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Warm-Up: To be turned in

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. Acid-Base and Redox Reactions. Acid-base Reactions. Arrhenius definition

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Warm-Up: To be turned in

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  1. Warm-Up: To be turned in Identify the type of reaction represented in the following equations: C10H8 + 12O2 ---> 10CO2 + 4H2O 8Fe + S8 ---> 8FeS NaOH + HCl NaCl + H2O

  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), H2SO4, HNO3, HClO4 • Ex. strong bases- all hydroxides • Weak acids/bases do not ionize completely aqueous solutions • Ex. Weak acids- HF, H3PO4, HCN, H2CO3 • Ex. Weak bases- NH3

  5. Acid-base Reactions • Acids and bases will combine in a double-replacement reaction to form water and a salt • HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) • Some acids will decompose to form a non-metal oxide and water • H2CO3(aq)  CO2(g) + H2O(l) • Acids can also undergo single-replacement by metals to form hydrogen gas and a salt • Mg(s) + 2HCl(aq) → H2(g) + MgCl2(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 H2O2, 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 F2(g) + 2NaCl(aq) → 2NaF(aq) + Cl2(g) reduced oxidized 0 +1 -1 +1 -1 0 Oxidizing agent Reducing agent

  11. Practice:Identify Oxidizing and Reducing Agents 2 Ag(s) + S(s)  Ag2S(s) 2 Ag(s) + Cu2+(aq)  2 Ag+(aq) + Cu(s)

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