Metals: Elements that tend to lose electrons to form positive ions (cations).
Examples: Iron (Fe), Copper (Cu), Gold (Au), Silver (Ag).
Non-metals: Elements that tend to gain electrons to form negative ions (anions).
Examples: Chlorine (Cl), Sulfur (S), Phosphorus (P).
Metalloids: Elements with properties of both metals and non-metals.
Examples: Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As).
Metals: “Malleable, shiny, conductive” – think of coins, jewelry, wires.
Non-metals: “Brittle, dull, insulating” – think of chalk, sulfur, oxygen.
Metalloids: “Stair-step line” on periodic table from B to Po (B, Si, Ge, As, Sb, Te).
Answer: Metalloids are elements that exhibit properties of both metals and non-metals. They are located along the “stair-step line” on the periodic table, starting at boron (B) and extending down to polonium (Po). Elements to the left of this line are metals, elements to the far right are non-metals, and elements along the line are metalloids.
Examples: Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Polonium (Po).
Period Number: Indicates the number of electron shells (principal quantum number, n).
Example: An element in 3rd period has 3 electron shells.
Group Number: Indicates the number of valence electrons.
Example: Group 1 elements have 1 valence electron, Group 2 have 2, etc.
Electronic Configuration: Distribution of electrons in atomic orbitals.
Example: Magnesium (Mg) in period 3, group 2: 1s² 2s² 2p⁶ 3s²
Period = Shells: Period number tells you how many electron shells an atom has.
Group = Valence Electrons: Group number (for main group elements) tells you how many electrons are in the outermost shell.
Configuration Pattern: Follow the order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.
| Element | Period | Group | Valence Electrons | Electronic Configuration |
|---|---|---|---|---|
| Sodium (Na) | 3 | 1 | 1 | 1s² 2s² 2p⁶ 3s¹ |
| Magnesium (Mg) | 3 | 2 | 2 | 1s² 2s² 2p⁶ 3s² |
| Aluminum (Al) | 3 | 13 | 3 | 1s² 2s² 2p⁶ 3s² 3p¹ |
| Silicon (Si) | 3 | 14 | 4 | 1s² 2s² 2p⁶ 3s² 3p² |
| Phosphorus (P) | 3 | 15 | 5 | 1s² 2s² 2p⁶ 3s² 3p³ |
| Sulfur (S) | 3 | 16 | 6 | 1s² 2s² 2p⁶ 3s² 3p⁴ |
Answer: The element is Magnesium (Mg).
Reasoning: Period 3 means it has 3 electron shells. Group 2 means it has 2 valence electrons.
Electronic Configuration: 1s² 2s² 2p⁶ 3s²
Explanation: The 2 valence electrons are in the 3s subshell (since it’s in the s-block).
Modern Periodic Law: Properties of elements are periodic functions of their atomic numbers.
Definition: Half the distance between two identical bonded atoms.
Trend across period: Decreases (left to right) due to increasing nuclear charge.
Trend down group: Increases due to additional electron shells.
Definition: Energy required to remove an electron from a gaseous atom.
Trend across period: Increases (left to right).
Trend down group: Decreases.
Definition: Energy change when an electron is added to a gaseous atom.
Trend across period: Becomes more negative (left to right).
Trend down group: Becomes less negative.
Definition: Ability of an atom to attract bonding electrons.
Trend across period: Increases (left to right).
Trend down group: Decreases.
Across Period (→): Atomic radius ↓, Ionization energy ↑, Electron affinity ↑ (more negative), Electronegativity ↑
Down Group (↓): Atomic radius ↑, Ionization energy ↓, Electron affinity ↓ (less negative), Electronegativity ↓
Memory phrase: “As you go RIGHT, things get TIGHT (smaller radius, harder to remove electrons). As you go DOWN, things get LOOSE (larger radius, easier to remove electrons).”
Answer:
Across a period (left to right): Atomic radius decreases because:
Down a group (top to bottom): Atomic radius increases because:
Sodium: Reacts vigorously with cold water
2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
Magnesium: Reacts slowly with cold water, vigorously with steam
Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g)
Mg(s) + H₂O(g) → MgO(s) + H₂(g)
Sodium: Burns with golden yellow flame, forms peroxide
2Na(s) + O₂(g) → Na₂O₂(s) (peroxide)
Magnesium: Burns with intense white flame, forms oxide
2Mg(s) + O₂(g) → 2MgO(s)
Sodium: Reacts exothermically, forms sodium chloride
2Na(s) + Cl₂(g) → 2NaCl(s)
Magnesium: Forms magnesium chloride
Mg(s) + Cl₂(g) → MgCl₂(s)
Sodium is more reactive than magnesium because it has lower ionization energy (easier to lose its single valence electron).
Color memory: Sodium = golden yellow flame, Magnesium = intense white flame.
Reactivity order: Na > Mg (sodium reacts vigorously with cold water, magnesium needs steam for vigorous reaction).
Answer:
Sodium is more reactive than magnesium due to:
With water:
With oxygen:
With chlorine:
Basic Oxides: Metal oxides that react with water to form bases
Examples: Na₂O, MgO (ionic, high melting points)
Acidic Oxides: Non-metal oxides that react with water to form acids
Examples: SO₂, SO₃, P₂O₅ (covalent, lower melting points)
Amphoteric Oxides: React with both acids and bases
Example: Al₂O₃ (aluminum oxide)
Neutral Chlorides: Form neutral solutions in water (pH ≈ 7)
Examples: NaCl, MgCl₂ (ionic, dissociate completely)
Acidic Chlorides: Hydrolyze in water to form acidic solutions
Examples: AlCl₃, SiCl₄, PCl₅ (covalent, undergo hydrolysis)
| Element | Oxide | Nature | Chloride | Nature in Water |
|---|---|---|---|---|
| Sodium (Na) | Na₂O | Basic | NaCl | Neutral (pH=7) |
| Magnesium (Mg) | MgO | Basic | MgCl₂ | Neutral (pH≈6.5) |
| Aluminum (Al) | Al₂O₃ | Amphoteric | AlCl₃ | Acidic (hydrolysis) |
| Silicon (Si) | SiO₂ | Acidic | SiCl₄ | Acidic (hydrolysis) |
| Phosphorus (P) | P₂O₅ | Acidic | PCl₅ | Acidic (hydrolysis) |
| Sulfur (S) | SO₃ | Acidic | SCl₂ | Acidic (hydrolysis) |
Oxidation Number: Formal charge on an atom in a molecule/ion.
Rule: For main group elements, maximum oxidation number = group number.
In Oxides: Increases from +1 (Na) to +6 (S)
In Chlorides: Increases from +1 (Na) to +5 (P)
Variable Oxidation States: Some elements show multiple oxidation states.
Examples:
Phosphorus: +3 (P₂O₃), +5 (P₂O₅)
Sulfur: +4 (SO₂), +6 (SO₃)
| Element | Oxide | Oxidation Number | Chloride | Oxidation Number |
|---|---|---|---|---|
| Sodium (Na) | Na₂O | +1 | NaCl | +1 |
| Magnesium (Mg) | MgO | +2 | MgCl₂ | +2 |
| Aluminum (Al) | Al₂O₃ | +3 | AlCl₃ | +3 |
| Silicon (Si) | SiO₂ | +4 | SiCl₄ | +4 |
| Phosphorus (P) | P₂O₅ | +5 | PCl₅ | +5 |
| Sulfur (S) | SO₃ | +6 | SCl₂ | +2 |
Answer: (C) Electron mass
Explanation: Electron affinity is affected by atomic size, nuclear charge, and electronic configuration, but not by the mass of the electron itself.
Answer: (B) Increases
Explanation: Metallic character increases down Group 1 because atomic size increases, making it easier to lose electrons.
Answer: (A) +1
Explanation: Sodium always has +1 oxidation state in its compounds because it loses its single valence electron.
Answer: (A) Sodium
Explanation: Sodium burns with a characteristic golden yellow flame in oxygen.
Answer: (C) SiCl₄
Explanation: SiCl₄ undergoes hydrolysis in water to form HCl and SiO₂, making the solution acidic.
Answer: (A) Equal to 7
Explanation: Neutral chlorides like NaCl dissociate in water without affecting the pH, resulting in pH ≈ 7.
Answer: (D) +5
Explanation: In PCl₅, phosphorus uses all 5 valence electrons for bonding, giving it an oxidation state of +5.
Answer: (C) Al₂O₃
Explanation: Aluminum oxide (Al₂O₃) reacts with both acids and bases, making it amphoteric.
Answer: (B) It increases from top to bottom in a group.
Explanation: Metallic character increases down a group due to increasing atomic size and decreasing ionization energy.
Answer: (A) Atomic radius
Explanation: Atomic radius increases down a group due to addition of new electron shells.
Answer: First ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous +1 ions.
Example: Na(g) → Na⁺(g) + e⁻ ΔH = +496 kJ/mol
Answer: Phosphorus has half-filled 3p subshell (3p³) which is particularly stable. Removing an electron from this stable configuration requires more energy. Sulfur has one extra electron in the 3p subshell (3p⁴), creating electron-electron repulsion. This makes it slightly easier to remove an electron from sulfur.