Physical Properties
of Water
Water is the strangest common substance on Earth. The fact that it sticks to itself, dissolves almost anything, holds heat better than rock, and floats when it freezes — none of that is normal. Almost every weird thing about it traces back to a single fact about a single tiny molecule.
Every winter, some lakes and ponds in the highveld go below 0 °C. Yet in spring, the fish are still alive. Why don't they freeze solid with the water around them?
Because ice floats. When the surface freezes, it forms an insulating lid on top — and the liquid water below stays around 4 °C, where fish can survive. If ice sank, lakes would freeze from the bottom up. Every fish in every cold-climate lake on Earth would die.
Almost every other liquid on the planet gets denser as it freezes. Water doesn't. Why? Because of one weird fact about the water molecule itself — and that one fact is the source of nearly everything water does.
One Bent Molecule. Eleven Strange Properties.
Water sticks to itself. That's the whole story.
A water molecule is two hydrogen atoms attached to one oxygen atom — but the three atoms aren't in a straight line. They form a bent shape, and the oxygen pulls the shared electrons closer to itself. That makes one side of the molecule slightly negative and the other side slightly positive: a polar molecule.
Polar molecules are sticky. Each water molecule attracts every neighbour, forming weak bonds called hydrogen bonds. Those tiny bonds — millions of them, constantly forming and breaking — are why ice floats, why oceans buffer climate, why droplets bead, why kelp can grow tall, and why your blood can carry oxygen. Same one fact. Eleven big consequences.
Inside the Water Molecule
Five things to notice about H₂O. Tap each highlighted part of the diagram to see what it is and why it matters. The faded molecule on the right is a neighbour — connected by the dashed hydrogen bond.
Local connection — Algoa Bay's mild summers
Walk the Donkin Reserve in Gqeberha at 4 pm in January. It's warm, but it's not extreme — usually around 24–26 °C, even on a peak summer day. Drive 200 km inland to Cradock or Graaff-Reinet at the same time, and the air can be 38 °C. The difference is one property of water: heat capacity. The ocean off Algoa Bay can absorb huge amounts of heat without changing temperature much, and that mass of cool water keeps the coastal air mild. Inland, no ocean, no buffer. The same property is why two-thirds of the heat from climate change is currently absorbed by the ocean — buying us time, but not forever.
Eleven Properties, One Source
Every property below is a consequence of polar bonding and hydrogen bonds between water molecules. Tested on every paper.
Three states
The only molecule on Earth that exists naturally as solid, liquid, and gas.
~ ice, sea, sea spray.Cohesion
Water sticks to itself. Highest cohesion of any non-metallic liquid.
~ raindrops are round; lakes have a smooth surface.Surface tension
The surface acts like elastic skin. Light insects can walk on it.
~ pond skaters; bluebottles riding the swell.Adhesion
Water sticks to other things. That's why a leaf or your wetsuit gets "wet".
~ kelp fronds glisten; dive masks fog.Capillary action
Water defies gravity in narrow spaces — adhesion + cohesion together.
~ mangroves push water up their roots; sponges suck it in.Viscosity
Resistance to flow. Water flows easily. Warmer = less viscous.
~ fish glide through it; cold water is "thicker".Universal solvent
Dissolves more substances than any other liquid. Salts, gases, nutrients.
~ sea water carries oxygen, salt, and life-sustaining minerals.Transparency
Light penetrates clear water for tens of metres.
~ kelp photosynthesises 30 m down off the Cape Peninsula.Incompressibility
Almost impossible to squeeze. Pressure is transferred, not absorbed.
~ a styrofoam cup at 4 000 m comes back thumbnail-sized.Salinity
The mass of dissolved salts in water. Average ocean ≈ 35 g/kg (35 ‰).
~ Antarctic sea ice freezing pushes salt down — drives global currents.Density
Pure water is densest at 4 °C. Ice is less dense than liquid water.
~ ice floats; cold water sinks; salty water sinks more.Heat, and the Energy Hidden Inside It
Of all common substances, water has the highest heat capacity. It can soak up enormous amounts of energy with barely any temperature change — and release it slowly. That's why oceans are climate machines. This section is where most of the marks in this topic live.
First: heat is not temperature
These two get swapped in exam answers constantly, and it costs marks. Temperature is how hot something is — the average kinetic energy of its molecules. Heat is the total thermal energy that moves from one thing to another.
The clean test: a kettle of boiling water and a bath of warm water. The kettle is at 100 °C and the bath at 40 °C, so the kettle has the higher temperature. But the bath holds far more water, so it contains far more heat energy — enough to warm a room all evening, which the kettle cannot do. Higher temperature ≠ more heat.
Specific heat capacity — the number CAPS wants
Specific heat (also called specific heat capacity) is the heat required to raise the temperature of one unit mass of a substance by one degree. For water that is 4.184 joules per gram per °C — which is the same thing as 1 calorie per gram per °C, because a calorie is defined as the energy that does exactly that job.
The two units, and why both appear in your notes
1 calorie = the heat needed to raise 1 g of water by 1 °C. 1 calorie = 4.184 joules. The joule is the SI unit; the calorie is the older one, and the Two Oceans booklet uses it throughout. Both are correct — just never mix them inside one calculation.
Compare with a metal. Water needs 4.184 J to warm 1 g by 1 °C. Copper needs only about 0.385 J — so the same energy that warms a gram of water by 1 °C would warm a gram of copper by nearly 11 °C. That is why a metal spoon in the sun burns your hand and the sea does not.
Your CAPS notes round copper to 0.3 J, which gives a ratio of about 14× instead of 11×. Either is accepted — the examinable point is that water's value is many times higher than a metal's, and the figure to remember for water is 4.184 J/g/°C.
Hydrogen bonds resist changes of state. To melt 1 g of ice takes 80 calories — even though the temperature stays at 0 °C the whole time. To boil 1 g of water takes 540 calories — even though the temperature stays at 100 °C. Those bonds have to be broken before the molecules can move apart. The energy used to break them is called latent heat — "hidden" heat that doesn't change the temperature reading.
The heating curve — read this graph until it is boring
Take 1 g of ice at −20 °C and add heat steadily. Plot temperature against heat added and you get the single most examinable graph in this topic. Tap each of the five numbered stages. Watch what happens to the line at 0 °C and 100 °C — it goes flat. You are adding energy and the temperature is not moving.
~ the two flat steps are latent heat · the graph is drawn to scale, so the boiling step really is that much longer
What an examiner will ask you about this graph
Almost every question on this curve is one of four. (1) "Explain why the line is horizontal at 0 °C and 100 °C" — because the energy is breaking hydrogen bonds, not raising kinetic energy, so there is no temperature change. (2) "Name the energy absorbed at each flat step" — latent heat of fusion at 0 °C, latent heat of vaporisation at 100 °C. (3) "Why is the second plateau so much longer?" — melting only loosens the lattice enough for molecules to slide past each other; boiling has to break the hydrogen bonds completely and separate the molecules into a gas. That takes 540 cal against 80. (4) "What state(s) exist at point X?" — on a plateau, both states are present at once.
The six changes of state — name them and give an example
CAPS asks for changes of state and examples. Water is the only substance on Earth that occurs naturally in all three states, so every one of these happens in the sea or the sky above it.
Melting (fusion)
Solid → liquid. Absorbs 80 cal/g.
~ Antarctic sea ice breaking up in spring.Freezing (solidification)
Liquid → solid. Releases 80 cal/g.
~ pack ice forming off Antarctica each winter.Evaporation (vaporisation)
Liquid → gas. Absorbs 540 cal/g.
~ the sea surface feeding moisture into the air over Algoa Bay.Condensation
Gas → liquid. Releases 540 cal/g.
~ sea fog rolling over Gqeberha; your dive mask fogging up.Sublimation
Solid → gas, skipping liquid entirely.
~ Antarctic surface ice vanishing into dry polar air.Deposition
Gas → solid, skipping liquid entirely.
~ frost forming on a research vessel's rigging.Coast vs inland — the same January day
Mild day, mild night
The Indian Ocean off Algoa Bay sits at ~21 °C. It absorbs heat all day without warming much, and releases it slowly overnight. Result: the air is buffered. Daytime peak ≈ 26 °C, nighttime low ≈ 18 °C. Eight-degree daily swing.
Scorching day, freezing night
No big body of water nearby. Soil and rock have low heat capacity — they heat up fast in the day and lose it fast at night. Daytime peak ≈ 38 °C, nighttime low ≈ 14 °C. Twenty-four-degree daily swing.
Why ice floats — the one fact that keeps fish alive
When water cools below 4 °C, hydrogen bonds force the molecules into an open lattice with extra space between them. The same number of molecules now takes up more volume — so ice is less dense than liquid water. This is the opposite of almost every other substance. Frozen lakes get an insulating lid of ice, with liquid water (around 4 °C) underneath. Without this property, the polar oceans would freeze from the seabed up, and complex life as we know it would never have evolved.
Water in the Air
CAPS attaches a note to the heat-capacity concept: "(Refer to the Water Cycle)". Here is why. The ocean does not just store heat — it moves heat, by evaporating in one place and condensing in another. Latent heat is the vehicle.
Evaporation cools. Condensation warms. That is the water cycle moving energy around the planet.
When a water molecule evaporates, it has to break its hydrogen bonds to escape the surface — and it takes that energy with it. The molecules left behind have less energy on average, so the surface it left is now cooler. This is evaporative cooling, and it is exactly why sweating works and why the sea surface sheds heat.
Later, somewhere else, that vapour condenses back into liquid — and every joule it carried is released again as latent heat, warming the air around it. So a molecule can pick up energy off Algoa Bay and drop it a thousand kilometres inland. The water cycle is a heat-delivery system.
Humidity and dew point
Air can hold water vapour — but only so much, and warm air holds more than cold air. That single fact gives you both definitions CAPS requires.
Humidity
The amount of water vapour held in the air. Usually given as relative humidity — a percentage of the maximum that air could hold at that temperature. 100% means the air is full.
~ a muggy February afternoon in Gqeberha: high humidity, sweat won't evaporate, so you can't cool down.Dew point
The temperature to which air must be cooled for its water vapour to start condensing into liquid. At the dew point, relative humidity has reached 100% and the air can hold no more.
~ cool the air below its dew point and you get dew on the grass, mist, or sea fog.Why? Air is mostly nitrogen (N₂, molar mass 28) and oxygen (O₂, 32). Water is H₂O — molar mass 18. Every water molecule that joins the air displaces a heavier N₂ or O₂ molecule, so the mixture gets lighter. This is part of why humid air rises and builds thunderclouds.
Local connection — the sea fog that shuts Gqeberha's airport
Algoa Bay gives you a live dew-point demonstration several times a year. Warm, moist air sits over the ocean. It drifts across the cooler Benguela-influenced water and coastal land at night, and its temperature drops. The moment it falls below its dew point, the vapour it is carrying condenses into billions of tiny droplets — advection fog — and the bay, the harbour and sometimes the runway at Chief Dawid Stuurman International disappear.
Nothing was added to the air. No water arrived. The air was simply cooled past the temperature at which it could hold what it already had. That is the dew point, and you can watch it happen from the beachfront.
Density — and How to Calculate It
Density is the property that drives every current in the ocean, and CAPS lists density calculations as one of its investigations. You need to be able to do the arithmetic, not just define the word.
The formula, and the two rearrangements
Density is mass per unit volume. Its symbol is the Greek letter rho (ρ).
Units. Mass in grams (g), volume in cubic centimetres (cm³), density in g/cm³. Note that 1 cm³ = 1 mℓ, so a measuring jug reading in millilitres is already giving you cm³.
The benchmark you must know: pure water has a density of 1 g/cm³ at 4 °C. That is not a coincidence — the gram was originally defined as the mass of 1 cm³ of water. Everything else in the ocean is measured against that 1.
Worked example 1 — is this sample fresh water or seawater?
A sample has a mass of 51.25 g and a volume of 50 cm³. Find its density and identify it.
ρ = m ÷ V = 51.25 ÷ 50 = 1.025 g/cm³
Pure water is 1.000 g/cm³, and average seawater is about 1.025 g/cm³. The sample is seawater. Notice how small the difference is — 2.5% — and yet that 2.5% is enough to sink one water mass beneath another and drive currents across an entire ocean basin.
Worked example 2 — going the other way
You need 2 000 cm³ (2 ℓ) of seawater for an experiment. What will it weigh?
m = ρ × V = 1.025 × 2 000 = 2 050 g = 2.05 kg
The same volume of pure water would be 2 000 g. The extra 50 g is the dissolved salt.
Three things change the density of water
Exam answers routinely name two and forget the third. CAPS lists all three.
1 · Temperature — big effect
Heat the water and the molecules move faster, hydrogen bonds break, the water expands — so density drops. Cool it and density rises… until 4 °C. Below that the open ice lattice starts forming and density falls again.
~ warm water floats on cold; cold water sinks. This is convection, and it is what stirs the ocean.2 · Salinity — big effect
Dissolved salt adds mass without adding much volume, so more salt = denser water. Average seawater ≈ 35 g of salt per kg of seawater (35 ‰), giving ≈ 1.025 g/cm³.
~ salinity also changes buoyancy (you float more easily in the sea) and forces marine life to osmoregulate.3 · Pressure — small effect
CAPS: "the effect of pressure on density is small." Because water is nearly incompressible, squeezing it barely packs the molecules closer. Do not answer that deep water is dense because of the pressure on it — it is dense because it is cold and salty.
~ a classic trap: pressure at 4 000 m is enormous, yet the water there is only ~2% denser than at the surface.Brine exclusion — how freezing seawater drives global currents
When seawater freezes, the growing ice crystal will not accept the salt — the lattice only has room for water molecules. The salt is squeezed out into the water below as concentrated brine, which is both very cold and very salty, and therefore very dense. It sinks.
Two consequences worth remembering. First, sea ice is mostly fresh — melt it and you can drink it, which surprises people. Second, that sinking brine off Antarctica is the engine of the global deep-water circulation that ventilates every ocean basin on Earth, including the water off Algoa Bay.
The correct figure is 35 g of salt per kilogram of seawater — written 35 g/kg, or 35 ‰ ("parts per thousand"), or 35 PSU. It is a ratio of masses, not a density.
In an exam: write 35 g/kg or 35 ‰. If a question hands you "35 g/cm³", answer it as salinity anyway — the marker wants the salinity concept, and the typo is theirs, not yours.
Pressure, and What It Does to a Diver
CAPS singles this one out: "When a person is diving, the pressure of water increases by 1 atmosphere for every 10 m of depth. This fact is of great importance to scuba divers." If you dive, you already feel this in your ears. Here is the physics behind it.
+1 atmosphere for every 10 m you descend. And you start at 1, not 0.
An atmosphere (atm) is the pressure the whole column of air above you exerts at sea level — about 1 kg per square centimetre. Standing on the beach you are already under 1 atm.
Seawater is roughly 800 times denser than air. So a mere 10 m column of seawater weighs as much as the entire ±500 km of atmosphere above it. Ten metres down, you have doubled the pressure on your body. That is the whole idea in one sentence.
Incompressible water, compressible everything-else
Here is the part that trips people up. All that pressure and the water itself barely changes at all — water is nearly incompressible. What gets crushed is anything containing air.
CAPS gives you the comparison directly: water is incompressible "unlike a diver's wetsuit (neoprene), which is compressed with depth." Neoprene is foamed rubber full of gas bubbles. At 30 m those bubbles are at a quarter of their surface volume, so the suit gets measurably thinner — and a diver loses both insulation and buoyancy on the way down. It is why divers add weight for a shallow dive and take it off for a deep one.
The famous demonstration: a styrofoam cup lowered to 4 000 m comes back the size of a thumbnail — the air in the foam is crushed out of it while the seawater around it is essentially unchanged. The cups pictured in your booklet were sent down from the SA Agulhas II by Seamester students in 2018 — a South African ship, on a South African research programme.
Viscosity and Dissolving
Both of these are named explicitly in the CAPS depth statement for this topic, and both carry a full investigation. They are short sections in most notes and disproportionately common in exams.
Viscosity — a liquid's resistance to flow
Viscosity is the measure of a liquid's resistance to flow. It comes from internal friction — molecules of the moving fluid dragging against each other. The more the molecules grip one another, the more the liquid resists being pushed around.
CAPS's own comparison: syrup at room temperature has a high viscosity — it resists flowing, and pours in a slow thick rope. Water has a much lower viscosity — its molecular structure produces very little internal friction, so it flows freely. Honey and cooking oil sit at the syrup end; ethanol is even runnier than water.
So cold water is "thicker" than warm water, and that is not a trivial fact in marine science: cold polar water is more viscous, which helps hold plankton up in the water column. A tiny drifter sinks more slowly through cold, viscous water than through warm, runny water — one of several reasons cold seas are so productive.
⚠️ Careful: your Two Oceans booklet says viscosity increases when heated. That is the wrong way round. The rule above is correct.
Dissolving — why "universal solvent" is an exaggeration
Water dissolves more substances than any other liquid on Earth, which earns it the nickname universal solvent. But the nickname oversells it, and the reason why is pure hydrogen bonding.
Hydrophilic — "water-loving"
Polar substances like salt and sugar dissolve because their polarity is stronger than the cohesive hydrogen bonds between water molecules. The water molecules are pulled away from each other and surround the solute particles instead.
~ salts, sugars, oxygen and CO₂ — which is why seawater can carry nutrients and gases at all.Hydrophobic — "water-fearing"
Non-polar substances have no charged poles, so they cannot compete with water's hydrogen bonding. Water's cohesion effectively pushes them out. They do not dissolve — they separate.
~ oil floating on the sea after a spill; the waxy coating on a kelp frond.So the honest statement is: water dissolves more substances than any other liquid — but it does not dissolve everything. If it did, an oil spill would simply disappear.
Two things you can actually see
Beading up
On a surface water does not stick to, cohesion wins and the droplet pulls itself into the smallest possible surface area — a bead. This is surface tension made visible.
~ rain beading on a waxed hull, or on a kelp frond.Meniscus
In a narrow tube, adhesion drags the edge of the water up the walls while cohesion holds the middle back — so the surface curves. That curve is the meniscus.
~ always read a measuring cylinder from the bottom of the meniscus, at eye level.In the Lab — the CAPS Investigations
CAPS sets five investigations for this topic. Every single one can be done at a kitchen table — the official materials list is salt, household liquids of different viscosities, dish-washing liquid, food colouring, warm water and ice. Nothing here needs a laboratory.
🔬 The five investigations
- Change of state, and examples. Take ice from freezer to boiling and record the temperature every 30 seconds — then plot your own version of the heating curve above and find the plateaus. 🟢 kitchen
- Density of water: cold versus warm. The food-colouring demonstration — warm dyed water and cold dyed water released into clear water, showing which rises and which sinks. 🟢 kitchen
- Density calculations. Measure the mass and volume of fresh water and of salt water you mix yourself, and calculate ρ = m/V for each. 🟢 kitchen
- Viscosity exploration. Race equal volumes of water, cooking oil, dish-washing liquid and syrup down a slope — then repeat one of them warm to test the temperature effect. 🟢 kitchen
- Surface tension exploration. How many drops of water will sit on a coin before it bursts? Then add one drop of dish-washing liquid and watch the surface tension collapse. 🟢 kitchen
Each one comes with a method, a results table to fill in, the analysis questions an examiner would ask, and a memo. If you cannot get a piece of equipment, every investigation also supplies a realistic set of results so you can still do the graphing and the analysis — which is the part that carries the marks.
📄 Practical sheet: Open the Topic 6 practical →
Key Vocabulary
Tap each card to flip it. CAPS names seven definitions for this topic outright — heat capacity, latent heat, latent heat of fusion, latent heat of vaporisation, dew point, humidity, and the boiling and freezing points of water. All seven are below, together with the nine other terms this topic runs on. Sixteen cards is a lot; this is an eight-hour block, the joint-longest in your year.
A molecule with uneven electron distribution — one side slightly negative, the other slightly positive. Water's bent shape makes it polar; CO₂ (linear) is not.
A weak attraction between the positive H of one water molecule and the negative O of another. Constantly forming and breaking — and the source of nearly every special property of water.
The attraction between like molecules. Water has the highest cohesion of any non-metallic liquid — that's why droplets form and lakes have a smooth surface.
The attraction between unlike substances. Why water "sticks" to glass, leaves and skin. Adhesion plus cohesion produces capillary action.
The elastic skin at the surface of water, caused by surface molecules being pulled inward by cohesion. Lets pond skaters walk on it; pulls droplets into spheres.
The spontaneous flow of water into a narrow space — a thin tube, a sponge, a plant root. Adhesion pulls water up the walls; cohesion drags more water along.
The "hidden" heat absorbed or released when water changes state, with no change in temperature. 80 cal/g for ice → liquid (fusion), 540 cal/g for liquid → steam (vaporisation).
Mass per unit volume (g/cm³), ρ = m/V. Pure water is densest at 4 °C, where it is 1 g/cm³; ice is less dense and floats. Salty water is denser than fresh — a key driver of ocean currents.
The heat required to raise the temperature of one unit mass of a substance by one degree. Water: 4.184 J/g/°C = 1 cal/g/°C — about 11× copper's. The highest of any common substance, and the reason the ocean buffers climate.
The energy required to change a substance from solid to liquid with no change in temperature. For water: 80 cal/g at 0 °C. The flat step at 0 °C on the heating curve. Released again, unchanged, when water freezes.
The energy required to change a substance from liquid to gas with no change in temperature. For water: 540 cal/g at 100 °C — nearly 7× the latent heat of fusion, because boiling must break the hydrogen bonds completely.
The temperature to which air must be cooled before its water vapour begins to condense into liquid. At the dew point relative humidity is 100%. Cool air below it and you get dew, mist, cloud — or Algoa Bay's sea fog.
The amount of water vapour present in the air. Relative humidity expresses it as a percentage of the maximum that air could hold at that temperature — and warm air can hold more than cold air.
Pure water: freezes at 0 °C, boils at 100 °C (at sea-level pressure). Seawater: freezes at about −2 °C, because dissolved salt obstructs the ice lattice forming.
A liquid's resistance to flow, caused by internal friction between its moving molecules. Syrup: high. Water: low. Heating lowers viscosity. Cold water is more viscous — which helps hold plankton up in the water column.
The mass of dissolved salts in a body of water. Average seawater ≈ 35 g/kg (35 ‰). It raises density, changes buoyancy, lowers the freezing point, and forces marine organisms to osmoregulate.
Quick Check
Question 1
Well done!
Move on to Topic 7 when you're ready.