Water

1. Water • Water is a polar molecule composed of two polar covalent O-H bonds in a bent or angular molecular geometry with two pairs of nonbonding electrons. • water has 4 pairs of electrons arranged tetrahedral around the central oxygen atom • 75% of the earth’s surface is covered with water; • about 97% of the total water available on Earth is salt water, about 2% is frozen as the polar ice caps, and the rest (1%) is fresh water. • The amount of water on this planet is fairly constant and cleans and replenishes itself via “The Hydrologic cycle”. • Water vapor in the atmosphere (clouds & such) returns to the earth via precipitation (rain, snow, etc.) where it flows into land pockets (oceans, lakes, or rivers), or it is absorbed into the ground, or it evaporates back up into the atmosphere (completing the cycle).

2. Water • Rainwater collects dust particles and gases as it travels from the atmosphere to the ground. • Gases like O2, N2, and CO2 all dissolve to some degree in rainwater. An equilibrium is established between dissolved CO2 in water with carbonic acid making rainwater (about pH 5) more acidic than pure water (pH 7). CO2 + H2O  H2CO3 • As water flows beneath or atop the surface of the planet, it readily dissolves many substances from the soil and rocks. • Some common dissolved substances are Na+, K+, Ca2+, Mg2+, Fe2+, Cl-, SO42-, and HCO3-. Ca2+, Mg2+, and Fe2+ salts are responsible for “Hard water” (these positive ions react with the negative ions in soap to form insoluble scum). Soft water contains soluble ions like sodium and potassium.

3. Water • There are different types of “Polluted” water: • Pathogenic (disease-causing) microorganisms like cholera, typhoid, hepatitis, and dysentery still effect over 70% of the world’s population. • Aerobic biodegradation (aerobic oxidation) happens when microorganisms break down organic material in the present of dissolved oxygen to produce, for example, CO2, PO43-, NO3- SO42-, and HCO3-. The measure of oxygen needed to degrade organic material is referred to as the BOD (biochemical oxygen demand). • Anaerobic decay happens when the oxygen is depleted. The microorganisms reduce organic material (instead of oxidizing it) to produce nasty smelling substances like CH4, NH3, H2S, and amines. No life (excepts anaerobic microorganisms) can exist in such water. • Certain bacteria in water breaks down organic matter and, in the process, depletes the dissolved oxygen (which marine life is dependent on) while enriching the amount of plant nutrients (PO43-, NO3-) present. These nutrients promote algae growth. If the concentration of plant nutrients (from natural and human contributions) is left unchecked, it can lead to an excess of algae which, as it dies, increases the BOD eventually leading to anaerobic biodegradation. This process called eutrophication. • Industrial waste like VOC’s (volatile organic compounds like trichloroethylene), heavy metal ions/compounds (like Hg, Pb, & Cd), and a number of organic and inorganic materials from LUST (leaking underground storage tanks). • Acid rain produced from dissolved SOx and NOx compounds from air pollution and acid mine drainage from mining operations.

4. Some Physical Properties of Water • Water is colorless, odorless, and tasteless. • The normal boiling point is 100oC and the normal melting point is 0oC. • The heat of vaporization (DHvap) is 2259 J/g or 540 cal/g and the heat of fusion (DHfus) is 335 J/g or 80 cal/g. • The vapor pressure of water at 20oC is 17.5 torr; this is relatively low when compared to volatile ethyl alcohol (43.9 torr) and very volatile ethyl ether (442.2 torr) • The density of water at 4.0oC is 1.0 g/mL; the density of ice at 0oC is 0.917 g/mL. • The specific heat of water is 1.0 cal/g oC or 4.184 J/g oC.

5. The Unusual Properties of Water • Water co-exists in all three states of matter naturally on earth. • The only common substance is a liquid at STP. • As a solid, it is less dense than its liquid form, that is “Ice floats”. Most substances contract upon solidifying. • It has a very high Heat Capacity. It stores a large amount of energy with very little atomic or molecular motion. • It requires a lot of heat energy (enthalpy) to change states. • It has a high boiling point for such a low molecular weight compound. • It is a universal solvent, as a good dissolving medium a large number of substances are soluble in water.

6. Water as a universal solvent • Water is called the universal solvent because of its ability to dissolve many substances. The general solubility rule is “like dissolves like”. Since water is a polar molecule it will dissolve other polar substances as well as ionic compounds. Water will not dissolve or mix with nonpolar substances therefore water is immiscible in nonpolar substances. • Description of how water dissolves an ionic salt (like NaCl) on the molecular level? Although the attractive force from the partial charge of a single polar molecule is not as strong as the charge from an ion, it is plausible that a multitude of polar molecules could react on a single ion effectively. The positive end (H+) of several water molecules are attracted to the negative end of the salt crystal (Cl-) while the negative end of several water molecules (O2-) are attracted to the positive end of the crystal (Na+). The ionic bonds of the crystal are weakened by the solvating effect of the water molecules and the ions break away from the bulk crystal. The large number of water molecules in the container prevent the salt ions from re-combining.

7. Why is Water so unusual? The fundamental explanation for water’s unusual properties relates to the polarity of its bonds. Polarity describes the partial charge associated with a bond or molecule. A polar bond or molecule has a charge distribution present (one end positively charged and the other end negatively charged) while a nonpolar bond or molecule has no distinct charge distribution (neutral). Water is composed of two polar covalent O-H bonds (the difference in electronegativity is 1.4) arranged in a “bent” molecular geometry. Each bond has a dipole moment pointing in an overall similar direction leading to the existence of an overall dipole moment. The oxygen atom pulls the pair of electrons closer towards itself (making it partially negative) and further from the hydrogen atoms (making them partially positive). - + This charge distribution allows the partially positive hydrogen atoms from one molecule to be attracted to the partially negative oxygen atom of another molecule. This strong interlocking network between neighboring molecules is called HYDROGEN BONDING.The ability to form strong hydrogen bonds is the main reason for water’s unusual properties.

8. PROPERTIES ASSOCIATED WITH WATER HYDRATES: Solids that contain water molecules as part of their crystalline structure. The water in the hydrate is known as the water of hydration or the water of crystallization. HYGROSCOPIC: A substance is hygroscopic if it readily absorbs water from the atmosphere and forms a hydrate. DELIQUESCENT: A substance is deliquescent if it absorbs water from the air until it forms a solution. DESICCANTS: Compounds that absorb water and are used as drying agents. EFFLORESCENCE: The process by which crystalline materials spontaneously lose water when exposed to air.

9. Water and the Changes of State The energy required to heat (or cool) a solid (or heat/cool a liquid or a gas) can be calculated using q = msDT. It requires additional energy to change states. The energy required to convert a specific amount of the solid to a liquid is known as the heat of fusion (q = DHfus) and the energy required to convert a specific amount of a liquid to a gas is the heat of vaporization (q = DHvap). The total amount of energy can be calculated from qT = q1 + q2 + q3... Heating curve for water Temperature oC

10. Water and the Changes of State Q. How many kilojoules of energy are needed to change 15.0 g of ice at -5.00oC to steam at 125.0 oC? The first step is to design a pathway: q1 = msDT for ice from -5.0 to 0.0 oC, the specific heat of ice is 4.213 J/g oC q2 = DHfus for ice to liquid at 0.0oC q3 = msDT for liquid 0.0oC to 100.0 oC q4 = DHvap for liquid to steam at 100.0oC q5 = msDT for steam 100.0 to 125.0 oC; the specific heat of steam is 1.900 J/g oC so qT = q1 + q2 + q3 + q4 + q5 The next step is to calculate each q: q1= (15.0 g) (4.213 J/g oC) (0.0 - (-5.0) oC) = 316 J q2 = (335 J / g) (15.0 g) = 5025 J q3= (15.0 g) (4.184 J/g oC) (100.0 - (0.0) oC) = 6276 J q4 = (2260 J / g) (15.0 g) = 33900 J q5= (15.0 g) (1.900 J/g oC) (110 - 100 oC) = 285 J qT = 316 J + 5025 J + 6276 J + 33900 J + 285 J = 45.8 kJ

11. PRACTICE PROBLEMS #21a Ice at 0oC contains less heat than liquid water at the same temperature. Heat must be added to convert ice to water, so the water will contain that much more additional heat energy. Also the liquid state is in motion much more than the solid state. An increase in motion can only be accomplished by an increase in energy. 1. Which contains less heat, ice at 0oC or water at 0oC? Explain your answer. 2. On the basis of KMT, explain why vapor pressure increases with temperature. 3. Write equations to show how the following metals react with water. a) aluminum b) calcium c) potassium d) iron According to the kinetic molecular theory, the vapor pressure of a liquid should increase with temperature because of the increase in collisions and kinetic energy that always accompanies an increase in heat energy (temperature). KEm = 3/2 RT. The increase in energy thus motion allows the liquid molecules to escape (overcome the surface tension and other cohesive forces maintaining the liquid state) from the surface of the liquid into the gas phase. a) Al (s) + 3H2O (g)  3H2 (g) + Al2O3 *requires steam b) Ca (s) + 2H2O  H2 (g) + Ca(OH)2 *slowly at ambient temperature c) 2K (s) + 2H2O  H2 (g) + 2KOH + heat * vigorous at ambient temperature d) 3Fe (s) + 4H2O (g)  4H2 (g) + Fe3O4 *requires steam

12. PRACTICE PROBLEMS #21b See next slide for essay/answer 1. Explain the physical process of boiling. 2. Why does ice float in water? 3. Why does water have a relatively high boiling point? 4. Explain if ice will float in ethyl alcohol (d = 0.789 g/L)? 5. How much energy is needed to change 62.74 g of water at 15.00oC to steam at 103.0 oC? 6. Magnesium carbonate, MgCO3, forms a hydrate containing 39.1 % water of hydration. Calculate the formula of this hydrate. 1.645 x 105 J or 3.931 x 104 cal MgCO3 . 3 H2O

13. 1. Explain the physical process of boiling. At room temperature the water molecules have enough energy to allow the particles to move past each other but not enough to escape the surface tension. As the temperature of water increases, the heat energy (from the burner) is transferred to kinetic energy (for the molecules) leading to an increase in the molecular motion of the molecules. This action results in an increase in the vapor pressure above the surface of the liquid. When the vapor pressure of the water equals the external pressure, boiling begins. Now a sufficient amount of the molecules have enough energy to resist the attractive forces. Bubbles of vapor are formed throughout the liquid and these bubbles rise to the surface to escape.

14. 2. Why does ice float in water? Ice floats in its own liquid due to the intermolecular force, hydrogen bonding. As water freezes, the molecular motion of the molecules slow down and the partial positive end (hydrogen) of one water molecule is attracted to the partial negative end (oxygen) of another water molecule. Combine this event with the bent shape of water and the molecules become arranged in a 3-D hexagonal array. This array creates pockets of vacuum (empty space) in the lattice structure as well as a decrease in the number of molecules per unit volume. The mass is directly related to the number of molecules therefore, in the solid state, since there are less particles then there must be less mass per unit volume therefore the solid is less dense than the liquid.

15. 3. Why does water have a relatively high boiling point? Water has a relatively high boiling point because of the amount of intermolecular forces present. Water experiences LDF (London Dispersion Forces) and d-d (dipole-dipole) forces, along with the additional attractive force, Hydrogen bonding. A large amount of heat energy is required to break all of these forces in order for a phase transition to occur, thus the high boiling point. 4. Explain if ice will float in ethyl alcohol (d = 0.789 g/L)? Ice would not float in pure ethyl alcohol because the density of water is 1.000 g/mL which is greater than 0.789 g/mL for ethyl alcohol. Yet since ethyl alcohol also undergoes a small degree of hydrogen bonding, the sinking effect is not as dramatic as it would be with a nonpolar substance.

16. GROUP STUDY PROBLEMS #21 Short Essay 1. Can ice be colder than 0.0oC? Justify your answer. 2. Why does a boiling liquid maintain a constant temperature when heat is continually being added? 3. Why does a lake freeze from the top down? Math 1. Suppose 50.0 g of ice at 0.0oC are added to 285g of water at 22.0oC. Is there sufficient ice to lower the temperature of the system to 0.0oc and still have ice remaining? Show all work. 2. A mixture of 70.0 mL of hydrogen and 50.0 mL of oxygen is ignited to form water. Does any gas remain unreacted? 3. A 25.0 g sample of a hydrate of FePO4 was heated until all the water was driven off. The mass of anhydrous sample is 16.9 g. What is the formula of the hydrate?