Chemistry 12
UNIT IV - ACIDS AND BASES ver 1.2
I. CHARACTERISTICS OF ACIDS AND BASES
II. THE ARRHENIUS THEORY OF ACIDS AND BASES
Named after Svente Arrhenius who received the Nobel prize in Chemistry in 1903 for his work with electrolytes. He was one of the first chemists to develop a complete definition of acids and bases.
ARRHENIUS THEORY
ACID
BASE
SALT
To simplify his theory:
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NEUTRALIZATION REACTION REVIEW
Ex. Write the neutralization reaction between H3PO4 and calcium hydroxide.
Try: Write the neutralization reaction between magnesium hydroxide and hydrogen iodide.
Consider the following neutralization reaction:
The complete ionic equation would be:
The net ionic equation would be:
DESCRIPTIVE DEFINITIONS OF ACIDS AND BASES
All acids have certain properties in common, as do all bases.
ACIDS
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BASES
Litmus is a special type of paper that turns colour in the presence of acids or bases, important to know and easily memorized by:
COMMON ACIDS
1. SULPHURIC
ACID
2. HYDROCHLORIC
ACID
3. NITRIC
ACID
4. ACETIC
ACID
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COMMON BASES
1. SODIUM
HYDROXIDE
2. POTASSIUM
HYDROXIDE
3. AMMONIA
THE NATURE OF H+
Hydrogen Atom Water
Now, dissociation equations for Acids look slightly different:
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III. BRONSTED-LOWRY THEORY OF ACIDS AND BASES
The theory set forth by Arrhenius worked well, however it did not take into consideration the equilibrium reactions that occurred in acids and bases, so another theory was needed. The Bronsted-Lowry Theory is more general and incorporates Arrhenius’ theory into a larger scheme.
BRONSTED-LOWRY THEORY
To state it more simply,
Consider the following typical Bronsted-Lowry acid base equation:
The trick to determining which is the acid and which is the base is to look at the reactant side of the equation. Pick one of the chemical species. Find a similar looking species (having a similar composition) on the product side.
Ex. Identify the Acid and the Base in the following reaction:
CH3COOH + H2O ↔ CH3COO- + H3O+
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Consider the original reaction again:
NH3 + H2O ↔ NH4+ + OH-
Forward reaction:
But look at the reverse reaction:
Ex.
The substances that differ from each other by only one proton are referred to as CONJUGATE ACID BASE PAIRS. In any Bronsted-Lowry equation there are two conjugate acid base pairs.
Ex.
Ex. What is the conjugate base of:
a. H2O c. HN3
b. HNO2 d. HSO4-
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Ex. What is the conjugate acid of:
a. OH- c. PO43-
b.HC2O4- d. HSO4-
Ex. Identify the conjugate acid base pairs and state which is the acid and which is the base:
a. HNO3 + H2O ↔ NO3- + H3O+
b. HCO3- + SO32- ↔ CO32- + HSO3-
Ex. Write the Bronsted-Lowry equations for the reaction between:
a. HCN- and F-
b. NO2- and HSO3-
WATER
Consider the following two Bronsted-Lowry equilibrium:
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MONOPROTIC
ACID
DIPROTIC
ACID
TRIPROTIC
ACID
Note that diprotic and triprotic acids that have donated a proton can be AMPHIPROTIC
RECOGNIZING AMPHIPROTIC SUBSTANCES
IV. STRENGTHS OF ACIDS AND BASES
STRONG ACIDS
AND BASES
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WEAK ACIDS
ANDS BASES
Don’t confuse the terms STRONG and CONCENTRATED
The relative strength of an acid or base can be determined by looking at the appropriate chart (Relative Strengths of Bronsted-Lowry Acids and Bases)
STRONG ACIDS
STRONG BASES
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WEAK ACIDS
Ex. Rank the following acids in strength, from strongest to weakest:
H2O HI HF H2CO3 H3PO4
WEAK BASES
Ex: Rank the following bases in order from strongest to weakest:
C2O42- OH- NH3 SO42- CO32-
Note – the relationship between conjugate pairs and relative strength.
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IMPORTANT
LEVELLING EFFECT
Consider the following example. 1.0 M aqueous solutions were prepared of four different acids.
1.0 M HClO4 produced:
1.0 M HCl produced:
1.0 M HF produced:
1.0 M CH3COOH produced:
The strong acids:
The weak acids:
Strong acids are 100 % dissociated in water to form H3O+ and an anion. Water is said to have LEVELLED all the strong acids to the same strength.
Note- recall the electrolytes conduct electricity due to ions being present in solution. The greater the
concentration of ions, the more conductive the solution. So,
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V. EQUILIBRIUM CONSTANT FOR THE IONIZATION OF WATER
A solution can be classified as acidic, basic, or neutral based on the relative concentrations of H3O+ and OH-.
IONIZATION OF WATER
The reaction between a Strong acid and Strong base
Even if no acid or base is present, pure water will always contain small amounts of H3O+ and OH-.
This is a result of collisions between water molecules.
An equilibrium constant for this reaction can be written as:
Since water is a pure liquid, the concentration of H2O is a constant at a given temperature, and so it is incorporated into the Keq value.
The value of Kw varies only with temperature. Unless otherwise stated, assume that:
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Ex. What is the [H3O+] in pure water?
Ex. What is the [OH-] in 0.25 M HCl?
Try: What is the [H3O+] in 0.0075 M NaOH?
VI. pH and pOH
When working with dilute solutions of strong or weak acids or bases, the [H3O+] or [OH-] can be very small (often less than ______). With such small concentrations, it can often be difficult to compare them, so new units were developed to make it easier for comparisons to be made. The new units were pH and pOH.
Quick math lesson – logarithms
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CONVERTING FROM [H3O+] and [OH-] to pH and pOH.
Ex. If the hydronium ion concentration is 4.67 x 10-5 M, what is the pH?
Ex. If the hydroxide ion concentration is 2.83 x 10-6 M, what is the pOH?
Consider the following:
pH = - log[H3O+]
To perform the calculation on a calculator:
CONVERTING FROM pH and pOH to [H3O+] and [OH-].
Ex. If the pH is 3.17, what is the hydronium ion concentration?
Ex. If the pOH = 5.32, what is the hydroxide ion concentration?
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There is a very important relationship between pH, pOH, and Kw.
Start with the Kw expression.
Using the following chart (showing the relationships between H3O+, OH-, pH, and pOH, you can work back and forth between any of the values.
PH FACTORS TO REMEMBER
1. It is an important fact to realize that since pH and pOH are logarithmic scales, a difference in one
pH or pOH unit is equivalent to a 10 fold difference in concentration of the ion.
pH scale
[H3O+]
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2. pH and significant figures are slightly different than significant figures in other calculations.
Ex. How many significant figures in pH = 4.14?
Take the following example:
[H3O+] = 5.28 x 10-5 M
3. Since pH and pOH are the negative of the exponent,
Low pH and pOH values
High pH and pOH values
4. At 25oC
5. It is possible to have negative pH values, although this will only occur in concentrated strong acids.
Recall that the pH scale was developed to express and compare SMALL [H3O+], so negative pH
values are of little use.
Ex.
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VII. MIXING STRONG ACIDS AND BASES
When a strong acid and base are mixed, a neutralization reaction occurs. The resulting solution may be acidic, basic, or neutral depending on the relative amounts of each of the reactants that were used.
(recall limiting reactant calculations from Chemistry 11)
Ex. What is the pH that results when 25.0 mL of 0.250 M HCl is mixed with 35.0 mL of
0.200 M NaOH?
Try: What is the resulting pH of a solution made by mixing 45.0 mL of 0.450 M KOH with 75.0 mL
of 0.275 M HClO4?
CALCULATING MOLES OF ACID NEEDED TO REACH A DESIRED pH
Ex. How many moles of HCl(g) must be added to 40.0 mL of 0.180 M NaOH to produce a solution
that has a pH of 12.500? Assume that there is no change in volume when the HCl is added.
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VIII. ACID AND BASE EQUILIBRIUM CONSTANTS
Weak acids and bases are represented as equilibrium systems as they do not completely dissociate.
ACID
IONIZATION
CONSTANT
To find the Ka values, look down the LEFT side (the acid side) of the Strengths of Acids table and find the species that you are looking for that is acting as an acid.
Ka values for a STRONG acid are not listed in the table of Strengths of Acids. Since they are 100 % ionized, the concentration of the un-ionized acid in the denominator of the Ka expression would be zero.
BASE
IONIZATION
CONSTANT
The table of Strengths of acids does not list the values of Kb, but these can be calculated using the Ka values given in the table.
There is a very important relationship existing between Ka and Kb for CONJUGATE ACID BASE PAIRS.
Consider the following reaction:
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CALCULATING Kb from Ka
Ex. Calculate the Kb for C2O42-.
Try: Calculate the Kb for NO2-.
Note – because of the relationship between the acid and base ionization constants:
IX. RELATIVE STRENGTHS OF ACIDS AND BASES
If solutions containing H2CO3 and SO32- are mixed, the SO32- can only act as a base since it has no protons that it can donate.
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Note that all Bronsted-Lowry reactions that are studied in Chemistry 12 will only involve the transfer of a single proton. There won’t be two or three proton transfers such as :
If solutions containing amphiprotic ions such as HCO3- and H2PO4- are mixed:
In the following equilibrium:
H2CO3 + SO32- ↔ HCO3- + HSO3-
There are two conjugate acid base pairs and there is a ‘competition’ set up as to which of the acids (H2CO3 or HSO3-) will donate their proton.
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X. SALT HYDROLYSIS
Salts consist of a cation (positive ion) and an anion (negative ion). Many salt contain an ion that can react with water.
HYDROLYSIS
In this section we will deal only with the reactions between ions and water. Reactions between ions may occur, but they are not being considered at the moment.
Before looking at ions that will react we will first look at ions that won’t react with water:
Recall SPECTATOR
IONS
When considering the hydrolysis of ions, the following ions do NOT hydrolyze:
SPECTATOR
CATIONS
SPECTATOR
ANIONS
When an ion hydrolyzes it is simply acting as a Bronsted-Lowry acid or base with water.
ANIONIC HYDROLYSIS
If the anion (- ion) of the salt hydrolyzes:
CATIONIC HYDROLYSIS
If the cation (+ ion) of the salt hydrolyzes:
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PROCEDURE TO DETERMINE THE BEHAVIOUR OF A SALT IN WATER
Ex. Predict if each of the following salts will hydrolyze in water. If so, write the hydrolysis equation
for the reaction.
a. NaCl
b. NH4Cl
c. KF
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d. NH4NO2
e. NaHCO3
METAL IONS
Some metal ions tend to form hydrated complexes, where water molecules join up and form a polyatomic ion consisting of a central metal ion with water molecules attached. This is because:
Any time you see a metal ion with water molecules attached, it will be located on the Table of Acid Strengths. NEVER split off the water molecules and attempt to show the metal ion separately.
Ex.
Metal ions from Group 1 and 2 (except Be2+) do not hydrolyze. The most common metal ions that you will see in a hydrolysis reaction are:
METAL OXIDES
When a metal oxide is added to water, there is an initial dissociation of ions.
Ex.
Ex.
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The metal ions are spectator ions and the oxide ion is a ______. The reaction for the
hydrolysis of the oxide ion is:
Since both metal and hydroxide ions are present in solution, the hydrolysis can be written as:
NON-METAL OXIDES
When a nonmetal oxide reacts with water, the water tends to bond to the existing oxide
molecule to produce ______.
Ex.
Ex.
Ex.
SUMMARY
XI. CALCULATIONS INVOLVING Ka
The following information is involved in a Ka problem:
1.
2.
3.
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So, there are only three types of Ka problems that can be asked:
1.
2.
3.
USING INITIAL CONCETRATION OF ACID AND Ka TO DETERMINE [H3O+] OR pH
Ex. Calculate the pH of a 0.75 M acetic acid solution
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CALCULATING Ka FROM [H3O+] or pH
Ex. If the pH of 0.100 M HCHO2 is 2.38 at 25oC, calculate Ka.
FINDING THE INTIIAL CONCENTRATION OF A WEAK ACID
Ex. What mass of NH4Cl will produce 1.50 L of a solution having a pH of 4.75?
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XII. CALCULATIONS INVOLVING Kb
When a weak base is put into water, some of the base ionizes. As a result of the ionization, a certain amount of hydroxide will be produced. The smaller the value of Kb for the particular base, the less ionization that occurs, the less OH- formed.
Calculations involving weak bases are similar to the calculations involving weak acids, with two important differences:
USING INITIAL CONCETRATION OF BASE AND Kb TO DETERMINE pH
Ex. Calculate the pH of 0.10 M NaCN.
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CALCULATING Kb FROM pOH or pH
Ex. The pOH of a 0.50 M solution of the weak acid HA is 10.64. What is Kb for the
conjugate base A-?
XIII. ACID BASE TITRATIONS
Recall: A titration is a process in which a measured amount of a solution is reacted with a known volume of another solution (and one of the solutions has an unknown concentration) until the ‘equivalence point’ (also known as stoichiometric point) is reached.
EQUIVALENCE
POINT
Recall: MOLARITY
ALL titration problems involve at least FIVE PARAMETERS