Chemical Changes

1. 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)

3. 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)

5. Balancing a Chemical Equation • 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) • 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

6. 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

8. 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)

10. 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

11. 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)

13. 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

15. 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

19. 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