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Draw the electronic diagrams for the bonding in O2 and H2O.

Chemistry

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Here are the electronic diagrams showing the bonding in each molecule/compound:

  • Key for diagrams:
    • ∙\bullet represents valence electrons from one atom.
    • ×\times represents valence electrons from another atom.
    • For ionic compounds, ×\times also represents electrons gained from another atom.

Part a· 2 steps

Oxygen (O2\text{O}_2)

Oxygen is a diatomic molecule with a double covalent bond. Each oxygen atom has 6 valence electrons and needs 2 more to achieve a stable octet. They share two pairs of electrons.

  1. Identify valence electrons.

    Each Oxygen atom (O) has 6 valence electrons.
    Total valence electrons = 2×6=122 \times 6 = 12.

  2. Form a double bond and distribute remaining electrons.

    Two oxygen atoms share two pairs of electrons (a double bond). Each oxygen atom then has two lone pairs.

    ×××O=××∙∙O∙××\begin{array}{c} \times \times \\ \times \text{O} \underset{\bullet \bullet}{\stackrel{\times \times}{=}} \text{O} \bullet \\ \times \times \end{array}

    Each oxygen atom has 4 shared electrons and 4 non-shared electrons (2 lone pairs), totaling 8 electrons in its valence shell.
    The bonding in oxygen is a double covalent bond.

Part b· 2 steps

Water (H2O\text{H}_2\text{O})

Water is a covalent molecule where oxygen shares electrons with two hydrogen atoms.

  1. Identify valence electrons.

    Oxygen (O) has 6 valence electrons.
    Each Hydrogen (H) has 1 valence electron.
    Total valence electrons = 6+(2×1)=86 + (2 \times 1) = 8.

  2. Form single bonds and distribute remaining electrons.

    The oxygen atom forms a single covalent bond with each hydrogen atom. The remaining electrons on oxygen form two lone pairs.

    ×H×∙O×∙H∙∙∙∙\begin{array}{c} \times \\ \text{H} \underset{\bullet}{\times} \text{O} \underset{\bullet}{\times} \text{H} \\ \bullet \bullet \\ \bullet \bullet \end{array}

    Each hydrogen atom achieves a stable duet (2 electrons), and the oxygen atom achieves a stable octet (8 electrons: 4 shared, 4 lone pair).
    The bonding in water consists of two single covalent bonds.

Part c· 2 steps

Ammonia (NH3\text{NH}_3)

Ammonia is a covalent molecule where nitrogen shares electrons with three hydrogen atoms.

  1. Identify valence electrons.

    Nitrogen (N) has 5 valence electrons.
    Each Hydrogen (H) has 1 valence electron.
    Total valence electrons = 5+(3×1)=85 + (3 \times 1) = 8.

  2. Form single bonds and distribute remaining electrons.

    The nitrogen atom forms a single covalent bond with each of the three hydrogen atoms. The remaining two electrons on nitrogen form one lone pair.

    ×H×∙N×∙H∙∙×∙H\begin{array}{c} \times \\ \text{H} \underset{\bullet}{\times} \text{N} \underset{\bullet}{\times} \text{H} \\ \bullet \bullet \\ \underset{\bullet}{\times} \\ \text{H} \end{array}

    Each hydrogen atom achieves a stable duet (2 electrons), and the nitrogen atom achieves a stable octet (8 electrons: 6 shared, 2 lone pair).
    The bonding in ammonia consists of three single covalent bonds.

Part d· 2 steps

Aluminium oxide (Al2O3\text{Al}_2\text{O}_3)

Aluminium oxide is an ionic compound formed by the transfer of electrons from aluminium (metal) to oxygen (non-metal).

  1. Identify valence electrons and ion formation.

    Aluminium (Al) is in Group 13, has 3 valence electrons, and tends to lose them to form Al3+\text{Al}^{3+} ions.
    Oxygen (O) is in Group 16, has 6 valence electrons, and tends to gain 2 electrons to form O2−\text{O}^{2-} ions.

  2. Show electron transfer and resulting ions.

    To balance the charges, two aluminium atoms will lose a total of 6 electrons (2×3=62 \times 3 = 6), and three oxygen atoms will gain a total of 6 electrons (3×2=63 \times 2 = 6).

    Initial atoms:

    Al⋅⋅⋅Al⋅⋅⋅\text{Al} \cdot \cdot \cdot \quad \quad \text{Al} \cdot \cdot \cdot ⋅⋅O⋅⋅⋅⋅O⋅⋅⋅⋅O⋅⋅\cdot \cdot \text{O} \cdot \cdot \quad \quad \cdot \cdot \text{O} \cdot \cdot \quad \quad \cdot \cdot \text{O} \cdot \cdot ⋅⋅⋅⋅⋅⋅\quad \quad \cdot \cdot \quad \quad \quad \quad \cdot \cdot \quad \quad \quad \quad \cdot \cdot

    Electron transfer:
    Each Al atom transfers its 3 valence electrons. Each O atom accepts 2 electrons.

    Resulting ions:

    2×[Al]3+3×[×××O×××]2−2 \times [\text{Al}]^{3+} \quad \quad 3 \times \left[ \begin{array}{c} \times \times \\ \times \text{O} \times \\ \times \times \end{array} \right]^{2-}

    The bonding in aluminium oxide is ionic bonding.

Part e· 2 steps

Magnesium sulphide (MgS\text{MgS})

Magnesium sulphide is an ionic compound formed by the transfer of electrons from magnesium (metal) to sulphur (non-metal).

  1. Identify valence electrons and ion formation.

    Magnesium (Mg) is in Group 2, has 2 valence electrons, and tends to lose them to form Mg2+\text{Mg}^{2+} ions.
    Sulphur (S) is in Group 16, has 6 valence electrons, and tends to gain 2 electrons to form S2−\text{S}^{2-} ions.

  2. Show electron transfer and resulting ions.

    One magnesium atom transfers its 2 valence electrons to one sulphur atom.

    Initial atoms:

    Mg⋅⋅⋅⋅S⋅⋅\text{Mg} \cdot \cdot \quad \quad \cdot \cdot \text{S} \cdot \cdot ⋅⋅\quad \quad \cdot \cdot

    Electron transfer:

    Mg⋅⋅+⋅⋅S⋅⋅⟶[Mg]2++[×××S×××]2−\text{Mg} \cdot \cdot + \cdot \cdot \text{S} \cdot \cdot \longrightarrow [\text{Mg}]^{2+} + \left[ \begin{array}{c} \times \times \\ \times \text{S} \times \\ \times \times \end{array} \right]^{2-}

    The bonding in magnesium sulphide is ionic bonding.

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