Bonding

LEWIS STRUCTURES/DIAGRAMS

We use Lewis Dot Diagrams to represent the bonding pairs and lone pairs of electrons in a molecule and an ion. We only focus on the valence electrons, as these are the ones that influence the shape of the ion/molecule, so affect their properties the most.

Simple Lewis Dot Diagrams

In a lot of molecules, the atoms automatically obey the Octet Rule once we include the lone pairs of electrons.

More Complex Lewis Dot Diagrams

These molecules (and ions) do not obey the Octet Rule once we assign the remaining electrons as lone pairs. Therefore, we need to shift pairs of electrons around to create double (or triple) bonds.

Expanded Octet

Some elements can break the Octet Rule by having more than eight electrons. We call this having an expanded octet. It can happen because electrons can be "promoted" into the 3d orbitals to accommodate bonding. There are a couple of "rules" around this:

The atom must already have electrons in both the 3s and 3p sub-levels.
The atoms must be the central atom (it would be anyway, as this increases its valency)

Polyatomic Ions

We can even draw Lewis Diagrams for Ions. We just need to allow for the addition of electrons (anions) or removal of electrons (cations).

There are also a couple of special cases that we should be aware of - sulfate and phosphate. Have a go at drawing these, knowing they have expanded octets around the central atom. Then get your diagrams checked...

FORCES BETWEEN MOLECULES

Once we know if a substance is ionic, polar or non-polar, we can describe the forces (bonds) between the particles in more detail. These forces can then be related to the properties of that substance, especially their:

state of matter at room temperature
melting point
boiling point
enthalpy of vaporisation
enthalpy of fusion

Ionic Compounds

Ionic compounds are made up of cations and anions, held together in a 3-dimensional lattice by strong, directional ionic bonds. However, not all ionic bonds are equal...

Smaller ions have stronger ionic bonds than larger ions (of the same charge).

Ions with a higher charge have stronger bonds than ions with a lower charge (of a similar radius).

We need to consider these two factors (together), when inferring the relative strength of ionic bonds.

The strength of an ionic bond can be inferred from a substance's melting point, boiling point and/or solubility in water.

The higher the m.p./b.p., the stronger the ionic bond. This is because melting and boiling break ionic bonds to turn the substance into liquids or gases, respectively.

The higher the solubility in water, the weaker the ionic bond. This is because dissolving requires the ions to have a strong enough attraction to the water molecules to break the ionic bonds holding them in the ionic lattice

More Detail...

If you want to know more about this, and the link it has to Coulomb's Law, check out this video. It is a bit of "extra for experts" and is not assessed in Level 3 NCEA.

Temporary Dipole Forces

Non-polar molecules can form temporary dipoles. These dipoles come and go (are not permanent), but do allow for intermolecular bonding. The frequency of these bonds help us explain characteristics like solubility, melting point and boiling point (and sublimation point for iodine and carbon dioxide).

Temporary dipoles are more likely to form in molecules with large electron clouds. This is hard to infer quickly, so we use relative molar mass (MR) to infer the size of the electron cloud. Molecules with larger MR will usually have larger electron clouds.

The intermolecular bond between neighbouring non-polar molecules is called a "temporary dipole-temporary dipole force". It is usually due to an instantaneous dipole in one molecule and an induced dipole in the other molecule.

The intermolecular bond between a polar molecule and a non-polar molecule is called a "permanent dipole-temporary dipole force". It is due to a permanent dipole in one molecule (the polar molecule) and an induced dipole in the other molecule. (the non-polar molecule).

Bonding Overview

Khan Academy does- a really good job of summarising the forces between molecules. This section is sharing their videos and notes for you.

Some of the terminology is different to what we will use for NCEA:

London dispersion forces = temporary dipole-temporary dipole forces

dipole = permanent dipole

Intramolecular and intermolecular forces (article) | Khan AcademyLearn for free about math, art, computer programming, economics, physics, chemistry, biology, medicine, finance, history, and more. Khan Academy is a nonprofit with the mission of providing a free, world-class education for anyone, anywhere.

PROPERTIES

Melting and Boiling Points

The nature of the particles, and therefore the strength of forces between the particles, can be used to explain trends in melting and boiling points. Melting and boiling require the breaking of bonds. The stronger the bonds/forces between the particles, the higher the melting and boiling points will be.

The water example has an interesting outcome when it comes to boiling. While it lowers the melting point of water (ice), it increases the boiling point of water. This is because liquid water primarily contains only permanent dipole-permanent dipole forces (a type of intermolecular bonds). Salt water contains ions, dissolved in (surrounded by) water. These have a stronger ionic-permanent dipole forces between the particles. These take more energy to break than hydrogen bonds and a lot more energy to break than permanent dipole-permanent dipole forces.

This is why we often add salt to vegetables when we cook them. The water can get a lot hotter before boiling, so cooks the vegetables more quickly. It also has the added bonus of seasoning the food (adding flavour).

Dissolving

Dissolving is actually a difficult concept to explain. Why do some solutes dissolve in some solvents but not in others? Why are some ionic substances very soluble in water while others appear to be insoluble? Why are some fluids (liquids and gases) miscible in water, while others are not. Why is something like iodine sparingly soluble? Surely it should either be soluble or insoluble!!

Dissolving requires an attraction between the solvent and solute particles to do two things:

Break bonds between solute particles
Break bonds between solvent particles (in the liquid phase)

Dissolving Molecular Substances

For molecular solutes, we have an easy "rule" to remember: "Like Dissolves Like".

Polar solvents will dissolve polar solutes.
Non-polar solvents will dissolve non-polar solutes.

Dissolving Ionic Substances

Ionic solutes can be explained by looking at the strength of the ionic bond. We learned that this is determined by:

ionic radius
ionic charge

Water (and other polar solvents) have permanent dipoles, so can form bonds, called ionic-permanent dipole forces. Multiple ionic-permanent dipole forces allow for the ionic bond to be overcome, so for dissolving to occur.

Try to use this understanding to explain things like why iodine (very large molar mass) will dissolve in water (sparingly) despite being non-polar. Why is its solubility so much higher in cyclohexane and cyclohexene?

Sublimation

Sublimation occurs when a solid goes directly to the gas phase, without first becoming a liquid. It is most common with solid carbon dioxide (dry ice) and iodine.

At a particle level, why does this happen?

Both iodine and dry ice are non-polar molecules. Therefore, the forces between the molecules are temporary dipole-temporary dipole forces (very weak intermolecular bonds). However, their respective electron clouds are large enough to generate enough temporary dipoles to form a 3D lattice.

Image Source: https://socratic.org/questions/552d8428581e2a12d7a7324d

However, with such weak forces between these molecules, a small amount of (heat) energy is enough to break multiple intermolecular bonds. The molecules are released and moving too fast (and too far away from other molecules) to form new, temporary intermolecular bonds, required to exist in the liquid phase. Once the bonds in the solid are broken, the entire structure is compromised, so every molecule "escapes" the 3D lattice. It becomes a gas. The particles are moving too quickly and are too far apart to form the bonds required to hold them in the liquid phase.

Iodine is the only one of these that is stable as a solid at room temperature, so we will explore that further. It is good to compare to bromine (liquid at room temperature, but readily becomes a gas as well), as they are both halogens (in Group 17).

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