Acids, Bases, pH, and Buffers Flashcards

Acid-Base Homeostasis, pH, and Buffers Notes

Importance of Maintaining Blood pH

• Proper organ function requires blood pH to be between 7.35 and 7.45.
• Acidosis: pH < 7.35
• Alkalosis: pH > 7.45

Acids and Bases: Aqueous Solutions

• Acids produce hydronium ions (H_3O^+).
• Bases produce hydroxide ions (OH^–).

Ionic Compounds Containing Hydroxide Ions (Arrhenius Bases)

• Group 1A hydroxide salts (e.g., NaOH, KOH):
  • Soluble in water.
  • Strong electrolytes: Dissolve entirely to produce ions in water.
• Group 2A hydroxide salts (e.g., Ca(OH)2, Mg(OH)2):
  • Slightly soluble in water.
  • Weak electrolytes: Dissolve slightly to produce few ions in water.

Brønsted–Lowry Theory of Acids and Bases

• Acid: A proton (H^+) donor.
• Base: A proton (H^+) acceptor.
• Amphoteric: A substance that can act as both an acid and a base (e.g., water).
• Example: Ammonia (NH_3) acts as a base by accepting a proton from a water molecule.

Conjugate Acid-Base Pairs

• A conjugate acid-base pair consists of an acid and its conjugate base or a base and its conjugate acid.

Predicting Formula of Conjugate Base

1. Start with the formula and charge of the acid.
2. Decrease the number of H atoms by 1.
3. Decrease the charge by 1.

Predicting Formula of Conjugate Acid

1. Start with the formula and charge of the base.
2. Add an H to the formula.
3. Add +1 to the charge.

Strengths of Acids and Bases

• The strength of an acid depends on the extent to which the acid donates a proton to water.
• The strength of a base depends on the extent to which the base accepts a proton from water.

Strong Acids

• Dissociate completely in water to form a conjugate base and hydronium ions.
• HCl is the only strong acid produced in the human body.

Common Strong Acids

• Nitric acid (HNO_3)
• Sulfuric acid (H2SO4)
• Perchloric acid (HClO_4)
• Hydrochloric acid (HCl)
• Hydrobromic acid (HBr)
• Hydroiodic acid (HI)
• chloric (HClO_3)

Strong Bases

• Completely dissociate in aqueous solution, producing hydroxide ions and cations (conjugate acids).
• Strong electrolytes if soluble.
• Examples: (Groups 1A and 2A Metal Ions with Hydroxide Ions)
  • Ba(OH)_2 (Barium hydroxide)
  • Ca(OH)_2 (Calcium hydroxide)
  • LiOH (Lithium hydroxide)
  • KOH (Potassium hydroxide)
  • NaOH (Sodium hydroxide)
  • Sr(OH)_2 (Strontium hydroxide)

Weak Acids and Weak Bases

• Characterized by little dissociation in water; produce few ions in aqueous solution; are weak electrolytes.
• Reaction with water is reversible; at equilibrium, reactants > products.

Examples of Weak Acids

• Organic acids: acetic acid (HC2H3O2) and citric acid (H3C6H5O_7)
• Inorganic acids: carbonic acid (H2CO3), phosphoric acid (H3PO4), and hydrofluoric acid (HF)
• Any acid NOT listed as one of the six common strong acids.

Weak Bases

• 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–}).

Le Châtelier’s Principle

• Very important in weak-acid and weak-base reactions.
• Biochemical pathways – reactions are approaching equilibrium.
• Reaction shifts in the 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.

pH

• Measures the concentration of hydronium ions (H_3O^+) in an aqueous solution.
• [H3O^+] = molar concentration (mole/L) of H3O^+.
• [OH^–] = molar concentration of OH^–.
• [H_3O^+] > [OH^–]: acidic solution
• [H_3O^+] = [OH^–]: neutral solution
• [H_3O^+] < [OH^–]: basic solution
• Blood pH: 7.35–7.45 (Lower = acidosis; higher = alkalosis)

pH of Some Body Fluids

• 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

K_w and the Autoionization of Water

• Autoionization of water occurs when very few water molecules react with other water molecules to produce hydronium ions and hydroxide ions.
• In pure water at 25 °C:
  • [H_3O^+] = 1.0 \,x\, 10^{-7} M
  • [OH^–] = 1.0 \,x\,10^{-7} M
  • Kw = [H3O^+][OH^–] = (1 \,x\, 10^{-7})(1\,x\, 10^{-7}) = 1 \,x\, 10^{-14}
• K_w = ion-product constant for water.
• K_w = 10^{-14}

Calculating Hydroxide and Hydronium Ion Concentrations Using the Ion-Product Constant

• 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, and vice versa.
• Adding acid to water: [H_3O^+] increases from 1 \,x\, 10^{−7} M to 1 \,x\, 10^{−6} M.
• [H_3O^+] \,x\, [OH^–] = 1 \,x\, 10^{−14}
• 1 \,x\, 10^{−6} \,x\, 1 \,x\, 10^{−8} = 1 \,x\, 10^{−14}

Relationship between pH and Concentration of H_3O^+ and OH^−

• [H_3O^+] = \frac{1 \,x\, 10^{-14}}{[OH^–]}
• [OH^–] = \frac{1 \,x\, 10^{-14}}{[H_3O^+]}

Calculating pH from the Hydronium Ion Concentration

• pH of an aqueous solution = the negative logarithm of the hydronium ion concentration.
• pH = -log[H_3O^+]
• For pure water:
  • pH = -log[1 \,x\, 10^{-7}] = 7.0

pH Scale

• pH decreases as hydronium ion increases.
• The pH scale is logarithmic; each unit is a 10-fold change in hydronium and hydroxide ion concentrations.

Acid-Base Neutralization Reactions

• 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.
• If an equal number 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.

Net Reaction for Reaction of Any Acid and Hydroxide Ions

• H^+ from any acid combines with OH^− from any base to form neutral water.

Neutralization Reactions with Polyprotic Acids

• Acids with more than one acidic proton are known as polyprotic acids (e.g., sulfuric acid, H2SO4; phosphoric acid, H3PO4; citric acid, H3C6H5O7).
• 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.

Hydrogen Carbonate and Carbonate Ion-Containing Bases

• Ionic compounds containing hydrogen carbonate, HCO3^–, and carbonate ion, CO3^{2–}:
  • Are weak bases.
  • Are frequently used in neutralization reactions.
  • Can react with acids to produce water, an ionic compound, and CO_2 gas.
• Net Reactions:
  • Between hydrogen carbonate ion and acids: H^+ (aq) + HCO3^− (aq) → H2O (l) + CO_2 (g)
  • Between carbonate ion and acids: 2H^+ (aq) + CO3^{2−} (aq) → H2O (l) + CO_2 (g)

Buffers

• A buffer is a solution that resists changes in pH when small amounts of acid or base are added.
• A buffer contains approximately equal concentrations of a weak acid and its conjugate base, thus resisting changes in pH when a small amount of acid or base is added.
• Components: Weak acid + its conjugate base in similar concentrations.
• pH range maintained by buffer depends on the identity of the weak acid.
• Buffer capacity depends on the concentration of the buffer components.

Common Buffers

• Acetic acid/acetate buffer (CH_3COOH): Common laboratory buffer.
• Carbonic acid/hydrogen carbonate buffer (H2CO3): Extracellular buffer.
• Dihydrogen phosphate/hydrogen phosphate buffer: Intracellular buffer and common laboratory buffer.

The Buffer Equilibrium

• 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 the 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^+]

The Carbonic Acid/Hydrogen Carbonate Buffer in the Blood

• Primary blood buffer = combination of carbonic acid (H2CO3) + hydrogen carbonate (HCO_3^–).

Carbonic Acid Forms from CO_2 and Water

• Combination reaction; reversible:
  • CO2 (g) + H2O (l) ⇌ H2CO3 (aq)

Central Role of H2CO3

• Combination/decomposition reaction:
  • CO2 (g) + H2O (l) ⇌ H2CO3 (aq) ⇌ H3O^+ (aq) + HCO3^− (aq)

Adding H_3O^+

• 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 CO_2.

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 CO_2.

Chemistry in Medicine: Acid-Base Homeostasis, Acidosis, and Alkalosis

• Acid-Base homeostasis is maintained in the blood by:
  • Buffers
  • Regulation of breathing (lung ventilation)
  • Absorption and release of HCO_3^– by kidneys

Breathing Rate Affects CO_2 Levels in Blood

• 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 HCO_3^– into the blood.

Acidosis – Blood pH < 7.35

• Respiratory acidosis: breathing weak and shallow so CO_2 is not exhaled.
  • Causes a shift to the right, increasing [H_3O^+] and decreasing pH.
  • Caused by head injuries, emphysema, asthma, narcotic use.
• Metabolic Acidosis: kidneys do not release sufficient HCO_3^– to support the normal shift to the left.
  • Leads to increasing [H_3O^+] and decreasing pH.
  • Caused by kidney failure, uncontrolled diabetes, starvation.
  • Treatment: IV infusion of HCO_3^– causing a shift to the left.

Alkalosis – Blood pH > 7.45

• Respiratory alkalosis occurs when rapid breathing removes too much CO_2.
  • Causes a shift to the left, decreasing [H_3O^+] and increasing pH.
  • Hyperventilation caused by anxiety, altitude sickness, intense exercise.
• Metabolic acidosis—kidneys do not remove enough HCO_3^– from the blood
  • Causes a shift to the left, decreasing [H_3O^+] and increasing pH.
  • Caused by excessive vomiting, excess usage of antacids, and adrenal gland diseases.
  • Treatments:
    • Respiratory alkalosis: Breathing into a paper bag to increase blood CO_2 and shift equilibrium to the right.
    • Metabolic alkalosis: IV infusion of dilute HCl causing a shift to the right.