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O - THE OCEANS

The Oceans

Acids, Equilibria and Buffer Solutions

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Oceans of Energy

Energy, Entropy and Equilibrium

Entropy changes, the system and the surroundings

Properties of Water

Storing Carbon Dioxide

Solubility and Solubility Product

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The oceans play an essential part in the cycling of many chemicals throughout the earth.  Over 99% of all dissolved substances in seA^-water are ionic.  The most common dissolved ions are sodium, chloride, sulphate, magnesium, calcium and potassium.  Many of these ions come from the land.  They are leached from the soil by rainwater and washed into the sea by rivers.  The elements chlorine, bromine and sulphur however, are not obtained from the land.  These are obtained from lava found under the sea bed, generated by underwater volcanoes.  The sea is also a major source of food in the form of fish and other sea creatures.

Recently it has been realised that oceans play a major part in controlling the climate.  It has been found that 25% of all acidic pollution over French and German coasts is due to dimethyl sulphide which is produced by seaweeds and other marine algæ.  The dimethyl sulphide is volatile and is oxidised in the atmosphere to form acidic sulphur compounds.

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Oceans of Energy

Most of our energy comes from the Sun.  Each year, the Earth receives 3 x 10²⁴ Joules of solar energy.  Of this 47% is absorbed by the land and oceans, 23% is absorbed by the atmosphere and 30% is reflected.

If the Earth had no atmosphere or oceans, each part of it would settle down to a situation where the energy received from the Sun would be balanced by energy lost through radiation.  The presence of an atmosphere and oceans spreads out the heating effect of the Sun more evenly.  Currents in the sea and air take thermal energy from the tropics to the colder regions of the Earth.

The ocean/atmosphere is even more effective at spreading out energy.  Warm water not only circulates, but it can evaporate.  Energy is taken in when water evaporates and is given out when it condenses.  The tropics are cooled by evaporation, and currents in the atmosphere carry the water vapour to colder, high altitude regions where condensation releases energy.  High latitude regions receive more energy than is provided by the Sun alone.

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Energy, Entropy and Equilibrium

The simple molecular-kinetic theory explains the differences between solids, liquids and gases.

Solids.  The particles are in fixed positions with a regular lattice pattern.  The particles are close together and the volume depends only slightly on pressure and temperature.

Liquids.  The particles are moving around in a random arrangement.  The particles are close together and the volume depends only slightly on pressure and temperature.

Gases.  The particles are moving around in a random arrangement, widely separated.  The volume depends strongly on pressure and temperature.

The Ideal Gas Law:

tells us how p, V and T are related for a fixed number of moles, n, of gas.

If we have a gas in a can at room temperature and increase the temperature slightly, the molecules will gain a little extra energy.  On average the molecules will:

1. move more quickly - increase in translational energy,
2. rotate more rapidly - increase in rotational energy,
3. vibrate more violently - increase in vibrational energy.

Not every molecule gains the same amount of each of these energies.  Energy is quantised, and molecules are restricted to particular levels of these energies.  Molecules do not all have the same energy: they are spread out among the energy levels.  As molecules collide and exchange energy with each other, their energies change and they move up and down the energy levels.  The molecules are distributed among the energy levels in the way that gives the greatest entropy.  Entropy is a measure of the number of ways of arranging molecules and distributing their quanta of energy.  Entropy can also be visualised as a measure of the disorder of a system.  In general, gases have higher entropies than liquids, and liquids have higher entropies than solids.

Entropy is (commonly) given the symbol S; S is the entropy change for a process.  The Second Law of Thermodynamics says that for a spontaneous process to occur there must be an overall increase in the entropy, i.e. S > 0.

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Entropy changes, the system and the surroundings

If we consider the reaction between sodium and chlorine:

This involves a decrease in entropy, as the particles in a gas which are free to move become confined in an ionic solid.  We know that this reaction definitely occurs, and that it releases a considerable amount of energy, provided that there is a flame to supply the activation energy required for the reaction.  How can a change in which there is a decrease in entropy be spontaneous?

The answer is that we have only considered the entropy change of the system, S_sys.  We must also consider the change in entropy of the surroundings, S_surr.  Energy leaves the system to the surroundings during an exothermic reaction.  This produces a large increase in the entropy of the surroundings.  It is the total entropy change that determines whether the reaction is spontaneous or not.

To calculate the entropy change of the surroundings we need to know H^θ.

For the reaction between sodium and chlorine, H^θ = -411 kJ mol^-1, S_sys = -90.3J mol^-1K^-1.

Overall, the change in entropy is positive, hence this is a favourable process under standard conditions.

Similar principles apply to melting and freezing.

Both these changes are spontaneous: ice melts at 10^oC and water freezes at -10^oC of their own accord.  Both processes result in an increase in the total entropy change.

We can explain the processes of dissolving and crystallisation in a similar way.

If S_total = 0, there is no net change in either direction of a process: it is in a state of equilibrium.

At equilibrium S_total must be zero.

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Properties of Water

Water is unique when compared with similar compounds of low molecular mass.  Water has:

1. a high boiling point and enthalpy change of vaporisation
2. a greater specific heat capacity than almost any other liquid
3. a decrease in density when it freezes.

These properties can be explained in terms of hydrogen bonding.  The hydrogen bonds in water are particularly strong because there are two lone pairs of electrons on the oxygen atom, hence water is able to form two hydrogen bonds per molecule.  In order to change water into a vapour the hydrogen bonds have to be broken.  This requires quite a lot of energy: the latent enthalpy change of vaporisation.  Condensation on the other hand is an exothermic process.

Evaporation and condensation of water affect the temperature of different parts of the Earth in two ways:

• by transferring energy from low latitudes to high latitudes
• by warming the land through condensation of water which comes from oceans.

Water has a high specific heat capacity; it takes a lot of energy to raise its temperature which makes water is an excellent substance for absorbing and storing energy, again due to the hydrogen bonding.  Energy is used in overcoming the hydrogen bonds and is not available for raising the temperature.  Water is one of the best liquids for transporting energy, as shown by the movement of the Gulf Stream.

When water is cooled it contracts (this is normal for any liquid), however, below 4^oC it starts to expand, and there is a further expansion on freezing.  As it freezes it does so with an 'open' structure with the water molecules held in a tetrahedral arrangement by hydrogen bonds.  This increase in volume means there is a decrease in density which is why ice floats on water.

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Storing Carbon Dioxide

Carbon dioxide is soluble in water.  The solubility decreases with increasing temperature.  In fizzy drinks carbon dioxide is put in under pressure.  A high pressure and low temperature help to keep a drink fizzy.

Oceans also dissolve carbon dioxide.  Much of the excess carbon dioxide we release into the atmosphere from the combustion of fuels is absorbed by the oceans.  It is thought 35-50% is absorbed in this way.

An equilibrium is set up:

Also some carbon dioxide molecules react with water and are removed from the equilibrium.  More carbon dioxide therefore dissolves to maintain the equilibrium:

Any way of removing H⁺ or CO₃^2- ions from solution will cause more CO₂ to dissolve (the equilibria are shifted to the right).  As a result, through time, many marine organisms have evolved to build protective shells of insoluble calcium carbonate.

Hence shells provide another way of mopping up carbon dioxide and keeping the composition of the atmosphere constant.

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Solubility and Solubility Product

All chemical substances have a maximum solubility in water, depending upon the substance and the temperature of the solution.  The solubility of a salt is usually expressed in terms of the mass of a solute which will dissolve in 100g of the solvent at a particular temperature.  For sodium chloride the solubility is 36g per 100g of water, but for lead chloride it is only 0.99g per 100g of water.  Substances which are only very slightly soluble in a solvent are said to be sparingly soluble.

If an undissolved ionic compound is in contact with a saturated solution of its ions a solubility equilibrium is established.  The solid dissolves at the same rate that ions combine to precipitate fresh solid.

Consider a saturated solution of the sparingly soluble compound calcium carbonate in contact with an excess of solid calcium carbonate.

We can write an expression for the equilibrium constant:

Adding more solid will not cause the equilibrium to shift further to the right because the solution is saturated, so equilibrium is not affected by the amount of solid present.  It is an example of a heterogeneous equilibrium.  Therefore the mass of solid present can be taken as constant.  The equilibrium expression, therefore simplifies to:

K_sp stands for solubility product.

The concept of solubility product is only useful for sparingly soluble salts.  The fact that solubility product is constant for a given solute in water at a particular temperature provides a way of comparing the solubilities of sparingly soluble ionic substances.  A very small value for the solubility product of a substance implies that the substance is poorly soluble.

We can use K_sp to predict whether a precipitate will form from a solution.

Three situations exist:

• If [Ca²⁺(aq)][CO₃^2-(aq)]  <  K_sp, the solution is not saturated and more solute will dissolve.
• If [Ca²⁺(aq)][CO₃^2-(aq)]  =  K_sp, the solution is just saturated.
• If [Ca²⁺(aq)][CO₃^2-(aq)]  >  K_sp, precipitation will occur.

K_sp can be used to calculate the carbonate ion concentration in a saturated solution of calcium carbonate.

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