The pH scale and strong vs weak acids and bases

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A measure of acidity based on hydrogen-ion concentration: $\text{pH} = -\log_{10}[\text{H}^{+}]$.

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$\text{pH} = -\log_{10}[\text{H}^{+}]$ — given in the data booklet.

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$[\text{H}^{+}] = 10^{-\text{pH}}$ — the rearranged given equation.

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The ionic product of water, $K_{w} = [\text{H}^{+}][\text{OH}^{-}] = 1.0\times10^{-14}$ at 25 °C.

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pH < 7 acidic · pH = 7 neutral · pH > 7 basic (alkaline), at 25 °C.

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[H_{+}] changes by a factor of **10** (pH is a log scale).

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Strong = **fully** dissociated into ions; weak = only **partially** dissociated.

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No — strength is the **degree of dissociation**; concentration is the amount dissolved.

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Single arrow, e.g. $\text{HCl} \rightarrow \text{H}^{+} + \text{Cl}^{-}$ (full dissociation).

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Equilibrium arrows, e.g. $\text{CH}_{3}\text{COOH} \rightleftharpoons \text{H}^{+} + \text{CH}_{3}\text{COO}^{-}$ (partial).

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Strong acid has a **lower pH**, **higher conductivity** and a **faster** reaction (more H_{+} ions).

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It is fully dissociated, so it gives a **higher [H_{+}]**, and a higher [H_{+}] means a lower pH.