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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. SiO 2 and CO 2 ). 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
- Current Inorganic Chemistry Published/Hosted by Bentham Science Publishers. ISSN (printed): 1877-9441. ISSN (electronic): 1877-945X
- Canadian Journal of Chemistry ISSN (print): 0008 - 4042 Reference Journals
- International Journal of Inorganic Chemistry ISSN: 2090- 2026
- Journal of Biological Inorganic Chemistry Published/Hosted by Springer. ISSN (printed): 0949-8257. ISSN (electronic): 1432- 1327 Course text books
- Inorganic Chemistry (4th^ Edition) Gary L. Miessler and Donald A Tarr ,Prentice Hall (2010) ISBN – 10: 0136128661,ISBN – 13: 978 - 0136128663
- Chemical Structure and Bonding, Roger L. Dekock, Harry B. Gray, University Science Books, 2 nd^ Edition (1989) ISBN – 10: 093570261X,ISBN – 13: 978- Reference text books
- Inorganic Chemistry, 3rd^ Edition (2007) Catherine Housecroft and Alan G. Sharpe, Prentice Hall,ISBN – 10: 0131755536,ISBN – 13: 978 – 0131755536
- Concise Inorganic Chemistry,J. D. Lee (1999), Wiley – Blackwell, 5th^ Edition ISBN-10: 0632052937, ISBN-13: 978-
PERIODIC TRENDS OF ATOMIC PROPERTIES
Electronegativity
Electronegativity = is a measure of an atom’s ability to attract to itself the shared electrons in a covalent bond. The electronegativity decreases as you go down a group. This is due to the shielding effect where electrons in inner energy levels shield the outer electrons from the attractive force of the nucleus. Electronegativity increases across a period. This is because there is more electron- attracting power of the nucleus with the increasing nuclear charge while electrons enter the same energy level (electrons in the same energy level do not shield one another)
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 )
Electron affinity
The electron affinity ( EA) is the energy released when an electron is added to a gaseous
atom. If the process is exothermic, H EA 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:
CHEMICAL BONDING
A chemical bond is an attraction between atoms that allows the formation of chemical substances that contain two or more atoms. The tremendous variety of substances that we have is due to the ability of atoms to bond or combine with other atoms. The type of bonds they form determines many properties of substances. For example, properties such as melting point and boiling point, solubility, hardness, shape of molecule, types of reactions etc. depend on the type of bonds present in the compounds.
LEWIS STRUCTURES
Lewis structures are important to understand chemical bonding. Shows only the valence electrons in an atom
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, H 2 , can be represented as follows:
.
In a similar fashion, two fluorine atoms can pair electrons to form an F 2 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 NH 4 + 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 H 3 O+
H 2 O H^ H 3 O.
H O
H H
Carbon monoxide, CO
b) Melting and boiling point : although not generally as strong as in ionic compounds, the bonding in metals is relatively strong, and as a result the melting and boiling points of metals are relatively high.
Metal Na K Be Mg Melting point/ oC 98 64 1278 649 Boiling point/ oC 883 760 2970 1107
Smaller ions, and those with a high charge, attract the electrons more strongly and so have higher melting points than larger ions with a low charge. Na has smaller cations than K so has a higher melting and boiling point. Mg cations have a higher charge than Na so has a higher melting and boiling point.
c) Other physical properties : Since the bonding in metals is non-directional, it does not really matter how the cations are oriented relative to each other. The metal cations can be moved around and there will still be delocalized electrons available to hold the cations together. The metal cations can thus slip over each other fairly easily. As a result, metals tend to be soft, malleable and ductile.
Determining Whether a Bond Will Be Ionic, Covalent or Metallic
Whether an ionic bond or a covalent bond will form between two elements depends on their electronegativity : their ability to attract electrons. Electronegative + Electropositive = Ionic Electronegative + Electronegative = Covalent Electropositive + Electropositive = Metallic -A large difference in electronegativity between two elements will result in ionic bonds while, no difference or a small difference in electronegativity will result in covalent bonds as are found in molecular compounds. -To determine what type of bond exists between two atoms you subtract their respective electronegativities:
i. If the electronegativity difference is 0.2 or less, the bond is completely covalent. If the electronegativity difference is greater than 0.2 but less than 1.7 the bond is called a polar covalent bond****. ii. If the electronegativity difference is 1.7 or greater the bond is ionic.
Pauline Electronegativity Scale
When the electronegativity difference is in this range the atom with the greater electronegativity is strong enough to pull the bonding electrons so that they spend more time around it than the other atom, but is not strong enough to pull the bonding electrons away completely and form ions. These bonding arrangements form electric dipoles where one end has a slightly positive charge and the other end has a slightly negative charge. The greater the electronegativity difference, the greater the dipole: Electronegativity Electronegativity Compound Bond Element bonded to F Fluorine difference Bond type FF FF F 4.0 4.0 0.0 COVALENT │ ↑ F 2 O OF O 3.5 4.0 0.5 increasingly ionic │ │ │ NF 3 NF N 3.0 4.0 1.0 │ │ CF 4 CF C 2.5 4.0 1.5 │ │ BF 3 BF B 2.0 4.0 2.0 │ │ BeF 2 BeF Be 1.5 4.0 2.5 │ increasingly covalent
Instead of an octet the valence shell of Be contains only two electron pairs ( quartet ) Similar arguments can be applied to boron trichloride, BCl 3 , 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 BeCl 2 and BCl 3 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 BeCl 2 reacts with Cl–^ ions to form BeCl 4 – ;
BCl 3 reacts with NH 3 in the following way:
(b) Species with Expanded Octets Examples of molecules with more than an octet of electrons are phosphorus pentafluoride (PF 5 ) and sulfur hexafluoride (SF 6 ). Phosphorus pentafluoride is a gas at room temperature. It consists of PF 5 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
Instead of an octet the phosphorus atom has 10 ( decet ) in its valence shell. Sulfur hexafluoride (also a gas) consists of SF 6 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 3 s^23 p^6. However, the 3 d subshell is also available, and some of the 3 d orbitals may also be involved in bonding. This permits the extra pair of electrons to occupy the valence ( n = 3) shell of phosphorus in PF 5. Expansion of the valence shell is impossible for an atom in the second period because there is no such thing as a 2 d orbital. The valence ( n = 2) shell of nitrogen, for example, consists of the 2 s and 2 d subshells only. Thus nitrogen can form NF 3 (in which nitrogen has an octet) but not NF 5. Phosphorus, on the other hand, forms both PF 3 and PF 5 , the latter involving expansion of the valence shell to include part of the 3 d 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.
The structure is said to be a resonace hybrid of three equivalent structures. A species is said to be in resonace if it can be written in two or more structures that differ only in the arrangement of electrons.
The resonace hybrid is more stable than any contributing structures (resonsnce energy).
Other familiar examples of species that are best described by resonance include the carbonate (CO 3 2-) and the nitrate (NO 3 - ), SO3, etc
Carbondioxide The formal charges for the two Lewis electron structures of CO 2 are as follows:
Both Lewis structures have a net formal charge of zero, but the structure on the right has a +1 charge on the more electronegative atom (O). Thus the symmetrical Lewis structure on the left is predicted to be more stable.
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.
a b a a a a a EN EN
V N B EN
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
(^1 022). 12.^14. 0 0. 311
H
7 6 2 4.^0
F^ -0.
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
