1. Chapter 10
Chemical Reactions

2. Evidence of a Chemical Reaction
Energy release (heat or light)

3. Evidence of a Chemical Reaction
Color change

4. Evidence of a Chemical Reaction
Odor
gas bubbles
smoke formation

5. Evidence of a Chemical Reaction
Precipitation

6. Symbols used in Equations
+ separates two reactants or two products
→ “Yields” separated reactants from products
Designates a reactant in the solid state
Designates a reactant in the liquid state
(g) Designates a reactant in the gaseous state
(aq) Designates a reactant in the aqueous state
→ → indicates that heat is supplied to a reaction
→ Indicates that a catalyst is used (here, platinum)
heat Pt

7. Representing Chemical Equations
reactant 1 + reactant 2 → product 1 + product 2
produces
Product(s) must be different than the reactants

8. Representing Chemical Equations
Simply a “recipe” →
Eggs and butter and flour produce a cake
Word equations name the reactants and products.
What is missing from this recipe?
AMOUNTS!!!

9. Representing Chemical Equations→
Iron (II) and oxygen produce iron oxide
Fe (s) + O2 (g) → FeO (s)
What is missing?
Amounts!
How do you figure out the amounts needed and produced?

10. Balancing Equations Based on the Law of Conservation of Mass
“Matter is not created or destroyed in a reaction, just rearranged.”
Fe + O2 → Fe O
What is wrong here?
an oxygen atom is missing

11. Balancing Equations Atom inventory Fe + O2 → Fe O
Reactant side Product side
Fe O Fe O
Fe + O2 → Fe O
Final equation: Fe + O2 → Fe O 2 2 2 2 2 2

12. Balancing Equations Steps:
1. Write the skeleton equation (can’t escape formula writing)
2. Count atoms in reactants
3. Count atoms in products
4. Change coefficients to make atom counts match. Never change subscripts.
5. Reduce coefficients if necessary.

13. Practice Problems ____ AlBr3 + ____ K→____ KBr + ____ Al
____ P4 + ____ Br2 → ____ PBr3
____ P Br2 → 4 PBr3

14. C6H8O7(aq) + 3NaHCO3(aq) → 3H2O(l) + 3CO2(g) + Na3C6H5O7(aq)
citric acid + baking soda → water + carbon dioxide + sodium citrate
1. Calculate the molar mass for each compound in the equation.
2. If there is a coefficient, multiply the molar mass by that number.
3. Add the compounds on the left of the arrow.
4. Add the compounds on the right of the arrow.
5. How are the two sums related?

15. Symbols used in Equations
#1 separates two reactants or two products
#2 “Yields” separated reactants from products
#3 Designates a reactant in the solid state
Designates a reactant in the # 4 state
(g) Designates a reactant in the # 5 state
#6 Designates a reactant in the aqueous state
# 7 indicates that heat is supplied to a reaction
→ Indicates that a # 8 is used (here, platinum)
# 9 name the diatomic elements Pt

16. Classifying Chemical Reactions
Five general types
Synthesis one new product
Solid sodium and chlorine gas form solid sodium chloride 2Na(s) + Cl2(g) → 2NaCl(s) →

17. Classifying Chemical Reactions
Combustion Must have O2 as a reactant
2 Mg + O2 → 2 MgO →
Combustion reactions are exothermic; energy is produced.

18. Classifying Chemical Reactions
Combustion of hydrocarbons (compounds containing hydrogen and carbon)
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
Most fuels are hydrocarbons

19. Classifying Chemical Reactions
Decomposition Only one reactant
Hydrogen peroxide decomposes to release oxygen gas and leave water on your skin
H2O2(aq) → O2 (g) + H2O(l)
2 H2O2(aq) → O2 (g) + 2 H2O(l)

20. Classifying Chemical Reactions
Single replacement reactions One reactant is an element
Copper reacts with silver nitrate to produce silver and copper nitrate.
Cu(s) + AgNO3(aq) → Ag(s) + Cu(NO3)2(aq)

21. Single replacement reactions
Activity Series of Metals
Lithium Iron
Rubidium Nickel
Potassium Tin
Calcium Lead
Sodium Copper
Magnesium Silver
Aluminum Platinum
Manganese Gold
Zinc
Decreases
Decreases

22. Single replacement reactions
Halogens
Fluorine
Chlorine
Bromine
Iodine
Decreases

23. Classifying Chemical Reactions
Double replacement reactions both reactants are ionic compounds
A gas, liquid, or solid must be produced
Silver nitrate and sodium chloride react to produce sodium nitrate and silver chloride
AgNO3 (aq) + Na Cl (aq) → NaNO3 (aq) + AgCl (s)
Ag+
Cl-
Na NO3-
Na NO3-
AgCl

24. Summary Reaction Classification
Class of Reaction
Reactants
Probable products
Synthesis
Two or more substances
One compound
Combustion
Some substance and oxygen
An oxide
H20 and CO2 for hydrocarbon combustion
Decomposition
Two or more (elements or compounds)
Single-replacement
An element and a compound
A new compound and the replaced element
Double-replacement
Two compounds
Two different compounds including a s, l, g

25. Reactions in Aqueous Solutions
Review:
Solution – homogeneous mixture
Solute – substance dissolved
Solvent – most plentiful substance in the solution
Water is the “universal solvent”
Other substances can also be solvents

26. Reactions in Aqueous Solutions
How does water dissolve ionic compounds? Water is a polar molecule.

27. Reactions in Aqueous Solutions Reactions that form solids
Lead nitrate and potassium iodide
Two clear solutions – mix – precipitate
forms
Pb +2
NO3 - K+ I-

28. Reactions in Aqueous Solutions Reactions that form solids
What happened on an atomic level? →
Lead and iodide form a new compound
Nitrate and potassium are still dissolved ions – they did nothing but “watch” – spectator ions
NO3 –
Pb +2 K+ I- K+
NO3 –
PbI2

29. Reactions in Aqueous Solutions Reactions that form solids
Complete ionic equation
Pb NO3 - + K+ + I- → PbI2 + NO3 - + K+
Spectator ions
Net ionic equation
Pb I- → PbI2

30. Reactions in Aqueous Solutions Reactions that form liquids
HBr (aq) + NaOH(aq) → H2O (l) + NaBr(aq)
Complete ionic equation
.H+(aq) + Br-(aq) + Na+(aq) + OH -(aq) → H2O(l) + Na+(aq) + Br-(aq)
Net ionic equation
H+(aq) + OH -(aq) → H2O(l)

31. Reactions in Aqueous Solutions Reactions that form gases
2HI (aq) + Li2S (aq) → H2S (g) + 2LiI (aq)
Complete ionic equation
2H+ (aq) + 2I- (aq) + 2Li+ (aq) + S2-(aq) → H2S(g) + 2Li+(aq) + 2I- (aq)
Net ionic equation
2H+ (aq) + S2-(aq) → H2S(g)