Chemical Reactions and Equations:
(CP:Ch.11, TP:ch. 7)
Clues to determine if substances undergo chemical changes:
1. Precipitation 2. Energy 3. Odor 4. Gas released
5. Color Change
Changes in compounds, reactions, can be shown by chemical equations.
Reactants and products are shown on in equations.
Other items shown in an equation are:
Physical state, catalysts, “amounts” of substances.
Reactants are the substances that undergo a change or reaction are on the left side of the equation.
Products are the new substances that are formed and are written on the right side of the reaction.
Changes or reactions can be shown by word equations and chemical equations.
Word Eqs. - Common names:
Vinegar + baking soda  sodium acetate + water + carbon dioxide
1

Word Eqs. – Scientific names:
Acetic Acid + sodium hydrogen carbonate

Sodium acetate + dihydrogen monoxide + carbon dioxide
Chemical Eqs.
–
HC
2
H
3
O
2
+ NaHCO
3
 NaC
2
H
3
O
2
+ H
2
O + CO
2
Now add physical states:
HC
2
H
3
O
2
(aq) + NaHCO
3
(s)

NaC
2
H
3
O
2
(s) + H
2
O(l) + CO
2
(g)
Balancing equations – Why? Law of Conservation of mass. Matter is neither created nor destroyed during a
Chemical reaction.
HC
2
H
3
O
2
(aq) + NaHCO
3
(s) 
NaC
2
H
3
O
2
(s) + H
2
O(l) + CO
2
(g)
List elements on both sides.
Reactants Products
H 5 H 5
C 3 C 3
O 5 O 5
Na 1 Na 1
Is it balanced? Yes!
2

IF not follow the rules.
1.
Determine the correct formulas for all of the reactants and products.
2. Write the reactants on the left, products on the right separated by a  .
3. Count the number of atoms of each element in the reactants and products. Polyatomic ions appearing unchanged can be counted as a single unit.
4. Balance the elements one at a time by using coefficients. DO NOT CHANGE SUBSCRIPTS OF A
CHEMICAL FORMULA TO BALANCE AN
EQUATION.
TIP: In some equations, balance hydrogen and oxygen last!!
5. Check each element or polyatomic ion on each side to make sure equation is balanced.
6. Finally, make sure all coefficients are in the lowest possible ratio that balances.
___H
2
(g) + __ O
2
(g)

___H
2
O(l)
_2_H
2
(g) + __ O
2
(g)

_2_H
2
O(l)
3

Try:
NaOH(aq) + CO
2
(g)

Na
2
CO
3
(s) + H
2
O(l)
Na O H C Na O H C
1 3 1 1 2 4 2 1
2 NaOH(aq) +_CO
2
(g)  __Na
2
CO
3
(s) + __H
2
O(l)
Independent Practice: TP, p. 203, 7.7 a,c
CP, p.
Standard 4.1 and 4.2
Types of Reactions:
(TP: p.211; CP: p. 338-339)
1. Combination (Synthesis)
2. Decomposition
3. Single-Replacement
4. Double-Replacement
5. Combustion
6. Neutralization (Acid/base): (not in chapter)
4

1. Combination (Synthesis)
(use students to show relationships)
R + S

RS
2 Mg(s) + O
2
(g)

2 MgO(s)
TIP: What do you notice about the product?
2. Decomposition
RS

R + S
2HgO(s)  2Hg(l) + O
2
(g)
TIP: What do you notice about the reactant?
3. Single-Replacement
T + RS  TS + R
2K(s) + 2H
2
O(l)

2KOH(aq) + H
2
(g)
(HOH)
Tip: Is the cation or anion changing places?
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4. Double-Replacement
R + S - + T + U - 
R + U + T + S -
K
2
CO
3
(aq) + BaCl
2
 2KCl(aq) + BaCO
3
(s)
5. Combustion
C x
H y
+ (x+y/4)O
2

xCO
2
+ (y/2)H
2
O
CH
4
(g) + 2O
2
(g)  CO
2
(g) + 2H
2
O(g) (complete)

CO (g) + 2H
2
O(g) (incomplete)
6. Neutralization (Acid/Base) – same as double replacement except the reactants are acids and bases.
HA + BOH  BA + H
2
O
(salt)
HCl + KOH  KCl + H
2
O
6

How do you know if a compound is an acid?
How do you know if a compound is a base?
What is the exception?
Standard 4.2
Predict the Product:
1. Acid/Base Neutralizations – H
2
O (water) and salt
Ex:
NaOH + HCl

NaCl + H
2
O
2. Combustion Reactions - Carbon dioxide/carbon monoxide AND water
Ex:
C
5
H
12
+ O
2

CO
2
+ H
2
O
7

Nature of Reactions:
(TP:sect )
Reversible? Some are and some are not. Examples:
Ice melting (physical change)
Car battery (chemical change)
Fuel burning (chemical change)
Paint hardening (chemical change)
Reversible reactions reach an equilibrium (no net change occurs in the amount of reactants and products in a system)
How do you know if a reaction is not reversible?
If a gas is release or an insoluble precipitate is formed.
Exothermic Reaction – the energy of the reactants is greater than the product’s energy. Energy is given off during a chemical reaction. The Δ Heat is negative.
Example
2 Al(s) + Fe
2
O
3
(s)  2 Fe (s) + Al
2
O
3
(s) + 850 kJ
Δ H = -850 kJ (negative means it loses heat)
8

Endothermic Reaction – the energy of the products is higher than the energy of the reactants. Energy is taken in for the reaction to occur. Δ Heat is positive.
Example:
2 CO
2
(g) + 570 kJ

2 CO (g) + O
2
(g)
Δ H = +570 kJ
Many factors can influence the rate and the direction that a reaction occurs.
Factor Page # Example (CP)
1. activation energy 413(TP) cooking 543
2. concentration 391 (TP) 545
3. stirring/collision theory Kool-Aid 542
4. catalysts (enzymes) 415 Mufflers 546
5. inhibitors 547
6. temperature 545
7. Limiting Reagents 369
8. Particle Size 546
Predicting a Precipitate: Rules p. 386 (TP)/p. 344 (CP)
1. Salts of alkali metals and ammonia are soluble
2. Nitrate salts and chlorate salts are soluble
3. Sulfate salts are soluble
4. Chloride Salts are soluble
5. Carbonates, phosphates, chromates, sulfides, and hydroxides - most are insoluble
9

EQ: Why are reactions rates important in everyday life?
Name any of the stated factors that affect the following items.
Food: all
Film: Temperature - cooler
Batteries: Temperature - cooler
Digestive Aids: Enzymes
Catalytic Converters:
CP ONLY
Reactions in aqueous (aq) solutions:
Complete Ionic Equations shows dissolved ionic compounds as their free ions.
(refer to p. 225cp)
Net Ionic Equation:
Write the equation that indicates only those particles that actually take part in the reaction.
Ag + (aq) + Cl (aq)  AgCl (s)
What are spectator ions?
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