Periodicity and Group 2: The Alkaline Earth Metals

Periodic trends across period 3 are all about nuclear charge increasing whilst electrons stay in the same shell. This makes atomic radius decrease and first ionisation energy increase as you go from sodium to chlorine.

Melting points follow a clear pattern: metals like sodium and magnesium have giant metallic structures with strong bonding, silicon has a giant covalent structure $think diamond-like$, whilst phosphorus to argon are simple molecules held together by weak van der Waals forces.

Group 2 metals (magnesium to barium) are called alkaline earth metals because their oxides and hydroxides are alkaline. Going down the group, atomic radius increases due to more electron shells, first ionisation energy decreases due to increased shielding, and melting points decrease as metallic bonding weakens.

Key Exam Tip: Remember that beryllium isn't typical of Group 2, so it's often excluded from trend questions. Focus on Mg-Ba patterns.

Group 2 Reactions and Applications

Group 2 reactivity with water increases dramatically down the group. Magnesium barely reacts with cold water but produces magnesium oxide with steam. Calcium, strontium and barium react more vigorously with cold water, producing hydroxides and fizzing more violently as you go down.

The solubility trends are crucial to remember: hydroxides become more soluble down the group (Mg(OH)₂ is sparingly soluble, Ba(OH)₂ dissolves easily), whilst sulfates become less soluble (BaSO₄ is completely insoluble).

Practical applications are everywhere: magnesium extracts titanium metal because it's reactive enough to reduce TiCl₄, calcium compounds remove SO₂ from power station emissions, and barium sulfate is used in medical X-rays because it's completely insoluble and safe.

Real-World Connection: That "barium meal" patients drink before X-rays? It's BaSO₄ - perfectly safe because it won't dissolve in your body!

Group 7: The Halogens and Their Reactions

Halogen trends are the opposite of Group 2 in many ways. Electronegativity decreases down the group as atoms get bigger, boiling points increase due to stronger van der Waals forces, and oxidising ability decreases (fluorine is the most powerful oxidising agent).

Displacement reactions show these trends perfectly - chlorine displaces bromine and iodine from their compounds, bromine displaces iodine, but you can't test fluorine in water because it reacts too violently.

The reactions with concentrated sulfuric acid reveal reducing power: fluorine and chlorine just have acid-base reactions, bromine produces orange fumes and SO₂, whilst iodine goes further - producing purple iodine vapour, yellow sulfur, and even smelly hydrogen sulfide gas.

Testing halides uses silver nitrate: white precipitate for chloride, cream for bromide, pale yellow for iodide. If colours look similar, try ammonia - AgCl dissolves in dilute NH₃, AgBr needs concentrated NH₃, and AgI won't dissolve at all.

Lab Safety Note: These halogen reactions produce toxic gases - always work in a fume cupboard and know your observation colours for exams!

Period 3 Oxides and Their Behaviour

Period 3 oxides show a clear trend from basic to acidic as you move from metals to non-metals. Sodium and magnesium oxides are basic - they react with acids and dissolve in water to form alkaline solutions with pH values around 9-14.

Aluminium oxide is special because it's amphoteric - it reacts with both acids AND bases. This makes it incredibly useful industrially. Silicon dioxide is largely unreactive due to its giant covalent structure.

Non-metal oxides like P₄O₁₀, SO₂ and SO₃ are acidic oxides. They react vigorously with water to form strong acids (phosphoric, sulfurous, and sulfuric acids respectively) with very low pH values (0-3).

The melting point pattern reflects structure: ionic oxides have high melting points, SiO₂ has an extremely high melting point due to giant covalent bonding, whilst simple molecular oxides like SO₂ have low melting points.

Exam Strategy: Learn the pH values and whether oxides dissolve in water - this comes up frequently in multiple choice questions!

Transition Metals: Complex Ions and Ligands

Transition metals are defined as elements forming at least one stable ion with a partially filled d sub-shell. This excludes scandium and zinc from being "true" transition metals, despite their position in the periodic table.

Complex ions form when ligands (molecules or ions with lone pairs) form coordinate bonds with central metal ions. Coordination number tells you how many coordinate bonds surround the metal - usually 4 or 6.

Ligand exchange reactions are fascinating: water ligands can be replaced by ammonia without changing coordination number (both are similar sized and neutral), but chloride ligands are larger and often change coordination from 6 to 4.

Bidentate ligands like ethane-1,2-diamine can form two coordinate bonds per molecule, making complexes more stable. They "take up 1.5 seats" around the metal ion because they're bulkier than monodentate ligands.

Visual Learning Tip: Draw out these complex ions - seeing the 3D arrangements helps you understand why coordination numbers change with different ligands!

Transition Metal Colours and Catalysis

Colour in transition metals comes from d-electrons jumping between energy levels when they absorb specific wavelengths of visible light. The colour you see is what's left after certain wavelengths are absorbed.

Variable oxidation states make transition metals excellent catalysts. Vanadium shows +2, +3, +4, and +5 states with distinctive colours - perfect for demonstrating redox reactions in the lab.

Catalysis works in two ways: heterogeneous catalysts (like iron in the Haber process) are in a different phase from reactants, whilst homogeneous catalysts $like Fe²⁺/Fe³⁺ ions$ are in the same phase and work through intermediate species.

Redox titrations with transition metals are self-indicating. MnO₄⁻ (purple) turning colourless at the endpoint means no indicator needed - the manganese ion itself shows when the reaction's complete.

Lab Connection: That purple potassium manganate(VII) you use in titrations? It's the perfect example of transition metal colour and redox chemistry working together!

Transition Metal Ion Reactions

Aqueous metal ions behave as weak acids because water ligands can lose H⁺ ions. When you add NaOH dropwise, you get coloured precipitates: green for Fe²⁺, brown for Fe³⁺, blue for Cu²⁺, and white for Al³⁺.

Excess NaOH causes different behaviour: iron and copper precipitates remain unchanged, but aluminium dissolves to form a colourless solution because Al(OH)₃ is amphoteric.

Ammonia addition also gives precipitates initially, but excess NH₃ dissolves the copper precipitate to form the stunning deep blue $Cu(NH₃)₄(H₂O)₂$²⁺ complex - one of chemistry's most beautiful colour changes.

Carbonate reactions are trickier because CO₃²⁻ ions react with water around highly charged ions like Fe³⁺ and Al³⁺, producing bubbles of CO₂ and hydroxide precipitates instead of simple carbonates.

Practical Exam Focus: These precipitation reactions are perfect for identifying unknown metal ions - learn the colours and what happens with excess reagents!