The Oceans – SL notes.

O1 - They that go down to the sea in ships.

·         Useful things from the sea;

·         Fish and shellfish for food (for thousands of years!)

·         1500BC – 800BC – Tyrian Purple – a brilliant purple dye extracted from the shells of a particular species of marine snail and used to dye clothes of wealthy and powerful! (chemically similar to indigo)

·         Seaweed – used as food in Japan, as a source of chemicals or fertilisers. and also as a source of alkali (by burning it to produce kelp – Na₂CO₃, K₂CO₃) – which was used in soap and glass-making.

·         Today iodine is extracted from seaweed as well as a thickening agent called alginate used in foods.

·         Other used include maintaining the head on beer!, medical dressings, paper, textiles and pharmaceuticals.

·         Seaweed is a concentrates source of K, I, and Na.

·         Marine snails can contain 38% by mass of Br compared to 0.2% in sea water.

·         Br is rare on land so is also extracted from sea-water (minerals – M1)

·         Sea-water contains ionic compounds in solution – i.e. a mixture of ions.

·         The relative abundance of these ions in sea-water is remarkably consistent wherever it is sampled (see fig4 page 239)

·         Fig 5 page 239 shows where the ions in sea water come from;

·         Ca²⁺, Mg²⁺,CO₃^2-, and SiO₄^2-, are dissolved from soil by rivers etc. and carried to sea.

·         Cl, Br and S compounds are dissolved from  lava or gases released from volcanoes on the floor of the oceans (e.g. along the mid Atlantic ridge). When water seeps into cracks in solidified lava it dissolves minerals and forms hydrothermal vents, which are very rich sources of elements like chlorine, bromine and sulphur.

Salt of the Earth.

·         The balance of ions in the sea is kept constant by a complicated geochemical cycle we are just beginning to understand.

·         The composition of sea salt never varies but the concentration varies from place to place depending on several factors e.g;

·         Climate – in hot places evapⁿ is faster so concentration is higher.

·         Estuaries – fresh water entering sea reduces concentration.

·         Most abundant ions in sea water are Na⁺(aq) and Cl^-(aq)

·         These ions are present in all land creatures and are essential to our health – hence salt has always been regarded as a valuable commodity.

·         Salt was originally obtained from the sea for eating, flavouring or preserving food, but nowadays it is obtained from underground by mining.

Obtaining salt from sea water;

1)      Filter then heat in large stainless steel pans.

2)     As boils impurities rise to top and can be skimmed off.

3)     Simmer – calcium sulphate crystallises out first – forms scale on sides of vessel.

4)     As concⁿ ­ NaCl start to crystallise (as solubility is low)

5)     Heating is stopped before all water is evaporated and NaCl crystals collected. – This is so ‘bitter’ tasting magnesium salts do not crystallise out and spoil flavour (they are less abundant and more soluble).

·         Fish and minerals have always been taken from sea – we must be careful not to take too much!

·         We put our waste into sea – again we should be careful – heavy metal ions poison fish and nitrate pollution leads to rapid algae growth.

·         Sea is a potential source of medicines – some sea creatures have developed poisons as a defence mechanism – and these chemicals could be useful.

SL O2 – Wider still and deeper.

·         As well as providing food the oceans also play a vital part in controlling the climate on Earth.

·         The vastness of the oceans and their changing nature makes them difficult to study.

·         Satellites help (see fig 13 page 243 and read SL)

·         The depth of the ocean varies greatly – landscapes beneath the ocean are more extreme than those above! (see fig14 page 244)

Unravelling a complex system.

·         Oceans are involved in global processes and cycles – here is one example.

·         The sulphur cycle was being studied in the 1960’s because of the problem of acid rain or acid deposition.

·         Oxidised sulphur compounds are among the chemicals involved in acid rain.

·         The sums done from their measurements did not add up – there was a missing link in the sulphur cycle!

·         It is now known that dimethyl sulphide [(CH₃)₂S] is produced by seaweeds and other marine algae – this compound is volatile and gets into the atmosphere where it is oxidised.

·         This can account for 25% of all acidic pollution over Europe at certain times of the year (April and May)

·         This is one source of pollution we have no control over!

O4 – A Safe Place to Grow.

Storing Carbon Dioxide.

·         In gases solubility ­ as Temp ¯ and high pressure facilitates dissolving (see table 3 and fig 31 pages 252 and 253.

·         Low temp and high pressure encourage CO₂ to dissolve in the oceans, helping lower [CO₂] in the atmosphere.

·         Uptake of CO₂ by oceans is speeded up by marine life – see fig 32 page 253.

·         Table 3 shows that CO₂ is more soluble than O₂ or N₂. This is because CO₂ contains polar C=O bonds which can form hydrogen bonds with water

·         An equilibrium is set up;

CO₂(g)    «  CO₂(aq)

·         Also CO₂ can chemically react with water;

CO₂(aq)   +   H₂O (l)     «    HCO₃^-(aq)     +    H⁺(aq)

AND           HCO₃^-(aq)    «      H⁺(aq)    + CO₃^2-(aq)

·         Thus H⁺,  CO₃^2- and HCO₃^- are produced by these inter-linked reactions.

·         35-50%  of extra CO₂ in atmosphere from combustion of fuels is thought to be removed this way.

·         Deep ocean currents remove surface CO₂ and store it for 100’s of years.

·         Maximum removal occurs in coldest regions where CO₂ solubility is higher.

·         Thus currents, chemistry and marine life are effective at CO₂ removal.

Sinking shells.

·         Shells are involves in processes with CO₂.

·         3 equations above can be combined;

CO₂(g)  +  H₂O(l)   «   2H⁺(aq)    +    CO₃^2-(aq)

·         Le Chetelier’s priciple says that if H⁺ or CO₃^2- are removed then more CO₂ will dissolve to restore equilibrium.

·         Adding alkali would remove H⁺ and so encourage CO₂ to dissolve but this is not very practical in the oceans!

·         But marine life remove CO₃^2- to make their shells – this has the same effect on CO₂ – i.e. encourages it to dissolve.

·         Billions of years ago [CO₂] in the atmosphere was approx 35%

·         Once photosynthesis had evolved marine life had lots of CO₂ to use.

·         Shell production flourished – evidence of this can be seen in limestone and chalk rocks on land today.

·         CaCO₃ is insoluble in sea water so is suitable material for shells, but it is slightly soluble in pure water (it is a sparingly soluble solid). An equilibrium is set up;

CaCO₃(aq)     «     Ca²⁺(aq)      +      CO₃^2-(aq)

·         The equilibrium constant is called the solubility product (K_sp)

K_sp = [Ca²⁺(aq)]  [CO₃^2-(aq)]  =  5.0 x 10^-9 mol²dm^-6

·         When Ca²⁺(aq)  and CO₃^2-(aq) are mixed in solⁿ one of two things can happen;

·         CaCO₃(s) ppt forms – (if [Ca²⁺(aq)]  [CO₃^2-(aq)] > K_sp)

·         The ions simply remain in solution – (if [Ca²⁺(aq)]  [CO₃^2-(aq)]  < K_sp)

·         CaCO₃ is used to build shells because the [Ca²⁺(aq)] and [CO₃^2-(aq)] are high enough at the surface of the sea for the shells to stay insoluble. The shells are actually in equilibrium with the sea water!

·         At the bottom of the ocean pressure is higher and T is lower – so K_sp is greater and CaCO₃ is more soluble.

·         There is a continual drift of CaCO₃ downwards - marine snow – comprising of remains of dead organisms and waste from live ones!

·         But there are no shells on the ocean floor – they have all dissolved (See Fig 36 page 256) and deep sea creatures do not have protective coats of CaCO₃.

·         Also our limestone deposits must have formed in shallow seas – they could not have formed in deep ocean.

Life on Earth.

·         Early life on Earth evolved in the sea – and stayed there for most of history!

·         Life on land is fairly recent in geological time-scales (see table 4 page 258)

·         There are still bacteria (inside tube worms) in the sea which gain energy by using sulphate ions to oxidise methane and hydrogen sulphide from geothermal vents, these are similar to life as it was 3.5 billion years ago!!

·         The evolution of cyanobacteria allowed photosynthesis to take place and started O₂ production. The O₂ was used up by reducing agents in the ocean – just as well as cyanobacteria cannot tolerate oxygen!  Close relatives of these are still alive today – hiding away in places where there is no oxygen.

·         The atmosphere has altered drastically through history – but the oceans have changed very little!

Keeping things Steady.

·         The pH of the oceans has remained close to 8 for millions of years.

·         Carbon dioxide is only a weak acid, and the equilibrium;

CO₂(aq)  +   H₂O   «   H⁺(aq)    +    HCO₃^-(aq)

       Lies well over to the left hand side.

·         (NB – sometimes a solⁿ of CO₂(aq) is given formula H₂CO₃ – carbonic acid).

·         If a small amount of alkali is added then more CO₂ will react with the water to replace the H⁺ lost and restore the equilibrium – hence seawater can maintain a constant pH when small amounts of alkali added.

·         But what if the concⁿ of CO₂ in the atmosphere rose? Then the concⁿ of CO₂(aq) in the sea will also rise and the Concⁿ of H⁺ should also then rise – hence the pH of the sea will fall – But this does not happen – the sea acts as a buffer solution.

·         Buffers are solutions of weak acids and their salts. (see CI 8.3)

·         For sea water – a solution of carbonic acid – a supply of HCO₃^- ions are needed to remove any increase in H⁺ ions.

·         Figure 42 shows the processes which prevent the oceans from becoming acidic. The shells, chalk and limestone in the sea are the reservoir of anions needed to prevent changes in acidity.

·         If CO₂ in atmosphere rose dramatically solid CaCO₃ would dissolve to provide anions to remove extra H⁺ produced in oceans. Most of C would be in form of HCO₃^- and dissolved CO₂. The sea would be like a mixture of Perrier water and bicarbonate of soda and the white cliffs of Dover would dissolve!

·         Limestone could not therefore form in the early atmosphere of Earth – the earliest deposit are only 2 billion years old, by which time much of the CO₂ from the early atmosphere had been used up.

·         Ion exchange between H⁺ ions in sea water and Na⁺ and K⁺ ions in clay sediments does help control the pH of the oceans, but this can only occur at the bottom of the ocean.

·          The Carbon dioxide/calcium carbonate system made of the linked equilibria shown in fig 42 responds rapidly to changes in ion concentration and is the main system which keeps the pH of the oceans stable.