chapter 16. acid-base equilibrium (most important topic ... · chapter 16. acid-base equilibrium 1...

Chapter 16. Acid-Base Equilibrium

1

(Most important topic – connected to Ch 14, 16, 17, 5, 19, &

20)

16.1 Acids and Bases

SG Questions – Read p. 670-672. Answer the following questions on separate paper

1. What are the physical properties of acids?

2. What are the physical properties of bases?

3. How do acids react when dissolved in water?

4. How do bases react when dissolved in water?

5. How can you differentiate between Arrhenius acids and bases?

• Acids taste sour and cause certain dyes to change color. • Bases taste bitter and feel soapy. • Arrhenius concept of acids and bases: • An acid is a substance that, when dissolved in water, increases the concentration of H

+ ions.

• Example: HCl is an acid.

• An Arrhenius base is a substance that, when dissolved in water, increases the concentration of

OH–

ions.

• Example: NaOH is a base.

NOT ALL ACIDS BEGIN WITH HYDROGEN (H+) AND NOT ALL BASES END IN HYDROXIDE (OH-)

2

16.2 Bronsted-Lowry Acids and Bases


1. What occurs in acid - base reactions according to BrØnsted-Lowry?

2. How is a hydronium ion formed?

3. Why are H+ and H3O

+ interchangeable?

4. What is a BrØnsted-Lowry acid?

5. What is a BrØnsted-Lowry base?

6. What is always present in the transfer of protons?

7. Why are some substances amphiprotic?

8. What must occur in forward and reverse reactions to achieve acid - base

equilibrium?

9. What is a conjugate acid – base pair?

10. How is a conjugate base formed?

11. How is a conjugate acid formed?

12. How are the strength of acids and bases correlated to their conjugate pairs?

13. What is favored in acid – base equilibrium?

Brønsted-Lowry Acids and Bases

• We can use a broader, more general definition for acids and bases that is based on the fact that acid-base reactions involve proton transfers.

The H+ Ion in Water

• The H+(aq) ion = a proton with no surrounding valence electrons.

• H3O+(aq) - hydronium ion.

• Generally we use H+(aq) and H3O

+(aq) interchangeably.

Proton-Transfer Reactions

• • According to the Arrhenius definitions, an acid increases [H

+] and a base increases [OH

–].

• Another definition of acids and bases was proposed by Brønsted and Lowry.

• In the Brønsted-Lowry system, a Brønsted-Lowry acid is a species that donates H+ and a Brønsted-

Lowry base is a species that accepts H+.

• Therefore a Brønsted-Lowry base does not need to contain OH–.

• NH3 is a Brønsted-Lowry base, but not an Arrhenius base. • Consider NH3(aq) + H2O(l) NH4

+(aq) + OH

–(aq):

3

• H2O donates a proton to ammonia. • Therefore, water is acting as an acid. • NH3 accepts a proton from water. • Therefore, ammonia is acting as a base.

An amphiprotic substance can behave either as an acid or as a base. • water is an example of an amphiprotic species.

Conjugate Acid-Base Pairs

• Whatever is left of the acid after the proton is donated is called its conjugate base.

• Similarly, a conjugate acid is formed by adding a proton to the base. • Consider HX(aq) + H2O(l) H3O

+(aq) + X

–(aq):

HX and X

– differ only in the presence or absence of a proton.

• They are said to be a conjugate acid-base pair. • X

– is called the conjugate base.

• After HX (acid) loses its proton it is converted into X–

(base). • Therefore HX and X

– are a conjugate acid-base pair.

• After H2O (base) gains a proton it is converted into H3O+ (acid).

• H3O+ is the conjugate acid.

• Therefore, H2O and H3O+ are a conjugate acid-base pair.

4

Practice Problem #1: Identifying Conjugate Acids and Bases

a) What is the conjugate base of the following acids: HClO4, H2S, PH4+, HCO3

-

b) What is the conjugate acid of the following bases: CN-, SO4

-2, H2O, HCO3

- , HSO3

-, F

-,

PO4 -3

, CO

Practice Problem #2: Writing equations for acid-base reactions

a) Hydrogen sulfite ion (HSO3 –) is amphiprotic (means that it can act as an acid or base)

Write 2 equations for the reaction HSO3 – with water, in which the ion acts as an acid AND

base. In both equations, identify the conjugate acid-base pairs

Relative Strengths of Acids and Bases

• The stronger an acid is, the weaker its conjugate base will be.

• We can categorize acids and bases according to their behavior in water. • 1. Strong acids completely dissociates in aqueous solutions. • Their conjugate bases have negligible tendencies to become protonated.

• An example is HCl. • 2. Weak acids only partially dissociate in aqueous solution. • They exist in solution as a mixture of molecules and ions.

• These conjugate bases are weak bases. • Example: Acetic acid is a weak acid; acetate ion (conjugate base) is a weak base.

5

• 3. Substances with negligible acidity ( have H’s but done act like acids) do not transfer a proton to water.

• An example is CH4.

SEE P. 676 for Acid/Base Strengths to predict which side is favored in equilibrium!!!

• In every acid-base reaction, the position of the equilibrium favors the transfer of a proton from the

stronger acid to the stronger base.

***Look at the acid (or base) on each side of equation to see who is stronger = equilibrium shifts from the stronger acid to the weaker acid. (weaker side gets help from stronger side to re-establish equilibrium)

Practice Problem #3

For the following, use p. 676 to predict whether the equilibrium lies to the left (Kc < 1) or to the right (Kc >1)

a) HSO4 –

(aq) + CO3 -2

(aq) SO4 -2

(aq) + HCO3 – (aq)

b) HPO4 2–

(aq) + H2O (l) H2PO4 - (aq) + OH

– (aq)

c) NH4+

(aq) + OH- (aq) NH3 (aq) + H2O (l)

6

16.3 Autoionization of water (Kw)


1. How does water react in the presence of acids and bases?

2. What is autoionization?

3. How can the equilibrium-constant expression be used?

4. What is the value of the ion-product constant for water?

5. What must occur for a solution to be neutral?

6. What must the product of the H+ and OH

- concentrations always equal?

• In pure water the following equilibrium is established:

2H2O(l) H3O+(aq) + OH

–(aq)

• This process is called the autoionization of water.

The Ion Product of Water

• We can write an equilibrium constant expression for the autoionization of water. • Because H2O(l) is a pure liquid, we exclude it from the expression:

• Kw is called the ion-product constant.

• At 25oC the ion-product of water is:

NEED TO KNOW!!!

• A solution is neutral if [OH–] = [H3O

+].

• If the [H3O+] > [OH

–], the solution is acidic.

• If the [H3O+] < [OH

–], the solution is basic.

wc KOHOHK

3

CoonlyOHOHKw [email protected] 3

14

7

Practice Problem #4

Indicate whether the following solutions with each of the following concentration are neutral, acidic, or basic:

a) [H+] = 4 x 10-9

M

b) [OH-] = 1 x 10-7

M

Practice Problem #5

Calculate the [H+] concentration for the following at 25oC. Also, indicate whether the

following solutions are neutral, acidic, or basic:

a) [OH-] = 0.010 M

Practice Problem #6

Calculate the [OH-] concentration for the following at 25oC. Also, indicate whether the


a) [H+] = 2.1 x 10 -6

M

8

16.4 The pH Scale


1. How is the concentration of H

+ usually expressed?

2. What is the pH range for acidic solutions?

3. What is the pH range for basic solutions?

4. How are the concentrations of H+ and OH- related?

5. How does pOH differ from pH?

6. How is the pH of a solution measured?

pH scale: 0-14

0-6.99 Acidic (strong acids pH 0-2 and weak acids pH 3-6.9)

7.00 Neutral

7.01-14.00 Basic (weak bases pH 7.1-11 and strong bases pH 12-14)

pOH scale: 0-14

0-6.99 Basic

7.00 Neutral

7.01-14.00 Acidic

Calculating pH & [H+]: pH = log[H+] = log[H3O

+] [H+] = 10

-pH

• Thus, a change in [H

+] by a factor of 10 causes the pH to change

by 1 unit. • In neutral solutions at 25

oC, pH = 7.00

• In acidic solutions, [H

+] > 1.0 x 10

–7, so pH < 7.00.

• As the pH decreases, the acidity of the solution increases.

• In basic solutions, [H

+] < 1.0 x 10

–7, so pH > 7.00.

• As the pH increases, the basicity of the solution increases (acidity decreases).

9

SEE p. 682 for common acids and bases

Other “p” Scales

Calculating pOH & [OH-]: pOH = log[OH

–] [OH-] = 10

-pOH

pH & pOH have no units once you take the log of the [H+] & [OH-] concentrations • Recall that the value of Kw at 25

oC is 1.0 x 10

–14.

• Thus, we can describe a relationship between pH and pOH:

pH + pOH = 14.00

Practice Problem #7

Calculate the pH & pOHfor the following concentration at 25oC. Also, indicate whether the


a) [H+] = 5.6 x 10-6

M

10

Practice Problem #8

Calculate the [H+] concentration for the following at 25oC. Also, indicate whether the


a) pH = 3.76

b) pH = 9.18

Measuring pH

• The most accurate method to measure pH is to use a pH meter. • However, certain dyes change color as pH changes – called acid-base indicators.

• Indicators are less precise than pH meters. • Most acid-base indicators can exist as either an acid or a base.

• These two forms have different colors. • Thus, if we know the pH at which the indicator turns color, we can use this color change to

determine whether a solution has a higher or lower pH than this value. • Some natural products can be used as indicators. (Tea is colorless in acid and brown in base; red

cabbage extract is another natural indicator.)

Common Acid/Base Indicators (p. 684) – NEED TO KNOW!!!

Indicator Color Change pH range of color change

Methyl orange red to yellow 3 - 5

Methyl red red to yellow 4.5-6

Bromothymol blue yellow to blue 6-7.5

Phenolphthalein colorless to pink 8-10

Litmus (less precise) red to blue 6-8

11

pH paper (less precise) red to blue 2-12

16.5 Strong Acids and Bases


1. What are the seven most common strong acids?

2. What are the eight most common strong bases?

3. Why are the properties of strong acids and bases?

4. What types of arrows are used in equations involving strong acids and bases?

Strong Acids and Bases

Strong Acids (pH range 0-2)

• 7 strong acids: HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4. • Strong acids are strong electrolytes - dissociate (ionize) completely ***only one arrow for Strong Acids rxns*** • Example: Nitric acid ionizes completely in aqueous solution.

HNO3(aq) + H2O(l) H3O+(aq) + NO3

–(aq)

Or

12

• Since H+ and H3O

+ are used interchangeably, we write

HNO3(aq) H+(aq) + NO3

–(aq)

Strong Bases

• 8 strong bases: alkali metals (LiOH, NaOH, KOH, RbOH, CsOH) and heavier alkaline earth metals (Ca(OH)2 , Sr(OH)2 ,Ba(OH)2).

• Strong bases are strong electrolytes - dissociate (ionize) completely ***only one arrow for Strong Bases rxns*** For example:

NaOH(aq) Na+(aq) + OH

–(aq)

• Not all bases contain the OH

– ion.

• Ionic metal oxides, hydrides, and nitrides are basic.

O2–

(aq) + H2O(l) 2OH–(aq)

H–(aq) + H2O(l) H2(g) + OH

–(aq)

N3–

(aq) + 3H2O(l) NH3(aq) + 3OH–(aq)

Practice Problem #9: calculating pH of a strong base

Calculate the pH the following at 25oC.

a) 0.028M NaOH

b) 0.0011M Ca(OH)2

Practice Problem #10:

Calculate the concentration of the following at 25oC.

a) KOH for a pH = 11.89

13

b) Ca(OH)2 for a pH = 11.68

16.6 Weak Acids and Bases


1. What are the most acidic substances?

2. Why is water omitted from the equilibrium-constant expression?

3. What is the acid-dissociation constant?

4. What does a larger Ka value signify?

5. How is Ka calculated from pH?

6. How is perfect ionization used to measure acid strength?

7. What are polyprotic acids?

8. Why is it easier to remove the first proton from a polyprotic

Weak Acids

• Weak acids are only partially ionized in aqueous solution.

• There is a mixture of ions and un-ionized acid in solution. • Therefore, weak acids are in equilibrium:

HA(aq) + H2O(l) H3O+(aq) + A

–(aq)

• Or: HA(aq) H+

(aq) + A–(aq)

• We can write an equilibrium constant expression for this dissociation:

• Ka is called the acid-dissociation constant.

• Note that the subscript “a” indicates that this is the equilibrium constant for the dissociation of an acid.

• Note that [H2O] is omitted from the Ka expression. (H2O is a pure liquid.)

HA

AHK

HA

AOHK aa

or 3

14

• The larger the Ka, the stronger the acid. • If Ka >> 1, then the acid is completely ionized and the acid is a strong acid.

ICE IS BACK!!!

Calculating Ka from pH

• In order to find the value of Ka, we need to know all of the equilibrium concentrations. • The pH gives the equilibrium concentration of H

+.

• Thus, to find Ka we use the pH to find the equilibrium concentration of H+ and then use the

stoichiometric coefficients of the balanced equation to help us determine the equilibrium concentration of the other species.

• We then substitute these equilibrium concentrations into the equilibrium constant expression and solve for Ka.

Practice Problem #11

A 0.10M of formic acid (HCHO2) and the pH is measured to be 2.38 at 25oC.

a) Calculate the Ka of formic acid

b) What is the percentage of the acid is ionized in this solution?

.

15

Using Ka to Calculate pH

• Using Ka, we can calculate the concentration of H+ (and hence the pH).

• Write the balanced chemical equation clearly showing the equilibrium.

• Write the equilibrium expression. Look up the value for Ka (in a table). • Write down the initial and equilibrium concentrations for everything except pure water. • We usually assume that the equilibrium concentration of H

+ is x.

• Substitute into the equilibrium constant expression and solve. • Remember to convert x to pH if necessary.

• What do we do if we are faced with having to solve a quadratic equation in order to determine the

value of x? • Often this cannot be avoided. • However, if the Ka value is quite small, we find that we can make a simplifying assumption. • Assume that x is negligible compared with the initial concentration of the acid. • This will simplify the calculation.

• It is always necessary to check the validity of any assumption.

Check: to see how large it is compared with the initial

concentration. • If x is <5% of the initial concentration, the assumption is probably a good one. • If x>5% of the initial concentration, then it may be best to solve the quadratic equation or use

successive approximations.

• Percent ionization relates the equilibrium H

+ concentration, [H

+]equilibrium, to the initial HA

concentration, [HA]initial.

• The higher the percent ionization is, the stronger the

acid.

• However, we need to keep in mind that percent

ionization of a weak acid decreases as the molarity of the

solution increases.

100HA

ionization %mequilibriu

initial

H

16


Calculate the pH of a 0.20M solution of HCN.

Polyprotic Acids

• Polyprotic acids have more than one ionizable proton. • The protons are removed one reaction at a time. • Consider the weak acid, H2SO3 (sulfurous acid):

H2SO3(aq) H+(aq) + HSO3

–(aq) Ka1 = 1.7 x 10

–2

HSO3–(aq) H+

(aq) + SO32–

(aq) Ka2 = 6.4 x 10–8

• where Ka1 is the dissociation constant for the first proton released, Ka2 is for the second, etc.. • It is always easier to remove the first proton than the second proton in

a polyprotic acid. • Therefore, Ka1 > Ka2 > Ka3, etc.. • The majority of the H

+(aq) at equilibrium usually comes from the first ionization (i.e., the Ka1

equilibrium).

17

16.7 Weak Bases


1. How do weak bases react with water?

2. What does the base-dissociation constant always refer to?

3. What are amines?

4. What are the two types of weak bases?

Weak Bases

• Weak bases remove protons from substances.

• There is an equilibrium between the base and the resulting ions: Weak base + H2O(l) conjugate acid + OH

–(aq)

• Example: NH3(aq) + H2O(l) NH4

+(aq) + OH

–(aq).

• The base-dissociation constant, Kb, is defined as

• The larger Kb, the stronger the base.


Calculate the concentration of OH- in a 0.15M solution of NH3

][ 3

4

NH

OHNHKb

18

16.8 Relationship between Ka and Kb


Relationship Between Ka and Kb

1. What occurs to the conjugate base as the acid gets stronger?

2. How do Ka and Kb relate to Kw?

3. What is the product of the acid-dissociation and base dissociation constants?

4. What is the relationship between conjugate base pairs?

5. How can Kb be calculated from Ka?

• The net reaction is the autoionization of water.

H2O(l) H+(aq) + OH

–(aq)

• Recall that: Kw = [H

+][OH

–]

• We can use this information to write an expression that relates the values of Ka, Kb, and Kw for a

conjugate acid-base pair: Ka x Kb = Kw Kw = 1.00 x 10

-14

• Alternatively, we can express this as: pKa + pKb = pKw = 14.00 (at 25

oC)

pKa = -log Ka pKb = -log Kb

Important: The larger Ka (and the smaller pKa), the smaller Kb (and the larger pKb). • The stronger the acid, the weaker its conjugate base and vice versa.

19

16.9 Acid- Base Properties of Salt Solutions


1. What do we assume about salts?

2. What characterizes a hydrolysis reaction?

3. What are anions usually considered as in solutions?

4. What occurs when a solution has an ionizable proton?

5. How are polyatomic cations significant in reactions?

6. How do you determine which salts form acid, basic or neutral solutions?

• Nearly all salts are strong electrolytes. • Therefore, salts in solution exist entirely of ions. • Remember Acid-base reactions: Acid + Base salt + water • This process is called hydrolysis.

An Anion’s Ability to React with Water

• Anions (X-) from weak acids are basic. • They will cause an increase in pH. • Anions from strong acids are neutral. • They do not cause a change in pH.

• If Ka > Kb the anion (more acidic) will tend to decrease the pH. If Kb > Ka the anion (more basic) will tend to increase the pH.


List the following solutions in order of increasing pH:

Ba(C2H3O2)2 NH4Cl NH3CH3Br KNO3

20

Combined Effect of Cation and Anion in Solution

• Salts derived from a strong acid and a strong base are neutral. • Examples are NaCl and Ca(NO3)2.

• Salts derived from a strong base and a weak acid are basic. • Examples are NaClO and Ba(C2H3O2)2.

• Salts derived from a weak base and a strong acid are acidic. • An example is NH4Cl.

• Salts derived from a weak acid and a weak base can be either acidic or basic. • Equilibrium rules apply!

• We need to compare Ka and Kb for hydrolysis of the anion and the cation. • For example, consider NH4CN. • The Ka of NH4

+ is smaller than the Kb of CN

–, so the solution should be basic.


In each of the following, indicate which salt will form the more acidic solution

a) NaNO3 or Fe(NO3)3

b) KBr or KBrO

c) CH3NH3Cl or BaCl2

d) NH4NO3 or NH4NO3

21

16.10 Acid-Base Behavior & Chem Structure

SG Questions – Read p. 705-709 Answer the following questions on separate paper

1. What must occur when a substance is dissolved in water?

2. How does the chemical structure of a substance determine its behavior?

3. What factors determine whether a molecule containing an H-A bond will donate a

proton?

4. What determines the acidity of binary acids?

5. What are oxyacids?

6. What trends are associated with the acidity of a substance?

7. How does the strength of an acid increase?

8. What are carboxylic acids?

Factors That Affect Acid Strength

Binary Acids

• The H–X bond strength tends to decrease as the element X increases in size. • Acid strength increases down a group; base strength decreases down a group.

• H–X bond polarity is important in determining relative acid strength in any period of the periodic

table. • Acid strength increases and base strength decreases from left to right across a period as the

electronegativity of X increases.

Oxyacids

• Many acids contain one or more O–H bonds. • Acids that contain OH groups (and often additional oxygen atoms) bound to the central atom are

called oxyacids.

• All oxyacids have the general structure Y–O–H.

22

• For oxyacids with the same number of OH groups and the same number of oxygen atoms:

• acid strength increases with increasing electronegativity of the central atom, Y.

• Example: HClO > HBrO > HIO

• For oxyacids with the same central atom, Y:

• acid strength increases as the number of oxygen atoms attached to Y increases.

• Example: HClO4 > HClO3 > HClO2 > HClO


In each of the following, indicate which compound that leads to the more acidic (or less basic) solution

a) HBr or HF

b) PH3 or H2S

c) HNO2 or HNO3

d) H2SO3 or H2SeO3

chapter 16. acid-base equilibrium (most important topic ... · chapter 16. acid-base equilibrium 1...

Documents