecule of it is heavier than one-twelfth the mass of one atom of carbon-12. The relative molecular mass of an element or a compound is the sum of the relative atomic masses of all atoms in one molecule of that element or compound. Example Calculate the relative molecular mass of each of the following (i) limestone, CaCO3 (ii) magnesium tetraoxosulphate (VI), MgSO4 (iii) lead (II) trioxonitrate (V), Pb(NO3)2 (iv) oxygen molecule, O2 (v) potassium tetraoxomanganate (VII), KMnO4. (Ca = 40, C = 12, O = 16, Mg = 24, S = 33, Pb = 207, N = 14, K = 39, Mn = 55) Solution (I) CaCO3 Mr. = (1x40)+(1x12)+(3x16) = 40+12+48 = 100 (ii) MgSO4 (1x24)+(1x32)+(4x16) 24+32+64 = 120 (iii) Pb(NO3)2 (1x207)+(2x14)+(6x16) = 207+28+96 = 331 (iv) O2 2x16 = 32 (v) KMnO4 (1x39)+(1x55)+(4x16) = 39+55+64 = 158 Mole One mole of a substance is the amount containing as many elementary entities as the number of atoms in exactly 12 grams of carbon-12. Through experimental works, q mole of carbon-12 contains 6.02 X 10²³ atoms. This is called Avogadro number. Avogadro number= 6.02 X 10²³ The definition of a mole also gives the following relationship Number of mole = Also, Number of mole = Number of particles 6.02 X10²³ reacting mass molar mass Molar Mass The molar mass of any substance is the mass of one mole of that substance expressed in grams. It’s unit is gram per mole (g/mol or gmol-1). Since a mole of any substance can be represented by its formula, the molar mass of a substance is its formula mass expressed in grams. Calculations Involving Moles, Masses and Formulae 1. what is the mass of 3 moles of oxygen, O2? (O = 16) Solution Molar mass of O2 = 2 x 16 = 32 Mass of 3 moles of O2 = (3x32)g = 96g 2. How many atoms are there in 6g of C? (C = q2, 1 mole = 6.02 x 10²³) Solution 12g of C = 1 mole 6g of C = 6/12 =½ mole Using 𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒 = 𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒𝑠(𝑎𝑡𝑜𝑚𝑠) 6.02 𝑋1023 No of atoms = no of moles x 6.02x10²³ = ½ x 6.02x10²³ = 3.01x10²³ atoms. 3. How many moles are there in 20g of limestone CaCO3 (CaCO3 = 100) Solution Using no of mole = reacting mass molar mass = 20/100 = 0.2mole 4 Calculate the percentage by mass of nitrogen in sodium trioxonitrate (V), NaNO3 (Na = 23, N = 14, O = 16). Solution Molar mass of NaNO3 = 23+14+(3x16) = 85 % by mass of N = 14/85 x 100/1 = 1400/85 = 16.47% Exercise Calculate the percentage by mass of all the component elements in magnesium trioxonitrate (V), Mg(NO3 (Mg = 24, N = 14, O = 16). VALENCIES OF ELEMENTS AND RADICALS The valency of an element is a measure of its combining capacity with other atoms, which is determined by the number of electrons it can lose, gain, or share to achieve a stable electron configuration, typically a full outer shell (octet rule). How Valency is Determined Valency is directly related to the number of electrons in the outermost shell, known as valence electrons. The general rules for main-group elements are: • For elements with 1 to 4 valence electrons: The valency is usually equal to the number of valence electrons (e.g., Group 1 elements have a valency of 1). These elements tend to lose electrons to form positive ions. • For elements with more than 4 valence electrons: The valency is typically calculated as 8 minus the number of valence electrons (e.g., Oxygen in Group 16 has 6 valence electrons, so its valency is 8 - 6 = 2). These elements tend to gain or share electrons to complete their octet. • Noble Gases (Group 18): These elements have a full outer shell (8 valence electrons, or 2 for Helium) and are generally unreactive, so their valency is typically zero. Periodic Trends in Valency • Across a Period (left to right): The valency for main group elements first increases from 1 to 4 and then decreases to 0. • Down a Group (top to bottom): Elements within the same group generally have the same number of valence electrons and thus exhibit similar valencies. Key Points • Variable Valency: Many elements, especially transition metals (d-block) and fblock elements, can exhibit more than one valency depending on the compound they form. For example, iron (Fe) can have a valency of 2 or 3. • No Sign: Valency is a number without a positive or negative sign, unlike the oxidation number, which indicates the hypothetical charge an atom would have. Valencies of common elements are shown in the table below: Element Symbol Valency Element Symbol Valency Hydrogen H +1 Scandium Sc +3 Helium He 0 or nil Titanium Ti +2, +3, +4 Lithium Li +1 Vanadium V +2, +3, +4, +5 Beryllium Be +2 Chromium Cr +2, +3, +6 Boron B +3 Manganese Mn +2, +3, +4, +6, +7 Carbon C +2, +4 Iron Fe +2, +3 Nitrogen N -3, +5 Cobalt Co +2, +3 Oxygen O -2 Nickel Ni +2, +3 Fluorine F -1 Copper Cu +1, +2 Neon Ne 0 or nil Zinc Zn +2 Sodium Na +1 Barium Ba +2 Magnesium Mg +2 Silver Ag +1 Aluminum Al +3 Bromine Br -1 Silicon Si +2, +4 Iodine I -1 Phosphorus P -3, +5 Tin Sn +2, +4 Sulphur S -2, +4, +6 Lead Pb +2, +4 Chlorine Cl -1 Gold Au +1 Argon Ar 0 or nil Mercury Hg +1, +2 Potassium K +1 Calcium Ca +2 Note: • • The valency of hydrogen is always +1 except in hydrides where it is -1 The valency of oxygen is always -2 except in peroxides where it is -1 Valencies of common radicals are shown in the table below: Radical Formula Valency Ammonium NH4+ +1 Hydroxide OH- -1 Trioxonitrate (V) NO3- -1 Dioxonitrate (III) NO2- -1 Trioxocarbonate (IV) CO32- -2 Hydrogen trioxocarbonate (IV) HCO3- -1 Trioxosulphate (IV) SO32- -2 Hydrogen trioxosulphate (IV) HSO3- -1 Tetraoxosulphate (VI) SO42- -2 Hydrogen tetraoxosulphate (VI) HSO4- -1 Trioxophosphate (III) PO33- -3 Hydrogen trioxophosphate (III) HPO32- -2 Dihydrogen trioxophosphate (III) H2PO3- -1 Tetraoxophosphate (V) PO43- -3 Hydrogen tetraoxophosphate (V) HPO42- -2 Dihydrogen tetraoxophosphate (V) H2PO4- -1 Tetraoxomanganate (VII) MnO4- -1 Heptaoxodichromate (VI) Cr2O72- -2 Trioxochlorate (V) ClO3- -1 Oxochlorate (I) ClO- -1 Oxobromate (I) BrO- -1 Trioxobromate (V) BrO3- -1 Trioxoiodate (V) IO3- -1 Chemical Formulae Chemical formula it the symbolic representation of one mole of a compound. Chemical formula of compounds contain the symbols of the component elements written side by side. Chemical formulae are reduced from the valencies of component elements or radicals. Chemical formulae are written by exchanging the valencies of the component elements or radicals i.e. Ax + By —> AyBx Rules for Writing Chemical Formulae • • • Write out the symbols or formulae of the component elements or radicals Write out the valencies of the component elements or radicals Write the symbols close together, exchanging the valencies and writing the valencies as subscripts after the atom concerned ignoring the signs. • If a radical occurs more than once, it must be enclosed within parentheses (brackets) and the number written as subscripts after the parentheses. Examples Compound Component elements/radical Symbol/ Sodium hydroxide Sodium, hydroxide Na, OH- +1, -1 NaOH Water Hydrogen, oxygen H, O +1, -2 H2O Ammonia Nitrogen, hydrogen N, H -3, +1 NH3 +2, -1 Ca(NO3)2 Calcium (V) trioxonitrate Calcium, (V) Valencies Formula formula trioxonitrate Ca, NO3- Aluminum oxide Aluminum, oxygen Al, O +3, -2 Al2O3 Iron (III) chloride Iron, chlorine Fe, Cl +3, -1 FeCl3 Ammonium trioxocarbonate (IV) Ammonium, trioxocarbonate (IV) NH4+ SO42- , +1, +2 Iron (III) Iron, tetraoxosulphate Fe, SO42tetraoxosulphate (VI) (VI) +3, -2 (NH4)2SO4 Fe2(SO4)3 IUPAC NOMENCLATURE OF CHEMICAL COMPOUNDS IUPAC Nomenclature is the internationally accepted system of chemical compounds. This system of naming compound is determined by the constituent elements in the compound and oxidation number of necessary. Naming Binary Compounds Binary compounds are compounds containing only two elements. In naming such compounds, the electropositive element (cation) is named first while the electronegative element (anion) come last ending with -ide. Examples • • • • • MgO — magnesium oxide H2S — hydrogen sulphide CaC2 — calcium carbide NaH — sodium hydride KI — potassium iodide Note: when the two elements involved are Nonmetals, the rule does not apply e.g. H2O — water, NH3 — ammonia, PH3 — phosphine etc. Naming Radicals and Compounds Having Elements with Variable Oxidation Number In naming acid radicals having elements with variable Oxidation number, the Oxidation Number is indicated by Roman number in brackets e.g. • • • • NO3- — trioxonitrate (V) CaSO4 — calcium tetraoxosulphate (VI) KMnO4 — potassium tetraoxomanganate (VII) Fe2(SO4)3 — iron (III) tetraoxosulphate (VI) Naming Tertiary and Quaternary Compounds These are compounds containing more than two elements. Examples are oxo-acids, normal salts etc. • • • • • H2SO4 — Tetraoxosulphate (VI) acid HNO3 — Trioxonitrate (V) acid H3PO4 — Tetraoxophosphate (V) acid Na2SO4 — sodium tetraoxosulphate (VI) CuCO3 — Copper (II) trioxocarbonate (IV) Naming Compounds whose Cations has been Partially Replaced with Hydrogen ions Here, the cations are named first followed by hydrogen, then the radicals present indicating the oxidation number of the element that is combined with oxygen by Roman number in brackets e.g. • • • NaHSO4 — sodium hydrogen tetraoxosulphate (VI) KH2PO4 — potassium dihydrogen tetraoxophosphate (V) Ca(HCO3)2 — calcium hydrogen trioxocarbonate (IV) EMPIRICAL FORMULA AND MOLECULAR FORMULA Empirical formula of a compound is its simplest formula which gives the whole number ratio of the component elements. Molecular formula of a compound is one that gives the exact number of moles of atoms of the component elements in one mole of the compound. Molecular formula is calculated from the empirical formula and the relative molecular mass of the compound. Examples of empirical formula and molecular formula of some Compounds are shown in the table below: Compound Empirical formula Molecular formula Water H2O (H2O)n = H2O, n = 1 Hydrogen peroxide HO (HO)n. = H2O2 , n = 2 Propene CH2 (CH3)n = C3H6 , n = 3 Ethyne CH (CH)n = C2H2 , n = 2 Ethane CH3 (CH3)n = C2H6 , n = 2 Benzene CH (CH)n = C6H6 , n = 6 (CH3)n (CH)n (H2O)n. (H2O2)n C6H6 EXAMPLES 1. A certain compound on analysis yielded 2.00 g carbon. 0.34 g hydrogen and 2.67 g oxygen. If the relative molecular mass is 60, calculate it’s molecular formula. (C = 12,H = 1, O = 16). Solution: Component elements: C H. O 2.00. 0.34. 2.67 reacting mass 2.00 0.34 2.67 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 12 2 16 0.17 0.34 Reacting mass: No of mole = Mole ratio = 0.17 𝑛𝑜 𝑜𝑓 𝑚𝑜𝑙𝑒 0.17 0.34 0.17 𝑡ℎ𝑒 𝑠𝑚𝑎𝑙𝑙𝑒𝑠𝑡 0.17 0.17 0.17 2 1 1 :. The empirical formula of the compound is CH2O Molecular formula, (CH2O)n = 60 [(12x1)+(1x2)+(16x1)]n = 60 (12+2+16)n = 60 30n = 60 n =60/30 n=2 :. The molecular formula of the compound is C2H4O2 or CH3COOH 2. Determine the empirical formula of an oxide of nitrogen containing 70% oxygen. If the relative molecular mass of the oxide is 92, deduce it’s molecular formula (N = 14, O = 16). Solution % of oxygen = 70%; % of nitrogen = 100-70 = 30% Component elements: N O % composition: 30 70 No of mole = % compositin 30 70 molar mass 12 16 2.14 Mole ratio = 4.38 𝑛𝑜 𝑜𝑓 𝑚𝑜𝑙𝑒 2.14 4.38 𝑡ℎ𝑒 𝑠𝑚𝑎𝑙𝑙𝑒𝑠𝑡 2.14 2.14 1 2 :. The empirical formula of the compound is NO2 Molecular formula, (NO2)n = 92 [(14x1)+(16x2)]n = 92 (14+32)n = 93 46n = 92 n =92/46 n=2 The molecular formula of the oxide is N2O4 LAWS OF CHEMICAL COMBINATION The laws of chemical combination describe the general features of chemical changes. These laws are very important and form the basis of writing chemical formulae and equations. These laws include • • • • Law of conservation of mass Law of definite proportion or constant composition Law of multiple proportion Law of reciprocal proportion 1. Law of Conservation of Mass: This law was proposed by Antoine Lavoisier in 1789. The states that “matter is neither created nor destroyed during chemical reaction but changes from one form to another”. Chemical equations must be balanced to obey this law. 2. Law of Definite Proportion or Constant Composition: This law was proposed by Joseph Louis Proust in 1799. The law of definite proportion states that “all pure samples of a particular chemical compound contain similar elements combined in the same proportion by mass”. This law is based on the fact that when elements combine to form a compound, they do so in fixed proportion by mass so that all pure samples of a compound are identical in composition by mass irrespective of the method of preparation. 3. Law of Multiple Proportion: This law was proposed by John Dalton in 1803. The law of multiple states that “if two elements, A and B, combine to form more than one chemical compound, then the various masses of one element, A, which combine separately with a fixed mass of the other element, B, are in simple multiple ratio”. Examples of paired elements that form more than one compound include: • • • • • Copper and oxygen to form copper (II) oxide, CuO, and copper (I) oxide, Cu2O. Iron and oxygen to form iron (II) oxide, FeO, and iron (III) oxide, Fe2O3 Carbon and oxygen to form carbon (II) oxide, CO, and carbon (IV) oxide CO2 Iron and chlorine to form iron (II) chloride, FeCl2, and iron (III) chloride, FeCl3 Lead and sulphur to form lead (II) sulphide, PbS, and lead (IV) sulphide, PbS2 4. Law of Reciprocal Proportion: This law was proposed by Jeremias Richter in 1892. The law of reciprocal proportion states that “the masses of several elements A, B, C, which combine separately with a fixed mass of another element, D, are the same as or simple multiples of the masses in which A, B, C, themselves combine with one another”. For example, car Examples 1. Two different samples, A and B, of copper (II) oxide were obtained from different sources. When heated in a stream of hydrogen, they were reduced to yield the following results: Sample A B Mass of copper (II) oxide 3.55g 3.02g Mass of copper residue 2.81g 2.42g Proof that the figures above illustrate the law of definite proportion. Solution Sample A B Mass of copper (II) oxide 3.55g 3.02g Mass of copper residue 2.83g 2.42g 2.83 100 X 3.55 1 % of copper residue in copper (II) oxide = 79.7% 2.42 100 X 3.02 1 = 80.1% The percentage of copper residue in the two samples is approximately 80.0 irrespective of the method of preparation of the copper (II) oxide samples. Copper (II) oxide = copper + oxygen 100% = 80% 20% Ratio 4 : 1 :. In pure copper (II) oxide, copper and oxygen are always present in a definite proportion of approximately 4 : 1 by mass. Hence, the figures agree with the law of definite proportion. 2. Copper forms two different oxides, A and B, with oxygen. If 3.04g of oxide A and 1.91g of oxide B contain 2.55g and 2.38g of copper respectively, show that the figures agree with the law of multiple proportion. Write the formulae of the oxides. Solution Oxide A B Mass of oxide 3.04g 1.91g Mass of copper 2.55g 1.38g Mass of oxygen 3.04 – 2.55 = 0.49g 1.91 – 1.38 = 0.53g Calculate the various masses of copper which would combine separately with a fixed mass of oxygen says 1g. Oxide A 0.49g of oxygen ≡ 2.55g of copper to form oxide A :. 1.0g of oxygen ≡ 2.55 X 1.0 0.49 = 5.2g of copper Oxide B 0.53g of oxygen ≡ 1.38g of copper to form oxide B :. 1.0g of oxygen ≡ 1.38 X 1.0 Oxide Mass of copper Ratio of copper 0.53 = 2.6g of coper A. B 5.2g 2.6g 5.2 2.6 2.6 2.6 2 : 1 The masses of copper which combined with a fixed mass of oxygen in oxide A and B are in simple multiple ratio of 2 : 1 hence, the figures agree with the law of multiple proportion. Formula of oxide A is Cu2O and that of oxide B is CuO Practice Questions 1. A metal X forms two different chlorides, A and B. If 12.7g of chloride A and 16.3g of chloride B contain 7.1g and 10.7g of chlorine respectively, show that the figures agree with the law of multiple proportions. Write the formulae of the chlorides. 2. Chemical Equation Chemical equation is the symbolic representation of chemical reaction. The symbols or formulae of the reactants are written on the left hand side (LHS) while those of the products are written on the right hand side (RHS) of the equation. An arrow pointing towards the products is used to link the reactants and the products together. Symbols (s) for solid, (l) for liquid, (g) for gas and (aq) for aqueous are used to indicate the states of reactants and products . For example, the reaction between magnesium and hydrogen chloride to form magnesium chloride and hydrogen gas can be represented by the equation: Mg(s) + 2HCl(g) —> MgCl2(s) + H2(g). Balancing Chemical Equations Chemical equations must be balanced to obey the law of conservation of mass. An unbalanced equation would imply that atoms have been created or destroyed. Rules for Balancing Chemical Equations • • • • • Balance chemical equations by trial and error method The formulae of reactants and products are fixed and cannot be altered, hence never use subscripts to balance chemical equations. Balance chemical equations by placing appropriate number of moles (i.e. coefficient) in front of the reactants and products concerned. After balancing, if a fraction is obtained as the number of a reactant or product, multiply the whole equation by an appropriate number to convert the fraction to a whole number. Do an atom count to check that the equation is balanced. Examples Balance the chemical equations below 1. NH3(g) + O2(g) —> NO(g) + H2O(g) 2. N2(g) H2(g) —> NH3(g) 3. C2H6(g) + O2(g) —> CO2(g) + H2O(g) Solution 1. 2NH3(g) + 2½O2(g) —> 2NO(g) + 3H2O(g) … X 2 4NH3(g) + 5O2(g) —> 4NO(g) + 6H2O(g) 2. N2(g) 3H2(g) —> 2NH3(g) 3. C2H6(g) + 3½O2(g) —> 2CO2(g) + 3H2O(g) ….. X 2 2C2H6(g) + 7O2(g) —> 4CO2(g) + 6mH2O(g) CHEMICAL COMBINATION Chemical combination is the tendency of other elements to attain the stable duplet or octet structure possessed by rare gases. The two main types of chemical combination are: • • Electrovalent/Ionic combination; and Covalent combination Covalent combination is further divided into • • Ordinary covalent combination; and Coordinate or dative combination Electrovalent or Ionic Combination Ionic Combination involves the transfer of electrons from one atom (usually metallic) to another (usually nonmetallic). In ionic combination, one atom (usually metallic) acts as electron donor while the other (usually nonmetallic) acts as electron acceptor. The electrons involved reside in the outermost shell i.e. valence electrons. Elements whose atoms have only one, two or three valence electrons can donate these electrons to attain the stable duplet or octet structure of the rare gases. Similarly, elements whose atoms have 7, 6 or 5 valence electrons can accept 1, 2 or 3 electrons respectively to attain stable octet structure of the rare gases. In ionic combination, after donating their electrons, metallic particles become positively charged while nonmetallic particles become negatively charged after acquiring extra electrons. The number of protons is unaffected by the transfer of electrons during ionic combination