1. Kinetics and Equilibrium

2. Collision Theory
Basic concept of Kinetics- in order for a reaction to occur reactant particles must collide.
Particles must collide for a reaction to occur, but only a fraction of collisions have the correct factors such as angle to connect effectively to cause reactants to transform into products.
The amount of energy needed for this to occur is called the activation energy.
Certain factors affect the rate of reaction (collisions taking place)

3. Factors affecting rate of reaction
Nature of the reactants- Covalently bonded compounds take longer (more bonds to break) Than ionic bonds.
Concentration- increased amounts- faster rate
Surface Area- more surface area, greater chance of collision (better to have a fine powder than a lump of the same mass)

4. Pressure- only effects rates of reactions for gases- higher pressure more collisions, faster rate.
The presence of a catalyst- Catalyst is a substance that increases the rate of reaction, by providing a different/easier pathway for reaction to occur- catalyst stays unchanged.
Temperature-high temp, faster moving particles, more collisions-faster rate of reaction.

5. Enthalpy-also called the heat of reaction
Is equal to the heat lost or gained by a system under contant pressure.
∆H= H products – H reactants
-∆H means energy is released- the reactants have a higher energy than the products-exothermic.
+ ∆H means energy is absorbed- the reactants have a lower energy than the products- Endothermic

6. Reversible Reactions
A reaction that can occur in both the forward and reverse directions.
Reactants can form products but the products can also combine to form reactants.
We use a double headed arrow to show this.
Fe3O4(s) + 4H2(g)  3Fe(s) + 4H2O(g)
For any reaction the reverse reaction has an opposite sign for ΔH.