FORM SIX CHEMISTRY STUDY NOTES ACIDS AND BASE

FORM SIX CHEMISTRY STUDY NOTES ACIDS AND BASE

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FORM SIX CHEMISTRY STUDY NOTES ACIDS AND BASE

Arrhenius concept of acids and bases .
What is an acid? (according to Arrhenius concept of acids and bases)

Arrhenius considered that an acid is a substance which when dissolved in water dissociate to produce H⁺ions as the only positively charged ions i.e.

e.g

-He considered a base to be a substance which produce hydroxyl ions when dissolved in water as the only negatively charge ions i.e

e.g.

The neutralization of acid with a base yields to a salt and water.

e.g.
+ +

Ionic equation

According to Arrhenius, neutralization reaction is all about formation of water.

Weakness of Arrhenius equation

i) This concept is limited to water. It refers to H⁺and OH^– ions derived from water. A true general concept of acid and base should be appropriate to other solvent like liquid NH₃ and alcohols.

ii) The concept does not provide the room for acids and bases which do not contain H⁺ions and HO ^– ions.

Bronsted – Lowry concept of acids and bases

-Bronsted and Lowry proposed a theory of acids and bases applicable to all solvents.

– They proposed that, an acid is any substance that can donate a proton to any other substance.

e.g

– A base is a substance that can accept a proton from any other substance.

E.g.

^{Base Acid}

– They are called a monoprotic acid an acid which donates only one proton.

e.g. HNO₃, HCl

– Diprotic acid can donate two protons.

E.g. H₂SO₄

-A polyprotic acid is an acid that can donate more than one proton.

E.g. H₂SO₄, H₃PO₃, H₂C₂O₄

– A polyprotic base is one which can accept more than one proton.

E.g.

– A monoprotic base is one which can accept only one proton.

E.g.,

NOTE: HCl and Cl^–are acid-base conjugate pair. Another example is HNO₃ and

Amphoteric (amphiprotic) acids and bases.

These behave as bronsted – Lowry acids or bases

Conjugate acid base pair

-For every acid, there is a corresponding (conjugate) base to accept a proton

E.g.

Acid base

(Proton donor) (Proton acceptor)

HA and A^– are conjugate pair i.e.

HA is a conjugate acid of A^– and A^– is a conjugate base of HA.

®In a solution, there must be a base to accept a proton

E.g.

Acid Base Acid Base

E.g. ^–

A₁, B₁, and A₂, B₂, are acid base conjugate.

NOTE:

From Bronsted–Lowry concept of acid and bases, the stronger the acid, the weaker it’s conjugate base and the stronger the base, the weaker its conjugate acid.

-CH₃COOH is a weak acid, but its conjugate base i.e. is a strong base.

– H₂O is a weak base, but is strong acid.

Advantage of Bronsted–Lowry concept over Arrhenius

® It can apply to any solvent not necessarily. Here the definition of bases is much wider.

Weakness:

Since the concept is based on proton transfer, it does not consider other compounds which do not contain hydrogen i.e AlCl₃, BF₃, SO₃.

In contrast to Arrhenius theory, acid and bases are no longer related to salts (by neutralization).

Question 1:

A. Define

i) Conjugate acid-bases pair.

ii) Conjugate base.

B. For the following pairs, write down the equation to show the conjugate acid-bases pair.

i) /

ii)

iii)

iv) /

Question 2:

a. Write the formula and give the name of the conjugate base for each of the following acids.

ii)

b. Write the name and formula of the conjugate acids for each of the following bases.

i) NH₃

ii) Br ^–

iii) HS^–

Question3:
In each of the following acidbase reaction. Identify the acid and the base on the left and their conjugate partners on the right.

ANSWERS:

^{Acid Base Acid Base}

HCN, CN^– is a conjugate pair and NH₃, NH₄⁺ is a conjugate pair

b)HSO^–₄, SO₄^-2, is a conjugate pair

Acid Base Base Acid

[Al(H₂O]³⁺, [Al(H₂O)₅ OH]²⁺is a conjugate pair.

OH^–and O is a conjugate pair.

Q: Conjugate acid-base pair is a pair which shows that for every proton lost by acid, there is a corresponding base to accept it.

a) Conjugate base.

ii)

iii)

LEWIS CONCEPT OF ACIDS AND BASES.

– Lewis proposed an even broader concept of acids and bases focusing on electron transfer rather than total transfer.

– According to Lewis an acid is a substance that can accept a pair of electrons. Therefore, an acid is an electron pair acceptor.

– A base is a substance that can donate a pair of electrons i.e base is an electrons pair donor.

NOTE
Bronsted–Lowry acids e.g. HCl, H₂SO₄, HNO₃, are not Lewis acid.
Thus, an acid-base reaction can occur when a base provide a pair of electrons to share with an acid resulting into coordinate compound or complex.

Therefore, ammonia chloride ions (AlCl₃) are Lewis base, while H⁺, BF₃ are Lewis acids.

NOTE: A Bronsted–Lowry base (like NH₃) reacts by donating electron pair to a proton. Therefore Bronsted–Lowry bases are also Lewis bases

Reason:

This is because upon donating a pair of electron, it would have accepted a proton. Therefore, Bronsted – Lowry bases are also Lewis bases.

i . F^– fluoride ionSO^2-

ii. Sulphate ion

iii. NH⁺₄ ammonium ion

iv. HBr Hydrogen bromid

v. H₂S Hydrogen sulphide.

Reason:

These cannot accept a lone pair of electrons hence the Lewis concept of acid is usually used in special cases.

IONIC EQUILIBRIUM OF ACIDS AND BASES.

Most acids and bases are weak i.e does not ionize fully when dissolve in water. Thus a part from water equilibrium, they also establish equilibrium.

e.g.

Ammonia which is a typical weak base ionizes as follows:

But the ionization of weak acid/ bases generally occurs to a greater extent than that of water.

STRENGTH OF WEAK ACIDS AND BASES.

The position of equilibrium of a reaction between the acid and water varies from one weak acid to another. The further to the left it lies, the weaker the acid is

The equilibrium constant is written as;

But H₂0 is constant at constant temperature.

Putting the constant on the same side

Where Ka = dissociation/ ionization constant of an acid.

for

Similarly for weak base, the position of equilibrium varies from base to base. The further to the left it lies, the weaker the base is

Where K_b = dissociation/ ionization constant of a base.

The K_aand K_b values are used to determine the strength of acids and bases i.e K_a and K_bvalues are quite small for a very weak acid/ base reflecting very title ionization of these acids/ base in solution.

Example:

The K_b value for C₆H₅NH₂ is 4.17 x 10^-10 ,NH₃ is 1.78 x 10^-5 Indicate which base is stronger than the other.

NH₃ is stronger than C₆H₅NH₂

The K_b value for C₆H₅NH₂ is smaller than C₆H₅NH₂NH₃

The strength of weak acids and bases can also be determined from its degree of dissociation (Ostwald’s dilution law)

Since K_a and K_b values are inconvenient to handle usually pK_a and pK_b are used.

For example;

The lower the value for the stronger the acid base respectively and vice versa

THE RELATIONSHIP BETWEEN pk_a AND pk_θ FOR A CONJUGATE ACID–BASE PAIR.

Consider the equilibrium.

The product of K_a and K_bgives.

Example
Formic acid (HCOOH) has a K_aof 1.78 x 10_-14 moles. Calculate the [H₃O⁺] and the pH of 0.1M. Solution of HCOOH.

Solution

Since K_a value is small, the expression

NOTE:
The approximation is done when [HA]O is greater than 100 K_a

But if initial concentration [HA]O is less than 100 K_a, then the exact expression formed must be solved.
Calculate [H₃O⁺] and pH in which has K_a value of moles/ dm³.

Solution:

Question set 1:
Formic acid (HCOOH) has K_a of (HCOOH) 1.8 x 10^-14 moles dm³. If you have0.001 m a solution of the acids. What is the pH of this solution, what is the concentration of HCOOH at equilibrium?

2. If the acid HA is 2% ionized in solution of concentration 0.01m, calculate

a) K_a

b) pK_a

3.Calculate the degree of ionization is 9.37 in a 0.1M aqueous solution

_{4. Calculate pH of a 0.052m acetic acid solution if Ka is}

5. For a 0.1M solution of benzoic acid, calculate.

i) Concentration of ions and molecules in solution

ii) The degree of ionization of the acid.

_{iii) pH of the solution}

6). A hypothetical weak base (MOH) has Kb of for the reaction.

_{Calculate the equilibrium Concentration of MOH, M}⁺_{and OH}^–_{in a solution MOH.}

The weak base methylamine has Kb of 5 x 10^{– 4}_{.It reacts with water according to the equation.}

Calculate the equilibrium concentration of OH^–_{in a solution of base. What are the of the solution?}

Hydroxyl amine has a Kb of What are the of the base ?

A 0.1m solution of chloroacetic acid (ClCH2COOH) has a of 1.95. Calculate Ka for the acid.

FORM SIX CHEMISTRY STUDY NOTES ACIDS AND BASE

IONIC PRODUCT OF WATER AND PH.

– Water auto–ionizes i.e transfers a proton from one water molecule to another producing H₃O and OH^–

–

Base Acid

–

But the concentration of H₂O is much larger than the two ions and is constant at constant temperature. Since the concentration of is constant it is made part of the constant.

Where = ionic product of water constant.

In pure water, Kw = 1 x 10^-14 at 25°c. Since every one H₃O⁺ ion formed also one OH^– ion is formed, thus the concentrations are equal.

Since the concentrations are equal, this implies that pure water is neutral.

Scale

– Is a scale which shows degree of acidity or alkalinity of a solution:

= 7

I.e pH of 7 is neutral point.

Is the negative logarithm of base 10 hydrogen ion.

Is the negative logarithm of base 10 of hydroxyl ion concentration.

For pure water

= 7 pOH= 7

NOTE:

In acidic solution, the concentration of H₃O⁺ is greater than [0H^–].

i.e

In basic solution, the concentration of

i.e [

From

, Introducing negative log on both sides

–

Variation of of pure water with temperature.

The formation of H₃O⁺ and OH^– ions from water is an endothermic process i.e the forward reaction absorbs heat.

-According to Le Chatelier`s principle, when you increase the temp, the forward reaction is favoured, thus concentration of H₃O⁺and OH^–ions will increase but in equal amounts. Thus, the pH will drop but the water will not be acidic the pH scale will also change. It won`t remain as 1 to 14 and the neutral point will also shift. The direct effect on increasing temperature is to increase K_w.

The table below shows the effect of temperature on K_w and each value of K_w a new pH must be calculated.





		6.1449

From the table, pH of water fall as temperature increases. This does not mean that water becomes more acidic at higher temperature. This solution is only acidic, If the concentration of is great than

Example:

1. The value of K_w at physiological temperature of a body is . What is the pH at neutral point of water at this temperature.

Solution:

2. At , is 15. Find at O°c, if the [H⁺] is

Solution:

STRONG ACIDS AND BASES.

When an acid is added to water as an aqueous solution of HCl in addition to self ionization of water, the acid also ionizes.

Due to common ion effect of hydroxonium ion as HCl is fully ionized, it suppresses the ionization of water hence and [O ion from water will be less than

It is generally acceptable to consider the ionization of HCl to be the sole source of hydroxonium ions. This is also applicable in strong bases for OH^– ions.

Example1:

If 0.001m of NaOH is added to enough amount of 1L of water, what is the concentration of OH^– and H₃O⁺ ions?

Solution:

Example2:

Calculate the concentration of, and in a 100 cm³ sample of 0.015 m

NOTE:

HCl is the main source of H₃O⁺

NOTE: Molar concentrations are independent of solution volume i.e [H₃O⁺] in 0.015M HCl is the same whether we are describing 1L, 10L or 100cm³ thus the volume of acid is not involved in this calculation.

e.g. Calculate the [], and in 50 cm³ of

Solution:

_{0.01M 0.01M 2 X 0.01=0.02M}

From:

NEUTRALIZATION REACTION.

This is the reaction between H₃O⁺from an acid and OH^–ions from a base to form water. Therefore when solutions of acids and bases are mixed together, the chemical reaction must occur in which and combine to form water.

This reaction occur in order to maintain the required value of equilibrium constant K_w. The final solution can be acidic basic or neutral depending on the and after neutralization reaction.

Example:

What is the [H₃O⁺] obtained by mixing 100cm³ of 0.015m HCl and 50cm³ of 0.01m Ba(OH)₂ solution, Is the final solution acidic or basic?

Solution:

=[][]

Question:

1) What is the pH of a solution obtained by dissolving 312 cm³ of HCl, measured at 30°c at 340mmHg in 3.25lL of water?

Solution:

From

[] =1.7626

–

2) Calculate the pH of neutralization point when 40cm³ of 0.1m NaOH is mixed with 60cm³ of 0.1m HCl.

Solution:

3) Calculate the pH of the solution obtained when.

a) 1cm³ of 0.1m NaOH is added to 100cm³ of 0.001m HCl

b) 1cm³ of 0.1m NaOH is added to 100cm³ of 0.1m HCl

Solution:

–

Again

BUFFER SOLUTIONS:

A buffer solution is a solution which maintains its pH when small amount of an acid or alkali is added to it.
OR

Is the one that resist a change in pH when small amount of acid or alkali is added to the solution.

A buffer solution usually consists of a weak acid and one of its salt or a weak base and one of its salts.

Types of buffer solutions.

i) Acidic buffer solution.

This is the buffer solution which keeps the pH below 7.

They are formed by mixing a weak acid and its salt (of a strong base)

e.g.

How does the buffer system work?

Consider

( )

Since the salts is strong, it dissociates completely into increases the concentration of shifting the equilibrium to the left hand side suppressing the dissociation of acetic acid due to common ion effect. Hence is equal to the salt concentration and due to the common ion effect becomes equal to the initial concentration of the acid.

Therefore the solution will contain these important species.

i)A lot of unionized acid

ii) A lot of acetate ions from

iii) Enough to make the solution acidic.

When little H⁺are added, the following reaction occurs.

® The acetate ions concentration from the salt are large enough to consume the added hydrogen ions therefore there will be no accumulation of H⁺ in the solution.

® If OH^– are added the following reaction occur ;

This decrease in the solution, shifting the equilibrium to the right hand side to replace used to neutralize added. Therefore no accumulation of in the solution.