2.2 Water

Water is the foundation of life on Earth. Its unique properties, like high specific heat capacity and cohesive forces, enable crucial biological processes. From maintaining stable temperatures to enabling capillary action in plants, water's characteristics are essential for living organisms.

Because water is polar, it acts as a universal solvent, dissolving many substances vital for life. This polarity allows for the transport of nutrients and waste in organisms. Understanding water's role in biological systems is key to grasping how life functions at a molecular level.

Properties and Roles of Water in Biological Systems

Properties of water for life

Water has several properties that make it uniquely suited to support living systems. These all trace back to one feature: the hydrogen bonds that form between water molecules.

• High specific heat capacity
  • Water requires a large amount of energy to change temperature. Specific heat capacity is the amount of energy needed to raise 1 gram of a substance by 1°C, and water's value (1 cal/g/°C) is unusually high compared to most liquids.
  • This helps organisms maintain stable internal temperatures. Your body, which is roughly 60% water, resists rapid temperature swings for exactly this reason.
• Cohesive and adhesive properties
  • Cohesion is the attraction between water molecules due to hydrogen bonding. Water molecules "stick" to each other.
  • Adhesion is the attraction between water molecules and other surfaces, like the walls of a glass tube or a plant's xylem vessel.
  • Together, cohesion and adhesion enable capillary action, which is how water travels upward through narrow tubes against gravity.
• High surface tension
  • Surface tension results from cohesive forces pulling water molecules at the surface inward, creating a "film" effect.
  • This allows insects like water striders to walk across the surface without breaking through, and it's why water forms spherical droplets on a flat surface.
• Expands when frozen
  • Most substances become denser when they solidify, but water does the opposite. As water freezes, hydrogen bonds lock molecules into a crystalline structure that spaces them farther apart, making ice less dense than liquid water.
  • Ice floats as a result, forming an insulating layer on lakes and ponds that allows aquatic organisms like fish and frogs to survive winter underneath.

Water as a universal solvent

Water's polarity is what makes it such an effective solvent. Each water molecule has a partial negative charge near the oxygen atom and partial positive charges near the hydrogen atoms. These charged regions interact with and pull apart ionic and polar compounds.

• Hydrophilic substances dissolve readily in water.
  • Examples include ions ($Na^+$, $Cl^-$), sugars like glucose and fructose, and amino acids like glycine.
  • This is essential for cellular processes. Metabolic reactions like hydrolysis depend on substances being dissolved in the aqueous environment of the cell.
• Hydrophobic substances do not dissolve readily in water.
  • Examples include lipids (fats and oils) and nonpolar regions of certain proteins.
  • This property is critical for cell membrane structure. The phospholipid bilayer forms precisely because the hydrophobic fatty acid tails are repelled by water, while the hydrophilic phosphate heads face the aqueous environment. This arrangement creates selective permeability.

Cohesion and adhesion in organisms

These properties show up in several important biological processes:

• Capillary action in plants
  • Adhesion pulls water molecules along the walls of narrow xylem vessels, while cohesion keeps the column of water molecules connected to one another. Together, they move water and dissolved minerals upward from roots to leaves, even against gravity.
• Surface tension and insect locomotion
  • The cohesive forces at the water's surface are strong enough to support the weight of lightweight insects like water striders and fishing spiders, allowing them to walk across the surface without sinking.
• Transpiration and evaporative cooling
  • As water evaporates from stomata (pores on leaf surfaces), cohesion pulls the entire column of water upward through the plant. This process is called transpiration.
  • The evaporation itself also cools the plant, similar to how sweating cools your skin. Water's high heat of vaporization means that a lot of heat energy is carried away with each molecule that evaporates.

Solutions and solvents

• In any solution, the solvent is the substance doing the dissolving, and the solute is the substance being dissolved. In biological systems, water is almost always the solvent.
• Dissolved solutes in organisms include ions, sugars, amino acids, and gases like oxygen.
• Blood plasma is a water-based solution that carries nutrients, hormones, and waste products throughout the body. Cellular cytoplasm is also an aqueous solution where most biochemical reactions take place.

pH balance in biological systems

The pH scale measures the concentration of hydrogen ions ($H^+$) in a solution. It runs from 0 to 14.

• A pH of 7 is neutral (pure water).
• A pH below 7 is acidic (higher $H^+$ concentration). Examples: lemon juice (~pH 2), vinegar (~pH 3).
• A pH above 7 is basic/alkaline (lower $H^+$ concentration). Examples: baking soda (~pH 9), seawater (~pH 8).

Each step on the pH scale represents a tenfold change in $H^+$ concentration. A solution at pH 3 has ten times more $H^+$ than a solution at pH 4.

• Acids donate $H^+$ ions in solution, increasing $H^+$ concentration. Examples: hydrochloric acid ($HCl$) in your stomach, citric acid in citrus fruits.
• Bases accept $H^+$ ions in solution, decreasing $H^+$ concentration. Examples: sodium hydroxide ($NaOH$), ammonia ($NH_3$).

Buffers resist drastic changes in pH when acids or bases are added. They consist of a weak acid and its conjugate base, which work together to absorb excess $H^+$ or $OH^-$ ions.

• The bicarbonate buffer ($H_2CO_3 / HCO_3^-$) operates in blood.
• The phosphate buffer ($H_2PO_4^- / HPO_4^{2-}$) operates inside cells.

Acid-Base Balance and pH Regulation

pH regulation in living organisms

pH regulation matters because enzymes function optimally within narrow pH ranges, usually near neutral. Extreme pH changes can denature proteins, meaning they lose their three-dimensional shape and can no longer function. This would disrupt processes like DNA replication and metabolism.

Buffer systems in the human body

• Bicarbonate buffer system in blood
  • Maintains blood pH between 7.35 and 7.45, which is the narrow range needed for proper oxygen transport by hemoglobin.
  • Consists of carbonic acid ($H_2CO_3$) and bicarbonate ion ($HCO_3^-$) in equilibrium. If blood becomes too acidic, $HCO_3^-$ absorbs excess $H^+$. If blood becomes too basic, $H_2CO_3$ releases $H^+$.
• Phosphate buffer system in cells
  • Maintains intracellular pH near 7.2 for optimal enzyme function.
  • Consists of dihydrogen phosphate ($H_2PO_4^-$) and hydrogen phosphate ($HPO_4^{2-}$), working by the same principle as the bicarbonate system.

Respiratory and renal regulation of pH

Beyond chemical buffers, the body uses two organ systems to fine-tune blood pH:

Respiratory system (fast response, within minutes):

1. If blood pH drops too low (too acidic), breathing rate increases. This expels more $CO_2$, which shifts the equilibrium and raises pH.
2. If blood pH rises too high (too basic), breathing rate decreases. $CO_2$ accumulates, forming more carbonic acid, which lowers pH.

Renal system (slower response, over hours to days):

1. When blood is too acidic, the kidneys secrete excess $H^+$ into urine and reabsorb $HCO_3^-$ back into the blood, raising pH.
2. When blood is too basic, the kidneys excrete $HCO_3^-$ in urine and retain $H^+$, lowering pH.