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Chapter 17 Principles of Reactivity: Chemistry of Acids and Bases

Chapter 17 Principles of Reactivity: Chemistry of Acids and Bases. Acids & Bases.

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Chapter 17 Principles of Reactivity: Chemistry of Acids and Bases

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  1. Chapter 17 Principles of Reactivity: Chemistry of Acids and Bases

  2. Acids & Bases • Acids are some of peoples' favorite chemicals. Everyone's favorite soft drink is a dilute acid solution. Your own stomach contains the strong acid : HCl. Citrus fruits contain citric acid. If wine is too aged - exposed to oxygen, it turns sour - it forms acetic acid. Sulfuric Acid is the top commercially produced chemical in the United State. Although much of it is used in the steel and petroleum refining industries, several million tons of sulfuric acid are used to make Jello.

  3. ACIDS BASES PROPERTIES OF ACIDS AND BASES

  4. Arrhenius definition acid: produces hydronium ion (H3O+) in aqueous solution base: produces hydroxide ion (OH–) in aqueous solution

  5. H3O+ + F– HF + H2O acid: base: NH4+ + OH– NH3 + H2O Brønsted definition donates a proton (hydrogen ion, H+) accepts a proton A conjugate acid is formed by adding a proton to something. A conjugate base is formed by removing a proton from something. acid base conjugate acid conjugate base base acid conjugate acid conjugate base

  6. HA H3O+ + A– B BH+ + OH– Relative strengths of conjugate acid-base pairs  If HA is a stronger acid then A– is a weaker base. If HA is a weaker acid then A– is a stronger base.  If B is a stronger base then BH+ is a weaker acid. If B is a weaker base then BH+ is a stronger acid.

  7. You should memorize the names and formulas of the 6 STRONG ACIDs , i.e., HCl, HBr, HI, HClO4, HNO3, and H2SO4. • The organic acid present in vinegar, acetic acid, is a common WEAK ACID. • The common STRONG BASES contain the hydroxide ion (OH-). • Ammonia (NH3), a common WEAK BASE, is that smelly stuff your Grandma used in a dilute solution to clean windows

  8. Ion Product Constant of Water • Water is an important solvent. • Universal solvent • Biological solvent • Small size • Density of water is greater than ice • Very polar • Hydrogen Bonding

  9. Self-ionization of Water • Water is an amphiprotic substance that can act either as an acid or a base. • HC2H3O2(aq) + H2O(l) H3O+ + C2H3O2-(aq) • acidbaseacid base • H2O(l) + NH3(aq) NH4+(aq) + OH-(aq) • acidbaseacidbase

  10. Self-ionization of Water • When water molecules react with one another to form ions. • H2O(l) + H2O(l) H3O+(aq) + OH-(aq) • (10-7M) (10-7M) • Kw = [ H3O+ ] [ OH- ] • = 1.0 x 10-14 at 25oC • Note: [H2O] is constant and is already • included in Kw. ion product of water

  11. pH and pOH • We need to measure and use acids and bases over a very large concentration range. • pH and pOH are systems to keep track of these very large ranges. • pH = - log [H3O+] • pOH = - log [OH-] • pH + pOH = 14

  12. pH Calculations • Determine the following. pH = - log [H+] • pH of 6.7x10-3 M H+ • = 2.2 • pH of 5.2x10-12M H+ • = 11.3 • [H+], if the pH is 4.5 • = 3.2 x 10-5M H+

  13. pOH Examples • Determine the following. • pOH = - log [OH-] = 14 - pH • pOH of 1.7 x 10-4 M NaOH • pOH = 3.8 pH = 10.2 • pOH of 5.2 x 10-12M H+ • pOH = 2.7 pH = 11.3 • [OH-] , if the pH is 4.5 • pOH = 9.5 • [OH-] = 3.2 x 10-10M

  14. pH Scale • A log based scale used to keep track of the large change important to acids and bases. 14 7 0 10-14M 10-7M 1 M Very Neutral Very Basic Acidic When you add an acid, the pH gets smaller. When you add a base, the pH gets larger.

  15. pH of SomeCommon Materials • Substance pH • 1 M HCl 0.0 • Lemon juice 2.3 • Coffee 5.0 • Pure Water 7.0 • Blood 7.35 - 7.45 • Milk of Magnesia 10.5 • 1M NaOH 14.0

  16. [H3O+][A–] [HA][H2O] [H3O+][A–] [HA] ~ constant (55 M)  Kc·[H2O] = Ka = acid dissociation constant Definitions Ka, pKa HA + H2O H3O+ + A– Kc = pKa = -log Ka if pKa = 5 then Ka = 10–5 if pKa = 8 then Ka = 10–8 stronger acid weaker acid

  17. Acid Ionization Constant, Ka • Acid ionization constants let us define weak, moderate and strong acids. • Ka < 10-3; it is a weak acid. • Ka = 10-3 to 1; it is a moderate acid. • Ka > 1; it is a strong acid.

  18. [BH+][OH–] [B] [BH+][OH–] [B][H2O] ~ constant  Kc·[H2O] = Kb = base dissociation constant Definitions Kb, pKb B + H2O BH+ +OH– Kc = pKb = -log Kb if pKb = 4 then Kb = 10–4 if pKb = 9 then Kb = 10–9 stronger base weaker base

  19. Ka and Kb Values • For weak acids and bases: • Ka and Kb always have values that are smaller than one. • Acids with a larger Ka are stronger than ones with a smaller Ka. • Bases with a larger Kb are stronger than ones with a smaller Kb. • Ka x Kb = Kw • Most acids and bases are considered weak.

  20. pKa and pKb Concepts • The negative logarithms of Ka and Kb are useful in the same way as pH. • pKa = - log Ka • pKb = - log Kb • pKa + pKb = 14.00 • The larger that the value of pKa is, the weaker the acid. • The larger that the value of pKb is, the weaker the base.

  21. Kw = [H3O+][OH–] = 10–14 (constant at 25ºC) pKw = pH + pOH = 14  [H3O+] = 10–14 [OH–] Kw: autodissociation of water H2O + H2O H3O+ + OH– pH = 14 - pOH

  22. A– HA + OH– Kb = [HA][OH–] [A–] [HA][OH–] [A–] [H3O+][A–] [HA] [H3O+][A–] [HA] usually only one or the other given in a table Ka · Kb = · Ka · Kb = Kw Kw Ka Kw Kb or Kb = or Ka = for a conjugate acid-base pair Ka and Kb for conjugate acid-base pairs HA H3O+ + A– Ka = = [H3O+][OH–] = Kw

  23. log [OH–] -1.60 + 14 pH = 12.40 III. pH Calculations A. Strong acids and bases 100% dissociated  for strong acid: [H3O+]eq = [HA]I base: [OH–]eq = [B]i • e.g., 1.0 x 10–3M HCl • [H3O+] = pH = • [OH–] = pOH = • e.g., 2.5 x 10–2M NaOH • [OH–] = pOH = • [H3O+] = pH =

  24. HC2H3O2 H3O+ + C2H3O2– x2 (0.10 - x) 1.8 x 10–5 =  If [HA]i 400·Ka then x << [HA]i (If not, then have to solve quadratic.) 1.8 x 10–5 = x2 (0.10) III. pH Calculations B. Weak acids and bases [H3O+][A–] [HA] HA H3O+ + A– Ka = Solve equilibrium expressions [BH+][OH–] [B] B BH+ + OH– Kb = e.g., What is the pH of 0.10 M HC2H3O2? (Ka = 1.8 x 10–5) x = [H3O+] = 1.3 x 10–3M (assumption valid) pH = 2.87 assume x << 0.10

  25. Buffers • Solutions that resist change in pH when small amounts of acid or base are added. • Two types: • weak acid and its salt. • weak base and its salt. • HA(aq) + H2O(l) H3O+(aq) + A-(aq) • Add OH- Add H+ • shift to right shift to left • Based on Le Châtelier’s Principle.

  26. [H3O+][HA–] [H2A] H2A H3O+ + HA– Ka1 = [H3O+][A2–] [HA–] HA– H3O+ + A2– Ka2 = III. pH Calculations C. Polyprotic acids H2SO4, H2SO3, H2CO3, etc. Lose their protons in separate steps: (Usually, Ka1 >> Ka2) Assume: 1) [H2A], [H3O+], and [HA–] can be determined from the 1st step. (i.e., HA– dissociates only very little.) 2) [A2–] can be determined from the 2nd step.

  27. Buffers and Blood • Control of blood pH. • Oxygen is transported primarily by hemoglobin in the red blood cells. • CO2 transported both in plasma and the red blood cells. • CO2 (aq) + 2 H2O • H2CO3 (aq) • H3O+(aq) + HCO3-(aq) Carbonate Buffer

  28. Buffers and Blood • The amount of CO2 helps control blood pH. • Too much CO2 - Respiratory arrest, • pH goes down, acid level goes up. • acidosis • Solution - ventilate and give bicarbonate via IV. • Too little CO2 - Hyperventilation, anxiety, • pH goes up, acid level goes down. • alkalosis • Solution - re-breathe CO2 in paper bag to raise level.

  29. Quantitative Aspects of Buffers • Ka for a weak acid: • HA H+ + A- • Ka = [H+] [A-] • [HA] • Henderson-Hasselbalch Equation: • pH = pKa + log [anion] • [acid]

  30. Neutralization • The reaction of an acid with a base to produce a salt and water. • HCl + NaOH NaCl + H2O • We do this when we use antacids. • Neutralization can be used to determine the amount of acid or base in a sample. - titrations

  31. Titrations • Analytical methods based on measurement of volume. • If the concentration of an acid is known, the concentration of the base can be found. • If we know the concentration of the base, then we can determine the amount of acid. • All that is needed is some calibrated glassware and either an indicator or pH meter.

  32. TItrations Buret - volumetric glassware used for titrations. It allows you to add a known amount of your titrant to the solution you are testing. If a pH meter is used, the equivalence point can be measured. An indicator will give you the endpoint.

  33. Indicator Examples • Acid-base indicators are weak acids that undergo a color change at a known pH. phenolphthalein methyl red bromothymol blue

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