The enthalpy change of solution, $\Delta H^\circ_{\text{sol}}$, of the Group 2 sulfates becomes more endothermic further down the group. State and explain the change in solubility of the Group 2 sulfates down the group.
State the expression for $K_w$, the ionic product of water.
The numerical value of $K_w$ rises as temperature rises. Put a tick ($\checkmark$) in the correct column in each row to show the effect of increasing the temperature of water on the pH and on the ratio $[\text{H}^+] : [\text{OH}^-]$.
At 298 K, an aqueous solution of sodium hydroxide has a pH of 13.25. Calculate the concentration of this sodium hydroxide solution.
Buffer solutions are used to control the pH of a solution so that its pH remains within a narrow range. Write two equations to show how hydrogencarbonate ions, $\text{HCO}_3^-$, and carbonic acid, $\text{H}_2\text{CO}_3$, regulate the pH of blood.
The $K_{\text{a}}$ for ethanoic acid is $1.75 \times 10^{-5}\,\text{mol dm}^{-3}$ at $298\,\text{K}$. When ethanoic acid dissolves in water, an equilibrium mixture containing two acid-base pairs is formed. Write an equation for this equilibrium. In the boxes, label each species as acidic or basic to show how it behaves in this equilibrium.
A buffer solution was made by adding $30.0\,\text{cm}^3$ of $0.25\,\text{mol dm}^{-3}$ ethanoic acid, in excess, to $20.0\,\text{cm}^3$ of $0.15\,\text{mol dm}^{-3}$ sodium hydroxide. Calculate the pH of the buffer solution formed at $298\,\text{K}$. Give your answer to one decimal place.
Titration curves for two different acid-base reactions, $M$ and $N$, are shown. Use the titration curve for reaction $M$ to deduce the volume of acid added at the end-point for this titration.
The table gives some acid-base indicators. Name an appropriate indicator for each of the acid-base titrations $M$ and $N$. Explain your answers.