Chemistry 6.0

A. Definition: a process by which 1 or more substances, called reactants, are changed into 1 or more substances, called products, with different physical & chemical properties.
B. Evidence of a Chemical Reaction
1. Color change
2. Formation of a precipitate, ppt
3. Release of a gas
4. Energy change – heat, light, sound
5. Odor change
C. Reactions are started by the addition of energy

A. Form
1. Reactant + Reactant  Product + Product
2. Symbols: (s), (l), (g), (aq)
NR methanol

Interpretation of a
Balanced Equation
2Mg
(s)
+ O
2(g)
 2MgO
(s)
2 atoms of solid magnesium react with
1 molecule of oxygen gas to form
2 formula units of solid magnesium oxide
OR
2 moles of solid magnesium react with
1 moles of oxygen gas to form
2 moles of solid magnesium oxide

1. Exothermic reactions – release energy; energy a product
H
2
O
(g)
 H
2
O
(l)
+ 40.7 kJ
2. Endothermic reactions – absorbs energy; energy a reactant
H
2
O
(s)
+ 6.0 kJ  H
2
O
(l)

C. Characteristics of A Balanced
Chemical Equations
1. The equation must represent known facts. All substances have been identified.
2. The equation must contain the correct symbols and/or formulas for the reactants and products
3. Can be either a word equation or a formula equation
4. The law of conservation of mass must be satisfied.
This provides the basis for balancing chemical equations. 1 st formulated by Antoine Lavoisier
TOTAL MASS REACTANTS = TOTAL MASS PRODUCTS
Number of atoms of EACH element is the SAME on both sides of the equation.

1. Balance using coefficients after correct formulas are written.
Coefficients are usually the smallest whole number – required when interpreted at the molecular level
2. Balance atoms one at a time
3. Balance the atoms that are combined and appear only once on each side.
4. Balance polyatomics that appear on both sides
5. Balance H and O atoms last
NEVER CHANGE SUBSCRIPTS!!!
**Count atoms to be sure that the equation is balanced**

1. sodium + chlorine  sodium chloride
2. CH
3. K
(s)
4 (g)
+ O
2 (g)
 CO
2 (g)
+ H
2
O
(l)
+ H
2
O
(l)
 KOH
(aq)
+ H
2(g)
4. AgNO
3(aq)
+ Cu
(s)
 Cu(NO
3
)
2(aq)
+ Ag
(s)

• Two or more substances combine to form a more complex product.
A + B → AB (only ONE PRODUCT )
• A.K.A. Direct Combination Reactions, or composition reactions
• Ex. 8Fe + S
8
→ 8FeS
• Ex. CaO + H
2
O → Ca(OH)
2

2
2
 2

• Single Reactant breaks down to into a simpler substance.
AB → A + B (only ONE REACTANT )
• The opposite of a synthesis reaction.
• Ex. 2HgO → 2Hg + O
2
• Ex. CaCO
3
→ CaO + CO
2

• Atoms of one element replace atoms of another element in a compound.
A + BX → AX + B
• A more active element will replace a less active element. (See activity series)
• Ex. Fe + CuSO
4
• Ex. Mg + CuSO
4
→ FeSO
4
→ MgSO
4
+ Cu
+ Cu

• Atoms or ions from 2 different compounds replace each other.
AX + BY → AY + BX
• Ex. CaSO
4
+ 2HCl → CaCl
2
+ H
2
SO
4

• One substance reacts with oxygen to produce oxide compounds.
• Occurs when burning.
• Combustion reactions are often classified as synthesis reactions.
• These reactions are usually exothermic, releasing a large amount of energy as light, heat, or sound.

• When a hydrocarbon is involved in a combustion reaction, H
2
O and CO
2 are the products.
• Ex. CH
4
+ 2O
2
→ CO
2
+ 2H
2
O + 803 kJ
• Ex. S
8
+ 8O
2
→ 8SO
2

Pattern for prediction based on the kind of reactants
A. Combustion or Burning – complete combustion always produces carbon dioxide and water !
1. Hydrocarbons
C x
H y
+ O
2
2. Alcohols
 CO
2
C x
H y
OH + O
2
3. Sugars
+ H
2
O
 CO
2
+ H
2
O
C
6
H
C
12
12
O
6
H
22
O
11
+ O
2
+ O
2
 CO
2
+ H
2
O
 CO
2
+ H
2
O

2/more reactants  1 product
1. Element + Element  Compound
2
A + B  AB
Na + Cl
2
4 Al + 3 O
2
 2 NaCl
 2 Al
2
O
3

2. Compound + Compound  Compound
EXAMPLE 1: metal oxide + carbon dioxide
 a carbonate
CaO + CO
2

Ca CO
3
EXAMPLE 2: metal oxide + water
 a base (hydroxide)
Na
2
O + H
2
O

2 Na OH
EXAMPLE 3: nonmetal oxide + water
 an acid
SO
3
+ H
2
O

H
2
S O
4

1. Binary Compound  2 elements
AB  A + B
2 H
2
O  2 H
2
+ O
2
2 HgO  2 Hg + O
2

2. Ternary Compound

Compound + Element/Compound a. metal carbonate
 metal oxide + carbon dioxide
CaCO
3

CaO + CO
2 b. metal hydroxide
 metal oxide + water
(Except Group IA metals)
Mg(OH)
2

MgO + H
2
O c. metal chlorate
 metal chloride + oxygen
2KClO
3

2KCl + 3O
2 d. acids
 nonmetal oxide + water (HINT: MUST USE CHARGE OF H and O)
H
2
CO
3

CO
2
+ H
2
O e. Other 2H
2
O
2

2H
2
O + O
2

Single Replacement or
Single Displacement
Element + Compound  New Compound + New Element
1. Metals
A + B C  A C + B
Active metals displace less active metals or hydrogen from their compounds in aqueous solution. Refer to the Activity Series.
a. Zn + Cu SO
4
 Zn SO
4
+ Cu b. metal + H
2
O  metal hydroxide + H
2
Metals include the alkali metals and calcium.
2Na + 2H
2
O  2NaOH + H
2

• The reactivity of a metal is based on its ability to replace another in a compound.
• If a single replacement reaction occurs, the metal that “cuts in” is MORE reactive than the one that was removed or replaced.
• An activity series of metals is a listing that ranks metals according to their reactivity.
– The most active metal is at the TOP of the list
– The least active metal is at the BOTTOM of the list

The ACTIVITY SERIES is listed below : lithium potassium barium strontium calcium sodium magnesium aluminum manganese zinc iron cadmium cobalt nickel tin lead hydrogen copper silver mercury gold
The most active metal is
The least active metal is
Which is more active
IRON
LITHIUM
GOLD nickel or iron?

2

3
http://www.youtube.com/watch?v=QQ0eLyBz
KYs

Single Replacement or
Single Displacement
2. Nonmetals
D + E F  E D + F
Cl
2
+ 2 Na Br  2 Na Cl + Br
2
Many nonmetals displace less active nonmetals from combination with a metal or other cation. Order of decreasing activity is
F
2
 Cl
2
 Br
2
 I
2

Double Replacement or Double
Displacement or Metathesis:
Compound + Compound  New Compound + New Compound
A B + C D  A D + C B
Ag NO
3
+ Na Cl  Ag Cl + Na NO
3
Special Notes:
If NH
4
OH is one of the products, it breaks down into NH
3 and H
2
O.
If H
2
CO
3 is one of the products, it breaks down into CO
2 and H
2
O.

• If the reactants are both aqueous in a double replacement reaction, a precipitate or gas will form.
• Use solubility rules to determine the identity of the precipitate.
• AgNO
3 (aq)
+ NaCl
(aq)
 AgCl
(s)
+ NaNO
3 (aq)

Electrolyte vs. Nonelectrolyte

• Net ionic equations are equations that show only the soluble, strong electrolytes reacting
(these are represented as ions) and omit the
spectator ions, which go through the reaction unchanged.
– Substances that are aqueous break down into ions.
– Substances that pure solids, liquids, or gases do not break down in solution.
– Hint: Remember to check solubility rules to determine if a precipitate forms.

• Release heat into the surroundings
• Heat is a product of the reaction
• Combustion reactions are exothermic
• C
3
H
8
+ 5O
2
→ 3CO
2
+ 4H
2
O + 2043kJ

• Heat is absorbed by the reactants and stored in the chemical bonds of the product.
• Heat acts as a reactant.
• C + H
2
O + 113kJ → CO + H
2
• Motorcycle Helmet

• The heat absorbed or released during a reaction depends on a difference in a quantity called enthalpy. (Total energy content of a sample.)
• The symbol for enthalpy is H.
• When reactions take place at standard temperature and pressure, q = H.
• Stand. temp. = 25 ° CStand. Pres. = 1 atm
• Purest form of a substance = most stable form
• Enthalpy change at STP denoted H °

Type of reaction
∆H H products H reactants
Exothermic
-
< >
Endothermic + > <

Collision Theory
• Collision Theory – Molecules must collide with each other in order to react.
• Types of collisions
– Effective Collisions: Lead to the formation of products
– Ineffective Collisions: Do not lead to the formation of products
• Factors that determine effective collisions
– Energy of particles
– Orientation of particles

Ineffective Collision – Insufficient Energy - No Products
Effective Collision – Sufficient Energy – Forms Products

Effective collisions have enough energy, and the correct orientation to form products.
Ineffective collisions revert to the original products.

A
• Particles must possess a minimum amount of energy in order to react.
• Tied to collision theory
• Activation Energy: Energy needed to start a reaction; difference between the energy at the peak and the energy of the reactants

Exothermic
Reaction
Endothermic
Reaction

kJ
• What is the energy of the reactants?
• What is the energy of the products?
150 kJ
50 kJ
• Is the forward reaction exothermic or endothermic?
exothermic
• What is the ΔH for the forward reaction?
-100 kJ
• What is the ΔH for the reverse reaction?
100 kJ

• Bond energy is the strength of a chemical bond between atoms, expressed as the amount of energy required to break it apart. (Unit kJ/mol)
• It is as if the bonded atoms were glued together: the stronger the glue is, the more energy would be needed to break them apart. A higher bond energy, therefore, means a stronger bond.
• Ionic bonds are stronger than covalent bonds. Among covalent bonds, triple bonds are stronger than double bonds and double are stronger than single bonds.

Bond Energy for
Carbon Bonds
• Bond lengths (Å):
Single > double > triple
1.54
1.34
1.20
• Bond Energy (kJ/mol)
Single < double < triple
346 612 835

• On the REACTANT side, determine the number and type of each bond.
– Determine the total amount of energy needed to break them. ( endothermic – positive sign )
• On the PRODUCT side, determine the number and type of each bond.
– Determine the total amount of energy needed for bonds to form. ( exothermic – negative sign )
• Calculate the enthalpy change for the entire reaction.
– ∆H = reactants + products