Unit 4 Chemical Bonding
Ch. 8 and 9
Adapted from:http://www.sciencegeek.net/APchemistry/Powerpoints.shtml
Unit 3 Exam Analysis
 Grab out a different colored pen/pencil and a sheet of
paper.
 Ask questions, take notes on areas to review
(You will see this information on the term final!)
 Average
*Studying Tips (Blog)*
Unit 4 Topic Sheet!
 Please access the blog-under “Notes”
Why do atoms bond?
 What are some different types of bonds?
 How are the different types of bonds similar?
Different?
Ionic Bonding
 Electrons are transferred
 Electronegativity differences larger than 1.7
 The formation of ionic bonds is EXOTHERMIC.
 Let’s connect this idea to last unit.
Ionic Character
 Electronegativity
difference of 1.7 or
greater
AND
 Conduct electricity in
their molten or
aqueous state
Coulomb’s Law
 The energy of interaction between a pair of ions is
proportional to the product of their charges, divided by the
distance between their centers.
E = (2.31x10
-19
Q1Q2
J ·nm)(
)
r
Q1Q2
E µ(
)
r
 Energy is _____related to distance between charges.
 Energy is _____related to magnitude of charges.
Crystal Lattice (NaCl)
 Ionic compounds form solid crystals at ordinary
temperatures.
 Ionic compounds organize in a characteristic crystal
lattice of alternating positive and negative ions.
 All salts are ionic compounds and form crystals.
Ionic Compounds Properties
 Types of elements involved (Metal and Nonmetal)
 Polyatomic Ion Salts
 Structure: Crystalline solids (previous slide)
 Melting point: generally high (compared to covalent cpds)
 Boiling point: generally high (compared to covalent cpds)
 Electrical Conductivity: excellent conductors (molten and
aqueous state)
 Solubility in water: generally soluble
 Make connection to our “1” solubility rule
Ionic Radii
 Cations
 Lose electrons; same number of protons “stronger” effective nuclear
charge resulting in a smaller radius than neutral atom
 Anions
 Gain electrons; same number of protons “weaker” effective nuclear
charge resulting in a larger radius than neutral atom
 Isoelectronic ions
 Ions containing the same number of electrons
 O2-; F-; Na+; Mg2+; and Al3+
 All have Ne noble gas notation (same number of e- different numbers of
p+)
Application
 Order the following ions from smallest to largest. Be
able to EXPLAIN your ordering!
K+; Ca2+; Ar; S2-; Cl*Besides Ar what do we call this set of ions?
Term 1 Update….Finish STRONG!
 Term Finals March 11th and 12th!
 Begin Reviewing Unit Topic Lists, Notes, Practice Problems.
 26 Multiple Choice 5 Free Response
 Will have a review day Tuesday, March 10th
 Will include Unit 1-3 and covered parts of Unit 4.
 ALL Term 1 Materials need to be completed by March, 9th
(Monday)
 If you have missed an exam it will be taken Tuesday, March 10th during
the review day.
 All lab reports will need to be submitted by this date for credit.
 Final Lab of the Term will be next week.
 Your lowest lab grade will be dropped.
Lattice Energy
æ Q1Q2 ö
LatticeEnergy = k ç
÷
è r ø
 The energy released when an ionic solid is formed from
its gaseous ions.
 Many separate processes go into forming an ionic solid.
Lattice Energy equation modified from Coulomb’s Law
 k is a proportionality constant dependent upon the solid
 The higher the charge on each ion (i.e. F- vs O2-), the
greater the lattice energy will be.
 The smaller the the distance between the solid (r)
decreases the larger the LE
Formation of KCl ionic solid (salt) from
K(s) and Cl2(g)
1.
K(s) must form K(g) – Energy of Sublimation = +64kJ
2.
K(g) must form K+(g) – First Ionization Energy = +419 kJ
3.
½ Cl2 must form Cl(g) – Bond Energy for ½ mole =120 kJ
4.
Cl(g) must form Cl-(g) – Electron Affinity= - 349 kJ
5.
K+(g) must combine with Cl-(g) to form KCl(s) – Lattice Energy
= -690kJ
The net energy of formation equals the sum of the energy changes
= -436 kJ.
Creating Energy Diagrams
Application
Calculate the ΔHf° and Construct the Energy Diagram for the formation of
MgF2(s). (use previous example and page 342-345)
 Lattice Energy
-3916 kJ/mol
 First Ionization Energy of Mg
735 kJ/mol
 Second Ionization Energy of Mg
1445 kJ/mol
 Electron Affinity of F
-328 kJ/mol
(watch 2 moles!)
 Bond Energy of F2
154
kJ/mol
 Enthalpy of sublimation of Mg
150. kJ/mol
Covalent Bonding
 Electrons are SHARED
 Bonding models/representations
 Models are not reality
(limitations to models)
 Example: Methane CH4
Covalent Bonds
Two “types”
 Polar-Covalent Bonds
 UNEQUAL sharing of electrons
 Electronegativity difference between 0.3 and 1.7
*Recall the definition and trends of Electronegativity
 Non-polar-Covalent Bonds
 EQUAL sharing of electrons
 Electronegativity difference between 0 to 0.3
**Tug-of-war analogy**
Covalent Bonding Forces and Energy
 Figure 8.1 in text
Bond Analysis
A bond will form if the energy of the combination of atoms is
lower than that of the separated atoms.
 Bond length
 Distance between atoms where the energy is minimal.
 Bond strength
 Ionic; polar-covalent; non-polar covalent…
 Bond energy/stability (see previous diagram)
 Relationship between bond length and bond energy see next slide!
Bond Length and Energy Relationships
 Bond energy __________ and bond length _________.
 How does this relate to Coulomb’s Law?
 A triple bond has a ________bond length than a single
bond.
 The number of bonds between two atoms is _________
related to the length of the bond.
 The number of bonds between two atoms is
__________related to the bond energy.
Bonds between elements become shorter and stronger as
multiplicity increases.
Refresher: Bond Energy and Enthalpy
DH = SEBondsBroken - SEBondsFormed
Energy required
Energy released
 E= the bond energy per mole of bonds (found in table)
 Breaking bonds always requires energy (Endothermic)
 Forming bonds always releases energy (Exothermic)
The Octet Rule (Rule of 8)
 Combinations of elements tend to form so that each
atom, by gaining, losing, or sharing electrons, has an
octet of electrons in its highest occupied energy level.
 Lewis Structures of
 Monatomic Chlorine vs.
 Diatomic Chlorine
Octet Rule and Covalent Compounds
 Nonmetals, sharing electrons in highest energy level to
achieve access to an octet.
 Molecule is used exclusively for covalent bonding.
 The Diatomic Fluorine Molecule
 Orbital Notation and Lewis Structures to model
bonding.
Application
 Using orbital notation and Lewis structures
represent the Diatomic Oxygen Molecule.
 Using orbital notation and Lewis structures
represent the Diatomic Nitrogen Molecule.
Lewis Structures
Recall …
 Lewis structures show how valence electrons are
arranged among atoms in a molecule.
 Lewis structures reflect the central idea that stability of
a compound relates to noble gas electron configuration
(octet).
 Shared electron pairs are covalent bonds and can be
represented by two dots (:) or by a single line ( - )
Steps for Writing Lewis Structures
1. Sum the valence electrons from all the atoms.
Do not worry about keeping track of which electrons come from which
atoms. It is the total number of valence electrons contributed that is
important.
2. Use a pair of electrons to form a bond between each pair of bound
atoms. (Place atom wanting the most bonds in center or typically
least common element.)
3. Arrange the remaining electrons to satisfy the peripheral atoms
octet (for second row elements) or duet (for hydrogen) and then
place remaining electrons on central atom.
Model
 CH2Cl2
 H2CO
 NH3
 NH4+
Application
Draw Lewis Structures for the following molecules:
CSe2
ICN
AsF3
BH4-
Exceptions to the Octet Rule
 The second-row elements C, N, O, and F should always obey the octet
rule, will NEVER exceed octet. (Connection: 2s and 2p)
 The second-row elements B and Be often have fewer than 8 electrons in
their compounds. (Electron-deficient = very reactive)
 Third-row and heavier elements often satisfy the octet rule but can
exceed the octet rule by using their empty valence d orbitals.
 When writing Lewis structures satisfy the octet or duets first! Place
remaining valence electrons on the element(s) that has available d
orbitals (Period 3 and beyond).
 Localized Electron Model (LEM)
 A model which assumes that a molecule is composed of atoms that are
bound together by sharing pairs of electrons using the atomic orbitals of
the bound atom.
HONC Rule
Hydrogen (and Halogens) form one covalent bond.
Oxygen (and sulfur) form two covalent bonds.
 One double bond, or two single bonds
Nitrogen (and phosphorus) form three covalent bonds
 One triple bond, or three single bonds, or one double and one single
bond.
Carbon (and silicon) form four covalent bonds
 Two double bonds, four single bonds, a triple bond and a single bond,
or a double bond and two singles
Bond Polarity and Dipole Moments
 Recall: Polarity can be determined by the difference in
electronegativity.
 Result of polarity is a partial positive charge and partial negative
charge.
 The arrow points to the center of negative charge and the tail is at the
center of the positive charge.
 A dipole moment means that the molecule has two poles (two
opposite charges)
 Connection poles of a magnet
 2 atom molecules are simple to determine polarity.
 With 3 or more atoms in a molecule the individual dipoles can cancel
each other out. **Tug-of-war analogy**
Model- Formal Charge Discussion
 CO2
 XeF4
 IF2-1
Application
 SbCl3
 AlF63-
 PCl5
Individual Practice
 Ch. 8 Part 1
Starting on Page 382 in your textbook
 23, 25, 29, 35, 37, 38, 39, 41 (review),43, 45 (think about
what an energy diagram would look like/sketch), 53, 67,
68, 69, 71
Next…VSEPR!
 Valence
 Shell
 Electron
 Pair
 Repulsion
Geometries will need to be MEMORIZED!!