WEEK 3: ACIDS & BASES

Acids & Bases

• Acid = donates H + , which becomes H3 O + (hydronium ion)

• Base = accepts H+ , which reduces the concentration of H 3 O +

• (can also think of it as increasing the amount of OH- )

• HA = Acid (proton donor)

• A- = Base (proton acceptor)

• pH (potential of hydrogen)-> is a logarithmic measure of the

• [H +] in a solution.

If a compound can dissociate (dissolve/break into

smaller components) in water, then you can use its

chemical formula to identify if it is:

1. Acid (proton donor)

2. Base (proton acceptor)

• Organic acids and bases have specific acid and base

functional groups.

Two functional groups we will learn about can

donate protons in organic molecules: Organic acids

Organic bases

One functional group we will learn about is a

base in organic molecules

Acids and bases + pH

• pH is commonly expressed as –log[H + ]

• Pure water has [H + ]=10-7 and thus pH=7.

• Acids have high [H + ] pH is < 7

Bases have low [H + ] pH is > 7

• Acids produce H+ ions

  • increases H+ ion concentration in solution

• Bases bind to H+ ions

  •  So they “suck up” the H+ ions in solution, thereby reducing their concentration.

• pH affects solubility of many substances.

• pH affects structure and function of most proteins -

• including enzymes.

• Many cells and organisms (esp. plants and aquatic animals) can only survive in a specific pH environment.

Terms

• Ionised = has a charge due to loss or gain of H +

• Protonated = has the ‘proton’ (may be charged or uncharged)

– e.g. an acid that is protonated has not lost its proton yet, while a base that is protonated has already gained

• a proton

• Dissociated = not attached to the proton, may be charged or uncharged

• Acids can dissociate to give

– H + and a negative ion (A-) (anion)

– the anion (A-) is called the conjugate base

Conjugate acid-base pairs

• whenever an acid donates a proton, the acid changes into a base

• whenever a base accepts a proton, an acid is formed

• an acid and a base that differ only by the presence or absence of a proton are called a

conjugate acid-base pair

Ionisation and disassociation

Bases can ionise by accepting a proton

– they form a positively charged cation (BH +)

– the cation is the conjugate acid of the base

Acids + pH

Acids dissociate to give H +

– occurs more readily in basic/alkaline pHs

– in alkaline pH’s there is a shortage of H +

– this shortage encourages acids to dissociate/donate H+

Bases + pH

Bases can accept protons

– occurs more readily in acidic pHs

– in acidic pHs there is a lot of H +

– the large amount of protons enable bases to accept H+

pH

• pH is general

• measures amount of free H+ in solution

• only depends on H+

pKa

• • pKa is specific

  • # that describes the acidity of a particular molecule

  • the lower pKa, the stronger the acid and therefore its ability to donate its protons

• how will a molecule react in solution at a specific pH

• each molecule has its own pKa

pKa is important for organic acids

• Most acids are weak especially organic acids

• Weak acids do not completely ionise

pH vs pKa

• Lower pH = stronger acid

• Lower pKa = stronger acid

– but can have negative numbers or very big numbers

Relationship between pH and pKa

• The pH of the environment is exactly equal to the pKa of a drug molecule

– e.g. acetic acid has a pKa of 4.76

• At a pH of 4.76:

– exactly 50% of acetic acid is in the ionised form and 50% in the unionised form

For a weakly acidic drug:

• completely unionised at pH

• value more than 2 units below its pKa and completely ionised at pH values greater than 2 units above its pKa

• at a high pH, drug is completely ionised resulting in maximum solubility (in an aqueous environment)

• at low pH, less ionisation (more absorption expected in an acidic environment)

For weakly basic drug

• the opposite is true

• completely ionised at pH values more than 2 units below its pKa and completely unionised at

pH values greater than 2 units above its pKa

– absorption occurs in basic environment, intestine

• Both acidic and basic drugs are exactly 50% ionised when pH is equal to pka (pH = pKa)

pH>pKa

• Basic conditions

• Fewer H+ ions

• More A- (bases)

pH=pKa

• Not basic or acidic

• HA (acid) = A- (base

pH<pKa

• Acidic conditions

• Lots of H+

• Less A- (bases)

Using pKa

• pKa tells you when a drug will give/take a proton

at a certain pH

• Example

– pKa = 2 for ‘X’

– at pH = 2 à X is at equilibrium

– at pH = 1 (lots of protons)

(HX will dominate over X- )

– a t pH = 3

(X- will dominate over HX)

at any pH above 2:

‘X’ is a stable

negative ion (X- )

(it has dissociated

by donating its proton)

Therefore ‘X’ is a

strong acid

Henderson-Hasslebach equation

pH enhances or reduces the ionisation of drugs, thereby affecting solubility

• Ionisation can be predicted using the Henderson-Hasselbach equation

• This helps predict the absorption of drugs from different parts of the gastrointestinal tract

–acidic drugs: pH = pKa + Log[A-]/[HA]

–basic drug: pH = pKa + Log[base]/[salt]

where: [A-] = concentration of ionised form and [HA]= concentration of unionised form; [salt] =

conjugate form of the base

Ionisation and pKa matter

Because:

• Only the unionised form of a drug can cross biological

membranes

• providing the unionised form is lipophilic

• The ionised form is more water soluble

• required for drug administration and distribution in plasma

• pKa tells you what the pH has to be before a drug will give/take a proton

• The pH in the human body is not the same everywhere meaning that depending on where the drug is meant to go and what it does relies on the pKa in order to work

• Inflammation can affect how much the drug can get through

pH and solubility

For a solution of weakly acidic drug, a decrease in pH (addition of acid) increases the unionised form and precipitation occurs

• For a solution of weakly basic drug, an increase in pH (addition of base) increases the unionised form and precipitation occurs

• Precipitation is an example of chemical instability in liquid/solution dosage forms

Aromatic ring stabilises ions

Factors affecting the strength of an acid

•  Stability of conjugate base (anion), the stronger the

acid

• Any factor that stabilises the anion increases the acidity

of the group

– eg: resonance

• a way of describing bonding in chemistry. It describes how several equivalent Lewis structures contribute to one overall hybrid molecule

• Stable negative charge due to lower electron density on

the atom

Basicity

Lone pair available? -> base

– electron pair on the nitrogen is able to accept protons, but only a single pair = weak base

• Availability of electrons determines strength

– the more available the electrons, the more likely it is to accept a proton = stronger base

• So, an aromatic amino group is a much weaker base than an

aliphatic one (i.e. carbon atoms linked in an open chain)

Buffers

Definition: a solution that resists change in pH

– typically a mixture of the acid and base form of a chemical

– can be adjusted to a particular pH value

Blood: pH = 7.35-7.45

Too acidic? Increase respiration rate expelling CO 2 , driving reaction to

the left and reducing H + concentration.

Why use them?

– enzyme reactions and cell functions have optimum pH’s for performance

– important anytime the structure and/or activity of a biological material must be maintained

How buffers work

• Equilibrium between acid and base

• Example: acetate buffer

– CH 3 COOH « CH 3 COO - + H +

• If more H + is added to this solution, it simply shifts the equilibrium to the left, absorbing H + , so the [H + ] remains unchanged

• If H + is removed (e.g. by adding OH -) then the equilibrium shifts to the right, releasing H + to keep the pH constant

Equilibrium between acid and base

• Example: acetate buffer

– CH 3 COOH « CH 3 COO- + H +

• If more H + is added to this solution, it simply shifts the equilibrium to the left, absorbing H +, so the [H +] remains unchanged.

• If H + is removed (e.g. by adding OH-) then the equilibrium shifts to the right, releasing H + to keep the pH constant