Revision of Inorganic Chemistry Periodic table, properties of selected elements, their reactions and compounds.

CHEMISTRY S S 3: Revision of Inorganic Chemistry

Topic: Periodic Table, Properties of Selected Elements, their Reactions and Compounds

Welcome, SS3 Chemistry students! This revision note will help you refresh your knowledge of the Periodic Table and the essential chemistry of some key elements. Understanding these concepts is fundamental to inorganic chemistry.

Core Concepts

1. The Periodic Table: An Organised System

The Periodic Table is a systematic arrangement of elements based on their atomic number, electron configuration, and recurring chemical properties. It helps us predict how elements will behave.

• Periods: Horizontal rows (1-7). Elements in the same period have the same number of electron shells.
• Groups: Vertical columns (1-18). Elements in the same group have the same number of valence electrons and thus exhibit similar chemical properties.
• Blocks: Elements are also classified into s-, p-, d-, and f-blocks based on the subshell being filled with electrons.

Periodic Trends: These are predictable patterns in the properties of elements across periods and down groups.

• Atomic Radius: Generally decreases across a period (due to increasing nuclear charge pulling electrons closer) and increases down a group (due to increasing number of electron shells).
• Ionisation Energy: The energy required to remove an electron from a gaseous atom. Generally increases across a period and decreases down a group.
• Electronegativity: The ability of an atom to attract electrons in a chemical bond. Generally increases across a period and decreases down a group.
• Metallic Character: Decreases across a period (metals are on the left) and increases down a group.

2. Properties, Reactions, and Compounds of Selected Elements

Let's review some important groups and their characteristic chemistries.

a. Group 1: Alkali Metals (Li, Na, K, etc.) These are highly reactive metals with one valence electron. They readily lose this electron to form +1 ions.

• Properties: Soft, low melting points, highly reactive, strong reducing agents. React vigorously with water, oxygen, and halogens.
• Reactions:
  • With water: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) (produces hydrogen gas and a strong base)
  • With oxygen: 4Na(s) + O₂(g) → 2Na₂O(s) (forms oxides)
  • With halogens: 2Na(s) + Cl₂(g) → 2NaCl(s) (forms ionic halides)
• Compounds: Form soluble hydroxides (strong bases), stable oxides, and ionic salts.

b. Group 2: Alkaline Earth Metals (Be, Mg, Ca, etc.) These are reactive metals with two valence electrons, forming +2 ions.

• Properties: Harder and higher melting points than Group 1, less reactive than Group 1 but still quite reactive. Good reducing agents.
• Reactions: Similar to Group 1, but less vigorous.
  • With water: Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g) (calcium reacts steadily)
  • With oxygen: 2Mg(s) + O₂(g) → 2MgO(s) (forms stable oxides)
• Compounds: Form insoluble hydroxides (e.g., Mg(OH)₂), oxides, carbonates, and sulphates (solubility generally decreases down the group for sulphates).

c. Group 17: Halogens (F, Cl, Br, I) These are highly reactive non-metals with seven valence electrons. They readily gain one electron to form -1 ions.

• Properties: Diatomic molecules (F₂, Cl₂, Br₂, I₂), strong oxidising agents, reactivity decreases down the group.
• Reactions:
  • With metals: 2Na(s) + Cl₂(g) → 2NaCl(s)
  • With hydrogen: H₂(g) + Cl₂(g) → 2HCl(g) (forms hydrogen halides)
  • Displacement: A more reactive halogen can displace a less reactive halogen from its salt solution (e.g., Cl₂ + 2KBr → 2KCl + Br₂).
• Compounds: Form covalent hydrides (e.g., HCl, HBr) and ionic halides with metals.

d. Transition Metals (e.g., Iron, Copper, Zinc) Located in the d-block, these metals exhibit unique properties.

• Properties: High melting points, high densities, variable oxidation states, form coloured compounds, good catalysts, form complex ions.
• Reactions: Often involve redox processes. For example, Iron can exist as Fe²⁺ and Fe³⁺.
  • Fe²⁺ can be oxidised to Fe³⁺: 2Fe²⁺ + Cl₂ → 2Fe³⁺ + 2Cl⁻
  • Fe³⁺ can be reduced to Fe²⁺: 2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂
• Compounds: Form various oxides, hydroxides, and salts which are often coloured due to d-orbital electron transitions. Many are important industrial catalysts.

Real-World Example

The reactivity of Group 17 elements (halogens) is very useful. Chlorine is widely used in water treatment to kill harmful bacteria, making tap water safe to drink. This application relies on chlorine's strong oxidizing ability to disinfect water.

Practice/Reflection Question

Consider the elements Magnesium (Mg) and Chlorine (Cl).

1. In which groups and periods are they located?
2. Describe one key difference in their metallic character and reactivity.
3. Write a balanced chemical equation for their reaction to form a compound.

Home Practice Idea: Create flashcards for each of the discussed groups, listing their general properties, common reactions, and examples of compounds. Include one real-life use for an element from each group.