Chemistry

What is a Base



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Defining bases and skinning cats have one thing in common – they each can be done in more than one way.  Bases are typically defined either by the use of a relative scale, or by their chemical behavior.

The most familiar way for most people to recognize a base is to see how it affects the pH of water.  The pH scale is numerical, with values greater than seven corresponding to a basic pH.  If a chemical is dissolved in water (which has a neutral pH of seven), and the pH is more than seven, the chemical that was dissolved is called a base.  The pH may be measured using a pH meter (which gives a numerical value) or by using a chemical indicator.  Chemical indicators will change color at certain pH values.  Litmus paper is a common example, and turns blue when exposed to a basic solution.  Phenolphthalein is another popular indicator, turning a pink to magenta color in basic solutions. 

The use of the pH scale to define bases has a limited usefulness.  It applies mostly to solutions of chemicals in water.  While this encompasses most of daily experiences, it does not describe fully the nature of the chemical world, which includes solids, gases and solvents other than water.

To more completely define bases, chemists examine their chemical behavior.  Several definitions have emerged over time, with three being well-known, and named after the chemist who proposed the definition.

An Arrhenius base is a chemical that has an –OH group (a hydroxyl group) that can dissociate (separate) from the chemical when dissolved in water, forming a hydroxide ion.  This is an easily grasped concept, and applies to well-known bases like sodium hydroxide (NaOH) or potassium hydroxide (KOH).  This definition agrees well with the pH scale, but is limited to hydroxide-containing molecules only.  It cannot address other molecules that turn water basic, such as ammonia.

A Bronstead-Lowry base is a chemical that can act as a “proton acceptor”.  In chemistry parlance, “proton” in this case refers to the nucleus of a hydrogen atom.  (The electron has been stripped away to form the hydrogen ion – a single, positively charged proton.)  For a molecule or ion to be a proton acceptor (a base), it must have an electronegative atom with a lone pair of electrons available to form a bond with the proton.  The hydroxide ion fits this description well – the oxygen atom in the negatively charged OH ion has three lone pairs of electrons.  This definition also captures ammonia, in which the nitrogen atom has a lone pair of electrons.  The addition of a proton to these substances results in the formation of water (H2O) and the ammonium ion (NH4+), respectively.  As you can see, this definition is more widely applicable than the Arrhenius definition.  It also has its limitations, however.  Outside of aqueous (water-based) environments, acid-base chemistry does not always revolve around proton exchange. 

Lewis bases are known as electron donors.  As in the Bronstead-Lowry definition, these bases have an electronegative atom with a lone pair of electrons.  The difference is that the lone pair may be used to form a bond to any electron-poor atom, not just the hydrogen ion.  As a result, any Bronstead-Lowry base is a Lewis base, but so are chemicals that form complexes with metal ions, and any other chemical species that is ready to share a pair of its electrons.  The Lewis definition is the broadest, most universally applicable definition.  Its very breadth is perhaps its only limitation, as it is often more convenient to refer to a smaller set of chemicals when dealing with bases. 

As with many things, context defines which definition will be used.  Gardeners and swimmers will likely be satisfied with a pH reading.  Young science students are typically introduced to acid-base chemistry using pH, indicators, and the Arrhenius definition.  A full understanding of aqueous chemistry typically requires the Bronstead-Lowry definition, while the Lewis definition is necessary for full mastery of the subject.

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