1. Chapter 8 Chemical Equations and Reactions

2. 8-1: Describing Chemical Reactions A. Indications of a Chemical Reaction 1)Evolution of energy as heat and light 2)Production of a gas 3)Precipitate formation 4)Color change

3. B. Chemical Equations 1.) Definition of chemical equation: A shorthand method of showing the changes that take place in a chemical reaction. It shows chemical formulas of the reactants and products, as well as the relative amounts of reactants and products (if it doesn’t show the relative amounts, it is called a formula equation). Ex. H 2 + O 2 → H 2 O (formula equation) 2H 2 + O 2 → 2H 2 O (chemical equation)

4. 2.) Writing chemical equations:  Correct chemical formulas must be used for compounds  Elements are written as single atoms, except for 7 that are diatomic (H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 )  The physical state of a substance is given in parenthesis after the chemical formula. Solid = (s) or ↓, liquid = (l), gas = (g) or ↑, aqueous (a solution in water) = (aq)

5.  A catalyst (a substance that changes the rate of a reaction, but is not changed itself) is written above the →  Heat added to the reaction is also written above the → as the word heat or the symbol Δ  Other symbols are given on page 266

6. 3.) Balancing chemical equations: Atoms are neither lost nor created during a reaction, they are just rearranged (laws of conservation of matter and mass). Each side of an equation must have the same number of each kind of atom. Ex. H 2 + O 2 → H 2 O (not balanced) To balance, use coefficients (#s placed in front of chemical formulas to show the # of units involved). Ex. 2H 2 + O 2 → 2H 2 O (balanced)

7. Hints: If a polyatomic ion appears on both sides of the equation, balance it as a unit Balance H and O last

8. 8-2: Types of Chemical Reactions overview of reaction types 1) Synthesis – 2 or more substances react to form a single substance. Ex. Al + O 2 → Al 2 O 3 (not balanced) C + O 2 → CO 2 CaO + H 2 O → Ca(OH) 2 (a metal oxide and water form a base) SO 2 + H 2 O → H 2 SO 3 (a nonmetal oxide and water form an acid)

9. Calcium and oxygen combine to make calcium oxide

10. 2) Decomposition – a single compound is broken down into 2 or more simpler products. Most require energy. Ex. HgO → Hg + O 2 (not balanced) H 2 O 2 → H 2 O + O 2 (not balanced) CO → C + O 2 (not balanced)

11. 3) Single-Replacement – a free element replaces an element in a compound. Metals replace other metals or hydrogen. Nonmetals replace nonmetals. Ex. Cu + AgNO 3 → Ag + Cu(NO 3 ) 2 (not balanced) Fe + CuSO 4 → FeSO 4 + Cu

12. The free element must be more reactive than the element it replaces in order for the reaction to occur. Activity series of metals = a table that lists metals in order of decreasing reactivity (pg. 286). A free metal can only replace a metal listed below it from a compound. For nonmetals, the activity of the halogens is also on pg. 286.

13. 4) Double-Replacement – the ions of 2 compounds exchange places to form 2 new compounds. Ex. AgNO 3 + NaCl → AgCl + NaNO 3 FeS + HCl → FeCl 2 + H 2 S (not balanced) One of the products is usually either: a precipitate a gas that bubbles out, or a molecular compound such as water

14. 5) Combustion– a substance reacts with oxygen, releasing heat and light. Ex. CH 4 + O 2 → CO 2 + H 2 O (not balanced) These reactions often involve hydrocarbons (compounds of C and H) which are used as fuel. The products are CO 2 and H 2 O.

15. 6) Oxidation / Reduction reactions (redox) [from section 19.1] – any chemical reaction in which elements undergo changes in oxidation number (which means there’s a transfer of electrons). Oxidation = an increase in oxidation number (loss of electrons) Ex. Na  Na + + e - Reduction = a decrease in oxidation number (gain of electrons) Ex. Cl 2 + 2e -  2Cl - These 2 processes must occur together.

16. Ex. MnO 4 - + 5Fe 2+ + 8H +  Mn 2+ + 5Fe 3+ + 4H 2 O All single replacement reactions, combustion reactions, and some of the other types, are redox reactions. Double replacement reactions are NOT redox reactions

17. Chapter 17 Reaction Process and Rate 1)Reaction Process: a. Reaction Mechanism = the step by step sequence of reactions that “add up” to the overall reaction Ex. H 2 + I 2 → 2HI This reaction occurs in several steps. b. Collision Theory – states that particles will react if they:  Collide  collide with enough energy  collide with a favorable orientation

18. c. Activation Energy = the minimum energy colliding particles must have in order to react.

19. 2) Reaction rate = the change in concentration of reactants per unit time. The factors that affect reaction rate include: temperature: ↑ temp., ↑ rate (more collisions have enough energy to get over the activation energy barrier) concentration: ↑ conc., ↑ rate (collision frequency increases) particle size: as particle size ↓, surface area and rate ↑. A good way to ↑ surface area is to dissolve the reactants. catalyst: they lower the activation energy needed to react.

20. 18.1: Chemical Equilibrium Reversible Reactions  Definition: a reaction in which products → reactants as well as reactants → products. Example: 2HI H 2 + I 2  If you start with just HI, the reaction goes forward until the concentration of products builds up. Then the reverse reaction occurs.  Chemical Equilibrium = the point at which the forward and reverse reactions are taking place at the same rate (there is no change in the amounts of reactants or products)

22.  Equilibrium position = the relative amounts of reactants and products at equilibrium. In some reactions, the products are favored. In some, the reactants are favored.  Equilibrium expression: for the reaction – aA + bB cC + dD  K = the equilibrium constant. It depends on the particular reaction and the temperature. – If K>1, products are favored. – If K<1, reactants are favored