Chemistry (9701) | Brønsted–Lowry theory of acids and bases

Chemistry 9701 Study Notes: Brønsted–Lowry Theory of Acids and Bases (Equilibria 7.2)

Welcome to the world of acids and bases! This chapter moves beyond simple definitions to explore why some acids are fierce and others are mild. Understanding the Brønsted–Lowry theory is essential because it gives us the language to describe reactions involving proton transfer, which is fundamental to almost all aqueous chemistry. Don't worry if this seems tricky at first; we'll break down the concepts using clear language and helpful analogies!

1. Defining Common Acids and Alkalis (Syllabus 7.2.1 & 7.2.2)

Before diving into the theory, let's quickly recall the substances you need to know:

Acids (The Proton Donors)

• Hydrochloric acid (\(\text{HCl}\))
• Sulfuric acid (\(\text{H}_2\text{SO}_4\))
• Nitric acid (\(\text{HNO}_3\))
• Ethanoic acid (\(\text{CH}_3\text{COOH}\)) - This is vinegar! It’s our key example of a weak acid.

Alkalis (The Soluble Bases/Proton Acceptors)

• Sodium hydroxide (\(\text{NaOH}\))
• Potassium hydroxide (\(\text{KOH}\))
• Ammonia (\(\text{NH}_3\)) - Our key example of a weak base.

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2. The Brønsted–Lowry Theory: The Proton Handshake (Syllabus 7.2.3)

The Brønsted–Lowry definition is simple and powerful. It focuses entirely on the movement of a proton (\(\text{H}^+\)).

Key Definitions

• Brønsted–Lowry Acid: A species that is a proton donor. (It gives away an \(\text{H}^+\)).
• Brønsted–Lowry Base: A species that is a proton acceptor. (It takes an \(\text{H}^+\)).

Analogy: Think of the proton (\(\text{H}^+\)) as a hot potato being passed between two molecules. The acid throws the potato (donates the proton), and the base catches it (accepts the proton).

Example: Reaction of \(\text{HCl}\) and \(\text{H}_2\text{O}\)

\[\text{HCl}(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightarrow \text{Cl}^-(\text{aq}) + \text{H}_3\text{O}^+(\text{aq})\]

• \(\text{HCl}\) is the acid (it donated \(\text{H}^+\)).
• \(\text{H}_2\text{O}\) is the base (it accepted \(\text{H}^+\) to become the hydroxonium ion, \(\text{H}_3\text{O}^+\)).

Did you know?

The \(\text{H}^+\) ion is often called a proton because a hydrogen atom (\(\text{}^1\text{H}\)) consists of just one proton and one electron. When it loses the electron to become \(\text{H}^+\), all that remains is the proton!

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3. Conjugate Acid-Base Pairs

When an acid reacts, it forms a base, and when a base reacts, it forms an acid. These related species are called conjugate pairs (Syllabus 7.2.2).

Key Rule for Conjugate Pairs

A conjugate acid-base pair only differs by one proton (\(\text{H}^+\)).

Acid \(\rightleftharpoons\) Conjugate Base + \(\text{H}^+\)
Base + \(\text{H}^+\) \(\rightleftharpoons\) Conjugate Acid

Example: Ammonia in Water

The equilibrium reaction for ammonia (\(\text{NH}_3\)) in water:

\[\text{NH}_3(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{NH}_4^+(\text{aq}) + \text{OH}^-(\text{aq})\]

Let’s identify the pairs:

1. \(\text{NH}_3\) (Base) accepts \(\text{H}^+\) to form \(\text{NH}_4^+\) (Conjugate Acid). (Pair 1)
2. \(\text{H}_2\text{O}\) (Acid) donates \(\text{H}^+\) to form \(\text{OH}^-\) (Conjugate Base). (Pair 2)

Important Note: Notice that water (\(\text{H}_2\text{O}\)) acted as an acid here. Substances that can act as both an acid and a base (depending on what they react with) are called amphoteric or amphiprotic.

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4. Strength of Acids and Bases (Dissociation) (Syllabus 7.2.4)

The strength of an acid or base is defined by the extent of its dissociation (or ionisation) in water.

A. Strong Acids and Bases

Strong acids and strong bases are assumed to be fully dissociated (100% ionised) in aqueous solution.

• Strong Acid Example (\(\text{HCl}\)): The arrow is one-way, meaning nearly every molecule breaks apart: \[\text{HCl}(\text{aq}) \rightarrow \text{H}^+(\text{aq}) + \text{Cl}^-(\text{aq})\]
• Strong Base Example (\(\text{NaOH}\)): The ions fully separate: \[\text{NaOH}(\text{aq}) \rightarrow \text{Na}^+(\text{aq}) + \text{OH}^-(\text{aq})\]

B. Weak Acids and Bases

Weak acids and weak bases are partially dissociated (only a small fraction ionises) in aqueous solution. An equilibrium is established.

• Weak Acid Example (\(\text{CH}_3\text{COOH}\)): The equilibrium lies mostly to the left: \[\text{CH}_3\text{COOH}(\text{aq}) \rightleftharpoons \text{H}^+(\text{aq}) + \text{CH}_3\text{COO}^-(\text{aq})\]
• Weak Base Example (\(\text{NH}_3\)): \[\text{NH}_3(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{NH}_4^+(\text{aq}) + \text{OH}^-(\text{aq})\]

Memory Aid: Think of the "W" in "Weak" standing for "Wobbly equilibrium" that favours the reactants.

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5. Qualitative Differences: Strong vs. Weak (Syllabus 7.2.5 & 7.2.6)

You must be able to explain the physical and chemical differences between strong and weak acids, even if they have the same concentration (e.g., 1.0 mol \(\text{dm}^{-3}\) \(\text{HCl}\) versus 1.0 mol \(\text{dm}^{-3}\) \(\text{CH}_3\text{COOH}\)).

Difference 1: \(\text{pH}\) Value (Syllabus 7.2.5)

Recall the \(\text{pH}\) scale:

• Acidic solutions: \(\text{pH} < 7\)
• Pure Water (Neutral): \(\text{pH} = 7\)
• Alkaline solutions: \(\text{pH} > 7\)

Because strong acids fully dissociate, they produce a much higher concentration of \(\text{H}^+(\text{aq})\) ions than weak acids of the same concentration. This means:

• Strong Acid: Has a much lower \(\text{pH}\) (more acidic). (e.g., \(\text{pH}\) 1)
• Weak Acid: Has a higher \(\text{pH}\) (less acidic). (e.g., \(\text{pH}\) 3 or 4)

Detection Method: Use a \(\text{pH}\) meter or Universal Indicator.

Difference 2: Electrical Conductivity

Electrical conductivity depends on the concentration of mobile ions in solution.

• Strong Acid/Base: High concentration of ions due to full dissociation/ionisation \(\rightarrow\) High conductivity.
• Weak Acid/Base: Low concentration of ions due to partial dissociation \(\rightarrow\) Low conductivity.

Difference 3: Reaction with Reactive Metals

Acids react with reactive metals (like Mg or Zn) to produce \(\text{H}_2\) gas. This reaction depends on the concentration of \(\text{H}^+(\text{aq})\) ions.

\[\text{Mg}(\text{s}) + 2\text{H}^+(\text{aq}) \rightarrow \text{Mg}^{2+}(\text{aq}) + \text{H}_2(\text{g})\]

• Strong Acid: High \(\text{H}^+\) concentration \(\rightarrow\) Fast rate of reaction (vigorous bubbling).
• Weak Acid: Low \(\text{H}^+\) concentration \(\rightarrow\) Slow rate of reaction (slow bubbling).

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6. Neutralisation and Salt Formation (Syllabus 7.2.7 & 7.2.8)

Neutralisation is one of the most important acid-base reactions.

The Overall Neutralisation Process

When an acid reacts with a base (or alkali), they produce a salt and water (Syllabus 7.2.8).

Example: \(\text{NaOH}\) and \(\text{HCl}\) \[\text{HCl}(\text{aq}) + \text{NaOH}(\text{aq}) \rightarrow \text{NaCl}(\text{aq}) + \text{H}_2\text{O}(\text{l})\]

The Essential Net Ionic Equation (Syllabus 7.2.7)

Regardless of whether the acid or alkali is strong or weak, the fundamental chemical event in neutralisation is the reaction between the aqueous hydrogen ion and the aqueous hydroxide ion to form water:

\[\mathbf{H^+(\text{aq}) + OH^-(\text{aq}) \rightarrow H_2O(\text{l})}\]

Note: The other ions (like \(\text{Na}^+\) and \(\text{Cl}^-\)) are spectator ions and combine to form the salt.

Key Takeaway: Neutralisation is fundamentally the formation of the neutral water molecule from the acidic proton and the basic hydroxide ion.

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7. pH Titration Curves and Indicator Selection (Syllabus 7.2.9 & 7.2.10)

Titration curves show how the \(\text{pH}\) changes as an acid is gradually added to a base (or vice versa). You must be able to sketch and interpret four main types of curves.

Key Features of Titration Curves

1. Starting \(\text{pH}\): Determined by the substance in the flask (high for strong base, low for strong acid).
2. Buffering Region: Only visible when a weak component is present (flatter region).
3. Equivalence Point: The point where the acid and base have reacted exactly stoichiometrically.
4. Vertical Region (Inflection Point): The rapid jump/drop in \(\text{pH}\) near the equivalence point.

The Four Types of Titration Curves (Base in the Flask, Acid Added)

(Assuming titrating base with acid)

1. Strong Acid vs. Strong Base (SA/SB):
   • Starts high (e.g., \(\text{pH}\) 13).
   • Equivalence point is exactly at \(\text{pH} = 7\).
   • The vertical region is very long (e.g., from \(\text{pH}\) 3 to \(\text{pH}\) 11).
2. Strong Acid vs. Weak Base (SA/WB):
   • Starts high but below SB (e.g., \(\text{pH}\) 10).
   • Equivalence point is below \(\text{pH} = 7\) (acidic salt formed).
   • Vertical region is shorter and starts lower.
3. Weak Acid vs. Strong Base (WA/SB):
   • Starts high (e.g., \(\text{pH}\) 13).
   • Equivalence point is above \(\text{pH} = 7\) (basic salt formed).
   • Vertical region is shorter and ends higher.
4. Weak Acid vs. Weak Base (WA/WB):
   • No sharp vertical region. The \(\text{pH}\) changes gradually.
   • (This type is rarely used for titration calculations because the end point is too hard to determine accurately.)

Selecting a Suitable Indicator (Syllabus 7.2.10)

An indicator is a weak acid/base whose color changes over a specific \(\text{pH}\) range.

The crucial principle for selecting an indicator is that its colour change must occur entirely within the vertical region of the titration curve.

Since the vertical region represents the sharp change in \(\text{pH}\) at the equivalence point, if the indicator changes color during this jump, it will give an accurate endpoint.

• For SA/SB titrations: The vertical region is huge, so most indicators (like Methyl Orange or Phenolphthalein) are suitable.
• For WA/SB titrations (Eq. point > 7): Need an indicator that changes color in the basic range (e.g., Phenolphthalein, range ~8.3–10.0).
• For SA/WB titrations (Eq. point < 7): Need an indicator that changes color in the acidic range (e.g., Methyl Orange, range ~3.1–4.4).