Equilibrium 2

As a review from the first semester of general chemistry, the list of strong acids and bases are the most important as these are among the most common acids and bases. Know their formulas. Remember that the word "strong" means complete ionization.

1. (a) Give the conjugate base of the following Bronsted-Lowry acids: (i) HCOOH, (ii) HPO42-. (b) Give the conjugate acid of the following Bronsted-Lowry bases: (i) SO42-, (ii) CH3NH2.

2. (a) Write an equation for the reaction in which H2C6H7O5- (aq) acts as a base in H2O (l). (b) Write an equation for the reaction in which H2C6H7O5- (aq) acts as an acid in water. (c) Identify the conjugate acid-base pairs in both equations.

3. (a) A solution is labeled 0.035 M Sr(OH)2. What is the [OH-] for the solution? (b) Is the following statement true or false? Because Mg(OH)2 is not very soluble, it cannot be a strong base. Explain.

HCl is an acid and water is a base on the reactant side. If this reaction were to reverse itself, the hydronium ion could give up that extra H+ and chloride could accept it. Therefore, we call hydronium the conjugate acid of the base, water, and chloride ion is the conjugate base of the acid, HCl. However, because the acid that generated this equation to begin with is so strong, the conjugates are too weak to reverse the reaction. This example can be seen above

For weak acids and bases, the conjugates produced are stronger than the weaker acids or bases that produced them. In most of these weak acids and bases, it is the reverse reaction that is more favored and, therefore, the rule of thumb of favoring the generation of a weaker side still holds. The fact that these weak acids and bases generate an equilibrium system says that the reverse reaction is feasible and, in fact, occurs at the same rate as the forward reaction. In acid/base land that means that the conjugates produced do indeed have formidable acid/base properties themselves. This example is pictured below.

As can be seen from the picture below, all species are present in a solution of a polyprotic acid such as citric acid. The table below also shows the individual Ka's for each proton in a polyprotic acid. On the equilibrium 3 smore, polyprotic acids will be revisited pertaining to titration curves for these acids. Each proton is removed sequentially and as seen by the values for Ka1 vs Ka2 (and Ka3), each proton becomes harder to remove as removing a postive species from an anion becomes increasingly difficult.

• Strong acids are completely dissociated in water.

• Their conjugate bases are quite weak.

• Weak acids only dissociate partially in water.

• Their conjugate bases are weak bases.

• In any acid–base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base.

• Acetate is a stronger base than H2O, so the equilibrium favors the left side (K < 1).

• The more polar the H–X bond and/or the weaker the H–X bond, the more acidic the compound.

• So acidity increases from left to right across a row and from top to bottom down a group.

• In oxyacids, in which an –OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid.
  

• For a series of oxyacids, acidity increases with the number of oxygens.
  

• Resonance in the conjugate bases of carboxylic acids stabilizes the base and makes the conjugate acid more acidic.
  

• Cations with acidic protons (like NH4+) will lower the pH of a solution.

• Most metal cations that are hydrated in solution also lower the pH of the solution.

• Greater charge and smaller size make a cation more acidic.

1. (a) Which of the following is a stronger Bronsted-Lowry acid, HClO3 or HClO2? (b) Which is the stronger Bronsted-Lowry base, HS- or HSO4-? Explain your choices.

2. Explain the following observations: (a) HCl is a stronger acid than H2S; (b) H3PO4 is a stronger acid than H3AsO4; (c) HBrO3 is a stronger acid than HBrO2; (d) H2C2O4 is a stronger acid than HC2O4-; (e) benzoic acid (C6H5COOH) is stronger acid than phenol (C6H5OH).

3. Based on their compositions and structures and on conjugate acid-base relationships, select the stronger base in each of the following pairs: (a) NO3- or NO2-, (b) PO43- or AsO43-, (c) HCO3- or CO32-.

4. Identify the Lewis acid and Lewis base in each of the following reactions:

(a) HNO2 (aq) + OH- (aq) ↔ NO2- (aq) + H2O (l)

(b) FeBr3 (s) + Br- (aq) ↔ FeBr4- (aq)

(c) Zn2+ (aq) + 4NH3 (aq) ↔ Zn(NH3)42+ (aq)

(d) SO2 (g) + H2O (l) ↔ H2SO3 (aq)

5. Which member of each pair produces the more acidic aqueous solution: (a) ZnBr2 or CdCl2, (b) CuCl or Cu(NO3)2, (c) Ca(NO3)2 or NiBr2? Explain.

• As we have seen, water is amphoteric.

• In pure water, a few molecules act as bases and a few act as acids.

This is referred to as autoionization.

• The equilibrium expression for this process is

Kc = [H3O+] [OH−]

• This special equilibrium constant is referred to as the ion product constant for water, Kw.

• At 25°C, Kw = 1.0 ⋅ 10−14

In pure water, hydronium and hydroxide concentrations are equal and we call pure water neutral. When [H3O+] > [OH-] then the solution (now water is not pure ) is said to be acidic and when [H3O+] < [OH-], the solution is said to be basic. The concentration of [H3O+] or [OH-] in an acidic or basic solution will be determined by the stoichiometry of the acid or base. For example, if the concentration of HCl is 1 x 10-3 M, then the hydronium ion concentration in this solution will also be 1 x 10-3 M. If you need the hydroxide ion concentration for the same solution, use the Kw relationship: Kw = [H3O+][OH-] and the hydroxide concentration for this example is

1 x 10-14/ 1 x 10-3 = 1 x 10-11 M. See second chart below (memorize this one). It is important to realize, however, that this simple relationship between acid/base concentration and hydronium/hydroxide concentrations only works for STRONG ACIDS AND BASES. Because concentrations like 1 x 10-8 and so on are hard to write and conceptualize for many, a log of the Kw expression gives us an easier scale to use. The important thing to realize is that because the pH scale is logarithmic, every change in value of 1 on the pH scale really represents a change of 10 in concentration.

1. Calculate the pH of each of the following strong acid solutions: (a) 0.0167 M HNO3, (b) 0.225 g of HClO3 in 2.00 L of solution, (c) 15.00 mL of 1.00 M HCl diluted to 0.500 L, (d) a mixture formed by adding 50.0 mL of 0.020 M HCl to 125 mL of 0.010 M HI.

2. Calculate [OH-] and pH for each of the following strong base solutions: (a) 0.182 M KOH, (b) 3.165 g of KOH in 500.0 mL of solution, (c) 10.0 mL of 0.0105 M Ca(OH)2 diluted to 500.0 mL, (d) a solution formed by mixing 20.0 mL of 0.015 M Ba(OH)2 with 40.0 mL of 8.2 x 10-3 M NaOH.

3. Carbon dioxide in the atmosphere dissolves in raindrops to produce carbonic acid (H2CO3), causing the pH of clean, unpolluted rain to range from about 5.2 to 5.6. What are the ranges of [H+] and [OH-] in the raindrops?

The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate Ka for formic acid at this temperature.

To calculate Ka, we need the equilibrium concentrations of all three things.

ICE table!! The pH will lead you to the equilibrium concentration of H3O+ AND the conjugate base. These concentrations are the "x" in the ICE table. This concentration is also the amount of the acid (or base) that dissociated.

As with all percentages, it is part/total x 100. The total is the initial concentration of the acid or base and the "part" is the amount of hydronium or hydroxide at equilibrium which can be obtained from the pH.

1.Set up ICE table. This time you have "x" for intial concentration of hydronium or hydroxide.

2. Using initial concentration of acid or base and Ka (or Kb), solve for x using equilibrium expression.

3. THE GOOD NEWS: You can almost always use the "approximation" that [HA] - x ~ [HA] due to the small values of Ka and Kb for most weak acids and bases.

4. If you are dealing with an acid, you are one step from pH once you have solved for x. For bases, it is two steps-see calculation table above.

Polyprotic acids have more than one acidic proton.

If the difference between the Ka for the first dissociation and subsequent Ka values is 103 or more, the pH generally depends only on the first dissociation.

1. Phenylacetic acid (C6H5CH2COOH) is one of the substances that accumulates in the blood of people with phenylketonuria, an inherited disorder that can cause mental retardation and even death. A 0.085 M solution of C6H5CH2COOH has a pH of 2.68. Calculate the Ka for this acid.

2. A 0.100 M solution of bromoacetic acid (BrCH2COOH) is 13.2% ionized. Calculate [H+], [BrCH2COO-], [BrCH2COOH] and Ka for this acid.

3. Determine the pH of each of the following solutions (Ka and Kb values are given in Appendix D): (a) 0.095 M hypochlorous acid, (b) 0.0085 M hydrazine, (c) 0.165 M hydroxylamine.

4. Write the chemical equation and the Kb expression for the reaction of each of the following bases with water: (a) propylamine, C3H7NH2, (b) monohydrogen phosphate ion, HPO42-, (c) benzoate ion, C6H5CO2-.

5. Codeine (C18H21NO3) is a weak organic base. A 5.0 x 10-3 M solution of codeine has a pH of 9.95. Calculate the value of Kb for this substance. What is the pKb for this base?

• Anions are bases.

• As such, they can react with water in a hydrolysis reaction to form OH− and the conjugate acid.

• Cations with acidic protons (like NH4+) will lower the pH of a solution.Most metal cations that are hydrated in solution also lower the pH of the solution.

1. Use the acid-dissociation constants in Table 16.3 to arrange these oxyanions from strongest to weakest base: SO42-, CO32-, SO32, and PO43-.

2. Predict whether aqueous solutions of the following substances are acidic, basic, or neutral: (a) AlCl3, (b) NaBr, (c) NaClO, (d) [CH3NH3]NO3, (e) Na2SO3.

3. An unknown salt is either KBr, NH4Cl, KCN, or K2CO3. If a 0.100 M solution of the salt is neutral, what is the identity of the salt.

4. Using data from Appendix D, calculate [OH-] and pH for each of the following solutions: (a) 0.105 M NaF, (b) 0.035 M Na2S.

When an acid and a base combine, the product is most often water and a salt. Instead of forming a precipitate, molecular water is formed which drives these reactions forward. The process for writing the reaction is the same. As the picture below represents, acid + base reactions are often called neutralization reactions which can be misleading. Not all acid + base reactions will end with a neutral pH. You will learn later that the pH at the endpoint (titration term) depends on the identity of the salt.

HCl (aq) + NaOH (aq) → NaCl(aq) + H2O

This video starts with going over factors affecting acid/base strength, Lewis acids and bases. Naming of acids is revisited along the way. There are six additional problems. The questions for these problems can be found in the document linked below.

Practice covering almost the whole smore

More practice-shorter assignment than the one above

Even more practice-about the same length as #2 above

Chemistry is a subject that must be practiced everyday if possible. Take advantage of the practice in Mastering Chemistry to give you the practice you need to be successful. DO NOT PROCRASTINATE! Email me with questions!!

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