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 equation4 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 Equation1. Write correct formulas for reactants2. Count atoms on both sides (is it balanced?)3. 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 + H2O3 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 + 4H2O3 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