An Online Introduction

to Advanced Biology

 
   

Terms and Concepts

 
 

SUBSITE THREE

CHAPTER 3 -

Chemistry:  The Properties of Water

 
   

 

 
 

WATER - THE ELIXIR OF LIFE

 
 


Life on Earth could not exist without water.  Discussions about possible life on Mars or on some of the moons of Jupiter all connect to the possible presence of water there.  Why is water so important?

 
     

 

 

WATER AS SUPPORT FOR LIVING SYSTEMS

 
 


The bipolar nature of water molecules was discussed in the section on hydrogen bonding, and much of what water can do that other liquids can't is based upon water's polar nature.  First, water is important as the support liquid for the "soup" of ions and molecules that absolutely must be free to move around in order to produce the complicated chemistry of Life.  Ions and certain parts of molecules attract one end or the other of water molecules, making it possible for them to dissolve.  The "layer" of water molecules surrounding dissolved ions and molecules is called a hydration shell When things dissolve, they are called solutes and the supporting liquid is the solvent, and water is an excellent solvent.  Living cells are a combination of water, lots of various materials dissolved in the water (that is all a solution), and oily membrane barriers that allow isolation of different chemical processes, including the broad isolation of a cell from its surroundings.

A side effect of water molecules' attraction for other molecules (called adhesion) is the ability of water to climb small tubes by sticking to the walls.  Water molecules' powerful attraction for each other (called cohesion) allows columns of water to hold together even in the tallest of trees.  Materials with an attraction for water molecules are generally water-soluble, and are called hydrophilic;  materials that do not have this attraction are not soluble and are called hydrophobic;  these are used in the cell membrane barriers already discussed, among other things.

 
     

 

 

WATER AS AN ENVIRONMENTAL STABILIZER

 
 


Water as a medium on this planet also serves a role environmentally.  It is very stable as temperatures change because of the hydrogen bonds between the water molecules themselves.  If you think of heat as energy that increases the motion of atoms and molecules, you won't be far from a technical definition.  When water absorbs heat, the molecules do move faster but the "glue" of the hydrogen bonds slows the process.  Heat is transmitted slowly through water, and large bodies of water take a long time to heat up.  That energy does not quickly come back out if the temperature around the water drops, either, so masses of water cool more slowly than other liquids would.

Masses of water also are stable in that individual molecules that have picked up heat / motion have a hard time getting free into the atmosphere, or evaporating.  There are two reasons for this:  the first is that heat is passed around fairly evenly (if slowly) and so the number of molecules going fast enough to zip free rises slowly;  the second is that hydrogen bonding at the water-atmosphere surface pulls molecules at the surface more tightly together (this produces the surface tension that lets you put slightly more water into a glass than its volume - you've seen the "dome" of water on top? - or that bugs walk across without sinking) and leaves less room for potential evaporating molecules to squeeze through.  This also explains why evaporation is a cooling process (useful in sweating, or dogs panting):  what is lost during evaporation are the very fastest, very hottest molecules, leaving what's behind cooler.

Connected to the above property is the wide range of temperature in which water is liquid;  although life is tricky at extremes near freezing and boiling, it is possible, as long as water remains liquid.

Another fairly unique property of water is how it solidifies:  if water cools, its molecules move more slowly, collide more rarely, and tend to pack more closely together.  Like most substances, water gets more dense as it cools.  However, when too crowded, at about 4 C, the repulsions among the tightly-packed bipolar molecules cause them to slip into an arrangement which, as the temperature drops, actually pushes them further apart into kind of a crystal arrangement.  We all know that ice floats;  what this means is that water in its solid form is less dense than water in its liquid form.  If ice did not float, it would freeze, sink, and expose more surface to freeze, and sink, and frozen bodies of water would be frozen solid from bottom to surface, a very poor environment for living things and a difficult task to thaw.  In fact, floating ice acts as an insulator to the water underneath it.  The thicker the ice, the harder it is for the water to lose heat and freeze, so very few deep bodies of water, even in the coldest climates, are totally frozen.

 
     

 

 

WATER IONIZATION AND pH

 
 


It's easy to think of molecules as solid objects whizzing around space, but with all of their parts they are constantly shaking and may even pop apart occasionally.  In a typical container of pure water, .0000001 (10-7) of the H2O molecule pairs have collided and come apart at any given moment into two ions, H+ (hydrogen ion, actually an individual proton, and which tends  to associate with water molecules as a hydronium ion, H3O+), and OH- (hydroxide ion).  They do not stay apart indefinitely, and while some water molecules are coming apart, others are reforming.  While apart, though, either of these ions can react with other materials, and being tiny and charged, they can get into and react with many sorts of materials.  In pure water, the fact that the two ions are in balance and that there aren't many of them limits what they can do - water has little chemical effect and is considered neutral.  However, any material dissolved in the water that changes the balance of hydrogen or hydroxide ions will also change the chemistry of the solution.  A solution rich in hydrogen ions is an acid, while one rich in hydroxide ions is a base.

The pH of a solution is based on the concentration of hydrogen ions (pH = proportion of Hydrogen), and the numbers of the pH scale come from the decimals of that proportion.  Pure water, with a hydrogen ion concentration of 10-7, has a pH of seven (a scale with negative numbers was, apparently,  just too confusing);  anything that releases more hydrogen ions changes the concentration and drops the pH below seven and is considered an acid.  A material that lowers the hydrogen ion concentration will raise the hydroxide concentration and the pH above seven and is considered a base (older terminology would have called it an alkali).  From this you should see that every unit change on a pH scale is actually a tenfold change in hydrogen ion concentration.  Both of the ions are small and chemically unstable, so the more there are, the more they can disrupt the chemistry of other materials, especially large biological molecules whose shapes depend upon hydrogen bonds - an environment with lots of charged particles can disrupt those weak charge attractions, which is why stomach acid is an important first step in breaking down the molecules in food.  Strangely enough, dissolved in our stomach acid is a large protein-digesting molecule, pepsin, which can only keep its active shape in the presence of lots of hydrogen ions, another example of real life appearing to break the rules.

 
     

Links -

An advanced look at the principles of water ionization.

Recent water discoveries.

Lots of water links.

A very different look at water, under an alias (not to be taken seriously!).

 

     
 

Terms and Concepts

In the order they were covered.

 

Water as Solvent  
Hydration Shell  
Solutes  
Solvents  
Solutions  
Adhesion  
Cohesion  
Hydrophilic  
Hydrophobic  
Water and Heat  
Heat  
Evaporation  
Surface Tension  
Water as Ice  
Ionization of Water  
Hydrogen, Hydronium, & Hydroxide Ions  
Acid - Neutral - Base  
pH  
Hydrogen Ions & pH  
pH Scale  
   

 
     
 

 
 

GO ON TO THE NEXT CHAPTER - ORGANIC CHEMISTRY INTRODUCTION

 
     

 

Online Introduction to Biology (Advanced)

Copyright 2003 - 2011, Michael McDarby.

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