Acid and base characteristics Notes

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:

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

Acids React with Metals

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

Acids React with Carbonates and Hydrogencarbonates

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

Acids React with Oxides

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

Bases React with Acids and Non-Metal Oxides

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

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

How Do These Reactions Link to Proton Donors and Acceptors?

The Brønsted–Lowry theory defines:

1. Acid = proton donor (donates H⁺)
2. 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:

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

So:

1. HCl acts as an acid (proton donor).
2. 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:

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

1. HCl donates H⁺ ions (acts as a proton donor).
2. 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:

1. The acid donates H⁺.
2. 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:

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

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

How Does Understanding Acid/Base Characteristics Help with Safety?

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

Dangers of Acids and Bases

1. Strong acids (e.g. HCl, H₂SO₄, HNO₃) can cause severe burns and damage metals, skin, eyes.
2. Strong bases (e.g. NaOH, KOH) can be just as dangerous, even if they feel “soapy”.
3. Even weak acids and bases can be harmful in high concentration or with long contact.
4. 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

1. 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.
2. 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.

Industrial Uses and Safety

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