BIOLOGY II
CHAPTER 3 NOTES
Water
contributes to the fitness of the environment to support life.
·
Life
on earth probably evolved in water.
·
Living
cells are 70%-95% H2O
·
Water
covers about ¾ of the earth.
·
In
nature, water naturally exists in all three physical states of matter-solid,
liquid and gas.
Water’s extraordinary properties are emergent properties resulting from water’s structure and molecular interactions.
I.
The polarity of
water molecules results in hydrogen bonding.
Water is a polar molecule. Its polar bonds and asymmetrical shape give water molecules opposite charges on opposite sides.
What does it mean to be polar?
·
Four
valence orbitals of O point to corners of a tetrahedron.
·
2
corners are orbitals with unshared pairs of electrons and weak negative charge.
·
2
corners are occupied by H atoms which are in polar covalent bonds with O. Oxygen is so electronegative, that shared
electrons spend more time around the O causing a weak postive charge near H’s.
Hydrogen bonding orders water into a higher level of structural organization.
·
The
polar molecules of water are held together by hydrogen bonds.
·
Positively
charged H of one molecule is attracted to the negatively charged O of another
water molecule.
·
Each
water molecule can form a maximum of four hydrogen bonds with neighboring water
molecules.
Water
has extraordinary properties that emerge as a consequence of its polarity and
hydrogen-bonding. Some of these properties are that water:
Ñ Has cohesive behavior
Ñ Has a high heat of
vaporization and cools surfaces as it evaporates.
Ñ Expands when it freezes
Ñ Is a versatile solvent
II.
Organisms depend on the
cohesion of water molecules.
Cohesion= Phenomenon of a substance being held together
by
hydrogen
bonds.
·
Though
hydrogen bonds are transient, enough water molecules are hydrogen bonded at any
given time to give water more structure than other liquids.
·
Contributes
to upward water transport (in plants) by holding the water column
together.
Adhesion of water to vessel walls
counteracts the downward
pull of gravity.(clinging
of one substance to another)
Surface
tension-
Measure of how difficult it is to stretch or break
the surface of a liquid.
·
Water
has a greater surface tension than most liquids; function of the fact that the
air/H2O interface, surface water molecules are hydrogen bonded to each other
and to the water molecules below.
·
Causes
H20 to bead (shape with smallest area to
volume ratio and allows maximum hydrogen bonding).
Example: soapy water bubble lab in Bio 1
H2O stabilizes temps.
III.
Water contributes to Earth’s habitability by moderating
temperatures
A. Heat and Temperature.
Kinetic energy= the energy of motion.
Heat=
Total kinetic energy due to molecular motion in a body of matter.
Calorie(cal)=Amount of heat it takes to raise the temperature of one gram of water
by one degree Celsius.
Kilocalorie(kcal or Cal)=Amount of heat required to raise the temperature of
one kilogram of water by one degree Celsius(1000 cal).
Temperature=Measure
of heat intensity due to the average kinetic energy of molecules in a body of
matter.
Joule= energy unit
1J=0.239cal
1cal=4.184J
*Know
conversions
|
Celsius Scale at Sea Level |
Scale Conversion |
|
100˚C (212˚F)= water boils 37˚ C (98.6˚F)= human body temperature 23˚ C (72˚F) = room temperature 0˚C (32˚F)= water
freezes |
˚C = 5(˚F-32) 9 ˚F= 9˚C +32 5 ˚K=˚C+273 |
B. Water’s High Specific Heat
Water has a high specific heat, which means that it
resists temperature changes when it absorbs or releases heat.
Specific heat= Amount of heat that must be absorbed or lost for one gram of a
substance to change its temperature by one degree Celsius.
Specific heat of water= One calorie per gram per degree Celsius (1 cal/g/˚C). (constant)
Ñ As a result of hydrogen bonding among water
molecules, it takes a relatively large heat
loss or gain for each 1°C change in temperature.
Ñ Hydrogen bonds must absorb heat to break, and
they release heat when they form.
Ñ Much absorbed heat energy is used to disrupt
hydrogen bonds before water molecules can move faster(increase temperature)
A large body of water
can act as a heat sink- absorbing heat from
sunlight during the day and summer and releasing heat during the night and
winter as the water gradually cools.
As a result:
Ñ Water, which covers three-fourths of the planet,
keeps temperature fluctuations within a range suitable for life.
Ñ Coastal areas have milder climates than
inland.(Florida)
Ñ The marine environment has a relatively stable temperature.
Why does H2O not flame?
C. Evaporative cooling
Vaporization-(evaporation)= Transformation of a liquid to a gas.
Ñ Molecules with enough kinetic energy to overcome the
mutual attraction of molecules in a liquid, can escape into the air.
Heat of vaporation- Quantity of heat a liquid must absorb for 1 g to be
converted to the gaseous state.
Ñ For water molecules to evaporate, hydrogen bonds
must be broken which requires heat energy.
Ñ Water has a relatively high heat of vaporization at
the boiling point(540 cal/g or 2260 J/g; Joule=0.239 cal).
Evaporative cooling- Cooling of a liquid’s surface when a liquid
evaporates.
Ñ The surface molecules with the highest kinetic
energy are most likely to escape into gaseous form; the average kinetic energy
of the remaining surface molecules is thus lower.
Ñ Sweating does not make you cool…..the sweat
evaporating does!
Ñ Which evaporates first?…..Water on skin or alcohol
on skin?
Water’s high heat of
vaporization:
1.Moderates the earth’s climate.
Ñ Solar heat absorbed by tropical seas dissipates when
surface water evaporates (evaporative
cooling).
Ñ As moist tropical air moves poleward, water vapor
releases heat as it condenses into rain.
2.Stabilizes temperature in aquatic ecosystems (evaporative cooling).
3.Helps organisms from overheating by evaporative cooling.
I. Oceans and lakes don’t freeze solid because ice floats
Because of hydrogen bonding, water is less dense as
a solid than it is as a liquid.
Consequently, ice floats.
Ñ Water is densest at 4°C.
Ñ Water contrasts as it cools to 4°C.
Ñ As water cools from 4°C to freezing(0°C), it expands and becomes less dense than liquid water(ice floats)
Ñ When water begins to freeze, the molecules do not
have enough kinetic energy to break hydrogen bonds.
Ñ As the crystalline lattice forms, each water
molecule forms a maximum of 4 hydrogen bonds, which keeps water molecules
farther apart than they would be in the liquid state.
Expansion of water contributes to the fitness of the environment for life:
iPrevents deep bodies of
water from freezing solid from the bottom.
iSince ice is less dense, it
forms on the surface first. As water
freezes
it releases heat to the water below and insulates it.
iMakes the transitions
between seasons less abrupt. As water
freezes, hydrogen bonds form
releasing heat. As ice melts,
hydrogen bonds break
absorbing heat.
Solution-A liquid that is a homogeneous mixture of two or
more
substances.
Solvent-Dissolving agent of a solution.
Solute-Substance dissolved in a solution.
Aqueous solution- Solution in which water is the solvent.
Water is a versatile solvent owing to the polarity
of the water molecule.
1.Ionic compounds dissolve
in water.
Ñ Charged regions of polar
water molecules have an electrical attraction to charged ions.
Ñ Water surrounds individual
ions, separating and
Hydrophilic shielding them from one
another.
2.
Polar
compounds in general, are water soluble.
·
Charged
regions of polar water molecules have an affinity for oppositely charged
regions of other polar moledules.
Hydrophobic
3.
Nonpolar compounds (which have symmetric
distribution in charge) are NOT water
–soluble.
A.
Hydrophilic and Hydrophobic
Substances
Ionic and polar substances are hydrophilic, but nonpolar compounds are hydrophobic.
Hydrophilic=(Hydro=water;philo=loving) Property of having an
affinity
for water.
iSome large hydrophilic
molecules can absorb water without
dissolving.
Hydrophobic=(Hydro=water; phobos=fearing) Property of not
having an
affinity for water, and thus not being
water-soluble.
Most biochemical reactions involve solutes dissolved
in water. There are two important
quantitative properties of aqueous solutions:
·
solute
concentration
·
pH.
Molecular weight-Sum of the weight of all atoms in a
molecule(expressed in daltons)
Mole- Amount of a substance that has a mass in grams
numerically
equivalent to its molecular weight in
daltons.
For example, to determine a mole of
sucrose(C12H22O11):
iCalculate molecular weight:
C=12 dal 12dal * 12=144 dal
H= 1 dal
1 dal *22= 22 dal
O=16 dal 16 dal *11= 176 dal
342 dal
i Express it in grams(342 g).
Molarity-Number of moles of solute per liter of solution.
iFor example, to make a 1M
sucrose solution, weigh out 342g of
sucrose
and add water up to 1 L.
Advantage of measuring in
moles:
iRescales weighing of single
molecules in daltons to grams, which is
more
practical for laboratory use.
iA mole of one substance has the same number
of molecules as a
mole of
any other substance(6.02*10>23; Avogadro’s
number)…..constant- the amount of molecules in a mole.
iAllows one to combine substances in fixed
ratios of molecules.
A.
Dissociation of Water
Molecules
Occasionally, the hydrogen atom that is shared in a hydrogen bond between two water molecules, shifts from the oxygen atom to which it is covalently bonded to the unshared orbitals of the oxygen atom to which it is hydrogen bonded.
iOnly a hydrogen
ion(proton with a +1 charge) is actually
transferred.
iTransferred proton binds to an unshared
orbital of the second
water molecule creating a hydronium
ion(H3O+).
iWater molecule that lost a proton has a net
negative charge
and is called a hydroxide ion(OH-).
H2O+H20àH3O(+)+OH(-)
iBy convention, ionization of H20 is expressed
as the
dissociation into H+ and OH-.
H2OàH(+)+OH(-)
iReaction is reversible.
iAt equilibrium, most of the H20 is not
ionized.
At equilibrium in pure water 25iC:
iNumber of H+ ions= number of OH- ions.
i[H+] = [OH+] = 1 M=10(-7) M
10,000,000
iNote
that brackets indicate molar
concentratoin.
NOTE: This is a good place to point out how few water molecules are actually dissociated (only 1 out of 554,000,000
molecules.)
Acid |
Base |
|
Substance
that increases the relative [H+] of
a solution. Also
removes OH+ because it tends to combine with H+ to form H2O. For
example: (in water) HCl
à H(+) + Cl(-) |
Substance
that reduces the relative [H+} of a
solution. May
alternately increase [OH]. For
example: A
base may reduce [H+] directly: NH3 + H[+] à NH4(+) A
base may reduce [H+] indirectly: NaOHà Na(+) + OH(-) OH(-)
+ H[+] à H2O |
A solution which:
i[H+] = [OH-] is a neutral solution.
i[H+] > [OH-] is an acidic solution.
i[H+] < [OH-] is a basic solution.
Strong acids and bases dissociate completely in water.
iFor example, HCl and NaOH.
iSingle arrows indicate complete dissociation.
NaOH à Na(+) + OH(-)
Weak acids and bases dissociate only partially and
reversibly.
iFor example, NH3 (ammonia) and H2Co3
(carbonic acid)
iDouble arrows indicate a reversible reaction;
at equilibrium there
will be a fixed ratio of reactants
and products.
H2CO3 ®
Hco3(-) + H(+)
Carbonic acid Biocarbinate
ion Hydrogen ion
In any aqueous solution:
i[H+] [OH-]=1.0 x 10(-14)
For example:
i In a neutral solution, [H+] = 10(-7)M
and [OH-]=10(-7) M.
iIn an acidic solution where the [H+] =
10(-5)M, the [OH-]=
10(-9) M.
iIn a basic solution where the [H+]=10(-9) M,
the [OH-]=
10(-5) M.
pH scale = Scale used to measure degree of acidity. It
ranges from 0 to 14.
pH= Negative log(10) of the [H+] expressed in moles
per liter.
ipH of 7 is a neutral solution.
ipH <7 is an acidic soution
ipH >7 is basic solution
iMost biological fluids are
within the pH range of 6 to 8. There
are some expceptions such as stomach acid with pH = 1.5.
iEach pH unit represents a
tenfold difference (scale is logarithmic), so a slight change in actual[H+].
Example: How
much greater is the [H+] in a solution with pH
2 than
in a solution with pH 6?
Ans: pH 2= [H+] of 1.0 x 10(>-2)= 1 M
100
pH 6= [H+] of 1.0 x 10(>-6)= 1
M
100
10,000 times greater [molar
concentration]
By minimizing wide fluctuations in pH, buffers help organisms maintain the pH of body fluids within the narrow range necessary for life. (usually pH 6-8). (7.4) for human blood.
Buffer- Substance that prevents
large sudden changes in pH.
Ñ Are combinations of H+ donor and H+ acceptor forms of weak acids or bases.
Ñ Work by accepting H+ ions from solution when they are in excess, and by donating H+ ions to the solution when they have been depleted.
Ñ Most buffers work in acid/base pairs
IV.
Acid precipitation
threatens the fitness of the environment
Acid precipitation- Rain, snow or fog more strongly acidic than
pH 5.6.
Ñ Has been recorded as low as pH 1.5 in West Virginia
Ñ Occurs when sulfur oxides and nitrogen oxides in the
atmosphere react with water in the air to form acids
which fall to Earth in precipitation.
Ñ Major oxide source is the combustion of fossil fuels by industry and cars.
Acid Rain affects the fitness of the
environment to support life.
Ñ Lowers soil pH which affects mineral solubility. May leach out necessary mineral nutrients and increase the concentration of minerals that are potentially toxic to vegetation in higher concentration(e.g. aluminum). This is contributing to the decline of some European and North America Forests.
Ñ Lowers the pH of lakes and ponds, and runoff carries leached out soil minerals into aquatic ecosystems. This adversely affects aquatic life. For example: In the Western Adirondack Mountains, there are lakes with a pH <5 that have no fish.
What can be done to reduce the problem?
Ñ Add industrial pollution controls.
Ñ Develop and use antipollution devices
Ñ Increase involvement of voters, customers, politicians and business leaders.