As Level - Blog Posts

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Water: Making a Splash

You don’t have to be a genius to know that water is essential for life. After all, we’re made up of it, we sweat it, we drink it, some people even opt to give birth in it. But what is it about two hydrogens and an oxygen which make it so sensational?

The answer is to do with water’s structure. A H2O molecule is covalently bonded, which means each atom shares electrons. In this case, the covalent bonds are between two hydrogen atoms and one oxygen atom. Oxygen is cool because it is highly electronegative. Electronegativity is the ability for one atom to “pull” the electrons towards it in a covalent bond. Since oxygen is highly electronegative, it pulls the electrons in the bond towards it which gives the oxygen a slight negative charge because of the electron proximity. This is represented by  δ- (delta negative). The hydrogen is therefore δ+ (delta positive) and has a slight positive charge. Overall, the molecule is said to be polar, or to be dipolar in nature, because there is a difference in charge across the molecule.

Water being a dipole gives it different properties, which you need to know about if you are sitting the AS or A level biology exam. 

A quick note on hydrogen bonding…

Being a dipole, water has areas of different charge. When many molecules come together, hydrogen bonds can form between H+ on one molecule and O- on another, shown in the diagram with a dashed line. 

It is hydrogen bonds which give water a property called surface tension. Water has a high tendency to ‘stick together’, called cohesion. This is important in water transport through the xylem in later units. Surface tension is a bit like a “skin” because it can allow small organisms to walk along it. It occurs because water molecules on the surface bond to their neighbours much like throughout the whole liquid, but since one side is exposed to air and cannot form hydrogen bonds upwards, they will form stronger ones with the molecules beside them. The net attraction is downwards.

Water is good as a temperature buffer too. Heating a substance makes its particles gain more kinetic energy and therefore the overall temperature rises since particles are moving faster. With water, the temperature doesn’t rise as much as other liquids do. This is because it takes more heat energy to raise the temperature of water by 1 degree - it has a high specific heat capacity due to the many hydrogen bonds that have to be broken (even though they are weak on their own). It takes a lot of heat energy for water to raise its temperature significantly. 

This is useful in organisms because our cells are mostly water, which can absorb heat energy without raising our temperature very much. Therefore it “buffers” or reduces heat changes. Seas, lakes and oceans are all good environments to live in because they do not change temperature as quickly as air. Aquatic organisms have an environment with less temperature fluctuation than land organisms.

Having a high latent heat of vaporisation means water can cool down organisms by evaporating a small amount of water. Evaporation is when water becomes a gas due to the large amount of KE. Fast-moving molecules are removed when this occurs and take their energy with them, therefore decreasing the amount of energy left behind and cooling it. Sweat is a good example of high latent heat of vaporisation. A small quantity of water is removed with a large cooling effect, meaning temperature is stabilised without losing a lot of water.

Water is also a good solvent (a substance which can dissolve other substances) and this is due to more hydrogen bonding. Water’s charges of H+ and O- are attracted to the positive and negative charges on molecules and therefore solutes such as NaCl are split into Na+ and Cl-, then spread out. Solvent properties are important in transport (such as blood plasma dissolving glucose, vitamins, urea etc), metabolic reactions, urine production and mineral transportation through the xylem and phloem in plants.

Water molecules can also take place in metabolic reactions. Hydrolysis reactions involve breaking down the covalent bonds between hydrogen and oxygen and making new ones, for example, in digestion. Condensation reactions produce water as a byproduct e.g. the formation of phosphodiester bonds. Water is referred to as a metabolite.

Summary

Water is a dipole due to the slight opposite charges on oxygen and hydrogen atoms.

Hydrogen bonds form between hydrogens on one water molecule and oxygens on another. 

Because of this, water has the tendency to stick to itself - cohesion. Cohesion is the reason for surface tension.

Water is a good temperature buffer because of its high specific heat capacity. It takes a lot of energy to raise the temperature by a degree.

Water has a high latent heat of vaporisation so evaporating a little has a large cooling effect.

Water is a good solvent because of how the hydrogen bonds attract charged molecules and separate them. This is useful for transporting solutions.

Water is a metabolite important for hydrolysis reactions and produced in condensation reactions.

Happy studying!

Enthalpy - a thermodynamic property

When I first learned about enthalpy, I was shocked - it felt more like a physics lesson than a chemistry lesson. The thought of learning more about thermodynamics than my basic understanding from my many science lessons in lower school made me bored out of my mind. But enthalpy is actually pretty interesting, once you get your head around it…

Reactions which release heat to their surroundings are described to be exothermic. These are reactions like combustion reactions, oxidation reactions and neutralisation reactions. Endothermic reactions take in heat from their surroundings, such as in thermal decomposition. Reversible reactions are endothermic in one direction and exothermic in the other.

These facts are important when you start to look at enthalpy. Enthalpy is basically a thermodynamic property linked to internal energy, represented by a capital H. This is pretty much the energy released in bond breaking and made in bond making. We usually measure a change in enthalpy, represented by ∆H.  ∆H = enthalpy of the products (H1) - enthalpy of the reactants (H2). This is because we cannot measure enthalpy directly.

In exothermic reactions,  ∆H is negative whereas in endothermic reactions,  ∆H is positive.

∆H is always measured under standard conditions of 298K and 100kPa. 

In reversible reactions, the ∆H value is the same numerical value forwards and backwards but the sign is reversed. For example, in a forward exothermic reaction, the  ∆H value would be -ve but in the backwards reaction (endothermic) the  ∆H would be +ve. 

Reaction profiles are diagrams of enthalpy levels of reactants and products in a chemical reaction. X axis is enthalpy rather than ∆H and the Y axis is the progress of reaction, reaction coordinate or extent of reaction. Two horizontal lines show the enthalpy of reactants and products with the reactants on the left and the products on the right. These should be labelled with their names or formulae. 

In an endothermic reaction, product lines are higher enthalpy values than reactants. In an exothermic reaction, product lines are lower enthalpy values than reactants. The difference between product and reactant lines is labelled as  ∆H. Values are measured in kJ mol-1. 

Reaction pathways are shown with lines from the reactants to the products on enthalpy level diagrams. This shows the “journey” that the enthalpy takes during a reaction. They require an input of energy to break bonds before new bonds can form the products. The activation energy is the peak of the pathway above the enthalpy of reactants. It is the minimum amount of energy that reactants must have to react. 

Standard enthalpy values are the ∆H values for enthalpy changes of specific reactions measured under standard conditions, represented by ⊖. There are three of these:

1. Standard enthalpy of reaction ( ΔHr⊖ )

The enthalpy change when substances react under standard conditions in quantities given by the equation for the reaction.

2. Standard enthalpy of formation ( ΔfH⊖ )

The enthalpy change when 1 mole of a compound is formed from its constitutent elements with all reactants and products in standard states under standard conditions.

The enthalpy of formation for an element is zero is it is in it’s standard state for example, O2 enthalpy is zero.

3. Standard enthalpy of combustion ( ΔcH⊖ )

The enthalpy change when 1 mole of a substance is burned completely in excess oxygen with all reactants and products in their standard states under standard conditions.

Values for standard enthalpy of formation and combustion must be kept to per mole of what they refer.

Summary

Reactions which release heat to their surroundings are described to be exothermic. Endothermic reactions take in heat from their surroundings, such as in thermal decomposition. 

Reversible reactions are endothermic in one direction and exothermic in the other.

Enthalpy is a thermodynamic property linked to internal energy, represented by a capital H. We usually measure a change in enthalpy, represented by ∆H. 

 ∆H = enthalpy of the products (H1) - enthalpy of the reactants (H2). We cannot measure enthalpy directly.

In exothermic reactions,  ∆H is negative whereas in endothermic reactions,  ∆H is positive.

∆H is always measured under standard conditions of 298K and 100kPa. 

In reversible reactions, the ∆H value is the same numerical value forwards and backwards but the sign is reversed. 

Reaction profiles are diagrams of enthalpy levels of reactants and products in a chemical reaction. They 

In an endothermic reaction, product lines are higher enthalpy values than reactants. In an exothermic reaction, product lines are lower enthalpy values than reactants.

 The difference between product and reactant lines is labelled as  ∆H. 

Values are measured in kJ mol-1.

Reaction pathways are shown with lines from the reactants to the products on enthalpy level diagrams. They plot enthalpy against reaction progress.

Reactions require an input of energy to break bonds before new bonds can form the products. The activation energy is the peak of the pathway above the enthalpy of reactants. It is the minimum amount of energy that reactants must have to react.

Standard enthalpy values are the ∆H values for enthalpy changes of specific reactions measured under standard conditions, represented by ⊖. 

Standard enthalpy of reaction ( ΔHr⊖ ) is the enthalpy change when substances react under standard conditions in quantities given by the equation for the reaction.

Standard enthalpy of formation ( ΔfH⊖ ) is the enthalpy change when 1 mole of a compound is formed from its constitutent elements with all reactants and products in standard states under standard conditions.

The enthalpy of formation for an element is zero is it is in it’s standard state.

Standard enthalpy of combustion ( ΔcH⊖ ) is the enthalpy change when 1 mole of a substance is burned completely in excess oxygen with all reactants and products in their standard states under standard conditions.

Values for standard enthalpy of formation and combustion must be kept to per mole of what they refer.

Happy studying!

Nomenclature - what in the organic chemistry is it?

Organic chemistry is so widely studied it requires a standard system for naming compounds, developed by IUPAC. Nomenclature is simply naming these organic compounds.

So, you want to be an organic chemist? Well, it starts here. Are you ready?

(psst… once you’ve learnt this theory, try a quiz here!)

1. Count your longest continuous chain of carbons.

Bear in mind that some chains may be bent. You’re looking for the longest chain of subsequent carbon atoms. This number correlates to root names that indicate the carbon chain length, listed below:

The second part of naming your base comes from the bonding in the chain. Is it purely single bonds or are there double bonds in there? If you are familiar with carbon chemistry, you’ll know that saturated hydrocarbons are called alkanes and unsaturated hydrocarbons are called alkenes. Therefore, the syllable -ane is used when it has only single bonds and the syllable -ene is used when it has some double bonds. For example:

Sometimes carbon chains exist in rings rather than chains. These have the prefix of -cyclo.

2. Identify your side chains attached to this main carbon and name them.

Side chains are added as prefixes to the root names. Sometimes called substituents, these are basically anything that comes off the carbon chain. Examples of the prefixes are listed below:

There are other prefixes such as fluoro (-F) and chloro (-Cl) which can describe what is coming off the chain.

3. Identify where each side chain is attached and indicate the position by adding a number to the name. 

We aim to have numbers as small as possible. For example, if bromine is on the second carbon of a 5-carbon saturated chain, we number it as 2-bromopentane instead of 4-bromopentane, since it would essentially be 2-bromopentane if it was flipped. Locant is the term used for the number which describes the position of the substitute group, e.g. the ‘2′ in 2-chlorobutane is the locant.

Sometimes there are two or more side chains e.g. a methyl group and a chlorine attached to a pentane. In these cases, these rules apply:

1. Names are written alphabetically.

2. A separate number is needed for each side chain or group.

3. Hyphens are used to separate numbers and letters.

This would be named 2-chloro-3-methyl-pentane. This is because the longest chain of carbons is 5 (pentane), the chlorine is on the second carbon (2-chloro) and the methyl group is on the third carbon (3-methyl). It is 2-chloro rather than 4-chloro as we aim to have as small as numbers as possible.

Another variation of this step to be aware of is how many of the same side chains or groups there are, for example, having two methyl groups would be dimethyl rather than solely methyl. Each group must also be given numbers separated by commas to show where each one is located. 

The list of these prefixes is found here:

Convention does not usually require mono- to go before a single group or side chain.

4. Number the positions of double bonds if applicable.

Alkenes and other compounds have double bonds. These must be indicated with numbers. For example, pent-2-ene shows that the double bond is between carbon 2 and carbon 3. The number goes in the middle of the original root name e.g. butene, pentene.

(!) Below is a list of functional groups that you may need to study for the AS and A Level chemistry exams. “R” represents misc. carbons. It is important to know that some groups are more prioritised than naming. From the most to least priority: carboxylic acid, ester, acyl chloride, nitrile, aldehyde, ketone, alcohol, amine, alkene, halogenalkane. It is worthwhile learning these.

bigger version here (I suggest downloading and printing it)

But wait, there’s more:

Here are some things to bear in mind when naming organic compounds:

1. The letter ‘e’ is removed when there are two vowels together e.g. propanone rather than propaneone. The ‘e’ isn’t removed when it is next to consonant, e.g. propanenitrile isn’t propannitrile.

2. When compounds contain two different, one is named as part of the unbranched chain and the other is named as a substituent. Which way round this goes depends on the priority. 

SUMMARY

Count your longest continuous chain of carbons.

Chains may be bent. You’re looking for the longest chain of subsequent carbon atoms. This number correlates to root names that indicate the carbon chain length, e.g. pentane.

The second part of naming your base comes from the bonding in the chain. Is it purely single bonds or are there double bonds in there? The syllable -ane is used when it has only single bonds and the syllable -ene is used when it has some double bonds.

Rings have the prefix of -cyclo.

Identify your side chains attached to this main carbon and name them.

Side chains are added as prefixes to the root names. Sometimes called substituents, these are basically anything that comes off the carbon chain. 

There are other prefixes such as fluoro (-F) and chloro (-Cl) which can describe what is coming off the chain.

Identify where each side chain is attached and indicate the position by adding a number to the name.

We aim to have numbers as small as possible. Locant is the term used for the number which describes the position of the substitute group, e.g. the ‘2′ in 2-chlorobutane is the locant.

Sometimes there are two or more side chains e.g. a methyl group and a chlorine attached to a pentane. In these cases, names are written alphabetically, a separate number is needed for each side chain or group and hyphens are used to separate numbers and letters.

When there are two or more of the same side chains or substituent groups, these must also be given numbers separated by commas to show where each one is located.

Number the positions of double bonds if applicable.

Alkenes and other compounds have double bonds. These must be indicated with numbers. The number goes in the middle of the original root name e.g. butene, pentene.

It is worthwhile learning the other functional groups that can be added on.They have varying priorities.

The letter ‘e’ is removed when there are two vowels together e.g. propanone rather than propaneone. The ‘e’ isn’t removed when it is next to consonant, e.g. propanenitrile isn’t propannitrile.

When compounds contain two different, one is named as part of the unbranched chain and the other is named as a substituent. Which way round this goes depends on the priority.

Happy studying guys!

Halogenoalkanes

Halogenoalkanes are a homologous series of saturated carbon compounds that contain one or more halogen atoms. They are used as refrigerants, solvents, flame retardants, anaesthetics and pharmaceuticals but their use has been restricted in recent years due to their link to pollution and the destruction of the ozone layer.  

They contain the functional group C-X where X represents a halogen atom, F,Cl, Br or I. The general formula of the series is CnH2n+1X.

The C-X bond is polar because the halogen atom is more electronegative than the C atom. The electronegativity decreases as you go down group 7 therefore the bond becomes less polar. Flourine has a 4.0 EN whereas iodine has a 2.5 EN meaning it is almost non-polar.

The two types of intermolecular forces between halogenoalkane molecules are Van Der Waals and permanent dipole-dipole interactions. As the carbon chain length increases, the intermolecular forces (due to VDWs) increase as the relative atomic mass increases due to more electrons creating induced dipoles. Therefore the boiling point of the halogenoalkanes increases since more forces must be broken.  

Branched chains have lower boiling points than chains of the same length and halogen because the VDWs are working across a greater distance and are therefore weaker.

When the carbon chain length is kept the same, but the halogen atom is changed, despite the effect of the changing polar bond on the permanent dipole-dipole interactions, the changing VDWs have a greater effect on the boiling point. Therefore as RAM increases, the boiling point increases meaning an iodoalkane has a greater boiling point than a bromoalkane if they have the same carbon chain length.  

Halogenoalkanes are insoluble or only slightly soluable in water despite their polar nature. They are soluble in organic solvents such as ethanol and can be used as dry cleaning agents because they can mix with other hydrocarbons.

Summary

Halogenoalkanes are saturated carbon compounds with one or more halogen atoms. Their general formula is CnH2n+1X, where X is a halogen. Their functional group is therefore C-X.

They are used as refrigerants, solvents, pharmaceuticals and anaesthetics but have been restricted due to their link to the depletion of the ozone layer.

C-X bonds are polar due to the halogen being more electronegative than the carbon. The polarity of the bond decreases down group 7.

Van der Waals and permanent dipole-dipole interactions are the intermolecular forces in halogenoalkanes.

When carbon chain length increases, boiling points increase due to RAM increasing and the number of Van Der Waals increasing too.

In branched halogenoalkanes, Van Der Waals are working across a greater distance therefore attraction is weaker and boiling points are lower than an identical unbranched chain.

When the halogen is changed, the boiling point increases down the group due to the effect of a greater RAM  - more VDWs mean more intermolecular forces to break.  

Halogenoalkanes are insoluble in water but soluble in organic solvents like ethanol.

Bonus: free radical substitution reactions in the ozone layer

Ozone, O3, is an allotrope of oxygen that is usually found in the stratosphere above the surface of the Earth. The ozone layer prevents harmful rays of ultraviolet light from reaching the Earth by enhancing the absorption of UV light by nitrogen and oxygen. UV light causes sunburn, cataracts and skin cancer but is also essential in vitamin D production. Scientists have observed a depletion in the ozone layer protecting us and have linked it to photochemical chain reactions by halogen free radicals, sourced from halogenoalkanes which were used a solvents, propellants and refrigerants at the time.  

CFCs cause the greatest destruction due to their chlorine free radicals. CFCs – chloroflouroalkanes – were once valued for their lack of toxicity and their non-flammability. This stability means that they do not degrade and instead diffuse into the stratosphere where UV light breaks down the C-Cl bond and produces chlorine free radicals.

RCF2Cl UV light —> RCF2● + Cl●

Chlorine free radicals then react with ozone, decomposing it to form oxygen.

Cl● + O3 —> ClO● + O2

Chlorine radical is then reformed by reacting with more ozone molecules.

ClO● + O3 —-> 2O2 + Cl●

It is estimated that one chlorine free radical can decompose 100 000 molecules of ozone. The overall equation is:

2O3 —-> 3O2

200 countries pledged to phase of the production of ozone depleting agents in Montreal, leading to a search for alternatives. Chemists have developed and synthesised alternative chlorine-free compounds that do not deplete the ozone layer such as hydroflurocarbons (HFCs) like trifluromethane, CHF3.

SUMMARY

Ozone, found in the stratosphere, protects us from harmful UV light which can cause cataracts, skin cancer and sunburn.  

Ozone depletion has been linked to the use of halogenoalkanes due to their halogen free radicals.  

CFCs were good chemicals to use because they have low toxicity and were non-flammable. The fact they don’t degrade means they diffuse into the stratosphere. 

Chlorine free radicals are made when CFCs are broken down by UV light.

These go on to react with ozone to produce oxygen.

Chlorine free radicals are then reformed by reacting with more ozone.

It is a chain reaction that can deplete over 100 000 molecules of ozone.

There is a 200 country ban on their use and scientists have developed alternatives like hydrofluorocarbons to replace them

Happy studying!

Metallic Bonding

A short one to finish off my first ever mini-series on bonding – ionic, covalent and finally metallic. There are metallic and metallic compounds and elements but for the A Level exam, we must look at the bonding within metals themselves. Don’t worry – I saved the easiest to last!

Metals are most usually solid so have particles packed close together. These are in layers which mean that the outer electrons can move between them rather than being bound to particular atoms. These are referred to as delocalised electrons because of this.

It’s pretty common knowledge that metals are good conductors of heat and electricity and it’s these delocalised electrons that give them this property.

Metals are therefore without their electrons so become positive ions. The metallic bond is actually the attraction between delocalised electrons and positive metal ions in the lattice. And that’s pretty much metallic bonding, you just need to know the properties of metals which are touched upon at lower levels of education.

These are the properties of metals:

1.       High melting points

Metals have large regular structures with strong forces between the oppositely charged positive ions and negative electrons, meaning these must be overcome to melt the metal – this requires a large amount of heat energy. Transition metals tend to have higher melting points than the main group metals because they have large numbers of d-shell electrons which can become delocalised creating a stronger metallic bond. Melting points across a period increase because they can have progressively more delocalised electrons: Na+, Mg 2+ and Al 3+ for example.

2.       Heat conductivity

Heat is conducted if particles can move and knock against each other to pass it on. Delocalised electrons allow this to happen.  Silver is a particularly good conductor of heat.

3.       Electrical conductivity

Delocalised electrons can carry charge and move, the two requirements of electrical conductivity. Current can flow because of these delocalised electrons.

4.       Ductile and malleable

Metals can be stretched and hammered into shape, making them ideal for things such as wires. Layered lattices mean that layers can slide over each other without disrupting the bonding – it is all still held together by the delocalised electrons and their strong attraction to the positive metal ions.

5.       High densities

Being a solid, metal ions are packed closely together so they have a high density, which makes them ideal for musical instrument strings. These can withstand the frequency of vibration whilst also being thinner.

 SUMMARY

Metals are  solid so have particles packed close together. These are in layers which mean that the outer electrons can move between them rather than being bound to particular atoms. These are referred to as delocalised electrons because of this.

Metals are therefore without their electrons so become positive ions. The metallic bond is actually the attraction between delocalised electrons and positive metal ions in the lattice. 

Metals have high melting points.

Metals have large regular structures with strong forces between the oppositely charged positive ions and negative electrons, meaning these must be overcome to melt the metal – this requires a large amount of heat energy. Transition metals tend to have higher melting points than the main group metals because they have large numbers of d-shell electrons which can become delocalised creating a stronger metallic bond. 

Metals conduct heat.

Heat is conducted if particles can move and knock against each other to pass it on. Delocalised electrons allow this to happen.

Metals have good electrical conductivity

Delocalised electrons can carry charge and move, the two requirements of electrical conductivity. Current can flow because of these delocalised electrons.

Metals are ductile and malleable.

Metals can be stretched and hammered into shape, making them ideal for things such as wires. Layered lattices mean that layers can slide over each other without disrupting the bonding – it is all still held together by the delocalised electrons and their strong attraction to the positive metal ions.

Being a solid, metal ions are packed closely together so they have a high density. 

 Happy studying!