Problem 35 - Chemistry The Science in Context

Question

How many valence electrons does each of the following species contain? (a) \(\mathrm{BN} ;\) (b) \(\mathrm{HF} ;\) (c) \(\mathrm{OH}^{-} ;\) (d) \(\mathrm{CN}^{-}\).

Step-by-Step Solution

Answer

Answer: (a) 8, (b) 8, (c) 8, and (d) 10.

Step 1: Identify the valence electrons of individual elements in each species

Look up the periodic table to find the valence electrons of the individual elements for each given species. These electrons are found in their respective outermost (valence) shell. For species \(\mathrm{BN}\): - B (Boron) has 3 valence electrons - N (Nitrogen) has 5 valence electrons For species \(\mathrm{HF}\): - H (Hydrogen) has 1 valence electron - F (Fluorine) has 7 valence electrons For species \(\mathrm{OH}^-{}\): - O (Oxygen) has 6 valence electrons - H (Hydrogen) has 1 valence electron For species \(\mathrm{CN}^-{}\): - C (Carbon) has 4 valence electrons - N (Nitrogen) has 5 valence electrons

Step 2: Add/subtract electrons according to charges

For species with charges, we must account for the charge when counting the number of valence electrons. Add or subtract electrons as needed based on the charge. Species \(\mathrm{OH}^{-}\): As there is a negative charge, we need to add 1 electron. So the total number of valence electrons for this species is: \(\mathrm{OH}^{-}: 6 (\mathrm{Oxygen}) + 1 (\mathrm{Hydrogen}) + 1 (\mathrm{extra\;electron}) = 8\) Species \(\mathrm{CN}^{-}\): Similarly, we have a negative charge here as well and need to add 1 electron. The total count of valence electrons is: \(\mathrm{CN}^{-}: 4 (\mathrm{Carbon}) + 5 (\mathrm{Nitrogen}) + 1 (\mathrm{extra\;electron}) = 10\)

Step 3: Calculate valence electrons for each species

Now that we have the valence electrons of individual elements in each species and have accounted for any charges, we can calculate the total valence electrons for each species. (a) Species \(\mathrm{BN}\): Total valence electrons = 3 (Boron) + 5 (Nitrogen) = 8 (b) Species \(\mathrm{HF}\): Total valence electrons = 1 (Hydrogen) + 7 (Fluorine) = 8 (c) Species \(\mathrm{OH}^{-}\): Total valence electrons = 8 (from Step 2) (d) Species \(\mathrm{CN}^{-}\): Total valence electrons = 10 (from Step 2) So, the valence electrons for each species are: (a) 8, (b) 8, (c) 8, and (d) 10.

Key Concepts

Bonding Theories

Bonding theories help us understand how atoms attach to form molecules. These theories are fundamental in predicting the stability and shape of molecules. One of the central ideas is the sharing or transfer of electrons between atoms. This creates a bond, which can be either covalent, where electrons are shared, or ionic, where electrons are transferred from one atom to another.

• Covalent Bonds: Occur when atoms share pairs of electrons. This usually happens between nonmetals.
• Ionic Bonds: Form when electrons are transferred from one atom to another, leading to the attraction between positively and negatively charged ions.

Bonding theories guide us in understanding how atoms approach each other and how their interactions stabilize structural configurations. For example, in our exercise, species like \( ext{BN}\) and \( ext{HF}\) typically involve covalent bonding, where atoms share electrons to achieve a full outer shell.

Lewis Structures

Lewis structures are simple diagrams that represent the valence electron configuration of molecules. They help visualize the electrons involved in bonding and the lone pairs of electrons.

• Each dot represents a valence electron.
• Lines between atoms represent bonds (a pair of shared electrons).
• Lone pairs are shown as dots around individual atoms.

For example, in \( ext{OH}^-\), the Lewis structure shows a single bond between the hydrogen and oxygen, with oxygen having additional electrons to account for the negative charge. This straightforward representation makes it easier to understand the electronic and structural makeup of molecules. Lewis structures are particularly useful when analyzing how valence electrons are distributed within molecules, aiding in predicting molecular geometry and reactivity.

Periodic Table Groups

Periodic table groups categorize elements that share similar chemical properties, primarily because they have the same number of valence electrons in their outer shell. This regular pattern provides insight into the behavior of atoms in chemical reactions.

• Group 1: These elements have 1 valence electron and are very reactive metals.
• Group 17 (Halogens): Contain 7 valence electrons, making them reactive nonmetals.
• Group 18 (Noble Gases): These have 8 valence electrons, resulting in very low reactivity.

In the exercise, boron from \( ext{BN}\) is in Group 13, with 3 valence electrons. Nitrogen, in contrast, is in Group 15 with 5 valence electrons. These groupings help predict how each element will behave when forming bonds, as seen by their tendency to either gain, lose, or share electrons to achieve a full octet.

Molecular Charge Effects

Molecular charge effects refer to how the charge of a molecule influences the number and arrangement of its electrons. Charges can arise from the loss or gain of electrons, resulting in positively or negatively charged molecules (ions).

• Negative Charge: Indicates extra electrons. For example, in \( ext{OH}^-\) and \( ext{CN}^-\), additional electrons are added to account for the negative charge. This affects the total count of valence electrons and can influence reactivity and interaction with other molecules.
• Positive Charge: Means a deficit of electrons relative to the number of protons.

In our exercise, you adjusted the valence electron count to include these effects, ensuring that total electrons match the chemical formula of the ions. Understanding how charges affect a molecule is crucial for predicting how it will interact in chemical reactions, its stability, and its potential to form additional bonds.