Chapter 12 - Acids and Bases Study Guide PDF

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A very concrete and detailed study guide of chapter 12 test that covers everything from the book and class' lecture....

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Ch. 12 Acids and Bases I. General Properties A. Acids 1. Sour 2. React with some metals to form H2 gas 3. Causes certain dyes to change color (indicator paper) 4. Reacts with strong base to form salts and H2O B. Bases 1. Bitter 2. Slippery/soapy feeling 3. Causes certain dyes to change color 4. Reacts with (“dissolves”) oil and grease 5. Strong bases react with acids to form salts and water C. Brönsted-Lowry definition of acids and bases: 1. Acids are H+ donors 2. Bases are H+ acceptors E. We learned about some acids during nomenclature and double-replacement reactions 1. Molecular compounds that lose H+ in H2O 2. 1H has one electron and one proton, but no neutrons, H+ = proton 1

3. H+ does not float around freely in water Ex/ HCl + H2O  H3O+ + Cl- is more accurate 4. H3O+ is a hydronium ion (Note that HCl donated an H+ to the water) F. Acids ionize (donate H+) in water 1. Strong acids completely ionize (HCl, HBr, HI, H2SO4, HNO3, HClO4) 2. Weak acids partially ionize Ex/ HC2H3O2 + H2O ⇌ H3O+ + C2H3O2- (~5% ionization for this acid) acetic acid 4. Weak acid ionization is an equilibrium, with rxns going both directions, indicated by a double arrow 5. Some acids have one H+ to donate, others have more than one a. one H+ to donate: monoprotic (Ex/ HCl) b. two H+ to donate: diprotic, etc. (Ex/ H2SO4) c. >1 H+ to donate: (general) polyprotic d. H+’s that are donated are the acidic hydrogens G. Bases accept H+ 1. One of the most common types you’ll see are hydroxides with soluble cations (NaOH, KOH, etc.) 2. Strong bases are often ones which completely dissociate to OH- and a cation 3. Weak bases – an equilibrium rxn (Many weak bases contain a Nitrogen atom, but that is a general statement, not a rule.) Ex/ NH3 (g) + H2O ⇌ NH4+ (aq) + OH-(aq)

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H. Adding an acid and base together results in a neutralization reaction (type of doublereplacement) 1. Don’t always neutralize completely, but it is still a neutralization 2. Products of neutralizations are salt and water 3. Salt: cation from base and anion from acid 4. Salts will either be neutral or weaker acid/base 5. Need one H+ for every 1 OH- (mole to mole), so balance equations with that in mind 6. Acid + Base  Salt + Water a. Strong acid and strong base Ex/ 2HNO3 (aq) + Ca(OH)2 (aq)  Ca(NO3)2(aq) + 2 H2O(l) 1. Need 2 HNO3 for every 1 Ca(OH)2 (2 H+ and 2 OH-) 2. Strong acids dissociate completely in water, so we can write total and net ionic equations to denote this Total: 2H+(aq) + 2NO3-(aq) + Ca2+(aq) + 2OH-(aq) Ca2+(aq) + 2NO3-(aq) + 2H2O(l) Net: H+(aq) + OH-(aq)  H2O(l) b. Weak acid and strong base Ex/ HClO(aq) + LiOH(aq)  LiClO(aq) + H2O(l) 1. 1 H+ for each 1 OH2. Weak acids dissociate partially in water, there is still much of the molecular compound left, the ionic equations will show this Total: HClO(aq) + Li+(aq) + OH-(aq) ⇌ Li+(aq) + ClO-(aq) + H2O(l) Net: HClO(aq) + OH-(aq) ⇌ ClO-(aq) + H2O(l) c. Weak polyprotic acid and strong base 1. Every mole of OH- will remove/react with one H+ from the acid (each hydrogen can be neutralized one at a time) 2. Salts formed in the intermediate steps are called acid salts: an ionic compound containing an anion with one or more hydrogens that can be neutralized with a base Ex/ H3PO4(aq) + 3KOH(aq) ⇌ K3PO4(aq) + 3H2O(l) Reaction 1: H3PO4(aq) + KOH(aq) ⇌ KH2PO4(aq) + H2O(l) Reaction 2: KH2PO4(aq) + KOH(aq) ⇌ K2HPO4(aq) + H2O(l) Reaction 3: K2HPO4(aq) + KOH(aq) ⇌ K3PO4(aq) + H2O(l) d. Strong polyprotic acid and strong base 1. First rxn goes to completion, 100% 2. Following reactions reach equilibrium typical of weak acids Ex/ H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l) 3. First rxn produces NaHSO4 and H2O 4. Second rxn produces Na2SO4 and H2O

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II. Water A. Water can act as a weak acid or a weak base Ex/ H2O+ H2O ⇌ H3O+ + OHB. At 25oC, pure water has 1.0 x 10-7 M H3O+ + 1.0 x 10-7 M OHC. Also for all aqueous solutions (acidic, basic, neutral), [H3O+][OH-] = 1.0 x 10-14 M 1. Kw = Ion product of water = 1.0 x 10-14 M 2. Therefore, if [H3O+] goes up, [OH-] goes down 3. When [H3O+] = [OH-], it is a neutral solution (1.0 x 10-7 M each at 25oC) 4. When [H3O+] > [OH-], it is an acidic solution ([H3O+] > 10-7 M) 5. When [H3O+] < [OH-], it is a basic solution ([H3O+] < 10-7 M) D. pH Scale 1. pH scale is used to measure acidity of an aqueous solution 2. [H3O+] and [OH-] are small numbers, so we use a pH scale (instead of working with negative exponents) 3. It is a logarithmic scale (log scale) 4. pH is a log function a. pH = -log[H3O+] 1) lower the pH, the more acidic 2) pH < 7 is acidic 3) pH > 7 is basic 4) pH = 7 is neutral 5. We use the –log because we are working with very low concentrations Ex/ 10-7 = [H3O+] at neutrality, log (10-7) = -7 We don’t want to work with the negative numbers, so we use –log [H3O+] to get rid of the negative numbers Know how to go back and forth between [H3O+] and pH, [H3O+] = 10^-pH Ex/ pH = 8, [H3O+] = 1 x 10-8 M = inverse log or 10^-8 6. pOH = -log [OH-] 7. pH + pOH = 14 (at 25oC) 8. What affects pH? If you add acid to water, what determines the pH? a. Strength of acid b. Concentration of acid Ex/ What is the pH of a solution with 0.0030 mol HI dissolved in 250 mL of solution? What is [H3O+]? (HI is a strong acid) [H3O+] = 100% of [HI] = 0.0030 mol/0.250 L = 0.012 M pH = -log [H3O+] = - log(0.012) = 1.92

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Ex/ What is [H3O+] in a 0.100 M HC2H3O2 solution (weak acid) that is 1.34% ionized? What is the pH? HC2H3O2(aq) + H2O(l) ⇌ H3O+(aq) + C2H3O2-(aq) [H3O+] = 1.34% of [HC2H3O2] = 1.34% of 0.100 M = 0.00134 M pH = -log [H3O+] = -log (0.00134) = 2.87 E. Conjugate Acid/Base Pairs Ex/ HNO2 + H2O ⇌ H3O+ + NO2acid base 1. In reverse reaction, H3O+ acts as the acid a. H3O+ is called a conjugate acid b. NO2- is called a conjugate base 2. As written, HNO2 is an acid and NO2- is its conjugate base a. C.B. is what’s left after H+ is donated from an acid (acid – H+) b. C.A. is what’s formed after H+ is accepted by base (base + H+) Ex/ H3PO4 + NH3 ⇌ H2PO4- + NH4+ acid base CB CA III. Predicting direction of acid/base reactions A. Acids and bases react to form more of the weaker acid and the weaker base B. Stronger the acid, the weaker its conjugate base 1. Cases with strong acids are easy Ex/ HCl + H2O  H3O+ + ClHCl is a strong acid, so that means stronger than H3O+. Cl- is a very weak base (no ability to accept H+), so weaker base than water. Therefore, the reaction proceeds to the right. Cl- has to be a very weak base. Bases accept H+. If Cl- accepted an H+ then it would form HCl. But strong acids like HCl completely dissociate. So . . . Cl- must have no ability to accept an H+. 2. What about reactions of weak acids and weak bases? Ex/ H2CO3 + CN- ⇌ HCN + HCO3Which direction is favored? (need a chart or more information) Formation of weaker acid and weaker base H2CO3 is a stronger acid than HCN, so equilibrium favors the right CN- is a stronger base than HCO3-, so equilibrium favors the right The side with the weaker acid always has the weaker base!!! Ex/ HC2H3O2 + HSO3- ⇌ C2H3O2- + H2SO3 HC2H3O2 is a weaker acid than H2SO3, so equil favors the left HSO3- is a weaker base than C2H3O2-, so equil favors the left

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C. pKA 1. How can you determine acid strength without using a chart? 2. There is a term referred to as the acid dissociation constant or KA For the reaction: acid + H2O ⇌ H3O+ + CB 3. KA = [H3O+][CB]/[acid] 4. The bigger the value of KA, the stronger the acid 5. This term has exponents, so we use pKA = -log KA 6. The smaller the value of pKA, the stronger the acid (pH is a property of a solution, pKA is a property of an acid)

Stronger

Weaker

Acid H2SO3 HSO4HNO2 HF HC2H3O2 H2CO3 HClO NH4+ HCN

pKA 1.81 1.92 3.37 3.45 4.75 6.37 7.53 9.25 9.31

Ex/ HC2H3O2 + F- ⇌ 1. Complete the reaction HC2H3O2 + F- ⇌ C2H3O2- + HF 2. Name the acids HC2H3O2 and HF 3. Which is stronger? HF 4. Which side does equilibrium favor? Reaction favorably proceeds from right to left because HF is the stronger acid IV. Acid/Base Strength A. Review 1. Strong acids and strong bases completely donate or accept H+ when added to water 2. Weak acids and weak bases partially donate or accept H+ 3. Very weak acids and very weak bases (neutral) have essentially no ability to accept or donate H+. It is a term used in the text. In reality, these are not really acids or bases. B. Look at weak acids 1. When added to water, make the solution acidic 2. Weaker than H3O+ 3. In H2O, the rxn does not go 100%, but only partially 4. When a weak acid is added to water, [H3O+] increases compared to pure H2O, so pH is lowered, results in more acidic

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Ex/ HF + H2O ⇌ H3O+ + FH3O+ stronger than HF, so equil lies toward left, so [HF] > [F-] C. Look at weak bases 1. When added to water, make the solution basic 2. Weaker than OH3. In H2O, the rxn does not go 100%, but only partially 4. When a weak base is added to water, [OH-] increases compared to pure water, so increase pH, results in more basic Ex/ F- + H2O ⇌ OH- + HF OH- stronger than F-, so equil lies toward the left, so [F-] > [HF] D. Very weak acids and bases (neutral) 1. When salts made of very weak acids and bases are added to water, they do not alter pH 2. Very weak bases are the CB of strong acids (Cl-, NO3-, ClO4-) and very weak acids are the CA of strong bases (and alkali metals/alkaline earth metals, except Be2+) 3. When placed in water these materials do not accept or donate H+ 4. How do you recognize the “neutral salts”? a. Cl-, Br-, I-, ClO4-, NO3- with alkali metals/alkaline earth metals (except Be2+) b. Not HSO4-, Why? has an extra H+ to donate (diprotic) c. Ex/ KCl, Mg(ClO4)2, Ca(NO3)2, etc. V. Buffers A. Solutions which resist changes in pH B. Comprised of a weak acid and its conjugate base, or a weak base and its conjugate acid C. Buffers work by having a substance in solution that is available to react with any added H3O+ or OH- (weak bases and weak acids are logical candidates, strong acids and bases are not present in the molecular form) C. Reacts with added H+ and OH1. pH depends on [H3O+] 2. [H3O+] is related to [OH-] Ex/ H2O HCl + H2O  H3O+ + Clso ↓ pH

Ex/ Buffer H2CO3/HCO3HCl + HCO3-  H2CO3 + Clso ↓ pH a little

NaOH + H2O  Na+ + OH- + H2O so ↑ pH

NaOH + H2CO3 ⇌ Na+ + HCO3- + H2O so ↑ pH only a little

D. Buffers are vital to life and many reaction processes 1. pH of blood maintained by the carbonic acid-bicarbonate buffer 2. pH of cells maintained by the H2PO4--HPO42- buffer system

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E. When amount of acid or base added exceeds amounts of buffer, the pH will change rapidly F. Buffer capacity is the amount of acid or base that can be added w/o significantly changing pH (greater buffering capacity with higher concentrations of buffer acid and base) G. pH of a buffer solution is not necessarily neutral (pH 7), in fact is very rarely neutral. It depends on the buffer system (i.e., the acid and base) and buffers best at pH around the pKA

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