Acidity, basicity, and pKa

TAIBAH UNIVERSITY

COLLEGE OF SCIENCE

CHEMISTRY DEPARTMENT

Presented by:

Zubaydah Abdullah

* Acids and bases are obviously important

because many organic and biological reactions are catalysed by acids or bases.

!* pKa tells us how acidic (or not) a given hydrogen atom in a compound is

Types of Acids and Bases In the 1800s chemical concepts were based on the

reactions of aqueous solutions.

Svante Arrhenius developed a concept of acids and

bases relevant to reactions in H2O.

Arrhenius acid produces hydrogen ions in water.

Arrhenius base produce hydroxide ions in water.

Strong acids and bases are 100% dissociated.

A broader ,more modern concept of acids and bases was developed later.

The idea that acids are solutions containing a lot of H+ and bases are solutions containing a lot of OH- is not very useful in organic chemistry.

!

Bronsted-Lowry acid- can donate a proton.

Bronsted Lowry base can accept a proton.

Conjugate acid-base pairs.

A Brnsted-Lowry acid-base reaction results in the transfer of aproton from an acid to a base.

In an acid-base reaction, one bond is broken, and another one is formed

There is an inverse relationship between the strength of an acid and

the strength of its conjugate base: the stronger the acid, the weaker

its conjugate base

Conjugate acid- compound formed when an base gains a hydrogen ion.

Conjugate base compound formed when an acid loses a hydrogen ion.

Summary of

AcidBase Definitions

!

What is The pH Scale?

The pH scale is only a measure of how acidic or basic a solution is.

The pH scale is the concentration of hydrogen ions in a given substance.

[ ]+= HpH log

The pH scale depends on the concentration of acid

acidic solutions all have a pH of less than 7the lower the pH the more acidic the solution.

alkaline solutions all have pHs greater than 7the higher the pH, the more basic the solution.

The pH 7 is neither acidic nor alkaline but neutral.

Aqueous any strong acid has a lower pH than an equal concentration of aqueous any weak acid; because it is more fully dissociated and thereby produces more hydronium ions.

The pH scale also depends on the acid in question

For hydrochloric acid, the equilibrium lies well over to the right: in effect, HCl is completely dissociated.

Acetic acid is not fully dissociatedthe solution contains both acetic acid and acetate ions.

If a strong acid is added to water, the water acts as a base

and is protonated by the acid to become

!!!If we added a strong base to water, the base would deprotonate the water to give hydroxide ion, , and here the water would be acting as an acid.

!!!Such compounds that can act as either an acid or a base

are called amphoteric.

can behave as an acid or as a baseWater

The amino are amphoteric. Unlike water, these

compounds have separate acidic and basic groups built

into the same molecule.

!When amino acids are dissolved in water, the acidic end protonates the basic end to give a species with both a positive and a negative charge on it.

A neutral species that contains both a positive and a negative charge is called a zwitterion.

!

Pure water at 25 C has a pH of 7.00. This means that the concentration of hydronium ions in water must be

!

Hydronium ions in pure water can arise only from the self-dissociation or autoprotolysis of water.

!

! In this reaction, one molecule of water is acting as a base,

receiving a proton from the other, which in turn is acting as an acid by donating a proton

The ionization of water

Weak acid: a substance that dissociates only partially in water to produce H3O+ ions .

acetic acid, for example, is a weak acid; in water, acetic acid is incompletely ionized in aqueous solution.

!!

!! Weak base: a substance that dissociates only partially in water to

produce OH- ions.

ammonia, for example, is a weak base.

Acid and Base Strength

Acid strength is the tendency of an acid to donate a proton.

The more readily a compound donates a proton, the stronger an

acid it is.

When a Brnsted-Lowry acid HA is dissolved in water, an

acid-base reaction occurs, the position of equilibrium is measured by the equilibrium constant for this reaction Keq.

Stronger acids have larger Keq.

Acid and Base Strength

The concentration of water as a solvent does not change significantly when it is protonated.

with dilute solutions of acids wherever the equilibrium may be and a new equilibrium constant, Ka, is defined and called the acidity constant .

It is generally more convenient when describing acid

strength to use pKa values than K values. Like pH,

this is also expressed in a logarithmic form, pKa.

Because of the minus sign in this definition, the lower the pKa, the larger the equilibrium constant, Ka, is and hence the stronger the acid.

The pKa of the acid is the pH where it is exactly half dissociated. At pHs above the pKa, the acid HA exists as

in water; at pHs below the pKa, it exists as undissociated HA.

we can increase the solubility of a neutral acid in water by increasing the proportion of its conjugate base present. All we need to do is raise the pH.

organic bases such as amines can be dissolved by lowering the pH

In a mixture of two acids or two bases :

The ratio of Ka values gives us an indication of the equilibrium constant for the reaction between a base and an acid

The difference in pKas gives us the log of the equilibrium constant

The stronger the acid HA, the weaker its conjugate base, A

The stronger the base A, the weaker its conjugate acid AH

Acid and conjugate base strength

An acids pKa depends on the stability of its conjugate base.

The LOWER the pka the more acidic it is

The HIGHER the pka the

more basic it is.

Remember

The deference between Ka and pKa

- The strongest base in aqueous solution is OH and

the strongest acid in aqueous solution is H3O+.

Addition of stronger bases than OH just gives more OH by the deprotonation of water.

Addition of stronger acids than H3O+ just gives more H3O+ by protonation of water

The pH of pure water at 25C is 7.00 (not the pKa)

The pKa of H2O is 15.74

The pKa of H3O+ is 1.74

Remember

the more electronegative

the element

the more stable the conjugate

basethe stronger

the acid

Acid Strength

Bond strength AH. Clearly, the easier it is to break this bond, the stronger the acid.

!

The solvent. The better the solvent is at stabilizing the ions formed, the easier it is for the reaction to occur.

Factors affect the strength of an acid1) Electronegativity Electronegativity of the atom bearing the negative charge; within a

period.

The greater the electronegativity of the atom bearing the negative

charge, the more polarized the bond to H, H becomes more positive and the bond is easier to break; the greater the acidity of the acid HA, the less willing it is to share those electrons with a proton, so the weaker the base.

Electronegativity C < N < O < F

Stability

!Acidity

Electronegativity Increase

Acidity Increase

Basicity Increase

2) Size of the atomDown a column of the periodic table, the acidity of HA

increases as the size of A increases.

Acidity & Size Increase

Acidity H-F < H-Cl < H-Br < H-I

!!Stability

I-Br-Cl-F-

As size increases, Weak AH bonds is easier to break; is make stronger acids.

A larger size also stabilizes the anion

Factors affect the strength of an acid

Increasing Basicity

3) Resonance; delocalization of the negative charge the charge can be more delocalized, and this makes the anion more

stable. the greater the resonance stabilization of the anion, the more acidic

the compound.

If the negative charge can be delocalized on to more electronegative atoms such as oxygen or nitrogen, the conjugate base will be stabilized and hence the acid will be stronger, and it is going to be less basic than one with a more concentrated, localized charge

Factors affect the strength of an acid

4) The Inductive Effect

Any group that withdraws electrons will help to stabilize the conjugate base and therefore increase the strength of the acid.

An inductive effect is the pull of electron density through

bonds Caused by electronegativity differences in atoms,

Which cause polarization of the bond.

The reason for the increased acidity is that the electronegative chlorine atoms stabilize the negatively charged conjugate base.

Factors affect the strength of an acid

1. More electronegative atoms stabilize regions of high electron density by an electron withdrawing inductive effect.

2. The more electronegative the atom and the closer it is to the site of the negative charge, the greater the effect.

!!!!3. The acidity of HA increases with the presence of electron withdrawing groups in A

Factors affect the strength of an acid4) The Inductive Effect

Electron-donating groups decrease acidity The most common electron-donating groups encountered in

organic chemistry are the alkyl groups. These are weakly electron-releasing.

Factors affect the strength of an acid4) The Inductive Effect

Strong

Moderate

Weak

Electron-withdrawing

Electron-donating

Remember

5) HybridizationThe higher the percent of s-character of the hybrid orbital,

the closer the lone pair is held to the nucleus, and the more

stable the conjugate base.

Increasing Acidity & Stability

Factors affect the strength of an acid

Increasing Basicity

This makes an sp hybridized carbon less electron-donating

than an sp2 one, which in turn is less electron-donating than an sp3 carbon.

The more s-character an

orbital

The more it holds on to

the electrons in it

Factors affect the strength of an acid5) Hybridization

If we can delocalize the negative charge of a conjugate anion, the anion is more stable and consequently the acid is stronger.

Aromaticity increase acidity; because the conjugate base of the most acidic compound has more resonance structure over which spread its electron density.

Factors affect the strength of an acid5) Hybridization

The more stabilized the conjugate base, A, the stronger is

the acid, HA. Ways to stabilize A include:

! Having the charge on an electronegative element

Delocalizing the negative charge over other carbon atoms, or even

better, over more electronegative atoms .

Spreading out the charge over electron-withdrawing groups by

the polarization of bonds (inductive)

Having the negative charge in an orbital with more s-character

Becoming aromatic.

Summary

Summary

Nitrogen acids

Amines are not acidic, amides are weakly acidic (about the same as alcohols), and imides are definitely acidic (about the same as phenols).

Basicity

! A base is a substance that can accept a proton by donating a pair

of electrons. Some organic bases and their relative strengths in proton-transfer

reactions

Any substituent that increases the electron density on the nitrogen therefore raises the energy of the lone pair thus making it more available for protonation and increasing the basicity of the amine.

Neutral nitrogen bases

Amidines are stronger bases than amides or amines

Oxygen bases in general are so much weaker than their nitrogen because oxygen is more electronegative and wants to keep hold of its electrons.

Neutral oxygen bases

Note: that acidity and basicity are just the reverse of each other. AND Therefore, both are affected by the same factors, just in opposite ways.

The End