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. Nucleus = protons (positive) + neutrons (neutral)
Basic chemistry
Atomic structure
Nucleus = protons (positive) + neutrons (neutral)
Electrons in orbitals around nucleus (negative)

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

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-

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

9. Allows formation of H-bonding between water molecules
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

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 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

13. 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

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

15. 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

16. 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

17. Unusual thermal properties of H2O
High latent (hidden) heats of
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)

18. Fig. 5.6

19. 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

20. 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

21. 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

23. 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

24. Density of water
Fig. 5.10

25. 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

26. 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

27. 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

28. 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 = * Chlorinity (ppt)
Specific gravity (1.028 g/ml)
Hydrometer
Electrical conductivity
Salinometer

29. Pure water vs. seawater

30. 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 hypersaline conditions
Salinity may vary with seasons (dry/rain)

31. How to change salinity Add/remove water
Add/remove dissolved substances

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

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

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

35. 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

36. Residence time and steady state
Fig. 5.16

37. 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

38. Figure 5.17

39. 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. 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

43. 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

44. 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

45. 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

46. 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

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

49. 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.