CHMA10 TT1 Review Part 2
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CHMA10
Introductory Chemistry I
Agenda:
1. Module 9
2. 6.9-6.9
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TYPE OF CHEMICAL BONDING
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LEWIS BONDING THEORY
 Because valence electrons are held most loosely, and chemical bonding involves the transfer or
sharing of electrons between two or more atoms
 Valence electrons are most important in bonding
 Lewis theory focuses on the behavior of the valence electrons
 General Rules
I. Show bonding (b) and non-bonding (nb) e−, and formal charges
= − ( + 2)
-V is the number of valence electrons of the atom in in ground state
-N is the number of non-bonding valence electrons on this atom in the molecule
-B is the total number of electrons shared in covalent bonds with other atoms in the molecule.
II. Bonding e− can be involved in single, double, triple bonds
III. ‘Complete shells’ can be achieved by combination of bonding and nonbonding e−
• A complete shell is an octet for most elements, except hydrogen, which is satisfied with
1 electron pair.
IV. Exceptions, where the octet rule is not followed are:
• Beryllium can be satisfied with 2 electron pairs
• Boron/Aluminum can be satisfied with 3 electron pairs
• Elements in period 3 and beyond can have an expanded octet if involved as the central
atom
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DRAW LEWIS STRUCTURE
1. Count total # of valence e– including charge of structure
2. Draw skeletal structure (central and terminal atoms)
• Least electronegative atom is usually the central atom
• Hydrogen and Fluorine are always terminal
3. Complete the octet of terminal atoms (or 2e– for hydrogen)
4. Subtract all e– used in previous steps and place any remaining e– on central atom.
• This sometimes leads to an expanded octet
5. Calculate formal charges (FC) on each atom
= ( − – – −)
6. Minimize formal charges by creating multiple bonds using nonbonding electrons
• Typically happens when neighbouring atoms have opposite charges.
• Only minimize when it will not give an element from the 2nd period > 8 e–.
7. Ensure all atoms have an allowed electron count (e.g. C, N, O, F must obey octet rule, and only
Be and group 13 elements are allowed to have less than an octet)
EXERCISE
1. NOCl
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2. SO3
3. HClO3
4. NH4
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5. BrOF2+
RESONANCE 共振
EXERCISE
O3-
NO3-
NO2-
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ENERGY OF IONIC BOND FORMATION
 The ionization energy of the metal is endothermic
() → +() + 1 − ° = +496 /
 The electron affinity of the nonmetal is exothermic 122() + 1 − → −() ° = −244 /
 But the heat of formation of most ionic compounds is exothermic and generally large.
() + 122() → () ° = −411 /
 Crystal Lattice!!!
• The extra energy that is released comes from the formation of a structure in which every cation is
surrounded by anions, and vice versa
• there is no ionic molecule in the crystal
• The extra stability that accompanies the formation of the crystal lattice is measured as the lattice
energy
• Lattice energy depends directly on size of charges and inversely on distance between ions
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REVIEW OF ENTHALPY AND THERMODYNAMICS
 The value of ΔrH for a chemical reaction is the amount of heat absorbed or evolved in the reaction
under conditions of constant pressure.
 An endothermic reaction absorbs heat from the surroundings and has a positive ΔrH. An endothermic
reaction feels cold to the touch.
 An exothermic reaction gives off heat to the surroundings and has a negative ΔrH. An exothermic
reaction feels warm to the touch.
Endothermic Absorb heat ∆ Touch feel cold
Exothermic Give off heat ∆ Touch feel warm
 Relationships involving enthalpy
 Hess’s Law
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 the Born–Haber Cycle
for NaCl
() + 122() → ()
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EXERCISE
6. Use the data given below to construct a Born-Haber cycle to determine the electron affinity of Br.
ΔH°(kJ mol-1)
() → () 89
() → + () + − 419
2() → 2() 193
() + 1
2
2() → () -394
() → + () + −() 674
A) -885 kJ mol-1
B) -325 kJ mol-1
C) +367 kJ mol-1
D) -464 kJ mol-1
E) +246 kJ mol-1
7. Which of the following reactions is associated with the lattice energy of Li2O (ΔrH_lattice^° )?
A) Li2O(s) → 2Li+ (g) + O2– (g)
B) 2Li+ (aq) + O2– (aq) → Li2O(s)
C) 2Li+ (g) + O2– (g) → Li2O(s)
D) Li2O(s) → 2Li+ (aq) + O2– (aq)
E) 2Li(s) + O2(g) → Li2O(s)
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TRENDS IN LATTICE ENERGY
 Size:
• The force of attraction between charged particles decreases with increase of the distance between
them
• larger ion = weaker attraction
• weaker attraction = smaller lattice energy
 charge:
• larger charge = stronger attraction
• stronger attraction = larger lattice energy
EXERCISE
8. Order the following ionic compounds in order of increasing magnitude of lattice energy: MgS,
NaBr, LiBr, SrS
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INONIC BONDING: MODEL VS REALITY
1. Lewis theory predicts the number of electrons a metal atom should lose or a nonmetal atom
should gain in order to attain a stable electron arrangement
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2. Lewis theory predicts ionic compounds should have high melting points and boiling points.
 the larger the lattice energy, the higher the melting point
EXERCISE
9. Compare the melting point of KBr and MgO.
3. Lewis theory predicts that ionic solids should not conduct electricity.
4. Lewis theory predicts that both a liquid ionic compound and an ionic compound dissolved in
water should conduct electricity.
5. Lewis theory predicts that Ionic solids are brittle
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COVALENT BONDING
1. Overview
 When nonmetals bond together, it is better in terms of potential energy for the atoms to share
valence electrons
 Shared electrons hold the atoms together by attracting nuclei of both atoms
2. Covalent Bonding: Model vs Reality
1) Lewis Theory Predicts Electron Groups
• Lewis theory predicts there are regions of electrons in an atom
• Some regions result from placing shared pairs of valence electrons between bonding
nuclei (bonding pair)
• Other regions result from placing unshared valence electrons on a single nuclei (lone
pair)
2) Lewis theory predicts the melting and boiling points of molecular compounds should be
relatively low
• involves breaking the attractions between the molecules, but not the bonds between the
atoms
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3) Lewis theory predicts that neither molecular solids nor liquids should conduct electricity
• Molecular acids conduct electricity when dissolved in water, but not in the solid or liquid
state, due to them being ionized by the water
4) Lewis theory predicts covalently bonded compounds will be found as individual molecules
• rather than an array like ionic compounds
• the attractions between atoms are directional
5) Bond Strength
Lewis theory predicts that the more electrons two atoms share, the stronger the bond should
be.
• In general, triple bonds are stronger than double bonds, and double bonds are stronger
than single bonds
• however, Lewis theory would predict double bonds are twice as strong as single bonds,
but the reality is they are less than twice as strong
• Covalent bond types
6) Bond Length
Lewis theory predicts that the more electrons two atoms share, the shorter the bond should be.
• Bond length is determined by measuring the distance between the nuclei of bonded
atoms
• In general, triple bonds are shorter than double bonds, and double bonds are shorter than
single bonds
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EXERCISE
10. Which compound has the longest carbon-carbon bond length?
A) CH3CH3
B) CH2CH2
C) HCCH
D) All bond lengths are the same.
11. Which compound has the highest carbon-carbon bond strength?
A) CH3CH3
B) CH2CH2
C) HCCH
D) All bond strengths are the same.
BOND DIPOLE MOMENT
 Dipole moment, μ, is a measure of bond polarity
• Generally, the more opposite partial charge on two atoms and the larger the atoms are, the larger
the dipole moment
 Percent Ionic Character
• the percentage of a bond’s measured dipole moment compared to what it would be if the
electrons were completely transferred
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BOND ENERGY
 The amount of energy it takes to break one mole of a bond in a compound is called the bond
energy
• in the gas state
• homolytically–each atom gets ½ bonding electrons
 In general, the more electrons two atoms share, the stronger the covalent bond
 In general, the shorter the covalent bond, the stronger the bond
• bonds get weaker down the column
• bonds get stronger across the period
➢ Using Bond Energies to Estimate ΔH°rxn
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LIMITATIONS OF LEWIS THEORY
 Lewis theory generally predicts trends in properties, but does not give good quantitative
predictions
• e.g. bond strength and bond length
 Lewis theory gives good first approximations of the bond angles in molecules, but usually cannot
be used to get the actual angle
 Lewis theory cannot write one correct structure for many molecules where resonance is important
 Lewis theory often does not predict the correct magnetic behavior of molecules
• e.g. O2 is paramagnetic, though the Lewis structure predicts it is diamagnetic
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