Chapter 1–6: Water Properties and Solutions (Vocabulary)

Polar nature of water and hydrogen bonding

Water is a polar molecule due to unequal sharing of electrons (polar covalent bonds).
- Oxygen is more electronegative than hydrogen, giving the oxygen end a partial negative charge and the hydrogen ends a partial positive charge.
When two water molecules come together, they form hydrogen bonds:
- The negative (O) end of one molecule is attracted to the positive (H) end of another molecule.
- This bonding is responsible for many of water’s unique properties.
This polarity and hydrogen bonding lead to cohesion (water–water attraction) and adhesion (water–other substances attraction).

Cohesion, adhesion, and surface tension

Cohesion: attraction between like molecules (water–water).
Adhesion: attraction between unlike substances (water–cell walls, glass, etc.).
Surface tension:
- Defined as the measure of how difficult it is to stretch or break the surface of a liquid.
- Cohesion among water molecules creates a higher surface tension, effectively forming an invisible film on the surface.
- Analogy: gravy or other films on a surface illustrate surface tension – the surface behaves like a thin skin that resists breaking.
The combination of cohesion and adhesion explains phenomena like capillary action in plants (water moves up from roots to leaves).
In plants: cohesion causes water molecules to stick to one another; adhesion helps water climb against gravity by attaching to xylem cell walls; evaporation from leaves pulls the chain of water molecules upward (cohesion-tension mechanism).

Temperature, heat, and specific heat capacity

Water moderates temperature because it can absorb or release large amounts of heat with only a small change in its own temperature.
Specific heat capacity (definition): the amount of heat required to raise the temperature of 1 g of a substance by 1 °C.
- General formula: q = m \, c \, \Delta T
- For water: c_{\text{water}} \approx 4.184\ \frac{\text{J}}{\text{g} \cdot {}^{\circ}\text C}} \approx 1\ \frac{\text{cal}}{\text{g} \cdot {}^{\circ}\text C}
Calorie (historical unit in this context): 1\ \text{cal} = 1\ \frac{\text{g} \cdot {}^{\circ}\text C}{1} \text{(to raise 1 g of water by 1 °C)}
Beach example (beach temperature contrast): sand heats up quickly under solar radiation, while water remains cooler due to its high specific heat.
Heat capacity implications: high specific heat minimizes temperature fluctuations, benefiting aquatic life and organisms in fluctuating environments.
Evaporative cooling (energy loss during phase change):
- Latent heat of vaporization for water: L_v \approx 2260\ \frac{\text{J}}{\text{g}}
- Heat to vaporize 1 g of water: q = m \ L_v
- Evaporation cools the remaining surface; helps stabilize temperatures of organisms and bodies of water.
Calorie definition (revisited): 1\ \text{cal} = 1\ \frac{\text{g} \cdot {}^{\circ}\text C}{1}

Evaporation, evaporative cooling, and real-world implications

Evaporation: liquid water changing to a gas, requiring energy input.
Evaporative cooling: the surface cools as liquid water absorbs heat to become gas; the leaving vapor carries energy away, reducing the temperature of the remaining liquid.
Biological relevance: evaporative cooling helps regulate body temperatures in organisms; examples include moisture evaporation on skin or surfaces.
Ecological/climate relevance: water’s ability to absorb heat and buffer temperature helps stabilize climates and aquatic ecosystems.

Ice, density, and the frozen state

Ice is less dense than liquid water due to its hydrogen-bonded lattice structure, so ice floats.
- In ice, hydrogen bonds hold water molecules in a crystalline lattice with more open space, making ice less dense (density of ice is about 0.92 g/cm³ vs liquid water ~1.00 g/cm³).
- This lattice is described as more open and stable, with molecules held in place rather than moving freely.
In liquid water, molecules move freely and hydrogen bonds continually form and break, allowing higher density.
Climate relevance: many scientists are concerned about global warming’s impact on ice sheets and glaciers (described image shows a glacier with a marked retreat from the 1980s to 2010s).

Solutions, solutes, and solvents

Solvent vs solute:
- Solvent: the substance in which a solute dissolves (water in many biological contexts).
- Solute: the substance being dissolved (e.g., table salt NaCl).
Dissolution of NaCl in water:
- NaCl(s) dissociates into Na⁺(aq) and Cl⁻(aq) when dissolved.
- Water molecules surround and stabilize the ions via hydration shells, effectively separating the ions.
- Representations:
- Process: \mathrm{NaCl}{(s)} \rightarrow \mathrm{Na^+}{(aq)} + \mathrm{Cl^-}_{(aq)}
- In solution: \mathrm{NaCl}{(aq)} \text{ → } \mathrm{Na^+}{(aq)} + \mathrm{Cl^-}_{(aq)}

Acids, bases, and pH

Ions and charge: a charged particle (ion) has gained or lost electrons.
Hydrogen ions and hydroxide ions:
- Acids increase the concentration of H⁺ in solution; bases increase OH⁻ or decrease H⁺.
- In aqueous solutions, H⁺ is often represented as hydrated forms such as H₃O⁺; OH⁻ is a hydroxide ion.
pH scale basics:
- pH is a measure of hydrogen ion concentration: \mathrm{pH} = -\log_{10} [\mathrm{H^+}]
- Scale ranges from 0 to 14; 7 is neutral; below 7 is acidic; above 7 is basic (alkaline).
- The slide notes emphasize the relationship: as [H⁺] increases, pH decreases (more acidic).
- The arrow on the pH scale is conceptually shown as pointing downward for increasing acidity (higher H⁺) and upward for decreasing acidity (lower H⁺).
Water as a solvent: its polarity allows it to stabilize ions and polar molecules through hydration.

Connections and practical implications

Recurring theme: the polar nature of water and hydrogen bonding underlie cohesion, adhesion, surface tension, heat capacity, and solvent properties.
Biological relevance:
- Water transport in plants relies on cohesion and adhesion to move water from roots to leaves against gravity.
- Evaporative cooling helps organisms regulate temperature.
- The high specific heat and high heat capacity of water stabilize climates and aquatic ecosystems.
Environmental relevance:
- Ice vs liquid water densities influence buoyancy and seasonal water column mixing in lakes and oceans.
- Melting ice sheets and glaciers relate to global warming and sea-level changes.

Quick reference: key formulas and definitions

Specific heat and heat transfer:
- Definition: q = m \ c \ \Delta T
- For water: c_{\text{water}} \approx 4.184\ \frac{\text{J}}{\text{g} \cdot {}^{\circ}\text C} \approx 1\ \frac{\text{cal}}{\text{g} \cdot {}^{\circ}\text C}
Heat to raise 1 g of water by 1 °C (definition of a calorie): 1\ \text{cal} = 1\ \frac{\text{g} \cdot {}^{\circ}\text C}{1}
Latent heat of vaporization (evaporation): q = m \ Lv,\quad Lv \approx 2260\ \frac{\text{J}}{\text{g}}
Density relationships:
- Ice density: \rho_{\text{ice}} \approx 0.92\ \frac{\text{g}}{\text{cm}^3}
- Liquid water density: \rho_{\text{water}} \approx 1.00\ \frac{\text{g}}{\text{cm}^3}
pH and hydrogen ion concentration:
- \mathrm{pH} = -\log_{10} [\mathrm{H^+}]

Terms to memorize

Polar covalent bonds; electronegativity; partial charges (δ− on O, δ+ on H)
Hydrogen bond; hydrogen-bond network
Cohesion; adhesion; surface tension
Hydration shell; solvation; dissolution
Solvent vs solute
Specific heat; heat capacity; calorie
Evaporative cooling; latent heat of vaporization
Density of ice vs water; buoyancy
pH; acidity vs basicity; [H⁺], [OH⁻]
Hydration in ionic solutions (e.g., NaCl in water)