Chemical Changes
• Physical changes (Ch. 5) involve only changes in the
physical form of the substance, not in the atomic or
molecular make-up
• Chemical changes involve conversion into new
substances with new chemical properties
• Chemical changes can often be observed:
– Color change, precipitate (solid forms), bubbles, etc.
• In a chemical reaction, reactants go to products
• Atoms of reactants are recombined in products:
2Ag + S  Ag2S
(silver reacts with sulfur to form silver sulfide)
Chemical Equations
• Used to represent chemical reactions
• Like a recipe:
- Tells what you need to start with, and how much
- Also tells what you will make, and how much
• Example: 2H2 + O2  2H2O
• Number of each type of atom must be equal on the two
sides of the equation
4 H’s + 2 O’s = 4 H’s + 2 O’s
• Use coefficients to balance chemical equations
• Sometimes symbols are used to show physical state:
(s) = solid, (l) = liquid, (g) = gas and (aq) = aqueous (in
water)
• Example: C(s) + O2(g)  CO2(g)
Balancing a Chemical Equation
1.
2.
3.
Write correct formulas for reactants
Count atoms on both sides (is it balanced?)
Balance one element at a time (usually C first and O or H
last, but can be any order)
4. Count atoms again to check that it’s balanced
Example: Propane (C3H8) burns with oxygen to form carbon
dioxide and water. Write the balanced chemical equation.
C3H8 + O2  CO2 + H2O
3 C’s + 8 H’s + 2 O’s  1 C + 2 H’s + 3 O’s
C3H8 + O2  3CO2 + H2O
C3H8 + O2  3CO2 + 4H
C3H8 + 5O2  3CO2 + 4H2O
3 C’s + 8 H’s + 10 O’s = 3 C’s + 8 H’s + 10 O’s
Types of Reactions
• Reactions can be organized into 4 basic types:
combination, decomposition, replacement and combustion
- Combination reactions: 2 (or more) reactants combine
to form a single product
- Decomposition reactions: One reactant splits into 2 (or
more) products
- Replacement reactions: Elements are exchanged
between 2 reactants to form 2 products
- Combustion reactions: fuel + oxygen  products +
heat
• Reactions can be more than one type
Oxidation-Reduction (Redox) Reactions
• Some reactions are also categorized as redox reactions
• In these reactions the reactants exchange electrons
- Reduction = gain of electrons (GER)
- Oxidation = loss of electrons (LEO)
• Oxidation and reductions reactions are always coupled
(electrons gained = electrons lost)
• Example: Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
- Mg loses 2 electrons to become Mg2+ (Mg is oxidized)
- Each Cl gains an electron to become Cl- (Cl is reduced)
- H is not oxidized or reduced (no change in # of
electrons)
• Also, in general, gain of O or loss of H = oxidation and
gain of H or loss of O = reduction (in biological systems)
Energy in Chemical Reactions
• In order for a reaction to take place, the reactants must
contact each other with enough energy
• As reactants collide, bonds are broken and new bonds
are formed
• Example: 2H2 + O2  H2O (the H-H and O-O
bonds break and two new O-H bonds are formed)
• Between the reactants and the products there is a
“transition state” in which bonds are breaking and/or
forming (highest E point in reaction)
• Energy required to reach transition state is called
“activation energy” (EA)
• Transition state is always highest E (higher than
reactants and products) because it takes E to break
bonds and E is released when bonds are formed
Exothermic and Endothermic Reactions
• The difference in energy between the reactants and the
products is called the “heat of reaction”
• Heat of reaction can be heat released or heat
consumed, depending on the reaction
• Reactions that release heat are exothermic
CH4 + 2O2  CO2 + 2H2O + heat (213 kcal)
• Reactions that consume heat are endothermic
H2 + I2 + heat (12 kcal)  2HI
• Energy diagrams are used to show energy changes
during a chemical reaction (E vs. reaction progress)
Reaction Rates
• Reaction rate = how fast a reaction goes from
reactants to products
• Rate is based on activation energy and not on heat of
reaction (lower EA = faster reaction)
• Reaction rates are affected by such factors as :
- Reactant concentration
(more reactants =
more collisions = faster reaction)
- Temperature (at higher T reactants collide
more often at higher E = faster reaction)
- Catalyst (addition of a catalyst lowers the
activation energy = faster reaction)
- a catalyst makes the transition state more stable,
so it takes less EA to reach it
Chemical Equilibrium
• Some chemical reactions are reversible (products can
also go to reactants)
• Example: N2(g) + O2(g)  2NO(g)
- Forward reaction = N2 + O2  2NO
- Reverse reaction = 2NO  N2 + O2
• When rate of forward reaction = rate of reverse
reaction, chemical equilibrium has been reached
• When at equilibrium:
- If more products exist in reaction mixture, then
reaction favors products
- If more reactants exist in reaction mixture, then
reaction favors reactants
LeChâtelier’s Principle
• The equilibrium can be shifted towards more products
or more reactants by placing a “stress” on the system
• Add reactants or remove products and equilibrium is
shifted towards products
• Add products or remove reactants and equilibrium is
shifted towards reactants
• Heat is also considered a reactant (endothermic
reactions) or a product (exothermic reactions)
• Example: C(s) + H2O(g) + heat  CO(g) + H2(g)
- Add heat: equilibrium shifts towards products
- Remove H2(g): equilibrium shifts towards products
- Remove H2O(g): equilibrium shifts towards reactants