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# acids and bases

***

## **theories**

#### *arrhenius theory*

* an acid has an h in its formula and dissolves in water to form hydrogen ions, $$\mathrm{H}^{+}\_{(aq)}.$$
  * *e.g.* $$\mathrm{HNO\_3}*{(aq)} \rightarrow \mathrm{H}^{+}*{(aq)} + \mathrm{NO\_3}^{-}\_{(aq)}$$
* a base has an oh in its formula and dissolves in water to form hydroxide ions, $$\mathrm{OH}^{-}(aq).$$
  * *e.g.* $$\mathrm{Ca(OH)*{2,(s)}} \rightarrow \mathrm{Ca}^{2+}*{(aq)} + 2\mathrm{OH}^{-}\_{(aq)}$$

#### *davy theory*

* an acid is a substance that reacts with a metal to produce hydrogen gas

#### *brønsted lowry theory*

* an acid donates a proton, $$\mathrm{H}^{+}$$
* a base accepts a proton, $$\mathrm{H}^{+}$$
* brønsted-lowry acids and bases cannot exist in isolation
  * if an acid donates a proton, it must donate it to a base and vice versa

## **electrolytes**

* substances that dissociate/ionise in water to form a solution that conducts electricity
* there are strong electrolytes which completely dissociate in water and exist as ions in solution
  * $$\mathrm{NaCl}*{(s)} \rightarrow \mathrm{Na}^{+}*{(aq)} + \mathrm{Cl}^{-}\_{(aq)}$$
  * dissociation involves two ions splitting into two separate ions
* there are also weak electrolytes which partially ionise in water and exist as both molecules and the composite ions of the molecule in solution
  * $$\mathrm{NH\_3}*{(aq)} + \mathrm{H\_2O}(l) \rightleftharpoons \mathrm{NH\_4}^{+}*{(aq)} + \mathrm{OH}^{-}\_{(aq)}$$
  * ionisation involves a non-ionic molecule splitting into ions

## **hydrolysis of salts**

* salts are strong electrolytes and so they completely dissociate into ions
  * *e.g.* $$\mathrm{Na\_2CO\_3}*{(s)} \rightarrow 2\mathrm{Na}^{+}*{(aq)} + \mathrm{CO\_3}^{2-}\_{(aq)}$$
* it is then possible for some ions to undergo a further reaction with water
  * *e.g.* $$\mathrm{CO\_3^{2-}} + \mathrm{H\_2O} \rightleftharpoons \mathrm{HCO\_3^{-}} + \mathrm{OH^{-}}$$

#### *neutral salts*

* neutral salts will contain the negative ion from a strong monoprotic acid
* ions like $$\mathrm{Cl^{-}}$$ and $$\mathrm{NO\_3^{-}}$$ will not undergo hydrolysis
* cations of group 1 and 2 metals will not hydrolyse in water

#### *basic salts*

* the negative ion that is the conjugate base of a weak acid such as the ethanoate ion, $$\mathrm{CH\_3COO^{-}},$$ or $$\mathrm{HCO\_3^{-}},$$ undergo hydrolysis to make a basic solution
  * $$\mathrm{CH\_3COO^{-}} + \mathrm{H\_2O} \rightleftharpoons \mathrm{CH\_3COOH} + \mathrm{OH^{-}}$$
  * the formation of the hydroxide ion produces a basic solution
* *acidic salts:*
  * the cation of the conjugate acid of a weak base, such as the ammonium ion, $$\mathrm{NH\_4^{+}},$$ produces an acidic solution
    * $$\mathrm{NH\_4^{+}} + \mathrm{H\_2O} \rightleftharpoons \mathrm{NH\_3} + \mathrm{H\_3O^{+}}$$

## **pH**

* water will self-ionise and undergo the following equilibrium reaction:
  * $$2\mathrm{H\_2O}(l) \rightleftharpoons \mathrm{H\_3O^{+}}*{(aq)} + \mathrm{OH^{-}}*{(aq)}$$
* $$K\_w = \[\mathrm{H\_3O^{+}}]\[\mathrm{OH^{-}}]$$
  * at 25°C, the value of $$K\_w$$ is $$10^{-14}$$
* if the temperature of water increases, the self-ionisation of water increases
  * $$\text{heat} + \mathrm{H\_2O}(l) \rightleftharpoons \mathrm{H}^{+}*{(aq)} + \mathrm{OH^{-}}*{(aq)}$$
    * endothermic as $$\Delta H > 0$$, leading to an increase in $$\[\mathrm{H}^{+}]$$ and $$\[\mathrm{OH^{-}}]$$ and thus also in $$K\_w$$
    * as $$\mathrm{pH} = -\log\_{10}\[\mathrm{H}^{+}],$$ this causes an increase in measured pH at neutral

## **buffers**

* buffers are solutions that resist changes in pH when small amounts of an acid or base are added

#### *addition of* $${H\_3}O^+$$

* on addition of $$\mathrm{H\_3O^{+}},$$ the concentration of $$\mathrm{H\_3O^{+}}$$ increases
* this increases the rate of the reverse reaction as there is an increase in the frequency of collisions between products
* this means more products are consumed than produced and this decreases the concentration of $$\mathrm{H\_3O^{+}}$$ until equilibrium is reestablished so only a small increase in the concentration of $$\mathrm{H\_3O^{+}}$$ occurs
* as $$\mathrm{pH} = -\log\_{10}\[\mathrm{H\_3O^{+}}],$$ pH only decreases slightly
* $$\text{base} + \mathrm{H\_3O^{+}} \rightarrow \text{conjugate acid} + \mathrm{H\_2O}$$

#### *addition of* $$OH^{-}$$

* $$\mathrm{OH^{-}}$$ reacts directly with the weak acid when added
* $$\text{weak acid} + \mathrm{OH^{-}} \rightarrow \text{conjugate base} + \mathrm{H\_2O} \quad (\text{equilibrium shifts left})$$
* this causes a slight decrease in the concentration of $$\mathrm{H\_3O^{+}}$$ as the equilibrium shifts a small amount to the left
* as $$\mathrm{pH} = -\log\_{10}\[\mathrm{H\_3O^{+}}],$$ pH only increases slightly

## **ocean issues**

#### *reactions*

1. $$\mathrm{CO\_2}(g) \rightleftharpoons \mathrm{CO\_2}\_{(aq)}$$
2. $$\mathrm{CO\_2}*{(aq)} + \mathrm{H\_2O}(l) \rightleftharpoons \mathrm{H\_2CO\_3}*{(aq)}$$
3. $$\mathrm{H\_2CO\_3}*{(aq)} + \mathrm{H\_2O}(l) \rightleftharpoons \mathrm{HCO\_3^{-}}*{(aq)} + \mathrm{H\_3O^{+}}\_{(aq)}\quad (K\_1 = 2.50\times10^{-4}\ \text{at }25^\circ\mathrm{C})$$
4. $$\mathrm{HCO\_3^{-}}*{(aq)} + \mathrm{H\_2O}(l) \rightleftharpoons \mathrm{CO\_3^{2-}}*{(aq)} + \mathrm{H\_3O^{+}}\_{(aq)}\quad (K\_2 = 4.69\times10^{-11}\ \text{at }25^\circ\mathrm{C})$$

#### *ocean acidification*

* le châtelier's principle states that if a system at equilibrium is subjected to a change in conditions, the system will adjust to re-establish equilibrium in such a way as to partially counteract the imposed change
* an increase in concentration of atmospheric carbon dioxide would cause equilibrium 1 to shift to the right as it dissolves into the seawater, increasing $$\[\mathrm{CO\_2}]$$
* this increase causes equilibrium 2 to shift to the right and increase $$\[\mathrm{H\_2CO\_3}]$$
* by equilibrium 3, $$\mathrm{H\_2CO\_3}$$ can be ionised to form hydrogen carbonate and hydronium ions. as $$\[\mathrm{H\_2CO\_3}]$$ is increased the concentration of both of these substances is also increased
* as $$\mathrm{pH} = -\log\_{10}\[\mathrm{H\_3O^{+}}],$$ this increase in $$\[\mathrm{H\_3O^{+}}]$$ results in a decrease in pH and hence an increase in acidity of seawater

#### *shell thinning*

* because $$K\_3$$ is much larger than $$K\_4,$$ equilibrium 3 lies much further to the right than equilibrium 4 and thus $$\[\mathrm{HCO\_3^{-}}]$$ is much larger than $$\[\mathrm{CO\_3^{2-}}]$$ and $$\[\mathrm{H\_3O^{+}}]$$
* as ocean acidification increases the concentration of $$\mathrm{H\_3O^{+}}$$ ions in seawater, equilibrium 4 is driven towards the left causing a decrease in $$\[\mathrm{CO\_3^{2-}}]$$ in seawater
* due to this, the calcification process is slowed as less carbonate ions are being deposited in the shells of marine animals, causing them to grow thinner and less protective shells