Water, Acids, Bases, and pH

Core Concepts

Water's polarity

A water molecule (H2O) consists of one oxygen atom covalently bonded to two hydrogen atoms. Oxygen is highly electronegative — it attracts shared electrons much more strongly than hydrogen does. This creates a polar covalent bond: the oxygen end of the molecule carries a partial negative charge, and each hydrogen carries a partial positive charge. The molecule is V-shaped (bent), so the charges are distributed asymmetrically, making water a polar molecule.

Hydrogen bonding in water

Because of its polarity, each water molecule can form hydrogen bonds with up to four neighboring water molecules. At any given moment in liquid water, molecules are constantly forming and breaking hydrogen bonds. This dynamic network of hydrogen bonds is responsible for water's distinctive properties.

Cohesion and adhesion

Cohesion is the attraction between molecules of the same substance. Water molecules cohere to each other through hydrogen bonding. Cohesion produces surface tension — a "skin" at the water's surface that allows small insects to walk on water.

Adhesion is the attraction between molecules of different substances. Water adheres to polar surfaces, such as the walls of plant xylem vessels. The combination of cohesion and adhesion enables water to move upward against gravity in narrow tubes (capillary action), which is essential for water transport in plants.

High specific heat

Specific heat is the amount of energy required to raise the temperature of 1 gram of a substance by 1 degree Celsius. Water has an unusually high specific heat because much of the added energy goes into breaking hydrogen bonds rather than increasing molecular motion. This means water resists temperature changes, helping organisms maintain stable internal temperatures and moderating Earth's climate.

Evaporative cooling

When water evaporates, the molecules with the highest kinetic energy escape as vapor, leaving behind molecules with lower average kinetic energy — and thus a lower temperature. This evaporative cooling is why sweating cools the body and why evaporation from leaves helps protect plants from overheating.

Ice density

Most substances are denser as solids than as liquids. Water is an exception: ice is less dense than liquid water, so it floats. This occurs because, as water freezes, the hydrogen bonds lock molecules into a crystalline lattice that holds them farther apart than in the liquid state. Because ice floats, it insulates the water below, allowing aquatic life to survive beneath the frozen surface of lakes and oceans.

Water as a solvent

Water's polarity makes it an excellent solvent — a substance that dissolves other substances. When ionic compounds (like NaCl) are placed in water, the partial charges on water molecules surround and separate the ions, dissolving the compound. Polar molecules (like sugars) also dissolve in water because they can form hydrogen bonds with water molecules.

Hydrophilic and hydrophobic substances

• Hydrophilic ("water-loving") substances are polar or ionic and dissolve readily in water. Examples: salts, sugars, many proteins.

• Hydrophobic ("water-fearing") substances are nonpolar and do not dissolve in water. Examples: oils, fats, hydrocarbons.

The hydrophobic effect — the tendency of nonpolar molecules to cluster together in water — is a major force driving protein folding, membrane formation, and other biological processes.

Water dissociation and pH

A small fraction of water molecules dissociate into hydrogen ions (H+) and hydroxide ions (OH−):

H2O ⇌ H+ + OH−

In pure water at 25°C, the concentration of H+ equals the concentration of OH−, each being 1 × 10−7 moles per liter (M). The pH scale is a logarithmic measure of H+ concentration:

pH = −log[H+]

• Neutral: pH = 7 ([H+] = [OH−] = 10−7 M)

• Acidic: pH < 7 (higher [H+] than [OH−])

• Basic (alkaline): pH > 7 (lower [H+] than [OH−])

Because pH is logarithmic, each unit change represents a tenfold change in H+ concentration. A solution of pH 4 has ten times the H+ concentration of pH 5 and one hundred times that of pH 6.

Acids, bases, and buffers

• Acid: A substance that increases the H+ concentration of a solution (donates H+). Example: HCl dissociates into H+ and Cl−.

• Base: A substance that decreases the H+ concentration of a solution (accepts H+ or releases OH−). Example: NaOH dissociates into Na+ and OH−; the OH− combines with H+ to form water.

• Buffer: A substance that resists changes in pH by accepting or releasing H+ as needed. Buffers typically consist of a weak acid and its conjugate base. The carbonic acid–bicarbonate buffer system helps maintain blood pH near 7.4. Biological fluids contain buffers to keep pH within ranges that proteins and other molecules can tolerate.

Why pH matters for biology

Most biological processes are exquisitely sensitive to pH. Enzymes have optimal pH ranges; even modest deviations can denature proteins and disrupt metabolism. Human blood pH is maintained between approximately 7.35 and 7.45; values outside this range can be life-threatening. The interior of lysosomes is maintained at around pH 5, optimal for the digestive enzymes within. Different cellular compartments maintain different pH values suited to their functions.

ELI Example

Water is like dancers with magnetic hands — constantly pairing positive to negative, briefly connecting, then moving on. Heat the dance floor, and the fastest dancers break away (evaporation), taking energy with them and cooling those who remain.

Do Not Confuse

Lab Link

Understanding pH and buffers is critical for enzyme laboratories. Most enzymes have a pH optimum, and activity drops sharply outside that range. When investigating enzyme activity, controlling pH is as important as controlling temperature. Buffers are used to maintain consistent pH across experimental treatments. Additionally, pH indicators are used in several biomolecule tests and in photosynthesis investigations that track CO2 consumption (CO2 dissolving in water produces carbonic acid, lowering pH).

High-Yield Memory Anchors

• Water = polar molecule → hydrogen bonds → all water's special properties.

• Cohesion = water sticking to water. Adhesion = water sticking to other things.

• Ice floats because H-bonds in ice hold molecules farther apart.

• pH is logarithmic: each step = 10× change.

• Buffers resist pH change by absorbing or releasing H+.

Quick Check

Q1 (Foundational): List four properties of water that result from hydrogen bonding, and briefly explain the biological significance of each.

Q2 (Application): A student measures the pH of a solution and finds it to be 3. She then dilutes the solution tenfold. What is the expected new pH? Explain your reasoning.

Q3 (Comparison/Reasoning): Compare how NaCl (table salt) and a droplet of oil behave when placed in water. Explain the behavior of each in terms of polarity and the types of chemical bonds involved.

Quick Check Answers

A1: Four properties: (1) Cohesion — enables water transport in plants via capillary action. (2) High specific heat — stabilizes temperatures in organisms and ecosystems. (3) Evaporative cooling — protects organisms from overheating (sweating, transpiration). (4) Ice floats — insulates bodies of water, protecting aquatic life in winter. (Other valid answers: adhesion, surface tension, solvent properties.)

A2: pH 3 means [H+] = 10−3 M. A tenfold dilution reduces [H+] to 10−4 M. The new pH would be 4, because pH = −log(10−4) = 4. Each unit increase in pH corresponds to a tenfold decrease in H+ concentration.

A3: NaCl dissolves in water because it is an ionic compound. Water's polar molecules surround the Na+ and Cl− ions, separating them and keeping them in solution (the ions become hydrated). NaCl is hydrophilic. Oil does not dissolve in water because oil consists of nonpolar hydrocarbon molecules. There are no partial charges for water to interact with, and the oil molecules are excluded from the hydrogen-bond network of water. Oil is hydrophobic. The oil molecules cluster together (the hydrophobic effect) rather than mixing with water.

Chapter Summary

Water's polarity enables hydrogen bonding, producing cohesion, adhesion, high specific heat, evaporative cooling, and the property that ice floats. Water is an excellent solvent for polar substances. pH measures H+ concentration logarithmically; buffers resist pH changes.

Common Mistakes

Mistake: "Individual hydrogen bonds between water molecules are permanent."

Reality: In liquid water, hydrogen bonds constantly form and break. Each bond lasts only a tiny fraction of a second, but at any moment, many water molecules are hydrogen-bonded to their neighbors.

Mistake: "Water's high specific heat means it gets hot quickly."

Reality: The opposite — high specific heat means water resists temperature change. It takes a relatively large amount of energy to raise water's temperature.

Mistake: "A solution with pH 6 is only slightly more acidic than one with pH 7."

Reality: Because pH is logarithmic, pH 6 has ten times the H+ concentration of pH 7 — a substantial difference.

Mistake: "Neutral pH means no H+ or OH− ions are present."

Reality: Even in pure water at pH 7, H+ and OH− are present at 10−7 M. Neutral means [H+] = [OH−], not that neither ion exists.