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Acids, Bases, and Acid-Base Equilibria
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Acid-Base Theories and Relative Strengths Arrhenius Theory of acids and bases acid – produces H + ions base – produces OH - ions Strong acids and bases ionize completely Problems with this theory: It’s restricted to aqueous solutions and it doesn’t include bases like NH 3 which don’t directly ionize to yield OH -. Does predict formation of salt and water.
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Bronsted-Lowry Theory of Acids and Bases Acids are proton (H + ) donors (A Brønsted–Lowry acid must have at least one removable (acidic) proton (H + ) to donate) Bases are proton (H + ) acceptors (A Brønsted–Lowry base must have at least one nonbonding pair of electrons to accept a proton (H + ))
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Conjugate acid: formed when a base accepts a proton Conjugate base: formed when an acid donates a proton Conjugate acid-base pair: an acid and a base that differ by only one H + On each side of the equation, there’s an acid and a base
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Amphiprotic – a substance that can act as an acid or a base Ex. H 2 O As a base → As an acid →
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The relative strengths of acids and bases: – The stronger an acid is, the weaker its conjugate base is – The stronger a base is, the weaker its conjugate acid is In every acid-base reaction, equilibrium favors transfer of the proton from the stronger acid to the stronger base to form the weaker acid and the weaker base. See Zumdahl sect. 14.2
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The Self-ionization of Water K w (ion product constant of water) = H 3 O + x OH - = 1 x 10 -14 at 25 o C H 3 O + = OH - = 1 x 10 -7 for pure water (at 25°C)
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If a solution is neutral, [H + ] = [OH – ]. If a solution is acidic, [H + ] > [OH – ]. If a solution is basic, [H + ] < [OH – ]. To calculate the concentration of H + or OH - when you only know one of them, use the equation H 3 O + x OH - = 1 x 10 -14
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The pH Scale pH: it’s a logarithmic scale: pH = -log H 3 O + pOH = -log OH - pK w = -log K w = 14 (at 25°C) pK w = pH + pOH = 14 Neutral solutions have a pH = 7 Acidic solutions have a pH<7 Basic solutions have a pH>7 Note: when the pH changes by 1, the H 3 O + changes 10 fold.
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Measuring pH: 2 main ways of measuring pH – 1)with a pH meter – has electrodes that indicate small changes in voltage to detect pH. 2)With acid-base indicators (substances that can exist in either an acid or a base form, and are different colors at different pHs) sample rxn: HIn + H 2 O H 3 O + + In - Works according to LeChatelier’s Principle (more H+ shifts equilibrium to left, less shifts to right; shift Indicated by color change)
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Strong Acids and Bases Strong Acids Remember that the seven strong acids are HCl, HBr, HI, HNO 3, H 2 SO 4, HClO 3, and HClO 4. These are strong electrolytes and exist totally as ions in aqueous solution; e.g., HA + H 2 O → H 3 O + + A – So, for the monoprotic strong acids, [H 3 O + ] = [acid] Just use acid concentration for H+ concentration.
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Strong Bases Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal (Ca 2+, Sr 2+, and Ba 2+ ) hydroxides. Again, these substances dissociate completely in aqueous solution; e.g., MOH(aq) → M + (aq) + OH – (aq) or M(OH) 2 (aq) → M 2+ (aq) + 2 OH – (aq) For the alkali metal hyroxides: pOH = -log[OH - ] = -log[base] pH = 14 – pOH For the alkaline earth metal hydroxides: pOH = -log[OH - ] = -log{2x[base]}
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Weak Acids For a weak acid, the equation for its dissociation is HA(aq) + H 2 O(l) ⇌ H 3 O + (aq) + A – (aq) Since it is an equilibrium, there is an equilibrium constant related to it, called the acid-dissociation constant, K a : The greater the K a, the stronger the acid.
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Calculating pH from k a and initial concentration: Compare Ka to Kw (Ka should be larger, otherwise pH = 7) Ka = [H+][A-]/[HA]. Ka is given. H+ A- HA Init 0 0 Given (G) Change +x +x G-x End x x Assume G (small dissociation, < 5%) Solve (Ka(HA)) 1/2 = x ( Hope M acid /K a > 100, othewise use quadratic eqn to find x ) -log [x] = pH Calculating Percent Ionization: Percent ionization is a measure of acid strength (the stronger the acid, the greater the % ionization) Percent ionization = 100
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Polyprotic Acids: Monoprotic acid – 1 ionizable H Polyprotic acid – more than 1 ionizable H. The ionizations occur separately. Each step has its own K a. The first K a is the largest. If the K a values for the first and second dissociation differ by a factor of 10 3 or more, the pH generally depends only on the first dissociation.
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Weak Bases Like weak acids, weak bases have an equilibrium constant called the base dissociation constant (K b ). Equilibrium calculations work the same as for acids, using the base dissociation constant instead. Since “x” in this case is the [OH - ] at equilibrium, taking the -log of it gives you the pOH.
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Relationship between K a and K b For a conjugate acid–base pair, K a and K b are related in this way: K a × K b = K w Therefore, if you know one of them, you can calculate the other. Also, pK a + pK b = pK w = 14 for a conjugate acid-base pair
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Acid-Base Properties of Salt Solutions Combined effect of cation and anion – possible combinations: 1)Salts containing ions from strong acids and strong bases form neutral solutions. (ex. NaCl, KNO 3, BaI 2 ) 2)Salts containing ions from weak acids and strong bases form basic solutions. (ex. Na 2 CO 3, KNO 2, NaCH 3 COO) 3)Salts containing ions from strong acids and weak bases form acidic solns. (ex. NH 4 Cl, NH 4 NO 3 ) 4)Salts containing ions from weak acids and weak bases (or containing ions like Fe 3+ ) –depends on the relative acid and base strength (ex. NH 4 CN, NH 4 NO 2, CrF 3 ). Look for larger Ka or Kb.
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Acid-Base Behavior and Chemical Structure Factors that Affect Acid Strength: 1)The H—A bond must be polar with δ + on the H atom and δ – on the A atom 2)Bond strength: Weaker bonds can be broken more easily, making the acid stronger. 3)Stability of A – : More stable anion means stronger acid. The strength of an acid is often a combination of all three factors.
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Binary Acids: Binary acids consist of H and one other element. Within a group, H—A bond strength is generally the most important factor (acid strength increases going down a group). Within a period, bond polarity is the most important factor to determine acid strength (acid strength increases going across a period)
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Oxyacids: acids in which OH groups and possible additional oxygen atoms are bound to a central nonmetal atom. As the electronegativity of the nonmetal increases, the O-H bond gets weaker, and the acid gets stronger Also, as additional atoms high in electronegativity bond to the center atom, the acid strength increases:
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Carboxylic Acids: have a –COOH (carboxyl) group. The electronegativity of the R group attached to the carboxyl group determines the strength. The greater the electronegativity, the stronger the acid. K a = 1.6 x 10 -5 K a = 2.3 x 10 -1
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