A Level - Intermolecular Forces

Intermolecular forces are weak bonds which form between molecules. They are different to the bonds between atoms within molecules - these are intramolecular forces (intra- meaning inside). All molecules and even the monoatomic noble gases will experience these forces, though in some cases, you'll need to lower the temperature in order to reduce the kinetic energy to such a level that the molecules can 'stick' together (i.e. the intermolecular forces can properly take hold).

There are different types of intermolecular forces, but all of them are based on electrostatic attraction; the force of attraction between positive and negative charges.

At A Level, there are three types of intermolecular forces that you need to be aware of:

• Hydrogen bonds
• Permanent dipole-dipole interactions
• London forces (sometimes called Van der Wall's, depending on your specification)

Of these, hydrogen bonds are the strongest and they only occur between molecules which have a O, N or F atom bonded to a hydrogen atom. Because of the difference in electronegativity of these atoms, there is a dipole present across the bond. As the oxygen is more electronegative, there is a partial positive charge on the hydrogen and a partial negative charge on the other atom.

When these two partial charges interact, a hydrogen bond is formed, represented here by a dashed straight line, as shown here in water:

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If the above was a student answer to the question 'draw a diagram to show how molecules of water undergo hydrogen bonding', have a think about what this student may have missed.

From an exam perspective, it is important to remember to draw the dipoles and any lone pairs, as well as the hydrogen bond itself as this is where students very often lose marks - and easy marks too! The lone pairs is what the student would have missed if the diagram above was an exam answer - they'd thrown away easy marks! The diagram below would be a superior answer:

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The comparative strength of hydrogen bonds, against to the other types of intermolecular forces experienced by other simple covalent molecules is what gives water an unusually high melting and boiling point as they require more energy to overcome. Hydrogen bonds also play a part in reducing the density of water as it transitions form a liquid to a solid; they hold the water molecules apart in an open lattice, as depicted below:

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After hydrogen bonding, the next strongest intermolecular force is permanent dipole-dipole interactions. These interactions are between dipoles which have irreversibly formed due to a difference in electronegativity between the two bonded atoms in a molecule. One single molecule may contain more than one permanent dipole. 

An example of such a molecule would be chloromethane, in which the chlorine atom is significantly more electronegative than the carbon atom, resulting a permanent dipole. The partial positive charge on the carbon can interact with a partial negative charge on the chlorine atom of another molecule, creating a weak intermolecular bond. Whilst they are weak, may of these bonds can require a good amount of energy to overcome, thus raising the boiling point of a substance. 

Finally, we have London Forces. These forces are the weakest of the three intermolecular forces studies at A Level. They arise due to a asymmetrical distribution of a electrons in a species, which gives rise to an instantaneous (sometimes termed 'transient') dipole. Once this dipole forms, it creates dipoles in other nearby species and these are aptly referred to as 'induced dipoles', as shown below:

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When these dipoles interact, they produce London forces, which, although incredibly weak, when in number they can raise the boiling point of a molecule by raising the energy required to overcome the intermolecular forces and bring about a change of state.

So, to summarise:

• Molecules with O - H, N - H and F - H bonds form hydrogen bonds, the strongest type of intermolecular force. 
• The permanent dipoles within molecules can interact and these forces are weaker than hydrogen bonding. 
• Instantaneous dipoles can induce dipoles in atoms of other molecules, forming London Forces - the weakest of the intermolecular forces. 
• The stronger or more numerous the intermolecular forces, the greater the melting/boiling point of a chemical. 

A Level - Electronegativity: An Atomic Tug Of War!

Covalent bonds aren't all purely covalent. The electron pair which resides between the nuclei of two atoms, held in place by strong electrostatic forces, doesn't always stay exactly in the middle of the two atoms. 

You can imagine that each of the two nuclei exert a force of attraction upon the bonded pair. and this force is different for different elements. The ability of an atom to draw a pair of bonded electrons within a covalent bond towards itself is known as electronegativity. 

The above diagram illustrates, in a silly way, that some atoms are much better at drawing electrons towards them than others. This means that the bonding electrons spend more of their time closer to the more electronegative atom within the bond, creating a partial charge (dipole):

Where the electrons have been drawn towards the more electronegative atom, a partial negative charge is created, represented here by the δ- symbol. In turn, where the electrons have been drawn away from one atom, this leaves it with a partial positive charge, denoted by the δ+ symbol. Collectively, these two charges make up a dipole. 

Electronegativity is influenced by three main factors:

• Atomic Radius - the smaller an atom, the closer the bonding electrons will be to the nucleus and therefore the greater the electrostatic attraction will be. 
• Nuclear Charge - the more protons in the nucleus, the greater the charge of the nucleus and the electrostatic attraction. 
• Shielding - Inner electron shells shield the nuclear charge from the bonding pair, reducing the effective nuclear charge (attraction felt by the bonding pair of electrons).

The electronegativity of elements increases as you go from left to right and from the bottom to the top of the periodic table, like so:

Note that group 0 are not accounted for with electronegativity as they rarely bond to anything. 

The electronegativity of different elements was quantified by American Chemist Linus Pauling, in the Pauling scale, as shown below:

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So, now we return to the start of this post, where I stated that most covalent bonds are not purely covalent. Bonds actually fall onto one of three categories:

• pure covalent
• polar covalent 
• ionic

So, in summary:

• Electronegativity is the ability of an atom to draw a bonding pair of electrons from a covalent bond towards itself.
• Electronegativity increases progressively across and up the periodic table. 
• A difference in electronegativity can create a polar bond with a dipole. 

A Level - spd Notation of Transition Elements

The transition elements all have valence electrons in a d sub-shell, which is part of the reason that they make such good catalysts and form coloured solutions of their ions. 

Depending upon your exam board, a transition element is either defined as an element which forms ions with an incomplete d sub-shell or a element with an incomplete d sub-shell. Personally, I settle with the former definition, and for that reason both zinc (Zn) and scandium (Sc) are not classed as transition elements.

The electron configuration of scandium (Sc, above) shows that the 3d sub-shell contains only one electron. Scandium only forms Sc⁺ ions by losing the 3d electron as it has the highest energy, thus forming an ion without an incomplete d sub shell - not a transition element!

The story is similar for zinc (Zn, above), which forms Zn²⁺ ions, by losing the two electrons in the 4s, forming an ion with a complete 3d sub-shell - not a transition element! It loses the 4s electrons and not the 3d because it more energetically beneficial for it to have an empty 4s sub-shell and a full 3d sub-shell than it is for it to have a full 4s sub-shell and a partially-full 3d sub-shell. 

Now we have covered the elements in the series that are exempt from the title of 'transition element', let's turn out attention to Ti - Zn. By the time we get to chromium (Cr), there are 4 electrons in the 3d sub-shell and 2 still in the 4s:

However, the 3d sub-shell is really close to being half-full which would impart some stability (release some energy). As the energy gap between 3d and 4s is only very small, the energy required to promote a 4s electron to the 3d sub-shell in order to half-fill it is comparable to the energy released by half-filling it, therefore, this is exactly what happens:

Thus making the electron configuration of chromium 1s²2s²2p⁶3s²3p⁶4s¹3d⁵.

A similar phenomenon happens for copper, where the 3d can be completely filled by promoting a 4s electron:

This makes copper's electron configuration 1s²2s²2p⁶3s²3p⁶4s¹3d¹⁰.

So, in summary:

• Zinc and scandium are sometimes not considered transition metals as their ions do not have an incomplete d sub-shell. 
• Chromium promotes a 4s electron to half-fill the 3d sub-shell. 
• Copper promotes a 4s electron to fill the 3-d sub-shell completely. 

GCSE - Oxidation and Reduction

Oxidation and reduction are terms used in chemistry to describe what has happened to an element in a chemical reaction in terms fo loss or gain of electrons. 

A element which has been oxidised in a chemical reaction has lost one or more electrons. An element which has been reduced in a chemical reaction had gained one or more electrons. A good example of both processes happening simultaneously would be would be the reaction of magnesium with chlorine:

Mg + Cl₂ à MgCl₂

In the above reaction, magnesium, as a group 2 element, loses two electrons to get a full outer shell, which the chlorine atoms each gain one of these two electrons to complete each of their outer shells. Magnesium has lost electrons and has been oxidised and chlorine had gained electrons and been reduced. As oxidation and reduction are occurring within the same reaction, we cal this type of reaction a 'redox' reaction. 

A simple way to remember this is OIL RIG:

O - xidation 
I - s
L - oss

R - eduction
I - s
G - ain

A common exam question surrounding oxidation and reduction is:

1. Explain in terms of electron why X has been oxidised.

Provided you know your oxidation and reduction, these questions can be easy marks. All you need to say is that 'X has been oxidised as it has lost one or more electrons'.

So there, you have it, a brief summary of oxidation and reduction!

GCSE - Exam Technique: Explaining Reactivity Trends

When answering exam questions on reactivity trends within a group, you must address the following points if you want to maximise your marks:

• How the number of shells changes down the group 
• How the number of shells affects the atomic radius and shielding
• How atomic radius and shielding affects the attraction between the nucleus and outer shell electrons
• How the changing force of attraction affects how easily electrons are lost or gained. 

As you go down group 1, the reactivity of the elements increases; Li < Na < K. In order to explain this properly, we need to hit all of the points above. As an example, I will use an exam question taken from the aQA exam question bank, Exambuilder and created a model answer:

(ii)     Potassium is more reactive than sodium.

Explain why, in terms of electronic structure.

               Potassium has more shells than fluorine. This makes its atomic                         radius greater and also the effect of shielding increases, reducing                       the attraction between the nucleus and the outer shell. This                               makes the electron in the outer shell of potassium easier to                               remove than in sodium. 

As you descend group 7, the elements get less reactive F > Cl > Br > I. Again, it is vitally important to hit those key points again - remember, you're trying to show the examiner all that you know! Here's another question from AQA Exambuilder, along with a model answer to illustrate this point:

(c)   Explain, in terms of electrons, why fluorine is the most reactive element in Group 7.

Of all the group 7 elements, fluorine has the fewest number of shells, which reduces the atomic radius and the effect of shielding. This means that in fluorine, the force of attraction between the nucleus and the outer shell is greatest, so it is easier for fluorine to gain an eighth electron to fill its outer shel than it is for the rest of group 7. 

It's really important to do practice paper questions as part of your revision - I cannot stress that enough! Your exam technique is a skill that will develop over time, but practicing writing answers and reviewing them with the mark scheme will make a big difference to your marks because you'll build an idea of what the examiner is looking for in certain types of questions. If you ask your teacher for exam questions, they swill likely be more than happy to supply you with some!

A Level - Electron Configuration and spd Notation

Starting where the last blog left off, you hopefully remember that electrons are pair with opposite spin in orbitals, which sit within sub-shells, within shells described by the principle quantum number (n). We can show electron configurations on an energy diagram as shown below:

When it comes to assigning electrons to orbitals, the simplest atom is hydrogen (H), which has one electron in the 1s sub-shell, as shown below:

The single electron is placed in the lowest energy orbital (1s) first, and denoted here by an arrow, with the direction, up or down, showing the spin. In helium (He), the second electron has opposite spin (down-spin) and is paired with another electron in the 1s:

By the time we get to nitrogen, each p orbital contains one electron, and the 2p sub-shell is half-full. This imparts some stability, which is something that ties into ionisation energies:

Oxygen has one more electron, which is paired with one of the electrons already in the 2p sub-shell with opposite spin: 

The diagrams above are a good way of visualising the arrangement of orbitals, however, it would be extremely lengthy to draw out every time. We usually represent the electron configuration of a substance spd notation, as shown below:

So this means that the spd notation of oxygen is 1s²2s²2p⁴.

The periodic table can be separated into three blocks; s, p and d blocks. The s-block is comprised of elements who's last sub-shell is an s sub-shell and so on...

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So, if we pick out phosphorus (P) for example, as we know it is in the p block, its last occupied sub-shell would be the 3p and as it is three places into the p block, the 3p sub-shell would have an occupancy of 3, so 3p³. Overall the spd notation of phosphorus would be 1s²2s²2p⁶3s²3p³.

In summary:

• Electrons fill the lowest energy sub-shell first.
• Electrons fill sub-shells until they are half full and then pair with electrons of opposite spin to fill the sub-shell.
• An element's position in the periodic table can tell you its electron configuration.