Atomic Properties: Periodic Trends in Electronegativity, Ionization, and Electron Affinity, Exercises of Chemistry

Download Atomic Properties: Periodic Trends in Electronegativity, Ionization, and Electron Affinity and more Exercises Chemistry in PDF only on Docsity! Page 1 of 22 SCH 2101 CHEMICAL BONDING AND STRUCTURE (45 hours) Course Description  Some atomic properties and their variations across the period and down groups. Atomic size, Ionization energy, electronegativity, electron affinity.  Qualitative treatment of bonding in terms of dot and cross formula. Deviations from the octet rule. Bond types; covalent, dative, ionic and metallic bonding  Atomic properties and their effects on bonding types.  Qualitative treatment of resonance.  Valency bond theory and hybridization  Intermolecular forces- and hydrogen bonding Van-der-Waal's. Hybridization of atomic orbitals and shapes of simple molecules. Relation between structure and physical properties (e.g. SiO2 and CO2). Acids and bases. Practicals will be on further work on acid- base and redox titrations. Teaching Methodology: Lectures, Tutorials and practicals Course assessment Written CATS 30%, final written examination 70% Course Journals 1. Current Inorganic Chemistry Published/Hosted by Bentham Science Publishers. ISSN (printed): 1877-9441. ISSN (electronic): 1877-945X 2. Canadian Journal of Chemistry ISSN (print): 0008-4042 Reference Journals 1. International Journal of Inorganic Chemistry ISSN: 2090-2026 2. Journal of Biological Inorganic Chemistry Published/Hosted by Springer. ISSN (printed): 0949-8257. ISSN (electronic): 1432-1327 Course text books 1. Inorganic Chemistry (4th Edition) Gary L. Miessler and Donald A Tarr ,Prentice Hall (2010) ISBN – 10: 0136128661,ISBN – 13: 978 - 0136128663 2. Chemical Structure and Bonding, Roger L. Dekock, Harry B. Gray, University Science Books, 2nd Edition (1989) ISBN – 10: 093570261X,ISBN – 13: 978-0935702613 Reference text books Page 2 of 22 1. Inorganic Chemistry, 3rd Edition (2007) Catherine Housecroft and Alan G. Sharpe, Prentice Hall,ISBN – 10: 0131755536,ISBN – 13: 978 – 0131755536 2. Concise Inorganic Chemistry,J. D. Lee (1999), Wiley – Blackwell, 5th Edition ISBN-10: 0632052937, ISBN-13: 978-0632052936 Page 5 of 22  Elements with high electronegativities readily gain electrons to form anions e.g F-, O2- while those of low electronegativites readily lose electrons to form cations e.g Na+, Ca2+.  The American chemist Linus Pauling developed a convenient measure of electronegativity called the Pauling Electronegativity Scale, in which the values range from 0.7 (lowest) for Fr to 4.0 (highest) for F. Ionization Energy  Ionization energy is the minimum energy required to remove an electron from an isolated gaseous atom in its ground state.  Units is kilojoules per mole (kJ/mol)  Na(g) Na+ + e- + ΔE (positive value = Endothermic)  The value of ionization energy is a measure of how tightly the valence electron is held in the atom. The higher the ionization energy, the more difficult it is to remove the electron.  Metals are characterized by low ionization energies and tend to lose electrons to form cations (Electropositive) Page 6 of 22  Nonmetals have high ionization energies and tend to gain electrons to form anions (electronegative). Factors affecting ionization energy i. Nuclear charge – The larger the nuclear charge, the greater the ionization energy. Nuclear charge increases across the period, hence I.E. ii. Shielding effect – The greater the shielding effect the less the ionization energy. Shielding effect increases down a group. iii. Radius – The greater the distance between the nucleus and the outer electrons of an atom, the less the ionization energy. iv. Sublevel/orbital – The ease with which an electron is removed depends on the nature of the orbital. S> p> d > f. This is related to the shape of the orbital. e.g the first I.E of Mg is higher than that of Al because it is easier to remove a p-electrom in Al than an s-electron in Mg. v. An electron from a full or half-full sublevel requires additional energy to be removed. Half-filled or completely filled configurations are more stable. This accounts for the exceptionally high I.E of noble gases. I.E of P (haff-filled p-level i.e 3p3) is higher than that of S (3p4) yet expected to increase across the period. Page 7 of 22 Electron affinity  The electron affinity (EA) is the energy released when an electron is added to a gaseous atom. If the process is exothermic, HEA is (-) and (+) endothermic. F(g) + e- F- (g) + HEA  The halogens have the most exothermic electron affinities of all the elements.  As one progresses from left to right across a period, the electron affinity will increase, due to the larger attraction from the nucleus. Down a group, the electron affinity decreases because of a large increase in the atomic radius, electron-electron repulsion and the shielding effect of inner electrons against the valence electrons of the atom. Atomic Radius  The atomic radius is simply the distance from the nucleus to the outermost electron. Since the position of the outermost electron can never be known precisely, the atomic radius is usually defined as half the distance between the nuclei of two bonded atoms of the same element:  Atomic radius increases down a group and decreases across a period.  The following graphic shows the trend in atomic radius: Page 10 of 22  The elements of Groups I, II and III can form the electronic structure of an inert gas by losing their outer 1, 2 and 3 (valence) electrons. (The resulting species are positively charged ions.)  elements of Groups V, VI and VII form noble gas structure by gaining 1,2 and 3 electrons (by formation of negatively charged ions).  The compounds formed involve electrostatic attraction (electrovalent bonds) of oppositely charged species called ions.  The electrovalent bond is the result of electrostatic attraction between ions of opposite charge. This attractive force accounts for the stability of these compounds, typified by NaF, LiCl, CaO, and KCl.  The electrostatic forces are active in all directions; they attract oppositely charged species and thus can form regular arrays, resulting in ordered lattice structures, i.e. the solid state. COVALENT BONDING  A covalent bond is a form of chemical bonding that is characterized by the sharing of pairs of electrons between atoms. Each of the combining atoms contribute one electron to the shared pair.  Pure covalent bonding only occurs when two nonmetal atoms of the same kind bind to each other.  Lewis structures are used to represent molecules. The covalent bond in the hydrogen molecule, H2, can be represented as follows: . Page 11 of 22  In a similar fashion, two fluorine atoms can pair electrons to form an F2 molecule. Some molecules require more than single bonds to provide each atom with the required Octet eg CO-ORDINATE (DATIVE COVALENT) BONDING A co-ordinate bond (also called a dative covalent bond) is a covalent bond in which the shared pair of electrons is contributed by only one atom. Ammonium ion NH4 + Formed by the transfer of a hydrogen ion to the lone pair of electrons on the ammonia molecule In simple diagrams, a co-ordinate bond is shown by an arrow. The arrow points from the atom donating the lone pair to the atom accepting it. Hydroxonium ion H3O+ H2O H H3O . H O H H Carbon monoxide, CO Page 12 of 22 Carbon monoxide can be thought of as having two ordinary covalent bonds between the carbon and the oxygen plus a co-ordinate bond using a lone pair on the oxygen atom. C O Metallic bonding Metals are formed from elements on the left hand side of the periodic table. Having generally low electronegativity they tend to lose their valence electrons easily. When we have a macroscopic collection of the same or similar type of metallic atoms, the valence electrons are detached from the atoms but not held by any of the other atoms. In other words, these valence electrons are free from any particular atom and are only held collectively by the entire assemblage of atoms. Valence electrons are detached from atoms, and spread in an 'electron sea' that holds the positive ions together. In a metal the ion cores are held more or less at fixed places in an ordered, or crystal, lattice. The valence electrons are free to move about under applied stimulation, such as electric fields or heat. Properties of metals a) Electrical conductivity: since the electrons in a metal are delocalised, they are free to move throughout the crystal in a certain direction when a potential difference is applied and metals can thus conduct electricity in the solid state. The delocalised electron system is still present in the liquid state, so metals can also conduct electricity well in the liquid state. Page 15 of 22 ↓ │ LiF LiF Li 1.0 4.0 3.0 IONIC Assignment Classify each of bonds as ionic, covalent or polar covalent using the Pauling Electronegativity Scale Al–I; Br–Cl; B–F; Ca–Cl; Mg–Cl; K–S Exceptions to the Octet Rule (a) Electron Deficient Species Good examples of the first type of exception are provided by BeCl2 and BCl3. Beryllium dichloride, BeCl2, is a covalent rather than an ionic substance. Solid BeCl2 has a relatively Complex structure at room temperature, but when it is heated to 750°C, a vapor which consists of separate BeCl2 molecules is obtained. Since Cl atoms do not readily form multiple bonds, we expect the Be atom to be joined to each Cl atom by a single bond. The structure is Page 16 of 22 Instead of an octet the valence shell of Be contains only two electron pairs (quartet) Similar arguments can be applied to boron trichloride, BCl3, which is a stable gas at room temperature. We are forced to write its structure as in which the valence shell of boron has only three pairs of electrons (sextet). Molecules such as BeCl2 and BCl3 are referred to as electron deficient because some atoms do not have complete octets. Electron-deficient molecules typically react with species containing lone pairs, acquiring octets by formation of coordinate covalent bonds. Thus BeCl2 reacts with Cl– ions to form BeCl4 – ; BCl3 reacts with NH3 in the following way: (b) Species with Expanded Octets Examples of molecules with more than an octet of electrons are phosphorus pentafluoride (PF5) and sulfur hexafluoride (SF6). Phosphorus pentafluoride is a gas at room temperature. It consists of PF5 molecules in which each fluorine atom is bonded to the phosphorus atom. Since each bond corresponds to a shared pair of electrons, the Lewis structure is Page 17 of 22 Instead of an octet the phosphorus atom has 10 (decet) in its valence shell. Sulfur hexafluoride (also a gas) consists of SF6 molecules. Its structure is Here the sulfur atom has six electron pairs (duodecet) in its valence shell. An atom like phosphorus or sulfur which has more than an octet is said to have expanded its valence shell. This can only occur when the valence shell has enough orbitals to accommodate the extra electrons. For example, in the case of phosphorus, the valence shell has a principal quantum number n = 3. An octet would be 3s23p6. However, the 3d subshell is also available, and some of the 3d orbitals may also be involved in bonding. This permits the extra pair of electrons to occupy the valence (n = 3) shell of phosphorus in PF5. Expansion of the valence shell is impossible for an atom in the second period because there is no such thing as a 2d orbital. The valence (n = 2) shell of nitrogen, for example, consists of the 2s and 2d subshells only. Thus nitrogen can form NF3 (in which nitrogen has an octet) but not NF5. Phosphorus, on the other hand, forms both PF3 and PF5, the latter involving expansion of the valence shell to include part of the 3d subshell. (c) Free Radicals The majority of molecules or complex ions discussed in general chemistry courses are demonstrated to have pairs of electrons. However, there are a few stable molecules which contain an odd number of electrons. These molecules, called "free radicals", contain at least one unpaired electron, a clear violation of the octet rule. Free radicals play many important roles a wide range of applied chemistry fields, including biology, medicine, and astrochemistry. Three well-known examples of such molecules are nitrogen (II) oxide, nitrogen (IV) oxide, and chlorine dioxide. The most plausible Lewis structures for these molecules are Free radicals are usually more reactive than the average molecule in which all electrons are paired. In particular they tend to combine with other molecules so that their unpaired electron finds a partner of opposite spin. Page 20 of 22 CALCULATION OF PARTIAL AND FORMAL CHARGE (i) Partial Charge  The partial charge on an atom is determined by the difference between the electronegativities of the atoms that form the covalent bond.  The difference in electronegativity between the two atoms has no effect on the nonbonding electrons; only affects the distribution of the bonding (shared electrons)  The calculation of partial charge therefore involves the electronegativities of the two atoms (ENa and ENb) that form the covalent bond.           ba a aaaa ENEN EN BNV Where Va = no of valence electrons in a, Na is the number of nonbonding electron in a, Ba is the number of bonding electrons in a and ENa and ENb are the electron affinities of a and b, respectively. Consider HF 311.0 0.41.2 1.2 201        H         0.41.2 0.4 267F -0.311 NB.  The magnitude of partial charge is the same in each atom, only the sign of the charge differs i.e –ve in the more electronegative atom and +ve at the less electronegative atom.  This means that fluorine atom has 31.1% more electron density than hydrogen The sum of the partial charges is zero, so that the molecule is electrically neutral. Exercise Confirm that partial charge = 0.11 for HCl and 0.08 for HBr  Page 21 of 22 (ii) Formal Charge  Formal charge helps us identify which atoms are more likely to carry a significant positive or negative charge.  For the formal charge on atom a 2 a aaa B NVFC  Where Va = valence electrons on the neutral atoms of the element, Na = number of nonboding electrons on the element and Ba = the bonding electrons on the atom.  Consider the case of the N2O, nitrous oxide gas. There are three possible structures: Which is the best possible representation of the molecule?  The best structure is considered to be the one in which; (i) The atoms have the smallest formal charge (ii) The negative formal charges are on the more electronegative atoms. Calculation of formal charges (i) N1 = 5-4-4/2 = -1. N2 = 5-0-8/2 = +1. O = 6-4-4/2 = 0 (ii) N1 = 5-2-6/2 = 0; N2 = 5-0-8/2 = +1; O = 6-6-2/2 = -1 Page 22 of 22 (iii) N1 = 5-6-2/2= -2; N2 = 5-0-8/2 = +1; O = 6-2-6/2 = +1 -The sum of the formal charges must equal the overall charge on the molecule = 0 for this molecule Structure 3 has several flaws (i) Puts a +ve charge on the more electronegative atom (ii) It has the largest formal charge Thus (a) and (b) are the plausible resonance structures of N2O. Thiocyanate  The thiocyanate ion (SCN−), has 3 possible resonance structures; Calculate the formal charges on each atom and show that structure (b) is the more stable structure. Exercise i. Draw three Lewis structures for CNO− and use formal charges to predict which is more stable. (Note: N is the central atom.) ii. Using arguments based on formal charges, explain why the most feasible Lewis structure for SO4 2− has two sulfur–oxygen double bonds.