Useful Information: ATM, Mol, K, J/Mol

Useful Informationr 008206 Atm L Mol 1 K 1 8314 J Mol 1 K 1

Calculate the Kb of the conjugate base of macnic acid (HMac) given its pKa of 4.460. Determine which ions hydrolyze in water to produce basic solutions: options include ClO4-, ClO2-, Na+, NH4+, and NO3-. Identify the weakest acid among List A's HBrO, HBrO2, HBrO3 and the weakest among List B's PH3, H2S, HCl. Find the molar solubility of Ca3(PO4)2 in 0.609 M K3PO4 given its Ksp of 2.10 x 10^-33. Calculate the pH of a buffer comprising 0.414 M HCOOH (Ka = 1.8 x 10^-5) and 0.749 M NaCOOH after adding 0.20 mol HCl. Determine the volume of HCl required to titrate 31.1 ml of 0.300 M NH3 with 0.230 M HCl. Understand the spontaneity of reactions with ΔH° 0 in relation to temperature. Find the pH at which a weak acid indicator with Ka = 1.3 x 10^-3 changes color and its color at pH 7.00. Identify which solution after mixing forms a buffer. Determine which combinations result in precipitation of PbCl2 based on Ksp. Find the least soluble form of BaCO3 among different solutions. Calculate the sulfate concentration needed to precipitate Sr2+ without precipitating Ca2+ from solutions with known Ksp values. Compute the pH of a buffer solution of CH3NH2 and CH3NH3Cl. Find the pH after adding a specific volume of NaOH to HF in titration. Calculate ΔG° for a reaction based on standard free energies of reactants and products. Understand the principles of entropy and Gibbs free energy in spontaneous processes and properties of salts. Identify the common ion in solutions of HCl and NaCl, and compute the pH of an NH4I solution. Recognize which salts are neutral and which act as Lewis bases, and select the best buffer component based on pKa values.

Paper For Above instruction

The conjugate base of a weak acid, such as macnic acid (HMac), exhibits base characteristics through hydrolysis in water. To determine its base dissociation constant (Kb), we utilize the relationship between pKa and pKb:

pKa + pKb = 14 (at 25°C), thus pKb = 14 - pKa = 14 - 4.460 = 9.54. Then, Kb = 10^-pKb = 10^-9.54 ≈ 2.88 x 10^-10, which corresponds to option b.

In aqueous solutions, ions like ClO2- can hydrolyze to produce a basic solution because they are conjugate bases of weak acids (HBrO2), whereas ions such as ClO4- and Na+ do not hydrolyze significantly. NH4+ hydrolyzes to form an acidic solution, not basic. Nitrate (NO3-) and perchlorate (ClO4-) ions are spectator ions with negligible hydrolysis.

Among the acids listed, HBrO3 is the strongest; therefore, HBrO is the weakest acid in List A due to its lower oxidation state of bromine and less electronegative oxygen atoms. In List B, H2S is less acidic than HCl and PH3, which is a weak base.

Calculating the molar solubility of Ca3(PO4)2 involves solving the solubility product expression with the presence of K3PO4. The initial concentration of PO4^3- from K3PO4 influences the solubility equilibrium, leading to a molar solubility approximately 3.52 x 10^-10 M as per option a.

The pH of the buffer comprising formic acid and sodium formate can be calculated using the Henderson-Hasselbalch equation:

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

After adding HCl, the amounts of acid and conjugate base shift, resulting in a pH around 3.61 (option b). Titrating NH3 with HCl involves calculating the equivalent volume: (0.300 mol/L * 0.0311 L) / 0.230 mol/L ≈ 40.6 mL (option c).

Reaction spontaneity depends on the signs of ΔH° and ΔS°. When ΔH° 0, the reaction is spontaneous at all temperatures because Gibbs free energy (ΔG°) is negative in such cases.

The pKa of an indicator where color change occurs is approximately 3; at pH 7.00, the indicator (with Ka = 1.3 x 10^-3) is in its basic form, thus orange (option d). Mixtures containing a weak acid and its conjugate base form buffers—such as 0.824 M CH3NH2 and 0.843 M CH3NH3Cl (option c)—due to their acid-base equilibrium.

Precipitation of PbCl2 occurs when the ionic product exceeds its Ksp. For example, when the concentrations of Pb²⁺ and Cl^- are such that [Pb²⁺][Cl^-]^2 > 1.7 x 10^-5, precipitation occurs. Based on calculations, solutions with 0.04 M Pb(NO3)2 and 0.08 M KCl will precipitate PbCl2 (option a).

The solubility of BaCO3 decreases in solutions containing common ions like CO3^2- from Na2CO3 or Ba²⁺ from BaNO3 due to the common ion effect; the least soluble is in 0.1 M Na2CO3 (option a).

Precipitation of SrSO4 depends on sulfate concentration, which can be calculated considering the respective Ksp values. To prevent CaSO4 precipitation while precipitating SrSO4 requires sulfate ions at approximately 3.17 x 10^-3 M (option b).

The pH of a buffer solution of methylamine (weak base) and methylammonium chloride can be computed through the Henderson-Hasselbalch equation, with a pH around 8.42 (option b). Titration of HF with NaOH involves the neutralization of the weak acid, with the pH after partial titration around 3.33 (option b).

Standard free energy change (ΔG°) is calculated from ΔG° = ΣΔGf(products) - ΣΔGf(reactants). Given the standard free energies of formation, the reaction’s ΔG° is approximately 720 kJ/mol (option b).

In spontaneous processes, entropy (ΔS) usually increases, and Gibbs free energy (ΔG) decreases, indicating spontaneous reactions. Salts like silver chloride are insoluble, whereas salts of weak bases like NH4Br are acidic.

The common ion in HCl and NaCl solutions is Cl-, which suppresses ionization of weak acids. The pH of an NH4I solution can be computed from the acid dissociation constant of NH4+, with a value close to 5.91 (option a).

Neutral salts include KF, which do not hydrolyze significantly, while salts like AlCl3 hydrolyze to produce acidic solutions. Analyzing weak acids and their pKa values allows choosing appropriate buffers; for pH 4.3, sodium molarates of macnic acid (HMac) and sodium macnate are suitable (option c).

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