Acids and Bases 2016: Class Notes and Examples Name ______Per. _____
1. Arrhenius concept of Acids and Bases – the oldest and more restrictive (less useful) definition.
An acid is a substance that donates hydrogen ions (H+) in solution and a base donates hydroxide ions (OH-) in solution.
Examples of hydrochloric acid reacting with water and sodium hydroxide dissociating in water.
2. BrØnsted-Lowry Acid and Base Model – a more general and much more useful approach!
An acid is a substance that can donate a proton while a base is a substance that can accept a proton.
Example of an acid with water: HNO3 (introducing the hydronium ion: H3O+)
Example of a base with water: NH3
*[Note: the ammonia would not be classified as a base under the Arrhenius definition.]
3. The Nature of Acids and Bases
Acids will react with most metals to produce hydrogen gas and ions in solution.
Example: HCl reacting with magnesium metal:
*Note…acids taste sour and bases taste bitter…if you are of the mind to taste them…and we do! Which one do you feel is more often associated with poisons found in nature?
4. Conjugate Acid-Base Pairs
A conjugate acid-base pair consists of two substances related to each other the by donating and accepting of a single proton (where the acid form of the pair contains a proton that is missing in the base form).
Examples: Acetic acid (HC2H3O2) with water:
Ammonia (NH3) with water:
-the ammonia acts as the base in accepting the proton from the water, while water acts as an acid!
-the ammonium ion (NH4+) is now capable of donating a proton to some other ion (which is why it is called the conjugate acid) while the hydroxide ion is now able to accept a proton (the conjugate base)!
5. Strength of Acids and Bases
a. Strong acids react completely with water to form ions in solution…100% dissociation! The six most common strong acids are HCl, HBr, HI, HNO3, HClO4, H2SO4. (Note: sulfuric acid is considered to be a diprotic acid while the others in this list are monoprotic. Other examples of polyprotic acids: H2CO3; H3PO4; H2S)
Example:
b. Weak acids only dissociate to a small degree and then reach an equilibrium. (Most often these dissociate ≤1%.)
Example:
c. Strong and weak bases behave in a similar manner producing hydroxide ions. All metals attached to hydroxide are considered strong bases…ammonia and other amines are the common weak bases.
6. pH scale
This is a logarithmic scale (power of 10) due to the immense concentration differences encountered when dealing with acid/base reactions.
Equation: pH = -log[H+]*the brackets[ ] around H+ means concentration in molarity of H+ ions!
The pH changes by 1 for every power of 10 change in the concentration of H+. For example, a solution of pH 3 has an H+ concentration 10 times that of a solution of pH 4 and 100 times that of a solution of pH 5. Also note that because
pH = -log [H+], the pH decreases as H+ increases.
· pH scale range…is it 0-14 (actually no!) A pH less than 7 is considered acidic and greater than 7 is considered basic.
· Sig. Figs. for pH/pOH: the number of sig figs = the number of decimal places in the pH/pOH value…feel free to ignore this if your teacher tells you it is OK!
Examples: [H+] = .0100(3 s.f.) so the pH = 2.000(3 decimal places) –or- [H+] = .010(2s.f.) so the pH = 2.00(2 decimal places)
a) Example pH calculation:
More Examples:
(1) What is the pH of a 0.0010 M of HCl solution?
(2) What is the pH of a 0.010 M HI solution?
(3) What is the pH of a 1.0 x 10-6 M solution of HNO3?
(4) What is the pH of a 1.5 x 10-5 M solution of HNO3?
You should also be able to reverse this calculation to determine the concentration of H+ when given the pH!
[Take the inv. log of the negative pH value. Practice this calculation on #1-4 above as you already have the answer!]
(5) What is the [H3O+] of a pH 5.00 solution?
b) pOH calculation: Equation: pOH = -log[OH-]
Example:
More Examples:
(1) What is the pH of a 0.0010 M NaOH solution?
(2) What is the pH of a 0.010 M KOH solution?
(3) What is the pH of a 1.0 x 10-6 M solution of Ba(OH)2? (caution…this particular base produces two OH- ions per formula unit)
(4) What is the pH of a 1.5 x 10-5 M solution of HNO3?
Important connections b/t pH and pOH…using water as the solvent allows us to use the following connection:
The concentrations of H+ and OH- in room temperature water are equal to each other and are 1.00 x 10-7. If you determine the pH and pOH of pure water, you will get a 7 for each…thus for all aqueous solutions, pH + pOH = 14! This is a very easy method of determining pOH from pH or vice versa…as well as [H+] or [OH-]!
c) Example Calculations cont.
(1) What is the pOH of a solution whose pH = 3.0?
(2) What is the pH of a solution whose pOH is 8.50?
(3) A solution has a [OH-] = 0.50M. Is it basic or acidic? What is the pOH, pH and [H+]?
d) What are the pH and pOH of a solution made by dissolving 0.20 g NaOH in enough H2O to make 500. mL of solution?
Step 1: Determine how the pH will change by finding the pH (pOH) affecting species.
Step 2: Determine the molarity of the solution.
Step 3: Calculate either the pH or pOH.
5. Titration Reaction
A titration reaction is a lab reaction carried out to determine the concentration of an unknown acid or base solution. The reaction between a strong acid and base results in a neutralization when equal concentrations have reacted and the product is water and a salt. The point when the acid and base exactly neutralize each other is called the equivalence point or end point if you are in the lab.
Equation:
In this reaction, the ratio of HCl to NaOH is 1:1. This means that one mole of acid will react with one mole of base.
Thus, 0.25 mol HCl will neutralize 0.25 mole NaOH.
Example Titration Calculations:
1. It takes 30. mL of 0.50 M HCl to neutralize 1.0 g of an unknown base. How many moles of base do you have? What is the molar mass of the unknown base?
i) How many moles of base? (Find the number of moles of acid used)
ii) Molar mass of the unknown base?
2. What is the concentration of a KOH solution when 90.0 mL of 0.50 M HCl is needed to neutralize 35.0 mL of the base? (Short method of solving when the acid/base are reacting in a 1:1 stoichiometric ratio: MaVa = MbVb)
Step 1:
Step 2:
Step 3:
Redo #2 above by utilizing the short-cut method of solving when the acid/base are reacting in a 1:1 stoichiometric ratio: Short-cut Equation: MaVa = MbVb (note-this method will not work if the ratios in the balanced equation are not 1:1)
Sample Calculations:
3. What is the concentration of a Mg(OH)2 solution when 90.0 mL of 0.50 M HCl is needed to neutralize 35.0 mL of the base?
4. What is the concentration of a H2SO4 solution when 35.5 mL of 1.50 M Al(OH)3 is needed to neutralize 35.5 mL of the acid?
5. Extension Problem/Final Exam Review: A solution of sodium carbonate, a weak base, is made by dissolving 45.5 g of this solid into enough water to make 150. mL of solution. Hydrochloric acid is then added until all of the base has reacted with the acid and the bubbling has stopped. This required 65.0 mL of the acid solution to reach this endpoint.
a. Determine the molarity of the hydrochloric acid solution.
b. Determine the volume of gas produced if the lab pressure was 98.7 kPa and the temperature was held at 25.6C. The vapor pressure of water is 24.617torr at this temperature.
c. If the temperature was reduced to 0.00C and the pressure was raised to 101.325 kPa, what volume would the gas now occupy?