Bonding Types

We have met types of bonding already, but we need to be more specific about "inter-molecular bonding". Once we classify what type of bonding a substance can make, we can use this to explain physical properties such as melting point trends or molar heat of vaporisation trends.

Polarity

This is a very quick overview for how to work out if a molecule is polar or not:

Shapes of Ions/Molecules

We were introduced to the shapes of molecules/ions with 5 or 6 regions of electron density.

Source: http://figures.boundless.com/10397/full/vsepr-geometries.png

Lewis Dot Diagrams

Start with a recap of Level 2 Chemistry:

Then, how about some examples which break the Octet Rule:

Finally, Lewis Dot Diagrams for polyatomic ions:

Transition Metals

Transition metals are a great context to show our understanding of atomic structure and periodic trends. The transition metals are found in the "d block" of the Periodic Table of Elements.

They have some important (but not necessarily unique) properties:

The following website has an excellent summary of these key properties: http://www.chemguide.co.uk/inorganic/transition/features.html

Variable Oxidation States/Numbers

In all metals, the s-orbital electrons are involved in bonding. This is usually an ionic bond (when it is with another element). As d-orbital electrons can also be involved in bonding, there are many permutations for the number of electrons being involved in bonding, so often more than one possible oxidation state/number.

Zinc is an exception, as its ion has a full 3d-orbital. This makes these electrons unavailable for bonding, similar to the p-orbital electrons of Noble Gases being unavailable for bonding.

This is summarised well in this website: http://www.s-cool.co.uk/a-level/chemistry/transition-metals/revise-it/variable-oxidation-states

Coloured Compounds

Substances are coloured when they absorb part of the electromagnetic spectrum (light) and reflect the remainder. Transition metals' ions are usually coloured because they have incomplete d-orbitals. When bound to other ions/molecules, the d-orbital splits into two groups (with slightly different energy levels). The energy levels of these is so close that white light has enough energy to excited electrons from one energy level to the other. The remaining wavelengths/frequencies are the colour we see.

SOURCE: http://www.chemguide.co.uk/inorganic/complexions/absorbyellow.gif

Zinc is an exception, as its ion has a full 3d-orbital. This makes it impossible for the electrons to move into an excited state due to electromagnetic radiation (light) as they would need to "jump" into a completely different orbital or energy level.

This is summarised well in this website: http://www.s-cool.co.uk/a-level/chemistry/transition-metals/revise-it/coloured-compounds

Catalysts

Simplistically, the density of positive charge in transition metal ions, and their ability to make dative bonds (due to incomplete d-orbitals), make these useful catalysts.

Alternatively, the variable oxidation states of transition metals allow them to be involved in the reaction (alternate pathway provided) then be recovered in their original oxidation state. While being involved, the transition metal is not "used up", so this is still considered to be a catalyst.

There are some excellent examples in these websites: http://www.sky-web.pwp.blueyonder.co.uk/Science/transitionmetalcatalysts.htm and http://www.4college.co.uk/a/ss/catalyst.php

Complex Ions

Explaining the details of this is actually very complicated. However, the reason transition metals can form complex ions is very similar to their ability to be catalysts; a dative bond can be formed with molecules/ions (ligands) due to the charge density of the cation. d-orbitals are involved in this, but the d-orbitals do not need to be incomplete.

This is well explained in this website: http://www.s-cool.co.uk/a-level/chemistry/transition-metals/revise-it/complex-ions

A more in-depth explanation can be found on these pages:

1. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch12/valence.php
2. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch12/crystal.php
3. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch12/ligand.php

Periodic Trends

This week, we need to use our understanding of atomic structure to justify the trends in certain Periodic Trends:

1. atomic radius
2. ionic radius
3. electronegativity
4. 1st ionisation energy

For example:

Electron Configuration and Hund's Rule

We spent more time on using the spdf electron configuration. We also looked at applying Hund's Rule, which helps us understand such things as the electron configuration of Cu:

Using the Periodic Table

In previous years, we focused primarily on using the first 20 elements to understand how the Periodic Table is arranged. From this, we could infer the valency of and ionic charge formed by each element. While this was a good model, it fell down once we got to Sc (#21) and the other Transition Metals.

In this lesson, we were introduced to a more robust model of the atom, and a better way to note electron configurations. The "spdf notation" helps us better explain the properties of the elements.

We also heated some ionic compounds in Bunsen flames, to show that certain elements (and their ions) create very characteristic colours. The electrons move into a much higher Energy Level (excited state), then emit wavelengths of light when they return to their natural Energy Level (ground state).