CHAPTER 5 Water and Seawater

1. CHAPTER 5 Water and Seawater

2. Chemistry scares people, but it doesn’t have to! • Understanding chemistry is very important for understanding aspects of oceanography and the living organisms that inhabit the oceans

3. Basic chemistry • Atomic structure • Nucleus = protons (positive) + neutrons (neutral) • Electrons in orbitals around nucleus (negative) http://www.rstp.uwaterloo.ca/manual/matter/graphic/atom.jpg

4. Electrons (negative charge) • Found in shells around nucleus • 1st shell can hold 2 electrons; 2nd , 3rd, and all other shells can hold 8 electrons • Not all atoms have outer shells that are completely filled • Atoms bond with other atoms to fill outer shell

5. Chemical bonds • Attractive force that holds atoms together, interaction of those electrons • Three major types • Ionic bonds • Covalent bonds • Hydrogen bonds http://w3.dwm.ks.edu.tw/bio/activelearner/02

6. http://serc.carleton.edu/images/usingdata/nasaimages

7. Ionic bonds • Atoms “exchange” electrons to fill outer shell • becomes positive ion if lose electron • becomes negative ion if gain electron • + & – ions attracted to each other • NaCl  Na+ + Cl- http://www.physicalgeography.net/fundamentals/images http://www.msnucleus.org/membership/html/k-6/rc/minerals/3

8. Covalent bonds • Atoms “share” electrons to fill outer shell • For example: • H (hydrogen) has one electron, needs 1 more • O (oxygen) has 6 electrons in outer shell, needs two electrons • Therefore, oxygen and 2 hydrogens bond to form water • Covalent bonds are stronger because there is sharing of the electrons http://ghs.gresham.k12.or.us/science/ps/sci/ibbio/chem/notes/chpt2

9. Polar covalent bonds • Electrons not equally distributed in molecule • Water is a polar molecule • O strongly attracts electrons  slightly negative • H slightly positive • Think of oxygen as being the “bully” – it’s larger so it pulls the electrons towards it’s nucleus more often • Allows formation of H-bonding between water molecules http://www.mie.utoronto.ca/labs/lcdlab/biopic/fig

10. H2O molecule • Two hydrogen H and one oxygen O atoms bonded by sharing electrons • Polar covalent bond

11. Hydrogen bonding as a result of polar covalent bonds • Polarity • small negative charge at O end • small positive charge at H end • Attraction between + and – ends of water molecules to each other or other ions • Happens because of the polar covalent bond Fig. 5.3

12. Hydrogen bonds are weak bonds between +/- ends of poles causes water molecules to "stick" together – cohesion • Gives water important distinct properties • H2O molecule forms H-bonds w/ up to fourother water molecules, depending on temperature http://www.nyu.edu/pages/mathmol/modules/water

13. Hydrogen bonding and water • Hydrogen bonds are weaker than covalent bonds but still strong enough to result in unique properties of water • Cohesion = sticks to other water molecules • Adhesion = sticks to other types of molecules • High surface tension http://faculty.uca.edu/~benw/biol1400 http://ucsu.colorado.edu/~meiercl/photography

14. Hydrogen bonding and water • H-bonds absorb red light, reflect blue light blue color • High solubility of chemical compounds in water • Solid, liquid, gas at Earth’s surface, not all substances can say the same • Unusual thermal properties • Unusual density http://www.pacific-promotion.com.fr/Phototek

15. Water exists naturally in all 3 states on Earth’s surface

16. Changes of state due to adding or subtracting heat • Heat is energy of moving molecules • calorie is amount of heat needed to raise the temperature of 1 gram of water by 1o C • Temperature is measurement of average kinetic energy

17. Unusual thermal properties of H2O • H2O has high boiling point • H2O has high freezing point • Most H2O is in liquid form of water on Earth’s surface • VERY important for life

18. Unusual thermal properties of H2O • High latent (hidden) heatsof • Vaporization • Melting/freezing • H-bonds holding water together require extra energy (heat) to break bonds •  change states without change in temperature (a to b, c to d in figure)

19. Fig. 5.6

20. Unusual thermal properties of H2O • Water high heat capacity (specific heat) • Amount of heat required to raise temperature of 1 gram of any substance 1o C • Water can take in/lose lots of heat without changing temperature – must break H-bonds • On the other hand, rocks have low heat capacity • Rocks quickly change temperature as they gain/lose heat

21. Global thermostatic effects • Moderates temperature on Earth’s surface – water temp less variable and less extreme than air temperatures • Equatorial oceans (hot) don’t boil • Polar oceans (cold) don’t freeze solid http://www.goredsea.com/media/images/EN

22. Global thermostatic effects • Marine effect • Oceans moderate temperature changes day/night; different seasons • Continental effect • Land areas have greater range of temperatures day/night and during different seasons • Look at the differences between coastal Florida compared to Orlando

24. Density of water • Density of water increases as temperature decreases down to 4oC • From 4oC to 0oC density of water decreases as temperature decreases • Density of ice is less than density of water • VERY unique property • Think about it – most solids are MORE dense then their liquids

25. Density of water Fig. 5.10

26. Density of water • Dissolved solids reduce freezing point of water • As water freezes, the crystalline structure “pushes out” much of the dissolved solids • Creates icy “slush” and surrounding waters become saltier • Putting salt on icy roads melts ice • Salt lowers freezing point of water on roads allowing it to remain liquid at colder temps

27. Water = Life • Summary: • Unique properties of water that make life possible • High heat capacity and specific heat • Moderates climates • Keeps equatorial regions from boiling and pole regions from freezing solid • High latent heat – when undergoing change of state, large amount of heat is absorbed or released • Sweat evaporating from your skin draws heat from your body, keep you cool • Ice is less dense than liquid water • Cohesion • Water moving up xylem in plants • Surface tension – allows plankton to stay near surface of water

28. Salinity • Six elements make up 99% of dissolved solids in seawater – from erosion of land, volcanism • Total amount of solid material dissolved in water- Traditional definition • Typical salinity is 3.5% or 35o/oo • o/oo or parts per thousand (ppt) = grams of salt per kilogram of water (g/Kg ) • Adding salts changes many properties of water Fig. 5.12

29. Measuring salinity • Evaporation • Chemical analysis - titration • Principle of constant proportions • Major dissolved constituents in same proportion regardless of total salinity • Measure amount of halogens (Cl, Br, I, F) (chlorinity) • Salinity = 1.80655 * Chlorinity (ppt) • Specific gravity (1.028 g/ml) • Hydrometer • Electrical conductivity • Salinometer http://iodeweb5.vliz.be/oceanteacher/resources/other/AndersonBook/images/salmeter.jpg

30. Pure water vs. seawater

31. Salinity variations • Open ocean salinity 33 to 38 o/oo • However, coastal areas salinity varies more widely • Influx of freshwater lowers salinity or creates brackish conditions • Greater rate of evaporation raises salinity or creates hypersalineconditions • Salinity may vary with seasons (dry/rain) http://farm1.static.flickr.com/58/186020843_205a03e35e.jpg?v=0

32. How to change salinity • Add/remove water • Add/remove dissolved substances

33. Processes that add/subtract salinity from oceans Salinity increases through: • Precipitation (rain or snow) • Runoff (river flow) • Melting icebergs • Melting sea ice • Evaporation • Formation of sea ice Salinity decreases through: Floating in the Dead Sea

34. Hydrologic cycle describes recycling of water near Earth’s surface Fig. 5.15

35. Processes that add/subtract dissolved substances • River flow • Volcanic eruptions • Atmosphere • Biologic interactions • Salt spray • Chemical reactions at seawater-sea floor interface • Biologic interactions • Evaporite formation • Adsorption • Physical attachment to sinking clay or biological particles Dissolved substances increases through: Decreases through:

36. Residence time • Average length of time a substance remains dissolved in seawater • Ions with long residence time are in high concentration in seawater (Na+, Cl-) • Ions with short residence time are in low concentration in seawater  percipitate out (K+, Ca2+ ) • Steady state condition

37. Residence time and steady state Fig. 5.16

38. pH – concentration of H+ ions • Acid releases H+ when dissolved in water (HCl, H2SO4) • Alkaline (or base) releases OH- (NaOH) • pH scale measures the hydrogen ion concentration, logarithmic scale 0-14 • Low pH value, acid • High pH value, alkaline (basic) • pH 7 = neutral

39. Figure 5.17

40. Carbonate buffering • Keeps ocean pH about same (8.1, slightly alkaline) • pH too high, carbonic acid releases H+ • pH too low, bicarbonate combines with H+ • Precipitation/dissolution of calcium carbonate CaCO3 buffers ocean pH (CaCO3 Ca+ + CO3-) • CO3- bonds with H ions created when CO2 interacts with H2O • Oceans can absorb CO2 from atmosphere without much change in pH

42. Carbonate buffering Too acidic removes H+ Too basic adds H+ Fig. 5.18

43. Surface ocean variation of salinity • Surface salinity varies primarily with latitude • Polar regions: salinity lower • lots of rain/snow and runoff • Low temps, not a lot of evaporation • Mid-latitudes: higher salinity • because of evaporation (dry areas) • Equator: salinity slightly lower than mid-latitudes • due to lots of rain despite high evaporation

44. Deep ocean variation of salinity • Surface ocean salinity is variable • Due to occurrences at surface – rain, evaporation, etc • Deeper ocean salinity is nearly the same (polar source regions for deeper ocean water) • Halocline, rapid change of salinity with depth

45. Density of seawater • 1.022 to 1.030 g/cm3 surface seawater • Saltwater more dense than pure water • That is why you can float better in saltwater • Ocean layered according to density • Density seawater controlled by temperature, salinity, and pressure • Most important influence is temperature • Density increases with decreasing temperature

46. Density of seawater • Overall, temp has greatest effect on density • However, salinity greatest influence on density in polar oceans • polar ocean is isothermal (same temperature as depth increases) • Currents from lower latitudes bring higher salinity water into polar areas • But polar waters are overall isothermal AND isopycnal http://www.waterencyclopedia.com/images/wsci_03_img0394.jpg

47. Density versus depth • Pycnocline, abrupt change of density with depth • Thermocline, abrupt change of temperature with depth • Density differences cause a layered ocean • Mixed surface water • Pycnocline and thermocline • Deep water

49. Desalination processes • Remove salt from seawater • Distillation – most common process, but energetically costly • Reverse osmosis – flimsy membranes

50. Misconceptions – What have we learned that make these statement false? • Increases in global temperatures in the atmosphere and the consequent warming of the oceans will only create a problem for people living along the coast. • Water exists in the ground in actual rivers or lakes that are constantly renewed. • People drink bottle water because it is better for our health; the safety of tap water is below consumption standards.