Jim’s Discussion Notes

WATER AND THE FITNESS OF

THE ENVIRONMENT

Introduction

· Because water is the substance that makes possible life as we know it on Earth, astronomers hope to find evidence of water on newly discovered planets orbiting distant stars.

· Life on Earth began in water and evolved there for 3 billion years before spreading onto land.

· Even terrestrial organisms are tied to water.

· Most cells are surrounded by water and cells are about 70-95% water.

· Water exists in three possible states: ice, liquid, and vapor.

A. The Effects of Water’s Polarity

1. The polarity of water molecules results in hydrogen bonding

· In a water molecule two hydrogen atoms form single polar covalent bonds with an oxygen atom.

· Because oxygen is more electronegative, the region around oxygen has a partial negative charge.

· The region near the two hydrogen atoms has a partial positive charge.

· A water molecule is a polar molecule with opposite ends of the molecule with opposite charges.

· Water has a variety of unusual properties because of attractions between these polar molecules.

· The slightly negative regions of one molecule are attracted to the slightly positive regions of nearby molecules, forming a hydrogen bond.

· Each water molecule can form hydrogen bonds with up to four neighbors.

2. Organisms depend on the cohesion of water molecules

· The hydrogen bonds joining water molecules are weak, about 1/20th as strong as covalent bonds.

· They form, break, and reform with great frequency.

· At any instant, a substantial percentage of all water molecules are bonded to their neighbors, creating a high level of structure.

· Hydrogen bonds hold the substance together, a phenomenon called cohesion.

· Cohesion among water molecules plays a key role in the transport of water against gravity in plants.

· Water that evaporates from a leaf is replaced by water from vessels in the leaf.

· Hydrogen bonds cause water molecules leaving the veins to tug on molecules further down.

· This upward pull is transmitted to the roots.

· Adhesion, clinging of one substance to another, contributes too, as water adheres to the wall of the vessels.

· Surface tension, a measure of the force necessary to stretch or break the surface of a liquid, is related to cohesion.

· Water has a greater surface tension than most other liquids because hydrogen bonds among surface water molecules resist stretching or breaking the surface.

· Water behaves as if covered by an invisible film.

· Some animals can stand, walk, or run on water without breaking the surface.

3. Water moderates temperatures on Earth

· Water stabilizes air temperatures by absorbing heat from warmer air and releasing heat to cooler air.

· Water can absorb or release relatively large amounts of heat with only a slight change in its own temperature.

· Atoms and molecules have kinetic energy, the energy of motion, because they are always moving.

· The faster that a molecule moves, the more kinetic energy that it has.

· Heat is a measure of the total quantity of kinetic energy due to molecular motion in a body of matter.

· Temperature measures the intensity of heat due to the average kinetic energy of molecules.

· As the average speed of molecules increases, a thermometer will record an increase in temperature.

· Heat and temperature are related, but not identical.

· When two objects of different temperature meet, heat passes from the warmer to the cooler until the two are the same temperature.

· Molecules in the cooler object speed up at the expense of kinetic energy of the warmer object.

· Ice cubes cool a drink by absorbing heat as the ice melts.

· In most biological settings, temperature is measured on the Celsius scale (^oC).

· At sea level, water freezes at O ^oC and boils at 100^oC.

· Human body temperature averages 37 ^oC.

· While there are several ways to measure heat energy, one convenient unit is the calorie (cal).

· One calorie is the amount of heat energy necessary to raise the temperature of one g of water by 1^oC.

· In many biological processes, the kilocalorie (kcal), is more convenient.

· A kilocalorie is the amount of heat energy necessary to raise the temperature of 1000g of water by 1^oC.

· Another common energy unit, the joule (J), is equivalent to 0.239 cal.

· Water stabilizes temperature because it has a high specific heat.

· The specific heat of a substance is the amount of heat that must be absorbed or lost for 1g of that substance to change its temperature by 1^oC.

· By definition, the specific heat of water is 1 cal per gram per degree Celcius or 1 cal/g/^oC.

· Water has a high specific heat compared to other substances.

· For example, ethyl alcohol has a specific heat of 0.6 cal/g/^oC.

· The specific heat of iron is 1/10th that of water.

· Water resists changes in temperature because it takes a lot of energy to speed up its molecules.

· Viewed from a different perspective, it absorbs or releases a relatively large quantity of heat for each degree of change.

· Water’s high specific heat is due to hydrogen bonding.

· Heat must be absorbed to break hydrogen bonds and is released when hydrogen bonds form.

· Investment of one calorie of heat causes relatively little change to the temperature of water because much of the energy is used to disrupt hydrogen bonds, not move molecules faster.

· The impact of water’s high specific heat ranges from the level of the whole environment of Earth to that of individual organisms.

· A large body of water can absorb a large amount of heat from the sun in daytime and during the summer, while warming only a few degrees.

· At night and during the winter, the warm water will warm cooler air.

· Therefore, ocean temperatures and coastal land areas have more stable temperatures than inland areas.

· The water that dominates the composition of biological organisms moderates changes in temperature better than if composed of a liquid with a lower specific heat.

· The transformation of a molecule from a liquid to a gas is called vaporization or evaporation.

· This occurs when the molecule moves fast enough that it can overcome the attraction of other molecules in the liquid.

· Even in a low temperature liquid (low average kinetic energy), some molecules are moving fast enough to evaporate.

· Heating a liquid increases the average kinetic energy and increases the rate of evaporation.

· Heat of vaporization is the quantity of heat that a liquid must absorb for 1 g of it to be converted from the liquid to the gaseous state.

· Water has a relatively high heat of vaporization, requiring about 580 cal of heat to evaporate 1g of water at room temperature.

· This is double the heat required to vaporize the same quantity of alcohol or ammonia.

· This is because hydrogen bonds must be broken before a water molecule can evaporate from the liquid.

· Water’s high heat of vaporization moderates climate by absorbing heat in the tropics via evaporation and releasing it at higher latitudes as rain.

· As a liquid evaporates, the surface of the liquid that remains behind cools - evaporative cooling.

· This occurs because the most energetic molecules are the most likely to evaporate, leaving the lower kinetic energy molecules behind.

· Evaporative cooling moderates temperature in lakes and ponds and prevents terrestrial organisms from overheating.

· Evaporation of water from the leaves of plants or the skin of humans removes excess heat.

4. Oceans and lakes don’t freeze solid because ice floats

· Water is unusual because it is less dense as a solid than as a liquid.

· Most materials contract as they solidify, but water expands.

· At temperatures above 4^oC, water behaves like other liquids, expanding when it warms and contracting when it cools.

· Water begins to freeze when its molecules are no longer moving vigorously enough to break their hydrogen bonds.

· When water reaches 0^oC, water becomes locked into a crystalline lattice with each molecule bonded to the maximum of four partners.

· As ice starts to melt, some of the hydrogen bonds break and some water molecules can slip closer together than they can while in the ice state.

· Ice is about 10% less dense than water at 4^oC.

· Therefore, ice floats on the cool water below.

· This oddity has important consequences for life.

· If ice sank, eventually all ponds, lakes, and even the ocean would freeze solid.

· During the summer, only the upper few inches of the ocean would thaw.

· Instead, the surface layer of ice insulates liquid water below, preventing it from freezing and allowing life to exist under the frozen surface.

5. Water is the solvent of life

· A liquid that is a completely homogeneous mixture of two or more substances is called a solution.

· A sugar cube in a glass of water will eventually dissolve to form a uniform mixture of sugar and water.

· The dissolving agent is the solvent and the substance that is dissolved is the solute.

· In our example, water is the solvent and sugar the solute.

· In an aqueous solution, water is the solvent.

· Water is not a universal solvent, but it is very versatile because of the polarity of water molecules.

· Water is an effective solvent because it so readily forms hydrogen bonds with charged and polar covalent molecules.

· For example, when a crystal of salt (NaCl) is placed in water, the Na⁺ cations form hydrogen bonds with partial negative oxygen regions of water molecules.

· The Cl^- anions form hydrogen bonds with the partial positive hydrogen regions of water molecules.

· Each dissolved ion is surrounded by a sphere of water molecules, a hydration shell.

· Eventually, water dissolves all the ions, resulting in a solution with two solutes, sodium and chloride.

· Polar molecules are also soluble in water because they can also form hydrogen bonds with water.

· Even large molecules, like proteins, can dissolve in water if they have ionic and polar regions.

· Any substance that has an affinity for water is hydrophilic.

· These substances are dominated by ionic or polar bonds.

· This term includes substances that do not dissolve because their molecules are too large and too tightly held together.

· For example, cotton is hydrophilic because it has numerous polar covalent bonds in cellulose, its major constituent.

· Water molecules form hydrogen bonds in these areas.

· Substances that have no affinity for water are hydrophobic.

· These substances are dominated by non-ionic and nonpolar covalent bonds.

· Because there are no consistent regions with partial or full charges, water molecules cannot form hydrogen bonds with these molecules.

· Oils, such as vegetable oil, are hydrophobic because the dominant bonds, carbon-carbon and carbon-hydrogen, exhibit equal or near equal sharing of electrons.

· Hydrophobic molecules are major ingredients of cell membranes.

· Biological chemistry is “wet” chemistry with most reactions involving solutes dissolved in water.

· Chemical reactions depend on collisions of molecules and therefore on the number of molecules available.

· Counting individual or even collections of molecules is not practical.

· Instead, we can use the concept of a mole to convert weight of a substance to the number of molecules in that substance and vice versa.

· A mole (mol) is equal in number to the molecular weight of a substance, but upscaled from daltons to units of grams.

· To illustrate, how could we measure out a mole of table sugar - sucrose (C₁₂H₂₂O₁₁)?

· A carbon atom weighs 12 daltons, hydrogen 1 dalton, and oxygen 16 daltons.

· One molecule of sucrose would weigh 342 daltons, the sum of weights of all the atoms in sucrose or the molecular weight of sucrose.

· To get one mole of sucrose we would weigh out 342g.

· The advantage of using moles as a measurement is that a mole of one substance has the same number of molecules as a mole of any other substance.

· If substance A has a molecular weight of 10 daltons and substance B has a molecular weight of 100 daltons, then we know that 10g of A has the same number of molecules as 100g of substance B.

· The actual number of molecules in a mole is called Avogadro’s number, 6.02 x 10²³.

· A mole of sucrose contains 6.02 x 10²³ molecules and weighs 342g, while a mole of ethyl alcohol (C₂H₆O) also contains 6.02 x 10²³ molecules but weighs only 46g because the molecules are smaller.

· In “wet” chemistry, we are typically combining solutions or measuring the quantities of materials in aqueous solutions.

· The concentration of a material in solution is called its molarity.

· A one molar solution has one mole of a substance dissolved in one liter of solvent, typically water.

· To make a 1 molar (1M) solution of sucrose we would slowly add water to 342g of sucrose until the total volume was 1 liter and all the sugar was dissolved.

B. The Dissociation of Water Molecules

· Occasionally, a hydrogen atom shared by two water molecules shifts from one molecule to the other.

· The hydrogen atom leaves its electron behind and is transferred as a single proton - a hydrogen ion (H⁺).

· The water molecule that lost a proton is now a hydroxide ion (OH^-).

· The water molecule with the extra proton is a hydronium ion (H₃O⁺).

· A simpler way to view this process is that a water molecule dissociates into a hydrogen ion and a hydroxide ion:

· H₂O <=> H⁺ + OH^-

· This reaction is reversible.

· At equilibrium the concentration of water molecules greatly exceeds that of H⁺ and OH^-.

· In pure water only one water molecule in every 554 million is dissociated.

· At equilibrium the concentration of H⁺ or OH^- is 10^-7M (25°C).

· Because hydrogen and hydroxide ions are very reactive, changes in their concentrations can drastically affect the proteins and other molecules of a cell.

· Adding certain solutes, called acids and bases, disrupts the equilibrium and modifies the concentrations of hydrogen and hydroxide ions.

· The pH scale is used to describe how acidic or basic (the opposite of acidic) a solution is.

1. Organisms are sensitive to changes in pH

· An acid is a substance that increases the hydrogen ion concentration in a solution.

· When hydrochloric acid is added to water, hydrogen ions dissociate from chloride ions:

· HCl -> H⁺ + Cl^-

· Addition of an acid makes a solution more acidic.

· Any substance that reduces the hydrogen ion concentration in a solution is a base.

· Some bases reduce H⁺ directly by accepting hydrogen ions.

· Ammonia (NH₃) acts as a base when the nitrogen’s unshared electron pair attracts a hydrogen ion from the solution, creating an ammonium in (NH₄⁺).

· NH₃ + H⁺ <=> NH₄⁺

· Other bases reduce H+ indirectly by dissociating to OH- that combines with H⁺ to form water.

· NaOH -> Na⁺ + OH^- OH^- + H⁺ -> H₂O

· Solutions with more OH^- than H⁺ are basic solutions.

· Some acids and bases (HCl and NaOH) are strong acids or bases.

· These molecules dissociate completely in water.

· Other acids and bases (NH₃) are weak acids or bases.

· For these molecules, the binding and release of hydrogen ions are reversible.

· At equilibrium there will be a fixed ratio of products to reactants.

· Carbonic acid (H₂CO₃) is a weak acid:

· H₂CO₃ <=> HCO₃^- + H⁺

· At equilibrium, 1% of the molecules will be dissociated.

· In any solution the product of their H⁺ and OH^- concentrations is constant at 10^-14.

· [H⁺] [OH^-] = 10^-14

· In a neutral solution, [H⁺] = 10^-7 M and [OH^-] = 10^-7 M

· Adding acid to a solution shifts the balance between H⁺ and OH^- toward H+ and leads to a decline in OH^-.

· If [H⁺] = 10^-5 M, then [OH^-] = 10^-9 M

· Hydroxide concentrations decline because some of the additional acid combines with hydroxide to form water.

· Adding a base does the opposite, increasing OH^- concentration and dropping H⁺ concentration.

· The H⁺ and OH^- concentrations of solutions can vary by a factor of 100 trillion or more.

· To express this variation more conveniently, the H⁺ and OH^- concentrations are typically expressed via the pH scale.

· The pH scale, ranging from 1 to 14, compresses the range of concentrations by employing logarithms.

• pH = - log [H⁺] or [H⁺] = 10^-pH

· In a neutral solution [H⁺] = 10^-7 M, and the pH = 7.

· Values for pH decline as [H⁺] increase.

· While the pH scale is based on [H⁺], values for [OH^-] can be easily calculated from the product relationship.

· The pH of a neutral solution is 7.

· Acidic solutions have pH values less than 7 and basic solutions have pH values more than 7.

· Most biological fluids have pH values in the range of 6 to 8.

· However, pH values in the human stomach can reach 2.

· Each pH unit represents a tenfold difference in H⁺ and OH^- concentrations.

· A small change in pH actually indicates a substantial change in H⁺ and OH^- concentrations.

· The chemical processes in the cell can be disrupted by changes to the H⁺ and OH^- concentrations away from their normal values near pH 7.

· To maintain cellular pH values at a constant level, biological fluids have buffers.

· Buffers resist changes to the pH of a solution when H⁺ or OH^- is added to the solution.

· Buffers accept hydrogen ions from the solution when they are in excess and donate hydrogen ions when they have been depleted.

· Buffers typically consist of a weak acid and its corresponding base.

· One important buffer in human blood and other biological solutions is carbonic acid.

· The chemical equilibrium between carbonic acid and bicarbonate acts at a pH regulator.

· The equilibrium shifts left or right as other metabolic processes add or remove H⁺ from the solution.

2. Acid precipitation threatens the fitness of the environment

· Acid precipitation is a serious assault on water quality and therefore the environment for all life where this problem occurs.

· Uncontaminated rain has a slightly acidic pH of 5.6.

· The acid is a product of the formation of carbonic acid from carbon dioxide and water.

· Acid precipitation occurs when rain, snow, or fog has a pH that is more acidic than 5.6.

· Acid precipitation is caused primarily by sulfur oxides and nitrogen oxides in the atmosphere.

· These molecules react with water to form strong acids.

· These fall to the surface with rain or snow.

· The major source of these oxides is the burning of fossil fuels (coal, oil, and gas) in factories and automobiles.

· The presence of tall smokestacks allows this pollution to spread from its site of origin to contaminate relatively pristine areas.

· Rain in the Adirondack Mountains of upstate New York averages a pH of 4.2

· The effects of acids in lakes and streams is more pronounced in the spring during snowmelt.

· As the surface snows melt and drain down through the snow field, the meltwater accumulates acid and brings it into lakes and streams all at once.

· The pH of early meltwater may be as low as 3.

· Acid precipitation has a great impact on the eggs and the early developmental stages of aquatic organisms that are abundant in the spring.

· Thus, strong acidity can alter the structure of molecules and impact ecological communities.

· Direct impacts of acid precipitation on forests and terrestrial life are more controversial.

· However, acid precipitation can impact soils by affecting the solubility of soil minerals.

· Acid precipitation can wash away key soil buffers and plant nutrients (calcium and magnesium).

· It can also increase the solubility of compounds like aluminum to toxic levels.

· This has done major damage to forests in Europe and substantial damage of forests in North America.