Question

Draw Lewis structures for the following compounds, and predict which member of each pair will form \(\mathrm{H}\) bonds: (a) \(\mathrm{NH}_{3}\) or \(\mathrm{AsH}_{3}\) (b) \(\mathrm{CH}_{4}\) or \(\mathrm{H}_{2} \mathrm{O}\)

Short answer

NH\textsubscript{3} and H\textsubscript{2}O will form hydrogen bonds.

Step-by-step solution

- Drawing Lewis Structure for \(\text{NH}_3\)

Count the valence electrons for nitrogen (N) and hydrogen (H). Nitrogen has 5 valence electrons and each hydrogen has 1 valence electron. Draw the Lewis structure with nitrogen in the center and three hydrogen atoms connected to it by single bonds.

- Drawing Lewis Structure for \(\text{AsH}_3\)

Count the valence electrons for arsenic (As) and hydrogen (H). Arsenic has 5 valence electrons and each hydrogen has 1 valence electron. Draw the Lewis structure with arsenic in the center and three hydrogen atoms connected to it by single bonds.

- Predicting Hydrogen Bond Formation for \(\text{NH}_3\) and \(\text{AsH}_3\)

Hydrogen bonds generally form when hydrogen is bonded to highly electronegative atoms such as nitrogen, oxygen, or fluorine. Since nitrogen is more electronegative than arsenic, \(\text{NH}_3\) has the capacity to form hydrogen bonds, while \(\text{AsH}_3\) does not.

- Drawing Lewis Structure for \(\text{CH}_4\)

Count the valence electrons for carbon (C) and hydrogen (H). Carbon has 4 valence electrons and each hydrogen has 1 valence electron. Draw the Lewis structure with carbon in the center and four hydrogen atoms connected to it by single bonds.

- Drawing Lewis Structure for \(\text{H}_2 \text{O}\)

Count the valence electrons for oxygen (O) and hydrogen (H). Oxygen has 6 valence electrons and each hydrogen has 1 valence electron. Draw the Lewis structure with oxygen in the center, two lone pairs on the oxygen, and two hydrogen atoms connected to it by single bonds.

- Predicting Hydrogen Bond Formation for \(\text{CH}_4\) and \(\text{H}_2 \text{O}\)

Hydrogen bonds are expected when hydrogen is bonded to highly electronegative atoms like oxygen. Since oxygen is very electronegative, \(\text{H}_2 \text{O}\) can form hydrogen bonds, while \(\text{CH}_4\) cannot.

Key concepts

Lewis structures

Lewis structures are diagrams that show the bonding between atoms of a molecule and any lone pairs of electrons that may exist.
They are vital for visualizing the arrangement of valence electrons around atoms and to predict the geometry, polarity, and reactivity of the molecule.
For example, when drawing the Lewis structure of \(NH_3\), nitrogen is placed in the center with its three valence electrons making single bonds with three hydrogen atoms.
Additionally, the nitrogen atom has a lone pair of electrons.
This representation helps to understand the molecule's capability to engage in hydrogen bonding due to the lone pair and the polarity of the bonds.

Valence electrons

Valence electrons are the electrons in the outermost shell of an atom that are available to form bonds.
These electrons determine an atom's chemical properties and reactivity.

• Nitrogen has 5 valence electrons.
• Arsenic has 5 valence electrons.
• Carbon has 4 valence electrons.
• Oxygen has 6 valence electrons.

This information is crucial for drawing Lewis structures accurately.
For instance, \(NH_3\) and \(AsH_3\) both have central atoms with 5 valence electrons, allowing each to form three single bonds with hydrogen atoms.
Counting valence electrons accurately ensures that molecules are drawn to obey the octet rule, providing stability to the structure.

Hydrogen bonding

Hydrogen bonding is a type of strong intermolecular force that occurs when hydrogen is bonded to highly electronegative atoms like nitrogen, oxygen, or fluorine.
These bonds significantly influence the physical properties of compounds, such as boiling and melting points.
For example, \(NH_3\) can form hydrogen bonds because nitrogen is highly electronegative and has a lone pair of electrons.
On the other hand, \(AsH_3\) cannot form hydrogen bonds due to the lower electronegativity of arsenic.
Similarly, \(H_2O\) forms hydrogen bonds as oxygen is highly electronegative and has two lone pairs, while \(CH_4\) cannot, as carbon is not sufficiently electronegative.

Electronegativity

Electronegativity is the tendency of an atom to attract electrons towards itself in a chemical bond.
Highly electronegative atoms such as nitrogen, oxygen, and fluorine are key players in forming hydrogen bonds.
This property determines not only the type of bonds an atom can form but also the overall polarity of the molecule.
For example, nitrogen in \(NH_3\) has a higher electronegativity than arsenic in \(AsH_3\), meaning it attracts bonding electrons more strongly.
This difference explains why \(NH_3\) can form hydrogen bonds while \(AsH_3\) cannot.
In the case of \(H_2O\), oxygen's high electronegativity makes it capable of strong hydrogen bonding, unlike the carbon in \(CH_4\), which lacks sufficient electronegativity.