91Ó°ÊÓ

Calculate the \(\left[\mathrm{H}^{+}\right]\) in each of the following solutions, and indicate whether the solution is acidic or basic. a. \(\left[\mathrm{OH}^{-}\right]=5.99 \times 10^{-8} \mathrm{M}\) b. \(\left[\mathrm{OH}^{-}\right]=8.99 \times 10^{-6} \mathrm{M}\) c. \(\left[\mathrm{OH}^{-}\right]=7.00 \times 10^{-7} \mathrm{M}\) d. \(\left[\mathrm{OH}^{-}\right]=1.43 \times 10^{-12} \mathrm{M}\)

Calculate the \(\left[\mathrm{H}^{+}\right]\) in each of the following solutions, and indicate whether the solution is acidic, basic, or neutral. a. \(\left[\mathrm{OH}^{-}\right]=3.99 \times 10^{-5} \mathrm{M}\) b. \(\left[\mathrm{OH}^{-}\right]=2.91 \times 10^{-9} \mathrm{M}\) c. \(\left[\mathrm{OH}^{-}\right]=7.23 \times 10^{-2} \mathrm{M}\) d. \(\left[\mathrm{OH}^{-}\right]=9.11 \times 10^{-7} \mathrm{M}\)

Calculate the \(\left[\mathrm{OH}^{-}\right]\) in each of the following solutions, and indicate whether the solution is acidic, basic, or neutral. a. \(\left[\mathrm{H}^{+}\right]=8.89 \times 10^{-7} \mathrm{M}\) b. \(\left[\mathrm{H}^{+}\right]=1.19 \times 10^{-7} \mathrm{M}\) c. \(\left[\mathrm{H}^{+}\right]=7.00 \times 10^{-7} \mathrm{M}\) d. \(\left[\mathrm{H}^{+}\right]=1.00 \times 10^{-7} \mathrm{M}\)

Calculate the \(\left[\mathrm{OH}^{-}\right]\) in each of the following solutions, and indicate whether the solution is acidic or basic. a. \(\left[\mathrm{H}^{+}\right]=1.34 \times 10^{-2} \mathrm{M}\) b. \(\left[\mathrm{H}^{+}\right]=6.99 \times 10^{-7} \mathrm{M}\) c. \(\left[\mathrm{H}^{+}\right]=4.01 \times 10^{-9} \mathrm{M}\) d. \(\left[\mathrm{H}^{+}\right]=4.02 \times 10^{-13} \mathrm{M}\)

For each pair of concentrations, tell which represents the more acidic solution. a. \(\left[\mathrm{H}^{+}\right]=1.2 \times 10^{-3} \mathrm{M}\) or \(\left[\mathrm{H}^{+}\right]=4.5 \times 10^{-4} \mathrm{M}\) b. \(\left[\mathrm{H}^{+}\right]=2.6 \times 10^{-6} \mathrm{M}\) or \(\left[\mathrm{H}^{+}\right]=4.3 \times 10^{-8} \mathrm{M}\) c. \(\left[\mathrm{H}^{+}\right]=0.000010 \mathrm{M}\) or \(\left[\mathrm{H}^{+}\right]=0.0000010 \mathrm{M}\)

For each pair of concentrations, tell which represents the more basic solution. a. \(\left[\mathrm{H}^{+}\right]=1.59 \times 10^{-7} \mathrm{M}\) or \(\left[\mathrm{H}^{+}\right]=1.04 \times 10^{-8} \mathrm{M}\) b. \(\left[\mathrm{H}^{+}\right]=5.69 \times 10^{-8} \mathrm{M}\) or \(\left[\mathrm{OH}^{-}\right]=4.49 \times 10^{-6} \mathrm{M}\) c. \(\left[\mathrm{H}^{+}\right]=5.99 \times 10^{-8} \mathrm{M}\) or \(\left[\mathrm{OH}^{-}\right]=6.01 \times 10^{-7} \mathrm{M}\)

Why do scientists tend to express the acidity of a solution in terms of its \(\mathrm{pH},\) rather than in terms of the molarity of hydrogen ion present? How is \(\mathrm{pH}\) defined mathematically?

For a hydrogen ion concentration of \(2.33 \times 10^{-6} \mathrm{M}\) how many decimal places should we give when expressing the pH of the solution?

As the hydrogen ion concentration of a solution increases, does the pH of the solution increase or decrease? Explain.

Calculate the \(\mathrm{pH}\) corresponding to each of the hydrogen ion concentrations given below, and indicate whether each solution is acidic or basic. a. \(\left[\mathrm{H}^{+}\right]=4.97 \times 10^{-7} \mathrm{M}\) b. \(\left[\mathrm{H}^{+}\right]=1.01 \times 10^{-12} \mathrm{M}\) c. \(\left[\mathrm{H}^{+}\right]=2.49 \times 10^{-4} \mathrm{M}\) d. \(\left[\mathrm{H}^{+}\right]=1.00 \times 10^{-7} \mathrm{M}\)