Sciences - Co-ordinated (Double) (0654) | 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)

What is an Acid?

Acids are substances that produce hydrogen ions, $\text{H}^+$, when dissolved in water (in reality, they form $\text{H}_3\text{O}^+$, hydronium ions, but for IGCSE, thinking of them as $\text{H}^+$ ions is usually enough). 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 paper: Changes blue litmus $\rightarrow$ red.
• Methyl Orange: Changes yellow/orange $\rightarrow$ red.

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

(Note: The ionic equation for this reaction, $\text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O}$, is not required in this syllabus.)

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)

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

A base is defined as an oxide or hydroxide of a metal. Bases are chemical opposites of acids and are proton ($\text{H}^+$) acceptors.

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.

💡 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 = stronger acid. pH 1 is very strong.)
• pH > 7: Alkaline (Higher number = stronger alkali. pH 14 is very strong.)

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 exact 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 weak alkali like baking soda to neutralise bee stings (which are acidic).

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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 acids.
• 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 ($\text{ZnO}$).

Did you know? Aluminium oxide is amphoteric, which helps explain why aluminium metal resists corrosion — the protective oxide layer is chemically tough and can withstand a wider range of conditions than a simple basic oxide would.

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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., $\text{KNO}_3$).

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

Some salts trap water molecules within their crystal structure.

• Hydrated substance: A substance chemically combined with water molecules (often shown as dots, e.g., copper(II) sulfate pentahydrate, $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$). This is usually coloured (e.g., blue).
• Anhydrous substance: A substance containing no water of crystallisation. This is usually white or colourless (e.g., $\text{CuSO}_4$).

When you heat a hydrated salt, the water is driven off and the salt becomes anhydrous, usually causing a colour change. This process is reversible if you add water back.

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 $\text{NaOH}$).

Because both products ($\text{NaCl}$ and $\text{H}_2\text{O}$) are soluble, we cannot use filtration to remove excess reactants. We must measure the reactants perfectly using titration.

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 (the end-point).
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 most of the water, forming a saturated solution.
4. Cooling: Allow the saturated solution to cool slowly. The salt crystals will form (crystallisation).
5. Purification: Filter, wash the crystals with a little cold distilled water, and dry.

Method 2: Acid + Excess Insoluble Reactant

This is the easier method, used when the reactant (metal, insoluble base, or insoluble carbonate) is insoluble. This allows us to remove the excess reactant using simple filtration.

This method is used for making salts of metals like copper, zinc, magnesium, and iron (using their oxides, hydroxides, or carbonates).

Step-by-step:

1. Reaction: Add the insoluble reactant (e.g., excess $\text{CuO}$) to the hot acid (e.g., $\text{H}_2\text{SO}_4$). Stir until no more reactant dissolves (the excess ensures all acid is used up).
2. Filtration: Filter the mixture. The unreacted solid excess material is the residue, and the soluble salt solution (the filtrate) passes through.
3. Heating and Crystallisation: Heat the filtrate until saturated, then allow it to cool slowly to form crystals.
4. Purification: Filter, wash, and dry the crystals.

These insoluble reactants can be:

• Excess Metal (Core C7.3(1)(b))
• Excess Insoluble Base (Core C7.3(1)(c))
• Excess Insoluble Carbonate (Core C7.3(1)(d))

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

If you need to make a salt that does not dissolve in water (an insoluble salt), you cannot use the methods above, as heating would simply yield the powdered insoluble salt.

Insoluble salts are prepared by precipitation (also called double decomposition).

Step-by-step:

1. Mixing: Mix two different soluble salts solutions together.

   Example: To make insoluble Lead(II) sulfate ($\text{PbSO}_4$), mix soluble Lead(II) nitrate solution with soluble Sodium sulfate solution.

2. Precipitation: When mixed, the ions swap partners, and the desired insoluble salt forms immediately as a solid (a precipitate).
3. Filtration: Filter the mixture to separate the solid insoluble salt from the liquid.
4. Washing and Drying: Wash the precipitate with distilled water to remove any traces of the unwanted soluble salts, and then dry it.

You've made it through! Acids, bases, and salts might seem complicated due to the different reaction types, but if you memorize the three main reactions of acids and the two key methods for preparing soluble salts, you will be well prepared! Keep practicing those equations!