CHAPTER 5 Water and Seawater

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

Nucleus = protons (positive) + neutrons (neutral) 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

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

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

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

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-

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

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 http://www.mie.utoronto.ca/labs/lcdlab/biopic/fig

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

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

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

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

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

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

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

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)

Fig. 5.6

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

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

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

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

Density of water Fig. 5.10

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

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

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

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

Pure water vs. seawater

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)

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

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

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

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

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

Residence time and steady state Fig. 5.16

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

Figure 5.17

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

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

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

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

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

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

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

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.