Acids and Bases Revision Notes for VCE SSCE Chemistry

Acids and Bases

Introduction

Acids and bases are all around us in everyday life. You'll find them in many household products, from cleaning agents to food items. Understanding their properties and behaviour is essential for chemistry.

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Acids and bases aren't just confined to chemistry laboratories - they're present in everyday items like citrus fruits, vinegar, soap, and even the acid in your stomach that helps digest food. Recognizing these substances in daily life helps make chemistry concepts more relatable and practical.

In this topic, you'll learn about the Brønsted-Lowry theory, which explains how acids and bases behave. You'll discover how to write balanced equations for acid-base reactions and understand that some substances can act as either an acid or a base, depending on the reaction.

Properties of acids and bases

Common acids

Acids play important roles in our homes, in industry, and even in our bodies. For example, hydrochloric acid is present in your stomach to help digest proteins. The table below shows some common acids you might encounter.

Name	Formula	Uses
Hydrochloric acid	$\text{HCl}$	Present in stomach acid to help break down proteins; used as a cleaning agent for brickwork
Sulfuric acid	$\text{H}_2\text{SO}_4$	One of the most common chemicals manufactured; used in car batteries and in the manufacture of fertilisers and detergents
Nitric acid	$\text{HNO}_3$	Used in the manufacture of fertilisers, dyes and explosives
Ethanoic acid (acetic acid)	$\text{CH}_3\text{COOH}$	Found in vinegar; used as a preservative
Carbonic acid	$\text{H}_2\text{CO}_3$	Found in carbonated soft drinks and beer
Phosphoric acid	$\text{H}_3\text{PO}_4$	Used in some soft drinks and in the manufacture of fertilisers
Citric acid	$\text{C}_6\text{H}_8\text{O}_7$	Found in citrus fruits
Ascorbic acid	$\text{C}_6\text{H}_8\text{O}_6$	Found in citrus fruits (vitamin C)

Common bases

Many cleaning products contain bases. For instance, washing powders, oven cleaners, and floor cleaners often have bases as their active ingredients. Bases work well as cleaners because they react with fats and oils to create water-soluble soaps.

Bases that dissolve in water are called alkalis. Examples include ammonia ( $\text{NH}_3$ ) and sodium hydroxide ( $\text{NaOH}$ ). Some bases, like calcium carbonate, don't dissolve in water, so they're not classified as alkalis even though they can still react with acids.

Name	Formula	Uses
Sodium hydroxide (caustic soda)	$\text{NaOH}$	Used in drain and oven cleaners, and soap making
Ammonia	$\text{NH}_3$	Used in household cleaners, fertilisers and explosives
Calcium hydroxide	$\text{Ca(OH)}_2$	Found in cement and mortar; used in garden lime to adjust soil pH
Magnesium hydroxide	$\text{Mg(OH)}_2$	Key ingredient in some antacids, such as milk of magnesia, to overcome indigestion
Sodium carbonate	$\text{Na}_2\text{CO}_3$	Used in the manufacture of washing powder and glass

Comparing properties

All acids share certain characteristics, as do all bases. Understanding these common properties helps you identify and work with these substances safely.

Properties of acids	Properties of bases
Turn litmus indicator red	Turn litmus indicator blue
Tend to be corrosive*	Are corrosive, caustic** and slippery
Taste sour	Taste bitter
React with bases	React with acids
Solutions have a relatively low pH	Solutions have a relatively high pH
Solutions conduct an electric current	Solutions conduct an electric current

*Corrosive means able to dissolve or destroy the structure of an object.

**Caustic means able to burn or corrode organic tissue through chemical action.

chatImportant

Safety Warning: Strong bases feel slippery because they react with fats in your skin to produce soap. Always handle strong bases with care. For example, oven cleaners contain about 4% sodium hydroxide (also called caustic soda) and can cause burns. Never touch strong acids or bases directly, and always wear appropriate safety equipment when handling them.

Development of acid-base theories

Scientists have developed different ways to understand acids and bases over the centuries. Each new theory built upon previous observations and helped explain more about how these substances behave.

Early observations

In the seventeenth century, British scientist Robert Boyle described acids based on their taste, their ability to dissolve other substances, and how they changed the colour of certain plant extracts. He also noticed that alkalis (soluble bases) could reverse the colour changes that acids produced.

Lavoisier's oxygen theory

In the late eighteenth century, French chemist Antoine Lavoisier proposed that acidic properties came from the presence of oxygen. This explained why sulfuric acid ( $\text{H}_2\text{SO}_4$ ), nitric acid ( $\text{HNO}_3$ ), and phosphoric acid ( $\text{H}_3\text{PO}_4$ ) were acidic, but it couldn't explain why hydrochloric acid ( $\text{HCl}$ ) was acidic despite containing no oxygen.

Davy's hydrogen theory

Around 1810, Humphrey Davy suggested that acidic properties were actually associated with hydrogen, not oxygen. He reached this conclusion after producing hydrogen gas by reacting acids with metals. Davy also proposed that acids react with bases to form compounds called salts, plus water.

Arrhenius theory

In 1887, Swedish scientist Svante Arrhenius developed a more detailed theory:

Acids are substances that break apart (dissociate) and form ions (ionise) in water to produce hydrogen ions ( $\text{H}^+$ )
Bases dissociate in water to produce hydroxide ions ( $\text{OH}^-$ )

Brønsted-Lowry theory

In 1923, Danish chemist Johannes Nicolaus Brønsted and English chemist Thomas Martin Lowry independently developed the theory that now bears both their names. The Brønsted-Lowry theory is more comprehensive than earlier theories and can explain acid-base behaviours that Arrhenius's theory couldn't account for.

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The progression of acid-base theories shows how scientific understanding evolves. Each theory wasn't necessarily "wrong" - instead, each new theory built upon previous knowledge and extended our understanding to explain more situations. The Brønsted-Lowry theory is particularly useful because it applies to reactions beyond just aqueous solutions.

The Brønsted-Lowry theory

The Brønsted-Lowry theory provides our modern understanding of acids and bases. It's based on a simple but powerful concept: the transfer of protons.

Core definitions

chatImportant

The Brønsted-Lowry Theory - Key Definitions:

According to this theory:

Acids are proton donors - they give away a proton ( $\text{H}^+$ ) to another substance
Bases are proton acceptors - they receive a proton from an acid
An acid-base reaction involves the transfer of protons from an acid to a base

Remember: "Acids Donate, Bases Accept" (protons)

The hydronium ion

When acids dissolve in water, they donate protons to water molecules. This creates a special ion called the hydronium ion, written as $\text{H}_3\text{O}^+(\text{aq})$ .

infoNote

You might also see the hydronium ion written simply as $\text{H}^+(\text{aq})$ . Both notations represent the same thing, but writing $\text{H}_3\text{O}^+(\text{aq})$ makes it clearer that a proton transfer has occurred. In reality, a free proton ( $\text{H}^+$ ) cannot exist independently in solution - it immediately bonds to a water molecule to form the hydronium ion.

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Example: Hydrochloric Acid in Water

When hydrogen chloride gas dissolves in water, nearly all the molecules break apart and form ions. This creates hydrochloric acid:

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

In this reaction:

Each $\text{HCl}$ molecule donates a proton to a water molecule, so $\text{HCl}$ acts as an acid
Each water molecule accepts a proton from $\text{HCl}$ , so water acts as a base

Advantages of the Brønsted-Lowry model

The Brønsted-Lowry theory is more useful than earlier models because it's not limited to reactions in water. For example, it can describe reactions between gases:

In solution: $\text{HCl}(\text{aq}) + \text{NH}_3(\text{aq}) \rightarrow \text{NH}_4^+(\text{aq}) + \text{Cl}^-(\text{aq})$

As gases: $\text{HCl}(\text{g}) + \text{NH}_3(\text{g}) \rightarrow \text{NH}_4\text{Cl}(\text{s})$

Both reactions involve the same proton transfer from acid to base, so the Brønsted-Lowry theory classifies both as acid-base reactions.

Conjugate acid-base pairs

When an acid donates a proton, it becomes a new species that could potentially accept a proton back. These related species form what we call a conjugate acid-base pair.

Understanding conjugate pairs

chatImportant

Definition: A conjugate acid-base pair consists of two species whose formulas differ by exactly one proton ( $\text{H}^+$ ).

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Example: Identifying Conjugate Pairs in the HCl Reaction

Looking at the reaction of hydrochloric acid with water:

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

$\text{Cl}^-$ can be formed from $\text{HCl}$ by losing one proton, so $\text{Cl}^-$ is the conjugate base of HCl
$\text{HCl}$ is the conjugate acid of Cl $^-$
Together, $\text{HCl}$ and $\text{Cl}^-$ form a conjugate acid-base pair
Similarly, $\text{H}_3\text{O}^+$ and $\text{H}_2\text{O}$ also form a conjugate pair

By convention, we usually write the acid form of the conjugate pair first.

Common conjugate pairs

When an acid loses a proton, it forms its conjugate base:

When a base gains a proton, it forms its conjugate acid:

Ionisation reactions

When acids react with water, they produce hydronium ions ( $\text{H}_3\text{O}^+$ ). When bases react with water, they produce hydroxide ions ( $\text{OH}^-$ ). We call these ionisation reactions because ions are formed.

Writing equations for acid-base reactions

Chemical equations help us represent acid-base reactions clearly. When we show which species exist as ions, we call these ionic equations. For reactions in water, ions are in an aqueous state, shown by the symbol $(\text{aq})$ .

Understanding arrow symbols

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Arrow Notation in Chemical Equations:

Different arrow symbols tell us different things about how complete a reaction is:

Complete reactions use a forward arrow ( $\rightarrow$ ):

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

This indicates that virtually all the $\text{HCl}$ molecules react when sufficient water is present.

Incomplete reactions use a reversible arrow ( $\rightleftharpoons$ ):

This indicates that only a small proportion of reactant molecules are consumed, so the reaction mixture contains both reactants and products.

Amphiprotic substances

chatImportant

Definition: Amphiprotic substances can behave as either an acid or a base, depending on what they're reacting with. These substances can either donate or accept protons, depending on the reaction conditions.

Water as an amphiprotic substance

Water is the most important example of an amphiprotic substance. It can either donate or accept protons.

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Example: Water as an Amphiprotic Substance

Water acting as a base (accepting a proton from $\text{HCl}$ ):

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

Water acting as an acid (donating a proton to $\text{NH}_3$ ):

$\text{NH}_3(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightarrow \text{NH}_4^+(\text{aq}) + \text{OH}^-(\text{aq})$

Whether water acts as an acid or base depends on the other reactant. If the other substance is a stronger acid than water, water will act as a base. If the other substance is a stronger base than water, water will act as an acid.

Other amphiprotic substances

Many other substances are amphiprotic. They can donate a proton under some conditions or accept a proton under other conditions.

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Example: Hydrogen Carbonate Ion as Amphiprotic

The hydrogen carbonate ion ( $\text{HCO}_3^-$ ) is amphiprotic. In water, it can react in two ways:

As an acid (donating a proton):

$\text{HCO}_3^-(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{CO}_3^{2-}(\text{aq}) + \text{H}_3\text{O}^+(\text{aq})$

As a base (accepting a proton):

$\text{HCO}_3^-(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{H}_2\text{CO}_3(\text{aq}) + \text{OH}^-(\text{aq})$

The reversible arrows ( $\rightleftharpoons$ ) indicate that $\text{HCO}_3^-$ doesn't react completely as either an acid or a base.

infoNote

For most amphiprotic substances in water, one reaction usually dominates over the other. We can identify which reaction is dominant by measuring the pH of the solution.

Monoprotic acids

Monoprotic acids can donate only one proton per molecule. Common examples include:

Hydrochloric acid ( $\text{HCl}$ )
Hydrofluoric acid ( $\text{HF}$ )
Nitric acid ( $\text{HNO}_3$ )
Ethanoic acid ( $\text{CH}_3\text{COOH}$ )

The acidic proton

infoNote

Even though ethanoic acid contains four hydrogen atoms, each molecule can donate only one proton. This makes it monoprotic. Only the hydrogen atom bonded to the electronegative oxygen atom can be donated. We call this the acidic proton.

The acidic proton is part of a highly polar O-H bond. The other three hydrogen atoms are bonded to carbon and cannot be donated as protons.

Polyprotic acids

Polyprotic acids can donate more than one proton per molecule. The number of protons an acid can donate depends on its structure. Importantly, polyprotic acids don't donate all their protons at once. Instead, they donate them in stages when reacting with a base.

Diprotic acids

Diprotic acids can donate two protons. Examples include sulfuric acid ( $\text{H}_2\text{SO}_4$ ) and carbonic acid ( $\text{H}_2\text{CO}_3$ ).

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Example: Stepwise Ionisation of Sulfuric Acid (Diprotic)

A diprotic acid ionises in two stages. Let's look at sulfuric acid:

Stage 1:

$\text{H}_2\text{SO}_4(\text{l}) + \text{H}_2\text{O}(\text{l}) \rightarrow \text{HSO}_4^-(\text{aq}) + \text{H}_3\text{O}^+(\text{aq})$

Sulfuric acid is a strong acid, so this first stage occurs almost completely. Virtually no $\text{H}_2\text{SO}_4$ molecules remain in aqueous solution, which is why we use the forward arrow ( $\rightarrow$ ).

Stage 2:

$\text{HSO}_4^-(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{SO}_4^{2-}(\text{aq}) + \text{H}_3\text{O}^+(\text{aq})$

The hydrogen sulfate ion ( $\text{HSO}_4^-$ ) formed in stage 1 can also act as an acid. However, it's a weak acid, so only a small proportion reacts further in a 1.0 M solution. The reversible arrow ( $\rightleftharpoons$ ) indicates this incomplete reaction.

Therefore, a solution of sulfuric acid contains hydrogen ions, hydrogen sulfate ions, sulfate ions, and water molecules.

Triprotic acids

Triprotic acids can donate three protons. Examples include phosphoric acid ( $\text{H}_3\text{PO}_4$ ) and arsenic acid ( $\text{H}_3\text{AsO}_4$ ).

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Example: Stepwise Ionisation of Phosphoric Acid (Triprotic)

A triprotic acid ionises in three stages:

Stage 1:

$\text{H}_3\text{PO}_4(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{H}_2\text{PO}_4^-(\text{aq}) + \text{H}_3\text{O}^+(\text{aq})$

Stage 2:

$\text{H}_2\text{PO}_4^-(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{HPO}_4^{2-}(\text{aq}) + \text{H}_3\text{O}^+(\text{aq})$

Stage 3:

$\text{HPO}_4^{2-}(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{PO}_4^{3-}(\text{aq}) + \text{H}_3\text{O}^+(\text{aq})$

Phosphoric acid is a moderately weak acid. In a 1.0 M solution, only a small proportion of protons is donated at each stage. The extent of ionisation decreases progressively from stage 1 to stage 3. This happens because removing a positive proton from an already negatively charged species becomes more difficult at each step.

bookmarkSummary

Key Points to Remember:

Acids are proton donors and bases are proton acceptors according to the Brønsted-Lowry theory
A conjugate acid-base pair consists of two species that differ by exactly one proton ( $\text{H}^+$ )
Amphiprotic substances can act as either acids or bases depending on what they react with; water is the most common example
Monoprotic acids donate one proton, while polyprotic acids can donate two or more protons in successive stages
When acids react with water, they produce hydronium ions ( $\text{H}_3\text{O}^+$ ); when bases react with water, they produce hydroxide ions ( $\text{OH}^-$ )
The Brønsted-Lowry theory is more versatile than the Arrhenius theory because it applies to reactions beyond aqueous solutions
In polyprotic acids, each successive proton is harder to remove than the previous one

Acids and Bases (VCE SSCE Chemistry): Revision Notes

Acids and Bases

Introduction

Properties of acids and bases

Common acids

Common bases

Comparing properties

Development of acid-base theories

Early observations

Lavoisier's oxygen theory

Davy's hydrogen theory

Arrhenius theory

Brønsted-Lowry theory

The Brønsted-Lowry theory

Core definitions

The hydronium ion

Advantages of the Brønsted-Lowry model

Conjugate acid-base pairs

Understanding conjugate pairs

Common conjugate pairs

Ionisation reactions

Writing equations for acid-base reactions

Understanding arrow symbols

Amphiprotic substances

Water as an amphiprotic substance

Other amphiprotic substances

Monoprotic acids

The acidic proton

Polyprotic acids

Diprotic acids

Triprotic acids

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