CHM 1162 Part A & B Inorganic Chemistry PDF

Inorganic Chemistry Year 1: Organic and Analytical Chemistry Lecturer: Dr Consolee Sibosiko Lectures Part A & B Academic year 2023-2024  Part I: Inorganic Chemistry Knowledge and Understanding Having successfully completed the module, st...

Inorganic Chemistry Year 1: Organic and Analytical Chemistry Lecturer: Dr Consolee Sibosiko Lectures Part A & B Academic year 2023-2024  Part I: Inorganic Chemistry Knowledge and Understanding Having successfully completed the module, students should be able to demonstrate knowledge and understanding of: Part 1: Inorganic Chemistry 1. Chemical Bonding and properties; 2. Molecular Geometry: molecular shapes, VSERP model, and hybrides; 3. Atomic and molecular orbitals;  Part II: Transition Metals and Coordination Chemistry 1. Evolution of their physical and chemical properties: Hardness, melting point, Metal- metal effects, Chelating effect, Acid base theory, Oxidation states 2. d-block elements and their major compounds: Transition elements (3d, 4d and 5d), Types ligands: monodentante, bidentante, multidentante, Major compounds 3. Formation of Complexes and their nomenclature: Effective atomic number (Octet rule), Formation of complexes, Coordination Number, Naming of complexes 4. Structure and Isomerism of complexes : Stereoisomer (optical isomers and geometrical isomers), Structural isomerism (Coordination isomers, linkage isomers, ionisation isomers and hydrate isomers, coordinate isomers,.......) 5. Crystal field theory: General Introduction, Splitting diagram energy of Tetrahedral, Square Plan and Octahedral structure, Cristal Field Stabilisation Energy (strong and weak Field), Series Chemical Spectroscopy, Magnetic properties of complexes, Influence of ligands on the CFSE  Part III: Reactions of inorganic Complexes 1. Reactivity of complexes : Substitution reactions; Addition reactions; Elimination reactions; Redox reactions. 2. Rate and mechanism of coordination compounds: Labile and inert coordination compounds, Outer and inner sphere mechanism in octahedral, Outer and inner sphere mechanism in tetrahedral complexes, Outer and inner sphere mechanism in square Plan complexes, The kinetic effect 3. Applications of complex compound: Medicinal applications, Analysis applications, Catalyst applications, Environmental applications  General Introduction Aim This unit will allow you to be familiar with the basic concepts of inorganic chemistry. This unit will cover the following topics: The atomic models. The periodic table of elements. The evolution of properties in the periodic table. Learning outcomes Chemical bonding and properties; Molecular geometry: molecular shapes, VSERP model, and Hybrids; Atomic and molecular orbitals;  Essential ideas Knowledge of chemistry is central to understanding a wide range of scientific disciplines. This diagram shows just some of the interrelationships between chemistry and other fields.  General Introduction Definition: Inorganic chemistry is the study of the synthesis, reactions, structures and properties of compounds of the elements. It also studies the behavior of inorganic and organometallic compounds. This field covers all chemical compounds except the myriad organic compounds (carbon based compounds, usually containing C-H bonds). Inorganic Chemistry finds many applications such as: chemical industry, including catalysis, material science, pigments, coatings, medications, fuels, and agriculture.  General Introduction Many inorganic compounds are salts, consisting of cations and anions joined by ionic bonding. Examples: MgCl2, which consists of magnesium cations, Mg2+ and chloride anions Cl−; Na2O, which consists of sodium cations Na+ and oxide anions O2− In any salt, the proportions of the ions are such that the electric charges cancel out, so that the bulk compound is electrically neutral. Important classes of inorganic compounds: oxides, carbonates, sulfates and the halides. Many inorganic compounds are characterized by high melting points. Inorganic salts typically are poor conductors in the solid state. Another important feature is their solubility in water, and ease of crystallization. Where some salts (e.g. NaCl) are very soluble in water, others (e.g. SiO2) are not. The simplest inorganic reaction is double displacement when in mixing of two compounds the ions are swapped without a change in oxidation state. Example: HCl + AgNO3 AgCl + HNO3 In redox reactions one reactant, the oxidant, lowers its oxidation state and another reactant, the reductant, has its oxidation state increased. Example: Na(s) + HCl(aq) NaCl(aq) + H2(g) In a more general definition, an acid can be any chemical species capable of binding to electron pairs is called a Lewis acid; a Lewis base is any molecule that donates an electron pair. Inorganic compounds are found in nature as minerals. Inorganic compounds can be used as biomolecules: electrolytes (sodium chloride), energy storage (ATP) or in construction (the polyphosphate backbone in DNA). The first important man-made inorganic compound was ammonium nitrate for soil fertilization through the Haber process. Inorganic compounds are synthesized for use as catalysts such as vanadium(V) oxide and titanium(III) chloride, or as reagents in organic chemistry such as. lithium aluminium hydride Subdivisions of inorganic chemistry are organometallic chemistry, cluster chemistry and bioinorganic chemistry. Their classification focuses on: - the position in the periodic table of the heaviest element (the element with the highest atomic weight) in the compound. - grouping compounds by their structural similarities. Different classifications are: Coordination compounds Classical coordination compounds feature metals bond to "lone pairs" of the ligands such as H2O, NH3, Cl−, and CN−. In modern coordination compounds almost all organic and inorganic compounds can be used as ligands. One of the most important properties of metallic elements is their ability to act as Lewis acids that form complexes with a variety of Lewis bases. A metal complex consists of a central metal atom or ion that is bonded to one or more ligands (from the Latin ligare, meaning “to bind”), which are ions or molecules that contain one or more pairs of electrons that can be shared with the metal. Metal complexes can be neutral, such as [Co(NH3)3Cl3]; positively charged, such as [Nd(H2O)9]3+; or negatively charged, such as [UF8]4−. Electrically charged metal complexes are sometimes called complex ions. A coordination compound contains one or more metal complexes Coordination compounds are important for at least three reasons. First, most of the elements in the periodic table are metals, and almost all metals form complexes, so metal complexes are a feature of the chemistry of more than half the elements. Second, many industrial catalysts are metal complexes, and such catalysts are steadily becoming more important as a way to control reactivity. For example, a mixture of a titanium complex and an organometallic compound of aluminum is the catalyst used to produce most of the polyethylene and polypropylene “plastic” items we use every day. Finally, transition-metal complexes are essential in biochemistry. Examples include hemoglobin, an iron complex that transports oxygen in our blood; cytochromes, iron complexes that transfer electrons in our cells; and complexes of Fe, Zn, Cu, and Mo that are crucial components of certain enzymes, the catalysts for all biological reactions. Transition metal compounds Compounds containing metals from group 4 to 11 are considered transition metal compounds. Compounds with a metal from group 3 or 12 are sometimes also incorporated into this group, but also often classified as main group compounds. Transition metal compounds show a rich coordination chemistry, varying from tetrahedral for titanium (e.g. TiCl4) to square planar for some nickel complexes to octahedral for coordination complexes of cobalt. A range of transition metals can be found in biologically important compounds, such as iron in hemoglobin. Examples: Fe(CO)5 iron pentacarbonyl, TiCl4 titanium tetrachloride Organometallic compounds Organometallic compounds contain the M-C-H group. The metal (M) in these species can either be a main group element or a transition metal. Organometallic compounds are mainly considered a special category because organic ligands are often not sensitive to hydrolysis or oxidation. Organometallic compounds need more specialized preparative methods than was traditional in Werner-type complexes. Synthetic methodology, especially the ability to manipulate complexes in solvents of low coordinating power, enabled the exploration of very weakly coordinating ligands such as hydrocarbons, H2, and N2. Examples: Ferrocene Fe(C5H5)2, Molybdenum hexacarbonyl Mo(CO)6 Cluster compounds A cluster compound consists minimally of a triangular set of atoms that are directly bonded to each other. A metal-metal bonded dimetallic complexes are highly relevant to the area. Clusters occur in "pure" inorganic systems, organometallic chemistry, main group chemistry, and bioinorganic chemistry. Examples: Fe3(CO)12, B10H14, [Mo6Cl14]2−, B5H9, [Re(CO)3(Hbhp)]2 (with (μ-O)2Re2 moiety) Bioinorganic compounds These compounds occur in nature, but the subfield includes anthropogenic species, such as pollutants (e.g. methylmercury) and drugs (e.g. Cis platin). The field of biochemistry, includes many kinds of compounds, e.g. phosphates in DNA, metal complexes containing ligands that range from biological macromolecules, commonly peptides, to defined species such as humic acid, and to water (e.g. coordinated to gadolinium complexes employed for MRI). Bioinorganic chemistry focuses on electron- and energy- transfer in proteins relevant to respiration. Medicinal inorganic chemistry includes the study of both non- essential and essential elements with applications to diagnosis and therapies. Examples: hemoglobin, methylmercury H3C-Hg+X-, Solid state compounds This important area focuses on structure, bonding, and the physical properties of materials. In practice, solid state inorganic chemistry uses techniques such as crystallography to gain an understanding of the properties that result from collective interactions between the subunits of the solid. Solid state compounds are metals and their alloys or intermetallic derivatives. Related fields are condensed matter physics, mineralogy, and materials science. Examples: YBa2Cu3O7 (Barium copper Yttrium oxide).  Atom, element and compound Any atom is composed of a little nucleus surrounded by a "cloud" of electrons. In the nucleus there are protons and neutrons When an atom is defined by the number of protons contained in its nucleus, chemists refer to it as an element. All elements have a very specific identity that makes them unique from other elements. Compounds are composed of different type of atoms. More precisely, a compound is a chemical substance that consists of two or more elements.  Properties of matter/ Physical properties Intensive: A physical property that will be the same regardless of the amount of matter. – density: m/v – color: The pigment or shade – conductivity: Electricity to flow through the substance – malleability: if a substance can be flattened – luster: How shiny the substance looks Extensive: A physical property that will change if the amount of matter changes. – mass: How much matter in the sample – volume: How much space the sample takes up – length: How long the sample is  Physical changes Physical changes are change in which the matter's physical appearance is altered, but composition remains unchanged. (Change in state of matter). In a physical change, there is a change in the form of a substance, not in its chemical composition. The substances (atoms, molecules, or ions) present before or after a change are the same. In a physical change, there is a change in the form of a substance, not in its chemical composition. The substances (atoms, molecules, or ions) present before or after a change are the same. For e.g when water freezes or boils, it changes its state but remains water; it is still composed of H2O molecules; melting of a solid  Chemical properties and Chemical changes Any characteristic that gives a sample of matter the ability/inability to undergo a change that alters its composition. Chemical change is the change in which one or more kinds of matter are transformed to new kinds of matter with altered compositions. Is one in which one or more substances (the reactants) are transformed into one or more different substances (the products) with different properties and different composition. Eg: Magnesium + Oxygen --> Magnesium Oxide or 2 Mg + O2 --> 2 MgO  Atomic mass The modern system of atomic masses is based on carbon 12 as the standard. In this system, carbon twelve is assigned a mass of exactly 12 atomic mass units (amu), and the masses of other atoms are given relative to this standard. The average atomic mass is simply called atomic mass for an element. The relative atomic mass of a compound is the sum of the relative atomic mass of all atoms in the molecule of the considered compound. Ex: H2O: relative molar mass= 2*1+ 16 = 18grams Absolute atomic mass of carbon atom 12amu=12*1.66054*10 -24 g  Mole Because samples of matter typically contain many atoms, a unit of measure called the mole has been established to counting atoms. Mole is defined as the number equal to the number of carbon atoms in exactly 12 g of pure 12C. With the help of mass spectrometer, this number was found to be 6.02214 x 1023. This number is called Avogadro’s number to honor his contributions to chemistry. One mole of something consists of 6.02214 x 1023 units of that substance. Just as a dozen eggs is 12 eggs, a mole of eggs is 6.02214 x 1023 eggs.  Molar mass It is the mass in grams of one mole of the compound. The molar mass of a compound is obtained by summing the masses of the component atoms Molar mass is used to convert grams of a substance to moles and is used often in chemistry. The molar mass of an element is found on the periodic table and it is the element's atomic weight in grams/mole (g/mol). The atomic mass is the mass of one mole of that element.  If we know the mass of a substance, we can then determine how many moles are in the substance. Converting the mass, in grams, of a substance to moles requires a conversion factor of (one mole of substance/molar mass of substance). Using Avogadro's constant, it is also easy to calculate the number of atoms or molecules present in a substance. By multiplying the number of moles by Avogadro's constant, the mol units cancel out, leaving the number of atoms. Examples 1) Calculate the molar mass of potassium dichromate K2Cr2O7 2) How many atoms are in a 3.0 g sample of sodium (Na)? 3) Convert to moles and find the total number of atoms of a) 5.06 grams of Oxygen; b) 2.14 grams of K c) 0.134 Kg of Li 4) Convert the following to grams a) 4.5 mols of C; b) 7.1 mols of Al; c) 2.2 mols of Mg  Concentration Mass percent = (g of solute/g of solution)*100 Malarity (M) = (number of mole/L of solution) Molality (m) = (number of mole/Kg of solvent) Normality(N) = ( equivalent of solute/ L of solution) Mole fractional = number of mole of one component / total number of moles of all substances  Fundamental Chemistry Laws 1) Law of conservation of mass or Lavoisier’s law by Antoine Lavoisier (1743-1794), a French chemist. ‘Mass is neither created nor destroyed’. The law implies (requires) that during any chemical reaction, nuclear reaction, in an isolated system, the total mass of the reactants or starting materials must be equal to the mass of the products. 2) Law of definite proportion or Proust’s law by Joseph Proust (1754-1826), a Frenchman,. “a given compound always contains exactly the same proportion of elements by mass”.  3) Law of multiple proportions by John Dalton (1766-1844), an English schoolteacher. “When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1g of the first element can always be reduced to small whole numbers”. E.g. carbon forms two oxides by combining with oxygen in different proportions. A fixed mass of carbon, say 100 grams, may react with 133 grams of oxygen to produce one oxide, or with 266 grams of oxygen to produce the other. The ratio of the masses of oxygen that can react with 100 grams of carbon is 266:133 ≈ 2:1, a ratio of small whole numbers. The two oxides have one and two oxygen atoms respectively for each carbon atom. In modern notation the first is CO (carbon monoxide) and the second is CO2 (carbon dioxide).  4) Avogadro’s law by the Italian chemist in 1811“equal volumes of gases at the same temperature and pressure contain equal number of moles or molecules” V& n (T and P constant). V = a n (A = constant of proportionality). It follows that one mole of any gas at a given T and P has the same fixed volume called the molar gas volume. At standard temperature and pressure (STP), one mole of any gas occupies a volume of 22.4 litres. 5) Dalton’s law of partial pressures. In a mixture of gases, each component gas exerted a pressure as if it were alone in the container. The individual pressure of each gas in the mixture is defined as its partial pressure. By John Dalton in 1807. “the total pressure of a mixture of gases is equal to the sum of the partial pressures of all the gases present ” P = P + P + P + …(V and T are constant). P = n RT/V  Atomic Structure: Historical background Democritus was a Greek philosopher who was the first person to use the term atom (atomos: meaning indivisible). 400 BC He thought that if you take a piece of matter and divide it and continue to divide it you will eventually come to a point where you could not divide it any more. This fundamental or basic unit was what Democritus called an atom  Atomic Structure: Historical background He called this the theory of the universe: All matter consists of atoms, which are bits of matter too small to be seen. There is an empty space between atoms Atoms are completely solid Atoms have no internal structure Each atom (of a different substance) is different in size, weight and shape.  Scientist: John Dalton (1800’s) John Dalton was the first to adapt Democritus’ theory into the first modern atomic model. JOHN DALTON’S ATOMIC MODEL: 1. All matter consists of tiny particles called atoms 2.Atoms are indestructible and unchangeable 3.Elements are characterized by the weight of their atoms 4.When elements react, it is their atoms that have combined to form new compounds  Scientist: J.J Thomson (1890’s) J.J Thomson was a physicist who is credited for discovering the electron. He used his research on cathode ray tube technology in this discovery.  The charge is invisible, so to see where it traveled a fluorescent screen is placed at back of tube. Where the beam hits, a dot will appear on the screen. You could also use a fluorescent gas and the whole tube will light up. This beam will always travel straight if not interfered with. The deflection coils each have a specific charge. One is positive and the other is negative.  THOMSON’S ATOMIC MODEL Using what he had discovered, Thomson predicted what an atom should look like. These are the key points to Thomson’s Atomic Model: Because of its design this model is known as the plum pudding model Each atom is a sphere filled with positively charged ‘fluid’. This resembles the sticky jam part of a pudding.  Properties Fundamental of the particle Properties  Scientist: Ernest Rutherford (1910’s) Ernest Rutherford was not convinced about the model of the atom proposed by Thomson. He thus set up his now famous Gold Foil Experiment. – He fired alpha particles (positively charged) at a gold foil. – He measured the deflection as the particles came out the other side. – Most of the particles did not deflect at all. Every now and then a particle would deflect all the way back.  RUTHERFORD’S ATOMIC MODEL (THE PLANETARY MODEL) The nucleus of the atom is a dense mass of positively charged particles. The electrons orbit the nucleus  A problem raised was: Why are the negatively charged particles not attracted by the positively charged nucleus Rutherford stated that the atom was like a mini solar system and that the electrons orbited the nucleus in a wide orbit. That is why it is known as the planetary model.  Scientist: Niels Bohr (1910’s) Niels Bohr agreed with the planetary model of the atom, but also knew that it had a few flaws. Using his knowledge of energy and quantum physics he was able to perfect Rutherford’s model. He was able to answer why the electrons did not collapse into the nucleus. The energy difference can be calculated as the above Spectrum emission series  BOHR’S ATOMIC MODEL 1. Electrons orbit the nucleus in orbits that have a set size and energy. 2. The lower the energy of the electron, the lower the orbit. 3. This means that as electrons fill up the orbitals, they will fill the lower energy level first. 4. If that energy level is filled (or at capacity), a new energy level will begin. 5. Radiation is when an electron moves from one level to another. Problems with this theory: Electrons do not travel on a specific orbit or path.  Scientist: Erwin Schrödinger (1920’s) Erwin Schrödinger was a revolutionary physicist who used Heisenberg’s uncertainty principle to come up with the atomic model that we still use today. SCHRÖDINGER’S ATOMIC MODEL (THE CLOUD MODEL) 1. An electron does not travel in an exact orbit 2. We can predict where it will probably be 3. We cannot say for certain where it is, but only where it ought to be.  SUMMARY OF ATOM The smallest part of an element is called an atom Each atom (of an element) is different in structure from other atoms (of other elements) An atom can be divided in smaller subatomic particles: Protons, Electrons and Neutrons Atom 'structure Electromagnetic radiation Electromagnetic radiation Electromagnetic radiation  Quantum numbers The shape, size, and energy of each orbital is a function of 3 quantum numbers which describe the location of an electron within an atom or ion n (principal) ---> energy level l (orbital) ---> shape of orbital ml (magnetic) ---> designates a particular suborbital The fourth quantum number is not derived from the wave function s (spin) ---> spin of the electron (clockwise or counterclockwise: ½ or – ½)  Quantum numbers So… if two electrons are in the same place at the same time, they must be repelling, so at least the spin quantum number is different! The Pauli Exclusion Principle says that no two electrons within an atom (or ion) can have the same four quantum numbers. If two electrons are in the same energy level, the same sublevel, and the same orbital, they must repel. Think of the 4 quantum numbers as the address of an electron… Country > State > City > Street  Example Write n, l, ml and s quantum numbers for the 5 electrons of a boron atom Answer: 1,0,0,+1/2 1,0,0,-1/2 2,0,0,+1/2 2,0,0,-1/2 2,1,-1,+1/2 This example shows that no two electrons in one atom will have the same four quantum numbers. Quantum numbers  Types of orbitals The most probable area to find these electrons takes on a shape So far, we have 4 shapes. They are named s, p, d, and f. No more than 2 e- assigned to an orbital – one spins clockwise, one spins counterclockwise P orbitals d orbitals The shapes and labels of the five 3d orbitals f orbitals Electronic configuration Electronic configuration Energetic levels and sublevels of polyelectronic atoms  Electron configuration Orbitals and quantum numbers (Shells and sub shells) In atomic physics and quantum chemistry, the electron configuration is the distribution of electrons of an atom or molecule in atomic or molecular orbitals. Electron configuration was first conceived of under the Bohr model of the atom The electrons occupy a series of electron shells (numbered shell 1, shell 2 up to 7 or K, L up to Q). Each shell consists of one or more subshells (named s, p, d, f and g). Electrons fill the shells according to the Madelung rule or energy ordering rule. An electron shell is the set of allowed states electrons may occupy which share the same principal quantum number, n An atom's nth electron shell can accommodate 2n2 electrons, e.g. the first shell (n=1) can accommodate 2 electrons, the second shell 8 electrons, and the third shell 18 electrons. The factor of two arises because the allowed states are doubled due to electron spin-each atomic orbital admits up to two. electrons have opposite spin, one with a spin +1/2 (usually noted by an up-arrow) and one with a spin -1/2 (with a down-arrow). A subshell is the set of states defined by a common azimuthal quantum number, l, within a shell. The values l = 0, 1, 2, 3 correspond to the s, p, d, and f labels, respectively. The maximum number of electrons which can be placed in a subshell is given by 2(2l + 1). This gives two electrons in an s subshell, six electrons in a p subshell, ten electrons in a d subshell and fourteen electrons in an f subshell Example: The electronic configuration for neon is 1s2 2s2 2p6. Orbitals and the periodic table  Exceptions to the Aufbau Principle Remember d and f orbitals require large amounts of energy If we can’t fill these sublevels, then the next best thing is to be half full (one electron in each orbital in the sublevel) There are many exceptions, but the most common ones are d4 and d9 For the purposes of this class, we are going to assume that all atoms (or ions) that end in d4 or d9 are exceptions to the rule. This may or may not be true, it just depends on the atom.  Exceptions to the Aufbau Principle Ok, so this helps the d, but what about the poor s orbital that loses an electron? Remember, half full is good… and when an s loses 1, it too becomes half full! So… having the s half full and the d half full is usually lower in energy than having the s full and the d to have one empty orbital.  Exceptions to the Aufbau Principle d9 is one electron short of being full Just like d4, one of the closest s electrons will go into the d, this time making it d10 instead of d9. For example: Au would be [Xe] 6s2 4f14 5d9, but since this ends exactly with a d9 it is an exception to the rule. Thus, Au should be [Xe] 6s1 4f14 5d10. Procedure: Same as before! Find the closest s orbital. Steal one electron from it, and add it to the d.  Exceptions to the Aufbau Principle d4 is one electron short of being half full In order to become more stable (require less energy), one of the closest s electrons will actually go into the d, making it d5 instead of d4. For example: Cr would be [Ar] 4s2 3d4, but since this ends exactly with a d4 it is an exception to the rule. Thus, Cr should be [Ar] 4s1 3d5. Procedure: Find the closest s orbital. Steal one electron from it, and add it to the d. Lanthanides elements  EC of an atom in AO Eg:Using complete s, p, d, f subshell notation, predict the electron configurations of the following atoms or ions. (Note: The atomic number (Z) for S=16, Gd=64, Cs=55, and Cr=24) i) S- ii) Gd3+ iii) Cs iv) Cr Answer: i) [Ne] 3s²3p5 or 1s22s22p63s23p5 ii) [Xe] 4f7 or [Xe] 4f7 5d06s0 or 1s22s22p63s23p63d104s24p64d105s25p64f7 iii) [Xe] 6s1 or 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s25p6 6s1 iv) [Ar] 3d⁵4s¹ or 1s22s22p63s23p63d54s1 Periodic table  Periodic table variation Describe and explain the observed trends in atomic size, ionization energy, and electron affinity of the elements: The elements in groups (vertical columns) of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number and distribution of electrons in their valence shells. However, there are also other patterns in chemical properties on the periodic table. For example, as we move down a group, the metallic character of the atoms increases. Oxygen, at the top of group 16 (6A), is a colorless gas; in the middle of the group, selenium is a semiconducting solid; and, toward the bottom, polonium is a silver-grey solid that conducts electricity. As we go across a period from left to right, we add a proton to the nucleus and an electron to the valence shell with each successive element. As we go down the elements in a group, the number of electrons in the valence shell remains constant, but the principal quantum number increases by one each time. An understanding of the electronic structure of the elements allows us to examine some of the properties that govern their chemical behavior.  Periodic table variation (a) The radius of an atom is defined as one-half the distance between the nuclei in a molecule consisting of two identical atoms joined by a covalent bond. The atomic radius for the halogens increases down the group as n increases. (b) Covalent radii of the elements are shown to scale. The general trend is that radii increase down a group and decrease across a period Within each period, the trend in atomic radius decreases as Z increases; for example, from K to Kr. Within each group (e.g., the alkali metals shown in purple), the trend is that atomic radius increases as Z increases.  Periodic table variation As we move across a period from left to right, we generally find that each element has a smaller covalent radius than the element preceding it. This might seem counterintuitive because it implies that atoms with more electrons have a smaller atomic radius. This can be explained with the concept of effective nuclear charge, Zeff. This is the pull exerted on a specific electron by the nucleus, taking into account any electron–electron repulsions. For hydrogen, there is only one electron and so the nuclear charge (Z) and the effective nuclear charge (Zeff) are equal.  Periodic table variation For all other atoms, the inner electrons partially shield the outer electrons from the pull of the nucleus and thus: Zeff = Z − shielding Shielding is determined by the probability of another electron being between the electron of interest and the nucleus, as well as by the electron repulsions the electron of interest encounters.  Periodic table variation Core electrons are adept at shielding, while electrons in the same valence shell do not block the nuclear attraction experienced by each other as efficiently. Thus, each time we move from one element to the next across a period, Z increases by one, but the shielding increases only slightly. Thus, Zeff increases as we move from left to right across a period. The stronger pull (higher effective nuclear charge) experienced by electrons on the right side of the periodic table draws them closer to the nucleus, making the covalent radii smaller. Thus, as we would expect, the outermost or valence electrons are easiest to remove because they have the highest energies, are shielded more, and are farthest from the nucleus. As a general rule, when the representative elements form cations, they do so by the loss of the ns or np electrons that were added last in the Aufbau process. The transition elements, on the other hand, lose the ns electrons before they begin to lose the (n – 1)d electrons, even though the ns electrons are added first, according to the Aufbau principle.  Periodic table variation Predict the order of increasing covalent radius for Ge, Fl, Br, Kr. Solution Radius increases as we move down a group, so Ge < Fl (Note: Fl is the symbol for flerovium, element 114, NOT fluorine). Radius decreases as we move across a period, so Kr < Br < Ge. Putting the trends together, we obtain Kr < Br < Ge < Fl. Effective nuclear charge Z* The presence of other electrons around a nucleus “screens” an electron from the full charge of the nucleus. We can approximate the energy of the electrons by modifying the Bohr equation to account for the lower “effective” nuclear charge  Z *2  En = − R  2  n   Prediction  and Z* Slater’s rules for the prediction of  for an electron: 1. Group electron configuration as follows: (1s)(2s,2p)(3s,3p)(3d)(4s,4p)(4d)(4f)(5s,5p) etc. 2. Electrons to the right (in higher subshells and shells) of an electron do not shield it. 3. For ns or np valence electrons: a) Each other electron in the same group contributes 0.35 (0.30 for 1s) b) Each electron in an n-1 group contributes 0.85 c) Each electron in an n-2 or lower group contributes 1.00 4. For nd or nf valence electrons: a) Each other electron in the same group contributes 0.35 b) Each electron in a lower group (to the left) contributes 1.00  Prediction  and Z* Example 1 with a valence electron on oxygen: O, Z = 8 Electron configuration: 1s2 2s2 2p4 a) (1s2) (2s2 2p4) a)  = (2 * 0.85) + (5 * 0.35) = 3.45 1s 2s,2p Z* = Z -  Z* = 8 – 3.45 = 4.55 This electron is actually held with about 57% of the force that one would expect for a +8 nucleus.  Example/Practice Q1) Prediction  and Z* For N, O and F and compare their size Q2) Find Zeff for a 4s, 3d, 3p, and 3s electron in Fe Q3) Use Slater’s rule to estimate the effective nuclear charge of 2s electrons in Be Q4) Estimate the effective nuclear charge of a 1s electron in Lithium, Magnesium and Potassium  Prediction  and Z* Example 2 with two electrons for nickel: Ni, Z = 28 Electron configuration: 1s2 2s2 2p6 3s2 3p6 3d8 4s2 (1s2) (2s2 2p6) (3s2 3p6) (3d8) (4s2) For a 3d electron:  = (18 * 1.00) + (7 * 0.35) = 20.45 1s,2s,2p,3s,3p 3d Z* = Z -  Z* = 28 – 20.45 = 7.55 For a 4s electron:  = (10 * 1.00) + (16 * 0.85) + (1 * 0.35) = 23.95 1s,2s,2p 3s,3p,3d 4s Z* = Z -  Z* = 28 – 23.95 = 4.05  Variation in ionization energy The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy (IE1). The first ionization energy for an element, X, is the energy required to form a cation with +1 charge: X(g) ⟶ X+(g) + e− IE1 The energy required to remove the second most loosely bound electron is called the second ionization energy (IE2). X+(g) ⟶ X2+ (g) + e− IE2 The energy required to remove the third electron is the third ionization energy, and so on. Energy is always required to remove electrons from atoms or ions, so ionization processes are endothermic and IE values are always positive. For larger atoms, the most loosely bound electron is located farther from the nucleus and so is easier to remove. Thus, as size (atomic radius) increases, the ionization energy should decrease. Relating this logic to what we have just learned about radii, we would expect first ionization energies to decrease down a group and to increase across a period. The first ionization energy of the elements in the first five periods are plotted against their atomic number.  Variation in electron affinity The electron affinity (EA) is the energy change for the process of adding an electron to a gaseous atom to form an anion (negative ion). X(g) + e− ⟶ X−(g) EA1 As we might predict, it becomes easier to add an electron across a series of atoms as the effective nuclear charge of the atoms increases.  Chemical Bonding and Properties Molecules (and extended solids) are built from atoms that form chemical bonds. Theories of bonding seek to explain why molecules and solids form, what their structures are, why some are more stable than others, and how they react. Models such as Lewis dot structures and valence shell electron-pair repulsion (VSEPR) theory. When combined with a qualitative quantum mechanical description of bonding through the concepts of orbital hybridization and resonance, these simple models can help us understand a great deal about the structures, stabilities, and reactions of inorganic molecules.  Chemical Bonding and Properties A chemical bond is a lasting attraction between atoms that enables the formation of chemical compounds. The bond may result from the electrostatic force of attraction between atoms with opposite charges, or through the sharing of electrons as in the covalent bonds. Example. The element sodium is a silver-colored metal The element chlorine is a greenish-colored gas that is so poisonous that it was used as a weapon in World War I. When chemically bonded together, these two dangerous substances form the compound sodium chloride (NaCl), a compound so safe that we eat it every day - common table salt.  Assignment “There is no topic more fundamental to Chemistry than the nature of the chemical bond” Write an essay project of such statement  Chemical Bonding and Properties Ionic Bonding Covalent Bonding Lewis Symbols and Structures Formal Charges and Resonance Strengths of Ionic and Covalent Bonds Molecular Structure and Polarity  The Formation of Ionic Compounds Binary ionic compounds are composed of just two elements: A metal (which forms the cations) and a nonmetal (which forms the anions) For example, NaCl is a binary ionic compound. We can think about the formation of such compounds in terms of the periodic properties of the elements.  The Formation of Ionic Compounds Many metallic elements have relatively low ionization potentials and lose electrons easily. These elements lie to the left in a period or near the bottom of a group on the periodic table. Nonmetal atoms have relatively high electron affinities and thus readily gain electrons lost by metal atoms, thereby filling their valence shells. Nonmetallic elements are found in the upper-right corner of the periodic table.  The Formation of Ionic Compounds As all substances must be electrically neutral, the total number of positive charges on the cations of an ionic compound must equal the total number of negative charges on its anions The formula of an ionic compound represents the simplest ratio of the numbers of ions necessary to give identical numbers of positive and negative charges. For example, the formula for aluminum oxide, Al2O3, indicates that this ionic compound contains two aluminum cations, Al3+, for every three oxide anions, O2− [thus, (2 × +3) + (3 × –2) = 0].  The Formation of Ionic Compounds The atoms in sodium chloride (common table salt) are arranged to (a) maximize opposite charges interacting. The smaller spheres represent sodium ions, the larger ones represent chloride ions. In the expanded view (b), the geometry can be seen more clearly. Note that each ion is “bonded” to all of the surrounding ions six in this case.  The Formation of Ionic Compounds The strong electrostatic attraction between Na+ and Cl– ions holds them tightly together in solid NaCl. It requires 769 kJ of energy to dissociate one mole of solid NaCl into separate gaseous Na+ and Cl– ions: NaCl(s) ⟶ Na+(g) + Cl-(g) ΔH = 769 kJ  Electronic Structures of Cations When forming a cation, an atom of a main group element tends to lose all of its valence electrons, thus assuming the electronic structure of the noble gas that precedes it in the periodic table. For example, calcium is a group 2 element whose neutral atoms have 20 electrons and a ground state electron configuration of 1s22s22p63s23p64s2. When a Ca atom loses both of its valence electrons, the result is a cation with 18 electrons, a 2+ charge, and an electron configuration of 1s22s22p63s23p6. The Ca2+ ion is therefore isoelectronic with the noble gas Ar.  Electronic Structures of Cations Exceptions to the expected behavior involve elements toward the bottom of the groups. In addition to the expected ions Tl3+, Sn4+, Pb4+, and Bi5+, a partial loss of these atoms’ valence shell electrons can also lead to the formation of Tl+, Sn2+, Pb2+, and Bi3+ ions. The formation of these 1+, 2+, and 3+ cations is ascribed to the inert pair effect, which reflects the relatively low energy of the valence s-electron pair for atoms of the heavy elements of groups 13, 14, and 15.  Electronic Structures of Cations Transition and inner transition metal elements behave differently than main group elements. Most transition metal cations have 2+ or 3+ charges that result from the loss of their outermost s electron(s) first, sometimes followed by the loss of one or two d electrons from the next to outermost shell.  Electronic Structures of Cations For example, iron (1s22s22p63s23p63d64s2) forms the ion Fe2+ (1s22s22p63s23p63d6) by the loss of the 4s electron and the ion Fe3+ (1s22s22p63s23p63d5) by the loss of the 4s electron and one of the 3d electrons. Although the d orbitals of the transition elements are according to the Aufbau principle the last to fill when building up electron configurations, the outermost s electrons are the first to be lost when these atoms ionize. When the inner transition metals form ions, they usually have a 3+ charge, resulting from the loss of their outermost s electrons and a d or f electron.  Practice There are at least 14 elements categorized as “essential trace elements” for the human body. They are called “essential” because they are required for healthy bodily functions, “trace” because they are required only in small amounts, and “elements” in spite of the fact that they are really ions. Two of these essential trace elements, chromium and zinc, are required as Cr3+ and Zn2+. Write the electron configurations of these cations. Answer: Zn2+: [Ar]3d10 Cr3+: [Ar]3d3  Practice Selenium and iodine are two essential trace elements that form anions. Write the electron configurations of the anions. Solution Se2–: [Ar]3d104s24p6 I –: [Kr]4d105s25p6  Ionic Bonding Atoms gain or lose electrons to form ions with particularly stable electron configurations. The charges of cations formed by the representative metals may be determined readily because, with few exceptions, the electronic structures of these ions have either a noble gas configuration or a completely filled electron shell. The charges of anions formed by the nonmetals may also be readily determined because these ions form when nonmetal atoms gain enough electrons to fill their valence shells.  Covalent Bonding By the end of this section, you will be able to: Describe the formation of covalent bonds Define electronegativity and assess the polarity of covalent bonds Ionic bonding results from the electrostatic attraction of oppositely charged ions that are typically produced by the transfer of electrons between metallic and nonmetallic atoms. A different type of bonding results from the mutual attraction of atoms for a “shared” pair of electrons. Such bonds are called covalent bonds.  Covalent bonds Covalent bonds form when electrons are shared between atoms and are attracted by the nuclei of both atoms. In pure covalent bonds, the electrons are shared equally. In polar covalent bonds, the electrons are shared unequally, as one atom exerts a stronger force of attraction on the electrons than the other. The ability of an atom to attract a pair of electrons in a chemical bond is called its electronegativity.  Covalent bonds Covalent bonds are formed between two atoms when both have similar tendencies to attract electrons to themselves (i.e., when both atoms have identical or fairly similar ionization energies and electron affinities). For example, two hydrogen atoms bond covalently to form an H2 molecule; each hydrogen atom in the H2 molecule has two electrons stabilizing it, giving each atom the same number of valence electrons as the noble gas He.  Covalent bonds The difference in electronegativity between two atoms determines how polar a bond will be. In a diatomic molecule with two identical atoms, there is no difference in electronegativity, so the bond is nonpolar or pure covalent. When the electronegativity difference is very large, as is the case between metals and nonmetals, the bonding is characterized as ionic.  Electronegativity Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity. Electronegativity is a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity.  Electronegativity Electrons in a polar covalent bond are shifted toward the more electronegative atom; thus, the more electronegative atom is the one with the partial negative charge. The greater the difference in electronegativity, the more polarized the electron distribution and the larger the partial charges of the atoms.  Electronegativity versus Electron Affinity We must be careful not to confuse electronegativity and electron affinity. The electron affinity of an element is a measurable physical quantity, namely, the energy released or absorbed when an isolated gas-phase atom acquires an electron, measured in kJ/mol. Electronegativity, on the other hand describes how tightly an atom attracts electrons in a bond.  Quiz n02, 18th July 2023 Q1) Indicate whether the bonds between the following would be pure covalent, polar covalent or ionic: a) )O-H, b) Cs-Cl, c) H-Cl, d) Br-Br Q2) Place the following in order from lowest to the highest electronegativity: F, Nb, N, Si, Rb, Ca, Pt Electronegativity and Bond Type  Lewis Symbols and Structures Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each atom in a Lewis structure. Most structures, especially those containing second row elements obey the octet rule, in which every atom (except H) is surrounded by eight electrons. Exceptions to the octet rule occur for odd-electron molecules (free radicals), electron- deficient molecules, and hypervalent molecules.  Lewis Symbols We use Lewis symbols to describe valence electron configurations of atoms and monatomic ions. A Lewis symbol consists of an elemental symbol surrounded by one dot for each of its valence electrons: Lewis symbols can also be used to illustrate the formation of cations from atoms, as shown here for sodium and calcium:  Lewis Symbols Cations are formed when atoms lose electrons, represented by fewer Lewis dots, whereas anions are formed by atoms gaining electrons. The total number of electrons does not change.  Writing Lewis Structures with the Octet Rule For very simple molecules and molecular ions, we can write the Lewis structures by merely pairing up the unpaired electrons on the constituent atoms. See these examples:  Writing Lewis Structures with the Octet Rule For more complicated molecules and molecular ions, it is helpful to follow the step- by-step procedure outlined here: 1. Determine the total number of valence (outer shell) electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.  Writing Lewis Structures with the Octet Rule 2) Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (Generally, the least electronegative element should be placed in the center.) Connect each atom to the central atom with a single bond (one electron pair). 3) Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom.  Writing Lewis Structures with the Octet Rule 4) Place all remaining electrons on the central atom. 5) Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.  Writing Lewis Structures with the Octet Rule Let us determine the Lewis structures of SiH4, CHO2−, NO+, and OF2 as examples in following this procedure: 1. Determine the total number of valence (outer shell) electrons in the molecule or ion. For a molecule, we add the number of valence electrons on each atom in the molecule:  Writing Lewis Structures with the Octet Rule For a negative ion, such as CHO2−, we add the number of valence electrons on the atoms to the number of negative charges on the ion (one electron is gained for each single negative charge):  Writing Lewis Structures with the Octet Rule For a positive ion, such as NO+, we add the number of valence electrons on the atoms in the ion and then subtract the number of positive charges on the ion (one electron is lost for each single positive charge) from the total number of valence electrons:  Writing Lewis Structures with the Octet Rule Since OF2 is a neutral molecule, we simply add the number of valence electrons:  Writing Lewis Structures with the Octet Rule Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom and connecting each atom to the central atom with a single (one electron pair) bond. (Note that we denote ions with brackets around the structure, indicating the charge outside the brackets:)  Writing Lewis Structures with the Octet Rule Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen) to complete their valence shells with an octet of electrons. There are no remaining electrons on SiH4, so it is unchanged:  Practice NASA’s Cassini-Huygens mission detected a large cloud of toxic hydrogen cyanide (HCN) on Titan, one of Saturn’s moons. Titan also contains ethane (H3CCH3), acetylene (HCCH), and ammonia (NH3). What are the Lewis structures of these molecules?  Writing Lewis Structures: Octet Rule Violations Xenon is a noble gas, but it forms a number of stable compounds. We examined XeF4 earlier. What are the Lewis structures of XeF2 and XeF6? Answers a) b)  Formal Charges and Resonance In a Lewis structure, formal charges can be assigned to each atom by treating each bond as if one-half of the electrons are assigned to each atom. These hypothetical formal charges are a guide to determining the most appropriate Lewis structure. A structure in which the formal charges are as close to zero as possible is preferred. Resonance occurs in cases where two or more Lewis structures with identical arrangements of atoms but different distributions of electrons can be written. The actual distribution of electrons (the resonance hybrid) is an average of the distribution indicated by the individual Lewis structures (the resonance forms).  Formal Charges and Resonance Examples Eg 1) Eg 2)  Strengths of Ionic and Covalent Bonds The strength of a covalent bond is measured by its bond dissociation energy, that is, the amount of energy required to break that particular bond in a mole of molecules. Multiple bonds are stronger than single bonds between the same atoms. The enthalpy of a reaction can be estimated based on the energy input required to break bonds and the energy released when new bonds are formed. For ionic bonds, the lattice energy is the energy required to separate one mole of a compound into its gas phase ions. Lattice energy increases for ions with higher charges and shorter distances between ions. Lattice energies are often calculated using the Born-Haber cycle, a thermochemical cycle including all of the energetic steps involved in converting elements into an ionic compound.  Molecular Structure and Polarity VSEPR theory predicts the three-dimensional arrangement of atoms in a molecule. It states that valence electrons will assume an electron-pair geometry that minimizes repulsions between areas of high electron density (bonds and/ or lone pairs). Molecular structure, which refers only to the placement of atoms in a molecule and not the electrons, is equivalent to electron-pair geometry only when there are no lone electron pairs around the central atom.  Molecular Structure and Polarity A dipole moment measures a separation of charge. For one bond, the bond dipole moment is determined by the difference in electronegativity between the two atoms. For a molecule, the overall dipole moment is determined by both the individual bond moments and how these dipoles are arranged in the molecular structure. Polar molecules (those with an appreciable dipole moment) interact with electric fields, whereas nonpolar molecules do not.  VSEPR Theory Valence shell electron-pair repulsion theory (VSEPR theory) enables us to predict the molecular structure, including approximate bond angles around a central atom, of a molecule from an examination of the number of bonds and lone electron pairs in its Lewis structure. VSEPR theory predicts the arrangement of electron pairs around each central atom and, usually, the correct arrangement of atoms in a molecule.  VSEPR Theory As a simple example of VSEPR theory, let us predict the structure of a gaseous BeF2 molecule. Two regions of electron density around a central atom in a molecule form a linear geometry; three regions form a trigonal planar geometry; four regions form a tetrahedral geometry; five regions form a trigonal bipyramidal geometry; and six regions form an octahedral geometry.  VSEPR Theory The basic electron-pair geometries predicted by VSEPR theory maximize the space around any region of electron density (bonds or lone pairs).  VSEPR Theory The molecular structures are identical to the electron-pair geometries when there are no lone pairs present (first column). For a particular number of electron pairs (row), the molecular structures for one or more lone pairs are determined based on modifications of the corresponding electron-pair geometry.  Practice Predicting Electron-pair Geometry and Molecular Structure: CO2, BCl3, NH4+ H2O  Hydrogen bonding and Metallic Bonding Metallic bonding: Metallic bonding is the electromagnetic interaction between delocalized electrons, called conduction electrons and gathered in an "electron sea", and the metallic nuclei within metals. Understood as the sharing of "free" electrons among a lattice of positively charged ions (cations). Electrons move freely within the molecular orbitals, and so each electron becomes detached from its parent atom (electrons are delocalised.  The metal is held together by the strong forces of attraction between the positive nuclei and the delocalized electrons. Not all metals exhibit metallic bonding: one such example is the mercurous ion (Hg2+2), which forms covalent metal-metal bonds. This is sometimes described as "an array of positive ions in a sea of electrons".  A hydrogen bond is the attractive interaction of a hydrogen atom with an electronegative atom, such as nitrogen, oxygen or fluorine. The hydrogen must be covalently bonded to another electronegative atom to create the bond. These bonds can occur between molecules (intermolecularly), or within different parts of a single molecule (intramolecularly).  Molecular geometry Molecular geometry or molecular structure is the three- dimensional arrangement of the atoms that constitute a molecule. It determines several properties of a substance including its reactivity, polarity, phase of matter, color, magnetism, and biological activity. The molecular geometry can be determined by various spectroscopic methods and diffraction methods. IR, Microwave and Raman spectroscopy can give information about the molecule geometry from the details of the vibrational and rotational absorbances detected by these techniques. X-ray crystallography, neutron diffraction and electron diffraction can give molecular structure for crystalline solids based on the distance between nuclei and concentration of Gas electron diffraction can be used for small molecules in the gas phase. NMR and FRET methods can be used to determine complementary information including relative distances, dihedral angles, angles, and connectivity. Molecular geometries are best determined at low temperature because at higher temperatures the molecular structure is averaged over more accessible geometries. Larger molecules often exist in multiple stable geometries (conformational isomerism) that are close in energy on the potential energy surface. Geometries can also be computed by quantum chemistry methods to high accuracy. The molecular geometry can be different as a solid, in solution, and as a gas. The position of each atom is determined by the nature of the chemical bonds by which it is connected to its neighboring atoms. The molecular geometry can be described by the positions of atoms in space, evoking bond lengths of two joined atoms, bond angles of three connected atoms, and torsion angles (dihedral angles) of three consecutive bonds. Molecular Shapes There are six basic shape types for molecules Linear: In a linear model, atoms are connected in a straight line. The bond angles are set at 180°. A bond angle is very simply the geometric angle between two adjacent bonds. For example, carbon dioxide has a linear molecular shape. Trigonal planar: the trigonal planar shape are somewhat triangular and in one plane (meaning a flat surface). Consequently, the bond angles are set at 120°. An example of this is boron trifluoride (BF3). Tetrahedral: Tetra- signifies four, and -hedral relates to a surface, so tetrahedral almost literally means "four surfaces." This is when there are four bonds all on one central atom, with no extra unshared electron pairs. In accordance with the VSEPR (valence-shell electron pair repulsion theory), the bond angles between the electron bonds are across = 109.5°. An example of a tetrahedral molecule is methane (CH4). Octahedral: Octa- signifies eight, and -hedral relates to a surface, so octahedral almost literally means "eight surfaces." The bond angle is 90 degrees. An example of an octahedral molecule is sulfur hexafluoride (SF6). Pyramidal: Pyramidal-shaped molecules have pyramid-like shapes. Unlike the linear and trigonal planar shapes but similar to the tetrahedral orientation, pyramidal shapes requires three dimensions in order to fully separate the electrons. Here, there are only three pairs of bonded electrons, leaving one unshared lone pair. Lone pair - bond pair repulsions change the angle from the tetrahedral angle to a slightly lower value. An example is NH3 (ammonia). Bent: The final basic shape of a molecule is the non-linear shape, also known as bent or angular. One of the most unquestionably important molecules any chemist studies is water, or H2O. A water molecule has a non-linear shape because it has two pairs of bonded electrons and two unshared lone pairs. Like in the other arrangements, electrons must be spaced as far as possible. Lone pair - bond pair repulsions push the angle from the tetrahedral angle down to around 106°. Valence shell electron pair repulsion (VSEPR) theory is a model in chemistry used to predict the shape of individual molecules based upon the extent of electron-pair electrostatic repulsion. It is also named Gillespie-Nyholm theory after its two main developers. The acronym "VSEPR" is sometimes pronounced "vesper" for ease of pronunciation. The premise of VSEPR is that the valence electron pairs surrounding an atom mutually repel each other, and will therefore adopt an arrangement that minimizes this repulsion, thus determining the molecular geometry. The number of electron pairs surrounding an atom, both bonding and nonbonding, is called its steric number. The "AXE method" of electron counting is commonly used when applying the VSEPR theory. The A represents the central atom and always has an implied subscript one. The X represents the number of sigma bonds between the central atoms and outside atoms. Multiple covalent bonds (double, triple, etc) count as one X. The E represents the number of lone electron pairs surrounding the central atom. The sum of X and E, known as the steric number, is also associated with the total number of hybridized orbitals used by valence bond theory. Based on the steric number and distribution of X's and E's, VSEPR theory makes the predictions of the geometry of the molecules. Note that the geometries are named according to the atomic positions only and not the electron arrangement. example the description of AX2E1 as bent means that AX2 is a bent molecule without reference to the lone pair, although the lone pair helps to determine the geometry. Exceptions There are groups of compounds where VSEPR fails to predict the correct geometry. Many transition metal compounds do not have geometries explained by VSEPR which can be ascribed to there being no lone pairs in the valence shell and the interaction of core d electrons with the ligands. The structure of some of these compounds, including metal hydrides and alkyl complexes such as hexamethyltungsten, can be predicted correctly using the VALBOND theory (VBT), which is based on sd hybrid orbitals and the 3-center-4-electron bonding model. Crystal field theory is another theory that can often predict the geometry of coordination complexes. Exceptions There are groups of compounds where VSEPR fails to predict the correct geometry. Many transition metal compounds do not have geometries explained by VSEPR which can be ascribed to there being no lone pairs in the valence shell and the interaction of core d electrons with the ligands. The structure of some of these compounds, including metal hydrides and alkyl complexes such as hexamethyltungsten, can be predicted correctly using the VALBOND theory (VBT), which is based on sd hybrid orbitals and the 3-center-4-electron bonding model. Crystal field theory is another theory that can often predict the geometry of coordination complexes. VSEPR Table: The bond angles in the table below are ideal angles from the simple VSEPR theory, followed by the actual angle for the example given in the following column where this differs. For many cases, such as trigonal pyramidal and bent, the actual angle for the example differs from the ideal angle, but all examples differ by different amounts. For example, the angle in H2S (92°) differs from the tetrahedral angle by much more than the angle for H2O (104.5°) does. Hybridization and Molecular geometry Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy. The concept of orbital hybridization focuses on mixing atomic orbitals to form new hybrid orbitals. The obtained hybrid orbitals have different energies, shapes, etc., than the component atomic orbitals. Hybrid orbitals are very useful in the explanation of molecular geometry and atomic bonding properties.  sp3, sp2 and sp hybridization in carbon Carbon ground state configuration The carbon atom has six electron, thus it electronic configuration is 1s22s22p2 The diagram bellow shows the expected orbital notation of carbon in its ground state sp3 hybridization of carbon (C) in CH4  Overlapping and bond formation in CH4 Geometry: Tetrahedral translates into  sp2 hybridization in carbon (covalent bonding - double bonds) Overlapping and bond formation in C2H4 Geometry: Planar Ex: formation of ethene is CH2=CH2 (C2H4). In ethene, each carbon atom is sp2-hybridized. In this way six sp2-orbitals are generated (three for each carbon). One sp2-orbital of each carbon atom by overlapping forms a sigma bond between carbon atoms. Remaining two sp2-orbital of each atom overlap with 1s- orbital of hydrogen atom to produce four sigma bonds. pz un-hybrid orbital of each carbon atom by the parallel overlapping form a pi-bond between two carbon atoms. Geometry in ethene molecule is trigonal in which bond angles are120o. sp2 diagram  Sp hybrid orbitals The total of two orbitals is formed (sp) : 2s mixes with 2p orbital, with two remaining unchanged p-orbitals. This is the type of hybridization of carbon in ethyne ( C2H2). The chemical bonding in ethyne consists of sp-sp overlap between the two carbon atoms forming a σ bond and two additional π bonds formed by p-p overlap. Each carbon also bonds to hydrogen in sigma s-sp overlap at 1800 angles.  Overlapping and bond formation in C2H2 Geometry: Linear 0 180 H C C H 0 180 sp diagram  Discussion: Tetracyanoethylene has the skeleton shown below How many sigma and pi From its Lewis structure determine the following: bonds are in the molecule? How many of the atoms are sp2 hybridized? How many of the atoms are sp hybridized?  Lecture 5 Representative metals, metalloids and nonmetals  Assignment n0 1 A series of six elements called the metalloids separate the metals from the nonmetals in the periodic table. The metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium. These elements look metallic; however, they do not conduct electricity as well as metals so they are semiconductors. Halogens, Sulfur, Oxygen, group1, group2, Carbon, Nitrogen, Al and Carbon. Question: Describe the general preparation, properties, compounds, toxicity and uses of the mentioned metalloids. Note: Do not exceed 10 slides Lectures Part B  Coordination compounds 1. General introduction 2. Formation of Complexes 3. The coordination number 4. Naming Coordination Compounds 5. Isomerism 6. CFSE  Coordination compounds The coordination complex or metal complex is a structure consisting of a central atom or ion (usually metallic), bonded to a surrounding array of molecules or anions (ligands, complexing agents). The Ligand is called the donor atom. The Polydentate (multiple bonded) ligands can form a chelate complex.  Coordination compounds Let’s consider Cobalt(III) chloride ammonate  Coordination compounds In 1893 Werner produced a theory to explain the structures, formation and nature of bonding in the coordination compounds. This theory is known as Werner’s theory of coordination compounds. Werner was the first inorganic chemist to be awarded the Nobel Prize for chemistry in 1913. He studied many complex compounds obtained from the reaction between cobalt chloride and ammonia.  Werner's theory The central metals of coordination compounds exhibit two types of valencies Primary Valency Secondary valency  Werner's theory Primary valances are those which a metal exhibits in the formation of simple salts. e.g. CoCl3 , NaCl, CuSO4 etc. In modern terminology it represents the oxidation number of metal. e.g. the primary valances of Co in CoCl3 is 3, and oxidation state +3 Similarly, for NaCl, oxidation state of Na is +1, for CuSO4, oxidation state of Cu is +2 The primary valances are ionizable. These are written outside the coordination sphere. These are non-directional and do not give any geometry to complex compound Example: [Co(NH3)6] Cl3, number of primary valances 3, oxidation state +3  Werner's theory The secondary valency of metals is either by negative ions or neutral molecules or both. In modern terminology it represents the coordination number of the metal. Secondary valencies are written inside the coordination sphere. These are directional in nature and give definite geometry to the complex. These are non-ionizable. Example: [Co (NH3)6]Cl3 coordination number is 6.  Werner's theory Werner explained the Structure and properties of following four complexes of Co (III) chloride with ammonia.  Werner's theory Werner added the above four cobalt(III) complexes in table with an excess silver nitrate solution. Which resulted in different amounts of silver chloride precipitate.  From the table, primary and secondary valencies are shown in these pictures, (a) shows 3 primary valencies, (b) shows 2 primary valencies, (c)shows 1 primary valence and in (d) all valencies are secondary Practice  Hint: If the compound is an octahedral complex then the central metal atom in the complex must be attached to 6 ligands by the coordinate bond. The ions outside the coordination bracket will take place in the reaction.  Electrical Conductance Measurement The conductance of solution depends upon the numbers of charge particles present in that solution. [Co(NH3)6]Cl3 → [Co(NH3)6]+ + 3Cl– [Co(NH3)5Cl]Cl2 → [Co(NH3)5Cl]+ + 2Cl– [Co(NH3)4Cl2]Cl → [Co(NH3)4Cl2]+ + Cl– So that the molar conductance of the compound should be the following order which also satisfies the observed conductance value. [Co(NH3)6]Cl3 > [Co(NH3)5Cl]Cl2 > [Co(NH3)4Cl2]Cl > [Co(NH3)3Cl3]  Precipitation Reaction On the addition of silver nitrate solution, with chloride complex. The chloride ions which present outside the coordination sphere undergo precipitation reaction. As the number of chloride ions present outside the sphere increases, the number of formation of precipitates increases and vice versa.  Limitations of Werner's theory Though Werner explained some properties of the coordination compound, he failed to explain the colour of the coordinate compound. He could not explain the magnetic and optical properties of coordination compounds. He could not answer the question, why does the coordination sphere have a definite geometry  Frequently Asked Questions – FAQs Q1 Explain Werner’s theory of coordinate compounds with suitable examples. Q2 Explain the bonding in CoCl3.3NH3 and CoCl3.5NH3 on the basis of Werner’s theory. Q3 The complex studied by Werner had a composition corresponding to the formula PtCl4.2KCl from electrical conductance measurements, he determined that each formula unit contained three ions. He also found that silver nitrate did not give a precipitate of AgCl with this complex. Write the formula for his complex that agrees with information?  Frequently Asked Questions – FAQs Q4 What are the limitations of Werner’s theory? Sol/Werner failed to explain the colour of the coordinate compound. He could not explain the magnetic and optical properties of coordination compounds  Frequently Asked Questions – FAQs Q5 What is the primary valency according to Werner’s theory? Sol/Primary valencies are those that a metal exhibits in the formation of simple salts. It represents the oxidation state of metal. These are ionisable and written outside the coordination sphere.  Introduction to coordination compounds To better understand the formation of complexes, we must remember the concept of hybridization of d – orbital Hybridization involving d-orbital ❑sp3d-hybrid orbitals: ❑ This Hybridization involves the mixing of one s, three p and one d-orbital. These five orbitals hybridize to form five sp3d-hybrid orbitals. Hybridization involving d-orbital  Hybridization involving d-orbital sp3d-hybrid Example Phosphorus pentafluoride PF5 involves sp3d- Hybridization as described below: The outer electronic configuration of Phosphorus, the central atom, is 3s2 3p3 It has three unpaired electrons in the ground state:  PF5 The conventional way of explaining the bonding in compounds like PF5 is to say that one of the 3s electrons is promoted into a 3d orbital, and then the s, p and d electrons hybridize to give five lots of sp3d hybrid orbitals which are then used to form bonds with the five fluorines.  Octahedral Geometry sp3d2-hybrid orbitals / d2sp3 one s, three p and two d-orbitals get hybridized to form six sp3d2hybrid orbitals: Octahedral Geometry  3 Example/sp d 2  3 Example/sp d 2 To account for the hexavalency in SF6, one electron each from 3s and 3p orbitals is promoted to 3d orbitals. These six orbitals get hybridised to form six sp3d2hybrid orbitals. Each of these sp3d2 hybrid orbitals overlaps with 2p orbital of fluorine to form S-F bond. octahedral structure  3 sp d 2 and 2 d sp 3  3 sp d 2 and 2 d sp 3  3 sp d 2 and 2 d sp 3 In some cases, the same central atom can form either inner or outer complexes depending on the particular ligand and the manner in which its electrostatic field affects the relative energies of the different orbitals. Thus the hexacyanoiron(II) ion utilizes the iron 3d orbitals, whereas hexaaquoiron(II) achieves a lower energy by accepting two H2O molecules in its 4d orbitals.  3 sp d 2 and 2 d sp 3  3 sp d 2 and 2 d sp 3 Although this "inner-outer" hybrid model was instrumental in explaining the properties of transition metal complexes, it has now been replaced with a more comprehensive model known as ligand field theory which will be introduced in another lesson.  Square Planar Geometry Example: Geometry of [Ni(CN)4]2- The complex ion [Ni(CN)4]2- involves dsp2 Hybridization as explained below. The oxidation state of Nickel is +2. The outer electronic configuration of Ni2+ is 3d8: Square Planar of [Ni(CN)4 ]2-  Square Planar of [Ni(CN)4 ]2- The two unpaired d electrons are paired up making one 3d orbital empty. The four empty orbitals (one 3d, 4s and two 4p) hybridize to form four dsp2 hybrid orbitals, which point towards square planar arrangement. Each one of the four CN- groups donates lone pair of electrons to vacant hybrid orbitals forming four Ni-CN bonds. Thus, [Ni(CN)4]2- has square planar arrangement. Structure of [Ni(CN)4 ]2- [Pt(Cl)4 ]2-  2 dsp -hybrid orbitals In addition, dsp2 type of hybridization is also known particularly in case of transition metal ions. The orbitals involved in this type of Hybridization are dx2- y2, s and two p. The four dsp2 hybrid orbitals adopt square planar geometry.  Hybridation 3 sp d 3 sp3d3 hybrid orbitals in IF7 This involves the mixing of one s, three p and three d-orbitals forming seven sp3d3 hybrid orbitals having pentagonal bipyramidal geometry. The geometry of IF7 molecule can be explained on the basis of sp3d3 Hybridization. Hybridation 3 sp d 3  Hybridation 3 sp d 3 These seven orbitals are then hybridized to give seven sp3d3 hybrid orbitals. Each of these sp3d3 hybrid orbitals overlaps with 2p orbitals of fluorine to form IF7 molecule having pentagonal bipyramidal geometry.  Formation of Coordination Complex Many ways such as:  The coordination number The coordination number is the number of donor atoms bonded to the central metal atom/ion.  Naming Coordination Compounds The basic procedure for naming a complex: 1. When naming a complex ion the ligands are named before the metal ion. Write the names of the ligands in alphabetical order. (Numerical prefixes do not affect the order). 2. Multiple occurring monodentate ligands receive a prefix according to the number of occurrences: di-, tri-, tetra-, penta-, or hexa. Polydentate ligands (e.g., ethylenediamine, oxalate) receive bis-, tris-, tetrakis- , etc. Anions end in o. This replaces the final 'e' when the anion ends with '-ate', e.g. sulfate becomes sulfato. It replaces 'ide': cyanide becomes cyano.  Examples for naming [NiCl4]2− → [CuNH3Cl5]3− → [Cd(en)2(CN)2] → [Co(NH3)5Cl]SO4 →  Table: Name of Metals in Anionic Complexes How to name metals in anionic complex Names of Some Common Ligands  Example1 1. K2[Co(NH3)2Cl4] Cations: 2x K+ Anion: [Co(NH3)2Cl4]2- Ligands: NH3 is neutral, Cl- is negative The central atom: Co2+ Name is: Potassium diamminetetrachlorocobaltate(II)  Example 2 What is the name for K3Fe(CN)6? Charge of metal ion = charge of complex ion - total charge of all ligands = +3 - 6x(-1) = +3 Fe3+, or Iron(III) 6 CN ligands: Hexacyano Complex ion: Hexacyanoferrate(III) Counter ion: Potassium (3K+) K3Fe(CN)6 : Potassium hexacyanoferrate(III)  Example 3 What is the formula for triammine bromoplatinum(II) chloride? Complex ion: [Pt(NH3)3Br]+ Counter ion: Cl- Molecular is: [Pt(NH3)3Br]Cl  Example 4: What is the formula for potassium hexafluorocolbaltate(III)? Sol/ K3 [CoF6]  Exercises I. Give the systematic names for the following coordination compounds: 1. [Cr(NH3)3(H2O)3]Cl3 2. [Pt(NH3)5Cl]Br3 3. [Pt(H2NCH2CH2NH2)2Cl2]Cl2 4. [Co(H2NCH2CH2NH2)3]2(SO4)3 5. K4[Fe(CN)6] 6. Na2[NiCl4] 7. [ Pt(NH3)2Cl4 ] 8. [ Fe(CO)5 ] 9. (NH4)2[Ni(C2O4)2(H2O)2] 10. [Ag(NH3)2][Ag(CN)2]  Exercises II. Can you give the molecular formulas of the following coordination compounds? 1. Hexaammineiron(III) nitrate 2. Ammonium tetrachlorocuprate(II) 3. Sodium monochloropentacyanoferrate(III) 4. Potassium hexafluorocobaltate(III)  Isomerism Two principal types of isomerism: 1. Stereoisomerism. (Geometrical isomerism and Optical isomerism) 2. Structural Isomerism (Coordination isomerism, Ionization isomerism, Hydrate isomerism, Linkage isomerism) Isomerism  Isomerism Structural Isomerism  Structural Isomerism A. Coordination isomerism: Examples One isomer [Co(NH3)6] [Cr(C2O4)3] another isomer [Co(C2O4)3] [Cr(NH3)6] B. Ionization isomers: different ions in Solution One isomer [PtBr(NH3)3]NO2 →NO2- anions in solution another isomer [Pt(NH3)3(NO2)]Br→ Br- anions in solution  The Geometric isomers A. The Geometric isomers: cis and Trans For example, there are two isomers of square planar [Pt(NH3)2Cl2 ]: Cis/trans diamminedichloroplatinum(II), Stereoisomerism Octahedral structure  The Geometric isomers The intensity of the Trans effect (as measured by the increase in rate of substitution of the trans ligand) follows this sequence: F-, Cl-, I-, Br-, SCN-, OH-< H2O,