Chapter 16 Acids and Bases

Chapter 16 Acids and Bases
Chemistry, The Central Science, 11th edition Theodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten Chapter 16 Acids and Bases John D. Bookstaver St. Charles Community College Cottleville, MO  2009, Prentice-Hall, Inc.

Some Definitions Arrhenius
An acid is a substance that, when dissolved in water, increases the concentration of hydrogen ions. A base is a substance that, when dissolved in water, increases the concentration of hydroxide ions.  2009, Prentice-Hall, Inc.

Some Definitions Brønsted-Lowry An acid is a proton donor.
A base is a proton acceptor.  2009, Prentice-Hall, Inc.

A Brønsted-Lowry acid… …must have a removable (acidic) proton.
A Brønsted-Lowry base… …must have a pair of nonbonding electrons.  2009, Prentice-Hall, Inc.

If it can be either… HCO3- HSO4- H2O …it is amphiprotic.
 2009, Prentice-Hall, Inc.

Properties of Acids Corrosive Taste sour
Form solutions with a pH < 7 React with active metals to form a salt and hydrogen gas i.e. H2SO4(aq) + Mg(s)  MgSO4(q) + H2(g)  2009, Prentice-Hall, Inc.

Properties of Acids (cont)
React with bases such as metal oxides and hydroxides to form a salt and water i.e. 2HNO3(aq) + CuO(s)  Cu(NO3)2(aq) + H2O(l) React with metal carbonates and hydrogen carbonates to give a salt, water, and carbon dioxide i.e. 2HCl(aq) + ZnCO3(s)  ZnCl2(aq) + H2O(l) + CO2(g)  2009, Prentice-Hall, Inc.

Properties of Bases Most common are oxides, hydroxides, and carbonates of metals Also include substances such as ammonia and amines Also called alkalis Feel slippery Taste bitter  2009, Prentice-Hall, Inc.

Properties of Bases (cont)
Give a pH > 7 when soluble in water React with acids to form a salt i.e. CaO(s) + 2HCl(aq)  CaCl2(aq) + H2O(l)  2009, Prentice-Hall, Inc.

What Happens When an Acid Dissolves in Water?
Water acts as a Brønsted-Lowry base and abstracts a proton (H+) from the acid. As a result, the conjugate base of the acid and a hydronium ion are formed.  2009, Prentice-Hall, Inc.

Conjugate Acids and Bases
The term conjugate comes from the Latin word “conjugare,” meaning “to join together.” Reactions between acids and bases always yield their conjugate bases and acids.  2009, Prentice-Hall, Inc.

Acid and Base Strength 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.  2009, Prentice-Hall, Inc.

Acid and Base Strength Substances with negligible acidity do not dissociate in water. Their conjugate bases are exceedingly strong.  2009, Prentice-Hall, Inc.

Acid and Base Strength In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base. HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq) H2O is a much stronger base than Cl-, so the equilibrium lies so far to the right that K is not measured (K>>1).  2009, Prentice-Hall, Inc.

Acid and Base Strength In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base. CH3CO2H (aq) + H2O (l) H3O+ (aq) + CH3CO2- (aq) Acetate is a stronger base than H2O, so the equilibrium favors the left side (K<1).  2009, Prentice-Hall, Inc.

Autoionization of Water
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. H2O (l) + H2O (l) H3O+ (aq) + OH- (aq)  2009, Prentice-Hall, Inc.

Ion-Product Constant The equilibrium expression for this process is
AHL ONLY!! Ion-Product Constant 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  2009, Prentice-Hall, Inc.

pH pH is defined as the negative base-10 logarithm of the concentration of hydronium ion. pH = -log [H3O+]  2009, Prentice-Hall, Inc.

pH In pure water, Kw = [H3O+] [OH-] = 1.0  10-14
AHL ONLY!! pH In pure water, Kw = [H3O+] [OH-] = 1.0  10-14 Since in pure water [H3O+] = [OH-], [H3O+] =  = 1.0  10-7  2009, Prentice-Hall, Inc.

pH Therefore, in pure water, pH = -log (1.0  10-7) = 7.00
An acid has a higher [H3O+] than pure water, so its pH is <7. A base has a lower [H3O+] than pure water, so its pH is >7.  2009, Prentice-Hall, Inc.

pH These are the pH values for several common substances.

AHL ONLY!! Other “p” Scales The “p” in pH tells us to take the negative base-10 logarithm of the quantity (in this case, hydronium ions). Some similar examples are pOH: -log [OH-] pKw: -log Kw  2009, Prentice-Hall, Inc.

-log [H3O+] + -log [OH-] = -log Kw = 14.00
AHL ONLY!! Watch This! Because [H3O+] [OH-] = Kw = 1.0  10-14, we know that -log [H3O+] + -log [OH-] = -log Kw = 14.00 or, in other words, pH + pOH = pKw = 14.00  2009, Prentice-Hall, Inc.

How Do We Measure pH? For less accurate measurements, one can use
Litmus paper “Red” paper turns blue above ~pH = 8 “Blue” paper turns red below ~pH = 5 Or an indicator.  2009, Prentice-Hall, Inc.

How Do We Measure pH? For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.  2009, Prentice-Hall, Inc.

Strong Acids You will recall that the seven strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4. These are, by definition, strong electrolytes and exist totally as ions in aqueous solution. For the monoprotic strong acids, [H3O+] = [acid].  2009, Prentice-Hall, Inc.

Strong Bases Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal hydroxides (Ca2+, Sr2+, and Ba2+). Again, these substances dissociate completely in aqueous solution.  2009, Prentice-Hall, Inc.

Dissociation Constants
For a generalized acid dissociation, the equilibrium expression would be This equilibrium constant is called the acid-dissociation constant, Ka. HA (aq) + H2O (l) A- (aq) + H3O+ (aq) [H3O+] [A-] [HA] Kc =  2009, Prentice-Hall, Inc.

Dissociation Constants
The greater the value of Ka, the stronger is the acid.  2009, Prentice-Hall, Inc.

Calculating Ka from the pH
The pH of a 0.10 M solution of formic acid, HCOOH, at 25C is Calculate Ka for formic acid at this temperature. We know that [H3O+] [COO-] [HCOOH] Ka =  2009, Prentice-Hall, Inc.

The pH of a 0.10 M solution of formic acid, HCOOH, at 25C is Calculate Ka for formic acid at this temperature. To calculate Ka, we need the equilibrium concentrations of all three things. We can find [H3O+], which is the same as [HCOO-], from the pH.  2009, Prentice-Hall, Inc.

pH = -log [H3O+] 2.38 = -log [H3O+] -2.38 = log [H3O+] = 10log [H3O+] = [H3O+] 4.2  10-3 = [H3O+] = [HCOO-]  2009, Prentice-Hall, Inc.

Calculating Ka from pH Now we can set up a table… [HCOOH], M [H3O+], M
[HCOO-], M Initially 0.10 Change - 4.2  10-3 + 4.2  10-3 At Equilibrium  10-3 = = 0.10 4.2  10-3  2009, Prentice-Hall, Inc.

Calculating Ka from pH [4.2  10-3] [4.2  10-3] Ka = [0.10]
= 1.8  10-4  2009, Prentice-Hall, Inc.

Calculating Percent Ionization
[H3O+]eq [HA]initial Percent Ionization =  100 In this example [H3O+]eq = 4.2  10-3 M [HCOOH]initial = 0.10 M 4.2  10-3 0.10 Percent Ionization =  100 = 4.2%  2009, Prentice-Hall, Inc.

HC2H3O2 (aq) + H2O (l) H3O+ (aq) + C2H3O2- (aq)
Calculating pH from Ka Calculate the pH of a 0.30 M solution of acetic acid, HC2H3O2, at 25C. HC2H3O2 (aq) + H2O (l) H3O+ (aq) + C2H3O2- (aq) Ka for acetic acid at 25C is 1.8  10-5.  2009, Prentice-Hall, Inc.

Calculating pH from Ka The equilibrium constant expression is
[H3O+] [C2H3O2-] [HC2H3O2] Ka =  2009, Prentice-Hall, Inc.

Calculating pH from Ka We next set up a table…
[C2H3O2], M [H3O+], M [C2H3O2-], M Initially 0.30 Change -x +x At Equilibrium x  0.30 x We are assuming that x will be very small compared to 0.30 and can, therefore, be ignored.  2009, Prentice-Hall, Inc.

Calculating pH from Ka (x)2 (0.30) 1.8  10-5 =
Now, (x)2 (0.30) 1.8  10-5 = (1.8  10-5) (0.30) = x2 5.4  10-6 = x2 2.3  10-3 = x  2009, Prentice-Hall, Inc.

Calculating pH from Ka pH = -log [H3O+] pH = -log (2.3  10-3)

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.  2009, Prentice-Hall, Inc.

Bases react with water to produce hydroxide ion.
Weak Bases Bases react with water to produce hydroxide ion.  2009, Prentice-Hall, Inc.

Weak Bases The equilibrium constant expression for this reaction is
[HB] [OH-] [B-] Kb = where Kb is the base-dissociation constant.  2009, Prentice-Hall, Inc.

Kb can be used to find [OH-] and, through it, pH.
Weak Bases Kb can be used to find [OH-] and, through it, pH.  2009, Prentice-Hall, Inc.

pH of Basic Solutions What is the pH of a 0.15 M solution of NH3?
NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq) [NH4+] [OH-] [NH3] Kb = = 1.8  10-5  2009, Prentice-Hall, Inc.

pH of Basic Solutions Tabulate the data. [NH3], M [NH4+], M [OH-], M
Initially 0.15 At Equilibrium x  0.15 x  2009, Prentice-Hall, Inc.

pH of Basic Solutions (x)2 (0.15) 1.8  10-5 =

pH of Basic Solutions Therefore, [OH-] = 1.6  10-3 M
pOH = -log (1.6  10-3) pOH = 2.80 pH = pH = 11.20  2009, Prentice-Hall, Inc.

Ka and Kb Ka and Kb are related in this way: Ka  Kb = Kw
AHL ONLY!! Ka and Kb Ka and Kb are related in this way: Ka  Kb = Kw Therefore, if you know one of them, you can calculate the other.  2009, Prentice-Hall, Inc.

Reactions of Anions with Water
Anions are bases. As such, they can react with water in a hydrolysis reaction to form OH- and the conjugate acid: X- (aq) + H2O (l) HX (aq) + OH- (aq)  2009, Prentice-Hall, Inc.

Reactions of Cations with Water
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.  2009, Prentice-Hall, Inc.

Reactions of Cations with Water
Attraction between nonbonding electrons on oxygen and the metal causes a shift of the electron density in water. This makes the O-H bond more polar and the water more acidic. Greater charge and smaller size make a cation more acidic.  2009, Prentice-Hall, Inc.

Effect of Cations and Anions
An anion that is the conjugate base of a strong acid will not affect the pH. HCl + H2O  Cl- + H3O+ An anion that is the conjugate base of a weak acid will increase the pH. F- + H3O+ ↔ HF + H2O 3. A cation that is the conjugate acid of a weak base will decrease the pH. NH4+ H2O ↔ NH3 + H3O+  2009, Prentice-Hall, Inc.

Effect of Cations and Anions
Cations of the strong Arrhenius bases will not affect the pH. NaOH  Na+ + OH- Other metal ions will cause a decrease in pH. Fe3+ ↔ FeOH2+ + H+ When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the affect on pH depends on the Ka and Kb values.  2009, Prentice-Hall, Inc.

Factors Affecting Acid Strength
Organic compounds made of carbon, oxygen, and hydrogen are typically weak The hydrogen connected to the oxygen will ionize Hydrolysis of metal ions (section 16.11) causes acidic solutions – acid strength increases with increasing charge and decreasing radius  2009, Prentice-Hall, Inc.

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.  2009, Prentice-Hall, Inc.

In oxyacids, in which an -OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid.  2009, Prentice-Hall, Inc.

For a series of oxyacids, acidity increases with the number of oxygens.  2009, Prentice-Hall, Inc.

Resonance in the conjugate bases of carboxylic acids stabilizes the base and makes the conjugate acid more acidic.  2009, Prentice-Hall, Inc.

Factors Affecting Base Strength
Substances that react with water to form hydroxide are strong Most common of these is the oxide ion O2-(aq) + H2O(l)  2OH-(aq) ionic metal oxides (especially Na2O or CaO), ionic hydrides (H-), and nitrides (N3-) can all act as a strong base  2009, Prentice-Hall, Inc.

Any substance with an atom that has a nonbonding pair of electrons can act as a weak base – many of these will include a nitrogen or be classified as an amine (substances where one or more N—H bonds has been changed into a N—CHx bond, such as NH3 changing into NH2CH3 )  2009, Prentice-Hall, Inc.

Anions of weak acids can be weak bases because they will take a H+ from water to form the weak acid, leaving OH- in solution For example NaClO reacting with H2O will do this: ClO-(aq) + H2O(l)  HClO(aq) + OH-(aq) The Na+ spectates, so is removed  2009, Prentice-Hall, Inc.

Lewis Acids Lewis acids are defined as electron-pair acceptors.
Atoms with an empty valence orbital can be Lewis acids.  2009, Prentice-Hall, Inc.

Lewis Bases Lewis bases are defined as electron-pair donors.
Anything that could be a Brønsted-Lowry base is a Lewis base. Lewis bases can interact with things other than protons, however.  2009, Prentice-Hall, Inc.

AHL ONLY!! Salt Hydrolysis Salts are ionic compounds of cations from a base and anions from an acid that completely dissociate into ions in aqueous solutions If the acid or base is weak, the salts can affect the pH of the solution Cations can act as acids and anions as bases  2009, Prentice-Hall, Inc.

Salt Hydrolysis (cont)
AHL ONLY!! Salt Hydrolysis (cont) The stronger the conjugate acid/base the salt is derived from, the weaker the acid-base activity of the ion Cations from strong bases have little acid-base activity; anions from strong acids are the same This is why strong acids and strong bases make neutral solutions!  2009, Prentice-Hall, Inc.

AHL ONLY!! Salt Hydrolysis (cont) If the anion comes from a weak acid it will act like a weak base in solution – a strong base mixed with a weak acid will cause a pH > 7 If the cation comes from a weak base, the cation will act as a weak acid in solution – a strong acid with a weak base will form a pH < 7  2009, Prentice-Hall, Inc.

AHL ONLY!! Salt Hydrolysis (cont) Weak acids and weak bases form a solution pH that reflects the relative strength of the acid and the base – equal and opposite strength may cause a neutral, if the acid is stronger it will be acidic, and if the base is stronger it will be basic  2009, Prentice-Hall, Inc.

AHL ONLY!! Salt Hydrolysis Transition Metals and other small, highly charged cations These ions have a high charge density When these ions hydrate, the O-H bonds are weakened by the electron attracting power of the cation the hydroxide ions are stabilized Solutions of these ions are quite acidic i.e. [Fe(H2O)6]3+(aq)  [Fe(OH)(H2O)5]2+(aq) + H+(aq)  2009, Prentice-Hall, Inc.

The Common-Ion Effect Consider a solution of acetic acid:
If acetate ion is added to the solution, Le Châtelier says the equilibrium will shift to the left. CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO−(aq)  2009, Prentice-Hall, Inc.

The Common-Ion Effect “The extent of ionization of a weak electrolyte is decreased by adding to the solution a strong electrolyte that has an ion in common with the weak electrolyte.”  2009, Prentice-Hall, Inc.

The Common-Ion Effect Calculate the fluoride ion concentration and pH of a solution that is 0.20 M in HF and 0.10 M in HCl. Ka for HF is 6.8  10−4. [H3O+] [F−] [HF] Ka = = 6.8  10-4  2009, Prentice-Hall, Inc.

The Common-Ion Effect HF(aq) + H2O(l) H3O+(aq) + F−(aq)
Because HCl, a strong acid, is also present, the initial [H3O+] is not 0, but rather 0.10 M. [HF], M [H3O+], M [F−], M Initially 0.20 0.10 Change −x +x At Equilibrium 0.20 − x  0.20 x  0.10 x  2009, Prentice-Hall, Inc.

The Common-Ion Effect (0.10) (x) 6.8  10−4 = (0.20)
(0.20) (6.8  10−4) (0.10) = x 1.4  10−3 = x  2009, Prentice-Hall, Inc.

The Common-Ion Effect Therefore, [F−] = x = 1.4  10−3
[H3O+] = x =  10−3 = 0.10 M So, pH = −log (0.10) pH = 1.00  2009, Prentice-Hall, Inc.

Buffers Buffers are solutions of a weak conjugate acid-base pair.
AHL ONLY!! Buffers Buffers are solutions of a weak conjugate acid-base pair. They are particularly resistant to pH changes, even when strong acid or base is added.  2009, Prentice-Hall, Inc.

AHL ONLY!! Buffers If a small amount of hydroxide is added to an equimolar solution of HF in NaF, for example, the HF reacts with the OH− to make F− and water.  2009, Prentice-Hall, Inc.

AHL ONLY!! Buffers Similarly, if acid is added, the F− reacts with it to form HF and water.  2009, Prentice-Hall, Inc.

AHL ONLY!! Buffer Calculations Consider the equilibrium constant expression for the dissociation of a generic acid, HA: HA + H2O H3O+ + A− [H3O+] [A−] [HA] Ka =  2009, Prentice-Hall, Inc.

Buffer Calculations Rearranging slightly, this becomes
AHL ONLY!! Buffer Calculations Rearranging slightly, this becomes [A−] [HA] Ka = [H3O+] Taking the negative log of both side, we get base [A−] [HA] −log Ka = −log [H3O+] + −log pKa pH acid  2009, Prentice-Hall, Inc.

Buffer Calculations So pKa = pH − log [base] [acid]
AHL ONLY!! Buffer Calculations So pKa = pH − log [base] [acid] Rearranging, this becomes pH = pKa + log [base] [acid] This is the Henderson–Hasselbalch equation.  2009, Prentice-Hall, Inc.

Henderson–Hasselbalch Equation
AHL ONLY!! Henderson–Hasselbalch Equation What is the pH of a buffer that is 0.12 M in lactic acid, CH3CH(OH)COOH, and 0.10 M in sodium lactate? Ka for lactic acid is 1.4  10−4.  2009, Prentice-Hall, Inc.

Henderson–Hasselbalch Equation
AHL ONLY!! Henderson–Hasselbalch Equation pH = pKa + log [base] [acid] pH = −log (1.4  10−4) + log (0.10) (0.12) pH pH = 3.77 pH = (−0.08) pH  2009, Prentice-Hall, Inc.

AHL ONLY!! Titration In this technique a known concentration of base (or acid) is slowly added to a solution of acid (or base).  2009, Prentice-Hall, Inc.

AHL ONLY!! Titration A pH meter or indicators are used to determine when the solution has reached the equivalence point, at which the stoichiometric amount of acid equals that of base.  2009, Prentice-Hall, Inc.

Titration of a Strong Acid with a Strong Base
AHL ONLY!! From the start of the titration to near the equivalence point, the pH goes up slowly.  2009, Prentice-Hall, Inc.

AHL ONLY!! Just before (and after) the equivalence point, the pH increases rapidly.  2009, Prentice-Hall, Inc.

AHL ONLY!! At the equivalence point, moles acid = moles base, and the solution contains only water and the salt from the cation of the base and the anion of the acid.  2009, Prentice-Hall, Inc.

AHL ONLY!! As more base is added, the increase in pH again levels off.  2009, Prentice-Hall, Inc.

Titration of a Weak Acid with a Strong Base
AHL ONLY!! Unlike in the previous case, the conjugate base of the acid affects the pH when it is formed. At the equivalence point the pH is >7. Phenolphthalein is commonly used as an indicator in these titrations.  2009, Prentice-Hall, Inc.

AHL ONLY!! At each point below the equivalence point, the pH of the solution during titration is determined from the amounts of the acid and its conjugate base present at that particular time.  2009, Prentice-Hall, Inc.

AHL ONLY!! With weaker acids, the initial pH is higher and pH changes near the equivalence point are more subtle.  2009, Prentice-Hall, Inc.

Titration of a Weak Base with a Strong Acid
AHL ONLY!! The pH at the equivalence point in these titrations is < 7. Methyl red is the indicator of choice.  2009, Prentice-Hall, Inc.

Titrations of Polyprotic Acids
AHL ONLY!! Titrations of Polyprotic Acids When one titrates a polyprotic acid with a base there is an equivalence point for each dissociation.  2009, Prentice-Hall, Inc.

AHL ONLY!! Indicators Indicators are substances (usually organic dyes) that have different colors in acidic or alkaline solutions Are used to detect the endpoint of a titration, which is the point at which the indicator forms are in equilibrium with each other  2009, Prentice-Hall, Inc.

HIn(aq)  H+(aq) + In-(aq)
AHL ONLY!! Indicators (cont) For example, consider litmus: HIn(aq)  H+(aq) + In-(aq) Red Blue In acid solutions, the equilibrium is driven to the left, so the solution turns red In basic solutions, the equilibrium shifts to the right, so the solution turns blue  2009, Prentice-Hall, Inc.

Indicators (cont) The equilibrium equation is Ka = [H+][In-] [HIn]
AHL ONLY!! Indicators (cont) The equilibrium equation is Ka = [H+][In-] [HIn] [HIn] = [H+] (rearranged for the colored forms) [In-] Ka The rearranged equation shows that the pH AND the value of Ka determine the color change range  2009, Prentice-Hall, Inc.

AHL ONLY!! Indicators (cont) Two most commonly used indicators – methyl orange and phenolphthalein Property Phenolphthalein Methyl orange pKa 9.6 3.7 pH range 8.3 – 10.0 3.1 – 4.4 Color in acid Colorless Red Color in base Pink Yellow Useful for Titrations involving strong acids and bases Titrations involving strong acids  2009, Prentice-Hall, Inc.

AHL ONLY!! Indicators (cont) When pH = pKa the two colored forms will be equal to each other and the color change will be midway between the acid and base form So, phenolphthalein would look VERY pale pink And methyl orange, which is red in an acid and yellow in a base, would look orange   2009, Prentice-Hall, Inc.

AHL ONLY!! Indicators (cont) If the concentration of one form (for example HIn) is 10x greater than the other form (In-) the color will match the predominant species The pH at this point can be calculated like this: [H+] = Ka x [HIn] = Ka x 10/1 = 10Ka [In-] OR pH = pKa – 1  2009, Prentice-Hall, Inc.

AHL ONLY!! Indicators (cont) If the reverse were true, and In- was 10x greater than HIn, then the pH = pKa + 1 Most indicators, therefore, change color over a 2 pH unit range with the pKa value in the exact center of the change but intensity of the color rather than a direct color change must be considered for some cases (such as phenolphthalein)  2009, Prentice-Hall, Inc.

AHL ONLY!! Indicators (cont) The point at which an indicator changes color is referred to as the end point and the point at which we reach the ideal stoichiometric ratio is the equivalence point You pick the indicator that has a range that includes the pH at which you believe the equivalence point will be reached  2009, Prentice-Hall, Inc.

Chapter 16 Acids and Bases

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