Inorganic Chemistry: The Building Blocks of the Non-Living World

Welcome to the exciting world of Inorganic Chemistry! Don't worry if the name sounds complicated; it simply means the chemistry of substances that are not primarily based on carbon (that’s Organic Chemistry, which you'll study separately!).

In this chapter, we will master the behavior of key elements, understand how metals are extracted from the Earth, and learn the essential reactions involving acids, bases, and salts. These concepts are fundamental to many industrial and biological processes, so understanding them is crucial for your success in Science Double Award!

⭐ Quick Review: Core Concepts ⭐

Remember that the properties of an element are determined by its number of outer shell electrons. The groups in the Periodic Table tell us the number of these electrons.

Section 1: The Periodic Table – Key Groups

We focus on three main groups that show clear trends in their chemical and physical properties: Group 1, Group 7, and Group 0.

1.1 Group 1: The Alkali Metals

Group 1 elements (Lithium, Sodium, Potassium, etc.) are highly reactive metals because they all have one electron in their outer shell, which they desperately want to lose to become stable (\(+1\) ion).

  • Physical Properties: They are soft (can be cut with a knife!), have low density, and low melting points compared to typical metals.
  • Chemical Properties (Reactivity): They react vigorously with water, releasing hydrogen gas and forming an alkaline solution (metal hydroxide).

The Trend: Reactivity increases down Group 1 (K > Na > Li).

Why? As you move down the group, the outer electron is further away from the positively charged nucleus. This makes it easier for the atom to lose that electron, meaning it reacts more readily.

Common Mistake to Avoid: Students often think metals get less reactive as you go down the group, like non-metals. Remember: Group 1 metals want to lose electrons, so easier loss = higher reactivity!

1.2 Group 7: The Halogens

Group 7 elements (Fluorine, Chlorine, Bromine, Iodine) are non-metals. They all have seven electrons in their outer shell, which means they easily gain one electron to form a stable \(-1\) ion.

  • Physical Properties: They exist as diatomic molecules (e.g., \(Cl_2\), \(Br_2\)). They are often coloured and their state changes as you move down the group (Gas \(\rightarrow\) Liquid \(\rightarrow\) Solid).

The Trend: Reactivity decreases down Group 7 (F > Cl > Br > I).

Why? These elements want to gain an electron. As you move down the group, the incoming electron is further from the nucleus and shielded by more electron shells, making the attraction weaker. Therefore, it is harder to gain the electron, and reactivity decreases.

Displacement Reactions in Group 7

A more reactive halogen can displace a less reactive halogen from a solution of its salt.

Analogy: The strongest person always gets the seat! Chlorine (\(Cl_2\)) is stronger than Bromine (\(Br_2\)), so:

$$ \text{Chlorine} + \text{Potassium Bromide} \rightarrow \text{Potassium Chloride} + \text{Bromine} $$ $$ Cl_2 (aq) + 2KBr (aq) \rightarrow 2KCl (aq) + Br_2 (aq) $$

1.3 Group 0: The Noble Gases

Group 0 elements (Helium, Neon, Argon, etc.) are characterized by having a full outer shell of electrons (8 electrons, or 2 for Helium).

  • Key Property: They are inert (unreactive) because they do not need to gain, lose, or share electrons to be stable.
  • Uses: Argon is used in light bulbs (to stop the filament from reacting). Helium is used in balloons (because it is very light and safe).

Key Takeaway for Section 1: Groups 1 and 7 are reactive because they seek stability. Group 0 is stable and unreactive. Remember the opposite trends: Group 1 reactivity increases down the group; Group 7 reactivity decreases down the group.

Section 2: Metals, Reactivity, and Extraction

Understanding how reactive a metal is tells us how we can use it and, crucially, how we must extract it from its ore.

2.1 The Reactivity Series

The reactivity series lists metals in order of how easily they lose electrons (i.e., how reactive they are).

Mnemonic Aid (Must Know!):

Please Stop Calling Me A Careless Zebra Instead Try Learning How Copper Saves Gold

K (Potassium)
Na (Sodium)
Ca (Calcium)
Mg (Magnesium)
Al (Aluminium)
(C) Carbon (Reference Point)
Zn (Zinc)
Fe (Iron)
Sn (Tin)
Pb (Lead)
(H) Hydrogen (Reference Point)
Cu (Copper)
Ag (Silver)
Au (Gold)

Metals above Carbon are very reactive. Metals below Hydrogen do not react with dilute acids.

2.2 Metal Displacement Reactions

A more reactive metal will displace (replace) a less reactive metal from a solution of its salt. This is because the more reactive metal has a greater tendency to form positive ions.

Example: If you put Iron (Fe) into a Copper Sulfate (\(CuSO_4\)) solution:

Since Iron is higher in the series than Copper (Fe > Cu), Iron will displace the Copper: $$ Fe (s) + CuSO_4 (aq) \rightarrow FeSO_4 (aq) + Cu (s) $$

The blue colour of the copper sulfate solution fades as the iron forms iron sulfate, and solid copper coats the iron strip.

2.3 Methods of Metal Extraction

Metals exist naturally in ores, often combined with oxygen (oxides). Extracting them involves reduction (removing the oxygen). The method used depends entirely on the metal's reactivity.

1. Very Reactive Metals (K, Na, Ca, Mg, Al – Above Carbon):

  • These metals hold onto oxygen very strongly.
  • They must be extracted by electrolysis (using electricity to split the molten compound). This method is very expensive due to high energy costs.

2. Moderately Reactive Metals (Zn, Fe, Sn, Pb – Below Carbon, Above Hydrogen):

  • These metals can be reduced by heating their oxide ore with Carbon (usually in the form of Coke or Charcoal). Carbon acts as the reducing agent because it is more reactive than these metals.
  • $$ \text{Iron Oxide} + \text{Carbon} \rightarrow \text{Iron} + \text{Carbon Dioxide} $$

3. Low Reactivity Metals (Cu, Ag, Au – Below Hydrogen):

  • These metals are often found naturally in their uncombined form ('native state').
  • If they are found as oxides, they can often be extracted just by heating the ore alone.

Key Takeaway for Section 2: The higher the metal in the reactivity series, the harder and more expensive it is to extract (using electrolysis). Carbon is only effective for extracting metals less reactive than itself.

Section 3: Acids, Bases, and Salts

Acids and bases are fundamental chemical opposites. Their reaction—neutralization—is crucial for making salts.

3.1 Definitions and the pH Scale
  • Acids: Substances that produce hydrogen ions (\(H^+\)) when dissolved in water. (e.g., Hydrochloric Acid, \(HCl\)).
  • Bases: Substances that neutralize acids. Most bases are metal oxides or metal hydroxides.
  • Alkalis: Bases that are soluble in water. They produce hydroxide ions (\(OH^-\)) when dissolved in water. (e.g., Sodium Hydroxide, \(NaOH\)).

pH Scale: Measures the concentration of \(H^+\) ions.

  • pH 0–6: Acidic (lots of \(H^+\))
  • pH 7: Neutral (equal \(H^+\) and \(OH^-\))
  • pH 8–14: Alkaline (lots of \(OH^-\))
3.2 Neutralisation and Salt Formation

Neutralisation is the reaction between an acid and a base (or alkali) to form a salt and water.

The general reaction: $$ \text{Acid} + \text{Base/Alkali} \rightarrow \text{Salt} + \text{Water} $$

The salt produced takes its first name from the metal in the base/alkali and its second name from the acid:

  • Hydrochloric Acid (\(HCl\)) makes Chlorides.
  • Sulfuric Acid (\(H_2SO_4\)) makes Sulfates.
  • Nitric Acid (\(HNO_3\)) makes Nitrates.

The Ionic Equation for Neutralisation: $$ H^+ (aq) + OH^- (aq) \rightarrow H_2O (l) $$

Did you know? The tingling sensation you feel when stung by a wasp (alkaline venom) or a bee (acidic venom) can be neutralized by applying vinegar (acid) or baking soda (alkali) respectively!

3.3 The Three Key Reactions for Salt Preparation

You need to know three ways to produce a salt, depending on what type of substance you react the acid with:

Reaction 1: Acid + Metal

(Only works if the metal is reactive enough, but below Carbon, e.g., Zn, Fe) $$ \text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen Gas} $$ Example: \(Zn (s) + H_2SO_4 (aq) \rightarrow ZnSO_4 (aq) + H_2 (g)\)

Reaction 2: Acid + Base (Metal Oxide or Hydroxide)

$$ \text{Acid} + \text{Metal Oxide} \rightarrow \text{Salt} + \text{Water} $$ Example: \(CuO (s) + 2HCl (aq) \rightarrow CuCl_2 (aq) + H_2O (l)\)

Reaction 3: Acid + Carbonate

$$ \text{Acid} + \text{Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide Gas} $$ Example: \(MgCO_3 (s) + 2HNO_3 (aq) \rightarrow Mg(NO_3)_2 (aq) + H_2O (l) + CO_2 (g)\)

3.4 Preparing a Pure, Dry Sample of Soluble Salt

For the International GCSE, you must know the practical method for preparing a soluble salt using the Acid + Base method (Reaction 2), typically using an insoluble metal oxide and an acid.

Step-by-Step Method (Using copper oxide and sulfuric acid to make copper sulfate):

  1. Add Excess: Heat dilute sulfuric acid. Add the insoluble copper oxide (the base) to the hot acid, stirring constantly. Add excess copper oxide until no more reacts (the solid settles at the bottom).
    (Why excess? To ensure ALL the acid is neutralized. This is crucial!)
  2. Filter: Filter the mixture while it is still warm to remove the excess, unreacted copper oxide. The filtrate (the liquid that passes through) is pure copper sulfate solution.
  3. Evaporate: Heat the copper sulfate solution gently to evaporate about half of the water. Stop heating when crystals just begin to form around the edge of the solution (the crystallisation point).
  4. Crystallise: Place the concentrated solution in a cool place (like a windowsill) and allow it to cool slowly. Pure, hydrated copper sulfate crystals will form.
  5. Dry: Gently dab the crystals dry between sheets of filter paper or tissue.

Key Takeaway for Section 3: The product of acid-base neutralisation is a salt and water. The practical preparation of salts relies on using excess insoluble reactant, filtration, and crystallization.