Chemical Bonding

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-

Q: How do we know whether ionic or covalent?

A: Compare electronegativities of the atoms.

In general:

Small difference = covalent bondMedium difference = polar covalentLarge difference = ionic bond

Covalent Bonding & Molecular Compounds

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

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

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

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

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

Drawing Lewis Structures

1. Determine type and number of atoms

2. Find total number of valence electrons

3. Arrange atoms with carbon in the middle; if no carbon present, least electronegative element in the middle; hydrogen is never the central atom

4. Add unshared pairs of electrons forming octets

Ex – NH3, H2S, CO2, HCN

Resonance Structures

Bonding in molecules or ions that cannot be correctly represented by a single Lewis structure

Ex – ozone

Ionic Bonding & Ionic Compounds

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)

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

Ionic Bonding

Electrostatic forces hold ionic compounds together

Makes ionic compound very hard but brittle

Strong attractions raise boiling and melting points

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

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

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

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

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

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