1. Chapter 10 Chemical Reactions
Chemistry 1 – Notes #2b
Chapter 10
Chemical Reactions

2. 10.1 Reactions and Equations
Chemical Reaction – process by which atoms are rearranged to form different substances.
How do you know when a chemical reaction has occurred?
Evidence
temperature change
color change
odor
bubbles (gasses are being formed)
solid forming

3. Representing Chemical Reactions
Reactant – starting substances
Products – substances formed during reaction
Symbol
Meaning +
Separate two or more reactants or products = “reacts with”
Separates reactants from products = “to produce”
(s)
Identifies solid state
(l)
Identifies liquid state
(g)
Identifies gaseous state
(aq)
Identifies water soluntion

4. Representing Chemical Reactions
reactant1 + reactant product1 + product 2
word equations:
iron(s) + chlorine(g) iron(III)chloride(s)
skeleton equations
Fe(s) + Cl2(g) FeCl3(s)
Practice page 279

5. Balancing Chemical Reactions
Law of conservation of mass
Steps for balancing are on page 281 in the textbook
(Metals First, then non-metals, save H & O for last)
Only Use coefficients for balancing equations
Balancing Equations is like solving a puzzle

6. 10.2 Classifying Chemical Reactions
5 types of chemical reactions
Synthesis
Combustion
Decomposition
Single Replacement
Double Replacement
some reactions may fit into more than one category

7. Synthesis Two or more substances react to produce a single product.
A + B  AB
These are some examples:
K + Cl2 ---> KCl N2 + H2 ---> NH3
CaO + H2O ---> Ca(OH)2

8. Combustion
Oxygen combines with a substance and releases energy in the form of light
The reactions can also be classified as one of the other types of reactions.
These are some examples:
C + O2 ---> CO2 (also synthesis)
2H2 + O2 ---> 2H2O (also synthesis)
CH4 + 2O2 ---> CO2 + 2H2O (also replacement)

9. Decomposition
A single compound breaks down into two or more elements or new compounds
Decomposition of water
often requires an energy source
AB  A + B
These are some examples:
2 H2O ---> 2 H2 + O2
2 NaN3 ---> 2 Na + 3N2

10. Single Replacement
Atoms of one element replace the atoms of another element in a compound
A + BX  AX + B
These are some examples:
2 Li + 2 H2O ---> 2 LiOH + H2
Cu + 2 AgNO3 ---> 2 Ag + Cu(NO3)2
See Fig on page 288

11. Double Replacement An exchange between two compounds AX + BY  AY + BX
Double Replacement: Evolution of a Gas
When a solid is produced during a chemical reaction in solution, the solid is called a precipitate
These are some examples:
Ca(OH)2 + 2 HCl ---> CaCl H2O
2 NaOH + CuCl2 ---> 2 NaCl + Cu(OH)2
See Table 10-2 on page 290

12. 10.3 Reactions in Aqueous Solutions
Vocabulary to Know
Solution – homogeneous mixture
Solute – a substance dissolved in a solution
Solvent – the substance that dissolves a solute to form a solution
Aqueous solution – a solution in which the solvent is water.

13. 2NaOH(aq) + CuCl2(aq)  2NaCl(aq) + Cu(OH)2 (s)
Complete ion equation – equation that shows all of the particles in a solution as they realistically exist.
2Na+(aq) + 2OH-(aq) + Cu2+(aq) + 2Cl-(aq)
2Na+(aq) + 2Cl-(aq) +Cu(OH)2(s)
Spectator ions – ions that do not participate in the reaction
Net ionic equation
becomes
2OH-(aq) + Cu2+(aq)  Cu(OH)2(s)

14. Double-displacement reactions in aqueous solutions
Reactions that form precipitates
write the complete ionic and net ionic equations for
KI(aq) + AgNO3(aq)  KNO3(aq) + AgI(s)
Reactions that form water
HBr(aq) + NaOH(aq)  H2O(l) + NaBr(aq)
Reactions that form gases
2HI(aq) + Li2S(aq)  H2S(g) + 2LiI(aq)