Chemistry Grade 12 Past Paper PDF

G. C. E. (Advanced Level) CHEMISTRY Grade 12 Unit 1: Atomic Structure Unit 2: Structure and Bonding Unit 3: Chemical Calculations Department of Science Faculty of Science and Technology National Institute of Education www.nie.lk...

G. C. E. (Advanced Level) CHEMISTRY Grade 12 Unit 1: Atomic Structure Unit 2: Structure and Bonding Unit 3: Chemical Calculations Department of Science Faculty of Science and Technology National Institute of Education www.nie.lk i Chemistry Resource Book Grade 12 © National Institute of Education First Print – 2018 Department of Science Faculty of Science and Technology National Institute of Education Sri Lanka Published by: Press National Institute of Education Maharagama Sri Lanka ii Message from the Director General The National Institute of Education takes opportune steps from time to time for the development of quality in education. Preparation of supplementary resource books for respective subjects is one such initiative. Supplementary resource books have been composed by a team of curriculum developers of the National Institute of Education, subject experts from the national universities and experienced teachers from the school system. Because these resource books have been written so that they are in line with the G. C. E. (A/L) new syllabus implemented in 2017, students can broaden their understanding of the subject matter by referring these books while teachers can refer them in order to plan more effective learning teaching activities. I wish to express my sincere gratitude to the staff members of the National Institute of Education and external subject experts who made their academic contribution to make this material available to you. Dr. (Mrs.) T. A. R. J. Gunasekara Director General National Institute of Education Maharagama. iii Message from the Director Since 2017, a rationalized curriculum, which is an updated version of the previous curriculum is in effect for the G.C.E (A/L) in the general education system of Sri Lanka. In this new curriculum cycle, revisions were made in the subject content, mode of delivery and curricular materials of the G.C.E. (A/L) Physics, Chemistry and Biology. Several alterations in the learning teaching sequence were also made. A new Teachers’ Guide was introduced in place of the previous Teacher’s Instruction Manual. In concurrence to that, certain changes in the learning teaching methodology, evaluation and assessment are expected. The newly introduced Teachers’ Guide provides learning outcomes, a guideline for teachers to mould the learning events, assessment and evaluation. When implementing the previous curricula, the use of internationally recognized standard textbooks published in English was imperative for the Advanced Level science subjects. Due to the contradictions of facts related to the subject matter between different textbooks and inclusion of the content beyond the limits of the local curriculum, the usage of those books was not convenient for both teachers and students. This book comes to you as an attempt to overcome that issue. As this book is available in Sinhala, Tamil, and English, the book offers students an opportunity to refer the relevant subject content in their mother tongue as well as in English within the limits of the local curriculum. It also provides both students and teachers a source of reliable information expected by the curriculum instead of various information gathered from the other sources. This book authored by subject experts from the universities and experienced subject teachers is presented to you followed by the approval of the Academic Affairs Board and the Council of the National Institute of Education. Thus, it can be recommended as a material of high standard. Dr. A. D. A. De Silva Director Department of Science iv  Guidance Dr. (Mrs.) T. A. R. J. Gunasekara Director General National Institute of Education Supervision Dr. A. D. A. De Silva Director, Department of Science National Institute of Education Mr. R. S. J. P. Uduporuwa Former Director, Department of Science National Institute of Education Subject Leader Mrs. M. S. Wickramasinghe Assistant Lecturer, Department of Science National Institute of Education Internal Editorial Panel Mr. L. K. Waduge Senior Lecturer, Department of Science Mrs. G. G. P. S. Perera Assistant Lecturer, Department of Science Mr. V. Rajudevan Assistant Lecturer, Department of Science Writing Panel Dr. Russel C. L. de Silva - Senior Lecturer, Department of Chemistry, University of Kelaniya (Unit 1) Dr. M.A.B. Prasantha - Senior Lecturer, Department of Chemistry, University of Sri Jayewardenepura (Unit 2) Dr. M.N. Kaumal - Senior Lecturer, Department of Chemistry, University of Colombo (Unit 3) External Editorial Panel Prof. S. P. Deraniyagala - Senior Professor, Department of Chemistry, University of Sri Jayewardenepura Prof. M. D. P. De Costa - Senior Professor, Department of Chemistry, University of Colombo Prof. H. M. D. N. Priyantha - Senior Professor, Department of Chemistry, University of Peradeniya Prof. Sudantha Liyanage - Dean, Faculty of Applied Sciences, University of Sri Jayewardenepura Mr. K. D. Bandula Kumara - Deputy Commissioner, Education Publication Department, Ministry of Education Mrs. Deepika Nethsinghe - SLTS-1 (Rtd), Ladies College, Colombo 07 v  Mrs. Muditha Athukorala - SLTS-1, Prajapathi Balika Vidyalaya, Horana Miss. C. A. N. Perera - SLTS-1, Princess of Wales’, Moratuwa Mrs. V. K. W. D. Salika Madavi - SLTS-1, Muslim Ladies College, Colombo 04 Mrs. H.M.D.D. D. Manike - SLTS-1, Viharamhadevi Balika Vidyalaya, Kiribathgoda Mr. S. Thillainathan - SLTS-1 (Rtd), Hindu Ladies College, Colombo 06 Miss. S. Veluppillai - SLTS-1 (Rtd), Hindu Ladies College, Colombo 06 Mrs. M. Thirunavukarasu - SLTS-1 (Rtd), Hindu Ladies College, Colombo 06 Mrs. S. Rajadurai - SLTS-1 (Rtd), St. Peters' College, Colombo 04 Language Editing Dr. Chandra Amarasekara Consultant, National Institute of Education Mr. M. A. P. Munasinghe Chief Project Officer (Rtd.), National Institute of Education Cover Page Mrs. R. R. K. Pathirana Technical Assitant, National Institute of Education Supporting Staff Mrs.Padma Weerawardana Mr. Mangala Welipitiya Mr. Ranjith Dayawansa vi  Content Message from the Director General ……………………………………………………………..…… iii Message from the Director ………………………………………………………………………........ iv Subject Committee………………………………………………………………………………..…... v 1.0 Atomic structure………………………………………………..………….....…01-42 1.1 The atomic theory of matter…………………………..…...………...…….………02 1.1.1 Properties of cathode rays (Experimental observations) 1.1.2 The nucleus of the atom 1.1.3 Properties of positive rays (Experimental observations) 1.1.4 Rutherford’s gold foil experiment 1.1.5 Atomic number, isotopes and mass number 1.1.6 The atomic mass scale 1.1.7 Average atomic mass and relative atomic mass of an element 1.1.8 Ions 1.2 Electromagnetic radiation and wave-like properties of matter………......….…..13 1.2.1 Quantization of energy 1.3 Electronic energy levels of atoms………………………………………...………...17 1.3.1 The hydrogen spectrum 1.3.2 Shapes of orbitals 1.3.3 Orbitals and quantum numbers 1.4 Electron configuration …………………………………………………...………...23 1.4.1 Aufbau principle 1.4.2 The Pauli exclusion principle 1.4.3 Hund's rule 1.4.4 Condensed electron configurations 1.5 Building of periodic table …………………………………………………….….…28 1.6 Periodic trends shown by s and p block elements …………………......……….…32 1.6.1 Sizes of atoms and ions 1.6.2 Ionization energy 1.6.3 Electron gain energy 1.6.4 Electronegativity 2.0 Structure and bonding………………………………………………..………...43-84 2.1 Covalent bonds……………………………………………...………………………44 2.1.1 Lewis dot diagrams and Lewis dot dash structures 2.2 Dative covalent bonds…………………………..…………………………………..51 2.3 Valance Shell Electron Pair Repulsion theory (VSEPR theory) ………………..52 2.3.1 Hybridization of atomic orbitals 2.3.2 Formation of double and triple bonds 2.3.3 Resonance structures 2.3.4 Effect of electronegativity and geometry on the polarity of molecules 2.3.5 Dipole moment 2.3.6 Factors affecting the magnitude of electronegativity 2.4 Ionic bonds/ ionic interactions …………………...……………………...…….…..75 2.5 Metallic bonds………………………………………..………………………….….78 vii 2.6 Secondary interactions………………...…………………………………………...79 3.0 Chemical calculations………………………………..……………..…….……87-118 3.1 Oxidation Number…………………………………………………...……..…….…88 3.1.1 Basic rules that applied in the determination of the oxidation states of an atom in a molecule/ polyatomic ion or in a compound 3.1.2 Use of oxidation states to understand electron transfer between atoms in redox reactions 3.2 Nomenclature of inorganic compounds ………………………………………..….93 3.2.1 Names of ionic compounds derived from monoatomic ions 3.2.2 Names of ionic compounds derived from elements that form more than one type of cation 3.2.3 Names of simple covalent compounds 3.2.4 Polyatomic ions 3.2.5 Inorganic acids 3.3 Atomic mass, mole and Avogadro constant ………………………………………97 3.3.1 The connection between atomic mass unit, moles and Avogadro constant 3.3.2 Calculation of average atomic mass of elements 3.3.3 Mole 3.3.4 Molar mass 3.4 Types of chemical formulae ………………………………………………………..99 3.4.1 Chemical calculations using chemical formulae 3.4.2 Determination of the formulae of a compound 3.4.3 Determination of molecular formula using the empirical formula mass and molecular mass 3.5 Composition of a substance in a mixture ……………………………………..…100 3.5.1 Composition given in fractions 3.5.2 Percentage composition in a solution 3.5.3 Molality 3.5.4 Molarity 3.6 Balancing chemical reactions ………………………………………………...…..105 3.6.1 Balancing a chemical reaction by inspection method 3.6.2 Balancing a chemical reaction by the redox method 3.6.3 Balancing simple nuclear reactions 3.7 Preparation of solutions …………………………………………………………..113 3.8 Calculations based on chemical reactions …………………………………...…..115 viii  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure 1.Atomic Structure Content 1.1 The Atomic Theory of Matter 1.3.2 Shapes of orbitals 1.1.1 Properties of cathode rays (Experimental 1.3.3 Orbitals and quantum numbers observations)  The principal quantum number (n) 1.1.2 The nucleus of the atom  The angular momentum quantum 1.1.3 Properties of positive rays (Experimental number (l) observations)  The magnetic quantum number (ml) 1.1.4 Rutherford’s gold foil experiment  The spin quantum number (ms) 1.1.5 Atomic number, isotopes and mass number 1.1.6 The atomic mass scale 1.4 Electron configuration 1.1.7 Average atomic mass and relative atomic 1.4.1 Aufbau principle mass of an element 1.4.2 The Pauli exclusion principle 1.1.8 Ions 1.4.3 Hund's rule 1.4.4 Condensed electron configurations 1.2 Electromagnetic radiation and wave- like properties of matter 1.5 Building of periodic table  Electromagnetic radiation · Properties  The long form of the periodic table [speed (c), wavelength (λ), frequency (ν), energy (E)] 1.6 Periodic trends shown by s and p block 1.2.1 Quantization of energy elements  Electromagnetic spectrum 1.6.1 Sizes of atoms and ions  c=νλ  van der Waals radius  E = h ν, λ = ℎ  Covalent radius 𝑚𝑉  Wave- particle dual nature of matter  Metallic radius  Periodic trends in atomic radii 1.3 Electronic energy levels of atoms  Electron configurations of ions  Variation of successive ionization  Periodic trends in ionic radii energies of elements 1.6.2 Ionization energy 1.3.1 The hydrogen spectrum  Periodic trends in first ionization energies  Existence of electrons in energy levels 1.6.3 Electron gain energy 1.6.4 Electronegativity 1  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Introduction Chemistry is the study of the properties and behaviour of matter. Matter is the physical material of the universe; it is anything that has mass and occupies space. Although the materials in our world vary greatly in their properties, everything is formed from only about 100 elements and, therefore, from only about 100 chemically different kinds of atoms. (118 elements have been discovered so far but the heavier atoms are short lived and not found naturally.) 1.1 The atomic theory of matter Philosophers from the earliest times speculated about the nature of the fundamental components from which the world is made. Empedocles (~ 440 BC) believed that the four elements-earth, fire, air and water made up all things. The Hindus believed that the four elements stated above makeup the world and space. However, Democritus (460– 370 BC) and other early Greek philosophers described the material world as being made up of tiny, invisible, indivisible particles that they called ‘atomos’, meaning “indivisible” or “uncuttable.” Later, however, Plato and Aristotle formulated the notion that there can be no ultimately indivisible particles, and the “atomic” view of matter faded for many centuries during which Aristotelean philosophy dominated the Western culture. It was in 1808 that an English scientist and school teacher, John Dalton (1766-1844), formulated a precise definition of the indivisible building blocks of matter that we call atoms. Dalton’s atomic theory was based on four postulates. 1. Elements are made out of extremely small, indivisible particles called atoms. 2. All atoms of a given element are identical in mass and size, but the atoms of one element are different from the atoms of all other elements. 3. Atoms of one element cannot be changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. 4. Compounds are formed by union of two or more atoms of different elements in a simple numerical ratio. Dalton’s atomic model is called the "Golf ball model". (a) (b) Figure1.1 (a) John Dalton and (b) the golf ball model 2  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Johnstone G. Stoney (1826-1911) named the fundamental unit carrying electricity as “electrons” in 1891 but did not have any experimental evidence of its existence. During the mid-1800s, scientists began to study electrical discharge through a glass tube pumped almost empty of air. This device was an invention of the British chemist and physicist Sir William Crookes (1832-1919) and is called Crookes tube or cathode ray tube. Figure 1.2 A cathode ray tube The experiment of Crookes and the others showed that when two electrodes are connected to a high-voltage source, the heated negatively charged plate, called the cathode, produced a stream of invisible radiation. Although the rays could not be seen, their presence was detected because they cause gases at low pressure to glow and which made other substances to fluoresce, or to give off light. The radiation emitted from the cathode was given the name 'cathode rays'. Later it was known that these rays could be deflected by a magnetic field and they carried a negative electrical charge. Some scientists felt that these rays were waves and others were inclined to think they were particles. The British scientist J. J. Thomson (1856–1940) observed that cathode rays are the same regardless of the identity of the cathode material or the gas in the tube. In 1897 he described cathode rays as streams of negatively charged particles. He used a cathode tube with an anode that had a hole at the centre. Using experimental measurements obtained from that cathode tube he then calculated a value of 1.76 x 108 coulombs per gram (C g-1) for the ratio of the electron’s electrical charge to its mass. Anode (+) Cathode (-) High voltage Fluorescent screen Figure 1.3 Thomson’s cathode ray tube 3  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure 1.1.1 Properties of cathode rays (Experimental observations)  Cathode rays travel in straight lines. When an opaque object like a metal cross is placed in the path of cathode rays in a discharge tube, a shadow with sharp edges of the metal cross is formed at the end opposite to the cathode. The placement of the shadow proves that cathode rays emit from the cathode and they travel in a straight line. Cathode Cathode Paddle rays rays wheel Cathode Cathode Anode Anode Shadow of (metal object) the metal Figure 1.4 Cathode ray properties  Cathode rays are a beam of particles having mass and possess kinetic energy. On placing a light paddle wheel in the path of cathode rays in a discharge tube, the blades of the paddle wheel rotate. This was considered evidence that electrons (cathode rays) have momentum. (However there is doubt on this conclusion as heating of the tube can also make the paddles move.)  When an electric field is applied in the path of cathode rays, they are deflected towards the positively charged plate. Hence the cathode rays are composed of negatively charged particles. They are affected by magnetic fields showing a deflection perpendicular to the magnetic field. The direction of deflection is similar to the deflection of any other negatively charged particles. Therefore electron can be concluded as a negatively charged particle too. Cathode Anode Cathode Electric field rays Figure 1.5 Interaction of cathode rays with external electrical fields 4  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure  The nature of the cathode rays does not depend on the nature of the gas taken in the discharge tube or the material of the cathode.  The ratio of the charge to mass (e/m ratio) of cathode ray particles obtained from different gases was found to be exactly the same. Electron s Positively charged sphere Figure 1.6 J. J. Thomson and his model Using his findings in 1899 J.J. Thomson postulated the “plum-pudding” model of atomic structure. In 1909, Robert Millikan (1868–1953) succeeded in measuring the charge of an electron as 1.602 x 10-19 C by performing the oil drop experiment. The mass of the electron could be calculated by using the experimental values for the charge of electron and Thomson’s charge-to-mass ratio. 1.602 ×10−19 C Electron mass = 1.76 ×108 C g−1 = 9.10 × 10−28 g Figure 1.7 Robert Millikan and mass of an electron This mass is about 1/1837 of a hydrogen atom which is the lightest atom. The relative charge of an electron is -1. 1.1.2 The nucleus of the atom The German physicist, Eugen Goldstein experimentally proved the existence of positive charges in matter. In his experiments, a perforated cathode was used in a discharge tube along with air at very low pressure. When a high voltage of about 10,000 volts was applied across the electrodes, a faint red glow was observed behind the perforated cathode. When the high voltage is applied, its electric field accelerates the small number of ions present 5  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure in the gas. These collide with atoms of the gas, knocking electrons off of them and creating more positive ions. These ions and electrons in turn strike more atoms, creating more positive ions. The positive ions are all attracted to the cathode, and some pass through the holes in the cathode. Goldstein called these positive rays "canal rays", because they were produced by the holes or channels in the cathode. Although the rays are not exactly formed at the positive electrode or anode, since they are formed away from the cathode close to the anode, they were also known as anode rays or positive rays. Figure 1.8 A cathode ray tube with a perforated cathode 1.1.3 Properties of positive rays (Experimental observations)  They travel in straight lines and cast a shadow of the objects placed in their way.  They can move a paddle wheel placed in their path.  These rays are positively charged and when an electric field is applied in the path of the rays, they are deflected towards the negative plate of an electric field.  The nature of the positive rays depends upon the gas taken in the discharge tube. Different gases give different types of positive rays, which contain particles having different masses and different charges. Therefore the e/m ratio is not constant for positive ray particles obtained from different gases. In 1907 a study of how this "ray" was deflected in a magnetic field, revealed that the particle making up the ray were not all the same mass. The lightest ones, formed when there was some hydrogen gas in the tube, were calculated to be about 1840 times heavier than an electron. The mass of any other positive particle is a multiplication of the mass of the lightest positive particle. Therefore it should be a subatomic particle. They were named as protons. The relative mass of a proton is 1, hence, the mass of a proton is 1.6 x 10-24 g or 1.007276 u (atomic mass units) or Da (Daltons). (The unit was earlier given the name amu.) The proton has a charge equal and opposite to that of an electron. Hence the absolute charge of a proton is 1.6 x 10-19 coulomb and its charge is positive. Proton is the smallest positive charge carrying particle in an atom and the relative charge of proton is +1. 6  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Following the discovery of radioactivity in 1896 by the French scientist Henri Becquerel (1852–1908), the British physicist Lord Ernest Rutherford (1871-1937) showed that radioactive materials produce three types of emissions alpha (), beta () and gamma (). The paths of  and  radiation are bent by an electric field. Alpha () rays consist of positively charged particles, called particles, and therefore are deflected towards the positively charged plate. Beta () rays or particles have the identity of electrons and are deflected towards the negatively charged plate. The third type of radioactive radiation consists of high–energy rays called gamma () rays. Like X rays,  rays have no charge and are not affected by an external electric or magnetic field. Lead block  rays  rays  rays Radioactive Electrically substance charged plates Slit Figure 1.9 Behaviour of alpha (), beta () and gamma () rays in an electric field (a) (b) Figure 1.10 (a) Henri Becquerel and (b) Lord Ernest Rutherford 1.1.4 Rutherford’s gold foil experiment In 1908-09, Rutherford together with his associate Johannes Hans Wilhelm Geiger (1882-1945) a German physicist and an undergraduate named Ernest Marsden, carried out a series of experiments using very thin foils of gold and other metals as targets for particles from a radioactive source. 7  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Source of alpha particles Deflected Thin gold foil Scattered  particles particles Fluorescent screen Beam of particles Circular fluorescent screen Most particles are undeflected Thin gold foil Figure 1.11 Rutherford’s gold foil experiment They observed that the majority of particles penetrated the foil either unreflected or only with a slight deflection. They also noticed that a few particles were scattered (or deflected) at a large angle. Very few  particles bounced back in the direction from which it came. To explain the results of the experiment, Rutherford devised a new model of atomic structure, suggesting that most of the atom must be empty. This structure would allow most of the  particles to pass through the gold foil with little or no deflection. The atom’s positive charges, Rutherford proposed, are all concentrated in the nucleus, a dense central core within the atom. Whenever an  particle came close to a nucleus in the scattering experiment, it experienced a large repulsive force and therefore a large deflection. Moreover, an  particle traveling directly toward a nucleus would experience an enormous repulsion that could completely reverse the direction of the moving particle. Nucleus Electron Figure 1.12 Rutherford’s model (1911) Subsequent studies, mainly based on mass spectroscopy revealed that the masses of atoms were much greater than the masses of protons and electrons present. Therefore another 8  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure subatomic particle should be present to contribute towards the mass of the atom. In1932 Sir James Chadwick (1891-1972) a British scientist discovered the ‘neutron’. The charge of a neutron is 0 (zero) and its mass is 1.6749 x 10-24g or 1.008665 amu. (a) (b) Figure 1.13 (a) James Chadwick and (b) Niels Bohr Since Rutherford’s time, physicists have learned more and more about atomic nuclei. In 1913 Niels Henrik David Bohr (1885-1962) a Danish physicist, combined the ideas at that time and suggested that the atomic nucleus was surrounded by electrons moving in orbit, like planets around the sun. He postulated that the electrons in order to remain in orbit, the electrostatic attraction between the nucleus and electron must be equal to the centrifugal force. In other words, the electrons have to travel in a constant speed around the nucleus keeping the distance from the nucleus constant. The model he introduced is known as the Rutherford–Bohr model or the Bohr model. Particles found in the nucleus are called nucleons, including the protons and neutrons in to the atom. A nuclide is the nucleus of an atom that has specific numbers of protons and neutrons (all nucleons). Therefore, nuclides are composite particles of nucleons. Nucleus Electrons Figure 1.14 The Bohr model 1.1.5 Atomic number, isotopes and mass number Henry Gwynn Jeffrey Moseley (1887-1915), an English physicist and a co-worker of Rutherford, found that the number of positive charges on the nucleus increases in atoms 9  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure by single electron units. The atoms of each element have a characteristic number of protons. The number of protons in an atom of any particular element is called that element’s atomic number. Atomic Number of protons Number of electrons number (Z) = in the nucleus = in an atom Since an atom has no net electrical charge, the number of electrons it contains is equal to the number of protons found in its nucleus. All atoms of carbon, for example, have six protons and six electrons, whereas all atoms of oxygen have eight protons and eight electrons. Thus, carbon has atomic number 6 and oxygen has atomic number 8. British scientists J. J. Thomson and Francis William Aston (1877-1945) perfected the ‘mass spectrometer' which they used in 1912-13 to discover the first isotopes (of neon). Atoms of a given element can differ in the number of neutrons they contain and therefore their mass can also vary. The number of protons plus neutrons (nuclide) in an atom is called its mass number. Mass number (A) = Number of protons (Z) + Number of neutrons In the atomic symbol used to represent a particular atom the mass number is given at the top left of the element symbol and the atomic number may be given at the bottom left. However, since the symbol also gives the same information, the atomic number usually may not be shown in the symbol. Figure 1.15 Atomic symbol of carbon 10  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Example 1.1 How many protons, neutrons and electron are there in an atom of 197Au? Answer: The superscript 197 is the mass number (protons + neutrons). According to the periodic table gold has atomic number 79. Hence, an atom of 197 Au has 79 protons, 79 electrons and 197-79 = 118 neutrons. Atoms with identical atomic numbers but different mass numbers (that is, same number of protons but different numbers of neutrons) are called isotopes of one another. For example, while most atoms of carbon have six neutrons, some have more. The carbon atoms containing six protons and six neutrons have a mass number 12 and are depicted as 12C, while atoms that contain six protons and seven neutrons have mass number 13 and are depicted as 13C. The atoms with six protons and eight neutrons have mass number 14 and are depicted as 14C. Isotopes of an element that are stable in nature are called stable isotopes and those that are not stable called radioisotopes. 1.1.6 The atomic mass scale Atoms are very small pieces of matter, however they have a mass. However, it is convenient to use the unified atomic mass unit (u) when dealing with these extremely small masses, where; 1 u or 1Da (earlier amu)  𝟏𝟐 𝐠 𝟔.𝟎𝟐𝟐𝟏𝟒 ×𝟏𝟎𝟐𝟑 × 𝟏 𝟏𝟐 = 1.66054 × 10−24 g 1 u = 1.66054 x 10-24 g and 1 g = 6.02214 x 1023 u or Da The unified atomic mass unit is defined as a mass of exactly 1/12 of a chemically unbound atom of the 12C isotope of carbon. In these units, a 1H atom has a mass of 1.0078 u or Da and a 16O atom has a mass of 15.9949 u or Da. 1.1.7 Average atomic mass and relative atomic mass of an element Most elements occur in nature as mixtures of isotopes. The mass of an atom can be given as relative atomic mass or atomic mass. It can be obtained by summing over the masses of its isotopes multiplied by their relative abundances: Average atomic mass = isotope mass) x (fractional isotope abundance)] 11  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Example 1.2 Naturally occurring carbon is composed of 98.93% 12C, 1.07% 13C and a negligible amount of 14C. The relative masses of two initial isotopes are 12 u (exactly) and 13.00335 u, respectively. Calculate the average atomic mass of carbon. Answer: (0.9893 × 12.00 u) + (0.0107 × 13.00335 u) = 12.01 u Relative atomic mass = 12.01 When the atomic mass is given as mass per mole of atoms (in units of g mol-1), it is known as the molar mass of the element or atom. Since 1 g = 6.02214 x 1023 u and 1 mole of atoms = 6.02214 x 1023 atoms, the molar mass of carbon will be 12.01 g mol-1. Relative atomic mass (Ar) is a dimensionless physical quantity. It is the ratio of the average mass of atoms of an element to 1⁄12 the mass of an atom of carbon-12 (known as the unified atomic mass unit). Hence the relative atomic mass of carbon will be 12.01. In periodic tables the relative atomic mass of the element is usually given below the symbol of the element. 1.1.8 Ions The nucleus of an atom is unchanged by chemical processes, but some atoms can readily gain or lose electrons. If electrons are removed from or added to an atom, a charged particle called an ion is formed. An ion with a positive charge is a cation and a negatively charged ion is an anion. 12  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure e.g.: The sodium atom, which has11 protons and 11 electrons can easily lose one electron. The resulting cation has 11 protons and 10 electrons, which means it has a net charge of +1. Figure 1.16 Ionization of sodium atom e.g.: Chlorine, with 17 protons and 17 electrons, can gain an electron in chemical reactions, producing the Cl- ion. Figure 1.17 Formation of chloride ion The net charge on an ion is represented by a superscript right to the atomic symbol. Hence the symbol for the ferric ion (an iron atom that has lost 3 electrons) will be: In addition to simple ions such as Na+ and Cl-, there are polyatomic ions, given as NH4+ (ammonium ion) and SO2− 4 (sulfate ion), which consist of joined atoms to form molecules carrying a net positive or negative charge. 1.2 Electromagnetic radiation and wave like properties of matter Much of our present understanding of the electronic structure of atoms has come from analysis of the light either emitted or absorbed by substances. Electromagnetic radiation (EMR) consists of electromagnetic waves, which are synchronized oscillations of electric and magnetic fields that propagate at the speed of light through a vacuum. The 13  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure oscillations of the two fields are perpendicular to each other and perpendicular to the direction of wave propagation. Figure 1.18 Electro-magnetic radiation The light we see with our eyes, visible light, is one type of electromagnetic radiation. All types of electromagnetic radiation move through a vacuum at a speed of 2.998 x108 m s-1, the speed of light (c) and have wave-like characteristics. Waves are periodic, which means that the pattern of peaks and troughs repeats itself at regular intervals. The distance between two adjacent peaks or between two adjacent troughs (distance between a cycle) is called the wavelength (). The number of complete wavelengths, or cycles, that pass a given point each second is the frequency () of the wave. Frequency is expressed in cycles per second using a unit called hertz (Hz). Since it is understood that cycles are involved, it may also be simply called “per second,” which is denoted by s- 1. Hence, c =   Figure 1.19 An electromagnetic wave     14  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Example 1.3 The yellow light given off by a sodium vapour lamp used for public lighting has a wavelength of 589 nm. Calculate the frequency of this radiation. Answer: c 3.00 × 108 m/s 1 nm  = ( ) ( −9 ) = 5.09 × 1014 s-1  589 nm 10 m Different types of electromagnetic radiations have different properties due to their different wavelengths. The display of electromagnetic radiation arranged in order of increasing wavelength is called the electromagnetic spectrum. Figure 1.20 The electromagnetic spectrum 1.2.1 Quantization of energy In 1900 a German physicist named Max Planck (1858–1947) proposed that energy is quantized; that is the energy can be either released or absorbed by atoms only in discrete quantities of some minimum size. Planck gave the name quantum (meaning “fixed amount”) to the smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation. He proposed that the energy E, of a single quantum equals a constant time the frequency of the radiation: E = h The constant h is called Planck constant and has a value of 6.626 x 10-34 J s. 15  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure (a) (b) Figure 1.21 (a) Albert Einstein and (b) Max Planck In 1905, Albert Einstein (1879–1955) extending Planck’s quantum theory, deduced that the radiant from a metal surface behaves like a stream of tiny energy packets. Each packet, which is like a “particle” of energy, is called a photon and each photon must have an energy equal to Planck constant times the frequency of the light: Energy of a photon = E = h Example 1.4 Calculate the energy of one photon of yellow light that has a wavelength of 589 nm. Answer: c  = 5.09 × 1014 s-1  E = hv = (6.626 × 10−34 J s × 5.09 × 1014 s −1 ) = 3.37 × 10−19 J If one photon of radiant energy supplies 3.37 x 10-19 J, Energy of one mole of photons = (6.02 x 1023 mol-1) (3.37 x10-19 J) = 2.03 x 105 J mol-1 of energy. In the years following the development of Bohr’s model for the hydrogen atom, scientists established that depending on the experimental circumstances, radiation appears to have a wave-like and a particle-like (photon) character. Louis de Broglie (1892–1987), extended the idea that if radiant energy could, under appropriate conditions, behave as though it were a stream of particles (photons), matter under appropriate conditions, possibly would show the properties of a wave. 16  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure De Broglie suggested that an electron moving about the nucleus of an atom behaves like a wave and therefore has a wavelength. He proposed that the wavelength of the electron, or of any other particle, depends on its mass, m, and on its velocity, : h h  = 𝒑 = 𝒎𝐯 where h is the Planck constant. The quantity 𝑚v for any object is called its momentum (p). Because de Broglie’s hypothesis is applicable to all matter, any object of mass (m) and velocity(v) would give rise to a characteristic matter wave. However, the wavelength associated with an object of ordinary size, such as a golf ball, is so tiny as to be completely unobservable. This is not so for an electron because its mass is so small. 1.3 Electronic energy levels of atoms The ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. The magnitude of the ionization energy tells us how much energy is required to remove an electron; the greater the ionization energy, the more difficult it is to remove an electron. The ionization energies for a given element increase as successive electrons are removed. This trend is because with each successive removal, an electron is being pulled from an increasingly more positive ion, requiring increasingly more energy. The sharp increase in ionization energy that occurs when an inner-shell electron is removed is a clear evidence for the fact that electrons are in discrete energy levels. Table 1.1 Successive values of ionization energies I, for the elements from sodium to argon (kJ mol-1) Element I1 I2 I3 I4 I5 I6 I7 Na 496 4562 Mg 738 1451 7733 (Inner-shell electrons) Al 578 1817 2745 11577 Si 786 1577 3232 4356 16091 P 1012 1907 2914 4964 6274 21267 S 1000 2252 3357 4556 7004 8496 27107 Cl 1251 2298 3822 5159 6542 9362 11018 Ar 1521 2666 3931 5771 7238 8781 11995 17  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure 1.3.1 The hydrogen spectrum Most common radiation sources, including light bulbs and stars, produce radiation containing many different wavelengths. A spectrum is produced when radiation from such sources is separated into its component wavelengths. This range of colours, containing light of all wavelengths, is called a continuous spectrum. Not all radiation sources produce a continuous spectrum. When a high voltage is applied to tubes that contain different gases under reduced pressure, the gases emit different colours of light. For example the light emitted by neon gas is the familiar red-orange glow of many “neon” lights. When light emitted from such tubes is passed through a prism, only a few wavelengths are present in the resultant spectrum. A spectrum containing radiation of only specific wavelengths is called a line spectrum. H 400 450 500 550 600 650 700 nm Figure 1.22 Line spectrum of hydrogen Scientists studied the line spectrum of hydrogen thoroughly in the mid-1800s. At that time, only four lines at wavelengths of 410 nm (violet), 434 nm (blue) 486 nm (blue- green) and 656 nm (red) were observed in the spectrum. Bohr's atomic model together with Planck’s idea that energies are quantized was capable of explaining the line spectrum of hydrogen. 18  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Continuum 0 kJ Pfund series Brackett series -82 kJ Energy (kJ mol-1) Paschen series -146 kJ Balmer series -328 kJ Lyman series -1312 kJ Figure 1.23 Possible electron emissions in hydrogen Each allowed orbit around the atom corresponds to a different value of n (n is a whole number). The radius of the orbit gets larger as n increases. Thus, the first allowed orbit (the one closest to the nucleus) has n = 1, the next allowed orbit (the one second closest to the nucleus) has n = 2, and so forth. The line spectra are the result of emission when electrons fall from an initial energy level (ni) to a final energy level (nf), so Ephoton = h = hc/  = -E = (Ef - Ei) for these transitions. E is negative for emission as nf is less than ni; the electron is falling in energy to a lower-energy orbit. The possible emissions results in the line spectra seen in hydrogen. 19  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure 0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 Frequency (Hz)/ 1015 Paschen series Balmer series Lyman series (infra-red) (partly visible) (ultra-violet) Figure 1.24 (a) The line spectrum of hydrogen Figure 1.24(a) shows the spectrum as it varies with frequency. However, the variation of wavelength has shown in Figure 1.24(b). ni = 2 ni = 4 ni = 3 ni = 6 ni = 5 ni = 4 Lyman series Balmer series Paschen series nf = 1 nf = 2 nf = 3 Wavelength () 1875 nm 365 nm 820 nm 656 nm 122 nm 91 nm Figure 1.24 (b) The line spectrum of hydrogen 1.3.2 Shapes of orbitals An electron’s probable location in space around an atom (shapes of orbitals) shows us how the electron density is distributed around the nucleus. The electron density for an s orbital is spherically symmetric and centered on the nucleus or in other words the s orbitals are spherical in shape. 1s 2s 3s Figure 1.25 Shapes of the s orbitals 20  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Each p subshell has three orbitals, corresponding to the three allowed values of ml: -1, 0 and +1. The electron density is not distributed spherically as in an s orbital. Instead, the density is concentrated in two regions on either side of the nucleus, separated by a node at the nucleus. This dumbbell-shaped orbital has two lobes. For each value of n, the three p orbitals have the same size and shape but differ from one another in spatial orientation. It is convenient to label these as px, py, and pz orbitals where the letter subscript indicates the Cartesian axis along which the orbital is oriented. Figure 1.26 Shapes of the p orbitals The different d orbitals in a given shell have different shapes and orientations in space. The shapes of the f orbitals are even more complicated than those of the d orbitals. 1.3.3 Orbitals and Quantum Numbers The Bohr model introduced a single quantum number, n, to describe an orbit. The quantum mechanical model uses three quantum numbers, n, l and ml, which result naturally from the mathematics used to describe an orbital that electrons occupy in an atom and ms that describes the spin of the electron. 1. The principal quantum number n, can have positive integral values 1, 2, 3,.... This quantum number defines the main energy level (electron shell) that the electron occupies in the atom. As n increases, the orbital becomes larger, and the electron spends more time further from the nucleus. 2. The angular momentum (or azimuthal) quantum number l, can have integral values from 0 to (n - 1) for each value of n. This quantum number defines the shape of the orbital. The value of l for a particular orbital is generally designated by the letters s, p, d and f corresponding to l values of 0, 1, 2 and 3 respectively. The set of orbitals that have the same n and l values is called a subshell. Each subshell is designated by a number (the value of n) and a letter (s, p, d or f corresponding to the 21  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure value of l). For example, the orbitals that have n = 3 and l = 2 are called 3d orbitals and are in the 3d sub shell. 3. The magnetic quantum number ml, can have integral values between -l and l, including zero. This quantum number describes the orientation of the orbital in space and the number of possible values for l denotes the number of orbitals in the subshell. For example for l = 2 the values possible will be 2, 1, 0, -1 and -2 denoting that a d subshell has five orbitals. 4. The spin quantum number ms. Two possible values are allowed for ms, + ½ or -1/2, which indicate the two opposite directions in which the electron can spin. A spinning charge produces a magnetic field. The two opposite directions of spin therefore produce oppositely directed magnetic fields. Table 1.2 Relationship between values n, l, and ml Number of Total number n Possible Subshell Possible orbitals in of orbitals in values of l designation values of ml subshell shell 1 0 1s 0 1 1 2 0 2s 0 1 4 1 2p -1, 0, 1 3 3 0 3s 0 1 9 1 3p -1, 0, 1 3 2 3d -2, -1, 0, 1, 2 5 4 0 4s 0 1 16 1 4p -1, 0, 1 3 2 4d -2, -1, 0, 1, 2 5 3 4f -3, -2, -1, 0, 7 1, 2, 3 The restrictions on possible values for quantum numbers give rise to the following very important observations: 1. The shell with principal quantum number n consists of exactly n subshells. Each subshell corresponds to a different allowed value of l from 0 to (n - 1). Thus, the first shell (n = 1) consists of only one subshell, the 1s (l = 0); the second shell (n = 2) consists of two subshells, the 2s (l = 0) and 2p (l = 1); the third shell consists of three subshells, 3s, 3p, and 3d, and so forth. 2. Each subshell consists of a specific number of orbitals. Each orbital corresponds to a different allowed value of ml. For a given value of l, there are (2l + 1) allowed values 22  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure of ml, ranging from -l to +l. Thus, each s (l = 0) subshell consists of one orbital; each p (l = 1) subshell consists of three orbitals; each d (l = 2) subshell consists of five orbitals, and so forth. 3. The total number of orbitals in a shell is n2, where n is the principal quantum number of the shell. The resulting number of orbitals for the shells 1, 4, 9 and 16 is related to a pattern seen in the periodic table: We see that the number of elements in the rows of the periodic table 2, 8, 18 and 32 equals twice these numbers. 1.4 Electron configuration When considering the electronic structures of atoms: In a many-electron atom, for a given value of n, the energy of an orbital increases with increasing value of l. For example, n = 3 orbitals increase in energy in the order 3s < 3p < 3d and all orbitals of a given subshell (such as the five 3d orbitals) have the same energy, just as they do in the hydrogen atom. Orbitals with the same energy are said to be degenerate. 1.4.1 The Aufbau principle The filling of electrons in an atom begins with the subshell lowest energy and continues upwards according to the “Aufbau principle” (The German word 'Aufbau' means 'building up'). Start here and move along the arrows one by one. Figure 1.27 Sequence of filling of electrons Thus the general order of energies of energy levels and sub energy levels are shown in Figure 1.28. 23  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure 5f 7s 6d 6p 6s 5d 4f 5p Potential Energy 5s 4d 4p 4s 3d 3p 3s 2p 2s 1s Figure 1.28 Order of the energy levels in an atom 1.4.2 The Pauli exclusion principle The Pauli Exclusion Principle, which was postulated by Wolfgang Pauli in 1925, states that no two electrons in an atom can have the same set of four quantum numbers n, l, ml, and ms. For a given orbital, the values of n, l, and ml are fixed. Thus, if we want to put more than one electron in an orbital and satisfy the Pauli Exclusion Principle, our only choice is to assign different ms values to the electrons. This indicates that an orbital can hold a maximum of two electrons and they must have opposite spins. This restriction allows us to index the electrons in an atom by giving their quantum numbers. Thus, each s subshell which consists of one orbital can hold a maximum of two electrons; each p subshell that consists of three orbitals can hold a maximum of six electrons; each d subshell which consists of five orbitals can hold a maximum of ten electrons, and so forth. Electrons are distributed among the various orbitals of an atom based on the relative energies of orbitals and the Pauli Exclusion Principle and this distribution is called the electron configuration of the atom. The most stable electron configuration, known as the ground state, is that in which the electrons are in the lowest possible energy states. According to the Pauli Exclusion Principle however, there can be only two electrons in any single orbital. Thus, the orbitals are filled in the order of increasing energy, with no more than two electrons per orbital. 24  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure For example, in the lithium atom, which has three electrons, the 1s orbital can accommodate two electrons. The third one goes into the next lowest-energy orbital, the 2s. Any electron configuration can be represented by writing the symbol for the occupied subshell and adding a superscript to indicate the number of electrons in that subshell. For example, for lithium we write 1s22s1 and for sodium 1s22s22p63s1. In another representation, called an orbital diagram, each orbital is denoted by a box or circle and each electron by a half arrow/ full arrow. A half/ full arrow pointing up represents an electron with a positive spin magnetic quantum number (ms = +1/2) and a half/ full arrow pointing down represents an electron with a negative spin magnetic quantum number (ms = -1/2). Li 1s 2s Electrons having opposite spins are said to be paired when they are in the same orbital. An unpaired electron is one not accompanied by a partner of opposite spin. In the lithium atom the two electrons in the 1s orbital are paired and the electron in the 2s orbital is unpaired. 1.4.3 Hund’s rule The Hund’s rule states that for degenerate orbitals, the lowest energy is attained when the number of electrons having the same spin is maximized. This means that electrons occupy orbitals singly to the maximum extent possible and that these single electrons in a given subshell all have the same spin magnetic quantum number. Electrons arranged in this way are said to have parallel spins. e.g. Carbon atom; the two 2p electrons singly occupy two of the three 2p orbitals and they will be parallel to each other so that they can have the same spin. 25  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Table 1.3 Electron distribution of some elements in second and third period Element Total Orbital diagram Electron configuration electrons 1s 2s 2p 3s Li 3 1s2 2s1 Be 4 1s2 2s2 B 5 1s2 2s2 2p1 C 6 1s2 2s2 2p2 N 7 1s2 2s2 2p3 Ne 10 1s2 2s2 2p6 Na 11 1s2 2s2 2p6 3s1 1.4.4 Condensed electron configurations The electron configuration (also known as the electron distribution) of sodium, atomic number 11, is written as 1s22s22p63s1. However, the filling of the 2p subshell is complete at neon, which has a stable configuration with eight electrons (an octet) in the outermost occupied shell. The next element, sodium, marks the beginning of a new row of the periodic table. Sodium has a single 3s electron beyond the stable configuration of neon. The configuration of sodium can therefore be abbreviated as [Ne]3s1. The part represented by the bracketed symbol is the noble-gas core of the atom. More usually, these inner-shell electrons are referred to as the core electrons. The electrons given after the noble-gas core are called the outer-shell electrons or valence shell electrons. The outer-shell electrons include the electrons involved in chemical bonding, as they are called the valence electrons. Similarly, phosphorous, which has 15 electrons can be represented as 1s22s22p63s23p3 or [Ne]3s23p3. 26  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Example 1.5 (a) Write the electron configuration for silicon, element 14, in its ground state. (b) How many unpaired electrons does a ground-state silicon atom possess? Answer: (a) 1s22s22p63s23p2 or [Ne]3s23p2 (b) two unpaired electrons. After the noble-gas element argon (1s22s22p63s23p6), according to Aufbau principle, it is not the 3d but 4s orbital that the next electron goes to. Hence, the element that follows Argon, which is potassium (K) has the electronic configuration of 1s22s22p63s23p64s1 or [Ar]4s1. Calcium, with 20 electrons hence is written as 1s22s22p63s23p6 4s2 or [Ar]4s2. Following the complete filling of the 4s orbital (this occurs in the calcium atom), the next set of orbitals to be filled is the 3d. 4s 3d Mn: [Ar] 3d5 4s2 or [Ar] Zn: [Ar] 3d10 4s2 or [Ar] Once all the 3d orbitals have been filled with two electrons each, the 4p orbitals begin to be occupied until the completed octet of outer electrons (4s24p6) is reached with krypton (Kr), atomic number 36, another of the noble gases. Elements with completely filled or precisely half-filled sub energy levels appear to be relatively more stable than elements with other electron configurations. Thus elements with configurations ending with s2, p6 and d10 will be more stable. e.g.: Zn: [Ar]3d104s2, Mg: [Ne]3s2, Ar: [Ne]3s23p6, N: [He]2s22p3 and Mn: [Ar]3d54s2 will be relatively stable atoms. 27  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure The electron configurations of certain elements appear to deviate from the rules of configuration discussed. For an example, the electron configuration of chromium (element 24) is [Ar]3d54s1 rather than [Ar]3d44s2 which we might expect. Similarly, the configuration of copper (element 29) is [Ar]3d104s1 instead of [Ar]3d94s2.This anomalous behaviour is largely a consequence of the closeness of the 3d and 4s orbital energies. It frequently occurs when there are enough electrons to form precisely half-filled sub energy levels (as in chromium) or a completely filled sub energy levels (as in copper), that would result in relatively stable configurations. (Note that the filling of the 3d orbitals occurs after filling of 4s. However, the electronic configuration is commonly written as 3d first and then 4s). 1.5 Building of the periodic table The discovery of chemical elements has been ongoing since ancient times. Certain elements, such as gold (Au), appear in nature in elemental form and were thus discovered thousands of years ago. In contrast, some elements, such as technetium (Tc), are radioactive and intrinsically unstable and were discovered after the development of technology during the twentieth century. As the number of known elements increased, scientists began classifying them. In 1869, Dmitri Ivanovich Mendeleev in Russia and Lothar Meyer in Germany published nearly identical classification schemes. Both noted that similar chemical and physical properties occur periodically when the elements are arranged in order of increasing atomic mass. Scientists at that time had no knowledge of atomic numbers. However with the introduction of the concept of atomic number the modern periodic table was constructed. : (a) (b) Figure 1.29 (a) Dmitri Mendeleev and (b) Lothar Meyer - 28 G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Figure 1.30 The periodic table of elements 29  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure The way in which the columns (groups) are labeled is somewhat arbitrary. The labeling scheme widely used in the past had Arabic numerals and A and B designations. Thus the numbers ran from 1A-8A and 1B-8B. The group headed by fluorine (F) hence was group 7A. In a similar convention, the columns are numbered using Roman numerals rather than Arabic numerals together with the letters A and B. In an effort to eliminate this confusion, the International Union of Pure and Applied Chemistry (IUPAC) has proposed a convention that numbers the groups from 1 to18 with no A or B designations, as shown in the Figure 1.30. The electron configurations of the elements correspond to their locations in the periodic table. The rows of the table are called periods and elements of the same row show trends in some of their properties. Elements in the same column of the table, which are called groups, have related outer- shell (valence) electron configurations. For example, all group 2 elements have an ns2 outer configuration, and all group 3 elements have an ns2np1 outer configuration, with the value of ‘n’ increasing as we move down each column. Table 1.4 Electron configurations of Group 2 and 13 elements Group 2 Group 13 Be [He]2s2 B [He]2s2 2p1 Mg [Ne]3s2 Al [Ne]3s23p1 Ca [Ar]4s2 Ga [Ar]4s24p1 Sr [Kr]5s2 In [Kr]5s25p1 Ba [Xe]6s2 Tl [Xe]6s26p1 Ra [Rn]7s2 Elements in a group in the periodic table, often exhibit similarities in physical and chemical properties. Table 1.5 Names of some groups in the periodic table Group Name Elements 1 Alkali metals Li, Na, K, Rb, Cs, Fr 2 Alkaline earth metals Be, Mg, Ca, Sr, Ba, Ra 16 Chalcogens O, S, Se, Te, Po 17 Halogens F, Cl, Br, I, At 18 Noble gases (rare gases) Ne, Ar, Kr, Xe, Rn 30  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Since the total number of orbitals in each shell equals n2: 1, 4, 9 and 16 respectively and because two electrons can be placed in each orbital, each shell accommodates up to 2n2 electrons: 2, 8, 18, or 32. The overall structure of the periodic table too reflects these electron numbers: Each row of the table has 2, 8, 18 or 32 elements in it. The periodic table can be further divided into four blocks based on the filling order of orbitals. 1 1s 1s 2 2s 2p 3 3s 3p 4 4s 3d 4p 5 5s 4d 5p 6 6s 5d 4f 6p 7 7s 6d 5f 7p s-block f- block d- block p- block Figure1.31 Regions of the periodic table The two columns of elements on the left known as the alkali metals (group 1) and alkaline earth metals (group 2), are those in which the valence s orbitals are being filled. These two columns make up the s block of the periodic table. The block on the far right with six columns (group 13 to group 18) comprises the p block, where the valence p orbitals are being filled. The s block and the p block elements together are the representative elements, sometimes called the main-group elements. The block before the p-block in the Figure 1.31 has ten columns containing the transition metals. However in general practice group 1 and 10 elements are not considered as transition metals. These are the elements in which the valence d orbitals are being filled and make up the d block. The elements in the two rows between and s block and d block, containing 14 columns are the ones in which the valence f orbitals are being filled and make up the f block. (However filling of electrons and hence their electron configurations are complicated). These elements are often referred to as the f block metals or inner transition elements. The number of columns in each block corresponds to the maximum number of electrons that can occupy each kind of subshell. Since that 2, 6, 10, and 14 are the numbers of electrons that can fill the s, p, d and f subshells, respectively, the s block has 2 columns, the p block has 6, the d block has 10 and the f block has 14. 31  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure 1.6 Periodic trends shown by s and p block elements Many properties of atoms depend on their electron configuration and on how strongly the outer electrons in the atoms are attracted to the nucleus. Coulomb’s law tells us that the strength of the interaction between two electrical charges depends on the magnitudes of the charges and on the distance between them. Therefore, the attractive force between an electron and the nucleus depends on the magnitude of the nuclear charge and on the average distance between the nucleus and the electron. The force increases as the nuclear charge increases and decreases as the electron moves further from the nucleus. In an atom that containing a large number of electrons, in addition to the attraction of each electron to the nucleus, each electron experiences a repulsion due to the proximity of other electrons. These electron–electron repulsions cancel some of the attraction of the electron to the nucleus so that the electron experiences less attraction than it would if the other electrons were not there. The outer electrons are said to be screened or shielded from the nucleus by the inner electrons close to the nucleus and this phenomenon is hence termed the screening effect or shielding effect of electrons. An electron, therefore, experiences a net attraction by the nucleus that is less than it would be in the absence of other electrons. This partially screened nuclear charge is termed the effective nuclear charge, Zeff. The effective nuclear charge is always less than the actual nuclear charge (Zeff < Z). For a valence electron, most of the shielding is due to the core electrons, which are much closer to the nucleus. As a result, the greater the number of core electrons and the higher the number of core shells, the greater will be the screening effect. The effective nuclear charge increases from left to right across any period of the periodic table. Although the number of core electrons stays the same across the period, the number of protons increases. The valence electrons added to counterbalance the increasing nuclear charge thus it screens ineffectively. Thus, Zeff increases steadily across the period. 1.6.1 Sizes of atoms and ions Atoms are not hard, spherical objects as many of us think. According to the quantum mechanical model, atoms do not have sharply defined boundaries. We can define atomic size in several ways, based on the distances between atoms in various situations. van der Waals radius The van der Waals radius (nonbonding atomic radius) refers to one half the distance between two equivalent non-bonded atoms in their most stable arrangement, that is, where attractive forces are maximum. 32  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Covalent radius A chemical bond is the attractive interaction between any two adjacent atoms in a molecule. The two bonded atoms are closer together than they would be in a nonbonding collision. The bonding atomic radius for any atom in a molecule is equal to half of the bond distance (the distance between two bonded atoms). The bonding atomic radius (also known as the covalent radius) is smaller than the nonbonding atomic radius. I I I van der Waals Covalent distance distance = 430 pm = 266 pm I van der Waals Covalent radius radius = 215 pm = 133 pm Figure 1.32 Covalent and van der Waals radius for Iodine (I2) Metallic radius Metal atoms in a metallic structure are bonded to each other by metallic bonds. Half of the bond distance between the nuclei of two adjacent metal atoms in a metallic structure is called the metallic radius. Metallic radius Metallic diameter Figure 1.33 The metallic radius 33  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Periodic trends in atomic radii The atomic sizes within the periodic table show two interesting trends: Within each group, the atomic radius tends to increase from top to bottom. This trend results primarily from the increase in the principal quantum number (n) of the outer electrons. As we go down a column, the outer electrons have a greater probability of being further away from the nucleus, causing the atomic radius to increase. Within each period, the atomic radius generally tends to decrease from left to right. The major factor influencing this trend is the increase in effective nuclear charge across a period. The increasing effective nuclear charge steadily draws the valence electrons closer to the nucleus, causing the atomic radius to decrease. Figure 1.34 (a) Trends in atomic radii in the periodic table in pm 34  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Figure 1.34 (b) Trends in atomic radii in the periodic table Electron configurations of ions When electrons are removed from an atom to form a cation, they are always removed first from the occupied orbitals having the highest principal quantum number, n. For example, when one electron is removed from a sodium atom (1s22s22p63s1), it is the 3s1 electron that is removed. Na(1s22s22p63s1or [Ne]3s1) Na+(1s22s22p6 or [Ne]) + e If there is more than one occupied subshell for a given value of n, the electrons are first removed from the orbital with the highest value of l. For example, a boron atom loses its 2p electron before it loses its 2s electrons. B(1s22s22p1) B+(1s22s2) + e B+(1s22s2) B3+(1s2) + 2e When two electrons are removed from Fe ([Ar]3d64s2), the 4s2 electrons are the ones that are removed even though in filling the 4s fills before the 3d. Fe([Ar] 3d64s2) Fe2+([Ar]3d6) + 2e If an additional electron is removed, forming Fe3+, it comes from a 3d orbital because all the orbitals with n = 4 are empty. 35  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Fe2+([Ar]3d6) Fe3+([Ar]3d5) + e Electrons added to an atom to form an anion are added to the empty or partially filled orbital having the highest value of n, which is the valence shell. For example, an electron added to a fluorine atom to form the F- ion goes into the one remaining vacancy in the 2p subshell. F(1s22s22p5) + e F-(1s22s22p6) Periodic Trends in Ionic Radii Like the size of an atom, the size of an ion depends on its nuclear charge, the number of electrons it possesses, and the orbitals in which the valence electrons reside. When a cation is formed from a neutral atom, electrons are removed from the occupied atomic orbitals that are the most spatially extended from the nucleus. Also, when a cation is formed the number of electron–electron repulsions is reduced. Therefore, cations are smaller than their parent atoms. Cation Anion Parent atom Figure 1.35 Radii of cations and anions compared to parent atoms in pm The opposite is true of anions. When electrons are added to an atom to form an anion, the increased electron–electron repulsions cause the electrons to spread out more in space. Thus, anions are larger than their parent atoms. For ions carrying the same charge (both positively or negatively charged ions), ionic radius increases as we move down a column in the periodic table. In other words, as the principal quantum number of the outermost occupied orbital of an ion increases, the radius of the ion increases. An isoelectronic series is a group of ions/ atoms containing the same number of electrons. For example, each ion/ atom in the isoelectronic series O2-, F-, Ne, Na+ and Mg2+ has a total number of 10 electrons. In any isoelectronic series the nuclear charge increases with 36  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure increasing atomic number in the series. Because the number of electrons remains constant, ionic radius decreases with increasing nuclear charge as the electrons are more strongly attracted to the nucleus. Electron configurations of the neutral atoms: 2s22p4 2s22p5 2s22p6 2s22p63s1 2s22p63s2 8+ 9+ 10 + 11 + 12 + O2− F− Ne Na+ Mg2+ Isoelectronic series: all these species have ten electrons: 1s22s22p6 O2− Figure 1.36 Radii in an isoelectronic series 1.6.2 Ionization energy As explained at the beginning of section 1.3, the ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. In general, the first ionization energy (I1) is the energy needed to remove the most loosely bound electron from a neutral gaseous atom. For example, the first ionization energy for the lithium atom is the energy required for the process; Li(g) Li+(g) + e The second ionization energy (I2) is the energy needed to remove the second loosely bound electron of an atom from a gaseous monovalent cation to form a gaseous divalent cation, and so forth, for successive removals of additional electrons. Thus, second ionisation energy for the lithium atom is the energy associated with the process; Li+(g) Li2+(g) + e The ionization energies for a given element increase as successive electrons are removed; I1 < I2< I3, and so forth. This trend is because with each successive removal, an electron is being pulled away from an increasingly more positive ion, requiring increasingly more energy. Furthermore, a sharp increase in ionization energy occurs when an inner-shell electron is removed compared to the removal of outer shell electrons. This is because the electrons of the inner shell are closer to the nucleus and hence are attracted to it more strongly. 37  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Frequently ionization energy is expressed in kJ mol-1 by considering one mole of atom/ ion. Periodic trends in the first ionization energies The first ionization energy generally increases as we move across a period. The alkali metals show the lowest ionization energy in each period, and the noble gases show the highest. Generally, the first ionization energy decreases as we move down any column in the periodic table. For example, the ionization energies of the group 1 elements (alkali metals) follow the following order Li >Na > K > Rb > Cs > Fr. The s and p block elements show a larger range of I1 values than do the transition metal elements. Generally, the ionization energies of the transition metals increases slightly from left to right in a period. The same factors that influence atomic size also influence ionization energies. The energy needed to remove an electron from the outermost occupied shell depends on both the effective nuclear charge and the average distance of the electron from the nucleus. Either increasing the effective nuclear charge or decreasing the distance from the nucleus increases the attraction between the electron and the nucleus. As this attraction increases, it becomes more difficult to remove the electron, and thus the ionization energy increases. Ionization energy increases 1 2 13 14 15 16 17 18 Ionization energy increases H H He 1312.0 1312.0 2372.3 Li Be B C N O F Ne 520.2 899.4 800.6 1086.4 1420.3 1313.9 1681.0 2080.6 Na Mg Al Si P S Cl Ar 495.8 737.7 577.6 786.4 1011.7 999.6 1251.1 1520.5 K Ca Ga Ge As Se Br Kr 418.8 589.8 578.8 762.1 947 940.9 1139.9 1360.7 Rb Sr In Sn Sb Te I Xe 403.0 549.5 558.3 708.6 833.7 869.2 1008.4 1170.4 Cs Ba Tl Pb Bi Po At Rn 375.7 508.1 595.4 722.9 710.6 821 -- 1047.8 Fr Ra -- 514.6 Unit kJ mol-1 Figure 1.37 Trends in the first ionization energies in the periodic table 38  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure 1s2 [He]2s2 2p6 [Ne]3s23p6 [Ar]3d10 4s2 4p6 2500 He Ionization energy (kJ mol-1) Ne 2000 Ar 1500 Kr Xe 1000 500 Li Na Rb K Cs 0 10 20 30 40 50 60 Atomic number [He]2s1 [Ne]3s1 [Ar]4s1 [Kr]5s1 [Xe]6s1 Figure 1.38 Variation of the first ionization energies with atomic number of elements The irregularities in the trend of the first ionization energy in a given period are small but still readily explained. Removal of electrons from a completely filled subshell (e.g. group 2, group 12 and group 18) or a half filled subshell (e.g. group 7 and group 15), which are generally stable, will require more energy and thus the ionization energy will be higher than is expected from the common trend. For example, in the second period neon with a completely filled shell has the highest first ionization energy. Beryllium with a complete s sub shell shows the first ionization energy greater than expected and it even exceeds the I1 of boron. Similarly nitrogen with an exactly half filled p subshell shows an I1 higher than what the common trend predicts. 39  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure He Filled shell 2400 1s2 2s2 2p6 Filled shell Ne 2000 Half -filled Ionization energy kJ/ mol-1 subshell 2s22p3 F 1600 N Filled subshell 1200 H 2s2 C O Be 800 Mg B 400 Li 2s1 Na 3s1 1 2 3 4 5 6 7 8 9 10 11 12 Atomic number Figure 1.39 Variation of the first ionization energies along the first and second periods 1.6.3 Electron gain energy The energy change that occurs when an electron is added to a gaseous atom is called the electron gain energy. For most atoms, energy is released when an electron is added. For example, electron gain energy for the chlorine atom shown in the process below is -349 kJ mol-1. The negative sign indicates that energy is released during the process. Cl(g) + e Cl-(g) EEG = -349 kJ mol-1 (EEG = Electron gain energy) However, there are few atoms with positive electron gain energy. E.g.: Be, N. This is because they possess relatively stable configuration Be (s2) and N (p3) and as a result adding of an electron would be somewhat difficult with electron – electron repulsion is the dominating factor. N(g) + e N-(g) EEG = +134 kJ mol-1 Be(g) + e Be-(g) EEG = +231 kJ mol-1 The electron gain energy become less positive across a period and more positive down a group. 40  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure As accepted internationally, EEG is used as a quantitative physical quantity to measure the attraction of an electron by an atom and related to electron affinity as follows. Electron gain energy (EEG) = - Electron affinity (EA) Thus the electron affinity of an atom can be closely related to EEG and defined as the change in energy when the anion of the element loses an electron in the gaseous phase (A-(g) A(g) + e; E = EA which is equal in magnitude and opposite in sign to EEG) become more positive across a period and less positive down a group. 1.6.4 Electronegativity It is defined as the ability of an atom in a molecule to attract electrons to itself. The greater an atom’s electronegativity, the greater its ability to attract electrons. The American chemist Linus Pauling (1901–1994) developed the first and most widely used electronegativity scale, which is known as the Pauling scale. Generally there is an increase in electronegativity from left to right across a period in the periodic table. With some exceptions (especially in the transition metals), electronegativity decreases with increasing atomic number in a group. Noble gases too have very small but non zero electronegativity according to the Pauling scale. In molecules, the difference between electronegativities of two atoms forming a bond will determine the ionic or covalent nature of the bond. Figure 1.40 Pauling electronegativity values and the trends in the periodic table 41  G.C.E. (A/L) CHEMISTRY : Unit 1 Atomic Structure Table 1.6 A summary of equations Equation Atomic number (Z) = Number of protons = Number of electrons in a atom Mass number (A) = Number of protons (Z) + Number of neutrons 1 u or Da = 1.66054 x 10-24 g and 1 g = 6.02214 x 1023 u or Da Atomic mass = isotope mass) x (fractional isotope abundance)] Speed of light = c =  x 108 m s-1 Energy of a photon = E = h constant h is called Planck constant and has a value of 6.626 x 10-34 J s 42  G.C.E. (A/L) CHEMISTRY UNIT 2 Structure and Bonding 2. Structure and Bonding Content 2.1 Covalent bonds 2.3.4 Effect of electronegativity and 2.1.1 Lewis dot diagrams and Lewis dot geometry for the polarity of molecules dash structures 2.3.5 Dipole moment 2.2 Dative covalent bonds 2.3.6 Factors affecting the magnitude of electronegativity 2.3 Valance shell electron pair repulsion theory (VSEPR) 2.4 Ionic bonds  Linear electron pair geometry  Trigonal planar electron pair geometry 2.5 Metallic bonds  Tetrahedral electron pair geometry 2.6 Secondary interactions  Trigonal bipyramid electron pair  Ion – dipole interactions geometry  Dipole – dipole interactions  Octahedral electron pair geometry Hydrogen bonding  Ion – induced dipole interactions 2.3.1 Hybridization of atomic orbitals  Dipole – induced dipole interaction 2.3.2 Formation of double and triple bonds  London interactions (forces) (Instantaneous induced dipole – 2.3.3 Resonance structures induced dipole interaction)  Characteristics of resonance  Formal charges  Rules to estimate relative stability of resonance structures 43  G.C.E. (A/L) CHEMISTRY UNIT 2 Structure and Bonding Introduction Chemical bonds and structure of molecules are conceptual models based on the modern atomic model, in order to explain the physical and chemical properties of matter. Many atoms do not have stable outermost valance shell configurations, therefore chemical bonds occur between atoms in order to achieve stability. The following table (Figure 2.1) explains how valence electrons participate in different types of chemical bonding using several acceptable models. Chemical bonds Sharing a pair of Metallic bonds Complete removal of electrons between Large numbers of electrons from an atom to two atoms metallic cations are form cations and acceptance stabilized by a cloud of of electrons by another atom many electrons to form anions Covalent bonds Dative covalent bonds Ionic bonds/ Ionic Sharing a pair of Sharing of an electron interactions/ Ionic electrons between two pair between two atoms attraction forces atoms in which each in which both electrons Formed due to atom contributes one are given by one atom electrostatic attractive electron forces between cations and anions Figure 2.1 Types of chemical bonds 2.1 Covalent bonds Covalent bonds are formed when a pair of electrons is shared between two atoms of the same element or different elements. The sharing pair of electrons contribute one electron from each atom to form the electron pair. Consequently, stable electron configurations are often achieved by both atoms in respect to the total number of electrons i