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Water and the Fitness of

the Environment

Water is the substance that makes possible life as we know it here on Earth

Earth = ¾ H2O

It is the only substance to exist in the natural environment in all three physical states of matter

Life on Earth began over 3 billion years long before it was able spread to land

All living organisms require water more than any other substance

Most cells are made up of a about 70-90% water and surrounded by it as well

POLARITY OF THE WATER MOLECULE

By itself, it seems small and quite insignificant

Electronegativity - oxygen

Type of bond - polar covalent

Hydrogen bonds

Why is it unlikely that two neighboring water molecules would be arranged like this??

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PROPERTIES OF WATER COHESION

Surface tension

ADHESION

HIGH SPECIFIC HEAT

HIGH HEAT OF VAPORIZATION

Evaporative cooling

COHESION Occurs because of the presence of hydrogen bonds

At any given moment, many molecules are linked by multiple hydrogen bonds

Cohesion contributes to the transport of water and dissolved nutrients against gravity in plants

Surface tension: related to cohesion

Measure of how difficult it is to break the surface of liquid

It makes water behave as if it were coated with a film

ADHESION

Clinging of one substance to another

Water, besides ‘sticking’ to other water molecules can also stick to other surfaces, like cell walls, through H bonds

This quality also helps water ‘climb’ capillary tubes (fine tubes).

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SPECIFIC HEAT Heat is a form of energy

Heat is the total amount of matter’s kinetic energy

Heat is partially dependent on the volume considered

Can be transferred between two bodies that are at different temperatures

Temperature is a measure of heat intensity

It represents the average kinetic energy of the molecules (regardless of the volume)

EX: glass of water in a pan vs. swimming pool

When two objects with different temperatures are brought together, heat passes from the warmer to the cooler, until the two are at the same temperature

Molecules in the cooler object speed up at the expense of the kinetic energy of the warmer object

Units of Temperature Celsius (˚C)

Calorie (cal): amount of heat required to raise the temp of 1 g of water by one degree Celsius

Kilocalorie (kcal)

1 J = 0.239 cal

1 cal = 4.184 J

Specific heat can be thought of as a measurement of how well a substance resists changing its temperature when it absorbs or releases heat

Amount of heat that must be absorbed or lost for 1 g of that substance to change its temperature by 1 ˚C

The specific heat of water is 1 cal per gram per degree Celsius (1 cal/g/˚C)

EX: specific heat of ethanol is 0.6 cal/g/˚C.

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When water is compared to other substances, it has a higher specific heat

This means that it will change its temperature less when it absorbs or loses a given amount of heat

EX: you can burn your fingers by touching the side of a pan on the stove while the water is still only warm.

Like many of its properties, water’s specific heat is due to the presence of hydrogen bonds

Much of the heat is used to disrupt hydrogen bonds before the water molecules can begin to move faster

This is how water moderates the air temperature: it absorbs the heat from air when it is warmer, and releases the stored heat back to the air when it is cooler

HEAT OF VAPORIZATION Amount of heat that must be absorbed to convert 1 g of

a liquid to a gas

EX: to evaporate 1 g of water at 25 ˚C you need about 580 cal of heat, nearly double the amount needed to vaporize a gram of alcohol or ammonia

Evaporative Cooling Evaporation: a very slow evaporation occurs even at

low temperatures because some molecules will be able to escape into the air

If a liquid is heated, the liquid evaporates more rapidly.

Evaporative cooling: while the molecules with the greatest kinetic energy leave to the gas state, the molecules left behind will have less kinetic energy

Effects: It moderates Earth’s climate: a considerable amount of

solar heat is absorbed by the seas

Moist tropical air circulates towards the poles releasing heat as it condenses to form rain

Explains the severity of steam burns

Lakes and ponds: they have stable temperature

Plants: evaporation of the leaves helps prevent overheating

Humans: dissipates body heat and helps prevent overheating

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ICE Water is less dense as a solid than as a liquid: ice floats

The fact that ice floats is important in the fitness of the environment

Floating ice insulates the liquid water below and prevents it from freezing

During summer only a few inches of ice would thaw

Water and the Fitness of

the Environment (part 2)

THE SOLVENT OF LIFE

Solution: a liquid that is a completely homogeneous mixture of two or more substances is called a solution:

Solvent: dissolving agent

Solute: substance that is dissolved

An aqueous solution is one in which water is the solvent (not a universal solvent, but a very versatile solvent)

Example: NaCl + H2O

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Ions are exposed to the solvent

Hydration shells

How many solutes?

Most ionic compounds dissolve in water

Ex

Seawater

Cytoplasm

Hydrophilic / Hydrophobic Hydrophilic (from the Greek hydro, water, and philios, loving) is any substance that has affinity for water

Hydrophilic substances have regions of partial positive an partial negative charges (H bonds)

Not all substances that are hydrophilic dissolve in water

Made of cellulose

Water can adhere to the fibers (wet), yet it does not dissolve

Cellulose is also present in the water-conducting cells in the plant

If the particles are small enough, they can instead remain suspended in the aqueous liquid forming a colloid (stable suspension of fine particles in a liquid)

Ex: milk, fog, whipped cream, mayonnaise, paint, ink, marshmallow, butter, pearls and opals

Hydrophobic (from the Greek hydro, water, and phobos, fearing): any substance that is mainly made up of nonpolar bonds and cannot mix stably with water-based substances.

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Solute Concentration in Aqueous Solutions Most of the chemical reactions in organisms involve

solutes dissolved in water. When carrying out experiments, we use mass to calculate the number of molecules.

From the Periodic Table of Elements we can obtain the atomic weight. By knowing the number of atoms present in specific molecule, we can calculate its molar mass.

Molar mass

Ex: H2O

H = 1.008 g/mol x 2 = 2.016 O = 16.00 g/mol x 1 = 16.00 = 18.016 g/mol

Ex: C12H22O11 (sucrose, table sugar)

C = 12.01 g/mol x 12 = 144.12 H = 1.008 g/mol x 22 = 22.176 O = 16.00 g/mol x 11 = 176

So, 1 mole of sugar weighs 342.296 grams.

MOLE

Chemical mass unit, defined as 6.022 x 1023 particles

These particles can be atoms, molecules, ions

Ex: one mole of marshmallows would have a mass of 17 exatonnes (= 1 quintillion or 1 billion-billion tons), enough the cover the whole surface of the Earth approx 100 Km deep (60 miles).

WHY MOLE?? A mole represents an exact number of objects

One mole of one substance will always contain the same number of molecules as a mole of another substance

Measuring in moles makes it convenient for scientists working in the laboratory to combine substances in fixed ratios of molecules.

MOLARITY It is the number of moles of solute per liter of solution

Ex: a 1M solution of sugar would have 1 MOLE of sugar in one LITER of water.

M = #moles / L

Ex: how do you make a 1 M solution of sucrose?

Molar Weight: 342.296 grams

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ACIDIC AND BASIC CONDITIONS About one in every 554 million water molecules

dissociates

This means that a H atom shifts from one molecule to the other

The hydrogen atom leaves its electron behind, and what is actually transferred is a Hydrogen Ion (H+), a single proton with a charge of 1+.

H+ does not exist on it’s own, but by convention H+ is used to represent the Hydronium ion

H+ / OH-

These ions are of biological and chemical importance because they are very reactive

Changes in their concentration can drastically affect a cell’s proteins and other complex molecules

The concentrations of H+ and OH- are equal in pure water, but this balance can be disrupted by adding certain solutes like acids and bases

ACIDS When an acid dissolves in water, it increases the

hydrogen ion concentration in the solution

Ex: hydrochloric acid

This dissociation results in an acidic solution, one that haves more H+ than OH-

HCl is a strong acid because it dissociates completely when mixed in water

BASES A substance that reduces the H+ concentration in a

solution is called a base

Strong bases, like NaOH, reduce the H+ indirectly by dissociating to form hydroxide ions

Weak base (reaction with double arrow): NH3

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pH SCALE In a neutral solution at room temp (25 ˚C):

[H+] = 10-7

[OH-] = 10-7

Brackets indicate molar concentration. The product of the H+ and OH- concentrations is constant at 10-14

So: [H+] [OH-] = 10-14

The pH scale compresses the range of H+ and OH-concentrations by employing logarithm.

The pH solution is defined as the negative logarithm (base 10) on the hydrogen ion concentration:

pH = -log [H+]

pH declines as H+ concentration increases

The pH of a neutral aqueous solution at 25 ˚C is 7, the mid point of the scale

Each pH unit represents a tenfold difference in H+ and OH- concentrations: a solution of pH 3 is a thousand times more acidic than a solution of pH 6

Less than 7 denotes an acidic solution

More than 7 is a basic solution

Most biological fluids are within the range pH 6-8

There are exceptions: digestive juice has a pH of about 2.

BUFFERS The internal pH of most living cells is close to 7

Even a slight change in pH can be harmful

The pH of human blood is very close to 7.4.

A person cannot survive for more than a few minutes if the blood pH drops to 7 or rises to 7.8

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If you add 0.01 mol of a strong acid to a liter of pure water, the pH would drop from 7.0 to 2.0. (10,000x)

If the same amount of acid were added to a liter of blood, the pH decreases only from 7.4 to 7.3 (0.1x)

This is because a chemical system in the blood allows to maintain a stable pH

The presence of substances called buffers allow for a relatively constant pH in biological fluids despite the addition of acids or bases

Buffers are substances that minimize changes in the concentrations of H+ and OH- in a solution

They do so by accepting H+ from the solution when they are in excess, or donating H+ to the solution when they have been depleted

Ex: carbonic acid / bicarbonate buffering in blood

CO2 + H2O ↔ H2CO3 ↔ HCO3– + H+

from cellular carbonic acid bicarbonate ion hydrogen ion

respiration

Bound by Hemoglobin at low pH