1. Warm-Up: To be turned in Identify the type of reaction represented in the following equations: C 10 H 8 + 12O 2 ---> 10CO 2 + 4H 2 O 8Fe + S 8 ---> 8FeS NaOH + HCl  NaCl + H 2 O

2. Acid-Base and Redox Reactions

3. Acid-base Reactions Arrhenius definition Acid- increases H + ion concentration in an aqueous solution Base- increases OH - ion concentration in an aqueous solution Brønsted-Lowery definition Acid- proton donor Base- proton acceptor Conjugate acid- base that has accepted a proton, becomes the acid in reverse reaction Conjugate base- acid that has donated a proton, becomes the base in the reverse reaction

4. Strong vs. Weak Acids/ Bases Strong acids/ bases completely ionize (form ions) aqueous solutions – Ex. Strong acids- all binary acids (except HF), H 2 SO 4, HNO 3, HClO 4 – Ex. strong bases- all hydroxides Weak acids/bases do not ionize completely aqueous solutions – Ex. Weak acids- HF, H 3 PO 4, HCN, H 2 CO 3 – Ex. Weak bases- NH 3

5. Acid-base Reactions Acids and bases will combine in a double- replacement reaction to form water and a salt – HCl (aq) + NaOH (aq)  H 2 O (l) + NaCl (aq) Some acids will decompose to form a non- metal oxide and water – H 2 CO 3(aq)  CO 2(g) + H 2 O (l) Acids can also undergo single-replacement by metals to form hydrogen gas and a salt – Mg (s) + 2HCl (aq) → H 2(g) + MgCl 2(aq)

6. Redox reactions Short for oxidation- reduction reactions Reactions that show the movement of electrons between substances – Oxidation Is Loss of electrons – Reduction Is Gain of electrons

7. Oxidation/reducing Agents Oxidation agent- substance which causes another to be oxidized – Reduced in the process Reducing agent- substance which causes another to be reduced – Oxidized in the process

8. Rules for determining oxidation states 1. The oxidation number of any uncombined element is 0. 2. The oxidation number of a monatomic ion equals the charge on the ion. 3. The more electronegative element in a binary compound is assigned the number equal to the charge it would have if it were an ion. 4. The oxidation number of fluorine in a compound is always −1. 5. Oxygen has an oxidation number of −2 unless it is combined with F, when it is +2, or it is in a peroxide, such as H 2 O 2, when it is −1.

9. Rules for determining oxidation states 6. The oxidation state of hydrogen in compounds is +1 unless it is combined with a metal, in which case it is −1. 7. In compounds, Group 1 and 2 elements and aluminum have oxidation numbers of +1, +2, and +3, respectively. 8. The sum of the oxidation numbers of all atoms in a neutral compound is 0. 9. The sum of the oxidation numbers of all atoms in a of polyatomic ion equals the charge of the ion.

10. Redox Example F 2(g) + 2NaCl (aq) → 2NaF (aq) + Cl 2(g) 0 +1 -1 +1 -1 0 reducedoxidized Reducing agent Oxidizing agent

11. Practice: Identify Oxidizing and Reducing Agents 2 Ag (s) + S (s)  Ag 2 S (s) 2 Ag (s) + Cu 2+ (aq)  2 Ag + (aq) + Cu (s)