Biology Chapter: Equilibrium and Water Properties

Water has an unusually high boiling point, melting point, heat of vaporization, and surface tension compared to other substances with similar molecular weights. This is because of strong intermolecular forces of attraction between water molecules.

The maximum density of water occurs in its liquid state, not its solid state (ice). This means that ice floats on water.

A permanent dipole occurs when two atoms in a molecule have significantly different electronegativity. This means one atom attracts electrons more strongly than the other, leading to a partial negative charge on one atom and a partial positive charge on the other.

Hydrogen bonding arises from the attraction between the partially positive hydrogen atom of one water molecule and the partially negative oxygen atom of another water molecule.

The solvent properties of water are due to its polar nature. Water can dissolve many substances because its polar molecules can surround and pull apart charged or polar molecules.

A difference in hydrogen ion concentration across a membrane represents a form of stored energy. This difference is crucial for biological energy transformations.

Water molecules are constantly moving and interacting, and some of these interactions cause H2O to ionize (split) into hydrogen ions (H+) and hydroxide ions (OH-).

pH is a measure of the acidity or alkalinity of a solution. It is defined as the negative logarithm (base 10) of the hydrogen ion concentration (H+).

The process where nonpolar solutes interact with water molecules, leading to a net decrease in order and an increase in disorder.

Molecules that have both strongly polar and strongly nonpolar groups, forming structures like micelles in water.

Properties of solutions that are primarily affected by the concentration of solute particles, regardless of their chemical nature.

The temperature at which a liquid turns into a solid. Solutes decrease the freezing point of water by disrupting the formation of ice crystals.

The temperature at which a liquid turns into a gas. Solutes increase the boiling point of water by making it harder for water molecules to escape into the gas phase.

The tendency of a liquid to evaporate. Solutes decrease the vapor pressure of water by reducing the number of water molecules at the surface.

The pressure that must be applied to a solution to prevent the inward flow of water across a semipermeable membrane. Solutes increase the osmotic pressure of water by drawing more water into the solution.

Cells store substances in polymeric form to minimize osmotic pressure. For example, glycogen or starch exerts less osmotic pressure than individual glucose molecules.

The process where water molecules spontaneously ionize into hydrogen ions (H+) and hydroxide ions (OH-) in a reversible reaction. This creates a balance, with a tiny but essential concentration of both ions.

A state where the rates of forward and reverse reactions are equal, leading to no net change in concentrations of reactants and products.

States that when a stress is applied to a system at equilibrium, the system shifts in a direction that relieves the stress, restoring a new equilibrium.

The binding of oxygen to hemoglobin in red blood cells to be transported throughout the body. This is an equilibrium process, with the oxygen binding and releasing depending on oxygen levels.

A state of oxygen deficiency in the body, often occurring at high altitudes where oxygen levels are lower. Symptoms include headache, fatigue, and shortness of breath.

The adaptation of the body to low oxygen levels by increasing red blood cell production, thus increasing the amount of hemoglobin available to carry oxygen. This happens at high altitudes.

The change in concentrations of reactants and products in a chemical reaction that occurs when a stress is applied to the system at equilibrium.

The ability of a system to oppose changes to its equilibrium state. When a stress is applied, the system will try to minimize the impact.

Brønsted-Lowry acids donate a proton (H+) to form their conjugate base, while Brønsted-Lowry bases accept a proton to form their conjugate acid.

Strong acids ionize completely in water, meaning they donate all their protons, and are represented by a single arrow in chemical equations.

Weak acids ionize partially in water, establishing an equilibrium between the acid and its conjugate base. This is represented by a double arrow in chemical equations.

A strong base completely ionizes in water, forming hydroxide ions (OH-), which can accept protons. Strong bases are typically hydroxides of group 1 and 2 metals.

The negative logarithm of the hydronium ion concentration ([H3O+]), used to measure the acidity of a solution.

The negative logarithm of the hydroxide ion concentration ([OH-]), used to measure the basicity of a solution.

The relationship between pH and pOH in aqueous solutions, where the sum is always 14.

The molarity of hydrogen ions ([H+]) can be calculated from the pH using the formula: [H+] = 10^(-pH)

A buffer system resists changes in pH by neutralizing added acids or bases.

The H2CO3/HCO3- buffer system, involving carbonic acid and bicarbonate, is a crucial buffer system in body fluids. It helps maintain the pH of blood within a narrow range, crucial for optimal bodily functions.

A condition where the blood pH drops below the normal range, often caused by an increase in CO2 levels.

A condition where the blood pH rises above the normal range, often caused by a decrease in CO2 levels.

The lungs play a critical role in regulating blood pH by eliminating CO2. This process is fast and efficient.

The kidneys help maintain a 20:1 ratio of bicarbonate to carbonic acid in the blood. They can also excrete excess acid or base to regulate blood pH, but this process is slower than the lungs' CO2 elimination.

The bicarbonate buffer relies on both the respiratory and urinary systems to function effectively. The respiratory system releases CO2 to lower pH while the kidneys excrete bicarbonate (HCO3-) to lower pH or excrete H+ and release CO2 to raise pH. This interconnected regulation ensures optimal blood pH levels.

The carbonic acid buffer system is limited by its reliance on the properly functioning respiratory system and the availability of bicarbonate ions. Additionally, its buffering power is weak due to its pKa being different from the pH of blood.

This buffer system is effective within cells (ICF) and kidney tubules due to higher phosphate concentrations, optimal pH around 6.8 and the constant production of acidic byproducts of metabolism.

Proteins act as buffers by accepting or releasing hydrogen ions (H+) through their amino acid components. Think of it as proteins acting like sponges soaking up excess acid or base.

Hemoglobin is an important protein buffer in the blood. It binds to carbon dioxide, transports oxygen and hydrogen ions, and helps maintain blood pH.

The lungs are responsible for the elimination of carbon dioxide (CO2), which is a major acid contributor. Increased breathing removes CO2, reducing acidity. It's a fast-acting buffer.

The kidneys are the most powerful pH regulator in the body. They can remove fixed acids, conserve and produce bicarbonate ions (HCO3-) and excrete bases from the blood.

Rapid breathing (hyperventilation) leads to the removal of CO2, reducing acid levels in the blood and increasing pH .

Slow breathing (hypoventilation) leads to the accumulation of CO2, increasing acid levels and lowering pH.

The kidneys compensate for the acid-base balance. They can excrete excess acids, reabsorb bicarbonate, and generate new bicarbonate to neutralize the acids.