Acidity, Basicity and pKa
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Transcript of Acidity, Basicity and pKa
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Acidity, basicity, and pKa
TAIBAH UNIVERSITY
COLLEGE OF SCIENCE
CHEMISTRY DEPARTMENT
Presented by:
Zubaydah Abdullah
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* 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
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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.
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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.
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Bronsted-Lowry acid- can donate a proton.
Bronsted Lowry base can accept a proton.
Conjugate acid-base pairs.
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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
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Conjugate acid- compound formed when an base gains a hydrogen ion.
Conjugate base compound formed when an acid loses a hydrogen ion.
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Summary of
AcidBase Definitions
!
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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
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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.
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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.
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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
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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.
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Pure water at 25 C has a pH of 7.00. This means that the concentration of hydronium ions in water must be
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Hydronium ions in pure water can arise only from the self-dissociation or autoprotolysis of water.
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! 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
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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.
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!! 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
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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
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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 .
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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.
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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
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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.
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The LOWER the pka the more acidic it is
The HIGHER the pka the
more basic it is.
Remember
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The deference between Ka and pKa
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- 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
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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.
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The solvent. The better the solvent is at stabilizing the ions formed, the easier it is for the reaction to occur.
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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
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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
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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
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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
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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
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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
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Strong
Moderate
Weak
Electron-withdrawing
Electron-donating
Remember
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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
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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
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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
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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
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Summary
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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).
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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
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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.
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The End