Chapter 15 Acids and Bases Week 1-1 (09)-1

Chapter 15 Acids and Bases Week 1-1 (09)-1 - CHAPTER 15...

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CHAPTER 15 ACID-BASE EQUILIBRIA Introduction to Acids, Bases and the Equilibrium Concept Water (H 2 O) – the most important molecule on earth. Even in pure water, there are small amounts of ions from the equilibrium below (“self-ionization of water” or “auto-ionization of water”). H 2 O (l) H + (aq) + OH - (aq) More accurately: H 2 O (l) + H 2 H 3 O + - H 3 O + (aq) = hydronium ion; often abbreviated as H + OH - = hydroxide ion [H 3 O + ] = [OH - ] in pure water Definitions of Acids and Bases 1) ACIDS: give [H 3 O + ] > [OH - ] in solution (vinegar, lemon juice) 2) BASES: 3 O + ] < [OH - ] in solution (bicarb) Understanding “acidity” and “basicity” involves knowing the [H + ] and [OH - ] concentrations in solution. Historically, the first definitions of acids and bases were the “ Arrhenius Definitions ”. 1) ACID = a substance with H in its formula, and which dissociates in water to give H 3 O + (aq) (= H + (aq)) Generic acid = HA (e.g., HCl, HNO 3 , H 2 SO 4 , etc.) 2) BASE = a substance with OH in its formula, and dissociates to yield OH - Generic base = MOH (e.g., NaOH, Ca(OH) 2 15-1
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Neutralization : the reaction between an acid and a base H + (aq) + OH - (aq) H 2 O (l); Strengths of Acids and Bases : (i.e., the amount of H + or OH - produced per mole of substance dissolved). S TRONG ACIDS AND BASES DISSOCIATE COMPLETELY (~100%) IN SOLUTION MUST KNOW THEM ! Strong Acids : HCl, HBr, HI, HClO 4 , HNO 3 , H 2 SO 4 - plus few rare ones Strong Bases : MOH and M(OH) 2 , where M = Li + , Na + , K + , Rb + , Cs + , Ca 2+ , Sr 2+ , Ba 2+ **A LL OTHER ACIDS AND BASES ARE WEAK ** Because strong acids and bases dissociate “completely” (i.e. ~100%), we do not consider them equilibria, (i.e., K c >>> 1) and we write them as a one-directional reaction. Strong Acid : HA (g or l) + H 2 O (l) H 3 O + (aq) + A - (aq) Strong Base : MOH (s) + H 2 M + - Weak Acids : dissociate only partially in solution; it is an equilibrium. HA (aq) + H 2 H 3 O + - **Amount of dissociation varies, depending on K of the acid ** Section 15.1 The Brønsted-Lowry Definition 15-2 kJ 55.9 - rxn H Δ =
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The Arrhenius definition of acids and bases doesn’t cover all possibilities, e.g., some bases do not contain OH in their formula. The Brønsted-Lowry definition is much better. Acid :an H + donor . Must contain H + in its formula (all Brønsted-Lowry acids are also Arrhenius acids). Base : an H + acceptor . Must contain a lone-pair capable of binding an H + (e.g., NH 3 , F - , OH - , etc.). Brønsted-Lowry bases are not necessarily Arrhenius bases, but Arrhenius bases contain the Brønsted-Lowry base, OH - . This gives acid/base reactions : something can act as a Bronsted-Lowry acid if there is a Bronsted-Lowry base to pick up the H + . **The acid is the H + donor; the base is the H + acceptor** e.g. HCl (g) + H 2 O (l) Cl - (aq) + H 3 O + (acid) (base) NH 3 (g) + H 2 NH 4 + + OH - HCl (g) + NH 3 NH 4 Cl (s) NH 4 + (aq) + H 2     NH 3 (g) + H 3 O + [Fe(H 2 O) 6 ] 3+ 2     [Fe(H 2 5 (OH)] 2+ 3 O + NOTE : H 2 O is amphiprotic – it can be a Bronsted acid or base.
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Chapter 15 Acids and Bases Week 1-1 (09)-1 - CHAPTER 15...

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