1. Chemical Changes
Physical changes (Ch. 5) involve only changes in the physical form of the substance, not in the atomic or molecular make-up
Chemical changes involve conversion into new substances with new chemical properties
Chemical changes can often be observed:
Color change, precipitate (solid forms), bubbles, etc.
In a chemical reaction, reactants go to products
Atoms of reactants are recombined in products:
2Ag + S  Ag2S
(silver reacts with sulfur to form silver sulfide)

3. Chemical Equations Used to represent chemical reactions Like a recipe:
- Tells what you need to start with, and how much
- Also tells what you will make, and how much
Example: 2H2 + O2  2H2O
Number of each type of atom must be equal on the two sides of the equation
4 H’s O’s = 4 H’s O’s
Use coefficients to balance chemical equations
Sometimes symbols are used to show physical state:
(s) = solid, (l) = liquid, (g) = gas and (aq) = aqueous (in water)
Example: C(s) + O2(g)  CO2(g)

5. Balancing a Chemical Equation
Write correct formulas for reactants
Count atoms on both sides (is it balanced?)
Balance one element at a time (usually C first and O or H last, but can be any order)
Count atoms again to check that it’s balanced
Example: Propane (C3H8) burns with oxygen to form carbon dioxide and water. Write the balanced chemical equation.
C3H8 + O2  CO2 + H2O
3 C’s H’s O’s  1 C H’s O’s
C3H8 + O2  3CO2 + H2O
C3H8 + O2  3CO H
C3H O2  3CO H2O
3 C’s H’s O’s = 3 C’s H’s O’s

6. Types of Reactions
Reactions can be organized into 4 basic types: combination, decomposition, replacement and combustion
- Combination reactions: 2 (or more) reactants combine to form a single product
- Decomposition reactions: One reactant splits into 2 (or more) products
- Replacement reactions: Elements are exchanged between 2 reactants to form 2 products
- Combustion reactions: fuel + oxygen  products + heat
Reactions can be more than one type

8. Oxidation-Reduction (Redox) Reactions
Some reactions are also categorized as redox reactions
In these reactions the reactants exchange electrons
- Reduction = gain of electrons (GER)
- Oxidation = loss of electrons (LEO)
Oxidation and reductions reactions are always coupled
(electrons gained = electrons lost)
Example: Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
- Mg loses 2 electrons to become Mg2+ (Mg is oxidized)
- Each Cl gains an electron to become Cl- (Cl is reduced)
- H is not oxidized or reduced (no change in # of electrons)
Also, in general, gain of O or loss of H = oxidation and gain of H or loss of O = reduction (in biological systems)

10. Energy in Chemical Reactions
In order for a reaction to take place, the reactants must contact each other with enough energy
As reactants collide, bonds are broken and new bonds are formed
Example: 2H2 + O2  H2O (the H-H and O-O bonds break and two new O-H bonds are formed)
Between the reactants and the products there is a “transition state” in which bonds are breaking and/or forming (highest E point in reaction)
Energy required to reach transition state is called “activation energy” (EA)
Transition state is always highest E (higher than reactants and products) because it takes E to break bonds and E is released when bonds are formed

11. Exothermic and Endothermic Reactions
The difference in energy between the reactants and the products is called the “heat of reaction”
Heat of reaction can be heat released or heat consumed, depending on the reaction
Reactions that release heat are exothermic
CH O2  CO H2O + heat (213 kcal)
Reactions that consume heat are endothermic
H2 + I2 + heat (12 kcal)  2HI
Energy diagrams are used to show energy changes during a chemical reaction (E vs. reaction progress)

13. Reaction Rates
Reaction rate = how fast a reaction goes from reactants to products
Rate is based on activation energy and not on heat of reaction (lower EA = faster reaction)
Reaction rates are affected by such factors as :
- Reactant concentration (more reactants = more collisions = faster reaction)
- Temperature (at higher T reactants collide more often at higher E = faster reaction)
- Catalyst (addition of a catalyst lowers the activation energy = faster reaction)
- a catalyst makes the transition state more stable, so it takes less EA to reach it

15. Chemical Equilibrium
Some chemical reactions are reversible (products can also go to reactants)
Example: N2(g) + O2(g)  2NO(g)
- Forward reaction = N2 + O2  2NO
- Reverse reaction = 2NO  N2 + O2
When rate of forward reaction = rate of reverse reaction, chemical equilibrium has been reached
When at equilibrium:
- If more products exist in reaction mixture, then reaction favors products
- If more reactants exist in reaction mixture, then reaction favors reactants

19. LeChâtelier’s Principle
The equilibrium can be shifted towards more products or more reactants by placing a “stress” on the system
Add reactants or remove products and equilibrium is shifted towards products
Add products or remove reactants and equilibrium is shifted towards reactants
Heat is also considered a reactant (endothermic reactions) or a product (exothermic reactions)
Example: C(s) + H2O(g) + heat  CO(g) + H2(g)
- Add heat: equilibrium shifts towards products
- Remove H2(g): equilibrium shifts towards products
- Remove H2O(g): equilibrium shifts towards reactants