Notes for Acids & Bases
Arrhenius theory on acids & bases:
Svante Arrhenius was the first to try to describe acids and bases. He suggested that acids are compounds that contain hydrogen and can dissolve in water to release hydrogen ions into solution. For example, hydrochloric acid (HCl) dissolves in water as follows:
HCl --> H+(aq) + Cl-(aq)
Arrhenius defined bases as substances that dissolve in water to release hydroxide ions (OH-) into solution. For example, a typical base according to the Arrhenius definition is sodium hydroxide (NaOH): NaOH --> Na+(aq) + OH-(aq)
The Arrhenius definition of acids and bases explains a number of things. It explains why all acids have similar properties to each other (and also why all bases are similar): because all acids release H+ into solution and all bases release OH-). The Arrhenius definition also explains why acids and bases counteract each other, or neutralize each other.
Neutralization
As you can see from the equations, acids release H+ into solution and bases release OH-. If we were to mix an acid and base together, the H+ ion would combine with the OH- ion to form H2O, or plain water: H+(aq) + OH-(aq) --> H2O
The neutralization reaction of an acid with a base will always produce water and a salt.
Acid + Base --> Water + Salt
HCl + NaOH --> H₂O+ NaCl
HBr + KOH --> H₂O+ KBr
But, Arrhenius theory has limits. It doesn't explain everything. It doesn't explain why some substances like baking soda (NaHCO3) can act like a base even though they do not contain hydroxide ions.
Bronsted - Lowery theory on acids & bases.
In 1923, the Danish scientist Joannes Bronsted developed another theory about acids & bases. In his theory, a proton, H⁺, is transferred from an acid to a base.
For the reaction: H₂O + HCl --> H₃O⁺ + Cl⁻
When HCl reacts with water, a proton is transferred. Acids donate protons and bases accept protons. In this reaction HCl is an acid and H₂O is a base.
A substance can be classified as a Bronsted-Lowry acid or base only for a specific reaction because some substances are amphoteric. Sometimes they act as a base and sometimes they act as an acid. Water is an example of this. In the above reaction, water is a base. But, water can also act as an acid:
NH₃ + H₂O --> NH₄ ⁺ + OH⁻ In this reaction water is donating a proton, it is an acid.
H₃O⁺ is called a hydronium ion
This also explains why baking soda (NaHCO3), for example, acts like a base. It is because it accepts a proton from an acid. This is the reaction: HCl + NaHCO3 --> H₂CO3 + NaCl
H₂CO3 is carbonic acid and in the above reaction, it then decomposes to water and carbon dioxide gas. This gas produces bubbles.
Acid base reactions are all reversible and they form acid bas equilibriums.
Strong & Weak Acids
Some acids are stronger than others. The difference between a strong acid and a weak acid is that strong acids more readily (easily) donate their protons. Although acid base reactions reach equilibrium, a strong acid will have more product than a weak acid. With a strong acid, the equilibrium favours the product side of the reaction. With very strong acids the reaction may go to completion.
The position at equilibrium is the percent ionization.
For the reaction of acetic acid and water:
HC₂H₃O₂ + H₂O <--> C₂H₃O⁻ + H₃O⁺
In the forward reaction, acid: HC₂H₃O₂, base: H₂O
In the reverse reaction: acid H₃O⁺ , base: C₂H₃O⁻
Short form for this reaction (remove the water from both sides of the equation: HC₂H₃O₂ <--> C₂H₃O₂⁻ + H⁺
Conjugate acid base pairs: In the above reaction, HC₂H₃O₂ and C₂H₃O₂⁻ are conjugate acid base pairs.
One is the acid in the forward direction and one is the base in the reverse direction.
Acetic acid (HC₂H₃O₂) is an acid. It's conjugate base is C₂H₃O₂⁻
They are very similar. The only difference is one proton.
Also, H₂O and H₃O⁺ are conjugate pairs in that reaction.
Autoionization of water
Water is never just a collection of water molecules. There is some electrical conductivity in pure water which suggests that there are some ions present. If 2 water molecules collide in the right way, then you can get autoionization of water, where H₂O --> OH⁻ + H⁺
This is also called the dissociation of water and in fact you can have
Kw = [H+][OH-] = 1.00x10-14 at 25C where Kw is the equilibrium constant for water.
Notice that is quite a low value of Kw so reactants are favoured quite a bit over products, i.e. very little of ionization of water occurs. The percent ionization is very small.
pH
Another thing that we will be looking at is pH where pH = -log[H+].
From our Kw, we know that the [H+] is 1.00 × 10-7 so we can calculate the pH of water ......
pH = -log[H+]= -log[1.00 × 10-7] = 7.00
Strong acids:
A strong acid is one that ionizes quantitatively (completely) in water. [Remember that a reaction that goes to completion, with no equilibrium, is a quantitative reaction]. The percent ionization is > 99% but for our calculations we will assume that it is 100%. There are not too many strong acids. Hydrochloric acid (HCl), hydrobromic acid (HBr), sulphuric acid (H₂SO₄), phosphoric acid (H₃PO₄) and nitric acid (HNO₃) are the main ones.
HCl --> H⁺ + Cl⁻ 100% ionization in H₂O .
If we have a solution of 0.1 M HCl, we really have a solution that is 0.1 M H⁺ + 0.1 M Cl⁻
Monoprotic acids – have one ionizable hydrogen atom (HCl)
Diprotic acids – have two ionizable hydrogen atoms. (H₂SO₄)
Triprotic acids – have three ionizable hydrogen atoms (H₃PO₄)
If an acid is strong, it’s conjugate base is weak.
Calculating [OH⁻] or [H⁺] after adding acid
Example: If a 0.15 mol/L solution of HCl at SATP has a [H⁺] of 0.15 mol/L what is the [OH⁻]?
1. Identify the substances that contribute to the [H⁺] and the [OH⁻]
2. Are there major entities/contributors? (Are there substances that contribute a lot to the [H⁺] and the [OH⁻]
3. Are there entities/substances that only contribute a very small amount to the [H⁺] and the [OH⁻]? If yes, these can be ignored because in most cases a major entity.
Since HCl is a strong acid, assume 100% ionization.
Therefore 0.15 mol/L solution of HCl will have 0.15 mol/L of H⁺ and 0.15 mol/L Cl⁻
Remember that the Kw of water: Kw = [H+][OH-] = 1.00x10-14
In water then [H+] = 1.00x10-7 and [OH-] = 1.00x10-7
BUT addition of HCl to the water will push the H₂O <--> OH⁻ + H⁺ equilibrium to the left (Le Chateliers principle). The concentrations of both OH⁻ and H⁺ (from water) will then be less than 1.00x10-7
Knowing that [H⁺] = 0.15 mol/L, you can use the equation Kw = [H+][OH-] = 1.00x10-14
to calculate the [OH-] and it is 6.7 x 10-14 mol/L (text pg 535 & 536)
Strong Bases
According to Arrhenius, a base is a substance that dissociates to increase the hydroxide concentration of a solution. Ionic hydroxides have varying solubility in water, but all of them are strong bases that dissociate quantitatively when they dissolve in water.
All of the hydroxides of Group 1 elements are strong bases – LiOH, NaOH, KOH, RbOH, and CsOH. When these bases dissolve in water, one mole of hydroxide ion is produced for every mole of metal hydroxide that dissolves in solution. Metal hydroxides are highly soluble in water.
Group 2 elements also form strong hydroxides – Mg(OH)₂, Ca(OH)₂, Ba(OH)₂ and Sr(OH)₂. When these bases dissolve in water, 2 moles of hydroxide ion are produced for every mole of metal hydroxide that dissolves in solution. These metal hydroxides are only slightly soluble in water. Their low solubilities makes them useful in medical applications (antacids).
Calculating [OH⁻] or [H⁺] after adding base
In a 0.025 mol/L aqueous solution of barium hydroxide, a strong base, what is the [H⁺]?
Ba(OH)₂ --> Ba²⁺ + 2 OH⁻
Since it is a strong base, Ba(OH)₂ should ionize completely. For every mole of Ba(OH)₂ there will be 2moles OH⁻.
Therefore, [OH-] is 0.050 mol/L.
To find the [H⁺] we have to look at the ionization of water AND the effect that Ba(OH)₂ has on that equilibrium, like we did when we found the [OH⁻] after adding HCl to water.
Kw = [H+][OH-] = 1.00x10-14
If [OH-] is 0.050 mol/L then the [H+] is 2.00 x 10-14 mol/L (text, page 538)
pOH
Just like the pH = -log [H+]
pOH = -log [OH-]
For the reaction that we just looked at, where we found that [OH-] was 0.050 mol/L, we can calculate the pOH
0.05 = 5 x 10-2, pOH = -log [OH-] = 1.3
There is a mathematical relationship between pH and pOH (at SATP)
pH + pOH = 14.
If pOH = 1.3, pH = 12.7
If you are interested in how the math works for this relationship it is the text on pages 542 & 543. You will NOT be tested on this math but you will need to know that pH + pOH = 14.
When pH = pOH the solution is said to be neutral. This is at pH = 7.
If pH > 7, basic solution (also called an alkaline solution) [OH-] > [H+]
If pH < 7, acidic solution, [H+] > [OH-]
The problems on page 545 - 549 are about calculating pH and pOH of strong acids and strong bases.
A summary of the equations used in looking at pH and pOH are on page 549.
Measuring pH
There are several ways to measure pH.
Many plant compounds (substances found in plants) and synthetic dyes change colour when mixed with an acid or base. These are called acid-base indicators. A common indicator is litmus. Litmus is a dye obtained from a plant called a lichen. It is used to make litmus paper, strips of which are used to measure pH. If the litmus paper comes into contact with neutral solutions, the litmus paper stays brown. If the solution is acidic, the litmus paper turns pink. If the solution is basic, the litmus paper turns blue.
A more accurate way to measure pH is with a pH meter. A pH meter measures the voltage generated by a pH dependent voltaic (or galvanic) cell.
Weak Acids & Bases
So far we have been talking about strong acids and bases - those that ionization completely. But, there are many common substances that are weak acids or weak bases. Vinegar (acetic acid) and Vitamin C (ascorbic acid) are weak acids.
--> Weak acids and bases do not ionize completely.
--> They are weak electrolytes [electrolytes are substances make solutions electrically conductive because they have free ions - most are in ionic solutions]
--> Most common acids are weak.
-->Many naturally occurring acids are weak acids - example: carboxylic acids [remember organic chemistry!]
--> Common inorganic acids include: hydrofluoric acid (HF), carbonic acid (H₂CO₃) hydrosulfuric acid (H₂S)
and boric acid (H₃BO₃)
--> Most weak acids ionize less than 50%, for example acetic acid only ionizes 1.3% at 25C and 0.10 mol/L concentration.
Weak bases have a lone pair of electrons to accept a proton from water, but it is a weak attraction for protons.
The general equation for a base can be written as
B: + H₂O <--> OH⁻ + HB⁺ where B: = base with pair of lone electrons
Calculating percent ionization of a weak acid from the pH.
If you know the pH of a solution, you can calculate the percent ionization.
Svante Arrhenius was the first to try to describe acids and bases. He suggested that acids are compounds that contain hydrogen and can dissolve in water to release hydrogen ions into solution. For example, hydrochloric acid (HCl) dissolves in water as follows:
HCl --> H+(aq) + Cl-(aq)
Arrhenius defined bases as substances that dissolve in water to release hydroxide ions (OH-) into solution. For example, a typical base according to the Arrhenius definition is sodium hydroxide (NaOH): NaOH --> Na+(aq) + OH-(aq)
The Arrhenius definition of acids and bases explains a number of things. It explains why all acids have similar properties to each other (and also why all bases are similar): because all acids release H+ into solution and all bases release OH-). The Arrhenius definition also explains why acids and bases counteract each other, or neutralize each other.
Neutralization
As you can see from the equations, acids release H+ into solution and bases release OH-. If we were to mix an acid and base together, the H+ ion would combine with the OH- ion to form H2O, or plain water: H+(aq) + OH-(aq) --> H2O
The neutralization reaction of an acid with a base will always produce water and a salt.
Acid + Base --> Water + Salt
HCl + NaOH --> H₂O+ NaCl
HBr + KOH --> H₂O+ KBr
But, Arrhenius theory has limits. It doesn't explain everything. It doesn't explain why some substances like baking soda (NaHCO3) can act like a base even though they do not contain hydroxide ions.
Bronsted - Lowery theory on acids & bases.
In 1923, the Danish scientist Joannes Bronsted developed another theory about acids & bases. In his theory, a proton, H⁺, is transferred from an acid to a base.
For the reaction: H₂O + HCl --> H₃O⁺ + Cl⁻
When HCl reacts with water, a proton is transferred. Acids donate protons and bases accept protons. In this reaction HCl is an acid and H₂O is a base.
A substance can be classified as a Bronsted-Lowry acid or base only for a specific reaction because some substances are amphoteric. Sometimes they act as a base and sometimes they act as an acid. Water is an example of this. In the above reaction, water is a base. But, water can also act as an acid:
NH₃ + H₂O --> NH₄ ⁺ + OH⁻ In this reaction water is donating a proton, it is an acid.
H₃O⁺ is called a hydronium ion
This also explains why baking soda (NaHCO3), for example, acts like a base. It is because it accepts a proton from an acid. This is the reaction: HCl + NaHCO3 --> H₂CO3 + NaCl
H₂CO3 is carbonic acid and in the above reaction, it then decomposes to water and carbon dioxide gas. This gas produces bubbles.
Acid base reactions are all reversible and they form acid bas equilibriums.
Strong & Weak Acids
Some acids are stronger than others. The difference between a strong acid and a weak acid is that strong acids more readily (easily) donate their protons. Although acid base reactions reach equilibrium, a strong acid will have more product than a weak acid. With a strong acid, the equilibrium favours the product side of the reaction. With very strong acids the reaction may go to completion.
The position at equilibrium is the percent ionization.
For the reaction of acetic acid and water:
HC₂H₃O₂ + H₂O <--> C₂H₃O⁻ + H₃O⁺
In the forward reaction, acid: HC₂H₃O₂, base: H₂O
In the reverse reaction: acid H₃O⁺ , base: C₂H₃O⁻
Short form for this reaction (remove the water from both sides of the equation: HC₂H₃O₂ <--> C₂H₃O₂⁻ + H⁺
Conjugate acid base pairs: In the above reaction, HC₂H₃O₂ and C₂H₃O₂⁻ are conjugate acid base pairs.
One is the acid in the forward direction and one is the base in the reverse direction.
Acetic acid (HC₂H₃O₂) is an acid. It's conjugate base is C₂H₃O₂⁻
They are very similar. The only difference is one proton.
Also, H₂O and H₃O⁺ are conjugate pairs in that reaction.
Autoionization of water
Water is never just a collection of water molecules. There is some electrical conductivity in pure water which suggests that there are some ions present. If 2 water molecules collide in the right way, then you can get autoionization of water, where H₂O --> OH⁻ + H⁺
This is also called the dissociation of water and in fact you can have
Kw = [H+][OH-] = 1.00x10-14 at 25C where Kw is the equilibrium constant for water.
Notice that is quite a low value of Kw so reactants are favoured quite a bit over products, i.e. very little of ionization of water occurs. The percent ionization is very small.
pH
Another thing that we will be looking at is pH where pH = -log[H+].
From our Kw, we know that the [H+] is 1.00 × 10-7 so we can calculate the pH of water ......
pH = -log[H+]= -log[1.00 × 10-7] = 7.00
Strong acids:
A strong acid is one that ionizes quantitatively (completely) in water. [Remember that a reaction that goes to completion, with no equilibrium, is a quantitative reaction]. The percent ionization is > 99% but for our calculations we will assume that it is 100%. There are not too many strong acids. Hydrochloric acid (HCl), hydrobromic acid (HBr), sulphuric acid (H₂SO₄), phosphoric acid (H₃PO₄) and nitric acid (HNO₃) are the main ones.
HCl --> H⁺ + Cl⁻ 100% ionization in H₂O .
If we have a solution of 0.1 M HCl, we really have a solution that is 0.1 M H⁺ + 0.1 M Cl⁻
Monoprotic acids – have one ionizable hydrogen atom (HCl)
Diprotic acids – have two ionizable hydrogen atoms. (H₂SO₄)
Triprotic acids – have three ionizable hydrogen atoms (H₃PO₄)
If an acid is strong, it’s conjugate base is weak.
Calculating [OH⁻] or [H⁺] after adding acid
Example: If a 0.15 mol/L solution of HCl at SATP has a [H⁺] of 0.15 mol/L what is the [OH⁻]?
1. Identify the substances that contribute to the [H⁺] and the [OH⁻]
2. Are there major entities/contributors? (Are there substances that contribute a lot to the [H⁺] and the [OH⁻]
3. Are there entities/substances that only contribute a very small amount to the [H⁺] and the [OH⁻]? If yes, these can be ignored because in most cases a major entity.
Since HCl is a strong acid, assume 100% ionization.
Therefore 0.15 mol/L solution of HCl will have 0.15 mol/L of H⁺ and 0.15 mol/L Cl⁻
Remember that the Kw of water: Kw = [H+][OH-] = 1.00x10-14
In water then [H+] = 1.00x10-7 and [OH-] = 1.00x10-7
BUT addition of HCl to the water will push the H₂O <--> OH⁻ + H⁺ equilibrium to the left (Le Chateliers principle). The concentrations of both OH⁻ and H⁺ (from water) will then be less than 1.00x10-7
Knowing that [H⁺] = 0.15 mol/L, you can use the equation Kw = [H+][OH-] = 1.00x10-14
to calculate the [OH-] and it is 6.7 x 10-14 mol/L (text pg 535 & 536)
Strong Bases
According to Arrhenius, a base is a substance that dissociates to increase the hydroxide concentration of a solution. Ionic hydroxides have varying solubility in water, but all of them are strong bases that dissociate quantitatively when they dissolve in water.
All of the hydroxides of Group 1 elements are strong bases – LiOH, NaOH, KOH, RbOH, and CsOH. When these bases dissolve in water, one mole of hydroxide ion is produced for every mole of metal hydroxide that dissolves in solution. Metal hydroxides are highly soluble in water.
Group 2 elements also form strong hydroxides – Mg(OH)₂, Ca(OH)₂, Ba(OH)₂ and Sr(OH)₂. When these bases dissolve in water, 2 moles of hydroxide ion are produced for every mole of metal hydroxide that dissolves in solution. These metal hydroxides are only slightly soluble in water. Their low solubilities makes them useful in medical applications (antacids).
Calculating [OH⁻] or [H⁺] after adding base
In a 0.025 mol/L aqueous solution of barium hydroxide, a strong base, what is the [H⁺]?
Ba(OH)₂ --> Ba²⁺ + 2 OH⁻
Since it is a strong base, Ba(OH)₂ should ionize completely. For every mole of Ba(OH)₂ there will be 2moles OH⁻.
Therefore, [OH-] is 0.050 mol/L.
To find the [H⁺] we have to look at the ionization of water AND the effect that Ba(OH)₂ has on that equilibrium, like we did when we found the [OH⁻] after adding HCl to water.
Kw = [H+][OH-] = 1.00x10-14
If [OH-] is 0.050 mol/L then the [H+] is 2.00 x 10-14 mol/L (text, page 538)
pOH
Just like the pH = -log [H+]
pOH = -log [OH-]
For the reaction that we just looked at, where we found that [OH-] was 0.050 mol/L, we can calculate the pOH
0.05 = 5 x 10-2, pOH = -log [OH-] = 1.3
There is a mathematical relationship between pH and pOH (at SATP)
pH + pOH = 14.
If pOH = 1.3, pH = 12.7
If you are interested in how the math works for this relationship it is the text on pages 542 & 543. You will NOT be tested on this math but you will need to know that pH + pOH = 14.
When pH = pOH the solution is said to be neutral. This is at pH = 7.
If pH > 7, basic solution (also called an alkaline solution) [OH-] > [H+]
If pH < 7, acidic solution, [H+] > [OH-]
The problems on page 545 - 549 are about calculating pH and pOH of strong acids and strong bases.
A summary of the equations used in looking at pH and pOH are on page 549.
Measuring pH
There are several ways to measure pH.
Many plant compounds (substances found in plants) and synthetic dyes change colour when mixed with an acid or base. These are called acid-base indicators. A common indicator is litmus. Litmus is a dye obtained from a plant called a lichen. It is used to make litmus paper, strips of which are used to measure pH. If the litmus paper comes into contact with neutral solutions, the litmus paper stays brown. If the solution is acidic, the litmus paper turns pink. If the solution is basic, the litmus paper turns blue.
A more accurate way to measure pH is with a pH meter. A pH meter measures the voltage generated by a pH dependent voltaic (or galvanic) cell.
Weak Acids & Bases
So far we have been talking about strong acids and bases - those that ionization completely. But, there are many common substances that are weak acids or weak bases. Vinegar (acetic acid) and Vitamin C (ascorbic acid) are weak acids.
--> Weak acids and bases do not ionize completely.
--> They are weak electrolytes [electrolytes are substances make solutions electrically conductive because they have free ions - most are in ionic solutions]
--> Most common acids are weak.
-->Many naturally occurring acids are weak acids - example: carboxylic acids [remember organic chemistry!]
--> Common inorganic acids include: hydrofluoric acid (HF), carbonic acid (H₂CO₃) hydrosulfuric acid (H₂S)
and boric acid (H₃BO₃)
--> Most weak acids ionize less than 50%, for example acetic acid only ionizes 1.3% at 25C and 0.10 mol/L concentration.
Weak bases have a lone pair of electrons to accept a proton from water, but it is a weak attraction for protons.
The general equation for a base can be written as
B: + H₂O <--> OH⁻ + HB⁺ where B: = base with pair of lone electrons
Calculating percent ionization of a weak acid from the pH.
If you know the pH of a solution, you can calculate the percent ionization.
