EXPERIMENT11
IntroductionAcidity and Basicity
Acids
Any solution that releases hydrogen ions when added to water and has a pH of less than 7.0 pH – it measures the acidity of a liquid by
measuring the concentration of hydrogen ions.
PROPERTIES OF ACIDS : Sour taste
Litmus paper blue turns red
Reactions with metal oxide and hydroxide
Neutralize bas forming water
Bases
Substances which combines with acid and also known as Alkaline substance.
Compound that furnishes the hydroxide ions.
Compound that gives or donate hydroxyl ions in water or other substances.
Hydroxyl – unit that composed of one or more atom of hydrogen and one of oxygen.
PROPERTIES OF BASES : Slippery
Soapy feeling and a biting , bitter taste
Red litmus blue , turn methyl organic from red to yellow
Turn phenolphthalein from colorless to red
ProceduresAcidity and Basicity
Using the pH Paper
pH paper in test solution• Dip• For 10
seconds
Color Chart• Match the
color obtained in the test sol’n
Record Repeat• With other
test solutions
Using the pH Meter
Buffer Solution• Immerse the
electrode• After that,
rinse the electrode with distilled water
• Wipe with tissue
First Solution• Dip the
electrode• Get the pH
reading• Record
Distilled Water• Rinse the
electrode• Wipe it with
tissue paper
Repeat• With other
test solutions Keep the electrode
immersed in distilled
water when not inuse.
Samples
Chemical Formula
Structural Formula
Functional Group
Acetic Acid CH3COOH Carboxyclic Acid
Monochloro- acetic acid
ClCH2COOHCarboxyclic Acid, Alkyl
Halides
Acetone (CH3)2CO Ketone
Acetamide CH3CONH2 Amides
Chemical Formula
Structural Formula
Functional Group
Glysine NH2CH2COOHAmines,
Carboxyxlic Acid
Lysine C6H14N2O2
Amines, Carboxyxlic
Acid
Isopropyl Alcohol
(CH3)2CHOH Alcohol
PhenolC6H5OH Alcohol,
Aromatic Compound
Chemical Formula Structural Formula
Commercial Vinegar
C2H4O2
Calamansi Juice
C6H8O7
Spoiled Milk C3H6O3
Arrhenius ConceptAcidity and Basicity
Svante Arrhenius, a Swedish chemist who received a Nobel prize in 1903 for his work on electrolytes, focused on what ions were formed when acids and bases dissolved in water.
One of the properties that acids and bases have in common is that they are electrolytes--they form ions when they dissolve in water.
He came up with the concept or idea that acids dissociated in water to give hydrogen ions (H+) and that bases dissociated in water to give hydroxide ions. (OH-)
Examples:HCl H+ + Cl-
An acid, like HCl, is something that dissociates in water to give hydrogen ion.
NaOH H+ + Cl-
A base, like NaOH, is something that dissociates in water to give hydroxide ion.
Arrhenius focused on the idea that acids and bases split into ions when they dissolved in water.
In a sense, the Arrhenius concept focuses on what the chemical contains or what is there in solution.
Brønsted-Lowry ConceptAcidity and Basicity
With the Brønsted-Lowry concept we usually refer to a hydrogen ion as a proton.
That is because a proton is all that is left when a hydrogen atom loses an electron to become an ion.
Brønsted Acids Proton (H+) Donor. When an acid reacts, the proton
is transferred from one chemical to another.
The chemical which accepts the proton is a base.
H+ Cl- +
+ Cl-
I DONATE!!
Note that in order for an acid to act like an acid, there needs to be something for it to react with. There needs to be something to take the proton.
There needs to be a base.
Brønsted Bases Proton Acceptor. Opposite of acids. Bases are basic because they take or
accept protons. Hydroxide (OH-) ion, for example can
accept a proton to form water.
Brønsted and Lowry realized that not all bases had to have a hydroxide ion. As long as something can accept a proton it is a base.
So anything, hydroxide or not, that can accept a proton is a base under the Brønsted-Lowry definition.
The water molecules that accept protons when HCl dissolves in water are acting as bases.
H+ Cl- +
+ Cl-
I DONATE!
I ACCEP
T!
Some additional examples of Brønsted-Lowry bases are shown accepting protons in these equations. These examples do not show the acids which are providing the protons.
Ammonia can accept or react with hydrogen ion to give ammonium ion (NH4
+)
NH3 + H+ NH4+
Carbonate ion (CO32- )can accept a
hydrogen ion, or accept a proton, to become bicarbonate ion (HCO3
-).
Also, water molecules, as mentioned before, can act as a base by accepting protons.
CO32- + H+ HCO3
-
H2O + H+ H3O+
Hydroxide, ammonia, carbonate and water are all
Brønsted-Lowry bases.
When a Brønsted-Lowry acid donates a proton, it forms the conjugate base of that acid.
When a base accepts a proton, it forms the conjugate acid of that base.
Conjugate base and acid are produced as products.
The formulas of a conjugate acid-base pair differ by one proton (H+)
Consider what happens when HCl(g) is bubbled through water, as shown by this equation:
HCl(g) + H2O(l) → Cl-(aq) + H3O+
(aq)
Conjugate acid-base pair
• The conjugate base of HCl is Cl-
• The conjugate acid of Cl- is HCl
Conjugate acid-base pair
• The conjugate base of H2O is H3O+
• The conjugate acid of H3O+ is H2O
To write the conjugate base of an acid,remove one proton from the acid formula:
Note that, by removing H+, the conjugate base becomes more negative than the acid by one minus charge.
H2O OH- (Conjugate base)
HNO3 NO3- (Conjugate base)
To write the conjugate acid of a base, add one proton to the formula of the base:
In each case the conjugate acid becomes more positive than the base by a +1 charge due to the addition of H+.
SO42- HSO4
- (Conjugate acid)
C2H3O2- HC2H3O2
(Conjugate acid)
Lewis ConceptAcidity and Basicity
The Lewis Concept as an Extension of the Brønsted Concept
An acid is an electron-pair acceptor.
A base is an electron-pair donor. An acid-base reaction is the
sharing of an electron pair with an acid by a base.
These three simple definitions constitute the heart of what is now known as the Lewis concept of acids and bases.
Experimentally and conceptually, they are an extension of the Brønsted definitions.
The Fundamental Lewis Acid-Base Reaction
the formation of a coordinate covalent bond between an acid and a base. The base is the electron-pair donor, the acid the acceptor.
The process is called neutralization, or simply coordination.
The product is a coordinated compound, coordinated complex, or adduct, made up of an acid portion and a base portion.
A typical, and oft-cited example is the reaction of the acid boron trifluoride with the base dimethyl ether to form the complex or adduct BF3CH3OCH3:
The coordinate molecule may be thought of as being made up of the acid portion BF3 and the base portion CH3OCH3
Classification of Lewis AcidsI. Simple Cations. Theoretically all simple cations are potential Lewis Acids, although their strength as acids varies within wide limits. In general, we can expect the acid strength or
coordinating ability of cations to increase with: an increase in positive charge on the ion, an increase in nuclear charge for atoms in any
horizontal period, a decrease in ionic radius, and a decrease in the number of shielding electron
shells.
This means that Lewis acidity of simple cations tends to increase for the elements from left to right and from bottom to top in the periodic table. (Periodic Trends)
II. Compounds Whose Central Atoms Has an Incomplete Octet. Among the most important Lewis Acids are compounds whose central atom has less than a full octet of electrons.
III. Compounds in Which the Octet of the Central Atom Can be Expanded.
Although carbon and silicon belong to the same family of elements, silicon tetrafluoride and silicon tetrachloride are tremendously more reactive than their carbon analogues, carbon tetrafluoride and carbon tetrachloride.
The explanation is straightforward- the silicon, with its vacant d orbitals, can act as a Lewis acid by expanding its octet.
This is illustrated by the reaction of silicon tetrafluoride with fluoride ion to form fluorosilicate ion:
With no available d orbitals, carbon cannot do this, in keeping with the fact that the elements in the first period of eight in the periodic table can accommodate no more than eight electrons in their valence shell.
Actually the silicon halides typify a large group of halides which, with vacant d orbitals, can expand their octets. Some examples are:
These halides tend to form adducts with halide ions and with organic bases such as ethers (R-O-R).
Halides of this type are vigourously hydrolyzed to form an oxy-acid (or oxide) of the central atom and the appropriate hydrogen halide. This reaction depends upon the ability of the halides to act as Lewis acids.
The first step in the removal of each halogen atom is undoubtedly the acid-base coordination of the acid halide with the base water. This is followed by elimination of the hydrogen halide from the adduct.
For the removal of the first chlorine in the hydrolysis of phosphorus trichloride, we believe the pathway or mechanism is
IV. Compounds Having Multiple-bonded Acid Centres.
There are many compounds, particularly organic, in which a multiple-bonded atom can accept a share in an electron pair with a synchronous shift in a pair of electrons of the multiple bond.
By a slight extension of the Lewis concept, we can classify such compounds as Lewis acids. Although the atom involved does not, in a strict sense, have an unfilled orbital nevertheless an orbital is made available as the incoming base forces the intramolecular electron pair shift.
A familiar example is carbon dioxide. Consider its neutralization by hydroxide ion to hydrogen carbonate in:
V. Elements with an Electron Sextet. To the extent that oxygen and sulfur
atoms participate directly in chemical reactions, they may be regarded as Lewis acids. On this basis, the oxidation with sulfur of sulfite to thiosulfate and of sulfide to polysulfide ion can be classified as acid-base reactions:
Classification of Lewis Base They are species that contain atoms with
lone pairs. Such atoms are called donor atoms. Thus H2O (donor atom, O), NH3 (donor atom, N), Cl-, and CH3CH2OH (donor atoms, O) are immediately recognized as Lewis bases.
Because the number of Lewis bases is almost unlimited, it is useful to categorize them based on the number of donatable nonbonding electron pairs that they contain.
monodentate
a Lewis base which can form only one bond to a Lewis acid
name means "one
tooth"
chelating / polydentate
has 2 or moredonor atoms spaced so that they can attach to the same Lewis acid. Normally the donor atoms must be "spaced"by 2 or 3 intervening atoms.
non-linear, often with 2 or 3 atoms separating the donor
atoms
Discussing the ResultsAcidity and Basicity
pH Paper Reading
pH Meter Reading
Acetic Acid 3 3.34
Monochloroacetic acid
1 1.96
• Acetic Acid has a higher pH reading than Monochloroacetic acid.
• But Monochloroacetic acid is more acidic than Acetic acid.
Why is that so??
Acid Strength is influenced by Inductive Effects
Carboxyclic acids are the strongest acids among compounds that contains only C, H and O.
Acid Strength is influenced by Inductive Effects Acetic acid provides the reference mark for
a typical carboxyclic acid and it has pKa value of 4.75. The pKa value of monochloroacetic acid is less than that for acetic acid.
Acid Strength is influenced by Inductive Effects The variation in acidity among
structurally similar compounds like these can be explained in the electronegativity values of the substituents.
These electronegativity difference manifest themselves by donating or withdrawing electrons through the bonds between atoms, an influence known as inductive effect.
Acid Strength is influenced by Inductive Effects In the comparison of acetic and
monochloroacetic acids, the argument goes like this:
“Chlorine, being electronegative, renders the carbon atom adjacent to the carbonyl
group partially positive.”
Acid Strength is influenced by Inductive Effects In turn, this withdraws electron density
from the carbonyl carbon atom and the oxygen atom, bearing the proton.
The effect is to weaken the O-H bond, making the proton more acidic.
Acid Strength is influenced by Inductive Effects Thus this inductive effect of the chlorine
atom makes monochloroacetic acid more acidic than acetic acid.
In general, an electron-withdrawing substituent near the COOH group increases the acidity of acetic acid.
Acid Strength is influenced by Inductive Effects The more electronegative a substituent,
the stronger the acid. Conversely, an electron-donating
substituent makes the acid less acidic than acetic acid.
Alkyl groups are the most common substituents that donate electrons.
pH Paper Reading
pH Meter Reading
Acetamide 9 8.19
Acetone 5 6.46
• Acetamide has a higher pH reading than Acetone.
• Applying our basic information, pH less than 7 is acidic while pH more than 7 is basic. To conlude, Acetamide is a base while Acetone is an acid.
But what makes them a base and an acid?
Amphoteric character (Amides) The presence of a lone pair of electrons
on Nitrogen atom should make acid amide basic in character. But actually, they are very feeble bases.
This is because the lone pairs of electrons on the nitrogen atom is involved in Resonance with carbonyl group and is therefore not available for protonation.
Amphoteric character (Amides)
However, under suitable conditions, acid amide can show basic or acidic character.
Amphoteric character (Amides)
The presence of lone pair of electrons on the nitrogen atom is resonating structure (I) makes it feebly basic.
Acid amide, therefor, act as a base.
Amphoteric character (Amides) Thus acetamide (a base) reacts with
hydrochloric acid (an acid) to form a salt.
Amphoteric character (Amides) Likewise, the positive charge on the
nitrogen atom shown in resonating structure (II)
Implies easy release of proton.
Amphoteric character (Amides) Acid amide, therefore, acts as an acid.
For example, acetamide behaving as an acid reacts with bases Na or HgO to form corresponding salt.
Acidity of Ketones Ketones are far more acidic (pKa ≈ 20)
than a regular alkane (pKa ≈ 50). This difference reflects resonance
stabilization of the enolate ion that is formed through dissociation.
The relative acidity of the α-hydrogen is important in the enolization reactions of ketones and other carbonyl compounds.
The acidity of the α-hydrogen also allows ketones and other carbonyl compounds to undergo nucleophilic reactions at that position, with either stoichiometric and catalytic base.
pH Paper Reading
pH Meter Reading
Glycine 5 6.15
Lysine 6 6.17
• Glycine has a lower pH reading than Lysine.
• But Glycine is more acidic than Lysine.
Why is that so??
Chemical Nature of the Amino Acids The amino acids found in proteins have
the following generalized structure:
Chemical Nature of the Amino Acids All peptides and polypeptides are
polymers of α-amino acids. Several other amino acids are found in
the body free or in combined states (i.e. not associated with peptides or proteins).
Acid-Base Properties of the Amino Acids The α-COOH and α-NH2 groups in amino
acids are capable of ionizing (as are the acidic and basic R-groups of the amino acids). As a result of their ionizability the following ionic equilibrium reactions may be written:
R-COOH <——> R-COO– + H+ R-NH3
+ <——> R-NH2 + H+
The equilibrium reactions, as written, demonstrate that amino acids contain at
least two weakly acidic groups.
However, the carboxyl group is a far stronger acid than the amino group. At physiological pH (around 7.4) the carboxyl group will be unprotonated and the amino group will be protonated.
An amino acid with no ionizable R-group would be electrically neutral at this pH. This species is termed a zwitterion.
pH Paper Reading
pH Meter Reading
Isopropyl Alcohol 4 3.40
Phenol 3 3.33
• Isopropyl Alcohol has a higher pH reading than Phenol.
• But Phenol is more acidic than Isopropyl Alcohol.
Why is that so??
Relative Acidities of Alcohols and Phenols The polar O-H bond of alcohols
makes them weak acids. By the Bronsted-Lowry definition, acids are hydrogen ion donors and bases are hydrogen ion acceptors in chemical reactions.
Strong acids are 100% ionized in water and weak acids are only partially ionized.
Weak acids establish an equilibrium in water between their ionized and unionized forms.
Phenols are one million to one billion times more acidic than alcohols and this is the characteristic property that distinguishes them. Phenols will react with the base sodium hydroxide but alcohols will not.
The acidity of phenols is explained by resonance stabilization of the phenoxide ion; the negative charge is dispersed throughout the benzene ring as opposed to being concentrated on the oxygen as it is in the alkoxide ion.
Electron-withdrawing groups on the benzene ring increase the acidity of phenols.
Chemical components that makes it Acidic.
Yes, spoiled milk is an acid. The lactic acid makes the milk acidic,
milk is said to be sour when it is at a pH level of 4.3-4.5 (acidic)
Spoiled Milk Lactic Acid C3H6O3
Chemical components that makes it Acidic.
Calamansi juice is an acid. It has mild sour taste. Sour taste is a characteristic of an acidic
property.
Calamansi JuiceCitric Acid C6H8O7
Ascorbic Acid C6H8O6
Chemical components that makes it Acidic.
Acetic acid is the source of the acidity in vinegar.
Acetic acid (ethanoic acid) is an organic acid (carboxylic acid) and is classified as a weak acid.
Commercial Vinegar
Acetic Acid (Ethanoic
acid)CH3COOH
Intramolecular H-Bond on acidity The effects of intramolecular hydrogen
bonding on acidity can be see not just on O-H and N-H, where acidity is greatly reduced, but also on certain C-H groups, which in some cases become the primary source of acidity.
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