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Understanding Acid-Base and Redox Reactions: Definitions and Examples

This resource provides a comprehensive overview of acid-base and redox reactions. It explains key definitions, including the Arrhenius and Brønsted-Lowry definitions of acids and bases, and discusses the characteristics of strong vs. weak acids and bases. The text also covers the nature of redox reactions, the movement of electrons, oxidation and reduction processes, and identifying oxidizing and reducing agents. Detailed examples illustrate each reaction type, aiding in the understanding of these essential chemical concepts.

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Understanding Acid-Base and Redox Reactions: Definitions and Examples

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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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