# 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. ## 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. | Electronegativity Difference ($\Delta EN$) | Bond Type | | :---------------------------------------- | :---------------- | | $> 1.8$ | Ionic | | $0.4 - 1.8$ | Polar Covalent | | $< 0.4$ | Non-polar Covalent | 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: $$ 4\mathrm{Na(s)} + \mathrm{O_2(g)} \rightarrow 2\mathrm{Na_2O(s)} $$ * In an excess of oxygen, it forms pale yellow sodium peroxide: $$ 2\mathrm{Na(s)} + \mathrm{O_2(g)} \rightarrow \mathrm{Na_2O_2(s)} $$ * **Magnesium (Mg):** Burns in oxygen with an intense white flame to produce solid magnesium oxide: $$ 2\mathrm{Mg(s)} + \mathrm{O_2(g)} \rightarrow 2\mathrm{MgO(s)} $$ **b) Reaction with Chlorine** * **Sodium (Na):** Burns in chlorine gas with a bright yellow flame to form sodium chloride: $$ 2\mathrm{Na(s)} + \mathrm{Cl_2(g)} \rightarrow 2\mathrm{NaCl(s)} $$ * **Magnesium (Mg):** Burns in chlorine gas with an intense white flame to form magnesium chloride: $$ \mathrm{Mg(s)} + \mathrm{Cl_2(g)} \rightarrow \mathrm{MgCl_2(s)} $$ **c) Reaction with Water** * **Sodium (Na):** Reacts violently and explosively with cold water to produce sodium hydroxide and hydrogen gas: $$ 2\mathrm{Na(s)} + 2\mathrm{H_2O(l)} \rightarrow 2\mathrm{NaOH(aq)} + \mathrm{H_2(g)} $$ * **Magnesium (Mg):** Reacts very slowly with cold water. However, it burns in steam to form magnesium oxide and hydrogen gas: $$ \mathrm{Mg(s)} + \mathrm{H_2O(g)} \rightarrow \mathrm{MgO(s)} + \mathrm{H_2(g)} $$ ## 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: $\mathrm{Na_2O}$ (+1), $\mathrm{MgO}$ (+2), $\mathrm{Al_2O_3}$ (+3), $\mathrm{SiO_2}$ (+4), $\mathrm{P_4O_{10}}$ (+5), $\mathrm{SO_3}$ (+6), $\mathrm{Cl_2O_7}$ (+7). * Some elements like Phosphorus (P) and Sulphur (S) can show variable oxidation numbers (e.g., S is +4 in $\mathrm{SO_2}$ and +6 in $\mathrm{SO_3}$). **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). | Group | IA | IIA | IIIA | IVA | VA | VIA | VIIA | | ------------------ | -------------------- | --------- | -------------------- | ----------- | --------------------- | ------------- | --------------------- | | **3rd Period** | Na | Mg | Al | Si | P | S | Cl | | **Oxide** | $\mathrm{Na_2O}$ | $\mathrm{MgO}$ | $\mathrm{Al_2O_3}$ | $\mathrm{SiO_2}$ | $\mathrm{P_4O_{10}}$ | $\mathrm{SO_3}$ | $\mathrm{Cl_2O_7}$ | | **Nature of Oxide** | **Strongly Basic** | **Basic** | **Amphoteric** | **Weakly Acidic** | **Acidic** | **Strongly Acidic** | **Very Strongly Acidic** | * **Basic Oxides:** React with acids. $$ \mathrm{Na_2O} + 2\mathrm{HCl} \rightarrow 2\mathrm{NaCl} + \mathrm{H_2O} $$ $$ \mathrm{MgO} + \mathrm{H_2SO_4} \rightarrow \mathrm{MgSO_4} + \mathrm{H_2O} $$ * **Amphoteric Oxides:** React with both acids and bases. Aluminum oxide is a classic example. * As a base: $\mathrm{Al_2O_3} + 6\mathrm{HCl} \rightarrow 2\mathrm{AlCl_3} + 3\mathrm{H_2O}$ * As an acid: $\mathrm{Al_2O_3} + 2\mathrm{NaOH} + 3\mathrm{H_2O} \rightarrow 2\mathrm{Na[Al(OH)_4]}$ (Sodium aluminate) **Reactions of Chlorides with Water** * **Ionic Chlorides** (e.g., $\mathrm{NaCl}$, $\mathrm{MgCl_2}$) dissolve in water to form neutral solutions (pH = 7). * **Covalent Chlorides** (e.g., $\mathrm{AlCl_3}$, $\mathrm{SiCl_4}$, $\mathrm{PCl_5}$) react vigorously with water (hydrolysis) to form strongly acidic solutions (pH << 7). $$ \mathrm{AlCl_3(s)} + 3\mathrm{H_2O(l)} \rightarrow \mathrm{Al(OH)_3(s)} + 3\mathrm{HCl(aq)} \quad (\mathrm{pH}=3) $$ $$ \mathrm{SiCl_4(l)} + 4\mathrm{H_2O(l)} \rightarrow \mathrm{Si(OH)_4(s)} + 4\mathrm{HCl(aq)} \quad (\mathrm{pH}=0) $$ $$ \mathrm{PCl_5(s)} + 4\mathrm{H_2O(l)} \rightarrow \mathrm{H_3PO_4(aq)} + 5\mathrm{HCl(aq)} \quad (\mathrm{pH}=0) $$ ## 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. ## 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 ($\mathrm{Br_2}$). * It combines with metals to form salts (e.g., $\mathrm{NaBr}$). * It forms an acidic oxide. * Its maximum oxidation state is +7.