Water

1. Water From Biology, Campbell 5th edition

2. Water’s polarity • A water molecule is shaped like a V, with two hydrogen atoms covalently bonded with an oxygen atom. • The oxygen atom is more electronegative than the hydrogen. • The electrons spend more time around the oxygen resulting in a polar molecule with a positive and a negative end.

3. O H H

4. Water molecules are electrically attracted to one another with the positive hydrogen ends of one molecule being attracted to the negative oxygen end of another water molecule. This attraction results in a hydrogen bond. Each water molecule can form hydrogen bonds to a maximum of four neighbors. Water’s polarity

6. Cohesion • Hydrogen bonds between water molecules are much weaker than the covalent bonds within the individual molecule. • Hydrogen bonds of liquid water will break and reform with great frequency. • Collectively the hydrogen bonds hold the substance together through cohesion. • When plants loose water through evaporation on their leaves it is replaced by water that is pulled up with hydrogen bonds tugging molecules upward.

7. Adhesion of water • Adhesion is the clinging of one substance to another. • Adhesion allows water in a plant counter the downward pull of gravity so it can get water to the top of the plant.

8. Surface tension of water • Surface tension is a measure of how difficult it is to stretch or break the surface of a liquid. • Water has a greater surface tension than most liquids.

9. In the space below, draw a picture showing 3 water molecules. Indicate H and O atoms, their partial charges, and show where Hydrogen bonds would occur. In the space below, list the three important properties of water we have learned so far….

10. Heat and temperature • 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 the molecules. • When the average speed of molecules increases, a thermometer records this as a rise in temperature. • Comparing a swimmer to the ocean, the swimmer has a higher temperature but the ocean has more heat due to its volume.

11. Measurements of heat and temperature • Temperature is measured in Celsius. • Heat is measured in calories which are the amount of heat energy it takes to raise the temperature of 1 g of water by 1 degree Celsius. • A kilocalorie is 1000 calories, it is what you read on food packages.

12. Specific heat • Specific heat is the amount of heat that must be absorbed or lost for 1 g of that substance to change its temperature by 1 degree Celsius. • Water has a specific heat of one calorie per gram per degree Celsius, which is much higher than most other substances.

13. Relevance of the specific heat of water • A large body of water can absorb and store a huge amount of heat from the sun in the daytime and during summer, while warming up only a few degrees. • At night and during the winter, the gradual cooling of water can warm the air. • Acts as a heat bank! • The high specific heat of water also tends to stabilize ocean temperatures creating a favorable environment for marine life.

14. Relevance of the specific heat of water • Because organisms are made primarily of water they are more able to resist changes in their own temperatures. • Water, in short, keeps temperature fluctuations on land and in water within limits that permit life.

15. Vaporization/Evaporation • When molecules of a liquid move fast enough they overcome their attractions to one another and enter the air as a gas. • The transformation from liquid to a gas is called vaporization, or evaporation. • Heat of vaporization is the quantity of heat a liquid must absorb for 1 g of it to be converted from the liquid to gas.

16. Water and heat of vaporization • Water has a high heat of vaporization due to hydrogen bonding. • The high heat of vaporization helps moderate the Earth’s climate because water absorbs solar heat, surface water evaporates, then the air circulates poleward and releases heat as it condenses to rain.

17. Evaporative cooling • As liquid evaporates, the surface of the liquid that remains behind cools down. • Evaporative cooling occurs because the hottest molecules leave as a gas. • Evaporative cooling of water again contributes to the stability of temperatures in lake and ponds, and it also prevents terrestrial organisms from overheating.

18. Aren’t we glad that ice floats!?! • Water is less dense in solid form, unlike most substances. • This phenomena is once again due to hydrogen bonding (love that hydrogen bonding!) • Water when frozen locks into a crystalline lattice with a molecule bonded to four others. This keeps the molecules apart from one another, hence it is less dense. • The ability of ice to float keeps bodies of water from freezing solid. Without it, life would be impossible on Earth.

19. Remember solutions? • Solutions are homogeneous mixtures of two or more substances. • The solvent is the dissolving agent. • The solute is what is dissolved. • An aqueous solution is one in which water is the solvent. • Water is an incredible solvent due to its polarity.

21. Hydrophilic and hydrophobic substances • Any substance that has an affinity for water is hydrophilic. (ex. Polar molecules, ionic molecules, cellulose) • Substance that do not have an affinity for water are hydrophobic. (non-ionic and nonpolar) • Thank goodness our cell membranes are hydrophobic or they might dissolve!

22. List 4 more important properties of water that we just learned… • 1. • 2. • 3. • 4.

23. Calculating the concentrations of solutes dissolve in aqueous solutions • To obtain a 1 mol of a substance, weigh out its molecular weight in grams. (Sucrose has a molecular weight of 342). • To make a 1 molar solution of sucrose, you would gradually add water and stir until completely dissolved and the total solution equals one liter. • Molarity is the number of moles of solute per liter of a solution. • Remember, according to Avogadro’s number, a one molar substance has the same amount of solute molecules in water, but weighs differently because different substances have different molecular weights.

24. Dissociation of water molecules • Occasionally, a hydrogen atom shared between two water molecules in a hydrogen bonds shifts from one molecule to the other. • The hydrogen molecule leaves its electron behind and what is transferred is a hydrogen ion (a single proton with a charge of +1) • The water molecule that lost a proton is now a hydroxide ion (OH-), which has a charge of –1. • The water molecule that the proton binds to makes a hydronium ion (H3O+). • However, the process can be simplified as the dissociation of a water molecule into a hydrogen ion and a hydroxide ion.

25. In pure water, only one molecule in every 554 million is dissociated, and there is only one ten-millionth of a mole of hydrogen ions per liter( 10 –7 M), and an equal number of hydroxide ions. • Hydrogen and hydroxide ions are very reactive. • Changes in their concentrations can drastically affect a cell’s proteins and other complex molecules. • By adding acids and bases, the balance of the two ions is disrupted.

26. Acids and bases • An acid is a substance that increases the H+ concentration of a solution. (ex. hydrochloric acid) • A base reduces the H+ concentration. (ex. Sodium hydroxide, ammonia)

27. The pH scale is used to measure the strength of acids and bases. It ranges from 0-14, with pure water (7) as being neutral, and the lower the number the more acidic, and the higher the number the more basic. pH Scale

28. The equation used to determine pH is a logarithmic scale with the equation: • pH = -log (H+) (H+ = hydrogen ion concentration) • For a neutral solution (H+) is 10-7 M resulting in –log 10-7 = -(-7)= 7 • As H+ concentration increases, pH declines. • Each pH unit represents a tenfold difference in H+ and OH- concentrations. A solution of pH 3 is a thousand times more acidic than a solution of pH 6.

29. Buffers • Buffers are substances that minimize changes in the concentrations of H+ and OH- ions in a solution. • They do this by accepting excess hydrogen ions when in excess and donating them when they’re depleted. • Ex. Blood and carbonic acid/bicarbonate ion.

30. Acid rain • Acid precipitation is rain, snow, or fog more acidic than pH 5.6. • It is caused by the presence in the atmosphere of sulfur oxides and nitrogen oxides. • These compounds react with water in the air to form strong acids which fall to earth with rain or snow. • Sources of these oxides are the burning of fossil fuels.