Acid-Base and Acid-Carbonate Reactions (HSC SSCE Chemistry): Revision Notes

Acid-Base and Acid-Carbonate Reactions

Introduction to acids and bases

Acids and bases are important chemical substances found in both household products and industrial processes. Understanding their properties and reactions is fundamental to chemistry. In this topic, we explore how these substances behave, how they react with each other, and their practical applications.

What are acids?

An acid is a substance that produces hydrogen ions ($\text{H}^+$) when dissolved in solution. However, hydrogen ions do not exist freely in aqueous solution. Instead, they attach to water molecules to form hydronium ions ($\text{H}_3\text{O}^+$). For convenience, we often write $\text{H}^+$ in equations, but it's important to remember the actual form is $\text{H}_3\text{O}^+$.

Common acids

The most frequently encountered acids in chemistry include:

• Hydrochloric acid ($\text{HCl}$) - found in stomach acid and used in industry
• Sulfuric acid ($\text{H}_2\text{SO}_4$) - used in car batteries and fertiliser production
• Nitric acid ($\text{HNO}_3$) - used in fertilisers and explosives
• Phosphoric acid ($\text{H}_3\text{PO}_4$) - found in soft drinks
• Acetic acid ($\text{CH}_3\text{COOH}$) - the main component of vinegar
• Carbonic acid ($\text{H}_2\text{CO}_3$) - forms when carbon dioxide dissolves in water

Everyday household substances containing acids include vinegar (approximately 6% acetic acid), lemons (containing citric acid), vitamin C, and aspirin.

Properties of acids

Acids share several characteristic properties that arise from the presence of hydrogen ions:

• Sour taste - although you should never taste chemicals in a laboratory setting
• Electrical conductivity - acids conduct electricity when dissolved in water due to the presence of ions
• Litmus test - acids change blue litmus paper to red

A useful memory aid for litmus testing is BAR: Blue in Acid goes Red.

What are bases?

A base is a substance that contains the hydroxide ion ($\text{OH}^-$), the oxide ion ($\text{O}^{2-}$), or produces the hydroxide ion when dissolved in solution. Metallic oxides and hydroxides are ionic compounds that contain these ions, making them bases.

Common bases

Frequently encountered bases include:

• Sodium hydroxide ($\text{NaOH}$) - used in drain cleaners and soap making
• Barium hydroxide ($\text{Ba(OH)}_2$)
• Magnesium oxide ($\text{MgO}$) - used in antacids
• Iron(III) oxide ($\text{Fe}_2\text{O}_3$) - rust
• Copper hydroxide ($\text{Cu(OH)}_2$)
• Ammonia ($\text{NH}_3$) - used in cleaning products

Alkalis

Alkalis are bases that dissolve in water. Not all bases are alkalis because some are insoluble. For example:

• Sodium hydroxide and potassium hydroxide are both bases and alkalis (soluble)
• Magnesium oxide and copper hydroxide are bases but not alkalis (insoluble)

Common household alkalis include sodium hydroxide (found in oven cleaners), sodium carbonate (washing soda), and magnesium hydroxide (found in antacids and toothpaste).

Properties of alkalis

Alkalis have distinct properties that help identify them:

• Soapy feel - they feel slippery to touch
• Bitter taste - though students should never taste chemicals
• Electrical conductivity - like acids, they conduct electricity in solution due to ions
• Litmus test - alkalis turn red litmus paper blue (the opposite of acids)

Neutralisation reactions

When an acid reacts with a base, they undergo a neutralisation reaction. This type of reaction produces ionic compounds called salts, typically along with water.

The general equation for neutralisation is:

$\text{acid} + \text{base} \rightarrow \text{salt} + \text{water}$

A salt can be defined in two ways:

1. An ionic compound formed when a base reacts with an acid
2. A compound formed when the hydrogen ion of an acid is replaced by a metal ion

Examples of neutralisation reactions

Ionic equations for neutralisation

In aqueous acid solutions, hydrogen ions ($\text{H}^+$) are present along with their corresponding anions (such as $\text{Cl}^-$, $\text{NO}_3^-$, or $\text{SO}_4^{2-}$). During neutralisation, bases react specifically with these hydrogen ions. The anions remain unchanged in solution and are called spectator ions because they don't participate in the actual reaction.

For the reaction between hydrochloric acid and sodium hydroxide, we can write a complete ionic equation:

$\text{H}^+ + \text{Cl}^- + \text{Na}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O(l)} + \text{Na}^+ + \text{Cl}^-$

Since both $\text{Na}^+$ and $\text{Cl}^-$ are spectator ions, the net ionic equation is:

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

Preparing salts

Salts can be prepared through neutralisation reactions. For example, to prepare zinc sulfate, we would react zinc hydroxide with sulfuric acid:

$\text{Zn(OH)}_2\text{(s)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + 2\text{H}_2\text{O(l)}$

Naming salts from acids

Understanding how to name salts and predict their formulae is an essential skill in chemistry. The relationship between acids and the salts they form follows predictable patterns.

Naming patterns

The metal ion (or cation) is named first, followed by the anion name derived from the acid:

Hydrohalic acids (hydrochloric, hydrobromic, hydroiodic) form salts called halides:

• Hydrochloric acid ($\text{HCl}$) forms chlorides (e.g., magnesium chloride, $\text{MgCl}_2$)
• Hydrobromic acid ($\text{HBr}$) forms bromides (e.g., potassium bromide, $\text{KBr}$)
• Hydroiodic acid ($\text{HI}$) forms iodides (e.g., silver iodide, $\text{AgI}$)

Oxyacids have oxygen attached to elements like sulfur, nitrogen, phosphorus, or carbon. When an oxyacid name ends in '-ic', the salt name ends in '-ate':

• Sulfuric acid ($\text{H}_2\text{SO}_4$) forms sulfates (e.g., aluminium sulfate, $\text{Al}_2(\text{SO}_4)_3$)
• Nitric acid ($\text{HNO}_3$) forms nitrates (e.g., lead nitrate, $\text{Pb(NO}_3)_2$)
• Carbonic acid ($\text{H}_2\text{CO}_3$) forms carbonates (e.g., potassium carbonate, $\text{K}_2\text{CO}_3$)
• Phosphoric acid ($\text{H}_3\text{PO}_4$) forms phosphates (e.g., sodium phosphate, $\text{Na}_3\text{PO}_4$)
• Acetic acid ($\text{CH}_3\text{COOH}$) forms acetates (e.g., silver acetate, $\text{Ag(CH}_3\text{COO)}$)

Anions formed from oxyacids are called oxyanions.

Determining salt formulae

To work out the formula of a salt, you need to:

1. Identify the acid and determine the charge on its anion (from the number of $\text{H}^+$ ions it loses)
2. Identify the charge on the metal ion (cation)
3. Combine them in a ratio where total positive charges equal total negative charges

Acid-carbonate reactions

Carbonates are salts of carbonic acid ($\text{H}_2\text{CO}_3$). When carbonates react with acids, they produce three products: a salt, carbon dioxide gas, and water. This is a characteristic reaction that makes carbonates easy to identify.

General pattern of acid-carbonate reactions

The general equation for these reactions is:

$\text{carbonate} + \text{acid} \rightarrow \text{salt} + \text{CO}_2\text{(g)} + \text{H}_2\text{O(l)}$

Examples of acid-carbonate reactions

Ionic equations for acid-carbonate reactions

In all these reactions, the hydrogen ion from the acid reacts with the carbonate. We can write partial ionic equations to show this:

$\text{CuCO}_3\text{(s)} + 2\text{H}^+\text{(aq)} \rightarrow \text{Cu}^{2+}\text{(aq)} + \text{CO}_2\text{(g)} + \text{H}_2\text{O(l)}$

$\text{K}_2\text{CO}_3\text{(aq)} + 2\text{H}^+\text{(aq)} \rightarrow 2\text{K}^+\text{(aq)} + \text{CO}_2\text{(g)} + \text{H}_2\text{O(l)}$

Whether the carbonate is in solution or solid form, it's the carbonate ion reacting with hydrogen ions. The general net ionic equation is:

$\text{CO}_3^{2-}\text{(s or aq)} + 2\text{H}^+\text{(aq)} \rightarrow \text{CO}_2\text{(g)} + \text{H}_2\text{O(l)}$

Test for carbon dioxide

The production of carbon dioxide in acid-carbonate reactions provides a useful test for identifying carbon dioxide gas in the laboratory.

Limewater test: The standard test involves bubbling the gas through clear limewater (calcium hydroxide solution). If carbon dioxide is present, the solution turns milky white due to the formation of insoluble calcium carbonate:

$\text{Ca(OH)}_2\text{(aq)} + \text{CO}_2\text{(g)} \rightarrow \text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)}$

Real-world application: cycad fruit processing

Aboriginal and Torres Strait Islander Peoples have used chemical processes for thousands of years, long before formal chemistry was established as a science. One fascinating example is the processing of cycad fruit to make it safe for consumption.

The challenge

Cycads are a type of palm that grows commonly in northern and north-eastern Australia. Whilst cycad fruit has been a traditional food source for Indigenous peoples for millennia, the fruit as harvested is poisonous and requires careful treatment to become edible.

Traditional processing methods

Indigenous peoples developed several sophisticated methods to detoxify cycad fruit, involving both physical and chemical processes:

Final preparation

After treatment, the seeds are typically ground into a meal and baked on hot stones to create a type of flat bread. This bread is nutritious, though relatively bland in flavour.

Significance

This development of effective chemical processes in ancient times demonstrates sophisticated understanding of chemistry principles, even without formal scientific knowledge. Similar trial-and-error chemical processes were developed by many cultures throughout history, including methods for extracting metals and creating alloys like copper, bronze, and iron.

Summary