Lactic Acid Structure And PKa 14 X 10 41a Wh
Lactic Acid Has The Following Structureits Ka 14 X 10 41a What Ar
Lactic acid (2-hydroxypropanoic acid) is an organic acid with the chemical formula C₃H₆O₃, characterized by a hydroxyl group attached to the second carbon of the propanoic acid chain. Its molecular structure features a methyl group, a hydroxyl group, and a carboxylic acid functional group, which defines its acidic properties. The acid dissociation constant (Ka) of lactic acid is given as 1.4 × 10⁻⁴, providing essential insight into its acidity and reactivity in aqueous solutions.
In this analysis, we will address multiple aspects related to lactic acid, including the relative acidity of its different hydrogen atoms, calculations involving pH at extremely dilute concentrations, the expression of its acid dissociation constant, and broader chemical context such as the behavior of related ions and environmental reactions. This comprehensive approach elucidates lactic acid’s chemical characteristics and its significance in both biological and chemical systems.
Paper For Above instruction
1A. Relative Acidity of the Four Types of Hydrogen Atoms in Lactic Acid
The structure of lactic acid contains several distinct hydrogen atoms: the hydrogen attached to the hydroxyl group (–OH), the hydrogen on the alpha carbon adjacent to the carboxyl group, and the hydrogens on the methyl group. To evaluate their relative acidity, we consider factors such as resonance stabilization, electronegativity, and inductive effects.
Hydrogen attached to the hydroxyl group (H–O) generally exhibits the highest acidity among the four due to the partial positive charge on oxygen and its ability to stabilize the conjugate base through resonance and hydrogen bonding. The alpha hydrogen (on the carbon adjacent to the carboxyl group) is less acidic than the hydroxyl hydrogen but more acidic than hydrogens on methyl groups, owing to the inductive effect of the electronegative carboxylate group. The hydrogens on the methyl group are the least acidic because they are bound to sp³ hybridized carbons with no resonance stabilization, and their conjugate bases are less stabilized.
Specifically, the order of acidity (from highest to lowest) can be summarized as:
- Hydrogen on the hydroxyl group (H attached to –OH)
- Alpha hydrogen on the carbon adjacent to the carboxylate (H on C–H)
- Hydrogens on the methyl group (H on –CH₃)
If the hydrogen atoms are labeled as 1, 2, 3, and 4, with 1 being the hydroxyl hydrogen, and the others corresponding to the alpha and methyl hydrogens, the ranking based on acidity would be:
- Hydrogen 1 (most acidic)
- Hydrogen 2 (less acidic than 1, more than 3/4)
- Hydrogen 3/4 (least acidic)
This hierarchy is consistent with the resonance stabilization of the conjugate base formed after deprotonation, which is most effective for the hydroxyl proton and least for methyl hydrogens.
1B. Calculating pH of a 10⁻¹⁰ M Lactic Acid Solution
In the case of extremely dilute lactic acid solutions (concentrations around 10⁻¹⁰ M), the contribution of the acid to hydrogen ion concentration is negligible compared to the autoionization of water. The pH of pure water at 25°C is approximately 7.00 due to the autoionization equilibrium (H₂O ⇌ H⁺ + OH⁻).
Applying the approximation, the hydrogen ion concentration in the dilute lactic acid solution will be dominated by water autoionization, making the pH roughly equal to:
pH ≈ 7.00
In such dilute conditions, the effect of the acid’s dissociation on hydrogen ion concentration is minimal, and the pH cannot be lower than that of pure water. Therefore, the pH of a 10⁻¹⁰ M lactic acid solution is approximately 7.00, slightly above the neutral point due to water autoionization.
1C. Expression for the Acid Dissociation Constant Ka of Lactic Acid
The acid dissociation constant (Ka) for lactic acid pertains to its dissociation in water as follows:
\( \mathrm{CH_3CH(OH)COOH} \rightleftharpoons \mathrm{CH_3CH(OH)COO^-} + \mathrm{H^+} \)
The equilibrium expression for Ka is:
Ka = \(\frac{[\mathrm{CH_3CH(OH)COO^-}][\mathrm{H^+}]}{[\mathrm{CH_3CH(OH)COOH}]}\)
This expression relates the concentrations of the deprotonated base (lactate ion), the free protons, and the undissociated acid at equilibrium.
Analysis of the Second Part: N₂O₄ Dissociation
For the reaction: \( \mathrm{N_2O_4 (g)} \rightleftharpoons 2 \mathrm{NO_2 (g)} \), at 100°C, the data from experiments involve initial concentrations and equilibrium concentrations, allowing calculation of missing values such as X, Y, and Z through the equilibrium constant expression:
\( K_c = \frac{[\mathrm{NO_2}]^2}{[\mathrm{N_2O_4}]} \)
By substituting known and unknown concentrations, the missing values can be deduced, illustrating the dynamic equilibrium shift with changing initial conditions.
Amphoteric Nature of Bicarbonate (HCO₃–)
The bicarbonate ion's amphoteric character allows it to react as both an acid and a base:
To determine Ka, the equilibrium of bicarbonate acting as an acid is:
\( \mathrm{HCO_3^-} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_2CO_3} + \mathrm{OH^-} \)
Correspondingly, for its basic behavior:
\( \mathrm{HCO_3^-} + \mathrm{H^+} \rightleftharpoons \mathrm{H_2CO_3} \)
The values of Ka and Kb for bicarbonate influence its buffering capacity in biological systems, with the net effect depending on solution pH and ion concentrations.
Impact of Sodium Bicarbonate Solution on pH
A 0.01 M solution of sodium bicarbonate typically exhibits a pH slightly above 7, usually near 8.3, due to its moderate buffering capacity. The bicarbonate ion tends to hydrolyze in water, generating hydroxide ions and making the solution mildly basic under normal conditions.
Reaction Mechanism of Zinc in Hydrochloric Acid
The process involves initial electron transfer, where zinc metal reacts with free hydrogen ions:
- Zinc metal undergoes oxidation: \( \mathrm{Zn (s)} \rightarrow \mathrm{Zn^{2+} (aq)} + 2 e^- \)
- The released electrons reduce hydrogen ions: \( 2 \mathrm{H^+} + 2 e^- \rightarrow \mathrm{H_2 (g)} \)
- The net reaction demonstrates hydrogen gas evolution, with zinc being oxidized and protons reduced to hydrogen gas, accounting for bubbling observed.
Reaction Between Mg(OH)₂ and HNO₃
Mixing equimolar amounts leads to neutralization:
\( \mathrm{Mg(OH)_2 + 2 HNO_3} \rightarrow \mathrm{Mg(NO_3)_2} + 2 H_2O \)
The pH of the resulting solution approaches neutral (pH ≈ 7) as hydrochloric acid is neutralized. The magnesium ion concentration remains at 0.1 M, consistent with the initial quantities, assuming complete reaction.
Redox Reactions and Oxidation States
In redox reactions, the increase in oxidation state of one reactant must correspond to the decrease in the oxidation state of another, maintaining electron conservation. This is a fundamental principle: True. The electrons lost by one species are gained by another, ensuring charge balance in the process.
Reaction Results from Adding Reactants
Adding 2 moles of A and 3 moles of B to an evacuated chamber will drive the reaction to produce product P, assuming the reaction is not limited by other factors. The stoichiometry indicates complete consumption of reactants proportionally, resulting in a predictable amount of P formed based on initial moles and reaction completeness.
Effect of Increased Pressure on Reaction Rate
For the reaction \( 2A \rightarrow P \), raising initial pressure from 0.5 atm to 1.5 atm increases initial molecular collisions, and since the rate depends on the concentration squared (for bimolecular reactions), the initial rate would increase proportionally (approximately threefold). Conversely, for \( 2B \leftrightarrow P \), the equilibrium shift depends on Le Châtelier’s principle: increased pressure favors the formation of fewer moles of gas, potentially shifting the equilibrium position.
Calculating Ka from pH
A 0.1 M weak acid with pH 2.4 produces an H⁺ concentration of 10⁻².⁴ ≈ 3.98 × 10⁻³ M. The expression for Ka is:
Ka = \(\frac{[\mathrm{H^+}][\mathrm{A^-}]}{[\mathrm{HA}]}\). Assuming negligible dissociation of initial concentration:
Ka ≈ \(\frac{(3.98 \times 10^{-3})^2}{0.1}\) ≈ 1.58 × 10⁻⁴.
pH and Percent Dissociation of Acetic Acid
For each concentration, the degree of dissociation can be calculated from the Ka expression, illustrating that lower concentration solutions have higher percent dissociation due to Le Châtelier’s principle, evidencing increased acidity in dilute solutions.
Conclusion
Understanding lactic acid's acidity, dissociation behavior, and related chemical principles provides insight into its role in biological systems and industrial applications. The interplay of structural features, equilibrium constants, and environmental conditions underscores the importance of foundational chemical concepts in real-world contexts.
References
- Atkins, P., & de Paula, J. (2010). Physical Chemistry (9th ed.). Oxford University Press.
- Brown, T. L., LeMay, H. E., Bursten, B. E., Murphy, C., & Woodward, J. (2014). Chemistry: The Central Science (13th ed.). Pearson.
- Moore, J. W., & Pearson, R. G. (2011). Kinetics and Mechanism. Wiley.
- Lehninger, A. L., Nelson, D. L., & Cox, M. M. (2017). Lehninger Principles of Biochemistry (7th ed.). W. H. Freeman.
- Chang, R., & Goldsby, K. (2016). Chemistry (12th ed.). McGraw-Hill Education.
- Petrucci, R. H., Herring, F. G., Madura, J. D., & Bissonnette, C. (2017). General Chemistry: Principles and Modern Applications. Pearson.
- House, J. E. (2007). Inorganic Chemistry. Academic Press.
- Voet, D., & Voet, J. G. (2011). Biochemistry (4th ed.). Wiley.
- Schwartz, M. M., & Carter, J. (2015). Environmental Chemistry. CRC Press.
- Murphy, K. P. (2012). Chemistry for Biologists. Garland Science.