8. Acids, Bases and Buffers

Question 1

Arrhenius Acid Base Definition

Answer 1

• an acid is a substance that produces H₃O⁺ ions in aqueous solution, and a base is a substance that produces OH^- ions in aqueous solution
• when an acid dissolves in water it reacts with the water to produce a hydronium ion H₃O⁺

H⁺_(aq) + H₂O_(l) → H₃O⁺_(aq)

• many bases are metal hydroxides - ionic solids - which dissolve in water to produce solvated ions

AOH_(s) →^H₂O A⁺_(aq) + OH^-_(aq)

• bases that are not hydroxides will react with water to produce hydroxide ions
• strong acids + bases ionize completely in aqueous solution
• weak acids + bases ionize only partially on aqueous solution

Question 2

Brønsted - Lowry Acid Base Definition

Answer 2

• acid is a proton donor
• base is a proton acceptor
• acid-base reaction is a proton transfer reaction
• monoprotic acid: acid that can give up only one proton
• diprotic acid: acid that can give up two protons
• triprotic acid: acid that can give up three protons
• amphiprotic: substance that can act as either acid or base
• acid can be positively charged, neutral or negatively charged
• base can be neutral or negatively charged
• a substance cannot be a Brønsted-Lowry acid unless it contains a hydrogen atom, but not all hydrogen atoms can be given up because a hydrogen must be bonded to a strongly electronegative atom, such as oxygen or a halogen, to be acidic

Question 3

Conjugate Acid - Base

Answer 3

• conjugate base: substance formed when an acid donates a proton to another molecule or ion
• conjugate acid: substance formed when a base accepts a proton
• conjugate acid-base pair: pair of molecules or ions that are related to one another by the gain or loss of a proton
• inverse relationship between the strength of an acid or base and its conjugate pair: the stronger the acid, the weaker its conjugate base, and conversely, the stronger the base, the weaker its conjugate acid

Question 4

Acid Base Equilibrium

Answer 4

In an acid base reaction, the equilibrium position always favours reaction of the stronger acid + base to form the weaker acid + base. Thus, at equilibrium, the major species present are the weaker acid + base.

Question 5

Acid Ionisation Constant

Answer 5

K_a

• also known as acid dissociation constant
• equilibrium constant for the ionization of an acid in aqueous solution to H₃O⁺ and its conjugate base

HA + H₂O ⇔ A^- + H₃O⁺

K_a = [A^-][H₃O⁺] / [HA]

• the weaker the acid, the smaller its K_a and the larger its pK_a

Question 6

Properties of Acids + Bases

Answer 6

Neutralization: acids + bases → neutral solutions

strong acids + active metals → H₂ + salt (redox reaction)

acids + metal hydroxides → salt + H₂O (salt is present as aqueous anions + cations)

strong acids + metal oxides → H₂O + soluble salt

strong acid + carbonate → carbon dioxide + H₂O

2H₃O⁺_(aq) + CO₃^2-_(aq) → H₂CO_3(aq) + 2H₂O_(l)

                  H<sub>2</sub>CO<sub>3(aq)</sub> →  CO<sub>2(g)</sub>  +  H<sub>2</sub>O<sub>(l)</sub>

_ _

2H₃O⁺_(aq) + CO₃^2-_(aq) → CO_2(g) + 3H₂O_(l)

strong acid + bicarbonate → carbon dioxide + H₂O

H₃O⁺_(aq) + HCO₃^-_(aq) → H₂CO_3(aq) + H₂O_(l)

                 H<sub>2</sub>CO<sub>3(aq)</sub> →  CO<sub>2(g)</sub>  +  H<sub>2</sub>O<sub>(l)</sub>

_ _

H₃O⁺_(aq) + HCO₃^-_(aq) → CO_2(g) + 2H₂O_(l)

amines + ammonia will react with any acid stronger than NH₄⁺ to form a salt - a particularly important reaction within the body

Question 7

Ion Product of Water

Answer 7

K_w

• equilibrium constant for the ionization of water

K_w = [H₃O⁺][OH^-] = 1.0 x 10^-14

• pure water at room temperature, K_w = 1.0 x 10^-14

[H₃O⁺] > 1.0 x 10^-7M and [OH^-] ^-7 is acidic

[OH^-] > 1.0 x 10^-7M and [H₃O⁺] ^-7 is basic

• higher [H₃O⁺] = more acidic

Question 8

pH + pOH

Answer 8

• a convenient way to designate the concentration of H₃O⁺ and OH^-

pH = -log[H₃O⁺]

pOH = -log[OH^-]

pH + pOH = 14

pH

pH > 7 solution basic

pH = 7 solution neutral

Question 9

Buffer

Answer 9

• mixture containing both weak acid + weak base
• capable of absorbing small additions of either strong acid or strong base with little change to pH
• weak base neutralises addition of small quantities of strong acid
• weak acid neutralises addition of small quantities of strong base
• buffer solutions can be prepared to maintain almost any pH

Question 10

Buffer Capacity

Answer 10

• amount of hydronium or hydroxide ions that a buffer can absorb without a significant change of pH
• the nature of the buffer capacity depends on both the pH relative to pK_a and the concentration of the buffer
• the closer the pH of the buffer is to the pK_a of the weak acid, the more symmetric the buffer capacity, meaning the buffer can resist a pH change with added acid or added base
• the greater the concentration of the weak acid and its conjugate base, the greater the buffer capacity
• an effective buffer has a pH equal to the pK_a of the weak acid +/- 1
• buffer pH = conjugate acid pK_a : equal capacity for addition of acid or base
• buffer pH _a : capacity greater toward addition of base
• buffer pH > conjugate acid pK_a : capacity greater toward addition of acid

Question 11

Henderson - Hasselbach Equation

Answer 11

• convenient way to calculate pH of a buffer when the concentrations of the weak acid and its conjugate base are not equal

For HA_{weak acid} + H₂O ⇔ A^-_{conj base} + H₃O⁺

pH = pK_a + log ([A^-] / [HA])

Question 12

Blood Buffers

Answer 12

• average pH of human blood is 7.4
• change larger than +/-0.10 pH unit may cause illness
• pH below 6.8 or above 7.8 may be deadly
• body uses 3 buffer systems: carbonate, phosphate and proteins

Question 13

Carbonate Blood Buffer

Answer 13

• most important buffer system
• weak acid is carbonic acid, H₂CO₃, with normal blood concentration of 0.0025M
• conjugate base is bicarbonate ion, HCO₃^-, with normal blood concentration of 0.025M
• [HCO₃^-] : [H₂CO₃] ratio is 10:1
• better buffer for acids (lowers the ratio thus improves buffer efficiency), than bases (raises the ratio thus decreasing buffer capacity)
• under normal conditions, larger amounts of acidic substances enter the blood
• 10:1 ratio is easily maintained as the body can quickly increase or decrease the amount of CO₂ entering the blood
• H₃O⁺ neutralized by HCO₃^- ion and OH^- neutralized by H₂CO₃

Question 14

Phosphate Blood Buffer

Answer 14

• second most important buffer system
• hydrogen phosphate ions, HPO₄^2-, and dihydrogen phosphate ions, H₂PO^₄-, in a ratio of 1.6:1

Question 15

Acidosis

Answer 15

• a condition whereby the pH of the blood is lower than 7.35, leading to depression of the nervous system resultng in dizziness, disorientation, fainting or coma
• respiratory acidosis results from a difficulty in breathing (hypoventilation) where the amount of oxygen that reaches the tissues is diminished and the amount of CO₂ that leaves the body via the lungs is also diminished
• metabolic acidosis can occur as a result of starvation or heavy exercise where the body burns its own fat and releases acidic compounds into the blood, or when muscles produce excessive amounts of lactic acid
• occassionally, both types of acidosis can occur. When cells are deprived of oxygen, respiratory acidosis results. These cells are unable to produce the energy required through aerobic (oxygen-requiring) pathways, so they must use the anaerobic (without oxygen) pathway - glycolysis - which produces lactic acid as an end product, leading to metabolic acidosis. Eventually the lack of oxygen must be recouped and the lactic acid removed. In extreme cases, the oxygen debt is too great and can cause death

Question 16

Alkalosis

Answer 16

• a condition whereby the pH of the blood is higher than 7.45, leading to overstimulation of the nervous system resulting in muscle cramps, dizziness and convulsions.
• it arises from hyperventilation where an excessive loss of CO₂ raises both the ratio of [HCO₃^-] : [H₂CO₃] and the pH
• short distance athletes use alkilosis to their advantage. By hyperventilating right before the start they force the removal of CO₂ from their lungs, causing more H₂CO₃ to dissociate to replace the lost CO₂ and therefore raising the blood pH. With a slightly higher blood pH, more lactic acid can be absorbed before the pH drops to where performance is impaired