A 1.00 L solution saturated at 25 °C with calcium oxalate (CaC2O4) contains 0.0061 g of CaC2O4. Calculate the Ksp for this salt at 25 °C.

Calculate the change in Gibbs free energy, ΔG, for the following reaction, H2(g) + Zn2+(aq) → Zn(s) + 2 H+(aq), where the standard reduction potential for the reduction of Zn2+ to Zn(s) is -0.76 V. Given the Faraday constant is 96,485 J/V-mol.

Predict whether the equivalence point for a titration of NH3 with HCl will be above, below, or at pH 7.0.

Given the standard reduction potentials for the following half reactions. Determine which combination of these half-cell reactions leads to the cell reaction with the largest positive cell potential. Ag+(aq) + e- → Ag(s) 0.80V Cu2+(aq) + 2e- → Cu(s) 0.34V Ni2+(aq) + 2e- → Ni(s) -0.28 V Cr3+(aq) + 3e- → Cr(s) -0.41 V

Consider the reaction, B(aq) + H2O(l) → BH+(aq) + OH-(aq). If the salt of BH+(aq) is added, how is the equilibrium constant affected?

The first-order decomposition of N2O5 was studied at a particular temperature and found to have a rate constant of 6.2 x 10-4s-1. What is the half-life for the decomposition under these conditions?

A buffer is made by adding 0.300 mol CH3COOH and 0.300 mol CH3COONa to enough water to make 1.000 L of solution. The pH of the buffer is 4.74. Calculate the pH of this solution after 15.0 mL of 4.0 M NaOH(aq) solution is added.

If the equilibrium concentration (also known as the molar solubility) of CaF2 is 1.24 x 10-3M at 35 °C, what is Ksp at this temperature?

Which would be correct expression of the solubility product constant, Ksp, for Fe(OH)3?

If the pH of a solution is 6.27, which of the following statements is true?