Chemical Bonding I: Basic Concepts Chapter 9 Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that participate in chemical bonding. Group e - configuration # of valence e - 1A ns 1 1 2A ns 2 2 3A ns 2 np 1 3 4A ns 2 np 2 4 5A ns 2 np 3 5 6A ns 2 np 4 6 7A ns 2 np 5 7 Lewis Dot Symbol
Lewis Symbols for Atoms Element symbol = nucleus + core electrons Valence electrons are drawn as dots around the symbol Up to 4 valence electrons are placed around the symbol one at a time; additional electrons are paired up The result is up to 4 pairs of electrons = octet NOTE: hydrogen can not have an octet. When forming bonds with other atoms, it can have a maximum of 2 electrons in its valence shell O Lewis Dot Symbols are not drawn for transition metals
Li(s) + ½ F 2 (g) LiF(s) After F 2 breaks apart into two neutral F atoms - Li Li + + e - 1s 2 2s 1 1s 2 = [He] e - + F F - Li + + F - Li + F - 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6 = [Ne] Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions. It is always endothermic. e.g. MgF 2 (s) Mg 2+ (g) + 2F - (g) E = k Q + Q - r Coulomb s Law Q + is the charge on the cation Q - is the charge on the anion r is the distance between the ions Lattice energy (E) increases as Q increases and/or as r decreases. cmpd MgF 2 MgO lattice energy 2957 Q= +2,-1 3938 Q= +2,-2 LiF LiCl 1036 853 r F < r Cl
Born-Haber Cycle for Determining Lattice Energy o H 5 < 0 exothermic opposite to lattice energy E o H 5 = -E o o o o o o o H overall = H rxn = H 1 + H 2 + H 3 + H 4 + H 5 o solve for H 5 = -E o o o o o o H rxn = H 1 + H 2 + H 3 + H 4 + H 5
A covalent bond is a chemical bond in which two or more electrons are shared by two atoms, resulting in an octet for both atoms. Lewis structure of F 2 Lewis structure of H 2 O
Double bond two atoms share two pairs of electrons CO 2 Triple bond two atoms share three pairs of electrons N 2 Guidelines for Drawing Lewis Structures (updated later on with the concept of formal charge ) 1. Hydrogen is always a terminal atom because it can form only one bond. 2. The CENTRAL ATOM usually has the lowest electron affinity (or electronegativity as defined later) 3. Arrange the atoms geometrically and symmetrically. e.g. CH 2 O
4. Sum up the total number of valence electrons (use the group number), and calculate the number of pairs. 5. Connect the atoms together so that each atom has an octet (except H). You may have to form multiple bonds. NF 3 HNO 3
Lewis Structures of Charged Species ClO - NO 2 + Lengths of Covalent Bonds Bond Type C-C C=C C C C-N C=N C N Bond Length (pm) 154 133 120 143 138 116 Bond Lengths Triple bond < Double Bond < Single Bond
Electronegativity is the ability of an atom to attract the electrons in a chemical bond towards itself. Electron Affinity - measurable, Cl is highest X (g) + e - X - (g) Electronegativity Pauling Scale (relative scale), F is highest
Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms electron poor region electron rich region e - poor H F H F δ + δ - e - rich Classification of bonds by difference in electronegativity Difference Bond Type 0 Covalent 2 Ionic 0 < and <2 Polar Covalent Increasing difference in electronegativity Covalent share e - Polar Covalent partial transfer of e - Ionic transfer e -
An atom s formal charge is the difference between the number of valence electrons surrounding an isolated atom, and the number of electrons assigned to that atom in a Lewis structure. N F N O H
formal charge on an atom in a Lewis structure = total number of valence electrons in - the free atom total number of nonbonding electrons - 1 2 ( ) total number of bonding electrons The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.
Formal Charge and Lewis Structures 1. For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. 2. Lewis structures with large formal charges are less plausible than those with small formal charges. 3. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. What is the most likely Lewis structure for CH 2 O?
A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. e.g. O 3 e.g. what are the resonance structures of the carbonate (CO 3 2- ) ion?
1. The Incomplete Octet (Central atom in Group 3A) e.g. BF 3 2. Odd-Electron Molecules NO NO 2
3. Expanded Octet central atom has greater than 8 valence electrons surrounding it. Occurs only with elements in row 3 and higher because they have available d-orbitals A. Covalent Molecules SF 6
B. Polyatomic Ions PO 4 3- SO 4 2-
The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy. Bond Energy H 2 (g) H (g) + H (g) H 0 = 436.4 kj Cl 2 (g) HCl (g) Cl (g) + Cl (g) H 0 = 242.7 kj H (g) + Cl (g) H 0 = 431.9 kj O 2 (g) O (g) + O (g) H 0 = 498.7 kj O O N 2 (g) N (g) + N (g) H 0 = 941.4 kj N N Bond Energies Single bond < Double bond < Triple bond
Average bond energy in polyatomic molecules H 2 O (g) H (g) +OH (g) H 0 = 502 kj OH (g) H (g) +O (g) H 0 = 427 kj Average OH bond energy = 502 + 427 2 = 464 kj Bond Energies (BE) and Enthalpy changes in reactions Imagine reaction proceeding by breaking all bonds in the reactants and then using the gaseous atoms to form all the bonds in the products. H 0 = total energy input total energy released = ΣBE(reactants) ΣBE(products)
Use bond energies to calculate the enthalpy change for: 2C 2 H 6 (g) + 7O 2 (g) 4CO 2 (g) + 6H 2 O(g) H 0 = ΣBE(reactants) ΣBE(products) Type of bonds broken Number of bonds broken Bond energy (kj/mol) Energy change (kj) Type of bonds formed Number of bonds formed Bond energy (kj/mol) Energy change (kj)