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Chemical principles: Acids, bases and pH

Introduction

 

When two water molecules are hydrogen bonded, sometimes the hydrogen involved will shift from one water molecule to the other. The result is a positively charged hydronium ion (H3O+) and a negatively charged hydroxide ion (OH-). Sometimes for simplicity it is described as the formation of an H+ and an OH- ion from a single water molecule, although in real life a free H+ ion would always attach itself to an available water molecule rather than float freely. In pure water this only happens occasionally by chance and nearly all the water molecules have no charge (H20).

 

In a solution containing other substances (acids and bases) the balance can shift in one direction or the other causing an increase or decrease in the concentration of hydronium ions. This is the pH of the solution. It is extremely important in biochemistry because hydronium and hydroxide ions are both very reactive. In particular they can interfere with bonding, changing the shapes of complex molecules like proteins and affecting how they interact with each other. Some biological processes need a lot of these ions available in order to complete reactions, for example hydrolases in the lysosomes need an acid environment to breakdown cell debris.

 



What is pH?

 

Mathematically the definition of pH is: pH = -log10[H3O+]

 

For the rest of the article we will refer to the hydronium ions as H+ for clarity. So when the concentration of H+ is high, it makes the pH low. A solution with a low pH (more H+ ions than OH- ions) is called acidic and one with high pH (more OH- ions than H+  ions) is called basic or alkaline, while pure water is neutral with a pH of 7.

 

There is a mathematical constant for this known as the ionic product of water, which can be shown as: Kw = [H+][OH-]

 

Useful pH equations



The pH Scale

 

Kw is a constant because of the tendency of hydroxide and hydronium ions to combine into water. Therefore as the concentration of H+ goes up, the concentration of OH- correspondingly goes down. This means that when these concentrations are multiplied together the product is always the same. If you know the value of Kw you can work out the concentration of H+ if you know the concentration of OH- (and vice versa) by dividing Kw by the known concentration.

 



 

The pH scale of 0-14 is a more convenient and accessible way of expressing the concentration of ions than in moles per litre. Each pH unit is a tenfold difference in concentration, so pH 2 is not twice as acidic as pH4 but a hundred times more acidic. Although pH is calculated using the equation above using H+ concentration, it also implies OH- concentration as the two values are linked.

 

Proton donors and acceptors



Strong and weak acids and bases

 

  • A strong acid or base will dissociate completely in water.
  • A weak acid or base dissociates reversibly and will reach a state of equilibrium between association and dissociation.

 

 

    The position of the equilibrium will be different for each acid; however, not all acids have the same “strength”. We are able to quantify an acids ability to dissociate (ie. where the position of the equilibrium will lie) using the acidity constant. If we use the above acid equilibrium then the equation for Ka is shown on the right.

     

    If the value for Ka >>1 for an acid then that acid is said to be a strong acid; the position of the equilibrium lies so that the ionic products are favoured. If there are more protons per volume (ie. a higher concentration of hydronium ions in solution) the pH of that solution will be quite low. Examples of these include HCl and H2SO4.



    However, if the value for Ka <<1 then the acid is called a weak acid, such as ethanoic acid, and the undissociated acid is favoured by the equilibrium. Therefore the pH will be slightly higher than those for strong acids. Because values for Ka can take values in a large range it is easier for us to use pKa which has no units and runs on a logarithmic scale. The relationship between Ka and pKa is analogous to the one between [H+] and pH:

     



     

    Whereas pH is seen as a bulk property of a solution, pKa is the probability that a hydrogen ion will leave a molecule. For a molecule like HCl there is only one hydrogen atom present to leave the molecule. However, when you look at more complex (particularly organic) molecules you will find there is more than one hydrogen atom which is able to dissociate from the molecule to contribute to the acidity of the solution. The different atoms surrounding these hydrogens (ie. their chemical environments) will determine how easily it will dissociate. The easier it is for a proton to leave the more acidic that hydrogen is and it takes a lower value for pKa.

     

    Useful facts

     

    • Acid Base Interactions: When acids and bases interact in solution they neutralize each other, or cancel each other out. For example where "A" is the acid and "B" is the base: HA (aq) + BOH (aq) → BA (aq) + H2O (l).
    • Buffers: Buffers are substances in a solution that make it resistant to pH change. The are used in living organisms where changes in pH can be extremely harmful and in biochemical experiments to ensure that the pH variable remains constant. Buffers work by accepting hydrogen ions at low pH and donating them at high pH.
    • Conjugate Acids and Bases: When an acid dissociates e.g. carbonic acid (H2CO3) it automatically produces a base, in this case a bicarbonate ion (HCO3-). The same is true when a base accepts a hydrogen ion. These are called conjugate acids and bases.

     

    References

     

    1. Bridge, G. (2012) Chemical Principles: Properties of water. Fastbleep.com Biochemistry Chapter.
    2. Campbell, N. and Reece, J. (2005) Biology 7th Edition. Pearson Education Inc.
    3. Berg, J., Tymoczko, J. and Stryer, L. (2002). Biochemistry. 5th Edition. New York, W.H. Freeman Publishing.
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