Importance of Water for Life

 

Just as it is a difficult task to concisely define what constitutes life on Earth, it is equally difficult to construct a list of environmental requirements that represents the bare essentials for life.  In general, astrobiologists are in agreement that life needs a liquid medium for transporting the molecules that are necessary for its existence.  For life on Earth, this liquid medium is water.  Many other liquids have been proposed as potential substitutes for water, such as ammonia (NH3), methane (CH4), or ethane (C2H6).  However, as we will see, water is a versatile and somewhat unique molecule that is extremely compatible for life on Earth, and probably other planets.

 

As we begin our discussion of the importance of water for life, it is important to recognize that water is possibly the most abundant liquid in the universe.  But does that necessarily make it the best candidate for life?  What are the properties of water that make it well suited?  Let’s begin by examining a single molecule of water.  Water is composed of two hydrogen atoms and a single oxygen atom; there are a total of two chemical bonds in the molecule.  Each of the hydrogen atoms is bound to the oxygen by a covalent bond – a bond formed by the sharing of one or more electrons between two atoms.  In the case of the oxygen-hydrogen covalent bond, the electrons are not shared equally between the two atoms.  In fact, oxygen is a strongly electronegative atom and can be thought of as a bit of an “electron hog”.  The oxygen will tug on electrons so that there will be a higher density of electrons around the oxygen as compared to the hydrogen.  As a result, if you were to examine just one of the covalent bonds in a molecule of water, you would notice that the oxygen atom would be slightly negative and the hydrogen atom would be slightly positive; we would say that the covalent bond is polar.  This may not seem important just yet, but the interactions of hydrogen and oxygen atoms between water molecules are the key to understanding the uniqueness of water. 

 

In a water molecule, the oxygen atom can be described as having a net negative charge and the hydrogen atom as having a net positive charge.  When two objects with opposite charge are very close to one another, they will experience an attractive force between them, much like when two magnets “stick” together.  On the flip side, two objects with similar charge will experience a repulsive force between them.  You can probably imagine that the hydrogen atoms of one water molecule, then, might be attracted to the oxygen atom of another water molecule.  The attraction between these two atoms of different molecules is an example of hydrogen bonding.  Hydrogen bonding isn’t as strong as bonds where electrons are shared or exchanged and can only happen when the molecules are fairly close together.  Nevertheless, it is water’s propensity to engage in hydrogen bonding that gives it some of its unique character.  Hydrogen bonding is responsible for the fact that ice in liquid water floats, that water remains liquid over a large range of temperatures, that water has a high specific heat, and that strands of DNA stay “zipped” together. 

 

Ever notice that a pond or lake during a cold winter never freezes solid?  Ever wonder why?  Well, you can blame it all on our friend the hydrogen bond.  As liquid water gets colder and colder, the number of hydrogen bonds between water molecules gradually increases.  When water freezes and becomes ice, the hydrogen bonds will hold the water molecules in such a way that they form a crystal lattice.  In this crystal lattice, the water molecules are actually more spread out than when they are in a liquid state, which means they take up more volume.  The actual mass of water molecules does not change.  If we increase the volume, yet maintain the mass, we decrease the density of a substance.  We can determine this fairly easily from the simple mathematical equation of d = m/V, where d represents density, m represents mass, and V represents volume.  Therefore, ice is LESS dense than water – making it float!  Let’s think about that pond in the cold winter.  As it starts to freeze, all of the ice floats to the top.  Eventually the pond might freeze over completely, sealing in a liquid oasis underneath where life can survive the winter.  This characteristic of water, becoming less dense when it freezes and therefore making it float, allows aquatic environments to survive from year to year.  Can you imagine what would happen to aquatic life if the pond froze solid every winter?

 

Hydrogen bonding is also responsible for two important properties of water, its ability to remain a liquid over a large temperature range and its high specific heat.  Why are these two properties important for life?  On Earth, a large percentage of living organisms is comprised of water.  Let’s imagine that instead of remaining liquid over a temperature range of 100 degrees Celsius, water was only liquid over a range of 20 degrees Celsius.  Over the course of a day, temperatures on Earth can vary by greater than 20 degrees.  For example, on an extreme day in Arizona, temperatures can go from around 20 degrees Celsius to 45 degrees Celsius.  If, instead of remaining liquid over that temperature range, water changed from a liquid to a gas and began to boil; the effects could be devastating for life.  Imagine, for instance, what would happen to the cells in your body, consisting mostly of water!  Or consider the opposite case of a temperature change resulting in water freezing.  If the liquid that life relies on were constantly frozen, how would life carry out the necessary chemical reactions?  Luckily, water is liquid over a large range of temperatures – 100 degrees Celsius.  This range is large enough to insure that the water in cells will neither freeze nor boil off and therefore make it difficult for life to exist.

 

Equally important to a large temperature range is the high specific heat of water.  Specific heat is just a fancy way of describing the amount of energy that is required to raise the temperature of a substance.  In the case of water, it is the amount of energy to raise the temperature of one gram of water by one degree Celsius, and is equal to 4.186 joules (J).  Every material has a unique specific heat that depends on its own chemical composition.  Water has a particularly high specific heat when compared to other liquids for life.  For example, the specific heat of ammonia (NH3) is 0.470 J/g*oC.  If our oceans were made of ammonia instead of water, it would take significantly less energy to change the temperature of the ocean.  This means that over the course of a normal year, there would be global temperature changes that could result in widespread flooding and glaciation on an annual basis!  The specific heat of water helps Earth maintain its fairly stable climate.

 

On a more elemental level, hydrogen bonding plays a vital role in the ability for life to reproduce and evolve.  The DNA in our cells is double stranded – there are two strings of nucleotides that are joined primarily by hydrogen bonds between the strands.  We mentioned earlier that hydrogen bonds, when considered as individual bonds, are very weak.  However, when we consider hundreds of hydrogen bonds together, a relatively strong and stable structure can result.  Nucleotides from one strand make two or three hydrogen bonds with complimentary nucleotides of the other strand, depending on the nucleotide pair.  The two strands together make up the helical DNA molecules that transfer genetic information from one generation to the next.  Without these molecules, life would have a very difficult time perpetuating itself on the planet. 

 

So far we have examined the nature of hydrogen bonding and how it contributes to the importance of water for life.  Water has other characteristics that also make it important to life on Earth.  First, water acts as an excellent solvent, dissolving a wide variety of materials.  As a solvent, water assists in transporting molecules within the cell.  In addition, water is what causes proteins to take on their unique three-dimensional shapes, allowing them to catalyze very specific chemical reactions within cells.  Finally, water can act as a shield, if enough of it is present.  For example, during the history of early Earth there was not a protective layer in the atmosphere that blocked harmful ultraviolet (UV) radiation.  Some astrobiologists contend that the first forms of life arose in the ocean, where depths of water could absorb the UV and protect budding life forms.   

 

Although astrobiologists are not limiting the existence of life to liquid water, due to the versatility of water it is likely that life on other planets might also utilize water as its “life liquid”.  Consequently, many of the exploratory missions inside and outside of our solar system focus on the discovery of past or present water on planetary bodies as an indicator for life.

 

 

 

 

 

 

This is a figure from the lab manual that Ed, Tim, and I published.  It is copyrighted, but I made it in Word so I know it is easy to recreate… and you could make something similar that is a LOT prettier!!

 

 

 

 

 

 

 

 

 

 

 

 


This image below serves as an example of how water molecules arrange themselves in a crystal lattice when it is frozen as ice.

 

Saved as:

ImportanceOfWaterForLife_image1.jpg

 

And here’s a chinsy version of the same thing from our lab manual… again, done with Word.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 


Cool animation for covalent bonding can be found at:

http://207.10.97.102/chemzone/lessons/03bonding/mleebonding/covalent_bonds.htm

 

As well as for ionic bonding:

http://207.10.97.102/chemzone/lessons/03bonding/mleebonding/ionic_bonds.htm