There are two very useful ways to define acids and bases.
- Brønsted–Lowry (general):
- An acid is a proton (H⁺) donor.
- A base is a proton (H⁺) acceptor.
- Arrhenius (in water):
- An acid is a substance that produces hydrogen ions, H⁺(aq), in water.
- A base is a substance that produces hydroxide ions, OH⁻(aq), in water.
In MYP/early IB, you mainly use the Arrhenius idea in water and the Brønsted–Lowry idea when thinking about proton transfer.
- The pH scale is a logarithmic scale used to measure how acidic or basic a solution is.
- pH < 7 → acidic
- pH = 7 → neutral
- pH > 7 → basic/alkaline
- Each step of 1 pH unit represents a 10× change in hydrogen ion concentration $[H^+]$.
- pH 3 is 10 times more acidic than pH 4.
- pH 2 is 100 times more acidic than pH 4.
- Very important: We never taste substances in the laboratory.
- Many acids and bases are corrosive and dangerous.
- Instead, we use observable properties and indicators.
Acids often:
- Have a pH less than 7.
- Are corrosive – they can react with some metals.
- React with metals to form hydrogen gas, e.g.:
$$Zn (s)+2HCl (aq)→ZnCl_2(aq)+H_2(g)$$
Bases (alkalis when dissolved in water) often:
- Have a pH greater than 7.
- Feel soapy/slippery if touched (but we do not test this on purpose!).
- Can be corrosive, especially strong bases like sodium hydroxide.
We use indicators – substances that change color depending on pH.
| Indicator | In acidic solution | In basic solution |
|---|
| Blue litmus | Turns red | Stays blue |
| Red litmus | Stays red | Turns blue |
| Methyl orange | Red | Yellow |
| Phenolphthalein | Colourless | Pink |
| Universal indicator | Red/orange/yellow (pH < 7) | Blue/purple (pH > 7) |
To understand their behaviour, we look at ions.
When acids dissolve in water, they ionise and produce hydrogen ions, $H^+(aq)$.
Hydrochloric acid:
$$\mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{H}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})$$
These H⁺ ions (often written as H₃O⁺ when bonded to water) are responsible for:
- Low pH
- Corrosive effects
- Reactions with metals and carbonates
Reaction with a carbonate:
$$2 \mathrm{HCl}(\mathrm{aq})+\mathrm{CaCO}_3(\mathrm{~s}) \rightarrow \mathrm{CaCl}_2(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})+\mathrm{CO}_2(\mathrm{~g})$$
You see fizzing because CO₂ gas is released.
When bases dissolve in water, they dissociate and produce hydroxide ions, OH⁻(aq).
Sodium hydroxide:
$$\mathrm{NaOH}(\mathrm{aq}) \rightarrow \mathrm{Na}^{+}(\mathrm{aq})+\mathrm{OH}^{-}(\mathrm{aq})$$
The OH⁻ ions give bases:
- High pH
- Slippery feel (like soap)
- Ability to neutralise acids
Reaction with ammonium salts (release of ammonia):
$$\mathrm{NH}_4^{+}(\mathrm{aq})+\mathrm{OH}^{-}(\mathrm{aq}) \rightarrow \mathrm{NH}_3(\mathrm{~g})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})$$
A sharp smell of ammonia is observed.
- When an acid and a base react, H⁺ from the acid and OH⁻ from the base combine to form water: $$\mathrm{H}^{+}(\mathrm{aq})+\mathrm{OH}^{-}(\mathrm{aq}) \rightarrow \mathrm{H}_2 \mathrm{O}(\mathrm{l})$$
- This is the key ionic idea behind neutralisation.
$$\mathrm{HCl}(\mathrm{aq})+\mathrm{NaOH}(\mathrm{aq}) \rightarrow \mathrm{NaCl}(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})$$
- Salt: NaCl
- Water: H₂O
- Solution becomes closer to pH 7.
Link to properties:
- The more H⁺ ions, the more acidic and corrosive the solution (lower pH).
- The more OH⁻ ions, the more basic and caustic the solution (higher pH).
- Indicators respond to these ions, changing colour.
Acids and bases are not just in the lab – they’re everywhere.
- Acids:
- Vinegar – contains ethanoic (acetic) acid, used in cooking and cleaning limescale.
- Citrus fruit juice (e.g. lemons, oranges) – contains citric acid; tastes sour.
- Carbonated drinks – contain carbonic acid (from dissolved CO₂) and often phosphoric acid.
- Bases:
- Baking soda (sodium hydrogencarbonate, NaHCO₃)
- Mild base used in baking (releases CO₂ when reacting with acids) and cleaning.
- Toothpaste
- Slightly basic to neutralise acids formed by bacteria in the mouth and protect tooth enamel.
- Soap and many cleaning products
- Often contain alkaline substances (like sodium hydroxide in soap-making) to remove grease and fats.
In bodily fluids (like blood), weak acids and bases act as buffers to keep pH within a narrow range so biochemical reactions can occur properly.
- Hydrochloric acid (HCl): Used to clean metals (remove rust and limescale) and in food processing.
- Sulfuric acid (H₂SO₄): Used in car batteries and in the manufacture of fertilisers.
- Sodium hydroxide (NaOH): Used in the manufacture of soap, paper, and in water treatment.
Strength vs Concentration of Acids and Bases - How could you safely determine whether an unknown solution is acidic or basic using indicators?
- Write the ionic equation for the neutralisation of nitric acid (HNO₃) with potassium hydroxide (KOH).
- Classify each as mostly acidic, basic, or neutral: lemon juice, blood, shampoo, drain cleaner, tap water.