properties of water and aqueous solutions

4.Properties of Water and Aqueous Solutions

Course: Introduction to Biotechnology

Instructor: Moman Niazi

Properties of Water• Polar molecule

• Cohesion and adhesion

• High specific heat

• Density – greatest at 4oC

• Universal solvent of life

Polarity of Water

• In a water molecule two hydrogen atoms form single polar covalent bonds with an oxygen atom. Gives water more structure than other liquids– Because oxygen is more electronegative, the

region around oxygen has a partial negative charge.

– The region near the two hydrogen atoms has a partial positive charge.

• A water molecule is a polar molecule with opposite ends of the molecule with opposite charges.

• Water has a variety of unusual properties because of attractions between these polar molecules.

– The slightly negative regions of one molecule are attracted to the slightly positive regions of nearby molecules, forming a hydrogen bond.

– Each water molecule can form hydrogen bonds with up to four neighbors.

Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings

Fig. 3.1

HYDROGEN BONDS• Hold water molecules

together

• Each water molecule can form a maximum of 4 hydrogen bonds

• The hydrogen bonds joining water molecules are weak, about 1/20th as strong as covalent bonds.

• They form, break, and reform with great frequency

• Extraordinary Properties that are a result of hydrogen bonds.– Cohesive behavior

– Resists changes in temperature

– High heat of vaporization

– Expands when it freezes

– Versatile solvent

Organisms Depend on Cohesion

• Cohesion is responsible for the transport of the water column in plants

• Cohesion among water molecules plays a key role in the transport of water against gravity in plants

• Adhesion, clinging of one substance to another, contributes too, as water adheres to the wall of the vessels.

Hydrogen bonds hold the substance together, a phenomenon called cohesion

• Surface tension, a measure of the force necessary to stretch or break the surface of a liquid, is related to cohesion.

– Water has a greater surface tension than most other liquids because hydrogen bonds among surface water molecules resist stretching or breaking the surface.

– Water behaves as if covered by an invisible film.

– Some animals can stand, walk, or run on water without breaking the surface.


Fig. 3.3

Moderates Temperatures on Earth

• What is kinetic energy?

• Heat?

• Temperature?

• Calorie?

• What is the difference in cal and Cal?

• What is specific heat?

Celsius Scale at Sea Level

100oC Water boils

37oC Human body temperature

23oC Room temperature

0oC Water freezes

Water stabilizes air temperatures by absorbing heat from warmer air and releasing heat to cooler air.Water can absorb or release relatively large amounts of heat with only a slight change in its own temperature.

Three-fourths of the earth is covered

by water. The water serves as a

large heat sink responsible for:

1. Prevention of temperature

fluctuations that are outside the

range suitable for life.

2. Coastal areas having a mild

climate

3. A stable marine environment

Specific Heat is the amount of heat that must be absorbed or lost for one gram of a substance to change

its temperature by 1oC.

Evaporative Cooling

• The cooling of a surface occurs when the liquid evaporates

• This is responsible for:

– Moderating earth’s climate

– Stabilizes temperature in aquatic ecosystems

– Preventing organisms from overheating

Density of Water

• Most dense at 4oC

• Contracts until 4oC

• Expands from 4oC to 0oC

The density of water:

1. Prevents water from freezing from the bottom up.

2. Ice forms on the surface first—the freezing of the

water releases heat to the water below creating

insulation.

3. Makes transition between season less abrupt.

– When water reaches 0oC, water becomes locked into a crystalline lattice with each molecule bonded to to the maximum of four partners.

– As ice starts to melt, some of the hydrogen bonds break and some water molecules can slip closer together than they can while in the ice state.

– Ice is about 10% less dense than water at 4oC.


Fig. 3.5

Solvent for Life

• Solution– Solute

– solvent

• Aqueous solution

• Hydrophilic– Ionic compounds

dissolve in water

– Polar molecules (generally) are water soluble

• Hydrophobic– Nonpolar compounds

Most biochemical reactions involve solutes dissolved in water.

• There are two important quantitative proprieties of aqueous solutions.

–1. Concentration

–2. pH

Concentration of a Solution

• Molecular weight – sum of the weights of all atoms in a molecule (daltons)

• Mole – amount of a substance that has a mass in grams numerically equivalent to its molecular weight in daltons.

• Avogadro’s number – 6.02 X 1023

– A mole of one substance has the same number of molecules as a mole of any other substance.

MolarityThe concentration of a material in solution is called its molarity.

A one molar solution has one mole of a substance dissolved in one liter of solvent, typically water.

Calculate a one molar solution of sucrose, C12H22O16.

C = 12 daltons

H = 1 dalton

O = 16 daltons

12 x12 = 144

1 x 22 = 22

16 x 11 = 176

342

For a 2M solution?

For a .05 M solution?

For a .2 M solution?

• Occasionally, a hydrogen atom shared by two water molecules shifts from one molecule to the other.

– The hydrogen atom leaves its electron behind and is transferred as a single proton - a hydrogen ion (H+).

– The water molecule that lost a proton is now a hydroxide ion (OH-).

– The water molecule with the extra proton is a hydronium ion (H3O+).

Dissociation of Water Molecules


Unnumbered Fig. 3.47

• A simpler way to view this process is that a water molecule dissociates into a hydrogen ion and a hydroxide ion:

– H2O <=> H+ + OH-

• This reaction is reversible.

• At equilibrium the concentration of water molecules greatly exceeds that of H+ and OH-.

• In pure water only one water molecule in every 554 million is dissociated.

– At equilibrium, the concentration of H+ or OH-

is 10-7M (25°C) .

Acids and Bases• An acid is a substance that

increases the hydrogen ion concentration in a solution.

• Any substance that reduces the hydrogen ion concentration in a solution is a base.– Some bases reduce H+ directly by

accepting hydrogen ions.

• Strong acids and bases complete dissociate in water.

• Weak acids and bases dissociate only partially and reversibly.

pH Scale• The pH scale in any aqueous solution :

– [ H+ ] [OH-] = 10-14

• Measures the degree of acidity (0 – 14)

• Most biologic fluids are in the pH range from 6 – 8

• Each pH unit represents a tenfold difference (scale is logarithmic)

– A small change in pH actually indicates a substantial change in H+ and OH-

concentrations.

Buffers• A substance that eliminates large sudden

changes in pH.• Buffers help organisms maintain the pH of

body fluids within the narrow range necessary for life. –Are combinations of H+ acceptors and

donors forms in a solution of weak acids or bases

–Work by accepting H+ from solutions when they are in excess and by donating H+ when they have been depleted.

Acid Precipitation• Rain, snow or fog with more strongly acidic than pH

of 5.6• West Virginia has recorded 1.5• East Tennessee reported 4.2 in 2000• Occurs when sulfur oxides and nitrogen oxides react

with water in the atmosphere– Lowers pH of soil which affects mineral solubility

– decline of forests– Lower pH of lakes and ponds – In the Western

Adirondack Mountains, there are lakes with a pH <5 that have no fish.

Copyright McGraw-Hill 2009 24

4.1 General Properties of Aqueous Solutions

• Solution - a homogeneous mixture

– Solute: the component that is dissolved

– Solvent: the component that does the dissolving

Generally, the component present in the greatest quantity is considered to be the solvent. Aqueous solutions are those in which water is the solvent.


• Electrolytes and Nonelectrolytes

– Electrolyte: substance that dissolved in water produces a solution that conducts electricity

• Contains ions

– Nonelectrolyte: substance that dissolved in water produces a solution that does not conduct electricity

• Does not contain ions


• Dissociation - ionic compounds separate into constituent ions when dissolved in solution

• Ionization - formation of ions by molecular compounds when dissolved


• Strong and weak electrolytes

– Strong Electrolyte: 100% dissociation

• All water soluble ionic compounds, strong acids and strong bases

– Weak electrolytes

• Partially ionized in solution

• Exist mostly as the molecular form in solution

• Weak acids and weak bases


• Examples of weak electrolytes

– Weak acids

HC2H3O2(aq) C2H3O2

(aq) + H+(aq)

– Weak bases

NH3 (aq) + H2O(l) NH4+

(aq) + OH(aq)

(Note: double arrows indicate a reaction that occurs in both directions - a state of dynamic equilibrium exists)


Method to Distinguish Types of Electrolytes

nonelectrolyte weak electrolyte strong electrolyte


Classify the following as nonelectrolyte,

weak electrolyte or strong electrolyte

– NaOH

strong electrolyte

– CH3OH

nonelectrolyte

– H2CO3

weak electrolyte


4.2 Precipitation Reactions• Precipitation (formation of a solid from two

aqueous solutions) occurs when product is insoluble

• Produce insoluble ionic compounds

• Double replacement (or metathesis reaction)

• Solubility is the maximum amount of a solid that can dissolve in a given amount of solvent at a specified temperature

• Prediction based on solubility rules