Hello, Future Chemist! Understanding Acids, Alkalis, and Titrations
Welcome to one of the most practical and important chapters in Inorganic Chemistry! Acids and alkalis (bases) are everywhere—in your kitchen, in your body, and in industrial processes. Understanding how they behave and react is fundamental to chemistry.
In these notes, we will break down the definitions of acids and alkalis, explore the useful pH scale, look at how acids react to form salts, and master the precise laboratory technique called titration. Don't worry if some terms look tricky; we'll use simple language and real-world analogies to make sure you understand every concept.
1. Defining Acids and Alkalis (Bases)
When we talk about acids and alkalis in Chemistry, we are usually talking about what happens when they dissolve in water.
What is an Acid?
An acid is a substance that produces hydrogen ions (\(H^+\)) when dissolved in water.
- These \(H^+\) ions are what cause the typical acidic properties, such as turning blue litmus paper red and having a sour taste.
- Example: Hydrochloric acid (\(\text{HCl}\)) dissolves in water to give \(H^+\) and \(\text{Cl}^-\) ions.
What is an Alkali (Base)?
A base is a substance that can neutralise an acid (often a metal oxide or hydroxide).
An alkali is simply a base that is soluble in water.
- An alkali produces hydroxide ions (\(OH^-\)) when dissolved in water.
- These \(OH^-\) ions are responsible for typical alkaline properties, such as feeling soapy and turning red litmus paper blue.
- Example: Sodium hydroxide (\(\text{NaOH}\)) dissolves in water to give \(\text{Na}^+\) and \(OH^-\) ions.
Quick Review:
Acids = Produce \(H^+\)
Alkalis = Produce \(OH^-\)
2. The pH Scale and Indicators
The pH scale is a numerical scale (usually 0 to 14) used to measure how acidic or alkaline a solution is.
Understanding pH Values
- pH 7: This is Neutral. Pure water is pH 7.
- pH < 7 (e.g., pH 1, 2, 3): This is Acidic. The lower the number, the stronger the acid.
- pH > 7 (e.g., pH 12, 13, 14): This is Alkaline (Basic). The higher the number, the stronger the alkali.
Did you know? The scale is based on the concentration of \(H^+\) ions. The higher the \(H^+\) concentration, the lower the pH.
Indicators
Indicators are dyes that change colour depending on whether they are in an acidic or alkaline solution. This allows us to determine the pH range of a substance.
Common Indicators
We use different indicators depending on the type of reaction (especially in titrations, which we will cover later).
| Indicator Name | Colour in Acid | Colour in Alkali | Neutral Colour |
|---|---|---|---|
| Litmus Paper | Red | Blue | Purple (if solution) |
| Methyl Orange | Red | Yellow | Orange |
| Phenolphthalein | Colourless | Pink/Magenta | Colourless |
Memory Aid for Phenolphthalein (P):
Think Pink for Phenolphthalein in the basic (alkaline) solution. In acid, it's nothing (colourless).
Universal Indicator: This is a mixture of several dyes that gives a different colour for every pH unit (0 to 14). It is very useful for estimating the pH, but less precise than a pH meter.
3. Neutralisation: The Core Reaction
The most important reaction involving acids and alkalis is neutralisation. This is the reaction between an acid and a base (or alkali) that results in a neutral solution (pH close to 7).
The Ionic Equation for Neutralisation
When any acid reacts with any alkali, the essential reaction is always the same:
Hydrogen ions react with Hydroxide ions to form water.
\[ H^+(aq) + OH^-(aq) \rightarrow H_2O(l) \]
The general equation for neutralisation is:
Acid + Base (or Alkali) \(\rightarrow\) Salt + Water
Analogy: Neutralisation is like combining two opposites (Acid = Sour, Alkali = Bitter/Soapy) to create something neutral (Water).
4. Reactions of Acids to Form Salts
A salt is a compound formed when the hydrogen ion (\(H^+\)) of an acid is replaced by a metal ion or ammonium ion (\(\text{NH}_4^+\)).
There are four main types of reaction that acids undergo to form salts. You must know these!
Reaction 1: Acid + Metal
When a reactive metal reacts with an acid, it produces a salt and hydrogen gas.
Acid + Reactive Metal \(\rightarrow\) Salt + Hydrogen Gas (\(H_2\))
- Example: \(\text{Zn}(s) + 2\text{HCl}(aq) \rightarrow \text{ZnCl}_2(aq) + H_2(g)\)
- Test for \(H_2\): Hold a lit splint near the gas—it produces a loud squeaky pop.
- Caution: Very unreactive metals (like copper) will not react with dilute acids. Very reactive metals (like sodium) react too violently and dangerously.
Reaction 2: Acid + Base (Metal Oxide)
Metal oxides are bases. This is a neutralisation reaction.
Acid + Metal Oxide \(\rightarrow\) Salt + Water
- Example: \(\text{H}_2\text{SO}_4(aq) + \text{CuO}(s) \rightarrow \text{CuSO}_4(aq) + H_2O(l)\)
Reaction 3: Acid + Alkali (Metal Hydroxide)
Metal hydroxides are alkalis (soluble bases). This is also a neutralisation reaction.
Acid + Metal Hydroxide \(\rightarrow\) Salt + Water
- Example: \(\text{HNO}_3(aq) + \text{KOH}(aq) \rightarrow \text{KNO}_3(aq) + H_2O(l)\)
Reaction 4: Acid + Carbonate (or Hydrogencarbonate)
Carbonates react with acids to produce three products: a salt, water, and carbon dioxide gas.
Acid + Metal Carbonate \(\rightarrow\) Salt + Water + Carbon Dioxide (\(\text{CO}_2\))
- Example: \(2\text{HNO}_3(aq) + \text{Na}_2\text{CO}_3(s) \rightarrow 2\text{NaNO}_3(aq) + H_2O(l) + \text{CO}_2(g)\)
- Test for \(\text{CO}_2\): Bubble the gas through limewater (calcium hydroxide solution). If \(\text{CO}_2\) is present, the limewater turns milky (cloudy).
5. Strong vs. Weak Acids and Alkalis
This is a crucial distinction that students often confuse with concentration. They are not the same!
Strength (Ionisation)
Strength refers to the degree to which an acid or alkali ionises (splits up) in water.
- Strong Acids/Alkalis: Substances that fully ionise in water. They release all their available \(H^+\) or \(OH^-\) ions.
- Examples: Hydrochloric acid (\(\text{HCl}\)), Sulfuric acid (\(\text{H}_2\text{SO}_4\)), Sodium hydroxide (\(\text{NaOH}\)).
- Weak Acids/Alkalis: Substances that only partially ionise in water. Most of the molecules remain whole.
- Examples: Ethanoic acid (vinegar), Carbonic acid, Aqueous ammonia.
Key Difference: A strong acid (like \(\text{HCl}\)) gives a much higher concentration of \(H^+\) ions in solution than a weak acid (like ethanoic acid) of the same concentration.
Concentration (Amount)
Concentration refers to how much solute (acid or alkali) is dissolved in a specific volume of solvent (water).
- A concentrated solution has a large amount of solute dissolved.
- A dilute solution has a small amount of solute dissolved.
Analogy: Imagine making cordial (squash).
Concentration relates to how much syrup you put in the jug (a lot = concentrated, a little = dilute).
Strength relates to the actual flavour quality of the syrup itself (Strong acid = very potent syrup, Weak acid = mild syrup).
You can have a concentrated weak acid (lots of mild vinegar) or a dilute strong acid (a tiny amount of potent \(\text{HCl}\) in lots of water).
6. Mastering Titrations
Titration is a highly precise experimental technique used to find the exact volume of one solution (often an alkali) that is needed to exactly neutralise a specific volume of another solution (often an acid).
Why do we Titrate?
The main purpose is to determine the unknown concentration of one of the solutions, provided the concentration of the other solution is known (this is called the standard solution).
Equipment Used
Titration requires specific, precise glassware:
- Burette: A long, graduated glass tube with a tap (stopcock) at the bottom. Used to add the solution of unknown concentration (or the standard solution) drop-by-drop. It allows for highly accurate volume measurement.
- Pipette: Used to measure and transfer a fixed, precise volume of the acid (or alkali) into the conical flask.
- Pipette Filler: Used safely with the pipette to draw up the liquid. Never use your mouth!
- Conical Flask: Used to hold the measured volume of solution and the indicator. It is swirled during the titration.
The Titration Procedure (Step-by-Step)
This method requires careful handling and precision for reliable results.
Step 1: Preparation
- Rinse the burette with the solution it will contain (e.g., the alkali) and then fill it above the zero mark.
- Run a small amount out to ensure the tip is filled and the meniscus is exactly on the zero mark (or any clearly readable mark).
Step 2: Measuring the Reactant
- Use the pipette to accurately measure a fixed volume (e.g., \(25.0 \text{ cm}^3\)) of the acid solution and transfer it into a clean conical flask.
- Add 2–3 drops of the appropriate indicator (e.g., phenolphthalein if titrating a strong acid with a strong alkali).
Step 3: The Rough Titration
- Slowly add the solution from the burette (the titrant) to the conical flask while continuously swirling the flask.
- Stop adding the solution the moment the indicator changes colour permanently. This is called the endpoint. This first titration is usually the "rough" one, done quickly to estimate the volume needed.
- Record the volume reading from the burette.
Step 4: Accurate Titrations (Concordant Results)
- Repeat Steps 2 and 3, but this time, slow down significantly as you approach the rough volume recorded previously.
- When near the endpoint, add the solution drop-by-drop. The moment the colour changes permanently (e.g., colourless to the faintest pink for phenolphthalein), stop.
- Record the final volume. The volume of solution added is called the titre.
- Repeat the titration until you have at least two volumes (titres) that are very close to each other (these are called concordant results).
Common Mistake to Avoid: When reading the volume on the burette, you must read the volume at the bottom of the curved surface (the meniscus) at eye level.
Calculating the Mean Titre
To ensure accuracy, discard the rough titre and calculate the average (mean) volume of the concordant (closest) accurate titres.
\[ \text{Mean Titre} = \frac{\text{Sum of concordant volumes}}{\text{Number of concordant volumes}} \]
This precise mean volume is then used in calculations to find the unknown concentration of the solution.
Summary of Key Takeaways
- Acids release \(H^+\); Alkalis release \(OH^-\) in water.
- Neutral solutions are pH 7. Acidic solutions are pH < 7. Alkaline solutions are pH > 7.
- The overall neutralisation reaction is \(H^+ + OH^- \rightarrow H_2O\).
- Acids react to form salts in four main ways: with metals, metal oxides, metal hydroxides, and metal carbonates.
- Strength (strong vs. weak) refers to the degree of ionisation. Concentration refers to the amount dissolved.
- Titration is a technique to find the exact volume needed for neutralisation, allowing us to calculate unknown concentrations.
You’ve covered the entire theory of acids, alkalis, and titrations! Now practice those titration steps and salt formation equations!