Chemistry Study Notes: Acids, Bases, and Salt Preparations

Welcome to one of the most fundamental and important chapters in Inorganic Chemistry! Acids, bases, and salts are everywhere—from the sourness of a lemon to the soap you use to wash your hands. Understanding how these chemicals interact is essential for almost every concept we cover later.

Don't worry if this seems tricky at first; we will break down the definitions and practical preparations into simple, easy-to-follow steps! Let's get started!

Section I: Defining Acids, Bases, and Alkalis

What are Acids?

The simplest way to define an acid (according to the IGCSE curriculum) is based on what it does when dissolved in water:

  • Acids are substances that produce hydrogen ions (\(H^+\)) when they dissolve in water.
  • It is the concentration of \(H^+\) ions that determines how strong or acidic a solution is.

Key Examples of Strong Acids:

  • Hydrochloric Acid (HCl)
  • Sulfuric Acid (\(H_2SO_4\))
  • Nitric Acid (\(HNO_3\))

Did you know? The acid in your stomach (gastric acid) is Hydrochloric Acid, which helps break down food and kills harmful bacteria!

What are Bases and Alkalis?

A Base is generally a substance that reacts with an acid to neutralise it. Bases are usually metal oxides or metal hydroxides.

An Alkali is a special type of base:

  • An alkali is a base that is soluble (dissolves) in water.
  • Alkalis produce hydroxide ions (\(OH^-\)) when they dissolve in water.

Memory Aid: Think "A" for Alkali and "A" for Aqua (water). Alkalis are bases that dissolve in water.

Key Examples of Alkalis (Soluble Bases):

  • Sodium Hydroxide (NaOH)
  • Potassium Hydroxide (KOH)
  • Ammonia solution (\(NH_3(aq)\))

Measuring Acidity: The pH Scale and Indicators

The pH Scale is a measure of the acidity or alkalinity of a solution. It runs from 0 to 14.

Quick Review of pH:

  • pH 0 to < 7: Acidic (Higher \(H^+\) concentration)
  • pH 7: Neutral (Pure water, equal amounts of \(H^+\) and \(OH^-\))
  • pH > 7 to 14: Alkaline (Higher \(OH^-\) concentration)
Indicators

Indicators are dyes that change colour depending on the pH of the solution. They are used to quickly test if a substance is acidic or alkaline.

Important Indicators to Know:

  1. Universal Indicator (UI): The most useful, as it shows a whole range of colours (Red for strong acid, Green for neutral, Purple/Violet for strong alkali).
  2. Litmus Paper: Simple test. Acids turn Blue litmus Red. Alkalis turn Red litmus Blue.
  3. Methyl Orange: Used often in titrations. Red in acid, Yellow in alkali.
  4. Phenolphthalein: Used often in titrations. Colourless in acid, Pink in alkali.

Key Takeaway I: Acids release \(H^+\). Alkalis release \(OH^-\) and are soluble bases. The pH scale measures their strength.

Section II: Key Reactions of Acids

Acids are highly reactive. You must know the general equations for the four main types of reactions acids undergo.

1. Reaction with Bases (Neutralisation)

This is the most important reaction. Neutralisation occurs when an acid reacts with a base (or an alkali) to produce a salt and water.

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

The name of the salt depends on the acid used:

  • Hydrochloric acid makes Chlorides.
  • Sulfuric acid makes Sulfates.
  • Nitric acid makes Nitrates.

Example:
\(HCl + NaOH \rightarrow NaCl + H_2O\)

The Ionic Equation for Neutralisation:
Regardless of the acid and alkali used, the essential reaction happening is:
\(H^+(aq) + OH^-(aq) \rightarrow H_2O(l)\)

2. Reaction with Metals

Acids react with metals that are more reactive than hydrogen (refer back to the reactivity series). This reaction produces a salt and hydrogen gas.

General Equation:
Acid + Metal \(\rightarrow\) Salt + Hydrogen gas

Example: Zinc reacting with Sulfuric Acid
\(Zn(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2(g)\)

Safety Note: We usually don't react acids with highly reactive metals like Sodium or Potassium, as the reaction is too explosive! We often use Magnesium or Zinc in the lab.

3. Reaction with Metal Carbonates

When an acid reacts with a metal carbonate, three products are always formed: a salt, water, and carbon dioxide gas.

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

Example: Calcium Carbonate (chalk/limestone) reacting with Hydrochloric Acid
\(2HCl(aq) + CaCO_3(s) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)\)

Testing for Carbon Dioxide: If the gas produced is bubbled through limewater (calcium hydroxide solution), the limewater turns milky (a white precipitate of calcium carbonate forms).

Common Mistake to Avoid: When reacting an acid with a metal, you get Hydrogen (\(H_2\)) gas. When reacting an acid with a carbonate, you get Carbon Dioxide (\(CO_2\)) gas. Don't confuse the two!

Section III: Preparation of Salts

A Salt is a chemical compound formed when the hydrogen ion (\(H^+\)) in an acid is replaced by a metal ion or the ammonium ion (\(NH_4^+\)).

In the IGCSE curriculum, you need to know three main methods for preparing salts, depending on whether the salt is soluble or insoluble in water.

Prerequisite Concept: Solubility Rules
While you don't need to memorise every rule, understanding these basic rules is vital for choosing the correct preparation method:

  • All Sodium (Na), Potassium (K), and Ammonium (\(NH_4\)) salts are soluble.
  • All Nitrates are soluble.
  • Most Chlorides are soluble (except Silver and Lead).
  • Most Sulfates are soluble (except Barium, Calcium, and Lead).
  • Most Carbonates are insoluble (except Sodium, Potassium, and Ammonium).

Method 1: Preparing Soluble Salts (Titration)

This method is used when both the starting base (or alkali) and the final salt are soluble. This method is essential for preparing salts of Sodium, Potassium, or Ammonium.

Because all reactants and products are dissolved, we cannot simply use excess reactant and filter; we must use Titration to find the exact amount needed for complete neutralisation.

Step-by-Step Titration Process:

  1. Measure a fixed volume of alkali (e.g., NaOH) into a conical flask using a pipette.
  2. Add a few drops of a suitable indicator (e.g., methyl orange).
  3. Slowly add the acid (e.g., HCl) from a burette until the indicator just changes colour (the end-point).
  4. Record the volume of acid used. This is called the titre.
  5. Repeat without the indicator: Perform the experiment again using the exact volumes measured in the first run, but this time, do not add the indicator. This ensures the salt solution is pure.
  6. Crystallisation: Gently heat the resulting salt solution (NaCl solution) to evaporate some water (until the saturation point/crystallisation point is reached).
  7. Allow the hot, saturated solution to cool slowly. Crystals will form.
  8. Filter and dry the crystals (e.g., between two pieces of filter paper).

Method 2: Preparing Soluble Salts (Excess Insoluble Reactant)

This method is used when the starting base is insoluble (a metal oxide, hydroxide, or carbonate) and the resulting salt is soluble. This method works for most metals (not Na, K, or \({NH_4}\)).

We use excess insoluble reactant to ensure that all the acid is reacted (neutralised).

Step-by-Step Process:

  1. Add the chosen acid (e.g., Sulfuric Acid) to a beaker and warm it gently.
  2. Add the insoluble reactant (e.g., excess Copper(II) Oxide) slowly while stirring until no more dissolves (the base is in excess).
  3. Filtration: Filter the mixture to remove the excess, unreacted insoluble base. The filtrate (the liquid that passes through) is the pure salt solution (e.g., Copper(II) Sulfate).
  4. Crystallisation: Heat the filtrate to evaporate some water until the point of crystallisation is reached.
  5. Cool the solution slowly to form crystals.
  6. Filter and dry the crystals.

Why this method works: Because the base is insoluble, we can easily remove the excess by filtering, guaranteeing a pure solution of the salt.

Method 3: Preparing Insoluble Salts (Precipitation)

This method is used when the desired salt is insoluble (e.g., Barium Sulfate or Lead Chloride). This method involves mixing two different soluble salts together.

When two solutions containing different ions are mixed, and the resulting combination forms an insoluble compound, that compound falls out of solution as a solid called a precipitate.

Step-by-Step Process:

  1. Choose two soluble compounds that contain the necessary ions. For example, to make insoluble Barium Sulfate (\(BaSO_4\)), mix:
    • A soluble Barium compound (e.g., Barium Nitrate, \(Ba(NO_3)_2\))
    • A soluble Sulfate compound (e.g., Sodium Sulfate, \(Na_2SO_4\))
  2. Mix the two solutions together. A cloudy white precipitate (the insoluble salt) instantly forms.
  3. Filtration: Filter the mixture to separate the solid precipitate from the liquid solution (which contains the other, soluble salt).
  4. Washing: Wash the precipitate several times with distilled water to remove any soluble impurities (the other salt).
  5. Drying: Dry the insoluble salt (precipitate) in an oven or by leaving it in a warm, dry place.

Example Reaction (Ionic Exchange):
\(Ba(NO_3)_2(aq) + Na_2SO_4(aq) \rightarrow BaSO_4(s) + 2NaNO_3(aq)\)

Key Takeaway III: Use Titration for soluble salts of Na, K, or \({NH_4}\). Use Excess Insoluble Reactant for other soluble salts. Use Precipitation for all insoluble salts.


You've covered the definitions, reactions, and practical methods! Keep reviewing the flowcharts for salt preparation—it's the key to mastering this topic! Well done!