Acids, bases and salts

🏷 Chapter C7: Acids, Bases and Salts 🔬

Hello future chemists! This chapter is all about one of the most fundamental groups of chemicals: acids, bases, and the compounds they create — salts. Don't worry, these aren't just abstract ideas; they are everywhere! From the vinegar you use in the kitchen to the cleaners under the sink, understanding this topic is key to explaining many everyday chemical reactions. Let's dive in!

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1. Defining Acids and Their Properties (Core & Supplement)

What is an Acid?

Acids are substances that produce hydrogen ions, \(\text{H}^+\), when dissolved in water. In terms of proton transfer, an acid is defined as a proton (\(\text{H}^+\)) donor. The presence of these \(\text{H}^+\) ions is what gives acids their characteristic properties.

Examples: Hydrochloric acid (\(\text{HCl}\)), Sulfuric acid (\(\text{H}_2\text{SO}_4\)), Citric acid (in lemons).

Characteristic Properties of Acids

• Taste: Sour (Think of lemons! Note: Never taste chemicals in a lab.)
• Corrosive: They can damage skin, tissue, and materials.
• Effect on Indicators: Acids change the colour of certain substances called indicators.

Quick Indicator Check (Core C7.1(2)):

• Litmus: Changes blue litmus \(\rightarrow\) red.
• Methyl Orange: Changes yellow/orange \(\rightarrow\) red.
• Thymolphthalein: Remains colourless in acid.

💡 Memory Aid: Acids turn litmus Red. (A Really Acidy Robot)

Reactions of Acids (Core C7.1(1))

A. Reaction with Metals (Only Reactive Metals)

Acids react with metals (like Magnesium, Zinc, Iron) to produce a salt and hydrogen gas.

Word Equation:
Acid + Metal \(\rightarrow\) Salt + Hydrogen

Example: Zinc metal reacting with Sulfuric acid:
\(\text{H}_2\text{SO}_4 \text{(aq)} + \text{Zn} \text{(s)} \rightarrow \text{ZnSO}_4 \text{(aq)} + \text{H}_2 \text{(g)}\)

🚨 Safety Tip: Hydrogen gas is flammable! We test for it using a lighted splint, which produces a characteristic squeaky pop sound.

B. Reaction with Bases (Neutralisation)

This is the classic neutralisation reaction, producing a salt and water.

Word Equation:
Acid + Base \(\rightarrow\) Salt + Water

Example: Hydrochloric acid reacting with Sodium hydroxide (a soluble base/alkali):
\(\text{HCl} \text{(aq)} + \text{NaOH} \text{(aq)} \rightarrow \text{NaCl} \text{(aq)} + \text{H}_2\text{O} \text{(l)}\)

The general ionic equation for neutralisation between an acid and an alkali is:
\(\text{H}^+ \text{(aq)} + \text{OH}^- \text{(aq)} \rightarrow \text{H}_2\text{O} \text{(l)}\)

C. Reaction with Carbonates

Acids react with carbonates (e.g., Calcium carbonate, \(\text{CaCO}_3\)) to produce three products: salt, water, and carbon dioxide gas.

Word Equation:
Acid + Carbonate \(\rightarrow\) Salt + Water + Carbon Dioxide

Example: Sulfuric acid reacting with Copper carbonate:
\(\text{H}_2\text{SO}_4 \text{(aq)} + \text{CuCO}_3 \text{(s)} \rightarrow \text{CuSO}_4 \text{(aq)} + \text{H}_2\text{O} \text{(l)} + \text{CO}_2 \text{(g)}\)

🚨 Why is this reaction important? This is the standard laboratory test for identifying carbonates (or for confirming the presence of Carbon Dioxide gas, which turns limewater milky).

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2. Bases, Alkalis, and Their Properties (Core & Supplement)

What is a Base? (Core C7.1(3))

A base is defined as an oxide or hydroxide of a metal. In terms of proton transfer, a base is defined as a proton (\(\text{H}^+\)) acceptor.

Examples: Copper oxide (\(\text{CuO}\)), Calcium oxide (\(\text{CaO}\)), Sodium hydroxide (\(\text{NaOH}\)).

What is an Alkali? (Core C7.1(3))

An alkali is a soluble base — a base that dissolves in water. When dissolved, alkalis produce hydroxide ions, \(\text{OH}^-\), in the aqueous solution.

All alkalis are bases, but not all bases are alkalis (since many bases, like \(\text{CuO}\), are insoluble).

Examples: Sodium hydroxide (\(\text{NaOH}\)), Potassium hydroxide (\(\text{KOH}\)).

Characteristic Properties of Bases and Alkalis

• Taste: Bitter (Think of soap, but don't taste it!)
• Feel: Soapy or slimy to the touch.
• Corrosive: Strong alkalis (like drain cleaner) can be just as corrosive as strong acids.

Effect on Indicators (Core C7.1(5))

Alkalis change the colour of indicators in the opposite way to acids:

• Litmus paper: Changes red litmus \(\rightarrow\) blue.
• Methyl Orange: Changes red/orange \(\rightarrow\) yellow.
• Thymolphthalein: Turns blue in alkali.

💡 Memory Aid: Bases turn litmus Blue. (Big Blue Bases)

Reactions of Bases and Alkalis (Core C7.1(4))

The most important reaction is neutralisation:

Word Equation:
Base (or Alkali) + Acid \(\rightarrow\) Salt + Water

Example: Copper oxide (a base) reacting with Hydrochloric acid:
\(\text{CuO} \text{(s)} + 2\text{HCl} \text{(aq)} \rightarrow \text{CuCl}_2 \text{(aq)} + \text{H}_2\text{O} \text{(l)}\)

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3. pH, Acidity, Alkalinity, and Neutrality (Core)

The pH scale is a measurement system used to determine how acidic or alkaline (basic) a solution is. The scale runs from 0 (very acidic) to 14 (very alkaline).

The pH Scale (Core C7.1(6))

• pH 7: Neutral (Neither acidic nor alkaline. Pure water is neutral.)
• pH < 7: Acidic (Lower number = higher acidity. pH 1 is strongly acidic.)
• pH > 7: Alkaline (Higher number = higher alkalinity. pH 14 is strongly alkaline.)

Using Universal Indicator

To compare the relative acidity and alkalinity of solutions, we use Universal Indicator. This is a mixture of several dyes that changes colour gradually across the entire pH range, allowing you to estimate the approximate pH value.

Approximate colours: Red/Orange (Strong Acid) \(\rightarrow\) Yellow (Weak Acid) \(\rightarrow\) Green (Neutral) \(\rightarrow\) Blue (Weak Alkali) \(\rightarrow\) Purple/Violet (Strong Alkali).

Neutralisation in Detail (Core C7.1(7))

Neutralisation is simply the reaction between the \(\text{H}^+\) ions from the acid and the \(\text{OH}^-\) ions from the alkali to form neutral water:

\(\text{H}^+ \text{(aq)} + \text{OH}^- \text{(aq)} \rightarrow \text{H}_2\text{O} \text{(l)}\)

Result: The product solution has a pH closer to 7 (neutral).

🧪 Useful applications:

• Treating acidic soil with lime (\(\text{CaO}\) or \(\text{Ca}(\text{OH})_2\)) to raise the pH.
• Using antacids (bases) to neutralise excess stomach acid.
• Applying a weak alkali like baking soda to neutralise acidic stings.

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4. Classifying Oxides (Core & Supplement)

Oxides are compounds containing oxygen combined with another element. We classify them based on how they react with acids and bases.

A. Acidic Oxides (Core C7.2(1))

• Nature: Oxides of non-metals.
• Property: They react with bases (alkalis) and dissolve in water to form acidic solutions.
• Examples: Sulfur dioxide (\(\text{SO}_2\)), Carbon dioxide (\(\text{CO}_2\)).

B. Basic Oxides (Core C7.2(1))

• Nature: Oxides of metals.
• Property: They react with acids to form a salt and water (acting as a base).
• Examples: Copper(II) oxide (\(\text{CuO}\)), Calcium oxide (\(\text{CaO}\)).

C. Amphoteric Oxides (Supplement C7.2(2) & C7.2(3))

Don't worry if this sounds complicated — "ampho" just means both.

• Definition: Amphoteric oxides are oxides that react with both acids and bases to produce a salt and water.
• Function: When reacting with an acid, they behave as a base. When reacting with a base, they behave as an acid.
• Examples: Aluminium oxide (\(\text{Al}_2\text{O}_3\)) and Zinc oxide (\(\text{ZnO}\)).

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5. Preparing and Naming Salts

What is a Salt?

A salt is an ionic compound formed when the hydrogen ion (\(\text{H}^+\)) in an acid is replaced by a metal ion or an ammonium ion (\(\text{NH}_4^+\)).

The name of the acid determines the second part of the salt name:

• Hydrochloric acid (\(\text{HCl}\)) makes Chlorides (e.g., \(\text{NaCl}\)).
• Sulfuric acid (\(\text{H}_2\text{SO}_4\)) makes Sulfates (e.g., \(\text{CuSO}_4\)).
• Nitric acid (\(\text{HNO}_3\)) makes Nitrates (e.g., \(\text{KNO}_3\)).

Solubility Rules of Common Salts

To choose the right method for preparing a salt, you need to know if it is soluble or insoluble:

• All nitrates are soluble.
• All sodium, potassium, and ammonium salts are soluble.
• Chlorides are soluble (except silver and lead chloride).
• Sulfates are soluble (except barium, lead, and calcium sulfate).
• Carbonates and hydroxides are mostly insoluble (except sodium, potassium, and ammonium).

Hydrated vs. Anhydrous (Core C7.3(2))

Some salts trap water molecules within their crystal structure.

• Hydrated substance: A substance chemically combined with water of crystallisation (e.g., copper(II) sulfate pentahydrate, \(\text{CuSO}_4 \cdot 5\text{H}_2\text{O}\)).
• Anhydrous substance: A substance containing no water of crystallisation (e.g., anhydrous \(\text{CuSO}_4\)).

6. Preparation of Soluble Salts (Core C7.3(1))

We use different methods to prepare soluble salts, depending on whether the starting materials are soluble or insoluble.

Method 1: Acid + Alkali (Titration)

This method is used when both the acid and the base/alkali are soluble (e.g., making Sodium Chloride, \(\text{NaCl}\), from \(\text{HCl}\) and \(\text{NaOH}\)).

Step-by-step:

1. Titration: Use an indicator (like methyl orange) to find the exact volume of acid needed to neutralise a known volume of alkali.
2. Repeat Without Indicator: Mix the exact volumes of acid and alkali determined in Step 1, but without the indicator (to ensure the final salt product is pure).
3. Heating and Crystallisation: Heat the resulting salt solution to evaporate some water until the crystallisation point is reached.
4. Cooling: Allow the saturated solution to cool slowly so salt crystals form.
5. Purification: Filter, wash the crystals with a little cold distilled water, and dry between filter papers.

Method 2: Acid + Excess Insoluble Reactant

This method is used when the reactant is insoluble (excess metal, insoluble base, or insoluble carbonate). The excess solid ensures all the acid reacts and can be easily filtered off.

Step-by-step:

1. Reaction: Add the insoluble reactant (e.g., excess \(\text{CuO}\)) to warm acid (e.g., \(\text{H}_2\text{SO}_4\)). Stir until no more reactant dissolves.
2. Filtration: Filter the mixture to remove the unreacted excess solid (residue). The filtrate is the pure salt solution.
3. Heating and Crystallisation: Heat the filtrate until saturated, then allow it to cool slowly to form crystals.
4. Purification: Filter, wash with a small amount of cold distilled water, and dry.

7. Preparation of Insoluble Salts (Supplement C7.3(3))

Insoluble salts are prepared by precipitation (mixing two different soluble salt solutions together).

Step-by-step:

1. Mixing: Mix two aqueous solutions of soluble salts (e.g., mix aqueous lead(II) nitrate with aqueous sodium sulfate to prepare insoluble lead(II) sulfate, \(\text{PbSO}_4\)).
2. Precipitation: An insoluble solid precipitate forms immediately.
3. Filtration: Filter the mixture to collect the precipitate on the filter paper.
4. Washing and Drying: Wash the precipitate thoroughly with distilled water to remove remaining soluble ions, and dry in a warm oven or between filter papers.