Properties of Covalent Compounds

When you hear the words "chemical compound" what do you think of? Most people would probably talk about man-made drugs or the weird words they can't pronounce in their food's ingredients list. However, pretty much any material this isn't a singular element is made up of chemical compounds.

In this article, we will be talking about a specific type of chemical compound: covalent compounds. We will be discussing what they are, the different types, and their common characteristics.

• This article covers covalent compounds and their properties.
• First, we will define what covalent compounds are.
• Next, we will look the different types of covalent bond.
• Then, we will learn the trends in covalent bond length.
• Thereafter, we will learn some common characteristics of covalent compounds.
• Lastly, we will look at some covalent compounds and their uses.

Covalent Compounds

Before we discuss their properties, let's first discuss what covalent compounds actually are.

As an example, here is a list of some covalent compounds:

• H₂O-Water

• SiO₂-Silicon dioxide (Silicon (Si) is a metalloid)

• NH₃-Ammonia

• F₂-Fluorine

Types of Covalent Bond

There are different types of covalent bond. These "types" can be broken up into two categories: categories based on number and categories based on electronegativity.

Let's break these types down based on category

Types of Covalent Bond: Numbers

There are three types of numbered covalent bonds:

• Single
• Double
• Triple

Numbered covalent bonds depend on two factors: the number of electrons shared and the types of orbital overlap.

In terms of electrons shared, each bond contains 2 electrons. Therefore, double bonds share 4 electrons in total, while triple bonds share six.

And now for orbital overlap:

There are 4 main types of orbitals, these are:

• S-orbitals

  • Contain 1 sub-orbital (have a total of 2 electrons)

• P-orbitals

  • Contain 3 sub-orbitals (have a total of 6 electrons, 2 each)

• D-orbitals

  • Contain 5 sub-orbitals (have a total of 10 electrons, 2 each)

• F-orbitals

  • Contain 7 sub-orbitals (have a total of 14 electrons, 2 each)

Below is what these orbitals look like:

Fig.1 The different orbital and suborbital shapes

Single covalent bonds are caused by direct orbital overlap. These bonds are also called sigma (σ) bonds. In double and triple bonds, the first of these bonds is a σ-bond, while the other(s) are pi (π) bonds. Π-bonds are caused by sideways overlap between orbitals.

Below is an example of both types of bonds:

Fig.2-Examples of sigma and pi bonding

On the top row are examples of sigma bonding, while the bottom row is pi-bonding. Pi-bonding can only occur between orbitals of p-orbital energy or higher (i.e. d or f), while sigma bonding can occur between any orbitals.

Here is what these bonds look like:

Types of Covalent Bond: Electronegativity

The second category of covalent bond is based on electronegativity.

Elements with the largest electronegativity are near the top right of the periodic table (fluorine) while elements with the smallest electronegativity are near the bottom left (francium), as shown below:

Fig.4-Table of electronegativities

The two types of covalent bonds in this category are:

• Non-polar covalent

• Polar covalent

Here, "polarity" refers to the difference in electronegativity between elements. When one element has a significantly higher electronegativity (>0.4), the bond is considered polar.

What happens is the electrons are attracted to this more electronegative element, which causes an uneven distribution of electrons. This in turn causes the side with more electrons to be slightly negatively charged (δ-), and the side with fewer electrons to be slightly positively charged (δ+)

For example, below is HF (hydrogen fluoride), which is a polar covalent compound:

Fig.5-Hydrogen fluoride has a polar covalent bond

The separation of these charges is called a dipole.

In non-polar covalent bonds, there is a small enough difference in electronegativity (<0.4), that is distribution of charge doesn't occur, so there is no polarity. An example of this would be F₂.

Determining Covalent Bond Length

Now, let's dive into bond length.

Covalent bond length is determined by bond order.

The higher the bond order, the shorter the bond. The reason why larger bonds are shorter is that the attractive forces between them are stronger.

When looking at diatomic (two-atom) compounds, the bond order is simply equal to the number of bonds (i.e. single=1, double=2, and triple=3). However, for compounds with more than two atoms, the bond order is equal to the total number of bonds minus the number of things bonded to that atom.

Let's do a quick example to explain:

Characteristics and Properties of Covalent Compounds

Now that we've covered the basics, we can finally talk about covalent compound properties!

Here are some of the common properties/characteristics of covalent compounds:

• Low melting and boiling points

  • While the bonds themselves are strong, the forces between molecules (called intermolecular forces) are weaker than those between ionic compounds, so they are easier to break/disrupt

• Poor conductors of electricity

  • Covalent compounds don't contain ions/charged particles, so they can't transport electrons well

• Soft and flexible

  • However, if the compounds are crystalline, this is not the case

• Nonpolar covalent compounds dissolve poorly in water

  • Water is a polar compound, and the rule for dissolving is "like dissolves like" (i.e. polar dissolves polar and non-polar dissolved nonpolar)

Uses of Covalent Compounds

There are a plethora of covalent compounds, and as such, there are a plethora of uses for them. Here are just some of the many covalent compounds and their uses:

• Sucrose (table sugar) (C₁₂H₂₂O₁₁) is a common sweetener is foods

• Water (H₂O) is a necessary compound for all life

• Ammonia (NH₃) is used in several types of cleaning products

• Methane (CH₄) is the main component in natural gas and can be used for things such as home heating and gas stoves