10.6 Trends in Metallic and Non-Metallic Behaviours

This section explores the trends in metallic and non-metallic properties of elements across the periodic table, focusing on how valence electrons influence chemical behavior, bonding, and reactivity.

10.6.1 Metallic Behaviour

Metals are located on the left side of the periodic table. Their defining chemical characteristic is the tendency to lose valence electrons to form positive ions (cations).

• Valence Electrons: Metals typically have 1-3 valence electrons.
• Ionization Energy: They have low ionization energies, meaning little energy is required to remove their outermost electrons.
• Electronegativity: Metals have low electronegativity, indicating a weak attraction for electrons.

This combination of factors makes it energetically favorable for metals to lose electrons and achieve a stable electron configuration.

Ionization Energy→

10.6.2 Non-Metallic Behaviour

Non-metals are found on the right side of the periodic table. Their chemical behavior is dominated by their tendency to gain or share electrons to achieve a stable electron configuration.

• Valence Electrons: Non-metals typically have 5-7 valence electrons and need only a few more to complete their valence shell.
• Ionization Energy: They have high ionization energies, making it difficult to remove their electrons.
• Electronegativity: Non-metals have high electronegativity, giving them a strong ability to attract electrons in a chemical bond.

10.6.3 Trends in Groups

The metallic and non-metallic character of elements changes predictably within groups.

• Down a Group: Metallic character increases down a group. As atomic radii increase and shielding effect becomes more pronounced, the outermost electrons are held less tightly and are easier to lose.
  • Group 1 (Alkali Metals): Strong metallic character that increases down the group.
  • Group 17 (Halogens): Strong non-metallic character that decreases down the group.
  • Group 14 (Carbon Family): This group clearly shows the transition from non-metal to metal.
    • Carbon (C): A typical non-metal.
    • Silicon (Si) and Germanium (Ge): Metalloids (properties of both metals and non-metals).
    • Tin (Sn) and Lead (Pb): Typical metals.

10.6.4 Electronegativity and Type of Bond

The difference in electronegativity ($\Delta EN$) between two bonded atoms can predict the type of chemical bond formed.

As the $\Delta EN$ increases, the ionic character of the bond increases.

• Example 1: Sodium Chloride (NaCl) $\Delta EN=EN(\mathrm{Cl})-EN(\mathrm{Na})=3.16-0.93=2.23$ Since $\Delta EN>1.8$, the bond is ionic.

• Example 2: Magnesium Oxide (MgO) $\Delta EN=EN(\mathrm{O})-EN(\mathrm{Mg})=3.44-1.31=2.13$ Since $\Delta EN>1.8$, the bond is ionic.

• Example 3: Carbon-Hydrogen (C-H) Bond $\Delta EN=EN(\mathrm{C})-EN(\mathrm{H})=2.55-2.2=0.35$ Since $\Delta EN<0.4$, the C-H bond is non-polar covalent.

• Example 4: Hydrogen Fluoride (H-F) $\Delta EN=EN(\mathrm{F})-EN(\mathrm{H})=3.98-2.2=1.78$ Since $0.4<\Delta EN<1.8$, the H-F bond is polar covalent.

10.6.5 Trends in Chemical Properties

The reactions of Group 1 (e.g., Sodium) and Group 2 (e.g., Magnesium) metals illustrate typical metallic behavior.

a) Reaction with Oxygen

• Sodium (Na): Reacts readily with oxygen.
  • In a limited supply of oxygen, it forms sodium oxide: $4\mathrm{Na}(\mathrm{s})+{\mathrm{O}}_2(\mathrm{g})\to 2\mathrm{N}{\mathrm{a}}_2\mathrm{O}(\mathrm{s})$
  • In an excess of oxygen, it forms pale yellow sodium peroxide: $2\mathrm{Na}(\mathrm{s})+{\mathrm{O}}_2(\mathrm{g})\to \mathrm{N}{\mathrm{a}}_2{\mathrm{O}}_2(\mathrm{s})$
• Magnesium (Mg): Burns in oxygen with an intense white flame to produce solid magnesium oxide: $2\mathrm{Mg}(\mathrm{s})+{\mathrm{O}}_2(\mathrm{g})\to 2\mathrm{MgO}(\mathrm{s})$

b) Reaction with Chlorine

• Sodium (Na): Burns in chlorine gas with a bright yellow flame to form sodium chloride: $2\mathrm{Na}(\mathrm{s})+\mathrm{C}{\mathrm{l}}_2(\mathrm{g})\to 2\mathrm{NaCl}(\mathrm{s})$
• Magnesium (Mg): Burns in chlorine gas with an intense white flame to form magnesium chloride: $\mathrm{Mg}(\mathrm{s})+\mathrm{C}{\mathrm{l}}_2(\mathrm{g})\to \mathrm{MgC}{\mathrm{l}}_2(\mathrm{s})$

c) Reaction with Water

• Sodium (Na): Reacts violently and explosively with cold water to produce sodium hydroxide and hydrogen gas: $2\mathrm{Na}(\mathrm{s})+2{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\to 2\mathrm{NaOH}(\mathrm{aq})+{\mathrm{H}}_2(\mathrm{g})$
• Magnesium (Mg): Reacts very slowly with cold water. However, it burns in steam to form magnesium oxide and hydrogen gas: $\mathrm{Mg}(\mathrm{s})+{\mathrm{H}}_2\mathrm{O}(\mathrm{g})\to \mathrm{MgO}(\mathrm{s})+{\mathrm{H}}_2(\mathrm{g})$

10.6.6 Variation in Properties of Oxides and Chlorides

Oxidation Number The maximum oxidation number of an element in its oxide or chloride corresponds to its number of valence electrons. Across Period 3 (Na, Mg, Al, Si, P, S, Cl), the maximum oxidation number increases from +1 to +7.

• Oxides: $\mathrm{N}{\mathrm{a}}_2\mathrm{O}$ (+1), $\mathrm{MgO}$ (+2), $\mathrm{A}{\mathrm{l}}_2{\mathrm{O}}_3$ (+3), $\mathrm{Si}{\mathrm{O}}_2$ (+4), ${\mathrm{P}}_4{\mathrm{O}}_{10}$ (+5), $\mathrm{S}{\mathrm{O}}_3$ (+6), $\mathrm{C}{\mathrm{l}}_2{\mathrm{O}}_7$ (+7).
• Some elements like Phosphorus (P) and Sulphur (S) can show variable oxidation numbers (e.g., S is +4 in $\mathrm{S}{\mathrm{O}}_2$ and +6 in $\mathrm{S}{\mathrm{O}}_3$).

Nature of Oxides

• Metal oxides are generally basic, forming bases in water (pH > 7).
• Non-metal oxides are generally acidic, forming acids in water (pH < 7).

• Basic Oxides: React with acids. $\mathrm{N}{\mathrm{a}}_2\mathrm{O}+2\mathrm{HCl}\to 2\mathrm{NaCl}+{\mathrm{H}}_2\mathrm{O}$ $\mathrm{MgO}+{\mathrm{H}}_2\mathrm{S}{\mathrm{O}}_4\to \mathrm{MgS}{\mathrm{O}}_4+{\mathrm{H}}_2\mathrm{O}$
• Amphoteric Oxides: React with both acids and bases. Aluminum oxide is a classic example.
  • As a base: $\mathrm{A}{\mathrm{l}}_2{\mathrm{O}}_3+6\mathrm{HCl}\to 2\mathrm{AlC}{\mathrm{l}}_3+3{\mathrm{H}}_2\mathrm{O}$
  • As an acid: $\mathrm{A}{\mathrm{l}}_2{\mathrm{O}}_3+2\mathrm{NaOH}+3{\mathrm{H}}_2\mathrm{O}\to 2\mathrm{Na}[\mathrm{Al}(\mathrm{OH}{)}_4]$ (Sodium aluminate)

Reactions of Chlorides with Water

• Ionic Chlorides (e.g., $\mathrm{NaCl}$, $\mathrm{MgC}{\mathrm{l}}_2$) dissolve in water to form neutral solutions (pH = 7).
• Covalent Chlorides (e.g., $\mathrm{AlC}{\mathrm{l}}_3$, $\mathrm{SiC}{\mathrm{l}}_4$, $\mathrm{PC}{\mathrm{l}}_5$) react vigorously with water (hydrolysis) to form strongly acidic solutions (pH << 7). $\mathrm{AlC}{\mathrm{l}}_3(\mathrm{s})+3{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\to \mathrm{Al}(\mathrm{OH}{)}_3(\mathrm{s})+3\mathrm{HCl}(\mathrm{aq})(\mathrm{pH}=3)$ $\mathrm{SiC}{\mathrm{l}}_4(\mathrm{l})+4{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\to \mathrm{Si}(\mathrm{OH}{)}_4(\mathrm{s})+4\mathrm{HCl}(\mathrm{aq})(\mathrm{pH}=0)$ $\mathrm{PC}{\mathrm{l}}_5(\mathrm{s})+4{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\to {\mathrm{H}}_3\mathrm{P}{\mathrm{O}}_4(\mathrm{aq})+5\mathrm{HCl}(\mathrm{aq})(\mathrm{pH}=0)$

10.6.9 Trends in Ionization Energies and Electron Affinities

• Ionization Energy (IE) in Group 1: Decreases down the group due to increased atomic radii and shielding effect, making it easier to remove the valence electron.
• Electron Affinity (EA) in Group 17: Generally decreases down the group for the same reasons (increased size and shielding weaken the nucleus's attraction for an incoming electron).
  • Exception: Fluorine has a lower electron affinity than Chlorine. Although F is smaller and more electronegative, its small size leads to significant electron-electron repulsion in its compact valence shell, making it less favorable to add another electron compared to Chlorine.

Halogens→

Worked Examples

Example 1: Identifying an Unknown Element

Suppose you have an unknown element with the following properties:

• Atomic Number: 35
• Goal: Determine its position in the periodic table and predict its properties.

Problem-Solving Strategy

1. Write the electronic configuration.
   • The element has 35 electrons.
   • Configuration: $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^5$
2. Use valence configuration to find its group and period.
   • The valence shell is the outermost shell, where $n=4$.
   • Valence configuration: $4s^{2}4p^5$.
   • The period is given by the principal quantum number of the valence shell, so it is in Period 4.
   • The number of valence electrons is $2+5=7$. This places it in Group 17 (or VIIA).
   • The element is Bromine (Br).

Predicted Properties:

• It is a non-metal.
• It exists as a diatomic molecule ($\mathrm{B}{\mathrm{r}}_2$).
• It combines with metals to form salts (e.g., $\mathrm{NaBr}$).
• It forms an acidic oxide.
• Its maximum oxidation state is +7.