Korean, Edit

Chapter 18. Acid-Base Reactions

Recommended post: 【Chemistry】 Chemistry Index

1. Definitions of Acids and Bases

2. Ionization Equilibrium

3. Acid Dissociation Constant (Ka) and pK_a

4. Buffer Solutions

5. Titration

6. Amino Acid Titration

1. Definitions of Acids and Bases

⑴ Everyday Acids and Bases

① General Properties of Acids

○ Sour Taste: Sour taste indicates a reaction with H+ ions.

○ Reaction with Metals Producing Hydrogen

○ pH < 7 at 25 ℃

② General Properties of Bases

○ Bitter Taste: Bitter taste indicates a reaction with hydroxide ions.

○ Slippery Feel

○ pH > 7 at 25 ℃

⑵ Arrhenius Acid-Base Theory

① Acid: A substance that donates hydrogen ions in an aqueous solution.

② Base: A substance that donates hydroxide ions in an aqueous solution.

③ Special Properties of Hydrogen

○ Hydrogen, being the lightest element, can participate in numerous chemical reactions.

④ Limitations

○ Cannot explain substances in a non-aqueous state.

○ Dissolution of NH3 in water exhibits basic properties but doesn’t satisfy the Arrhenius base definition.

⑶ Brønsted–Lowry Acid-Base Theory

① Proposed simultaneously by Brønsted of Denmark and Lowry of England in 1923.

② Acid: Proton donor (H+ donor)

③ Base: Proton acceptor (H+ acceptor)

④ Relativity: The concepts of acids and bases are relative.

○ Conjugate Base: A Brønsted–Lowry acid loses a proton to become a conjugate base.

○ Conjugate Acid: A Brønsted–Lowry base gains a proton to become a conjugate acid.

○ Amphiprotic: Substances like water, alcohols can act as either acids or bases under normal conditions.

⑤ Limitations

○ Acids must include acidic hydrogen atoms.

○ Substances that do not donate or accept hydrogen ions cannot be explained.

⑷ Lewis Acid-Base Theory: Describes all chemical reactions without oxidation state changes.

① Introduced by Lewis, an American physicist.

② Acid: Electron pair acceptor

③ Base: Electron pair donor, similar to nucleophiles

④ All Brønsted–Lowry acids are Lewis acids, but the reverse is not true.

○ Example 1: BF3 has an empty p orbital, making it a Lewis acid but not a Brønsted–Lowry acid.

○ Example 2: AlCl3

⑤ Lewis bases have a concept similar to Brønsted–Lowry bases.

⑥ Example: AlCl3 + NH3 → Cl3Al-NH3

○ The product satisfies the octet rule more than the reactant, favoring the forward reaction.

○ Lewis acid AlCl3 receives shared electron pairs from the Lewis base NH3, promoting the reaction.

⑸ Neutralization Reaction

① Reaction between acids and bases

② Acid-base reactions in aqueous solutions generate water and salts, releasing heat (heat of neutralization).

2. Ionization Equilibrium

⑴ Ionization Constant: Equilibrium constant for ionization reactions of acids and bases in aqueous solutions.

① Acid Dissociation Constant K_a

② Base Dissociation Constant K_b

⑵ Ionization of Water and pH

① For a conjugate acid-conjugate base pair, the following equation holds:

② pH + pOH = pKw

③ At neutral conditions (pH = pOH = 7) at 25 ℃, Kw = 10^-14

○ Water ionization is endothermic, so higher temperature increases Kw.

○ With increasing temperature, neutral pH decreases.

④ Useful Equation: pK_a + pKb = pKw

⑶ Degree of Ionization

① Experimental Results and Degree of Ionization

② Strong Acids and Strong Bases: Assumed to be 100% ionized

○ Leveling Effect: Strong acids stronger than the solvent’s conjugate acid and strong bases stronger than the solvent’s conjugate base are all converted.

○ Stronger acids than H3O+ cannot exist in water.

○ Stronger bases than OH- cannot exist in water.

○ H2SO4 is stronger than H3O+, so it is fully converted to HSO4- in solution.

○ Dissolution of H2SO4 in water cannot undergo a reaction with H2SO4 as the reactant, except for acid-base reactions with water.

○ Essence of the Leveling Effect

○ In the context of the leveling effect, consider HA1 as a strong acid and HA2 as H3O+.

○ Generally, the ratio of equilibrium constants is exponential, making their ratio very large or very small.

○ Since Ka1 > Ka2, Ka1 ÷ Ka2 is significantly large.

○ The third reaction has a very large equilibrium constant, so nearly 100% of it undergoes the forward reaction.

③ Weak Acids and Weak Bases: Assumed to be partially ionized

**3. Acid Dissociation Constant (Ka) and pK_a **

⑴ Using pK_a values to Compare Acidity and Basicity: pK_a = -logKa

⑵ Strength of Acids and Bases

① Larger Ka implies stronger acids, and larger Kb implies stronger bases.

② Since Ka × Kb is constant, if an acid is strong, its conjugate base is weak, and vice versa.

⑶ ΔpK_a

① ΔpK_a > 0: Forward reaction

② ΔpK_a < 0: Reverse reaction

③ ΔpK_a = 0: Equilibrium state

④ | ΔpK_a | ≤ 7: Equilibrium reaction between forward and reverse reactions

⑤ | ΔpK_a | > 7: Unidirectional reaction between forward and reverse reactions

4. Buffer Solutions

⑴ Definition: Solutions that maintain a nearly constant pH upon addition of small amounts of acid or base (e.g., carbonic acid buffer action)

⑵ Conditions: Solutions containing a weak acid and its conjugate base can act as buffer solutions.

① Method 1: Mixing a weak acid and a stronger

base with fewer moles

② Method 2: Mixing a weak base and a stronger acid with fewer moles

③ Mixing different weak acids or weak bases that are not conjugate acid-base pairs will not form buffer solutions.

○ Each chemical species progresses in a way that all are either conjugate acids or conjugate bases.

○ Since K = K1 ÷ K2, K is either much smaller or much larger than 1.

⑶ Buffer Capacity: The amount of H+ or OH- required to cause a significant change in pH in a unit volume of buffer solution.

① When the total amount of HA and A- is constant and only the ratio of [HA] to [A-] changes, the buffer capacity is greatest when [HA] = [A-].

○ [HA] > [A-]: Buffer capacity for bases is greater than that for acids.

○ [HA] < [A-]: Buffer capacity for acids is greater than that for bases.

② The larger the total amounts of HA and A-, the greater the buffer capacity.

③ Problems comparing buffer capacity with changing ratios and total amounts of [HA] and [A-] are not typically presented.

⑷ Buffer solutions work well when pK_a - 1 < pH < pK_a + 1 (i.e., pH change « 1) experimentally demonstrated.

⑸ Types of Buffer Solutions

① Carbonated Water: Mainly observed in the ocean, with pH around 8, where HCO3- is present. Since it’s a divalent cation, there are two buffering ranges.

② Phosphoric Acid: The second most commonly observed buffering system in nature after carbonated water. Since it’s a trivalent cation, there are three buffering ranges.

③ Protein Buffer System

④ Ammonia Buffer system

⑤ PBS (Phosphate Buffered Saline): A buffer frequently used in chemical experiments with a pH of around 7.4.

⑥ [DPBS](https://jb243.github.io/pages/445#:~:text=6.-,DPBS,-(Dulbecco%E2%80%99s%20Phosphate%20Buffer) (Dulbecco’s PBS): Removes Mg2+, Ca2+, and others involved in cell adhesion, making it easier to detach cells.

⑦ Tris Buffer (Tris-HCl buffer): Buffer inside test tubes

○ Weak base: pKb = 5.9

○ Tris(aq) + H+(aq) ⇄ TrisH+(aq)

⑧ HEPES Buffer: Buffer inside cells

⑨ STE

⑩ NaAc (Sodium Acetate)

⑪ NH4Ac (Ammonium Acetate)

⑫ MES (2-(N-morpholino)ethanesulfonic acid)

⑬ RIPA Buffer: Used for solubilizing cells during experiments

⑭ Loading Buffer: Used in gel electrophoresis (Laemmli sample buffer)

⑮ Transfer Buffer: Used in gel transfer

⑯ Sodium Cacodylate Buffer: Used in pre-processing bio-TEM samples

⑰ MOPS buffer

5. Titration

⑴ Titration: Experiment to measure the concentration of an unknown acid or base

① The titration solution should be a strong acid or base.

② Reason for ①: Complete neutralization should occur for accurate results.

⑵ Equivalence point: Point where the moles of acid (or base) within the unknown solution equal the moles of added base (or acid)

⑶ Indicator: Substance used to determine pH by color change

① End point: Also known as the pKIn where [HIn] = [In-]

○ Consideration of indicator’s color change range in titration experiments: pKIn - 1 ＜ Equivalence point pH ＜ pKIn + 1

○ All indicators used in acid-base titrations are weak acids or bases according to the derived equation

○ Equivalence point represents properties of the unknown reagent, while the end point represents properties of the indicator

② Common indicators

Table 1. Types of Indicators

③ Special indicators

○ Example 1. Starch Indicator: I- doesn’t color, but iodine complex (e.g., I3-) does

④ Principle of Indicator Color Change

○ π bonds can absorb visible light: The normal molecular vibrational frequency is in the ultraviolet range, but with broadened resonance, it extends to visible light

○ Longer conjugation leads to longer absorption wavelength

○ Absorption and observation colors based on absorption wavelength)

Figure 1. Shape change of phenolphthalein with respect to pH

Figure 2. Why phenolphthalein in its basic form absorbs visible light (particles inside the box)

○ However, under strongly acidic conditions, the resonance structure of phenolphthalein elongates, resulting in an orange-red color.

Figure 3. Indicator color change of phenolphthalein based on pH

⑷ Ionization equilibrium of monoprotic acids (1-proton acids)

① Reaction equation

② Henderson-Hasselbalch equation

③ When titrated with a strong base: Assume HA exists only in α mmol initially

○ Initial point

○ Half-equivalence point: When 0.5α mmol of strong base is added

○ Equivalence point: When α mmol of strong base is added

⑸ Ionization equilibrium of diprotic acids (2-proton acids) (e.g., H2CO3)

① Reaction equation

② Henderson-Hasselbalch equation

③ When titrated with a strong base: Assume H2A exists only in α mmol initially

Figure 4. Titration of diprotic acid with strong base

○ Initial point

○ First half-equivalence point: When 0.5α mmol of strong base is added

○ First equivalence point: α mmol of strong base added. Note: [H2A] and [A2-] are equal due to symmetry.

○ Second half-equivalence point: When 1.5α mmol of strong base is added

○ Second equivalence point: When 2α mmol of strong base is added

⑹ Ionization equilibrium of triprotic acids (3-proton acids) (e.g., H3PO4)

① Reaction equation

② Henderson-Hasselbalch equation

③ When titrated with a strong base: Assume H3A exists only in α mmol initially

○ Initial point

○ First half-equivalence point: When 0.5α mmol of strong base is added

○ First equivalence point: α mmol of strong base added

○ Second half-equivalence point: When 1.5α mmol of strong base is added. Note: [H3A] and [A3-] are equal due to symmetry.

○ Second equivalence point: When 2α mmol of strong base is added

○ Third half-equivalence point: When 2.5α mmol of strong base is added

○ Third equivalence point: When 3α mmol of strong base is added

6. Amino Acid Titration

⑴ Characteristics of amino acids

① Isoelectric point: pH where an amino acid is uncharged. Also the point of rapid pH change.

② Amino acids are zwitterions at neutral pH

⑵ Example 1. Neutral Amino Acid: Glycine (gly)

① Amino acid with -H as the functional group

② Carboxyl group: pK_a1 = 2.34

③ Amino group: pK_a2 = 9.60

④ Isoelectric point: When [NH3+-RH2-COOH] = [NH2-RH2-COO-]

⑤ Interpretation of Isoelectric Point

○ pH < pK_a1: NH3+-RH2-COOH dominant, NH3+-RH2-COO- weak

○ pK_a1 ＜ pH ＜ pK_a2: NH3+-RH2-COO- dominant, NH3+-RH2-COOH and NH2-RH2-COO- weak

○ pK_a2 ＜ pH: NH2-RH2-COO- dominant

○ Conclusion: Once the neutral molecule is identified, find the closest pK_a values before and after it, and take their average to obtain the isoelectric point (pI).

⑶ Example 2. Negatively Charged Amino Acids: Aspartic acid (Asp), Glutamic acid (Glu)

① Functional group pK_a is greater than 2.34

② Isoelectric point

③ Tip: The isoelectric point (pI) can be obtained by averaging the two closest pK_a values, as the pK_a of the functional group (pK_aR) is closer to pK_a1.

⑷ Example 3. Positively Charged Amino Acids: Arginine (Arg), Histidine (His), Lysine (Lys)

① Functional group pK_aR is less than 9.60

② Isoelectric point

③ Tip: The isoelectric point (pI) can be obtained by averaging the two closest pK_a values, as the pK_a of the functional group (pK_aR) is closer to pK_a2.

⑸ pK_a1, pK_a2, pK_aR, pI, Hydrophobicity index, Ratio in proteins for each amino acid

Table 2. pK_a1, pK_a2, pK_aR, pI, Hydrophobicity index, Ratio in proteins for each amino acid

⑹ Net Charge of Amino Acid Sequence, Isoelectric Point

Glu - His - Trp - Ser - Gly - Leu - Arg - Pro - Gly

① Calculating net charge at specific pH

○ Step 1. Compare pK_a values and pH for -NH3+ end, amino acid functional groups, and -COO- end, respectively.

○ -NH3+ end: pK_a = 9.60

○ Glu functional group: pK_a = 4.25

○ His functional group: pK_a = 6.00

○ Arg functional group: pK_a = 12.48

○ -COO- end: pK_a = 2.34

○ Step 2. All-or-none rule: For each amino acid, assume that the conjugate acid is present if pH < pK_a, and the conjugate base is present if pH > pK_a.

○ Step 3. Calculate charge for each amino acid and sum them up.

○ Note that it is not necessarily +1 for a conjugate acid and -1 for a conjugate base.

Table 3. Method for determining net charge at specific pH

② Isoelectric point

○ Method 1. Calculate isoelectric point using pK_a values

○ 1^st. Identify all ionizable groups: NH3+ end, Glu, His, Arg, COO- end in the above example.

○ 2^nd. Explore the interval where the net charge equals 0: 6.00 ＜ pH ＜ 9.60 in the above example.

○ 3^rd. Since the pI varies continuously with pH, average the two pK_a values that form the boundary of the interval

○ Method 2. Plot amino acid charge graph with respect to pH and find the point where the charge equals 0

Figure 5. Amino acid charge based on pH

Input: 2019.01.03 23:16

Modified: 2020.04.21 23:59