Welcome to Inorganic Chemistry!
Hello future scientists! This chapter, Inorganic Chemistry, is all about the chemicals that do not primarily involve carbon chains (like we see in organic chemistry). It covers incredibly important topics—from the air we breathe to the materials we build with, and even how we deal with pollution.
Don't worry if some terms look intimidating! We will break everything down into easy steps, use helpful analogies, and focus only on the essential concepts needed for your Science (Single Award) exam. Let's dive in!
1. The Atmosphere and Environmental Chemistry
1.1 The Composition of Clean Air
The air surrounding Earth is a mixture of gases. It’s important to know the approximate percentages of the main components:
- Nitrogen (\(\text{N}_2\)): About 78% (The largest component)
- Oxygen (\(\text{O}_2\)): About 21% (Essential for respiration and combustion)
- Argon and other Noble Gases: About 0.9%
- Carbon Dioxide (\(\text{CO}_2\)): About 0.04%
Quick Memory Tip: Think 78, 21, 1 (approximately 78% N, 21% O, and the remaining 1% is everything else).
1.2 Air Pollution and its Sources
Pollutants are harmful substances added to the atmosphere, usually from burning fuels (combustion) in vehicles or power stations.
Key Pollutants and Their Effects:
- Carbon Monoxide (\(\text{CO}\)):
Source: Incomplete combustion of carbon-based fuels (when there isn't enough oxygen).
Effect: It is a highly toxic, colourless, and odourless gas. It binds to haemoglobin in the blood, stopping it from carrying oxygen, leading to suffocation.
- Sulfur Dioxide (\(\text{SO}_2\)):
Source: Burning fossil fuels (especially coal) that contain sulfur impurities.
Effect: Causes respiratory problems and, crucially, contributes to acid rain.
- Oxides of Nitrogen (\(\text{NO}_x\), e.g., \(\text{NO}\) and \(\text{NO}_2\)):
Source: High temperatures inside vehicle engines cause nitrogen and oxygen from the air to react.
Effect: Causes respiratory problems and also contributes to acid rain.
What is Acid Rain?
When \(\text{SO}_2\) and \(\text{NO}_x\) gases dissolve in rainwater, they form dilute acids (sulfuric acid and nitric acid). Acid rain harms plants, damages buildings and statues made of limestone, and kills fish in lakes.
1.3 The Greenhouse Effect and Climate Change
Certain gases in the atmosphere trap heat radiated from the Earth's surface. This natural process is called the Greenhouse Effect, and it keeps the Earth warm enough to support life.
- Key Greenhouse Gases: Carbon Dioxide (\(\text{CO}_2\)), Methane (\(\text{CH}_4\)), and Water Vapour.
The Problem: Enhanced Greenhouse Effect
Human activities (like burning fossil fuels, deforestation, and farming) increase the concentration of greenhouse gases, especially \(\text{CO}_2\). This traps too much heat, leading to global warming and climate change.
Analogy: Imagine the Earth is wearing a blanket (the greenhouse gases). A normal blanket keeps you comfortable. But burning too much fossil fuel means you are adding more and more blankets, making the Earth dangerously hot.
Key Takeaway (Atmosphere): Combustion creates pollutants (\(\text{CO}\), \(\text{SO}_2\), \(\text{NO}_x\)). Increased \(\text{CO}_2\) is the main cause of the enhanced Greenhouse Effect.
2. Acids, Bases, and Salts
2.1 Defining Acids and Bases
Acids and bases are everywhere, from the vinegar in your kitchen to the antacids you take for a stomach ache.
- Acids: Substances that produce hydrogen ions (\(\text{H}^+\)) when dissolved in water. They have a pH less than 7. (Taste sour, highly corrosive.)
- Bases: Substances that neutralise acids. They accept \(\text{H}^+\) ions. Many bases are insoluble.
- Alkalis: Soluble bases. They produce hydroxide ions (\(\text{OH}^-\)) when dissolved in water. They have a pH greater than 7.
- Neutral Substances: Have a pH of exactly 7 (e.g., pure water).
2.2 The pH Scale and Indicators
The pH Scale runs from 0 to 14 and tells us how acidic or alkaline a solution is.
- pH 0–3: Strong Acid
- pH 4–6: Weak Acid
- pH 7: Neutral
- pH 8–10: Weak Alkali
- pH 11–14: Strong Alkali
Indicators are substances used to show the pH of a solution by changing colour:
- Litmus Paper: Red in acid, Blue in alkali.
- Universal Indicator (UI): Gives a whole range of colours corresponding to the numerical pH value (best for general use).
Common Mistake to Avoid: Students often mix up "Base" and "Alkali." Remember: All alkalis are bases, but not all bases are alkalis (because some bases are insoluble).
2.3 Neutralisation Reactions
Neutralisation is the reaction between an acid and a base (or alkali) to form salt and water.
The general equation is:
\( \text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water} \)
Example: Hydrochloric Acid and Sodium Hydroxide:
\( \text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l) \)
2.4 Reactions of Acids (Making Salts)
You need to know three key reaction types involving acids to make salts:
1. Acid + Metal
\( \text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen gas} \)
Mnemonic: MASH (Metal Acid Salt Hydrogen).
Test for Hydrogen: Collect the gas and hold a lighted splint near it. You hear a distinctive squeaky pop.
Note: Very unreactive metals (like gold) or very reactive metals (like sodium) are not usually used for this reaction.
2. Acid + Metal Carbonate
\( \text{Acid} + \text{Metal Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide gas} \)
Test for Carbon Dioxide: Bubble the gas through limewater (calcium hydroxide solution). The limewater turns cloudy (milky).
3. Acid + Metal Oxide or Hydroxide (Base)
This is the standard neutralisation reaction:
\( \text{Acid} + \text{Metal Oxide/Hydroxide} \rightarrow \text{Salt} + \text{Water} \)
Key Takeaway (Acids/Bases): pH determines acidity. Neutralisation always forms a salt and water. Remember the three salt-making reactions and how to test for the gases produced.
3. Metals and the Reactivity Series
3.1 Understanding the Reactivity Series
The Reactivity Series is a list of metals ordered by their tendency to react. The higher up the series, the more reactive the metal is.
Key Series Segment (Memorise this order!):
Most Reactive: Potassium, Sodium, Calcium, Magnesium, Aluminium, (Carbon), Zinc, Iron, Lead, (Hydrogen), Copper, Silver, Gold. (Least Reactive)
Why are Carbon and Hydrogen included? They are non-metals but are included to show which metals can be extracted using carbon and which can react with acid (above hydrogen).
Mnemonic for remembering the order:
Please Stop Calling Me A Cute Zebra In Low Heels Checking Silver Gold.
3.2 Displacement Reactions
A more reactive metal can displace (kick out) a less reactive metal from a solution of its salt.
Example: Magnesium is more reactive than Copper. If you put magnesium metal into a blue solution of copper sulfate:
\( \text{Mg}(s) + \text{CuSO}_4(aq) \rightarrow \text{MgSO}_4(aq) + \text{Cu}(s) \)
The magnesium displaces the copper, and brown copper metal is deposited.
3.3 Extraction of Metals
Most metals are found naturally as compounds in rocks called ores.
The method used to extract a metal depends on its position in the reactivity series:
1. Extraction using Heating with Carbon (Reduction)
Metals below Carbon (e.g., Zinc, Iron, Copper) can be extracted by heating their oxides with carbon (coke or charcoal).
The carbon is more reactive than these metals, so it takes the oxygen away (a process called reduction).
Example (Iron extraction):
\( \text{Iron oxide} + \text{Carbon} \rightarrow \text{Iron} + \text{Carbon Dioxide} \)
2. Extraction using Electrolysis
Metals above Carbon (e.g., Potassium, Sodium, Aluminium) are very reactive. They cannot be extracted using carbon because carbon is not reactive enough to displace them.
These metals must be extracted using electrolysis (passing an electric current through the molten ore).
Did You Know? Electrolysis is extremely energy intensive and therefore very expensive, which is why Aluminium is more costly to produce than Iron.
Key Takeaway (Reactivity): Reactivity dictates the extraction method. Metals below Carbon are extracted by reduction using Carbon; metals above Carbon require energy-intensive Electrolysis.
4. Limestone and Carbonates
4.1 What is Limestone?
Limestone is a naturally occurring rock, mainly composed of calcium carbonate (\(\text{CaCO}_3\)). It has many uses, including building materials, neutralising acid soil, and making cement.
4.2 Thermal Decomposition of Calcium Carbonate
When limestone (calcium carbonate) is heated strongly, it breaks down into two simpler substances. This process is called thermal decomposition (decomposition means breaking down; thermal means heat).
The reaction produces calcium oxide (known as quicklime) and carbon dioxide gas:
\( \text{Calcium carbonate} \xrightarrow{\text{heat}} \text{Calcium oxide} + \text{Carbon dioxide} \)
\( \text{CaCO}_3(s) \xrightarrow{\Delta} \text{CaO}(s) + \text{CO}_2(g) \)
The quicklime (\(\text{CaO}\)) produced is useful in its own right, especially in treating acidic soil or sewage sludge.
4.3 Reactions of Calcium Oxide and Hydroxide
Calcium oxide (\(\text{CaO}\)) reacts readily with water in a highly exothermic reaction (gives out a lot of heat) to form calcium hydroxide (\(\text{Ca(OH)}_2\)), also known as slaked lime:
\( \text{CaO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(s) \)
Calcium hydroxide is an important alkali used to neutralise acidic soils in agriculture. A solution of calcium hydroxide in water is called limewater, which we use to test for carbon dioxide.
Key Takeaway (Limestone): Limestone is calcium carbonate. Heating it (thermal decomposition) produces quicklime (\(\text{CaO}\)) and \(\text{CO}_2\). Quicklime reacts with water to form limewater, a key alkali.
Quick Review: Inorganic Chemistry Essentials
- Air: 78% N, 21% O. Pollutants like \(\text{SO}_2\) and \(\text{NO}_x\) cause acid rain.
- Greenhouse Gases: Primarily \(\text{CO}_2\), leading to climate change.
- Acids/Bases: pH < 7 (acid), pH > 7 (alkali). Neutralisation makes Salt + Water.
- Metals: Use Mnemonic for Reactivity Series. Extraction uses Carbon for low-reactive metals and Electrolysis for high-reactive metals.
- Limestone: \(\text{CaCO}_3\) thermally decomposes into \(\text{CaO}\) and \(\text{CO}_2\).
Keep practising those reaction equations and you'll master this chapter! You’ve got this!