Acids and Bases

Objectives

• This lecture outlines acid and base concepts/theories.
• Calculation of pH and acid/base dissociation constants will be covered.
• Determining acid and base strength is essential.
• Measuring pH is a key aspect.
• Buffers and their operation are another focus.
• pH of buffers and buffer capacity calculation is important.
• Physiological buffers and their mechanisms will be discussed.

Introduction

• Currently, there are four major concepts about acids and bases.
• These include the Traditional, Arrhenius, Bronsted-Lowry, and Lewis Concepts.

Traditional Concept

• Acids convert blue litmus paper to red and have a sour taste.
• Acids have a pH less than 7.
• Acids react with bases to form salt and water.
• Bases convert red litmus paper to blue with a bitter taste.
• Bases have a pH greater than 7.
• Bases react with acids to form salt and water.

Arrhenius Concept

• An acid releases hydrogen ions (H+) when dissolved in water, such as HCl(aq) → H+(aq) + Cl-(aq).
• A base releases hydroxide ions (OH-) when dissolved in water, such as NaOH(aq) → Na+(aq) + OH-(aq).

Arrhenius Concept Limitations

• The concept is limited to water-based solutions because free OH- and H+ do not exist in water.
• H+ forms H3O+ and OH- forms H3O2− in water.
• Some bases do not contain OH-, such as ammonia (NH3).

Bronsted-Lowry Concept

• An acid donates a proton (H+).
• A base accepts a proton.
• Acids and bases are identified as conjugate pairs.
• Consider a reaction: Acid₁ + Base₂ → Acid₂ + Base₁.
• Example: HCl + H₂O ⇔ H3O+ + Cl-.
• Conjugate acid-base pairs are acid-base pairs where the reaction members are formed when protons are gained or lost.
• A strong acid easily gives protons to form a weak conjugate base.
• A strong base accepts protons more easily to form a weak conjugate acid.
• A strong acid typically has a weak conjugate base pair, and a weak acid has a strong conjugate base pair, which also applies to bases.
• The formation of conjugate pairs is important when considering buffers and their mechanisms.

Bronsted-Lowry Conjugate Pairs

• This concept is applicable when considering the strength of acids and during pH and pKa measurement.

Bronsted-Lowry Acid and Base Classifications

• Acids and bases can be classified as monoprotic, capable of donating or accepting one proton.
• Acids and bases can be classified as polyprotic, capable of donating or accepting two or more protons.
• Amphiprotic substances can behave as both Bronsted-Lowry acids and bases.

Bronsted-Lowry Advantages

• It is not limited to aqueous solutions.
• It is simple to apply to acids and bases.
• The relative strength of acids and bases in aqueous solvents can develop buffers.
• It calculates pH, pOH, and pKa, which helps determine the characteristics and behaviors of acids and bases.

Bronsted-Lowry Limitations

• Reactions in non-protonic solvents such as COCl2, SO2, and N2O4 cannot be explained by the protonic definition.
• Substances like BF3 and AlCl3 do not have hydrogen; therefore, they cannot give a proton, but are known to behave as acids.

Lewis Concept

• Acids are electron pair acceptors.
• Bases are electron pair donors.
• Lewis acids and bases share an electron pair from the base, creating a covalent or coordinate bond.
• The result is called a complex.
• Cations and atoms with empty valence orbitals act like Lewis acids.
• Anions or molecules possessing a lone electron pair act as Lewis bases.
• During neutralization, a coordinate bond forms between an acid and a base.

Lewis Strengths

• All Bronsted-Lowry acids and bases are covered.
• A proton transfer or gain is accompanied by an electron-pair loss or donation.
• Reactions not involving proton transfer are also covered.

Lewis Limitations

• Substances not normally considered acids behave as Lewis acids, e.g., BF3.
• The concept does not provide a scale of acid and base strength.
• The strength of an acid or base varies based on the solvent and the type of reaction.

Relative Strength of Acids and Bases

• The strength refers to the efficiency with which acids or bases demonstrate their properties, whether they are monoprotic or polyprotic.
• Acid strength is the ease with which it donates protons.
• Base strength is the ease with which it accepts protons.
• This property is influenced by the surrounding environment.

Acid and Base Classes

• There are two general classes based on strength: strong and weak.
• Strong acids and bases are completely dissociated in aqueous media.
• Weak acids and bases are partially dissociated in aqueous media.
• The strength of an acid or base differs from its concentration.
• Concentration is a measure of moles dissolved per dm³, and some molecules may not dissociate into H+ or OH- ions.
• Strong dilute acids or bases and weak concentrated acids or bases can occur.
• Acid or base strength is typically determined by the dissociation constant "K".
• Ka represents the strength of an acid.
• Kb or Ka of its conjugate acid represents the strength of a base.

Calculating Strong Acid Strength

• An equation represents the dissociation of acid HA into H+ ion and A- ion.
• Ka is the acid dissociation or ionization constant.
• Acid strength is the concentration of H+ in an aqueous solution at a given temperature.
• H+ concentration is directly proportional to Ka; therefore, Ka measures acid strength.
• In aqueous solutions of strong acids, the acids are completely dissociated and the Ka is large.
• On the other hand, weak acids in aqueous solution are dissociated to a small extent, and the Ka is also small.
• The acid dissociation constant or Ka is large for strong acids and small for weak acids.
• Because a strong acid completely dissociates in an aqueous solution, the concentration of H+ ions equals that of the acid. For example, for 0.1 M HCl, [H+] = 0.1 M.
• For 0.1 M weak acids like ethanoic acid, [H+] << 0.1 M

Calculating Weak Acid Strengths

• Taking the initial acid concentration to be 1 at equilibrium, the following equations apply.
• [H+] = α
• [A-] = α
• [HA] = 1-α
• For Ka: K = α2 / (1-α)
• For weak acids, 1-α ≈ 1; therefore, K = α2
• For two different acids, 1 and 2, the dissociation constants are K₁ and K₂; the degrees of dissociation are α₁ and α₂.
• For acid 1: K₁ = α₁² (1)
• For acid 2: K₂ = α₂² (2)
• [H+] measures acid strength and depends on the degree of dissociation, α.
• The ratio of √K₁/K₂ gives the relative two acids strengths.

Calculating pH

• Because H+ and OH- concentrations in aqueous solutions are small, pH was introduced for convenient measurement.
• pH expresses the concentration of hydrogen or hydronium ions in a solution.
• pH = -log[H+]
• pOH expresses hydroxide ion concentration in a solution.
• pOH = -log[OH-]
• Since Kw = [H+][OH-], pKw = -log Kw,
• -log(1.0 x 10^-14) = 14
• pH + pOH = 14

Ion-Product Expression

• Kw = [H+][OH-]
• Kw = 1.0 x 10^-14 at 25°C
• For a Neutral Solution: [H+] = [OH-] = (1.0 x 10^-7)
• For an Acidic Solution: [H+] > [OH-]
• For a Basic Solution: [H+] < [OH-]

Measuring pH

• Adding pH indicator to a solution causes the indicator color to vary depending on the solution's pH.
• Universal indicators determine qualitative measurement.
• Indicators featuring strong color variability determine quantitative results.
• Combining indicators having multiple equilibriums with spectrophotometric methods can also produce precise measurements.
• Using pH or indicator paper causes a color change corresponding to a color key.
• Use a pH meter consisting of a voltameter and two electrodes.
• A standard electrode has a known potential.
• A special electrode (probe) is enclosed in a glass membrane that allows H+ ion migration. Also, the glass case contains a KCl reference solution.

Importance of Acids and Bases

• Acid-base neutralization reactions prepare salts and converts salts into forms suitable for effervescent mixtures.
• They are useful in analytical procedures, e.g., titration.
• They are used as therapeutic agents to control and adjust the pH of the GIT, body fluid, and urine.
• They create buffers.

Buffers: Maintaining pH

• Buffers manage pH in laboratory and industrial processes.
• Accomplished using buffer systems or buffer solutions.
• Buffer solutions maintain a stable pH upon adding small acid or base quantities.
• Buffers resist pH changes when adding acids or bases.
• Buffer solutions contain a weak acid and its conjugate base (salt) or a weak base and its salt.

Buffer Solution Types

• Acidic Buffer Solution: A solution containing a weak acid and its salt, for example, Acetic acid (CH3COOH) and Sodium acetate (CH3COONa).
• Basic Buffer Solution: A solution containing weak base and its salt, for example: Ammonium hydroxide (NH4OH) and Ammonium chloride (NH4Cl).

Buffer Mechanisms

• In non-buffer solutions, adding a small amount of acid changes the hydrogen ion concentration.
• Buffers prevent significant pH drops or rises.
• Buffer system components complement each other.
• When adding a small amount of acid to an acidic buffer, the added acid reacts with the conjugate base to form a weak acid, which cannot dissociate further.

Buffer Action with Added Base

• With small amounts of added base, the weak acid reacts, producing water and a conjugate base.
• The new result produces a negligible change to the pH.
• The outcome only occurs when adding minor quantities of acid or base.

Buffer Operation: Acetic Acid and Sodium Acetate (CH3COOH/CH3COONa)

• Reaction equations:
• CH3COOH ↔ H+ + CH3COO−
• CH3COONa → CH3COO− + Na+
• Because the salt is completely ionized, the solution provides CH3COO- ions in excess.
• The common ion effect suppresses the ionization of acetic acid.
• This reduces H+ ion concentration, increasing pH.
• e.g., 0.1 M acetic acid has a pH of 2.87, while 0.1 M acetic acid + 0.1 M sodium acetate solution (buffer) has a pH of 4.74.

Addition of Acid to Buffer

• An acid (HCl) to this buffer, increases H+, which the excess acetate ions counteract to form CH3COOH.
• The CH3COOH increase shifts the equilibrium right, slightly increasing the H+ concentration.
• A decrease in pH is observed from 4.74 to 4.66.
• This pH change qualifies as marginal, so, the pH of the buffer almost remains unchanged.

Adding Base to a Buffer

• Hydroxide ions (OH−) combine with (H+) from CH3COOH to form H2O.
• This shifts equilibrium to produce more H+, neutralizing excess OH− and restoring buffer pH.
• A new equilibrium forms, lowering [CH3COOH].
• [H+] is a bit lower, and the pH is a bit higher (4.74 to 4.83).
• The pH change is marginal, and the buffer pH remains stable.

Buffer Solution pH Calculation

• The hydrogen ion concentration, obtainable from the dissociation of a weak acid HA, is determined by the equation.
• HA ↔ H+ + A-
• Ka = [H+][A-]/[HA]
• [H+] = Ka X [HA]/[A-]
• The equilibrium concentration of [A−] is presumed equal to the salt’s initial concentration.

Henderson-Hasselbalch Equation

• When taking the negative log of both sides of the equation obtains an expression equal to pH and p.
• The equation provides pH = p.
• The equation helps expresses pH in terms of pKa and concentrations of acid and conjugate base.
• The equation to determine a basic buffer:
• pOH =p
• An equation applies to acidic buffers.

Henderson-Hasselbalch Significance

• This equation helps to determine buffer solution pH when concentrations are known.
• It also assists when preparing solutions with known pH.
• Maximum buffer action is attained when conjugate base concentration equals acid.
• Application applies to basic buffers.

Buffer Capacity (β)

• Knowledge of a buffer’s effectiveness is often a necessity on a quantitative level; Van Slyke introduced the term in 1922.
• Buffer capacity (β) measures acid or base volumes necessary to add to cause a unit change to the pH.
  • β= n/ΔpH, where,
    • n = Concentration of added acid or base
    • ΔpH = change in pH
• Buffer capacity is always positive.
• Buffer capacity measures how well a solution resists pH changes when adding acids or bases.
• The greater the buffer capacity, the less the change to pH.
• Concentration influences buffer capacity.
• Thus, the higher the buffering species concentration, the higher the buffer capacity.
• Buffer capacity increases, as 𝚫𝒑𝑯 decreases.

How to Select a Suitable Buffer

• Possess adequate buffer capacity in the required pH range.
• Be available in a high degree of purity.
• Be very water-soluble and impermeable to biological membranes.
• Be enzymatically and hydrolytically stable.
• The buffer should be nontoxic or contain biological inhibitors.
• Not participate in redox reactions.
• Not alter the other ingredients' solubility.
• The buffer should be safe.

Physiological Buffers: Maintaining Bodily pH Balance

• The body fluids must maintain a normal pH range because biochemical reactions are sensitive to acidity or alkalinity changes.
• For instance, low stomach pH permits pepsin's enzymatic activity, critical for dietary protein digestion.
• Saliva has a pH range of 5.4 to 7.5 needed for optimal functioning of salivary amylase, or ptyalin.
• Metabolic activity produces compounds able to alter blood pH in the body.
• Strong acids and bases constantly form but fluids inside and outside remain constant thanks to buffer systems.
• Most have a weak acid and conjugate base or the acid’s salt.

Physiological Buffers

• The body prevents rapid pH changes via buffer systems.
• The body fluid buffer systems are bicarbonate, phosphate, and protein (hemoglobin).

Bicarbonate Buffer System

• It is found in plasma and kidneys.
• It is a regulator of blood pH.
• If there are excess H+ ions in the blood, the bicarbonate (HCO3-) ion acts as a weak base and accepts H+ to form carbonic acid.
• H+ + HCO3- → H2CO3 → H2O + CO2
• With excess OH- ions, buffer system (H2CO3) components release H+ ions to maintain pH.
• H2CO3 → H+ + HCO3-

Phosphate Buffer System

• This buffer system can maintain a physiological pH of 7.4.
• Phosphate concentration is highest in intracellular fluids.
• The phosphate buffer system helps regulate pH of the cytosol.
• This system of monohydrogenphosphate (HPO4-2) and dihydrogenphosphate (H2PO4-) anions is present in cells and kidneys.
• HPO4-2 acts as a weak base by accepting a proton during excess H+ ion concentrations.
• On the other hand, the dihydrogenphosphate ion acts as the weak acid and neutralizes alkaline conditions.

Protein Buffer System

• This is the most copious buffer in body cells and plasma.
• Composed of amino acids with one COOH and one NH2 group.
• In an acidic situation, the amino (NH2) group acts as a base and accepts a proton: R-NH2 + H+ → R-NH3+.
• When in alkaline medium, the free COOH group liberates protons to neutralize: R-COOH → R-COO− + H+.
• These proteins provide acid and base functionality, due to amphoteric character.

Hemoglobin

• At a physiological pH, histidine and cysteine make for important amino acid buffers.
• Since hemoglobin registers as protein content due to its 37 histidine residues in structure, its performance helps as a physiological buffer.
• When CO2 enters erythrocytes from body cells, it merges rapidly with water when carbonic anhydrase acts upon it.

Physiological pH

• During alkaline conditions, carbonic acid releases into H+ and HCO3- ions while diffusing, creating a more plasma stable red blood cell buffer
• Concurrently, with less Hb, its carries negative load to carry oxygen while neutralizing high acidic excess erythrocyte-acid levels
• When present, the oxygen concentration forms an unstable equilibrium state where Hb concentrations rapidly combine deoxygenating cells while also delivering protination
• If the protons unite instead with the bicarbonate, there is acid action, the blood and other fluids in an organism
• Consequently, rapid respiration lowers H+ concentration to basic pH (alkalosis).
• Less respiration increases it to acidic pH (acidosis).