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ACIDS

ACIDS. 1) Sour taste: Lemon Juice – Citric acid. Vinegar – Acetic Acid. Stomach ulcers are aggravated by hydrochloric acid. HCl Dissolve active metals, usually liberating H 2 . 3) Corrosive – dissolve compounds that are otherwise hard to dissolve. Examples:

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ACIDS

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  1. ACIDS 1) Sour taste: Lemon Juice – Citric acid. Vinegar – Acetic Acid. Stomach ulcers are aggravated by hydrochloric acid. HCl Dissolve active metals, usually liberating H2. 3) Corrosive – dissolve compounds that are otherwise hard to dissolve. Examples: Precious metals such as gold (Au) dissolve in HNO3 + HCl (aqua regia). Hard water deposits dissolve in vinegar. 4) Turn litmus paper RED (low pH).

  2. BASES • 1) Bitter taste. • 2) Dissolve oil and grease. • Drano and lye soap contain NaOH. • Breaks ester and amide bonds • Slippery to the touch – dissolves hair and skin. e.g., soap: Na+ -OOC(CH2)16CH3 • 4) React with many metal ions to form precipitates. • Mg2+ + 2OH- Mg(OH)2 • Example: • Hard water (=Ca2+, Mg2+) + soap  White precipitate. (bathtub rings) • 5) Turn litmus paper BLUE (high pH)

  3. ARRHENIUS ACIDS AND BASES Arrhenius ACID: Any compound that releases H+ when dissolved in H2O. Example: HCl(g)  H+(aq) + Cl(aq) Arrhenius BASE: Any compound that releases OH- when dissolved in H2O. Example: KOH(s) + H2O (l) K+(aq) + OH(aq)

  4. BRØNSTED - LOWRY ACIDS AND BASES BrØnsted ACID: Any compound capable of donating a H+ ion. Example: HCl(g)  H+(aq) + Cl(aq) BrØnsted BASE: Any compound capable of accepting a H+ ion. Example: NH3(g) + H2O(l)  NH4+(aq) + OH(aq)

  5. WATER Water electrolyzes slightly to produce H+ and OH- reversibly. H2O H+ + OH- Autoionization of water Kw = [H+][OH-] = 1.0 x 10-14 at 25oC

  6. For pure water, [H+] = [OH-] = 10-7, so pH =7 Kw is constant even when [H+] and [OH-] are not equal Calculate [H+] in a 0.05 M Ca(OH)2 solution

  7. pH scale Most accurate method to measure pH is to use a pH meter.However, certain dyes change color as pH changes. These are indicators.HIn = H+ + In-Indicators are less precise than pH meters.Many indicators do not have a sharp color change as a function of pH. pH = -log10[H+] (low pH = acidic) pH + pOH = -log10[H+] + -log10[OH-] = 14 Measuring pH

  8. Which bulbs light up? SolutionStrong, weak, or non-electrolyte? Distilled water Tap water NaCl(aq) 1 M HCl (aq) 1 M CH3COOH (aq) Sugar CH3OH

  9. STRONG ACIDS Strong Acids dissociate completely when dissolved in water to form H+ and the corresponding BrØnsted base. HA  H+(aq) + A-(aq) Strong acids are strong electrolytes: COMPLETE dissociation into ions [H+]final = [HA]initial = CHA (If the analytical concentration, CHA, is less than 10-6M then the autoionization of water needs to be taken into account.)

  10. WEAK ACIDS • When dissolved in water weak acids only partially dissociate to form H+ and the corresponding base. • HA (aq) H+ (aq) + A- (aq) • Weak acids are weak electrolytes: • PARTIAL dissociation into ions • [H+]final < [HA]initial • Examples: CH3CO2H • HF • H3PO4 • Acid Dissociation Constant (Ka) <<1

  11. What is the [H+] of 0.10 M HI? What is the [H+] of 0.10 M acetic acid? Ka = 1.8 x10-5 1.8 x 10-5 M 4.2 x 10-3 M 1.8 x 10-6 M 1.3 x 10-5 M What is the pH? What is the % dissociation?

  12. % Dissociation of CH3CO2H CHA(M) [H+](M) % Dissoc. 10 0.013 0.13 1 0.004 0.4 0.1 0.0013 1.3 0.01 0.0004 4.0 0.001 0.00013 13.4

  13. OXYACIDS Many Brønsted acids consist of a central atom with several attached oxygen atoms. These are calledoxyacids. Acid strength increases with increasing oxidation number of the central atom: HOClO3 > HOClO2 > HOClO > HOCl General rule for uncharged oxyacids HxEOy: If y-x > 2 then strong(H2SO4, HNO3,…) If < 2 then weak(H2CO3, HBrO, HNO2,…)

  14. Increasing electronegativity of the central atom increases acid strength • HOCl > HOBr > HOI

  15. POLYPROTIC ACIDS Polyprotic acids are capable of donating more than one proton. Contain more than one ionizable proton. The Ka always gets smaller with each ionization Examples: H2CO3(aq) H+ (aq) + HCO3-(aq) Ka = 4.3 x 10-7 HCO3-(aq) H+ (aq) + CO32- (aq) Ka = 5.6 x 10-11 H3PO4 (aq) H+(aq) + H2 PO4- (aq) Ka = 7.5 x 10-3 H2PO4-(aq) H+ (aq) + HPO42- (aq) Ka = 6.2 x 10-8 HPO42-(aq) H+(aq) + PO43-(aq) Ka = 4.2 x 10-13

  16. What are the concentrations of H+, HCO3-, CO32- in 1 x 10-3 M H2CO3?

  17. Strong Acids Which one of the following are not strong acids?1.HNO3 5. HOBr 2. HF 6. HBr3. HClO3 7. HPO42-4. HClO4 8. H2SO3

  18. STRONG BASES Group I and II hydroxides (except Mg and Be). Arrhenius bases donate OH-. Brønsted bases accept H+ Examples: NaOH, KOH, Ca(OH)2 KOH + H2O  K+ (aq) + OH- (aq) Strong bases are strong electrolytes. [OH-] = Cbase

  19. WEAK BASES • When dissolved in water weak bases only partially react to form OH and the corresponding BrØnsted acid. • B + H2O HB+(aq) + OH(aq) • Weak bases are weak electrolytes: [OH-] < Cbase • Weak bases can be neutral • Example: NH3, amines • NH3 + H2O = NH4+(aq) + OH(aq) • Or Anions (any ion derived from a weak base)Example: F, NO2, CH3COO • F(aq) + H2O = HF(aq) + OH(aq) • Base Dissociation Constant Kb << 1

  20. What is the pH of 0.1 M NH3? Kb = 1.8 x 10-5 2.87 4.74 7.00 9.25 11.1

  21. CONJUGATE ACID BASE PAIRS • Differ only by the presence or absence of a proton (H+). • Conjugate Acid = Conjugate Base + H+ • Examples: • H3O+ / H2OH2O / OH • NH4+ / NH3 • HCl / Cl • Ka x Kb = constant = 1 x 10-14 • The conjugate of a weak acid is a weak base (and vice versa) • The conjugate of a strong acid is a spectator ion (example: Cl is the conjugate base of HCl). • The conjugate acid of OH (strong base) is water.

  22. When we add two reactions together, we multiply their equilibrium constants. For conjugate acid-base pairs: Ka x Kb = Kw = 1 x 1014 Larger Ka means smaller Kb The stronger the acid, weaker its conjugate base pKa = log Ka pKb = log Kb log ( Ka x Kb ) = log Kw = 14 log Ka log Kb = 14 pKa + pKb = 14

  23. Weaker acid stronger conjugate base H-F + OH- F- + H2O Weaker acid Ka = 10-14 Stronger acid 6.9 x 10-4 Weaker base Kb = 1.4 x 10-11 Stronger base

  24. ACETIC ACIDAcid: CH3COOH H+ + CH3COOBase: CH3COO + H2O CH3COOH + OH------------------------------------------------H2O H+ + OH-Ka =Kb =Kw =[H+][OH-] = Kax Kb = 1 x 10-14pKa + pKb = 14

  25. Hydrolysis: when a cation or anion reacts with H2O to form H+(aq) or OH(aq) • Will a salt be acidic or basic? • Salt derived from a strong acid and a strong base Neutral solution (pH = 7) Example: NaCl (from NaOH and HCl) • Salt derived form a weak acid and a strong base Basic solution (pH > 7) Examples: NaClO (NaOH and HClO) ClO (aq) + H2O HClO (aq) + OH(aq) (CH3COO)2Ba (Ba(OH)2 and CH3COOH) CH3COO(aq) + H2O CH3COOH(aq) +OH(aq)

  26. Salt derived from a strong acid and a weak base Acidic solution (pH <7) Example: NH4Cl (NH3 and HCl) NH4+ + H2O NH3 + H3O+ • Salt derived form a weak acid and a weak base pH depends on acid/base involved Example: NH4CN (NH4+ and CN)

  27. What is the pH of 0.02 M KN3 Ka (HN3) = 1.9 x 10-5 3.21 5.49 7.00 8.51 10.8

  28. LEWIS ACIDS • Any substance that can accept a pair of electrons. • Small cations • Molecules with unfilled octets • e.g. H+, BF3 • Examples of Lewis Acids: • Highly charged transition metal cations, e.g. Fe3+, Fe2+, Co3+ • Group III cations (Al3+, Ga3+) and compounds (AlCl3) • Smaller group II cations: Be2+ and Mg2+ LEWIS BASES • Any substance that can donate a pair of electrons. • Has lone pair electrons • May be neutral or anionic. • Examples: NH3, OH-, Brønsted bases, H2O, Cl-

  29. LEWIS CATIONS To compare acidity of Lewis acids, first compare charge. If charge is the same then compare size. Charge/Size Ratios Metal Ion Charge/Ionic radius (Å) Na+ 1.0 Li+ 1.5 Ca2+ 2.1 Mg2+ 3.1 Zn2+ 2.7 Cu2+ 2.8 Al3+ 6.7 Cr3+ 4.8 Fe3+ 4.7

  30. H : - O : + H HYDRATION Metal ions attract the lone pairs on the oxygen in water molecules. This is a Lewis acid – Lewis base reaction. Hydrated metal ions are acidic. Acidity increases with increasing charge/size ratio of the metal ions. Hydrolysis is a reaction that dissociates water: M(H2O)nz+ M(H2O)n-1(OH)(z-1)+ + H+ Fe(H2O)63+ Fe(H2O)5(OH)2+ + H+ (Ka=6.7 x 10-3) Mz+

  31. ACIDS AND BASES SO FAR • Arrhenius, Brønsted, and Lewis definitions • pH, pOH • 3) Acid and Base Dissociation Constants – Ka and Kb • 4) [H+] [OH-] = 1 x 10-14 = Ka x Kb • pH and % ionization calcn for strong and weak acids/bases • Conjugate Acid-Base Pairs: • Arrhenius • Bronsted-Lowry • Lewis • 7) Salts – Hydrolysis • 8) Structure Related to Acid-Base Properties (Oxyacids)

  32. YOU SHOULD KNOW GIVENFIND pH [H+], [OH-], pOH [H+] or [OH-] pH List of acids Weaker /Stronger List of pKa’s or Ka’s Weaker /Stronger Ka or pKa and [HX] pH, [H+], [OH-] pH and [HX] Ka Recall that a small Ka means high pKa, and both mean weak acid and not much dissociation.

  33. Acid/Base SALTS Review • 1) Which one of the following salts would have a basic aqueous solution? • 1. KF 3. NaI • 2. Al(NO3)3 4. NH4Br • 2) Arrange the following in the order of increasing base strength: • N3- NO3- HPO42- CN- • 3) Which of the following cannot act as a Lewis base? • 1. Cl- 4. NH3 • 2. OH-5. H+ • 3. CN-

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