Chemical Reactions

1. Chemical Reactions

2. Rates of Reactions • Not all chemical reactions occur at the same speed. • Rusting of Iron very slow, measured in weeks or months • Hydrogen reacting with Oxygen to form water, very quick, measured in nano or pico seconds • Defined as the rate of disappearance of a reactant or as the appearance of a product

3. Collision Theory • In order for chemicals to react the molecules must touch (collide) with each other with enough energy to break old bonds and build new bonds. They must also collide with the correct orientation. Example: head to head collision, or side to side

4. 4 factors (variables) that affect rate • 1) Nature of the reactant • Examples: • Size of particles (smaller particles react faster than larger chunks) Smaller particles have larger surface areas results in more opportunities for collisions • Shape of the molecule

5. 2) concentration • The more concentrated the solution the faster the rate of reaction • Because more molecules are crammed in the same amount of space, increasing collisions

6. 3) Temperature • Higher temperatures result in faster reactions • Lowering the temperature slows down the rate of reaction • Because – at higher temperatures molecules are moving faster, more collisions and collisions that are harder (more forceful)

7. 4) Catalysts • Catalysis – chemicals that speed up a reaction but do not become part of the products • Not all reactions are affected by catalysts but for those who are affected, adding a catalyst tremendously speeds up the reaction • Because the catalyst provides a surface for molecules to meet (collide)

8. Energy of reactions • Some chemical reactions require additional energy to occur (Endothermic Reactions) • Chemical Cold packs are an example • Some chemical reactions release additional energy (Exothermic Reactions) • Chemical hot packs are an example