1. CHAPTER 8
AP CHEMISTRY

2. COVALENT BONDING Polar bonds Nonpolar bonds
Uneven attraction for the shared electrons
Nonpolar bonds
Bonding electrons are shared equally by bonded atoms, balanced distribution of charge

3. CONTINUE
Molecule
Group of two or more atoms (same or different type of elements) held together by covalent bonds (sharing electrons) and can exist in nature
Diatomic - two atoms of the same element covalently bonded
Chemical formulas have symbols, numerical subscripts, and states the type of atoms and the number of each atom
Potential energy is at a minimum when attractive forces balance repulsive forces
A bond is formed if the potential energy is lower as a compound than as single atoms

4. OCTET RULE
Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet (eight) of electrons in its valence shell
Exceptions to the octet rule
Has less than a full octet around the central atom because the atom is too small (boron and beryllium)
Has more than a full octet (10 or 12) around the central atom, central atom is large with empty d and f sublevels
Central atom is the least electronegative atom
Central atom is a non-metal found in periods 3, 4, and 5 and are surrounded by halogens

5. LEWIS STRUCTURE Lone pair Lewis structure
Unshared pair of electrons, not involved in bonding and belongs to one atom
Oxygen has two lone pair electrons
Lewis structure
FORMULAS IN WHICH ATOMIC SYMBOLS REPRESENT NUCLEI AND INNER-SHELL ELECTRONS, DOTS OR DASHES BETWEEN TWO SYMBOLS REPRESENT BONDS AND DOTS REPRESENT UNSHARED ELECTRONS

6. STRUCTURAL FORMULA Arrangement of atoms and where the bonds are formed
Single bond: sharing two electrons
Double bond: sharing four electrons, higher bond energy than a single bond so bond length is shorter
Triple bond: sharing six electrons, higher bond energy than a double bond, so bond length is shorter

7. BONDS Bond energy Ionic bonds coulomb’s law
energy required to break a bond
Ionic bonds
very stable, high melting point
metal and nonmetal usually form these bonds
coulomb’s law
energy of interactions
E = 2.31 x J.nm (Q1Q2) r
Q1 and Q2 numerical charge

8. CONTINUE
A bond will form when the energy in a compound is lower than as separate atoms
this will minimize the sum of positive (repulsive) energy and negative (attractive) energy terms and is called the BOND LENGTH

9. ELECTRONEGATIVITY
Ability of an atom to attract shared electrons to itself
page 334
polarity
dipole moment >
б б- б- б- Б+

10. PREDICTING IONIC COMPOUNDS
Ca + O <=> Ca2+ + O 2-
calcium is transferring 2 electron‘s to the oxygen
positive ions are
negative ions are
Isoelectronic ions have the same electron configuration as the noble gas it is close to.

11. LATTICE ENERGY change in energy when gaseous ions form ionic solids
page 343
lattice energy = k(Q1Q2), k = proportionality
( r ) constant
covalent bonds with ionic characteristics
dipole moment gives a covalent bond a percent of ionic character
(measured dipole moment X-Y) X 100
(calculated dipole moment X+ Y-)
If the compound when melted can conduct an electrical current it is called a SALT

12. COVALENT BONDING forces which hold nonmetal atoms together
because this bond consists of sharing electrons it is very stable

13. LEWIS STRUCTURES
eight electrons in the valence shell. When there are eight electrons in the outer most energy level the atom is stable
number of valence electrons = 1 to 8
In the lewis structure a covalent bond is a line.
single bond = 2 electrons shared
double bond = 4 electrons shared
triple bond = 6 electrons shared

14. WRITING LEWIS STUCTURES
draw a skeleton structure by joining the atoms by single bonds
central atom written first in the formula
the following are usually terminal atoms
H, Br, F, I, Cl, O
make sure all bonds are accountable
determine the number of valence electrons that have not been accounted for

15. RESONANCE FORMS
the concept of resonance is involved whenever a single Lewis structure does not explain all bonds
Formal charge
the charge on an atom in a molecule or ion that is calculated by assuming that all lone pairs of electrons shared by an atom, belong to the atom
Cf = Ev - (Eu + 1/2Eb)
Cf = formal charge, Ev = # of valence electrons
Eu= # of unshared electrons, Eb= # of bonded electrons

16. EXCEPTIONS TO THE OCTET RULE
molecules that contain an odd number of valence electrons
species that contain an odd number of electrons are paramagnetic (attracted by a charge)
less than a full octet of electrons
expanded octet surrounded by more than four pairs of valence electrons

17. BOND ENERGY
enthalpy change, ΔH, to break a bond in a mole of gaseous substance
endothermic if ΔH > 0
exothermic if ΔH < 0
ΔH = (bond energies broken) -(bond energies formed)

18. CONTINUE bond strength and length bond polarity
as the number of bonds increase the length decreases making then stronger
bond polarity
nonpolar bond atoms share electrons equally
polar bond atoms share electrons unequally
ionic bond atoms do not share electrons