1. Chemical Bonding

2. Types of Chemical Bonds
Ionic Bonding – bonding as a result of atoms giving away or receiving e-; attraction between cations and anions
Covalent Bonding – bonding as a result of atoms sharing e-

3. Q: How do we know whether ionic or covalent?
A: Compare electronegativities of the atoms.
In general:
Small difference = covalent bond
Medium difference = polar covalent
Large difference = ionic bond

4. Covalent Bonding & Molecular Compounds

5. Covalent Bonding & Molecular Compounds
Molecule – a neutral group of atoms held together by covalent bonds
Molecular Compound – a compound whose simplest units are molecules
Molecular Formula – shows the type and number of atoms present
Diatomic molecule – molecule with two atoms

6. Forming Covalent Bonds
Atoms almost always favor an arrangement with lower potential energy
As atoms get closer, the positively charged nucleus is attracted to the negatively charged electron cloud

7. Covalent Bond Characteristics
Bond length – the average distance between two bonded atoms at minimum potential energy
Bond energy – the amount of energy needed to break a chemical bond; measured in kJ/mol
Overlapping orbitals

8. The Octet Rule
Chemical compounds tend to form so that each atom has an octet of electrons in its highest energy level
Exceptions to the rule to exist; called expanded octets

9. Electron Dot Notation & Lewis Structures
Electron configuration showing only the valence electrons around the chemical symbol
Unshared pair (lone pair) – a pair of electrons NOT involved in bonding and belong only to one atom

10. Drawing Lewis Structures
Determine type and number of atoms
Find total number of valence electrons
Arrange atoms with carbon in the middle; if no carbon present, least electronegative element in the middle; hydrogen is never the central atom
Add unshared pairs of electrons forming octets
Ex – NH3, H2S, CO2, HCN

11. Resonance Structures
Bonding in molecules or ions that cannot be correctly represented by a single Lewis structure
Ex – ozone

12. Ionic Bonding & Ionic Compounds

13. Ionic Compounds
Ionic compounds – made of positive and negative ions combined so the charges are equal (zero)
Formula unit – simplest collection of atoms wich form an ionic compound’s formula
Ex – sodium chloride (NaCl)

14. Ionic Bonding
Involve a transfer of electrons making one atom positively charged and the other negatively charged (cations and anions)
Form a crystal lattice structure which balances the charges of the ions
Lattice energy – amount of energy released when 1 mole of an ionic compound is formed from gaseous ions

15. Ionic Bonding Electrostatic forces hold ionic compounds together
Makes ionic compound very hard but brittle
Strong attractions raise boiling and melting points

16. Polyatomic Ions Charged group of covalently bonded atoms
Charge is on the collective group; not any particular atom
Caused by an excess or shortage of electrons
Lewis structures are written in brackets

17. Metallic Bonding
The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons
Outer-most electrons are free to move between overlapping orbitals
This accounts for metals ability to conduct heat and electricity

18. Metallic Bonding Can absorb a wide range of light wavelengths
Immediately lose energy and fall to lower level
Accounts for metals shiny appearance
Bonding is the same in all directions
Accounts for properties of malleability and ductility

19. Intermolecular Forces (Dipole-Dipole)
Created by equal but opposite charges separated by short distances
Strongest intermolecular force
Represented by an arrow in the negative direction
Polar molecules have dipole moments

20. Intermolecular Forces (Hydrogen Bonding)
The intermolecular force in which a hydrogen atom is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule
Creates a fairly strong dipole-dipole force

21. Intermolecular Forces (London Dispersion Forces)
Slight attraction between molecules caused by instantaneous dipole moments
Weakest intermolecular force
Only force acting between noble gases and non-polar molecules