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Importance of water
Structure of the H2O molecule
States of H2O, and energy transfer in changes of state
Density, and effects of temperature
Sea water -- the effect of dissolved salts
Powerpoint Lecture Slides
IMPORTANCE OF H2O
Most common substance at surface; rare on other planets
Essential for life
Absorbs (and releases) large amounts of solar energy
STRUCTURE OF H2O MOLECULE
Non-linear --> electrically "polar" --> additional attractive force
"Hydrogen-bonding" in water and ice
Strong solvent for other polar substances (ionic solids)
STATES (PHASES) OF H2O
Ice -- All H2O molecules are H-bonded in a regular crystalline structure
Water -- Free molecules and ice-like clusters
Vapor -- Free molecules only
ENERGY INVOLVED IN H2O PHASE CHANGES
Ice <-->Water -- 80 calories/gram at 0 deg C
... required (absorbed) to melt ice
... liberated when ice freezes
Water<-->Vapor -- 540 calories/gram at 100 deg C
... required (absorbed) to evaporate liquid water
... liberated when water vapor is condensed.
[Water evaporates and condenses at <100 deg C.
For example, the energy absorbed/liberated is 585 cal/g
at 20 deg C.]
Why so much energy? -- to break H-bonds in water and ice.
* evap., melting --> energy required to break H-bonds
* condensation, freezing --> energy liberated
when H-bonds form
Importance of H2O phase changes
Evap. of sea water --> water vapor in atm.
Condensation of water vapor: principal source for
HEAT CAPACITY (or specific heat)
How efficiently heat is stored in a substance
"Energy required to raise T of one gram of substance by 1°C."
Heat capacity of water = 1 calorie / g - °C ... very high!
Heat capacity of rocks and soils = 0.2 cal. / g - °C.
Consequences of difference between water and land
Bodies of water gain/lose much heat with little T change.
Land heats and cools much more.
DENSITY -- mass per unit volume, grams/cm3
Pure water (at 4°C) = 1.000 g/cm3
Ice = 0.92 g/cm3 . . . unusual behavior!
Sea water = 1.021 - 1.028 g/cm3
Large water bodies tend to be stratified according to density
Top layers -- least dense
Bottom layers -- most dense
TEMPERATURE EFFECT ON DENSITY OF PURE WATER
Maximum density at 4°C
Higher T --> density decreases (normal behavior)
Lower T (4°C to 0°C) -- density also decreases! Why?
Lakes during winter cooling
Surface freezes at 0°C
Bottom stays at 4°C and does not freeze -- good for fish!
Salinity - - dissolved salt content
Average S = 35 g/kg (p.p.t, o/oo)
Range of S (99% of all sea water) = 30 -37 o/oo
Dissolved salts change inter-molecular interactions and thus physical properties of sea water.
Boiling point = 103°C
Freezing point lowered
Sea water begins to freeze at -2°C; salt is excluded from ice.
Remaining water is saltier, freezes at lower T.
Temperature of maximum density "dissapears"
Density increases progressively to freezing point
Higher density promotes sinking of cold sea water
... increases as S increases.
... increases as T decreases.
Importance to deep circulation in oceans:
Deep-water "masses" form at surface
... cooling (T decreases)
... freezing & evaporation (S increases)
Sink to a level (depth) governed by density, and spread out
Detailed Notes Start Here
Although we all take it for granted, H2O is an exceptionally
important substance on our planet.
* As water, ice, and vapor in the atmosphere, H2O is the most common substance on Earth's surface; but it is rare on other planets in the Solar System.
* Liquid H2O (water) is essential for life, and the medium in which life first developed.
* H2O also has important (and unusual) physical and chemical properties:
- Water is an excellent "solvent" (it can dissolve large amounts of many different kinds of substances).
- Water in the oceans absorbs and stores large amounts of solar energy (heat).
- The fact that H2O can exist as a solid and gas as well as a liquid at Earth's surface is important in regulating the climate of Earth.
The H atoms are bonded to the O atom by strong covalent bonds.
This is common for many simple molecules made up of two or more
atoms. But unlike many molecules, H2O is non-linear; that is,
the angle between H-O bonds is 105 degrees (instead of 180 degrees
as in linear molecules like CO2). The non-linear results in an
unequal distribution of electrical charge in the molecule: the
"H end" is positively charged, and the "O end"
is negatively charged. The fact that H2O is an electrically "polar"
molecule accounts for many of its unusual properties:
* In water and ice, H2O molecules are held together by a type of chemical bonding called "hydrogen bonding" (or H-bonds) where opposite ends of molecules are oriented toward one another.
* Intra-molecular hydrogen bonding is surprisingly strong, equivalent to about 10% of the energy of covalent bonds between H and O atoms.
* The additional attractive force of hydrogen bonding accounts for:
- the molecular structure of crystalling ice and liquid water
- the strong ability of water to dissolve other polar substances, such as ionic solids (salts) - - D&D T-25 c, d
The states of H2O - - As everyone knows, H2O exists in three distinct states (or phases) - - Ross T-29, a
A significant amount of energy is liberated or absorbed in these phase changes.- Ross, T-31, right side
Ice <-->Water -- 80 calories/gram at 0 deg C
required (absorbed) to melt ice
liberated when ice freezes
Water<-->Vapor -- 540 calories/gram at 100 deg
required (absorbed) to evaporate liquid water
liberated when water vapor is condensed.
[Water evaporates and condenses at <100 deg C. For example, the energy
absorbed/liberated is 585 cal/g at 20 deg C.]
Why is so much energy involved in the phase changes of H2O? -- because of the strength of H-bonds in water and ice. A lot of energy is required to break H-bonds (melting, evaporation). Conversely, the same amounts of energy are liberated when H-bonds are formed (freezing, condensation).
Phase changes of H2O at Earth's surface are important processes:
* Evaporation of surface sea water is the principal source of water vapor in the atmosphere; it drives the hydrologic cycle
* Condensation of water vapor is the principal source of heating the atmosphere (even more than direct heating by solar radiation).
is a property of a substance that describes how efficiently the substance holds heat energy. Heat capacity is defined as "the energy required to raise the temperature of one gram of a substance by 1 deg C". The heat capacity of liquid water is 1 [calorie} / [gram - deg C]. That is, it takes one calorie of heat to raise the temperature of one gram of water by one degree C. A unique feature of water is that its heat capacity is the highest of all common liquids and solids. For example, the heat capacity of common rocks (such as basalt and granite) and soils is only about 0.2 [calorie] / [gram - deg C].
The large contrast in heat capacity between water and earth
materials has important consequences:
* Large bodies of water (lakes, ocean) can gain or lose large quantities of heat (thermal energy) with little change in temperature. [The high heat of evaporation contributes to this effect also.]
* In contrast, land surfaces heat and cool to a greater extent in response to the same input of solar energy.
Density - of a substance describes its mass per unit
volume and is usually expressed as grams per cubic centimeter
* The density of pure water (at 4 deg C) is 1.000.
* The density of ice is 0.92. This is another unusual property of H2O! For almost all other substances, the density of the solid phase is greater than that of the liquid phase. Why does this "density inversion" occur for H2O?
Large bodies of water tend to be "stratified" according to density. The least dense water is at the surface, and the most dense water is at the bottom; this is termed "stable density stratification." For example, the oceans are stratified by density, with surface waters at about 1.021-1.023 and bottom water at 1.028. -- D&D Fig. 7.3
The maximum density of pure water occurs at 4 deg C. At higher
temperatures, density decreases because of "thermal expansion"
(most liquids behave this way). At lower temperatures, density
also decreases. This is because more "open" ice-like
clusters of water molecules become important at <4 deg C. This
unusual density behavior at low temperatures has implications
for the cooling of a fresh-water lake during the winter. [Consider
* Ice forms at the surface of lakes when the temperature reaches 0 deg C, the freezing point.
* But the bottom of the lake -- the densest layers -- remain at 4 deg C and will not freeze.
* Only if we could vigorously stir the lake (not likely) could bottom waters reach the freezing point.
The total dissolved salt content of a parcel of sea water is its salinity. The average salinity of sea water is 35 grams of dissolved salt per kilogram of sea water [35 g/kg, or 35 parts per thousand (p.p.t. or o/oo). 99% of all sea water has a salinity in the range 30 to 37 g/kg.
The presence of dissolved salts (as charged ions) effects the
physical properties of sea water by altering interactions between
* Boiling point is elevated. Average sea water would boil at 103 deg C. (not very important because no place in the surface ocean is that hot).
* Freezing point is depressed
- Sea water begins to freeze at -2 deg C.
- Salt is excluded from ice.
- Unfrozen water is saltier and freezes at an even lower temperature. (There will always be some saline brine left at very low T)
* Temperature of maximum density is depressed
- Seawater does not have temperature of maximum density above the freezing point
- As sea water cools, it becomes progressively denser until it freezes.
- Progressive increase in density promotes sinking and vertical circulation.
- warm, surface waters1.021 g/cm^3
- cold, deep waters1.028 g/cm^3
As we shall discuss later, density of sea water increases as temperature decreases and as salinity increases. This is important to deep circulation in the oceans. Deep-water "masses" form at the surface of oceans by cooling and increase in salinity (freezing, evaporation). They sink to a level (depth) governed by density.
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