Study Guide:Acids and Bases
1.Compare and contrast acid and bases. List at least three differences between them.
2.List the names and formulae of two binary acids and two oxyacids. How is the naming oxyacids is different from naming binary acids?
3. What are Arrhenius acids and bases?
4. What are strong acids and bases? List five strong acids and five strong bases.
5. Write the neutralization reaction that occurs between a strong acid and a strong base taking strontium hydroxide and nitric acid as example. Write the net ionic equation for this reaction.
6.How is a hydronium ion formed? Show the formation of hydronium ion, when the two water molecules react with each other.
7. What is the pH scale? How does it work?
8. List iodic acid, hypoiodous acid and periodic acid in the order of increasing strength and give reasons for your answer.
9.What is the hydronium ion concentration of an aqueous solution that has a pH of 5.0?
+ - 10.Calculate [H3O ] and [OH ] for each of the following: a. 1X 10-4 M HCl
b. 1.0 X 10-4 M NaOH
1 | P a g e Chemistry Honors pH problems
1) Calculate the pH of the following solutions
a) 0.010 M HCl
b) 0.10 M H2SO4
c) 0.0050 M KOH
d) 0.00050 M Ba(OH)2
e) 0.020 M HNO3
f) 0.25 M HCl
g) 0.080 M NaOH
+ 2) Calculate the [H ] associated with the following pH values
a) 3.4
b) 8.2
c) 12.0
d) 4.8
e) 6.2 (pOH)
f) 5.2 (pOH)
g) 6.8 (pOH)
3) Calculate the pH of the resulting solutions
a) 12.5 ml of 0.020 M H2SO4 are diluted to 50.0 ml
b) 25.0 ml of 0.10 M HCl are mixed with 30.0 ml of 0.075 M NaOH
c) 10.0 ml of 0.20 M Ba(OH)2 are mixed with 5.0 ml of 0.50 M H2SO4
d) 10.0 ml of 0.10 M HCl, 5.0 ml of 0.10 M H2SO4 and 0.060 g of NaOH are mixed together and diluted to a final volume of 25.0 ml
e) 1.00 g of barium hydroxide dissolved to make 1.00 L of solution
f) 0.10 mole HCl diluted to give 100.0 ml of final solution
g) 0.0050 M Ba(OH)2
2 | P a g e Titration Problems Level 1
Adapted from http://www.mpcfaculty.net/mark_bishop/titration_solutions.pdf
1. The molarity of a hydrochloric acid solution can be determined by titrating a known volume of the solution with a sodium hydroxide solution of known concentration. If 14.7 mL of 0.102 M NaOH is required to titrate 25.00 mL of a hydrochloric acid, HCl, solution, what is the molarity of the hydrochloric acid?
2. If 36.2 mL of 0.152 M NaOH is required to neutralize 25.00 mL of an acetic acid, HC2H3O2, solution, what is the molarity of the acetic acid? (Obj #25)
3. The molarity of a sodium hydroxide solution can be determined by titrating a known volume of the solution with a hydrochloric acid solution of known concentration. If 19.1 mL of 0.118 M HCl is required to neutralize 25.00 mL of a sodium hydroxide solution, what is the molarity of the sodium hydroxide?
4. If 7.3 mL of 1.25 M HNO3 is required to neutralize 25.00 mL of a potassium hydroxide solution, what is the molarity of the potassium hydroxide?
5. If 12.0 mL of 1.34 M NaOH is required to neutralize 25.00 mL of a sulfuric acid, H2SO4, solution, what is the molarity of the sulfuric acid?
3 | P a g e Problems on Titration
1. 0.263 g of Na2CO3 requires 28.35 mL of aq HCl for titration to the equivalence point. What is the molar concentration of HCl? 0.175M
2. A 25.0 ml sample of vinegar requires 28.33 mL of a 0.953 M solution of NaOH for titration to the equivalence point. What mass in g of acetic acid is in the vinegar sample, and what is the concentration of acetic acid in vinegar?
3. A liter of approximately 0.2M HCl is prepared by diluting 17 mL of concentrated HCl solution to about 1L. The solution is standardized by titrating a 0.5015 g sample of pure fry Na2CO3. What is the exact concentration of the HCl solution if the titration requires 48.47 mL of acd to reach the end point. 0.1952M
4. Titration of a 40.00ml sample of a solution of H3PO4 requires 35.00 ml of 0.1500 M KOH to reach the end point. Determine the molar concentration of H3PO4 if the reaction is 2KOH + H3PO4 -- K2HPO4 + 2H2O 0.06562
4 | P a g e Review for Acid Base Quiz Chemistry Honors
1.Calculate the pH of the solution after 15.0 ml of 0.100 M NaOH has been added to the 25 ml of 0.100 M HCl solution. 1.602
-3 - 2.Calculate pH and [OH-] of a 5 X 10 M HClO4 solution. pH= 2.30, [OH-]= 2 x 10 12M
3.pH of a solution is 11.93. What is its [H+] ion concentration? [OH-]? POH? [H+]= 1.2 X10-12 M, [OH-]=8.5 X10-3 M, pOH= 2.07
4. If [nut buter]= 89.8 M, what is pNut butter? -1.95
5. An acid was titrated to the equivalence point with 0.0100 M KOH. It required 52.0 ml of KOH for 26.0 ml of acid. What was the acid molarity? 0.0200M
6. A beginning technician was told to standardize an NaOH solution. He carried out the titration and reported that 1.65 ml of base required 27 ml of 0.50 M acid. He therefore concluded that the NaOH was 0.818 M. Assume you are his supervisor. Check the results and criticize the work.
7. Write molecular, ionic and net ionic equations for the following: Aqueous Nitric acid reacts with barium hydroxide.
5 | P a g e Chapter 15 , 16 & 18.3 Review for Acids and Bases: ChemIsTry Honors
1. Write formulas for the following. Identify each as acid, base, or salt.
2. a. Magnesium hydroxide d. sulfurous acid b. hydrochloric acid e. sodium hypochlorite c. zinc nitrate f. potassium hydroxide.
3. Define an acid according to the following theories giving examples. Arrhenius theory
Bronsted-Lowry theory
Lewis theory
4. Identify the Lewis acid and Lewis base in the reaction of Ni2+ with four water ligands. 2+ 2+ Ni (aq) + 4H2O(l)à Ni(H2O) 4 (aq)
-2 5. Write equations showing that the hydrogen phosphate ion (HPO4 ) is amphoteric
6. Write the three equations for the stepwise ionization of phosphoric acid.
7. Use the Bronsted- Lowry and Lewis definitions of acid and bases to identify each reactant as an acid or a base.
a. KOH + HBr ßà KBr + H2O
+ b. HCl + H2O ßà Cl- + H3O
2. Use the phosphate buffer (H2PO4-/ HPO42-) to illustrate how a buffer system works. Show, by means of equations, how the pH of a solution can be kept almost constant when small amounts of acid or base are added.
3. How would the equilibrium between hypochlorous acid and the hypochlorite ion be affected by the addition of each? - - HOCl (aq) + OH (aq) ßà OCL (aq) + H2O (l) a. HCl b. NaOH
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4. Arrange these solutions in order of decreasing acidity. a. 0.1M NaOH b. 0.1M HCl c. 0.1M ammonium chloride d. 0.1M sodium ethanoate
5. Use the cyanate buffer HOCN/ OCN- to illustrate how a buffer system works. Show, by means of equations, how the pH of a solution can be kept almost constant when small amounts of acid or base are added.
+ - H (aq) + OCN (aq) HOCN(aq)
HOCN(aq) + OH- (aq) H2O(l) + OCN- (aq)
6. Suppose you slowly add 0.1M NaOH to 50.0mL of 0.1M HCl. What volume of NaOH must be added for the resulting solution to become neutral? Explain your reasoning.
For a strong acid (HCl) and strong base (NaOH), the concentration of H+ and OH- has to be the same. Since both the base and acid are at equal molarity then they would need the same volume. So it takes 50 mL of NaOH to neutralize 50 mL of HCl.
14.What is hydrolysis? Give an example of a salt that hydrolyses to give an acidic solution.
Hydrolysis is a reaction of water molecules and ions of a dissolved salt.
NH4Cl(aq) + H2O(l) NH4OH(aq) + HCl (aq) (salt) (water) (weak base) (strong acid)
So that means the overall reaction produces an acidic solution.
13. Identify whether the following salts produce acidic, basic or neutral solutions by writing the appropriate hydrolysis reactions.
a. Na C6H5CO2 ( Sodium Benzoate)
NaC6H5CO2 (aq) + H20(l) NaOH(aq) + C6H6CO2 (aq) Produces an basic solution with strong base NaOH
b. KCl
KCl(aq) + H20(l) KOH (aq) + HCl(aq) Produces a neutral solution due to the presence of strong base KOH and strong acid HCl.
c. CH3COOK ( Potassium Acetate)
CH3COOK(aq) + H2O(l) KOH (aq) + C2H4O2(aq) Produces an basic solution with the presence of strong base KOH.
7 | P a g e Weak Acid and Base and ksp problems: (OPTIONAL)
1. Calculate the hydronium ion concentration in a 0.045 M HIO solution.
2. What is the value of Kb for NH3(aq) if 0.1250M solution has the following equilibrium + concentrations ; [NH3 ] =
3. - + [OH ] = 1.478 x 10M, [NH3 ] = 0.1235M
4. Find the experimental ionization constant for a 0.0535M HClO solution that ionizes 2.35%
5. Calculate the percent ionization of a 0.075M HCN solution.
6. + Use the quadratic equation to calculate the [H3O ] in a 0.750M CH2ClCOOH solution. + + i. fCH2 ClCOOH(aq) + H2 O(l) à H3O (aq) + CH2 ClOO -(aq)
7. Given the following solubilities, determine the Ksp for each compound. -15 3 a) Bi2S3 1.70 x 10 mol/dm 3 b) Ca3(PO4)2 3.92 x 10-6 mol/dm
8. 3 Determine the mass of calcium fluoride, CaF2, that will dissolve in 100.0 dm of water. -11 Assume that there is no volume change. Ksp of CaF2 is 2.69 x 10
9. 3 Determine if a precipitate of silver chromate, Ag2CrO4, will form when 100.0 cm of 3 -12 0.100M AgNO3 are added to 100.0 cm of 0.350M K2CrO4. Ksp of Ag2CrO4 is 9.00 x 10
10. If a solution contains 1.00 x 10-2M chloride ions and 1.00 x 10-3M iodide ions, which will
precipitate first, AgCl or AgI, if a solution of AgNO3 is added one drop at a time? The Ksp -10 -16 of AgCl is 1.56 x 10 ; Ksp of AgI is 1.50 x 10
11. 3 In a laboratory experiment, 20.00 cm of NH3(aq) solution is titrated to the methyl orange endpoint using 15.65 cm3 of a 0.200M HCl solution. What is the concentration of the aqueous ammonia solution?
12. Suppose you mix 100.0 ml of 0.0200 M BaCl2 with 50.0 ml of 0.0300 M Na2SO4. Does -10 BaSO4 precipitate? ( Ksp= 1.1 X 10 )
13. -6 Calculate the solubility of PbBr2 in grams per liter given Ksp = 4.6 X 10
14. -5 .The molar solubility of Lead II Iodate, Pb(IO3)2 at 26 C is 4.0 X 10 . Determine the Ksp of Lead II Iodate. 15. The concentration of Calcium ion in blood plasma is 0.0025 M. If the concentration of Oxalate ion is 1 X 10-8 M. Will Calcium Oxalate, CaC2O4 precipitate? Ksp = 2.3 X 10-9.
8 | P a g e Study Guide and Lecture notes for Buffers, Ksp, Salt Hydrolysis Buffers Buffers are solutions that “resist” the change in pH, if a “small” amount of acid or base is added to it.
Importance of Buffer Action: - In our bodies H2CO3 / HCO3 buffer
Composition of Buffers Buffers are made of weak acid or weak bases and their conjugate base or acid. Buffers are a two- component system. They have Acidic component (this component helps neutralize the base) Basic Component (this component helps neutralize the acid)
Types of Buffer There are two types of buffers.
a. Acidic Buffer: Composed of weak acid and its conjugate base. Their buffering capacity is in - the acidic region. Ex: Acetic Acid/Acetate Buffer (H2CO3 / HCO3 )
b. Basic Buffer: Composed of a weak base and its conjugate acid. Basic buffers work best in + basic pH range. Ex. Ammonia/ Ammonium Buffer (NH3/ NH4 )
Action of buffer system
Acetic acid (CH3COOH) and Sodium acetate( NaCH3COO) system: In this example, acetic acid is the weak acid and sodium acetate is a salt of the weak acid with the acetate ion being a conjugate base of the weak acid. Let us see how a combination of acetic acid and acetate ion can resist a change in pH.
Addition of acid (H+) - + + CH3COO (aq) + H (or H3O )(aq) < ------> CH3COOH(aq) + H2O(l) Thus the conjugate base consumes the added acid.
HA A- Add H+ Acetic Ace Acid tate
Acidic and Basic Components of Acetic Acid/Acetate Buffer
9 | P a g e Addition of base (OH-) - - CH3COOH(aq) + OH (aq) < ------> CH3COO (aq) + H2O(l) Thus the acid consumes the added base.
Add OH- HA A-
Acidic and Basic Components of Acetic Acid/Acetate Buffer
Practice:
Given the NH3 and NH4OH system, explain how this system can act as a buffer by resisting change in pH?
Hydrolysis of salt solutions
Hydrolysis is the reaction between ions in a salt solution and water. For any salt, first determine the origin of the cation and the anion, example, for sodium chloride, the cation Na is from NaOH and the anion Cl is from HCl.
Second, determine if the acid and base are strong or weak Eg: NaOH is strong base HCl is a strong acid.
A solution of NaCl is” neutral” Acid Base PH of salt Strong Strong Neutral Strong Weak Acidic (0-7) Weak Strong Basic (7-14)
Weak Weak Depends on the Ka and Kb of acid and base
For each salt below, determine if the solution would be acidic, basic or neutral
a) CH3COONa b) NH4Cl
10 | P a g e c) KCl d) HCOONa e) K2SO4
Solubility Product
The solubility rules that you studies in chapter 14 give a qualitative account of solubility of ionic compounds but solubility product gives us a quantitative account of an ionic substance’s solubility. Precipitation/ dissolution reactions are very important in industry, medicine and our everyday life. Can you think of any examples?
For a sparingly soluble ionic compound, an equilibrium exists when this compound is dissolved in water, which is expressed by the following reaction:
+ - AgCl (s) Ag (aq) + Cl (aq)
K= [Ag +] [Cl-] , since AgCl is a solid, it will not be included in the equilibrium expression. [AgCl]
So, the equilibrium expression may be written as,
+ - Ksp = [Ag ] [Cl ]
Write Ksp expressions for MgF 2 and Ag2CO3.
Don’t confuse Ksp with solubility. Solubility of a substance is in mole/L or g/ml and shows how much of solid dissolves or dissociates in water, while Ksp shows the equilibrium between dissociated ions and undissolved solid. The higher the Ksp value, the greater the solubility of a substance.
Examples of Problems on Ksp
Calculating Ksp from solubility
1. Calculate Ksp for Ag2CrO4, if its solubility is 0.022 g/L. (Ans: 6.6 X 10^-5)
11 | P a g e Calculating Solubility given Ksp
0 2. Ksp for MgF2 is 6.4 X 10^-9 at 25 C. Calculate its solubility in mol/L and g/L. (Ans: 1.2 X 10^-3 M, 7.3 X 10^-2 g/L)
Predicting Precipitation, if two solutions are mixed (Q and Ksp Problems)
3. Does a precipitate form, if 100.0 ml of 0.0025 M AgNO3 and 150.0 ml of 0.0020 M NaBr are mixed? Ksp for AgBr= 5.0 X 10^-13
First find Q for AgBr and then compare if Q is bigger or smaller than Ksp. If Q is bigger than Ksp, then it precipitates and if Q is smaller than Ksp, then no precipitate happens.
Strong Acid and Strong Base Titrations
Observation Table
Titration 1 Titration 2 Acid Base Acid Base Final Volume (mL)
12 | P a g e Initial Volume (mL) Volume used (mL)
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