MYP Integrated Sciences Topic Acids and Bases…

What Defines an Acid or a Base?

There are two very useful ways to define acids and bases.

Brønsted–Lowry (general):
1. An acid is a proton (H⁺) donor.
2. A base is a proton (H⁺) acceptor.
Arrhenius (in water):
1. An acid is a substance that produces hydrogen ions, H⁺(aq), in water.
2. A base is a substance that produces hydroxide ions, OH⁻(aq), in water.

Note

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

The pH scale is a logarithmic scale used to measure how acidic or basic a solution is.
1. pH < 7 → acidic
2. pH = 7 → neutral
3. pH > 7 → basic/alkaline
Each step of 1 pH unit represents a 10× change in hydrogen ion concentration $[H^+]$.
1. pH 3 is 10 times more acidic than pH 4.
2. pH 2 is 100 times more acidic than pH 4.

How Do We Recognise Acids and Bases in the Lab (Without Tasting)?

Tip

Very important: We never taste substances in the laboratory.
Many acids and bases are corrosive and dangerous.
Instead, we use observable properties and indicators.

Observable Properties of Acids

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.:

$$Zn (s)+2HCl (aq)→ZnCl_2(aq)+H_2(g)$$

Observable Properties of Bases

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.

Using Indicators to Identify Acids and Bases

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)

How Do Acids and Bases Behave in Water?

To understand their behaviour, we look at ions.

Acids in Water – Source of H⁺

When acids dissolve in water, they ionise and produce hydrogen ions, $H^+(aq)$.

Example

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

Example

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.

Bases in Water – Source of OH⁻

When bases dissolve in water, they dissociate and produce hydroxide ions, OH⁻(aq).

Example

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

Example

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.

Neutralisation – H⁺ Meets OH⁻

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.

Example

$$\mathrm{HCl}(\mathrm{aq})+\mathrm{NaOH}(\mathrm{aq}) \rightarrow \mathrm{NaCl}(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})$$

Salt: NaCl
Water: H₂O
Solution becomes closer to pH 7.

Hint

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 in Everyday Life

Acids and bases are not just in the lab – they’re everywhere.

Household Products

Acids:
1. Vinegar – contains ethanoic (acetic) acid, used in cooking and cleaning limescale.
2. Citrus fruit juice (e.g. lemons, oranges) – contains citric acid; tastes sour.
3. Carbonated drinks – contain carbonic acid (from dissolved CO₂) and often phosphoric acid.
Bases:
1. Baking soda (sodium hydrogencarbonate, NaHCO₃)
  1. Mild base used in baking (releases CO₂ when reacting with acids) and cleaning.
2. Toothpaste
  1. Slightly basic to neutralise acids formed by bacteria in the mouth and protect tooth enamel.
3. Soap and many cleaning products
  1. Often contain alkaline substances (like sodium hydroxide in soap-making) to remove grease and fats.

Note

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.

Importance in Industry

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.

Active recall

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.

How Do Acids and Bases React with Metals, Carbonates and Oxides?

Acids and bases show predictable patterns of reaction with metals, carbonates and oxides. Recognising these patterns helps you:

predict products,
identify unknown substances,
and understand real-world processes like acid rain damage or antacid action.

Acids React with Metals

When an acid reacts with a reactive metal, the products are salt and hydrogen gas.
General word equation: $$\text{Metal} + \text{Acid} \rightarrow \text{Salt} + \text{Hydrogen gas}$$
Only sufficiently reactive metals (e.g. magnesium, zinc, iron) react with dilute acids in this way.
Less reactive metals like copper do not react with dilute acids under normal conditions.

Example

Magnesium reacts with hydrochloric acid:

Word equation: $$\text{Magnesium} + \text{Hydrochloric acid} \rightarrow \text{Magnesium chloride} + \text{Hydrogen gas}$$
Balanced symbol equation: $$\text {Mg}(s) + 2\mathrm{HCl}(aq) \rightarrow \mathrm{MgCl}_2(aq) + \mathrm{H}_2(g)$$
Magnesium replaces hydrogen ions in hydrochloric acid, forming magnesium chloride and hydrogen gas.

Acids React with Carbonates and Hydrogencarbonates

Acids react with carbonates (CO₃²⁻) and hydrogencarbonates (HCO₃⁻) to produce carbon dioxide, salt and water.
The general equations are: $$\text{Acid} + \text{Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon dioxide gas}$$ $$\text{Acid} + \text{Hydrogencarbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon dioxide gas}$$

Example

Calcium carbonate reacts with hydrochloric acid:

Word equation: $$\text{Hydrochloric acid} + \text{Calcium carbonate} \rightarrow \text{Calcium chloride} +$$ $$+ \text{Water} + \text{Carbon dioxide gas}$$
Balanced symbol equation: $$\mathrm {2HCl}(aq) + \mathrm{CaCO_3}(s) \rightarrow \mathrm{CaCl_2}(aq) + \mathrm{H_2O}(l) + \mathrm{CO_2(g)}$$
Calcium carbonate reacts to form calcium chloride, water, and carbon dioxide gas.

Note

Carbon dioxide gas is a byproduct of these reactions.

Acids React with Oxides

Many metal oxides (e.g. CuO, MgO, CaO) are basic oxides.
They react with acids in neutralisation reactions:
The general equation is: $$\text{Acid} + \text{Metal oxide} \rightarrow \text{Salt} + \text{Water}$$

Example

Copper(II) oxide reacts with sulfuric acid:

Word equation: $$\text{Sulfuric acid} + \text{Copper(II) oxide} \rightarrow \text{Copper(II) sulfate} + \text{Water}$$
Balanced symbol equation : $$\mathrm{H}_2\mathrm{SO}_4(aq) + \mathrm{CuO}(s) \rightarrow \mathrm{CuSO}_4(aq) + \mathrm{H}{2}\mathrm{O}(l)$$
Copper(II) oxide neutralizes sulfuric acid, forming copper(II) sulfate and water.

Bases React with Acids and Non-Metal Oxides

Bases react with acids in a neutralization reaction, producing a salt and water.
The general equation is: $$\mathrm{Acid\ +\ Base\ \rightarrow\ Salt\ +\ Water}$$

Example

Sodium hydroxide reacts with hydrochloric acid:

Word equation: $$\text{Hydrochloric acid} + \text{Sodium hydroxide} \rightarrow \text{Sodium chloride} + \text{Water}$$
Balanced symbol equation: $$\mathrm{HCl(aq)} + \mathrm{NaOH(aq)} \rightarrow \mathrm{NaCl(aq)} + \mathrm{H_2O(l)}$$

Bases can also react with non-metal oxides (which are acidic) to produce a salt and water.

Example

Magnesium hydroxide reacts with carbon dioxide:

Word Equation: $$\text{Magnesium hydroxide} + \text{Carbon dioxide} \rightarrow \text{Magnesium carbonate} + \text{Water}$$
Balanced Symbol Equation: $$\mathrm{Mg(OH}_2)(aq) + \mathrm{CO_2}(g) \rightarrow \mathrm{MgCO_3}(aq) + \mathrm{H_2O}(l)$$

Tip

Pattern summary:

Metal + acid → salt + hydrogen
Carbonate + acid → salt + water + carbon dioxide
Acid + base/metal oxide → salt + water (neutralisation)

These patterns are very useful for predicting products in reaction questions.

How Do These Reactions Link to Proton Donors and Acceptors?

The Brønsted–Lowry theory defines:

Acid = proton donor (donates H⁺)
Base = proton acceptor (accepts H⁺)

We can use this idea to explain what is happening in the reactions above at the particle level.

Metal + Acid

Take the reaction of magnesium with hydrochloric acid:

$$\mathrm{Mg}(\mathrm{~s})+2 \mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{Mg}^{2+}(\mathrm{aq})+2 \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{H}_2(\mathrm{~g})$$

Step-by-step:

HCl in water produces H⁺(aq) and Cl⁻(aq).
The acid donates H⁺ ions (protons).
Magnesium atoms lose electrons (they are oxidised) and become Mg²⁺.
The lost electrons reduce H⁺ to H₂ gas.

So:

HCl acts as an acid (proton donor).
The metal does not act as a base in the Brønsted–Lowry sense, but it provides electrons that reduce the protons.

Carbonate + Acid

Consider calcium carbonate reacting with hydrochloric acid:

$$\mathrm{CaCO}_3(\mathrm{~s})+2 \mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{Ca}^{2+}(\mathrm{aq})+2 \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})+\mathrm{CO}_2(\mathrm{~g})$$

HCl donates H⁺ ions (acts as a proton donor).
The carbonate ion (CO₃²⁻) acts as a base, accepting protons to eventually form carbonic acid, H₂CO₃, which breaks down into CO₂ + H₂O.

In simplified proton terms:

$$\mathrm{CO}_3^{2-}+2 \mathrm{H}^{+} \rightarrow \mathrm{H}_2 \mathrm{O}+\mathrm{CO}_2$$

So:

The acid donates H⁺.
The carbonate ion acts as a Brønsted–Lowry base, accepting H⁺.

Base (Hydroxide) + Acid

Take sodium hydroxide and hydrochloric acid:

$$\mathrm{NaOH}(\mathrm{aq})+\mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{NaCl}(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})$$

In ionic form:

$$\mathrm{Na}^{+}+\mathrm{OH}^{-}+\mathrm{H}^{+}+\mathrm{Cl}^{-} \rightarrow \mathrm{Na}^{+}+\mathrm{Cl}^{-}+\mathrm{H}_2 \mathrm{O}$$

Net ionic equation:

$$\mathrm{H}^{+} \text {(from the acid) }+\mathrm{OH}^{-} \text {(from the base) } \rightarrow \mathrm{H}_2 \mathrm{O}$$

The acid (HCl) acts as a proton donor.
The base (OH⁻ from NaOH) acts as a proton acceptor, forming water.

How Does Understanding Acid/Base Characteristics Help with Safety?

Many acids and bases are corrosive and can cause burns or damage if mishandled.
Understanding their properties helps us use them safely in labs and industry.

Dangers of Acids and Bases

Strong acids (e.g. HCl, H₂SO₄, HNO₃) can cause severe burns and damage metals, skin, eyes.
Strong bases (e.g. NaOH, KOH) can be just as dangerous, even if they feel “soapy”.
Even weak acids and bases can be harmful in high concentration or with long contact.
Some reactions release toxic gases (e.g. chlorine gas when strong oxidising acids react with chlorides; CO₂ in enclosed spaces).

Safe Handling Practices in the Lab

Personal Protective Equipment (PPE)
1. Wear safety goggles to protect your eyes.
2. Wear lab coats to protect skin and clothing.
3. Use gloves when handling corrosive solutions or cleaning up spills.
Working Safely
1. Always add acid to water, not water to acid, to avoid violent splashing.
2. Carry bottles with two hands and keep containers closed when not in use.
3. Use a fume hood when a reaction may release toxic or irritating gases.
4. Clearly label all containers with the name and hazard information.

Tip

Treat all acids and bases with respect, even if they are “weak” or “dilute”.

Industrial Uses and Safety

Acids and bases are vital in many industries, but they must be controlled carefully.

Example

Fertiliser production:
- Sulfuric acid is used to make fertilisers such as ammonium sulfate.
- Spills or leaks must be neutralised and contained to protect workers and the environment.
Wastewater treatment:
- Bases like calcium hydroxide are used to neutralise acidic waste before discharge.
- Careful pH control avoids over-neutralisation, which would make the water too basic.
Cleaning products:
- Acidic cleaners remove limescale (e.g. dilute HCl).
- Alkaline drain cleaners often contain NaOH to dissolve fats.
- Both can burn skin and eyes; correct PPE and storage are essential.
Overuse in agriculture
- Adding too much acid or base to soil can push pH too low or too high, harming crops.
- Understanding acid–base reactions helps farmers apply appropriate amounts of soil treatments like lime (CaCO₃) or ammonium-based fertilisers.

What Defines an Acid or a Base?

There are two very useful ways to define acids and bases.

Brønsted–Lowry (general):
1. An acid is a proton (H⁺) donor.
2. A base is a proton (H⁺) acceptor.
Arrhenius (in water):
1. An acid is a substance that produces hydrogen ions, H⁺(aq), in water.
2. A base is a substance that produces hydroxide ions, OH⁻(aq), in water.

Note

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

The pH scale is a logarithmic scale used to measure how acidic or basic a solution is.
1. pH < 7 → acidic
2. pH = 7 → neutral
3. pH > 7 → basic/alkaline
Each step of 1 pH unit represents a 10× change in hydrogen ion concentration $[H^+]$.
1. pH 3 is 10 times more acidic than pH 4.
2. pH 2 is 100 times more acidic than pH 4.

How Do We Recognise Acids and Bases in the Lab (Without Tasting)?

Tip

Very important: We never taste substances in the laboratory.
Many acids and bases are corrosive and dangerous.
Instead, we use observable properties and indicators.

Observable Properties of Acids

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.:

$$Zn (s)+2HCl (aq)→ZnCl_2(aq)+H_2(g)$$

Observable Properties of Bases

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.

Using Indicators to Identify Acids and Bases

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)

How Do Acids and Bases Behave in Water?

To understand their behaviour, we look at ions.

Acids in Water – Source of H⁺

When acids dissolve in water, they ionise and produce hydrogen ions, $H^+(aq)$.

Example

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

Example

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.

Bases in Water – Source of OH⁻

When bases dissolve in water, they dissociate and produce hydroxide ions, OH⁻(aq).

Example

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

Example

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.

Neutralisation – H⁺ Meets OH⁻

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.

Example

$$\mathrm{HCl}(\mathrm{aq})+\mathrm{NaOH}(\mathrm{aq}) \rightarrow \mathrm{NaCl}(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})$$

Salt: NaCl
Water: H₂O
Solution becomes closer to pH 7.

Hint

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 in Everyday Life

Acids and bases are not just in the lab – they’re everywhere.

Household Products

Acids:
1. Vinegar – contains ethanoic (acetic) acid, used in cooking and cleaning limescale.
2. Citrus fruit juice (e.g. lemons, oranges) – contains citric acid; tastes sour.
3. Carbonated drinks – contain carbonic acid (from dissolved CO₂) and often phosphoric acid.
Bases:
1. Baking soda (sodium hydrogencarbonate, NaHCO₃)
  1. Mild base used in baking (releases CO₂ when reacting with acids) and cleaning.
2. Toothpaste
  1. Slightly basic to neutralise acids formed by bacteria in the mouth and protect tooth enamel.
3. Soap and many cleaning products
  1. Often contain alkaline substances (like sodium hydroxide in soap-making) to remove grease and fats.

Note

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.

Importance in Industry

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.

Active recall

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.

How Do Acids and Bases React with Metals, Carbonates and Oxides?

Acids and bases show predictable patterns of reaction with metals, carbonates and oxides. Recognising these patterns helps you:

predict products,
identify unknown substances,
and understand real-world processes like acid rain damage or antacid action.

Acids React with Metals

When an acid reacts with a reactive metal, the products are salt and hydrogen gas.
General word equation: $$\text{Metal} + \text{Acid} \rightarrow \text{Salt} + \text{Hydrogen gas}$$
Only sufficiently reactive metals (e.g. magnesium, zinc, iron) react with dilute acids in this way.
Less reactive metals like copper do not react with dilute acids under normal conditions.

Example

Magnesium reacts with hydrochloric acid:

Word equation: $$\text{Magnesium} + \text{Hydrochloric acid} \rightarrow \text{Magnesium chloride} + \text{Hydrogen gas}$$
Balanced symbol equation: $$\text {Mg}(s) + 2\mathrm{HCl}(aq) \rightarrow \mathrm{MgCl}_2(aq) + \mathrm{H}_2(g)$$
Magnesium replaces hydrogen ions in hydrochloric acid, forming magnesium chloride and hydrogen gas.

Acids React with Carbonates and Hydrogencarbonates

Acids react with carbonates (CO₃²⁻) and hydrogencarbonates (HCO₃⁻) to produce carbon dioxide, salt and water.
The general equations are: $$\text{Acid} + \text{Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon dioxide gas}$$ $$\text{Acid} + \text{Hydrogencarbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon dioxide gas}$$

Example

Calcium carbonate reacts with hydrochloric acid:

Word equation: $$\text{Hydrochloric acid} + \text{Calcium carbonate} \rightarrow \text{Calcium chloride} +$$ $$+ \text{Water} + \text{Carbon dioxide gas}$$
Balanced symbol equation: $$\mathrm {2HCl}(aq) + \mathrm{CaCO_3}(s) \rightarrow \mathrm{CaCl_2}(aq) + \mathrm{H_2O}(l) + \mathrm{CO_2(g)}$$
Calcium carbonate reacts to form calcium chloride, water, and carbon dioxide gas.

Note

Carbon dioxide gas is a byproduct of these reactions.

Acids React with Oxides

Many metal oxides (e.g. CuO, MgO, CaO) are basic oxides.
They react with acids in neutralisation reactions:
The general equation is: $$\text{Acid} + \text{Metal oxide} \rightarrow \text{Salt} + \text{Water}$$

Example

Copper(II) oxide reacts with sulfuric acid:

Word equation: $$\text{Sulfuric acid} + \text{Copper(II) oxide} \rightarrow \text{Copper(II) sulfate} + \text{Water}$$
Balanced symbol equation : $$\mathrm{H}_2\mathrm{SO}_4(aq) + \mathrm{CuO}(s) \rightarrow \mathrm{CuSO}_4(aq) + \mathrm{H}{2}\mathrm{O}(l)$$
Copper(II) oxide neutralizes sulfuric acid, forming copper(II) sulfate and water.

Bases React with Acids and Non-Metal Oxides

Bases react with acids in a neutralization reaction, producing a salt and water.
The general equation is: $$\mathrm{Acid\ +\ Base\ \rightarrow\ Salt\ +\ Water}$$

Example

Sodium hydroxide reacts with hydrochloric acid:

Word equation: $$\text{Hydrochloric acid} + \text{Sodium hydroxide} \rightarrow \text{Sodium chloride} + \text{Water}$$
Balanced symbol equation: $$\mathrm{HCl(aq)} + \mathrm{NaOH(aq)} \rightarrow \mathrm{NaCl(aq)} + \mathrm{H_2O(l)}$$

Bases can also react with non-metal oxides (which are acidic) to produce a salt and water.

Example

Magnesium hydroxide reacts with carbon dioxide:

Word Equation: $$\text{Magnesium hydroxide} + \text{Carbon dioxide} \rightarrow \text{Magnesium carbonate} + \text{Water}$$
Balanced Symbol Equation: $$\mathrm{Mg(OH}_2)(aq) + \mathrm{CO_2}(g) \rightarrow \mathrm{MgCO_3}(aq) + \mathrm{H_2O}(l)$$

Tip

Pattern summary:

Metal + acid → salt + hydrogen
Carbonate + acid → salt + water + carbon dioxide
Acid + base/metal oxide → salt + water (neutralisation)

These patterns are very useful for predicting products in reaction questions.

How Do These Reactions Link to Proton Donors and Acceptors?

The Brønsted–Lowry theory defines:

Acid = proton donor (donates H⁺)
Base = proton acceptor (accepts H⁺)

We can use this idea to explain what is happening in the reactions above at the particle level.

Metal + Acid

Take the reaction of magnesium with hydrochloric acid:

$$\mathrm{Mg}(\mathrm{~s})+2 \mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{Mg}^{2+}(\mathrm{aq})+2 \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{H}_2(\mathrm{~g})$$

Step-by-step:

HCl in water produces H⁺(aq) and Cl⁻(aq).
The acid donates H⁺ ions (protons).
Magnesium atoms lose electrons (they are oxidised) and become Mg²⁺.
The lost electrons reduce H⁺ to H₂ gas.

So:

HCl acts as an acid (proton donor).
The metal does not act as a base in the Brønsted–Lowry sense, but it provides electrons that reduce the protons.

Carbonate + Acid

Consider calcium carbonate reacting with hydrochloric acid:

$$\mathrm{CaCO}_3(\mathrm{~s})+2 \mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{Ca}^{2+}(\mathrm{aq})+2 \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})+\mathrm{CO}_2(\mathrm{~g})$$

HCl donates H⁺ ions (acts as a proton donor).
The carbonate ion (CO₃²⁻) acts as a base, accepting protons to eventually form carbonic acid, H₂CO₃, which breaks down into CO₂ + H₂O.

In simplified proton terms:

$$\mathrm{CO}_3^{2-}+2 \mathrm{H}^{+} \rightarrow \mathrm{H}_2 \mathrm{O}+\mathrm{CO}_2$$

So:

The acid donates H⁺.
The carbonate ion acts as a Brønsted–Lowry base, accepting H⁺.

Base (Hydroxide) + Acid

Take sodium hydroxide and hydrochloric acid:

$$\mathrm{NaOH}(\mathrm{aq})+\mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{NaCl}(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})$$

In ionic form:

$$\mathrm{Na}^{+}+\mathrm{OH}^{-}+\mathrm{H}^{+}+\mathrm{Cl}^{-} \rightarrow \mathrm{Na}^{+}+\mathrm{Cl}^{-}+\mathrm{H}_2 \mathrm{O}$$

Net ionic equation:

$$\mathrm{H}^{+} \text {(from the acid) }+\mathrm{OH}^{-} \text {(from the base) } \rightarrow \mathrm{H}_2 \mathrm{O}$$

The acid (HCl) acts as a proton donor.
The base (OH⁻ from NaOH) acts as a proton acceptor, forming water.

How Does Understanding Acid/Base Characteristics Help with Safety?

Many acids and bases are corrosive and can cause burns or damage if mishandled.
Understanding their properties helps us use them safely in labs and industry.

Dangers of Acids and Bases

Strong acids (e.g. HCl, H₂SO₄, HNO₃) can cause severe burns and damage metals, skin, eyes.
Strong bases (e.g. NaOH, KOH) can be just as dangerous, even if they feel “soapy”.
Even weak acids and bases can be harmful in high concentration or with long contact.
Some reactions release toxic gases (e.g. chlorine gas when strong oxidising acids react with chlorides; CO₂ in enclosed spaces).

Safe Handling Practices in the Lab

Personal Protective Equipment (PPE)
1. Wear safety goggles to protect your eyes.
2. Wear lab coats to protect skin and clothing.
3. Use gloves when handling corrosive solutions or cleaning up spills.
Working Safely
1. Always add acid to water, not water to acid, to avoid violent splashing.
2. Carry bottles with two hands and keep containers closed when not in use.
3. Use a fume hood when a reaction may release toxic or irritating gases.
4. Clearly label all containers with the name and hazard information.

Tip

Treat all acids and bases with respect, even if they are “weak” or “dilute”.

Industrial Uses and Safety

Acids and bases are vital in many industries, but they must be controlled carefully.

Example

Fertiliser production:
- Sulfuric acid is used to make fertilisers such as ammonium sulfate.
- Spills or leaks must be neutralised and contained to protect workers and the environment.
Wastewater treatment:
- Bases like calcium hydroxide are used to neutralise acidic waste before discharge.
- Careful pH control avoids over-neutralisation, which would make the water too basic.
Cleaning products:
- Acidic cleaners remove limescale (e.g. dilute HCl).
- Alkaline drain cleaners often contain NaOH to dissolve fats.
- Both can burn skin and eyes; correct PPE and storage are essential.
Overuse in agriculture
- Adding too much acid or base to soil can push pH too low or too high, harming crops.
- Understanding acid–base reactions helps farmers apply appropriate amounts of soil treatments like lime (CaCO₃) or ammonium-based fertilisers.

MYP Integrated Sciences Acids and Bases Notes

Science skills

Scale and structure

Classification and organisation

Materials and bonding

Chemical reactions

Energy

Forces and motion

Thermal physics

Electricity and magnetism

Waves

Atoms and space

Genetics and evolution

Body systems

Ecosystems and environment

Human impact and sustainability

Biotechnology and the future

MYP Integrated Sciences Acids and Bases Notes

Science skills

Scale and structure

Classification and organisation

Materials and bonding

Chemical reactions

Energy

Forces and motion

Thermal physics

Electricity and magnetism

Waves

Atoms and space

Genetics and evolution

Body systems

Ecosystems and environment

Human impact and sustainability

Biotechnology and the future

What Defines an Acid or a Base?

The pH Scale

How Do We Recognise Acids and Bases in the Lab (Without Tasting)?

Observable Properties of Acids

Observable Properties of Bases

Using Indicators to Identify Acids and Bases

How Do Acids and Bases Behave in Water?

Acids in Water – Source of H⁺

Bases in Water – Source of OH⁻

Neutralisation – H⁺ Meets OH⁻

Acids and Bases in Everyday Life

Household Products

Importance in Industry

How Do Acids and Bases React with Metals, Carbonates and Oxides?

Acids React with Metals

Acids React with Carbonates and Hydrogencarbonates

Acids React with Oxides

Bases React with Acids and Non-Metal Oxides

How Do These Reactions Link to Proton Donors and Acceptors?

Metal + Acid

Carbonate + Acid

Base (Hydroxide) + Acid

How Does Understanding Acid/Base Characteristics Help with Safety?

Dangers of Acids and Bases

Safe Handling Practices in the Lab

Industrial Uses and Safety

What Defines an Acid or a Base?

The pH Scale

How Do We Recognise Acids and Bases in the Lab (Without Tasting)?

Observable Properties of Acids

Observable Properties of Bases

Using Indicators to Identify Acids and Bases

How Do Acids and Bases Behave in Water?

Acids in Water – Source of H⁺

Bases in Water – Source of OH⁻

Neutralisation – H⁺ Meets OH⁻

Acids and Bases in Everyday Life

Household Products

Importance in Industry

How Do Acids and Bases React with Metals, Carbonates and Oxides?

Acids React with Metals

Acids React with Carbonates and Hydrogencarbonates

Acids React with Oxides

Bases React with Acids and Non-Metal Oxides

How Do These Reactions Link to Proton Donors and Acceptors?

Metal + Acid