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CHP 9 SLIDES

Acids and Bases, pH, and

Buffers

Chapter 9

Guinn, Essentials of General, Organic, and Biochemistry, 3e,

9.1 Acids and Bases

9.2 pH

9.3 Acid-Base Neutralization

Reactions

9.4 Buffers

Metabolic Acidosis and Blood pH

A marathon runner stumbles into an aid station breathing

heavily and acting confused.

Nurse measures arterial blood pH and finds it below normal.

• Treatment – IV of saline plus oral sodium bicarbonate

• Result – runner recovers fairly quickly

Blood pH

Proper organ function

requires blood pH to be

between 7.35 and 7.45, acid-

base homeostasis.

• acidosis = pH < 7.35

• alkalosis = pH > 7.45

Cellular metabolism produces substances that affect blood pH

such as CO2 and H+ ions.

Blood pH maintained by

-increasing respiration to exhale more CO2, or

-shifting buffer equilibria to consume more H+

Acids and bases are found in many household

products and in various foods.

Acids produce hydronium ions, H3O+

H3O+ = polyatomic ion – water molecule, H2O, + proton, H+

Bases produce hydroxide ions, OH–

According to the Arrhenius definition, an acid is a substance

that produces hydronium ions, H3O+, and a base is a substance

that produces hydroxide ions, OH–, in aqueous solution.

Are Bases

Group 1A hydroxide salts:

NaOH, KOH

soluble in water

strong electrolytes

dissolve entirely to

produce ions in water

Group 2A hydroxide salts:

Ca(OH)2, Mg(OH)2

slightly soluble in water

weak electrolytes

dissolve slightly to produce

few ions in water

Soluble hydroxide ion salts (group 1A salts) are strong

electrolytes, and slightly soluble hydroxide ion salts (group 2A

salts) are weak electrolytes.

of Hydroxide Ions

Ionic compounds composed of hydroxide ion (OH–) are bases

because they produce solvated hydroxide ions in solution.

common product do they all produce in water?

(Answer)

hydronium ion, H3O+

compounds is a base ? Explain.

a. Ba(OH)2

b. NaBr

c. LiOH

d. CaBr2

compounds is a base ? Explain. (Answer)

Ionic compounds that are bases contain hydroxide ion, OH–

a. Ba(OH)2

b. NaBr

c. LiOH

d. CaBr2

Acid: a proton (H+) donor Base: a proton (H+) acceptor

Hydrochloric acid, HCl, donates a proton to a water molecule,

which acts as a base by accepting the proton.

Neutral acid loses proton to form conjugate base.

A conjugate base has one fewer H+ and, thus, a –1 charge

(compared to formula of acid).

HCl → Cl–

Water acts as a base, accepting proton to form conjugate acid.

A conjugate acid has one more H+ and, thus, a +1 charge

(compared to formula of base).

H2O → H3O+

According to the Brønsted-Lowry theory, when an acid is

added to water, the acid donates a proton to water to form

the conjugate base; and water accepts a proton to form

hydronium ion (H3O+), the conjugate acid of water.

Ammonia, NH3, acts as a base by accepting a proton from a

water molecule, which now acts as an acid.

(continued)

Neutral base accepts a proton to form the conjugate acid.

A conjugate acid has one more H+ and, thus, a +1 charge.

NH3 ⇄ NH4

Neutral water donates a proton to form conjugate base.

A conjugate base has one fewer H+ and, thus, a –1 charge.

H2O ⇄ OH–

According to the Brønsted-Lowry theory, a base accepts a

proton from water to produce the conjugate acid, and water

donates a proton to the base to form hydroxide ion (OH–).

According to the Brønsted-Lowry theory, water can accept or

donate a proton.

Water can act as an acid or a base.

Any substance that can act as an acid or a base is an

amphoteric compound.

Amphoteric compounds can act as either an acid or a base.

An acid and its conjugate base or a base and its conjugate acid

are called a conjugate acid-base pair.

An acid and its conjugate base are a conjugate acid-base pair.

A base and its conjugate acid are a conjugate acid-base pair.

Only certain H atoms can be donated, and these are usually

indicated at the beginning of a formula.

HCl hydrochloric acid Cl– chloride ion

HC2H3O2 acetic acid C2H3O2

– acetate ion

Start with

formula and

charge of the

acid.

Decrease

number of H

atoms by 1.

Decrease

charge by 1.

Start with

formula and

charge of base.

Add an H to the

formula.

Add +1 charge.

Base to Conjugate Acid

The conjugate base of an acid has one fewer hydrogen and a

charge one unit more negative than the acid. The conjugate

acid of a base has one more hydrogen and a charge one unit

more positive than the base.

a. H2SO4 + H2O → H3O+ + HSO4

–

b. HCO3

– + H2O ⇄ OH– + H2CO3

c. HI + H2O → I– + H3O+

(Answer a)

a. H2SO4 + H2O → H3O+ + HSO4

–

(Answer b)

a. H2SO4 + H2O → H3O+ + HSO4

–

b. HCO3

– + H2O ⇄ OH– + H2CO3

(Answer c)

a. H2SO4 + H2O → H3O+ + HSO4

–

b. HCO3

– + H2O ⇄ OH– + H2CO3

c. HI + H2O → I– + H3O+

2–, is an amphoteric ion.

What is its conjugate acid? What is its conjugate base?

2–, is an amphoteric ion.

What is its conjugate acid? What is its conjugate base?

(Answer part 1)

The conjugate acid has one more H+ and +1 charge compared

to HPO4

2–

H2PO4

–

2–, is an amphoteric ion.

What is its conjugate acid? What is its conjugate base?

(Answer part 2)

The conjugate acid has one more H+ and +1 charge compared

to HPO4

2–

H2PO4

–

The conjugate base has one less H+ and –1 charge compared

to HPO4

2–

PO4

3–

Strength of an acid depends on the extent to which the

acid donates a proton to water.

Strength of a base depends on the extent to which the

base accepts a proton from water.

Strong acids dissociate in water

completely to form a conjugate

base and hydronium ions.

HCl is the only strong acid

produced in human body.

Strong acids are highly corrosive

when concentrated, but are

safe to handle as dilute

solutions.

Acid Name Conjugate Base Name of Ion

−

HNO3 Nitric acid NO3

Nitrate ion

−

H2SO4 Sulfuric acid HSO4

Hydrogen sulfate ion

−

HClO4 Perchloric acid ClO4

Perchlorate ion

HCl Hydrochloric acid Cl−

Chloride ion

HBr Hydrobromic acid Br−

Bromide ion

HI Hydroiodic acid I−

Iodide ion

In solution, a strong acid contains mostly H3O+ and

the conjugate base, and very little of the acid. Strong

acids are strong electrolytes because they produce

many ions in solution.

(Groups 1A and 2A Metal Ions with Hydroxide Ions)

Completely dissociate in aqueous solution, producing hydroxide

ions and cations (conjugate acids).

Strong electrolytes if soluble

Strong bases are highly corrosive when concentrated; can be

handled safely as dilute solutions.

Formula Unit Name

Ba(OH)2 Barium hydroxide

Ca(OH)2 Calcium hydroxide

LiOH Lithium hydroxide

KOH Potassium hydroxide

NaOH Sodium hydroxide

Sr(OH)2 Strontium hydroxide

Strong bases include ionic compounds composed of hydroxide

ions and group 1A or 2A metal ions. Strong bases of soluble

ionic compounds are strong electrolytes.

a strong acid?

a. H2SO4

b. H2O

c. H2CO3

d. HF

a strong acid? (Answer)

a. H2SO4

b. H2O

c. H2CO3

d. HF

a. Ca(OH)2

b. NH3

c. NaOH

d. methylamine, CH3NH2

e. Ba(OH)2

a. Ca(OH)2

b. NH3

c. NaOH

d. methylamine, CH3NH2

e. Ba(OH)2

Characterized by little dissociation in water; produce few ions

in aqueous solution; are weak electrolytes.

Reaction with water is reversible. At equilibrium,

reactants > products.

Organic acids: acetic acid (HC2H3O2) and citric acid (H3C6H5O7)

Inorganic acids: carbonic acid (H2CO3), phosphoric acid

(H3PO4), and hydrofluoric acid (HF)

Any acid NOT listed in the Six Common Strong Acids

Nonbonding pair of electrons on N forms bond with H+ to form

conjugate acid.

Common organic weak bases include N with nonbonding electron

pair and N—C bonds.

Common inorganic weak bases: hydrogen carbonate ion (HCO3

–)

and carbonate ion (CO3

2–)

a. acetic acid, HC2H3O2

b. nitric acid, HNO3

c. hydrochloric acid, HCl

d. hydrofluoric acid, HF

a. acetic acid, HC2H3O2 weak

b. nitric acid, HNO3 strong

c. hydrochloric acid, HCl strong

d. hydrofluoric acid, HF weak

Reaction shifts in direction that counteracts a disturbance.

• Forward reaction enhanced (a shift to the right) until a new

equilibrium is attained.

• Reverse reaction enhanced (a shift to the left) until a new

equilibrium is attained.

Very important in weak-acid and weak-base reactions.

Biochemical pathways – reactions are approaching equilibrium.

a. b. c. Write the equilibrium reaction for ammonia, NH3, in

aqueous solution.

If additional hydroxide ions, OH–, are added, will the

reaction shift to the left or to the right? Explain.

If some ammonia molecules are removed from solution,

will the reaction shift to the left or to the right? Explain.

(Answer part a)

a. Write the equilibrium reaction for ammonia, NH3, in

aqueous solution.

(Answer part b)

b. If additional hydroxide ions, OH–, are added, will the

reaction shift to the left or to the right? Explain.

Adding hydroxide will enhance the reverse reaction and cause

a shift to the left.

(Answer part c)

c. If some ammonia molecules are removed from solution,

will the reaction shift to the left or to the right? Explain.

Removing ammonia molecules will decrease the rate of the

forward reaction and, thus, enhance the reverse reaction,

causing a shift to the left.

pH measures concentration of hydronium ions (H3O+) in

aqueous solution.

[H3O+] = molar concentration (mole/L) of H3O+

[OH–] = molar concentration of OH–

pH < 7 is acidic, [H3O+] > [OH–]

pH > 7 is basic (alkaline), [H3O+] < [OH–]

pH = 7 is neutral, [H3O+] = [OH–]

Blood pH: 7.35–7.45

• Lower = acidosis; higher = alkalosis

• Can be life-threatening

Other body fluids also have specific normal pH values.

Body Fluid Normal pH range

Gastric fluid 0.5–2

Urine 5–8

Saliva 6.5–7.5

Muscle cells 6.7–6.8

Arterial blood 7.35–7.45

Interstitial fluid 7.35

Intracellular fluid ≤7.0

Physiological pH is between 7.35 and 7.45. Acidosis occurs

when blood plasma pH is lower than 7.35, and alkalosis occurs

when the pH is above 7.45.

Autoionization of water occurs when very few water

molecules react with other water molecules to produce

hydronium ions and hydroxide ions.

(continued)

In pure water at 25 °C,

[H3O+] = 1.0 × 10–7 M

[OH–] = 1.0 × 10–7 M

𝐾! = H"O# × OH$

𝐾! = 1×10$% × 1×10$% = 1×10$&'

𝐾! = 10$&'

= ion-product constant for water

In pure water, the concentration of hydroxide and hydronium

ions is 1 × 10–7 M. The product of these two concentrations is

equal to the ion product constant, Kw. At room temperature,

Kw = 1 × 10–14. Thus, [H3O+] × [OH–] = 1 × 10–14

Concentrations Using the Ion-Product Constant

(part 1)

When acid or base is added to water, autoionization of water

shifts, but Kw is constant. Kw = 10–14

When hydronium ion increases, hydroxide must decrease.

[H3O+] /[OH-]

and vice versa.

Concentrations Using the Ion-Product Constant

(part 2)

Add acid to water: H!O" increases from 1×10#$ 𝑀 to

1×10#% 𝑀.

H!O" × OH#

= 1×10#&'

1×10#% × 1×10#( = 1×10#&'

Concentrations Using the Ion-Product Constant

(part 3)

H!O" × OH#

= 1×10#&'

1×10#&'

H!O" =

OH#

1×10#&'

OH#

H!O"

The ion product constant is H3O+ × 𝑂𝐻#

= 1×10#&' at 25

℃. Using this equation, [H3O+] can be calculated from [OH–],

or [OH–] can be calculated from [H3O+].

The Relationship

between pH and

the

Concentration of

H3O+ and OH−

Solution

aqueous solution with a hydronium ion concentration

of 1 × 10–3 M?

aqueous solution with a hydronium ion concentration

of 1 × 10–3 M? (Answer)

H!O" × OH#

1×10!"#

= 1×10#&'

1×10!"#

OH!

= 𝟏×𝟏𝟎!𝟏𝟏 𝑴

H$O%

1×10–3

aqueous solution with a hydroxide ion concentration of

1 × 10–10 M?

aqueous solution with a hydroxide ion concentration of

1 × 10–10 M? (Answer)

H!O" × OH#

1×10!"#

= 1 × 10#&'

1×10!"#

H$O% =

OH!

= 𝟏×𝟏𝟎!𝟒 𝑴

1×10–10

Concentration

pH of an aqueous solution = the logarithm of the hydronium

ion concentration multiplied by –1.

pH = −log"#[H$O%]

For pure water,

pH = −log"( 1×10!) = 7.0

pH decreases as hydronium ion increases.

[H3O+] = 1 × 10–6 has a pH = 6

[H3O+] = 1 × 10–4 has a pH = 4

Differ by 2 pH units, but

100 times different in

concentration!

pH scale is logarithmic; each unit is a 10-fold change in

hydronium and hydroxide ion concentrations

pH = –log[H3O+]. As the hydronium ion concentration increases

and the hydroxide ion concentration decreases, pH decreases.

The pH scale is a logarithmic scale.

ion concentration of 1 × 10–9 M? Is this solution acidic,

neutral, or basic?

ion concentration of 1 × 10–9 M? Is this solution acidic,

neutral, or basic? (Answer)

pH = – log[H3O+]

pH = – log 1×10!* = 𝟗. 𝟎

The solution is basic.

of a sample of stomach acid with a pH = 2?

(Answer)

pH = – log 𝑥 rearranges to – pH = log 𝑥

10#)* = 𝑥

𝑥 = 10#)* = 10#+

H3O +

= 𝟏×𝟏𝟎!𝟐 𝑴

When an acid is combined with a base, a neutral product is

formed and a neutralization reaction has occurred.

Neutralization reactions are double replacement reactions.

Antacids react with stomach acid (HCl) and neutralize the acid.

Ionic compounds containing:

• hydroxide ion, OH–

• hydrogen carbonate ion, HCO3

–

• carbonate ion, CO3

2–

The reaction between an acid and a base to form a neutral

product is known as a neutralization reaction.

Containing Hydroxide Ions

If an equal numbers of moles of NaOH and HCl are combined,

the result is a neutral solution.

Sodium and chloride ions are unchanged during the reaction;

they remain in solution as spectator ions.

Hydroxide Ions

OH&(aq) + H%(aq) → H'O(l)

hydroxide ion hydrogen ion water

(base) (acid) (neutral)

H+ from any acid combines with OH– from any base

to form neutral water.

Produces Water and a Salt

The reaction of an acid with a hydroxide ion-containing

base produces water and a dissolved salt. The net reaction

is H+ + OH– ® H2O. The solution is neutral when an equal

number of moles of hydroxide ions, OH–, and protons, H+

are combined.

Acids with more than one acidic proton are known as

polyprotic acids.

• sulfuric acid, H2SO4

• phosphoric acid, H3PO4

• citric acid, H3C6H5O7

(continued)

All acidic protons can be donated in a neutralization reaction.

One hydroxide ion is needed to neutralize each acidic proton.

Combine the spectator ions in a ratio that forms a neutral salt

for the ionic compound.

A polyprotic acid is an acid with two or more acidic

protons, H+. In a neutralization reaction, a polyprotic acid

donates all of its acidic protons to hydroxide ions to form

water molecules and a salt.

reaction between aqueous hydrobromic acid, HBr, and

aqueous potassium hydroxide, KOH.

a. b. c. d. e. Which solution is acidic? What ions are dissolved in the acidic

solution?

Which solution is basic? What ions are dissolved in the basic solution?

Write the complete balanced equation for the neutralization reaction.

Remember to first determine the correct formula unit for the ionic

compound formed.

What is the net reaction?

What is a spectator ion? What are the spectator ions in this

neutralization reaction?

reaction between aqueous hydrobromic acid, HBr, and

aqueous potassium hydroxide, KOH. (Answer a)

a. Which solution is acidic? HBr solution

What ions are dissolved in the acidic solution? H+ and Br–

reaction between aqueous hydrobromic acid, HBr, and

aqueous potassium hydroxide, KOH. (Answer b)

b. Which solution is basic?

KOH solution

What ions are dissolved in the basic solution?

K+ and OH–

reaction between aqueous hydrobromic acid, HBr, and

aqueous potassium hydroxide, KOH. (Answer c)

c. Write the complete balanced equation for the

neutralization reaction.

HBr (aq) + KOH (aq) → H2O (l) + KBr (aq)

reaction between aqueous hydrobromic acid, HBr, and

aqueous potassium hydroxide, KOH. (Answer d)

d. What is the net reaction?

OH– (aq) + H+ (aq) ® H2O (l)

reaction between aqueous hydrobromic acid, HBr, and

aqueous potassium hydroxide, KOH. (Answer e)

e. What is a spectator ion?

Spectator ions are ions that are unchanged during

the reaction.

What are the spectator ions in this neutralization

reaction? K+ and Br–

neutralization reaction of HF and Al(OH)3.

HF (aq) + Al(OH)3 (s) →

reaction of HF and Al(OH)3. (Answer)

3 HF (aq) + Al(OH)3 (s) → 3 H2O (l) + AlF3

Ion-Containing Bases

Ionic compounds containing hydrogen carbonate, HCO3

–, and

carbonate ion, CO3

–

• are weak bases,

• are frequently used in neutralization reactions, and

• can react with acids to produce water, an ionic

compound and CO2 gas.

HC2H3O2(aq) + NaHCO3 (aq) → H2O(l) + NaC2H3O2(aq) + CO2(g)

Reaction

The products of the reaction between an acid and a hydrogen

carbonate (HCO3

–) or carbonate ion (CO3

2–) containing base

are water, an ionic compound, and carbon dioxide gas.

Between hydrogen carbonate ion and acids:

H+ (aq) + HCO3- (aq) ® H2O (l) + CO2 (g)

Between carbonate ion and acids:

2H+ (aq) + CO3

2- (aq) ® H2O (l) + CO2 (g)

Producing Neutralization Reaction

The net reaction in a neutralization reaction involving

hydrogen carbonate ion (HCO3

–) is H+ + HCO3

– ® H2O + CO2,

and for carbonate ion (CO3

2–) it is 2 H+ + CO3

2– ® H2O + CO2.

reaction between hydroiodic acid, HI, and potassium hydrogen

carbonate, KHCO3. What are the spectator ions? What is the

net reaction?

reaction between hydroiodic acid, HI, and potassium hydrogen

carbonate, KHCO3. What are the spectator ions? What is the

net reaction? (Answer part 1)

Balanced equation:

HI (aq) + KHCO3 (aq) → H2O (l) + CO2 (g) + KI (aq)

reaction between hydroiodic acid, HI, and potassium hydrogen

carbonate, KHCO3. What are the spectator ions? What is the

net reaction? (Answer part 2)

Balanced equation:

HI (aq) + KHCO3 (aq) → H2O (l) + CO2 (g) + KI (aq)

Spectator ions: I– and K+

reaction between hydroiodic acid, HI, and potassium hydrogen

carbonate, KHCO3. What are the spectator ions? What is the

net reaction? (Answer part 3)

Balanced equation:

HI (aq) + KHCO3 (aq) → H2O (l) + CO2 (g) + KI (aq)

Spectator ions: I– and K+

Net reaction:

H+ (aq) + HCO3

– (aq) → H2O (l) + CO2 (g)

Buffer = solution that resists changes in pH when small

amounts of acid or base are added.

How acid-base homeostasis works to prevent acidosis and

alkalosis.

Weak acid + its conjugate base in similar concentrations.

pH range maintained by buffer depends on identity of the

weak acid.

Buffer capacity depends on the concentration of the buffer

components.

Buffer Weak Acid Conjugate

Base

Function

Acetic acid/acetate

buffer

CH3COOH CH3COO−

Common laboratory

buffer

−

Carbonic

acid/hydrogen

carbonate buffer

H2CO3 HCO3

Extracellular buffer

Dihydrogen

phosphate/hydrogen

phosphate buffer

−

H2PO4

HPO4

2−

Intracellular buffer and

common laboratory

buffer

A buffer contains approximately equal concentrations of a

weak acid and its conjugate base, therefore resisting changes

in pH when a small amount of acid or base is added.

Adding H3O+ increases the rate of the reverse reaction,

producing more HA and H2O, a shift to the left.

[H3O+] changes only a very little, resulting in a relatively

constant pH. pH = –log [H3O+]

Adding OH– decreases [H3O+], slowing the rate of reverse

reaction, producing more H3O+ and A –, a shift to the right.

Again, [H3O+] changes only a very little, resulting in a relatively

constant pH. pH = –log [H3O+]

A buffer resists changes in pH because the equilibrium shifts in

response to the addition or removal of H3O+, maintaining the

concentration of H3O+. The addition of hydroxide ion is

equivalent to the removal of hydronium ion, H3O+:

[H3O+] /[OH-]

Buffer range = pH 3.8 – pH 5.8

Acetic acid (HC2H3O) and sodium acetate (NaC2H3O2)

combined in equal numbers of moles

Buffer in the Blood

Primary blood buffer = combination of carbonic acid (H2CO3) +

hydrogen carbonate (HCO3

–)

Combination reaction; reversible

Equilibrium

Adding H3O+ causes a shift to the left in H2CO3/HCO3

– equilibrium.

[H2CO3] increases, causing a shift to the left in the CO2/H2CO3

equilibrium.

[CO2] increases, causing lungs to exhale more CO2.

Adding OH–

Adding OH– decreases H3O+ and causes a shift to the right in H2CO3/HCO3

–

equilibrium.

[H2CO3] decreases, causing a shift to the right in the CO2/H2CO3 equilibrium.

[CO2] decreases, causing lungs to exhale less CO2.

conjugate acid. For example, approximately equal amounts of

ammonia, NH3, and ammonium chloride, NH4

+, in solution

form a buffer.

a. b. c. d. e. Show the reversible equilibrium reaction by treating NH4

+ as

the acid. Label the buffer components.

How would this equilibrium shift if hydronium ions were

added?

How would this equilibrium shift if hydroxide ions were added?

Why does the pH not change significantly when hydronium or

hydroxide ions are added to the buffer?

If a large excess of hydronium ions or hydroxide ions were

added, exceeding the buffer capacity, would the pH change?

conjugate acid. For example, approximately equal amounts of

ammonia, NH3, and ammonium chloride, NH4

+, in solution

form a buffer. (Answer part a)

a. Show the reversible equilibrium reaction by treating NH4

as the acid. Label the buffer components.

NH4

+ (aq) + H2O (l) ⇄ NH3 (aq) + H3O+ (aq)

acid conjugate base

conjugate acid. For example, approximately equal amounts of

ammonia, NH3, and ammonium chloride, NH4

+, in solution

form a buffer. (Answer part b)

NH4

+ (aq) + H2O (l) ⇄ NH3 (aq) + H3O+ (aq)

acid conjugate base

b. How would this equilibrium shift if hydronium ions were

added?

Shift to the left

conjugate acid. For example, approximately equal amounts of

ammonia, NH3, and ammonium chloride, NH4

+, in solution

form a buffer. (Answer part c)

NH4

+ (aq) + H2O (l) ⇄ NH3 (aq) + H3O+ (aq)

acid conjugate base

c. How would this equilibrium shift if hydroxide ions were

added?

Shift to the right

conjugate acid. For example, approximately equal amounts of

ammonia, NH3, and ammonium chloride, NH4

+, in solution

form a buffer. (Answer part d)

NH4

+ (aq) + H2O (l) ⇄ NH3 (aq) + H3O+ (aq)

acid conjugate base

d. Why does the pH not change significantly when hydronium

or hydroxide ions are added to the buffer?

The addition of hydronium or hydroxide ions shifts the

equilibrium to a new position, causing only a small change in

the hydronium ion concentration and a small change in pH.

conjugate acid. For example, approximately equal amounts of

ammonia, NH3, and ammonium chloride, NH4

+, in solution

form a buffer. (Answer part e)

NH4

+ (aq) + H2O (l) ⇄ NH3 (aq) + H3O+ (aq)

acid conjugate base

e. If a large excess of hydronium ions or hydroxide ions were

added, exceeding the buffer capacity, would the pH change?

Yes, pH would decrease if excess hydronium ions or increase if

excess hydroxide ions were added.

Homeostasis, Acidosis, and Alkalosis

Acid-Base homeostasis is maintained in the blood by:

• buffers

• regulation of breathing (lung ventilation)

• absorption and release of HCO3

– by kidneys

(continued)

During exercise, [H3O+] increases, causing a shift to the left,

producing more H2CO3, shifting left again to produce more

CO2, which is exhaled by increasing ventilation.

Kidneys can support the shift to the left by releasing

additional HCO3

– into the blood.

Respiratory acidosis –

breathing weak and shallow so

CO2 is not exhaled

Causes a shift to the right,

increasing [H3O+] and

decreasing pH

Caused by head injuries,

emphysema, asthma, narcotic

use

Metabolic acidosis—kidneys do

not release sufficient HCO3

– to

support normal shift to the left

Leads to increasing [H3O+] and

decreasing pH

Caused by kidney failure,

uncontrolled diabetes,

starvation

Respiratory alkalosis occurs

when rapid breathing removes

too much CO2.

Causes a shift to the left,

decreasing [H3O+] and

increasing pH

Hyperventilation caused by

anxiety, altitude sickness,

intense exercise

Metabolic acidosis—kidneys do

not remove enough HCO3

– from

blood

Causes a shift to the left,

decreasing [H3O+] and

increasing pH

Caused by excessive vomiting,

excess usage of antacids, and

adrenal gland diseases

Respiratory or metabolic

acidosis: IV infusion of HCO3

–

causing a shift to left.

Respiratory alkalosis: breathing

into paper bag to increase

blood CO2 and shift

equilibrium to right

Metabolic alkalosis: IV infusion

of dilute HCl causing a shift to

right

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