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Help with 6 A-level Chemistry Concepts Many Students Don’t Understand – With Clear Explanations!|
About the author
Philippa Logan read Natural Sciences (specialising in Chemistry) at Newnham College, Cambridge.
A concept has a simple enough definition: ‘an abstract or general idea inferred or derived from specific instances’. The word itself is derived from the Latin conceptum, which means something received or conceived.
The study of chemistry at A level – or, indeed at any level – depends on the learning of many fundamental concepts in chemistry: nomenclature, atomic structure, energetics and kinetics, groups in the Periodic Table, and so on. When it comes to the exam, and throughout your course, it will make your life much easier if you understand the concepts rather than simply learning examples by heart. You can probably get away with rote learning to a certain extent, but you cannot learn examples for all possible situations: exam conditions will reveal whether you really understand the concepts, or whether you are attempting to apply your learning to the questions asked.
The concepts that each student finds difficult will naturally depend on the inclination of the student: those who hope to be going on to study chemical engineering will struggle with different concepts than someone who prefers physical chemistry; mathematically-inclined students will not struggle with the calculations of the more mathematical concepts in the same way as, perhaps, those students who prefer organic chemistry might.
That said, here are a few example concepts that A-level chemistry students may find particularly challenging.
The idea of equilibrium starts off straightforwardly enough.
The generic chemical equation is:
Reactants → Products
A + B → C + D
Some reactions are reversible. For example, if you heat blue hydrated copper sulphate, the water of crystallisation is driven off: the copper sulphate becomes anhydrous and turns white. If you then add water to the anhydrous copper sulphate, it turns back to blue again.
However, if you carry out this experiment in a CLOSED container, the products repeatedly react together to form the reactants, which then react to form the products, and so on. Eventually, the proportions of all the components becomes constant, but they still carry on reacting; this is called an equilibrium mixture, and the products are in dynamic equilibrium.
We write a reversible reaction as A + B ⇋ C + D
Some important points about equilibrium:
Kc = [C] [D]
More specifically, for a reversible reaction aA + bB ⇋ cC + dD, the equilibrium constant involves the ratio of the concentrations of the products to the concentrations of the reactants, each raised to their stoichiometric coefficient (the number of molecules of that substance in the balanced chemical equation), i.e.
Kc = [C]c [D]d
Still on the subject of equilibrium, it’s worth being clear about the concept involved in Le Chatelier’s Principle.
This states that:
If you disturb a system at equilibrium, the equilibrium moves in the direction that tends to reduce the disturbance.
In other words, it tells us in which direction the equilibrium moves, if there is any change to the conditions of the equilibrium mixture. This might be a change in the concentration of the reactants, the pressure, or the temperature. The equilibrium will move in the direction that opposes the change. We’ll look at some examples to illustrate this.
Let’s go back to our generic equation:
A + B ⇋ C + D
If we add more of reactant A (i.e. if we increase the concentration of A), then Le Chatelier’s Principle tells us that the equilibrium tries to counteract that by reducing the concentration of A. So the equilibrium moves to the right, to use up more B, which in turn produces more C and D.
Remember: changing the concentration of reactants, or products, has no effect on the rate constants of the reaction – and therefore no effect on the equilibrium constant.
We only need to consider pressure changes where the reaction involves gases. Here, we need to know the number of molecules of a gas on each side of the equilibrium equation. For instance,
2SO2(g) + O2(g) ⇋ 2SO3(g)
3 molecules 2 molecules
Increasing the pressure of a gas is the equivalent of increasing the concentration of a solution. So, if we increase the pressure on the above system, the position of the equilibrium will move so as to reduce the pressure. The equilibrium will move to the right, because there are fewer molecules on the right-hand side of the equation, and fewer molecules mean lower pressure.
Remember: increasing the pressure (or reducing the volume) has no effect on the equilibrium constant.
If the forward reaction is exothermic (i.e. if it gives out heat), then the reverse reaction will be endothermic (i.e. will take in heat).
Let’s say we increase the temperature of an equilibrium mixture whose forward reaction is exothermic. Le Chatelier’s Principle tells us that the equilibrium will move in the direction that will cool the system down – i.e. to the left. It might be helpful to think of heat as one of the reactants or products in a reaction. So, increasing the temperature of an exothermic reaction is like introducing another product of the reaction. The equilibrium will try to correct this imbalance, by converting some of the products back into reactants.
For example, take the reaction in which NO2 dimerises into N2O4.
2 NO2 (g) ⇋ N2O4 (g)
The reaction is exothermic. Increasing the temperature is equivalent to adding more of the product, so the equilibrium constant will decrease and the backward reaction will be more favoured.
Catalysts have no effect on the position of equilibrium, because they affect the forward and backward reactions to the same extent.
However, they do allow equilibrium to be reached more quickly – which is why they are useful in industry.
The term ‘redox’ stands for reduction-oxidation. The concept of oxidation states (or oxidation numbers) is all to do with the number of electrons lost or gained.
Oxidation is the loss of electrons – which is what happens when oxygen is added. For example:
Cu (s) + ½ O2 (g) → CuO (s)
The copper has lost two electrons, and has been oxidised, as the half equation shows:
Cu – 2e – → Cu2+
which we can also write as Cu → Cu2+ + 2e –
Reduction is the gain of electrons, and is what happens when oxygen is removed, e.g.:
CuO (s) + H2 (g) → Cu (s) + H2) (liq)
Here, the copper oxide has been reduced, and has gained two electrons. We can see this if we write CuO as Cu2+ O2- , and just consider the copper part of the molecule.
The half equation is:
Cu2+ + 2e– → Cu
One way of remembering which is which is OIL and RIG:
Oxidation Is Loss of electrons (OIL)
Reduction Is Gain of electrons (RIG)
but if you understand the concepts behind oxidation and reduction, you won’t need these acronyms.
Put another way:
These oxidation states are useful in helping us to balance redox reactions. For a redox equation to be balanced:
Here’s an example of a redox reaction that happens in a blast furnace, in the production of iron:
Fe2O3 (s) + 3CO (g) → 2 Fe (liq) + 3 CO2 (g)
The term pH stands for the power of hydrogen, and is a measure of the acidity of a solution.
More specifically, it is the concentration of H+ ions in a solution that determines its acidity.
pH = – log 10 [H+(aq)]
The logarithm of the concentration is used rather than the actual concentration, because the concentrations of H+ (aq)ions in solution are generally so small.
Now let’s look at a specific acid: ethanoic acid, CH3COOH (the acid in vinegar). This is a weak acid, because only about 4 in every 1,000 ethanoic acid molecules are dissolved into ions.
CH3COOH (aq) ⇋ H+ (aq) + CH3COO– (aq)
Before dissociation: 1,000 0 0
At equilibrium: 996 4 4
The equilibrium constant, Kc, is the ratio of the concentration of the products over the reactants (see the end of section 1 in this article). For weak acids, Kc is usually called the acid dissociation constant, Ka
So here, Ka = [CH3COO– (aq)]
[H+ (aq)] [CH3COOH (aq)]
This formula is the basis of all questions on pH and on buffer solutions.
Interpreting spectra is difficult. But it helps if you understand how nuclear magnetic resonance spectroscopy works, and the concepts behind it.
NMR depends on the concept of spin in materials. Many nuclei with odd mass numbers, such as 1H, 13C and 15N, have the property of spin, which gives them a magnetic field – just like a bar magnet.
If you put bar magnets in an external magnetic field, they will align themselves parallel to the field:
But imagine what would happen if the bar magnets aligned themselves anti-parallel to the external magnetic field.
In this case, the bar magnets have to be forced into position, because of the repulsive forces (like poles repel), and so this arrangement in Figure 2 has a higher energy than that in Figure 1. The stronger the external magnetic field and the stronger the bar magnets, the larger the energy difference between the parallel and the anti-parallel arrangements.
The situation is similar with nuclei that have spin. Most of the nuclei will be in the lower energy state (the parallel arrangement), though some will be in the higher energy state. However, if electromagnetic energy is supplied that is exactly equal to the energy gap between the two states, then some nuclei will flip into the higher energy (anti-parallel) state. This is called resonance. The energy used to supply this is supplied by a radio frequency source, and the resonances are detected by a radio receiver.
The interpretation of NMR spectra is all based on the energy needed to flip specific nuclei in specific environments, and is a powerful technique to help us determine the structure of molecules.
Isomers are compounds that have the same molecular formula, but their atoms are arranged differently. They occur frequently in organic chemistry.
The concept of structural isomers is relatively straightforward. Generally, the functional groups are attached to the main carbon chain at different points, or else they have a different arrangement of carbon atoms.
Stereoisomers are two or more compounds that have the same structural formula, but their bonds are arranged differently in space.
E-Z isomers (also known as geometric isomers or cis, trans isomers) are the easier type of stereoisomer to understand. They differ in the position of substituents at either end of a carbon-carbon double bond, e.g. E- 1,2 dichloroethene and Z-1,2 dichloroethene differ only in the position of the H and Cl atoms.
Optical isomers are not so easy to visualise; they differ in the three-dimensional positioning of their substituents. They will be mirror images of each other, but not identical, and cannot be superimposed.
Optical isomers were first distinguished using polarised light – hence their name. They are said to be chiral, and the two optical isomers for a molecule are called a pair of enantiomers.
It can be hard to understand why two compounds that are chemically identical, as optical isomers are, can have such dramatically different effects in biological systems. For instance, the thalidomide tragedy resulted from the use of both optical isomers of thalidomide made in a chemical reaction, and therefore in a 50:50 mix. The two are hard to separate and the mixture was used to create a drug for pregnant women: one relieves morning sickness, but the other – as it turned out – causes horrendous birth defects.
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