Campbell Biology · Chapter 3
Water and Life
pp. 46–57 · 3 sections
This chapter shows how one small, bent, polar molecule shapes the whole living world. Hydrogen bonds between water molecules explain why water sticks together, steadies temperatures, floats when frozen and dissolves so many things, and water's tendency to split into H⁺ and OH⁻ is the basis of pH. It's the foundation for Topic 1.1 and for everything later about membranes, enzymes and proton gradients.
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3.1 Why water is polar, and hydrogen bonds
pp. 46–47
This is the heart of Topic 1.1: be ready to explain how water's polarity leads to hydrogen bonds between molecules.
In the course: Topic 1.1 Structure of Water and Hydrogen Bonding (notes, videos and more questions)
Key points
- A water molecule (H₂O) is one oxygen atom joined to two hydrogen atoms by covalent bonds, and it's bent, with an angle of about 104.5° between the hydrogens.
- Oxygen is more electronegative than hydrogen, so the shared electrons spend more time near the oxygen. The oxygen end gets a partial negative charge (δ−) and each hydrogen gets a partial positive charge (δ+).
- Because the molecule is bent, those partial charges don't cancel out, so water as a whole is a polar molecule.
- A hydrogen bond is the pull between a δ+ hydrogen on one molecule and a δ− oxygen (or nitrogen) on another molecule or another part of a big molecule.
- Each water molecule can hydrogen-bond with up to four neighbors. In liquid water these bonds are weak and last only a tiny fraction of a second, but new ones form all the time.
- Hydrogen bonds are much weaker than the covalent bonds inside a molecule, but there are so many of them that together they explain almost all of water's special behavior.
Key terms (8)
- covalent bond
- A bond where two atoms share a pair of electrons. The O–H bonds inside a water molecule are covalent and strong.
- electronegativity
- How strongly an atom pulls shared electrons toward itself. Oxygen and nitrogen are high; hydrogen and carbon are lower.
- polar covalent bond
- A covalent bond where the electrons are shared unevenly, so one atom ends up slightly negative and the other slightly positive.
- nonpolar covalent bond
- A covalent bond where electrons are shared about equally, like the C–H bonds in fats and oils.
- partial charge
- A small charge (written δ+ or δ−) on part of a molecule caused by uneven electron sharing. It's smaller than the full charge on an ion.
- polar molecule
- A molecule with a slightly negative side and a slightly positive side because its partial charges don't cancel. Water is the classic example.
- hydrogen bond
- A weak attraction between a slightly positive hydrogen and a slightly negative oxygen or nitrogen. It forms between separate molecules or between different parts of one large molecule, like a protein or DNA.
- emergent property
- A property that shows up only when parts interact. One water molecule can't have surface tension; billions of hydrogen-bonded ones can.
Check yourself: 3.1 Why water is polar, and hydrogen bonds
4 questions on 3.1 Why water is polar, and hydrogen bonds. Pick an answer to see if you got it, and why.
Which statement best explains why a water molecule is polar?
A student builds a model showing a water molecule (H₂O) next to an ammonia molecule (NH₃), in which nitrogen carries a partial negative charge. Which feature of the model correctly represents a hydrogen bond between the two molecules?
A computer model follows one water molecule drifting alone in empty space, with no other molecules anywhere near it. Which property could this lone molecule still have?
Methane (CH₄, molecular mass 16 daltons) boils at about −162 °C, while water (H₂O, 18 daltons) boils at 100 °C. Which explanation best accounts for this large difference?
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3.2 Cohesion, heat control, floating ice and dissolving
pp. 47–52
Cohesion, adhesion, surface tension, specific heat and heat of vaporization are Topic 1.1; the mole and molarity math is background here, but molar concentration comes back in water potential (Topic 2.7).
In the course: Topic 1.1 Structure of Water and Hydrogen Bonding, Topic 2.7 Tonicity and Osmoregulation (notes, videos and more questions)
Key points
- Cohesion (water clinging to water) and adhesion (water clinging to other surfaces) both come from hydrogen bonds. In a plant, evaporation from leaves tugs on the water in the xylem, cohesion keeps that column from breaking, and water's attraction to the xylem walls helps hold it up against gravity. Surface tension, which is cohesion at the water–air boundary, makes the surface act like a thin skin that small insects can walk on.
- Thermal energy is the total kinetic energy of a sample's molecules, and heat is that energy moving from a warmer object to a cooler one. Temperature measures the molecules' average kinetic energy. One calorie raises 1 g of water by 1 °C, and 1 cal = 4.184 J.
- Water's specific heat is high (1 cal/g·°C) because a lot of the energy it takes in goes into breaking hydrogen bonds instead of speeding molecules up. That keeps the temperatures of organisms, lakes, oceans and coastlines fairly steady.
- Water's high heat of vaporization (about 580 cal per gram near room temperature) means evaporation carries off a lot of heat. Sweating and transpiration cool organisms this way.
- Ice floats because hydrogen bonds lock its molecules into an open lattice that is less dense than liquid water, which is densest near 4 °C. Floating ice works like a lid that shields the water underneath, so lakes don't freeze solid.
- Water dissolves ions and polar molecules (hydrophilic substances) by surrounding each one with a hydration shell, while nonpolar substances like oils are hydrophobic. A colloid forms when hydrophilic particles, such as very large proteins, are too big to make a true solution but stay evenly mixed in the water without settling.
- Concentration is counted in moles: 1 mol is 6.02 × 10²³ particles, and it weighs the molecular mass in grams. Molarity is moles of solute per liter of solution.
Key terms (15)
- cohesion
- Water molecules sticking to each other through hydrogen bonds. It helps hold a column of water together inside a plant's xylem.
- adhesion
- Water sticking to a different substance, like the walls of xylem cells or a glass tube. It helps water climb thin tubes against gravity.
- surface tension
- The tendency of a liquid's surface to resist being pushed in or pulled apart, so it acts like a thin elastic film. Water's is high because its surface molecules are held by hydrogen bonds to their neighbors.
- temperature
- How fast, on average, the molecules in something are moving (their average kinetic energy). A sparkler's spark is far hotter than a bathtub of warm water, yet the bathtub holds much more thermal energy.
- calorie (cal)
- The heat needed to warm 1 g of water by 1 °C. A food Calorie is really a kilocalorie (1,000 cal), and 1 cal equals 4.184 joules.
- specific heat
- The heat needed to warm or cool 1 gram of a substance by 1 °C. Water's is high, so it warms up and cools down slowly.
- heat of vaporization
- The heat needed to turn 1 g of a liquid into a gas. Water's is high because each escaping molecule has to break its hydrogen bonds.
- evaporative cooling
- The cooling of a surface as a liquid evaporates from it. The fastest, most energetic molecules leave first, so the ones left behind are cooler on average.
- solvent and solute
- In a solution, the solvent does the dissolving and the solute is what gets dissolved. In salt water, water is the solvent and salt is the solute. When water is the solvent, it's called an aqueous solution.
- hydration shell
- The cluster of water molecules that surrounds a dissolved ion or polar molecule, with their partial charges pointing toward it.
- hydrophilic
- 'Water-loving.' Describes ions and polar molecules that are attracted to water, whether or not they actually dissolve.
- hydrophobic
- 'Water-fearing.' Describes nonpolar substances, like oils and the tails of membrane lipids, that can't hydrogen-bond with water and don't mix with it.
- colloid
- A mixture in which particles too big to truly dissolve, like some large proteins, stay evenly spread through a liquid without settling out. The particles can be hydrophilic even though they don't dissolve.
- mole (mol)
- A counting unit for particles: 6.02 × 10²³ of them. A mole of a substance weighs its molecular mass in grams, so 1 mol of glucose is about 180 g.
- molarity (M)
- Concentration measured as moles of solute per liter of solution. A 1 M glucose solution has 1 mol (180 g) of glucose in every liter.
Check yourself: 3.2 Cohesion, heat control, floating ice and dissolving
4 questions on 3.2 Cohesion, heat control, floating ice and dissolving. Pick an answer to see if you got it, and why.
When the end of a thin glass tube is dipped into water, the water rises higher inside the tube than in a wider tube. Which explanation is best supported by the properties of water?
On a sunny day, 100 g of seawater and 100 g of dry beach sand each absorb 1,000 cal of heat from the sun. Water's specific heat is 1.00 cal/g·°C, and dry sand's is about 0.20 cal/g·°C. Assuming no heat is lost, about how much does the temperature of each sample rise?
On a hot, sunny afternoon, a plant closes its stomata (leaf pores) to save water, which sharply reduces evaporation from its leaves. Which change in the leaves is most likely, and why?
A student fills a plastic container with exactly 100 mL of liquid water at 4 °C and freezes it. The ice takes up about 109 mL. Which explanation best accounts for the increase in volume?
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3.3 Acids, bases, pH and buffers
pp. 52–56
No topic is only about pH, but the formula sheet gives you pH = −log[H⁺], and you use it with enzymes (Topic 3.2) and H⁺ gradients (Topic 3.5); buffers and acid rain are background, and ocean acidification is one of the human-caused disruptions in Topic 8.7.
In the course: Topic 3.2 Environmental Impacts on Enzyme Function, Topic 3.5 Cellular Respiration, Topic 8.7 Disruptions in Ecosystems (notes, videos and more questions)
Key points
- Water molecules occasionally split into a hydrogen ion (H⁺) and a hydroxide ion (OH⁻). The H⁺ is really attached to another water molecule as hydronium (H₃O⁺), but biologists just write H⁺. In pure water at 25 °C, [H⁺] = [OH⁻] = 10⁻⁷ M.
- An acid raises the H⁺ concentration of a solution. A base lowers it, either by grabbing H⁺ (like ammonia, NH₃) or by releasing OH⁻ that combines with H⁺ (like NaOH). Strong acids and bases split into ions completely in water; weak ones split only partly, and their reaction runs in both directions.
- In any watery solution at 25 °C, [H⁺] × [OH⁻] = 10⁻¹⁴, so knowing one tells you the other.
- pH = −log[H⁺]. A pH of 7 is neutral, below 7 is acidic and above 7 is basic. Each step of 1 pH unit is a tenfold change in [H⁺], so a solution at pH 5 has 100 times more H⁺ than one at pH 7.
- Most cells stay near pH 7, and human blood is held between about 7.35 and 7.45. A shift in pH changes the charges on proteins and can disrupt their shape, so enzymes may stop working.
- A buffer pairs a weak acid with its base. Add H⁺ and the base form binds it; remove H⁺ and the acid form gives some up, so the pH barely moves. In blood, carbonic acid (H₂CO₃) and bicarbonate (HCO₃⁻) form this kind of pair.
- Burning fossil fuels adds CO₂ to the air. Oceans absorb about a quarter of it, forming carbonic acid; surface ocean pH has dropped about 0.1 units since before industrial times (roughly 25–30% more H⁺), leaving less carbonate for corals and shell builders. Sulfur and nitrogen oxides from burning fuels cause acid rain.
Key terms (13)
- hydrogen ion (H⁺)
- A single proton with a +1 charge. In water it's always attached to a water molecule, but it's written H⁺ for short.
- hydronium ion (H₃O⁺)
- A water molecule that has picked up an extra H⁺. It's the form H⁺ actually takes in a watery solution.
- hydroxide ion (OH⁻)
- The ion left when a water molecule loses an H⁺. It has a −1 charge, and more of it makes a solution basic.
- dissociation of water
- The rare, reversible splitting of a water molecule into H⁺ (which joins another water molecule as H₃O⁺) and OH⁻. In pure water at 25 °C, each ion is at only 10⁻⁷ M.
- acid
- A substance that raises the H⁺ concentration of a solution, usually by releasing H⁺. Hydrochloric acid (HCl) is a strong acid.
- base
- A substance that lowers the H⁺ concentration of a solution, either by taking up H⁺ or by releasing OH⁻ that combines with H⁺.
- strong acid or base
- One that comes apart into ions completely in water, like HCl or NaOH. A weak acid or base comes apart only partly and reversibly.
- neutral solution
- A solution with equal H⁺ and OH⁻ concentrations. At 25 °C, its pH is 7.
- pH
- A number for how acidic a solution is: pH = −log[H⁺]. Lower pH means more H⁺, and each unit is a tenfold change.
- buffer
- A weak acid paired with its base. The base form binds added H⁺ and the acid form gives up H⁺ when it's scarce, so pH changes very little.
- carbonic acid–bicarbonate buffer
- The main buffer in your blood: H₂CO₃ ⇌ HCO₃⁻ + H⁺. It shifts left to soak up extra H⁺ and right to release H⁺ when it runs low.
- ocean acidification
- The drop in ocean pH as seawater absorbs extra CO₂ from the air. It lowers the carbonate that corals and shellfish need to build their skeletons and shells.
- acid precipitation
- Rain, snow or fog more acidic than normal rain (about pH 5.6), mostly from sulfur and nitrogen oxides released by burning fossil fuels.
Check yourself: 3.3 Acids, bases, pH and buffers
4 questions on 3.3 Acids, bases, pH and buffers. Pick an answer to see if you got it, and why.
After a meal, food dilutes the acid in the stomach. At 25 °C, a sample of stomach fluid taken after a meal has [H⁺] = 1 × 10⁻³ M. What are the sample's pH and [OH⁻]?
Pepsin, a protein-digesting enzyme in the stomach, works best near pH 2. Trypsin, which digests protein in the small intestine, works best near pH 8. The hydrogen ion concentration at pepsin's optimum is how many times the hydrogen ion concentration at trypsin's optimum?
Students added drops of 0.1 M NaOH, a strong base, to two beakers: one held 50 mL of distilled water, and the other held 50 mL of a phosphate buffer solution. They recorded the pH after every 5 drops. Distilled water: 7.0, 10.7, 11.0 and 11.2 after 0, 5, 10 and 15 drops. Buffer: 7.20, 7.21, 7.22 and 7.23 after 0, 5, 10 and 15 drops. Which conclusion is best supported by these data?
In blood, carbonic acid and bicarbonate are in equilibrium: H₂CO₃ ⇌ HCO₃⁻ + H⁺. During hard exercise, lactic acid from muscles adds H⁺ to the blood. How does this buffer system limit the change in blood pH?
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