1. Chapter 7 and 8: CHEMICAL BONDS
1. Ionic bonds
Today, we will learn about:
2. Metallic bonds
3. Covalent bonds
Introduction
Scale of electron sharing in bonds
3 types of covalent bonds (single, double and triple)
Polar covalent bonds and electronegativity
Non-polar covalent bonds
Coordinate covalent bonds
Examples of diatomic molecules and other common molecules with covalent bonds
Bond dissociation energy
VSPER theory and molecular shapes
Chemistry _ Notes
Dr. Chirie Sumanasekera
11/28/ 2017

2. 2. Metallic Bonds
Metallic bonds =Inside metals, positively charged (cations) swim in a sea of valence electrons. These electrons are shared by all neighboring atoms.
Metals are good electrical conductors as they have readily mobile electrons to produce a current
Metals form compact cubic crystals
The shared-electron mobility is also why metals are malleable and ductile
Alloys = mixtures of two or more metals (Ex: sterling silver, cast iron, stainless steel, surgical steel)

3. 3. Covalent bonds INTRO
Covalent bonds are found in molecular compounds (compounds that form molecules)
In covalent bonds, electrons are usually shared between non metal atom so that each can attain the noble gas electron configuration
There can be single, double or triple covalent bonds between two atoms
Based on the electronegativity values of each atom in a covalent bond, the resulting molecule may have a dipole (polarity based on negative and positive charged regions).
Covalent bonds and ionic bonds exist in a spectrum going from zero to extreme electronegativity

4. Types of Covalent Bonds
Sigma (Single) bond Cl H Cl
HCl H + H Cl
Double bond O O C O + C
CO2 2 O C O
Triple bond N2 N + N N N N N

5. Scale of Valence electron sharing Similar electronegativity
Large difference in electronegativity
Similar electronegativity
Unequal
Equal
Non-polar
Covalent bonds
Polar
Covalent
bonds
Ionic bonds

6. Covalent Bonds O
H2O 2 H + O H H H O H

7. Water: Polar Covalent Bonds
 +
 -
Electron Density
shifts towards the more electronegative oxygen atom O H O H O H
Oxygen is much more electronegative than Hydrogen, so Oxygen pulls the shared electrons towards itself
This greater electronegativity of oxygen compared to hydrogen in the covalent bonds of water molecules result in a dipole or polarity. The result is a shifting of the electron density towards Oxygen.
H gets a small positive charge Delta (+) and Oxygen gets a negative charge ( -) as electrons are less equally shared between O and H.

8. Nonpolar covalent bonds+ N N N N2 N N
If a molecule if formed by the same kind of atom or atoms with similar electronegativity then the covalent bond is non-polar (Ex: O2, N2, Cl2, O3, Br2)
Carbon and H makes non polar covalent molecules such as hydrocarbons (CH4 , C2H6 , C3H8 and so on.)

9. Coordinate covalent bonds
Both these electrons come from Oxygen
In typical covalent bond, each atom in the bond provides an electron for a single bond. But in coordinate covalent bonds this is not the case. Here, one atom contributes both of the shared electrons for one covalent bond.

10. Molecules: formed by covalent bonding (electron sharing) between atoms
Ionic bonded compounds: Do not form molecules. They are formed by atoms that stick close due to opposite charge (electrostatic) - attractions

11. **Bond length decreases when number of bonds increase
**Bond Strength increases when number of bonds increase

13. Electronegativity difference & bond types: Table 8.3 (p238)
Equal electron sharing
Unequal electron sharing
Increasing electronegativity
Non-polar
covalent bonds
Polar covalent bonds
Ionic bonds
2.0 and greater
Electronegativity difference:
Examples: H H Cl
Na+
Cl- H F

14. VSPER: Valence-Shell Electron Pair Repulsion theory:
VSPER Theory
VSPER: Valence-Shell Electron Pair Repulsion theory:
Repulsion between valence shell electron pairs cause molecular shapes to change so that valence electron pairs stay as far apart as possible
Three bonds
Four bonds
Five bonds
Six bonds
Two bonds
Pyramidal
Tetrahedral
Trigonal bipyramidal
Octahedral
Bent triatomic
180o
Trigonal planar
Linear triatomic
Square planar

15. O O H H Effect of unpaired electrons on molecular shape
Repulsion between unpaired electrons also effect 3D molecular shapes: Ex. H2O
In water, the two lone pair electron repulsion and the electron clouds of these electron pairs cause the molecule to have a tetrahedral rather than a triatomic shape.
Repulsive force between unpaired electrons O H O H

16. Diatomic Elements
These are elements that exist in nature as molecules made up of two identical atoms. These atoms have the same electronegativity so these molecules are non-polar!
Examples:
Hydrogen gas (H2)
Oxygen gas (O2)
Nitrogen gas (N2)
Gasses of Halogens (Cl2, Br2, I2)
***The subscript 2 is used to show the two identical atoms in the molecule
Side Note: this is very different from mono-atomic gases. All of the noble gases are mono atomic gases. They are not molecules. They are single atoms that exist as a gas

17. Common molecular compounds (Table 8.2, Page 224)
Nitrogen dioxide
NO2
Hydrogen peroxide
Name
Chemical formula
H2O2
Hydrochloric acid
Hydrogen cyanide
HCl
SO3
HCN
Sulfur trioxide
Nitrous Oxide
N2O
CH4
Methane
What do all these chemical formulas have in common based on the types of elements
You see? (Clue: what types of elements are they? Find their location in Periodic Table…)

18. Helpful videos

19. Increasing Bond strength (kJ/mol)
Covalent Bond Strength and Bond Dissociation Energy
Increasing Bond strength (kJ/mol)
Single bond
Double bond
Triple bond
When two H-atoms combine to form H2 gas by covalent bonding, a large amount of heat is released (435 kJ). To break this covalent bond, the same amount of energy is required. The energy required to break a covalent bond = Bond Dissociation Energy.
Bond dissociation energy depends on which molecules are involved the bonding and how many bonds are to be broken. If more electronegative atoms are present, the energy will be higher. It takes more energy to break a triple bond than a single bond. F H
+ 565 kJ of energy
kJ of energy