Year 12 Recap - Thermochemical Calculations

We need to remember and apply some key calculations in our thermochemistry:

ENERGY = AMOUNT (n) × ΔH

In Level 3, we may need to report extra information as well, such as the mass of one of the products or reactants.

Year 12 Recap - Making and Breaking Bonds

It takes energy to break chemical bonds. Therefore, this process is endothermic.

Conversely, when bonds form, energy is released. Bond formation is exothermic.

We can apply this to our understanding of melting (and boiling) points. We can also use this to understand what is happening at a particle level in an endothermic or exothermic chemical reaction.

Please note - the sciPad page are from the Year 12 book, not the Year 13 book

Year 12 Recap - Energy Changes

Some reactions feel hot; some feel cold. This means there are energy differences between the reactants and the products. We call this an enthalpy change.

If something feels hot, it is releasing energy. We call this exothermic.

If something feels cold, it is absorbing energy from our skin. We call this endothermic.

Please note - the sciPad page are from the Year 12 book, not the Year 13 book

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 pahse.

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).

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:

1. Break bonds between solute particles
2. 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?

Melting and Boiling Point

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).

Temporary Dipoles

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 (M_R) to infer the size of the electron cloud. Molecules with larger M_R 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).

Intermolecular Bonding Overview

Khan Academy does a really good job of summarising the forces between molecules. This blog post 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

Khan Academy Notes

(the heading is a hyperlink)

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.

Forces Between Particles

In Year 12, we explored four types of substances: metallic, ionic, covalent networks, and molecular. We had to explain the properties of these using the particles they were made up of, and the forces (bonds) between these particles. We simply take this a step further in Year 13, focusing on ionic and molecular compounds. We start to use our deeper knowledge of atomic structure and shapes/polarity to explain differences (and peculiarities) in melting point (and boiling point) and solubility.

This video is an overview of the learning ahead. Each concept will be explored in more detail over the next two weeks.