Water: particle theory, solubility, density and pressure

Kinetic particle theory and the changes of state in water

Everything in this topic comes from one model: all matter is made of particles that are always moving, and how much energy those particles have decides how they are arranged. In a solid the particles are held in a fixed, regular arrangement by strong forces of attraction and can only vibrate about fixed positions. In a liquid the particles are still close together and still attracting one another, but they have enough energy to slide past each other, so a liquid keeps its volume but takes the shape of its container. In a gas the particles are far apart, the forces between them are almost nothing, and they move quickly and randomly in straight lines, colliding with each other and with the walls. A gas has no fixed shape or volume, fills any container, and can be compressed because most of the space between its particles is empty.

Heating means transferring energy to the particles, and every change of state is that energy changing what the particles can do. Ice melts because the vibrations of its particles become violent enough to overcome the attractions holding the lattice together, so the particles break free of their fixed positions and can slide past one another. Water boils because particles throughout the liquid gain enough energy to overcome the remaining attractions completely and escape as a gas. Cooling does the reverse: gas particles lose energy, slow down, and the attractions between them pull them together into droplets (condensation); liquid particles lose energy until they can only vibrate, and the substance freezes.

Two details earn marks that most students miss. First, while ice is melting the temperature stays at 0 °C until the last ice has gone, and while water is boiling it stays at 100 °C at normal atmospheric pressure, because the energy being supplied is going into breaking attractions between particles rather than into making them move faster. Second, evaporation happens from the surface of liquid water at any temperature, not just at boiling point, and it is faster when the water is warmer, the surface is larger and the air above is moving, because the fastest particles escape and take their energy with them.

Notice what does not change during any of this. The water molecules themselves are identical in ice, in the sea and in a cloud; only their energy, their spacing and their freedom to move have changed. That is what makes a change of state a physical change, and it is why you can evaporate sea water and get the same water back.

• Solid: particles closely packed in a fixed regular lattice, vibrating only; fixed shape and volume.
• Liquid: particles close but irregular, able to slide past each other; fixed volume, takes the shape of the container.
• Gas: particles far apart, moving fast and randomly; no fixed shape or volume, easily compressed.
• Melting: energy overcomes the attractions holding the lattice, so particles can move past each other.
• Boiling: particles throughout the liquid have enough energy to escape as a gas; temperature stays constant during the change.
• Evaporation: the fastest surface particles escape below boiling point - faster when warm, windy or with a large surface area.
• Condensation and freezing: particles lose energy, attractions pull them together, and the arrangement becomes more ordered.

Atoms and ions: what sea water is actually made of

An atom has a tiny, dense nucleus containing protons (positive charge) and neutrons (no charge), surrounded by electrons (negative charge, almost no mass) arranged in shells. The number of protons is the atomic number and it decides which element the atom is; in a neutral atom the number of electrons equals the number of protons, so the positive and negative charges cancel out. The shells fill in a fixed order - 2 in the first, then 8, then 8 - and it is the outer shell that matters for bonding, because every atom is trying to reach a full outer shell.

An ion is an atom that has gained or lost electrons, so it is no longer neutral. Losing an electron leaves more protons than electrons and gives a positive ion: a sodium atom (2, 8, 1) loses its single outer electron and becomes Na+. Gaining an electron gives a negative ion: a chlorine atom (2, 8, 7) gains one and becomes Cl-, which now has the full outer shell of 8.

Sea water is a mixture, not a compound. It contains water itself, dissolved salts such as sodium chloride, magnesium sulfate and calcium carbonate, dissolved gases such as oxygen and carbon dioxide, and suspended and dissolved organic matter. It is a mixture because each substance keeps its own identity, the proportions vary from place to place and with depth, and the components can be separated by physical means - evaporation, distillation or filtration - without a chemical reaction. Evaporating a litre of sea water leaves the salts behind, which is exactly how sea salt is produced.

That distinction is worth stating precisely in an exam. A compound is two or more elements chemically joined in fixed proportions, like the water molecule itself. A mixture is two or more substances together in no fixed proportion, which is what a bucket of sea water is.

Covalent bonding: how the atoms inside a water molecule are held together

A covalent bond is a shared pair of electrons between two atoms. The shared pair is attracted to both nuclei at once, and that attraction is what holds the atoms together. In a water molecule, oxygen has 6 electrons in its outer shell and needs 2 more to be full; each hydrogen has 1 electron and needs 1 more. So the oxygen atom shares one pair of electrons with each of two hydrogen atoms, forming two single covalent bonds and giving the molecule the formula H2O. On a diagram the shared pair is drawn as two dots (or a dot and a cross) in the overlap between the circles, or simply as a line joining the atoms.

The shape matters as much as the bonding. Oxygen has two lone pairs of electrons that are not shared, and these push the two O-H bonds away from each other, so the molecule is bent (V-shaped) rather than straight. Remember that shape, because it is the reason one end of the molecule is slightly negative and the other slightly positive, and therefore the reason water can hydrogen bond at all.

You are also expected to recognise covalent molecules from diagrams. The clues are that the atoms are joined by shared pairs or lines, that there are no charges and no lattice, and that the elements involved are usually non-metals. Water is bent with two O-H bonds; carbon dioxide is linear, with a double bond to each oxygen (O=C=O); oxygen gas is two oxygen atoms joined by a double bond (O=O); sulfur dioxide has one sulfur joined to two oxygens by shared pairs and is bent; glucose (C6H12O6) is a six-membered ring of five carbon atoms and one oxygen atom, with hydrogen atoms and -OH groups attached.

Covalent molecular substances are usually gases, liquids or low-melting solids, because the forces between one molecule and the next are weak even though the bonds inside each molecule are strong. Water looks like an exception - it is a liquid over most of the range of ocean temperatures, not a gas - and the next section explains why: hydrogen bonds hold its molecules to one another.

Ionic bonding and the salts dissolved in sea water

Ionic bonding works by transfer rather than sharing. A sodium atom has one electron in its outer shell and a chlorine atom has seven. Sodium loses its outer electron to become Na+, with the stable arrangement 2, 8; chlorine gains that electron to become Cl-, with the arrangement 2, 8, 8. Both ions now have full outer shells. The positive sodium ion and the negative chloride ion then attract one another electrostatically, and because each ion attracts ions of the opposite charge in every direction, the result is a giant ionic lattice rather than a small molecule. The formula NaCl tells you only the ratio - one Na+ for every Cl- - not that the substance is made of pairs.

The properties follow straight from that structure. Ionic substances have high melting points, because melting means breaking strong attractions between ions in every direction. Solid sodium chloride does not conduct electricity, because its ions are locked in position, but molten sodium chloride does, and so does a solution of it, because the ions are then free to move and carry charge. That is why sea water conducts electricity and distilled water hardly does.

You are expected to identify ionic substances from diagrams, including sodium chloride and calcium carbonate. The clues are separate ions drawn with their charges (often in square brackets), a regular lattice, and no shared pairs and no lines joining atoms into molecules. Calcium carbonate is ionic too, made of Ca2+ ions and carbonate ions (CO3 2-), again one of each, so its formula is CaCO3.

Three salts appear by name in the syllabus and you should be able to give both the name and the formula of each without hesitating: sodium chloride, NaCl; magnesium sulfate, MgSO4, made of Mg2+ and SO4 2- ions; and calcium carbonate, CaCO3. All three are found in sea water, and calcium carbonate is also the material of coral skeletons, mollusc shells and the chalk and limestone of the sea bed.

Hydrogen bonding and the three properties it explains

A water molecule is polar. Oxygen holds the shared electrons more strongly than hydrogen does, so in each O-H bond the oxygen end is slightly negative and the hydrogen end is slightly positive. Because the molecule is bent, the whole molecule has a negative end (the oxygen) and a positive end (the two hydrogens). Now put two water molecules together: the slightly positive hydrogen of one is attracted to the slightly negative oxygen of its neighbour. That attraction is a hydrogen bond.

Each water molecule can form hydrogen bonds through its two hydrogens and its two lone pairs, so in ice it is bonded to four neighbours. A single hydrogen bond is only about a twentieth as strong as the covalent O-H bond inside the molecule, but there are enormous numbers of them, they break and re-form constantly, and together they are strong enough to give water its unusual behaviour. Keep the distinction clear: covalent bonds hold the atoms inside a water molecule together; hydrogen bonds hold one water molecule to the next.

Solvent action. Because water molecules are polar, they gather around ions and around other polar molecules - the slightly negative oxygens pointing at a Na+ ion, the slightly positive hydrogens pointing at a Cl- ion - and carry them away into solution. Substances that cannot hydrogen bond with water, such as oils, fats and waxes, are not pulled apart in this way and do not dissolve. This is why the sea can hold so much dissolved salt while oil spilled on it floats in a slick.

Density. In ice, hydrogen bonds hold every molecule at a fixed distance from four neighbours, building an open lattice with large spaces in it. When the ice melts, the lattice collapses and the molecules pack closer together, so liquid water is denser than ice. Almost every other substance is denser as a solid; water is not, which is why ice floats and why a pond freezes from the top down. Ice has a density of about 917 kg m-3 and sea water about 1025 kg m-3. Specific heat capacity. To raise the temperature of water, the hydrogen bonds between the molecules must be stretched and broken as well as the molecules made to move faster, so a large amount of energy produces only a small rise in temperature. Water therefore has a very high specific heat capacity and resists temperature change, which is why the sea stays cooler than the land on a hot day and warmer than it on a cold night.

• Why the bond forms: oxygen pulls the shared electrons towards itself, so hydrogen is slightly positive and oxygen slightly negative; the positive hydrogen of one molecule is attracted to the negative oxygen of another.
• How many: each molecule can hydrogen bond through its two hydrogens and two lone pairs, so four neighbours in ice.
• Strength: weak individually compared with a covalent bond, but so numerous that together they dominate water's behaviour.
• Solvent action: polar water molecules surround and separate ions and other polar molecules; non-polar oils and fats do not dissolve.
• Density: the open hydrogen-bonded lattice of ice makes ice less dense than liquid water, so ice floats.
• Specific heat capacity: energy is used breaking hydrogen bonds, so water needs a lot of energy to change temperature and resists change.

Solutions, solubility and salinity

The vocabulary is worth four easy marks and is often muddled. The solute is the substance that dissolves - the salt. The solvent is the liquid it dissolves in - the water. The solution is the mixture that results, and it is uniform throughout, so a sample taken from the top has the same concentration as one from the bottom. Solubility is the maximum mass of a solute that will dissolve in a given volume of solvent at a given temperature, usually quoted in grams per 100 g of water, and it is a property of that particular solute in that particular solvent at that temperature.

How sodium chloride dissolves. When salt is added to water, the polar water molecules surround the ions at the surface of the lattice: the slightly negative oxygen ends are attracted to the Na+ ions and the slightly positive hydrogen ends to the Cl- ions. That attraction is stronger than the attraction holding the ions in the lattice, so ions are pulled out of the lattice one at a time and become surrounded by water molecules, spread evenly through the liquid. The ions are still ions - they have been separated, not reacted - which is why the solution conducts electricity and why evaporating the water gives the original salt back.

Temperature and the solubility of salts. For most solid salts, raising the temperature of the water increases solubility, because the water molecules move faster and can break the lattice apart more quickly than ions can rejoin it. The size of the effect varies a lot: potassium nitrate becomes far more soluble in hot water, while sodium chloride's solubility changes only slightly, which is why heating water is not a good way to dissolve much more salt. A few salts, including calcium carbonate, are practically insoluble, and their solubility falls further as water warms - one reason coral reefs are built in warm water but do not dissolve away in it.

Salinity is the concentration of dissolved salts in sea water, and in this syllabus it is measured in parts per thousand (ppt), written with the symbol in the same way as a percentage but per 1000. Average open-ocean salinity is about 35 ppt, meaning 35 g of dissolved salts in every kilogram of sea water. Salinity is not the same everywhere, because it depends on the balance between water added and water removed. Evaporation removes pure water and leaves the salts behind, so it raises salinity - which is why the subtropics, where evaporation exceeds rainfall, are the saltiest open-ocean regions. Precipitation and surface run-off from rivers and land add fresh water and therefore lower salinity, which is why coastal water, estuaries and the wet tropics are less salty and why enclosed seas with large rivers, such as the Baltic, are brackish.

pH, and measuring it in the field

The pH scale runs from 0 to 14 and measures the concentration of hydrogen ions in a solution. A low pH means a high concentration of hydrogen ions and an acidic solution; a high pH means a low concentration of hydrogen ions and an alkaline solution. pH 7 is neutral, below 7 is acidic, above 7 is alkaline. The scale is logarithmic, so each whole unit represents a tenfold change in hydrogen ion concentration - a pH 5 solution has ten times the hydrogen ion concentration of a pH 6 solution. You are not required to do calculations with hydrogen ion concentration, but knowing that a small pH change is a big chemical change stops you writing that pH 8.0 and 8.3 are 'about the same'.

Typical sea water is slightly alkaline, around pH 8, because it is buffered by dissolved carbonate and bicarbonate. That is a narrow range of tolerance: marine organisms are adapted to it, and a fall of a few tenths of a pH unit is a significant change for anything that builds a calcium carbonate shell or skeleton.

You should be able to describe how pH is measured, and the three methods in the syllabus are not interchangeable. Litmus is the crudest: red litmus turns blue in an alkaline solution, blue litmus turns red in an acidic one. It tells you which side of neutral you are on and nothing more. Universal Indicator is a mixture of dyes that changes through a range of colours - red, orange, yellow, green, blue, purple - which you match to a colour chart to give an approximate pH, usually to the nearest whole unit. A pH probe or pH meter gives a numerical reading, often to one or two decimal places, and it is the method to choose when you need to detect a small change, compare samples precisely, or take a reading that does not depend on your judgement of a colour.

Two practical points come up in the paper. Calibrate a pH probe against buffer solutions of known pH before use and rinse it with distilled water between samples, or your readings will drift. And when you compare water samples, take the reading at the same temperature and as soon as possible after collection, because the pH of a water sample changes as it warms and equilibrates with the air.

Dissolved gases: oxygen's low solubility and what changes it

Oxygen has a low solubility in water. The oxygen molecule is a small, non-polar covalent molecule, so it cannot form hydrogen bonds with water and is not pulled into solution the way an ion is. Only a small amount dissolves - a few milligrams per litre in air-saturated water, which is far less oxygen than the same volume of air contains. That single fact shapes marine life: because there is so little oxygen in the water, every marine animal that breathes with gills is extracting oxygen from a dilute supply, and small changes in solubility matter enormously. Carbon dioxide behaves quite differently and is far more soluble, because it reacts with water as well as dissolving in it.

Temperature. As water gets warmer, gases become less soluble, because the dissolved gas molecules have more kinetic energy and escape from the surface faster than they dissolve back in. So warm water holds less oxygen than cold water, and the coldest polar seas are the most oxygen-rich. A fish in warm tropical surface water is living in water that holds perhaps half the oxygen of the Arctic at the same salinity, and its oxygen demand is higher because its metabolism is faster - which is why warm water plus a pulse of organic pollution can strip the oxygen out of a shallow bay and kill fish within hours.

Pressure. The solubility of a gas increases as pressure increases, so deeper water at higher water pressure can hold more dissolved gas, and water brought rapidly to the surface from depth may be supersaturated and fizz. Atmospheric pressure works the same way: a high atmospheric pressure pushes more gas into the surface water and a low pressure lets it escape, which is why oxygen readings at the surface drift with the weather. Salinity also matters: gases are less soluble in salt water than in fresh water at the same temperature, because the dissolved ions occupy water molecules that would otherwise have held gas, so sea water holds appreciably less oxygen than a freshwater lake at the same temperature.

Put the four together and you can predict where oxygen will be scarce. Deep water is under high pressure, which favours solubility, but it is cut off from the atmosphere and its oxygen is used up by respiration and by the decay of dead organisms, so oxygen falls to a minimum at mid-depths before rising again in the coldest bottom water. Warm, salty surface water in a sheltered lagoon on a still night, when photosynthesis has stopped but respiration continues, is the other danger zone. Intertidal animals cope by clamping their shells shut, using stored water, or switching to anaerobic respiration until the tide returns.

Density, floating ice and the layers of the ocean

Density is the mass of a defined volume of water, and it is calculated as density = mass / volume, with the mass in kg, the volume in m3 and the density therefore in kg m-3. Because the volume is often measured in cm3 in the laboratory, convert before you calculate: 250 cm3 is 250 x 10-6 m3, or 2.5 x 10-4 m3. If 250 cm3 of sea water has a mass of 256 g, then the density is 0.256 kg / 2.5 x 10-4 m3 = 1024 kg m-3 - and if you forget the conversion you will be out by a factor of a million.

Three factors change the density of sea water. Temperature is the most important in the surface ocean: as water warms the particles move faster and spread slightly further apart, so a given volume contains less mass and density falls. Salinity: dissolved ions add mass without adding much volume, so density rises as salinity rises - which is why the salty, cold water of the polar seas is the densest water in the ocean. Pressure: increasing pressure compresses water slightly, reducing its volume and raising its density, though because water is almost incompressible the effect is small compared with temperature and salinity over the depths of the ocean, and it matters mainly because it makes the very deep water denser still.

Ice floats because its density is lower than that of sea water - about 917 kg m-3 against about 1025 kg m-3 - because of the open hydrogen-bonded lattice described earlier. That has two consequences the syllabus asks for. As a thermal insulator, the floating ice and the snow on top of it separate the water beneath from the cold air above, greatly slowing heat loss, so the sea below the ice stays liquid and the ice does not keep thickening downwards; without it, lakes and polar seas would freeze solid from the bottom and kill almost everything in them. As a habitat, the ice is a surface to live on and under: ice algae grow on the underside where light penetrates, grazing krill and small crustaceans feed on them, fish and invertebrates shelter in the cracks and under the ice edge, and seals, penguins and polar bears use the ice as a platform to rest, breed and hunt. When the ice melts, that habitat disappears with it.

Density differences layer the ocean and keep the layers apart. At the surface, sunlight warms the water and wind and waves stir it, producing a surface (mixed) layer of roughly uniform temperature and salinity that is warm and therefore relatively light. Below it, temperature falls rapidly with depth through the thermocline; where salinity increases rapidly with depth, that layer is the halocline. Beneath both lies the deep ocean, cold, salty, dense and dark, at only a few degrees above zero, which has been out of contact with the atmosphere for a long time. Because warm water is lighter than cold water and fresh water lighter than salty water, these layers are stable and resist mixing - which is exactly why they persist. Mixing happens when something disturbs that stability: strong winds and winter storms deepen the mixed layer, and in high latitudes surface water that is chilled and made saltier by the formation of sea ice becomes dense enough to sink, carrying oxygen down and driving the deep circulation, while upwelling brings deep water back to the surface. Those density-driven movements are what distribute heat, oxygen and nutrients around the ocean and so help determine where organisms can live.