Draw the Lewis dot structures for H₂S, CH₃Br, and HCN.

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1. Draw the Lewis dot structures for each of the following molecules:

a) H₂S

Calculate total valence electrons.
Hydrogen (Group 1) has 1 valence electron. Sulfur (Group 16) has 6 valence electrons.
Total valence electrons = $(2 \times 1) + 6 = 8$ electrons.
Arrange atoms and form single bonds.
Sulfur is the central atom.
H-S-H (4 electrons used).
Distribute remaining electrons as lone pairs.
Remaining electrons = $8 - 4 = 4$ electrons. Place these on the central sulfur atom.
$\boxed{\text{H} : \underset{..}{\text{S}} : \text{H}}$
b) CH₃Br
Calculate total valence electrons.
Carbon (Group 14) has 4 valence electrons. Hydrogen (Group 1) has 1 valence electron. Bromine (Group 17) has 7 valence electrons.
Total valence electrons = $4 + (3 \times 1) + 7 = 14$ electrons.
Arrange atoms and form single bonds.
Carbon is the central atom.
C-H (3 bonds) and C-Br (1 bond) = 8 electrons used.
Distribute remaining electrons as lone pairs.
Remaining electrons = $14 - 8 = 6$ electrons. Place these on the more electronegative bromine atom.
$\boxed{\begin{matrix} & \text{H} & \\ & | & \\ \text{H} & - & \text{C} & - & \underset{..}{\text{Br}} : \\ & | & & & .. \\ & \text{H} & & & \end{matrix}}$
c) HCN
Calculate total valence electrons.
Hydrogen (Group 1) has 1 valence electron. Carbon (Group 14) has 4 valence electrons. Nitrogen (Group 15) has 5 valence electrons.
Total valence electrons = $1 + 4 + 5 = 10$ electrons.
Arrange atoms and form single bonds.
Carbon is the central atom.
H-C-N (4 electrons used).
Distribute remaining electrons as lone pairs and form multiple bonds to satisfy octets.
Remaining electrons = $10 - 4 = 6$ electrons. Place these on nitrogen.
Carbon needs 4 more electrons, and nitrogen has a full octet with 6 lone pair electrons. Move two lone pairs from nitrogen to form a triple bond with carbon.
$\boxed{\text{H} - \text{C} \equiv \underset{..}{\text{N}}}$
d) K₂O
Identify as an ionic compound (metal K and nonmetal O).
Potassium (Group 1) forms a K⁺ ion. Oxygen (Group 16) forms an O²⁻ ion.
Show the electron transfer and resulting ions with their charges and octets.
Two K atoms each lose 1 electron, and one O atom gains 2 electrons.
$\boxed{\text{K}^{+} \left[ \underset{..}{\overset{..}{\text{O}}} \right]^{2-} \text{K}^{+}}$
e) Al₂O₃
Identify as an ionic compound (metal Al and nonmetal O).
Aluminum (Group 13) forms an Al³⁺ ion. Oxygen (Group 16) forms an O²⁻ ion.
Determine the ratio of ions to balance charges.
To balance the charges, two Al³⁺ ions (total +6) combine with three O²⁻ ions (total -6).
Show the resulting ions with their charges and octets.
$\boxed{2 \text{Al}^{3+} \quad 3 \left[ \underset{..}{\overset{..}{\text{O}}} \right]^{2-}}$
2. In the Lewis structures listed below, M and X represent various elements in the third period of the periodic table. Write the formula of each compound using the chemical symbols of each element:

The third period elements are Na, Mg, Al, Si, P, S, Cl, Ar.

a) $\text{M}^{+} [\text{::X::}]^{-}$
Identify M and X based on their charges and electron configuration.
M has a +1 charge, indicating it is an alkali metal (Group 1). In the third period, this is Na (Sodium).
X has a -1 charge and an octet, indicating it is a halogen (Group 17). In the third period, this is Cl (Chlorine).
Write the chemical formula.
$\boxed{\text{NaCl}}$
b) $\text{M}^{2+} [\text{::X::}]^{2-}$
Identify M and X based on their charges and electron configuration.
M has a +2 charge, indicating it is an alkaline earth metal (Group 2). In the third period, this is Mg (Magnesium).
X has a -2 charge and an octet, indicating it is a chalcogen (Group 16). In the third period, this is S (Sulfur).
Write the chemical formula.
$\boxed{\text{MgS}}$
c) $\text{M}^{3+} [\text{::X::}]^{2-}_{2}$
Identify M and X based on their charges and electron configuration.
M has a +3 charge, indicating it is from Group 13. In the third period, this is Al (Aluminum).
X has a -2 charge and an octet, indicating it is a chalcogen (Group 16). In the third period, this is S (Sulfur).
Write the chemical formula, ensuring charge neutrality.
The given structure implies two X²⁻ ions for one M³⁺ ion, which is incorrect for charge balance. For M³⁺ and X²⁻, the formula should be M₂X₃. Assuming the structure meant to represent the ions that form M₂X₃.
$\boxed{\text{Al}_2\text{S}_3}$
d) $\text{M}^{2+} [\text{::X::}]^{-}_{2}$
Identify M and X based on their charges and electron configuration.
M has a +2 charge, indicating it is an alkaline earth metal (Group 2). In the third period, this is Mg (Magnesium).
X has a -1 charge and an octet, indicating it is a halogen (Group 17). In the third period, this is Cl (Chlorine).
Step 22

Draw the Lewis dot structures for H₂S, CH₃Br, and HCN.

Step-by-step solution

Calculate total valence electrons.

Arrange atoms and form single bonds.

Distribute remaining electrons as lone pairs.

Calculate total valence electrons.

Arrange atoms and form single bonds.

Distribute remaining electrons as lone pairs.

Calculate total valence electrons.

Arrange atoms and form single bonds.

Distribute remaining electrons as lone pairs and form multiple bonds to satisfy octets.

Identify as an ionic compound (metal K and nonmetal O).

Show the electron transfer and resulting ions with their charges and octets.

Identify as an ionic compound (metal Al and nonmetal O).

Determine the ratio of ions to balance charges.

Show the resulting ions with their charges and octets.

Identify M and X based on their charges and electron configuration.

Write the chemical formula.

Identify M and X based on their charges and electron configuration.

Write the chemical formula.

Identify M and X based on their charges and electron configuration.

Write the chemical formula, ensuring charge neutrality.

Identify M and X based on their charges and electron configuration.

Step 22

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