1. Chapter 11
Chemical Reactions

2. Physical Changes
Physical Change – A change during which some properties of a material change, but the composition of the material does not change.
Examples:
- Cutting, breaking, grinding, tearing
- phase changes (melting, boiling, freezing)
- dissolving

3. Chemical Changes
Chemical change – A change in which a substance undergoes a change in identity (composition).
Examples: burning, rusting, decomposing
Things that are NOT chemical changes:
Any phase change (evaporation, freezing…)
Dissolving

4. Signs of Chemical Changes
Several signs that a chemical change has occurred:
1) Heat, light or gas given off
2) Change in color or odor
3) A precipitate (solid product) is formed
(It will look cloudy.)
4) Sound
5) Bubbles (except boiling)

5. When signs of a chemical change can be deceiving.
Boiling - phase change from liquid to gas.
It is NOT a reaction. But you will see bubbles.
Diluting – adding water to lower the concentration. It will alter the color, but it is NOT a reaction.
Diluting is NOT a chemical change.
Two clear liquids making a yellow precipitate
IS a chemical reaction

6. Chemical Reactions
Chemical Reaction – A chemical change in which new substances are formed.
Chemical Equation – A symbolic representation of a chemical reaction.

7. Reaction Terms Reactants: Starting substances in a chemical reaction.
A + B  C + D
Products: Substances produced in a chemical reaction.

8. Reaction Terms Coefficient – Big number in front of a substance.
Tells the number of moles, molecules, or units of a substance.
3H2O
Subscript – Small number written below.
Tells the numbers of moles or atoms of a particular element.

9. Reaction Symbols  “yields” (produces, results in)
 reversible reaction, “equilibrium”
(s) solid
(l) liquid
(g) gas
(aq) aqueous – dissolved in water
reaction occurred by heating

10. Collision Theory
1) Molecules must collide in order to react. 2) When they collide, they have to have a) The right orientation b) Enough energy

11. Law of Conservation of Mass
Law of Conservation of Mass – Matter cannot be created or destroyed.
Every chemical equation must satisfy this law.
Therefore equations must be balanced so that:
total mass of reactants = total mass of products.

12. Balancing Equations Only COEFFICIENTS can be changed !
(NEVER the subscripts!)
Example:
Al(s) CuSO4(aq)  Cu(s) Al2(SO4)3(aq)

13. Balancing Examples H2 + O2  H2O Zn(OH)2 + H3PO4  Zn3(PO4)2 + H2O
Ag2S Al  Al2S Ag
_____ Na + _____ I2 _____ NaI

14. More Balancing Examples
___ Ca3(PO4)2 + ___ H2SO4 ___ CaSO4 + ___ H3PO4
___ KClO3  ___ KCl + ___ O2
SO O2  SO3
C3H O2  CO H2O

15. Types of Reactions Synthesis – 2 or more reactants, 1 product
Example: A + B  AB

16. Types of Reactions Decomposition: 1 reactant, 2 or more products
Example: AB  A + B

17. Types of Reactions
Single Replacement - a single element “switches places” with an element in a compound.
D BC  C BD
(nonmetal) (nonmetal)
A BC  B AC
(metal) (metal)

18. Types of Reactions
Double Replacement – Two compounds “switch partners”.
Don’t forget that in a compound the charges must be (+)(-).
Example: AB + CD 
AD + CB

19. Types of Reactions Combustion – Special type of reaction
CxHy + O2  CO2 + H2O
Watch for CO2 and H2O as products!

20. Predicting the Products – Single Replacement
To predict if a single replacement will react, use the ACTIVITY SERIES.
The free element MUST BE more reactive than the element in the compound.

21. Single Rep.
If the free element is MORE ACTIVE than the element in the compound, the reaction WILL
OCCUR.
Example: Cl2 + 2HBr 
If the free element is LESS ACTIVE than the element in the compound, the reaction WILL
NOT OCCUR.
Example: Br2 + 2HCl 
(DNR = Does Not React)
2HCl + Br2
DNR

22. Note
**Note – Don’t EVER bring a subscript across the arrow UNLESS it’s part of a polyatomic ion!!**

23. Activity Series Examples
Ag + ZnCl2 
Zn + 2AgNO3 
Zn + CuSO4 
NaCl + Li 
NaCl + I2 
LiOH + Na 
2NaI + Br2 
DNR
2Ag + Zn(NO3)2
Cu + ZnSO4
LiCl + Na
DNR
DNR
2NaBr + I2

24. Predicting the Products – Double Replacement
In order for a double replacement reaction to occur, one product MUST BE:
A gas
A liquid (water)
A precipitate (a solid)
If all products are aqueous (soluble), it WILL NOT occur.
Use the SOLUBILITY RULES:
Soluble = Aqueous
Insoluble = Solid precipitate

25. Solubility Rules

26. Double Replacement Examples
Ba(OH)2 + H3PO4 
K2SO4 + CaCl2 
FeBr2 + AlCl3
KOH + H2SO4 
(NH4)2CO3 + CaCl2 

27. Net Ionic Equations Rules:
Aqueous ionic compounds can be split into ions. (Don’t forget charges!)
Strong acids can be separated into ions.
Substances that are solids, liquids, or gases cannot be separated.
Spectator ions are removed from the ionic equation, leaving the net ionic equation.

28. If Aqueous… For all aqueous compounds: Step 1 – Split it up
Step 2 – Write charges for each thing
Step 3 – Write how many you have as a coefficient

29. Net Ionic Eqn. Example
Na2SO4 (aq) + BaCl2 (aq)  2NaCl (aq) + BaSO4 (s)
Ionic equation:
Spectator ions:
Net Ionic Equation:

30. Lab Tests – Burning splint
A burning splint can be used to test for:
Hydrogen (squeaky “pop” sound)
Oxygen (blow out the splint and it will reignite)
Because fire needs O2 to burn
Carbon dioxide (flame will go out)
Because CO2 smothers it

31. Lab test - Limewater and CO2
Clear, colorless limewater will turn a cloudy white if CO2 is added.