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What are Acids, Bases, and pH?

Updated on August 30, 2016

Acids, Bases, and pH

Below are some notes on acids, bases, and pH that can be used for an understanding or refresher that is helpful and useful for the sciences; specifically biology, chemistry, and biochemistry

What are Acids?

There are three definitions of acids (and bases) but here I will use the definition of the one that is emphasized at this point in my biochemistry textbook which is the Bronsted-Lowry definition.

Acid= a molecule that acts as a proton ([H+]) donor.

Bronsted-Lowry Acids and Bases

[H+] Donor
[H+] Acceptor

What are Bases?

Bases also have three definitions but here I will use the definition of that is emphasized at this point in my biochemistry textbook which is the Bronsted-Lowry definition.

Base= a molecule that acts as a proton ([H+]) acceptor.

Key Terms and Ideas

Acid Strength= the amount of [H+] released when a given amount of acid is dissolved in water.

Acid Dissociation Constant= a numerical value (Ka) given by an expression where HA (acid) is in equilibrium with H+ + A- (conjugate base), and where Ka=([H+][A-])/([HA]).

Ka and pKa

Ka= ([H+] [OH-]) / ([H2O]);(the concentration of hydrogen ions times the concentration of hydroxide ions divided by the concentration of water)

pKa= -log10Ka

More on Acids

In the expression for acid dissociation, the square brackets reflect molar concentration (concentration in moles per liter).

Each acid has a certain Ka at a given temperature.

Acids that are more dissolved have a larger Ka.

The greater the Ka, the stronger the acid.

pKa is a better measure of acid strength.

The lower the pKa, the stronger the acid.



H2O (l)=base

H3O+(aq)=Conjugate acid to H2O

A-(aq)=Conjugate base to HA

Acid-Base Reaction

HA(aq) + H2O(l) <--> H3O+(aq) + A-(aq)

What is pH?

pH describes the relationship between hydrogen ions ([H+]) and hydroxide ions ([OH-]) in aqueous solution.

pH connects the Ka of any weak acid with the pH of a solution that has both the acid and its conjugate base.

pH is useful when pH needs to be controlled for optimum reaction conditions.

pH Defined by Henderson-Hasselbach Equation

pH = pKa + log([A- ]/[HA])

*Where [HA] = concentration of acid, and [A- ] = concentration of the conjugate base.

*This equation is useful for predicting the properties of buffer solutions used to control the pH of reaction mixtures.

*When the pH value is equal to the pKa, the mixture will have equal concentrations of the weak acid and its conjugate base.

pH Defined by [H+]

pH = - log10[H+]

*Logarithm of base ten indicates a tenfold difference in [H+] per unit of pH.

pH and Solution Identity

In solutions:

Acidic= pH values lower than 7.

Basic= pH values higher than 7.

Neutral= pH value of 7.

Why Does pH Matter?

pH is important because most biological organisms, processes, enzymes, etc. need a specific range of pH values to function, survive, and thrive. Going outside of these ranges are detrimental to these organisms, processes, enzymes, etc. and can have a range of effects. These effects can fall anywhere from lowered function to death, depending on how far outside the pH range the pH is at and how long the pH has been at that outside value.

Kw Relationship

At 25oC in pure water:

Kw= [H+] [OH-]= (10-7) (10-7) = 10-14

*This relationship holds for any (neutral, acidic, basic, water, non-water) aqueous solution.

Ion Product Constant for Water

When given either the concentration of [H+] ions or [OH-] ions, Kw can be used to determine the concentration of the ion that is not given or known (either [H+] or [OH-]).

Source Information

The information used for this hub was taken from the following sources:

"Biochemistry" by Mary K. Campbell and Shawn O. Farrel; 7th edition.

My biochemistry lectures at school.

Knowledge and notes taken from previous courses in chemistry and biology. Some information for Kw was provided by my friend.


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