Chapter 23

Coordinate Covalent Bonding and

Lewis Acid Base Reactivity

Table of comparable charge density

Li^+ 1.7	Be^2+ 6.5
Na^+ 1.1	Mg^2+ 3.1										Al^3+ 6.0
K^+ 1.75	Ca^2+ 2.0		Ti^4+ 5.9	Cr^3+ 4.6	Mn^2+ 2.5	Fe^3+ 4.5	Co^2+ 2.4	Ni^2+ 2.9	Cu^2+ 2.9	Zn^2+ 2.7	Ga^3+ 4.8
Rb^+ .68	Sr^2+ 1.8	Y^3+ ^3.2	Zr^4+ 5.0						Ag^+ .79	Cd^2+ 2.1	In^3+ 3.7
Cs^+ .59	Ba^2+ 1.5	La^3+ 2.6						Pt^4+ 6.2		Hg^2+ 1.8	Tl^3+ 3.2

Transition metal ion	Hybridized state in complex	Number of ligands it attaches to	Geometric shape	Some ligands it typically bonds to
*Ag^+	sp	2	linear	CN^- NH3 Cl^-
*Al^3+	sp^3	4	tetrahedral	OH^-
Al^3+	sp^3d^2	6	octahedral	H2O
*Cr^3+	sp^3	4	tetrahedral	CN^- NH3 Cl^-OH^-
Cr^3+	sp^3d^2	6	octahedral	CN^- NH3 Cl^-
*Cu^2+	dsp^2	4	square planar	CN^- NH3 Cl^-
*Fe^3+	sp^3d^2	6	octahedral	SCN^- CN^- H2O
*Zn^2+	sp^3	4	tetrahedral	CN^- NH3 OH^-
Pt^4+	dsp^2	4	square planar	CN^- NH3 Cl^-
Pt^4+	sp^3d^2	6	octahedral	CN^- NH3 Cl^-

COMMON LEWIS BASES (LIGANDS)

Ligand	Name of Ligand
H2O	aquo
NH3	ammine
CN^-	cyano
Cl^-	chloro
OH^-	hydroxo
SO4^2-	sulfato
NO3^-	nitrato
C2O4^2-	oxalato
NO2^-	nitro
S2O3^-	thiosulfato

Iron	ferrate
Copper	cuprate
Tin	stannate
Silver	argentate
Lead	plumbate
Gold	aurate

Acid Base Comparative Strength

Acid	Formula	Conjugate base	Formula
Perchloric	HClO4	perchlorate ions	ClO4^-
Hydroiodic	HI	iodide ions	I^-
Hydrochloric Very	HCl	chloride ions Very	Cl^-
Hydrobromic Strong	HBr	bromide ions Weak	Br^-
Nitric	HNO3	nitrate ions	NO3^-
Sulfuric	H2SO4	hydrogen sulfate ions	HSO4^-
Hydronium ions	H3O^+	water	H2O
Sulfurous	H2SO3 (H2O + SO2)	hydrogen sulfite ions	HSO3^-
Hydrogen sulfate ions	HSO4^-	sulfate ions	SO4^2-
Phosphoric	H3PO4	dihydrogen phosphate ions	H2PO4^-
Hydrofluoric	HF	fluoride ions	F^-
Nitrous Moderate	HNO2	nitrite ions Weak	NO2^-
Acetic to Weak	HC2H3O2	acetate ions to Moderate	C2H3O2^-
Carbonic	H2CO3(H2O + CO2)	hydrogen carbonate ions	HCO3^-
Hydrogen sulfide	H2S	hydrogen sulfide ions	HS^-
Hydrogen sulfite ions	HSO3^-	sulfite ions	SO3^2-
Ammonium ions	NH4^+	ammonia	NH3
Hydrogen carbonate ions	HCO3^-	carbonate ions	CO3^2-
Hydrogen sulfide ions	HS^-	sulfide ions	S^2-
Water	H2O	hydroxide ions	OH^-
Hydroxide ions Very	OH^-	oxide ions Very	O^2^-
Ammonia Weak	NH3	amide ions Strong	NH2^-
Hydrogen	H2	hydride ions	H^-

RADICAL FORMULA	RADICAL NAME	ACID FORMULA	ACID NAME
ClO3^-	chlorate	HClO3	chloric acid
ClO4^-	perchlorate	HClO4	perchloric acid
ClO2^-	chlorite	HClO2	chlorous acid
ClO^-	hypochlorite	HClO	hypochlorous acid

23.2 Recognizing Lewis Acid Base Reactions

1. Draw the electron-dot structure for

a) NH₃ OH^- NH₄⁺ H₂O

b) Tell whether each can act as a Lewis base. Explain why or why not.

2. What is a

a) Lewis acid and how does it compare to a Bronsted acid?

b) Which of the following is not a Lewis acid?

Pt⁴⁺ HCl Na⁺ SO₂

3. Based on the examples shown in this unit use electron-dot structures to show the acid-base mechanism for the following reactions. Use an arrow to show the transfer of the Lewis acid to the lone pair of the Lewis base.

a) BF₃ + F^- ———> BF₄^- b) NH₄⁺ + OH^- ———> NH₃ + H₂O

c) BeF₂ + 2 F^- ———> BeF₄^2- d) CaO + CO₂ ———> CaCO₃

4. Which of the above acid-base reactions are (a) Bronsted type? b) Lewis type?

5. The structure of indium triiodide is shown below. It is evident from this structure that the molecular formula is In₂I₆.

a) Draw the Lewis electron dot structure from this model

b) What is the hybridization about In?

c) The In₂I₆ molecules dissociate in the gas phase into InI₃ molecules. Draw the Lewis structure for this molecule and indicate the hybridization about In.

d) What characteristic of the electronic structure of iodine in this compound is of importance in forming In₂I₆?

e) What characteristic of In in InI₃ is of importance?

6. Boric acid is a weak acid. The ions formed in a solution of H₃BO₃ in water are B(OH)₄^- and H₃O⁺. Show how the formation of these ions first involves a Lewis acid-base reaction and then a Bronsted acid-base reaction.

7. Predict the products, and their charge, of these Lewis acid-base reactions. The acid is not necessarily the first one listed. Products should have appropriate charges.

a) BH₃ + H^- b) MgO + CO₂

c) Au⁺ + 2 CN^- d) 2 F^- + GeF₄

8. Why would Li⁺ be regarded as a Lewis acid, but not Na⁺?

9. List two reasons why transition metal ions are Lewis acids.

10. Identify the Lewis acid and the Lewis base in each of the following reactions.

a) I₂ + I^- <——> I₃^-

b) Zn(OH)₂ + 2 OH^- <——> Zn(OH)₄^2-

c) Fe³⁺ + SCN^- <——> FeSCN²⁺

11. Ag⁺ will precipitate in the presence of Cl^- ion.

a) Write the equation for this process.

b) If excess Cl^- ions are added to the precipitate, a soluble complex ion forms. Write the equation for this.

c) Real neat question here: Compare Cl^- ion's strength in being a Bronsted base compared to a Lewis base.

23.3 Complex Ions or Coordination Compounds

23.4 Nomenclature of Complex Ions

12. What are

a) coordination compounds, or complex ions?

b) Why are transition metal ions capable of forming complex ions?

13. Write the balanced Lewis acid-base reaction for the formation of the complex ion containing

a) ammonia and zinc ion

b) iron(III) ion and thiocyanate ion

c) hydroxide ion and aluminum ion

d) cyanide ion and copper(II) ion

14. What is the term used to describe Lewis bases that are in coordination compounds, or complex ions?

15. Name these complex ions or coordination compounds

a) CuCl₄^2- b) [Al(H₂O)₆] Cl₃

c) Fe₄[Fe(CN)₆]₃ d) [CrCl₂(H₂O)₄]₃[Ni(CN)₆]

16. Write formulas for

a) hexaaquochromium(lll) chloride,

b) chloro-pentaaquochromium(lll)chloride,

c) sodium tetra-chlorodiaquochromate(lll),

17. What is the shape of each of these complex ions?

a) ZnCl₄^2- b) Pt(NH₃)₄²⁺

c) Fe(H₂O)₆²⁺ d) Ag(CN)₂^-

18. It is not clear at times whether chlorides are actually part of a complex ion. If they are, the association with the Lewis acid is covalent. If not, simply adding Ag⁺(aq) to the system will cause these loose chlorides to precipitate as AgCl. For instance, one mole of the Pt(NH₃)₂Cl₄ complex is added to excess AgNO₃(aq) and 2 moles of AgCl form. The actual formula of the complex salt must be

[Pt(NH₃)₂Cl₂] Cl₂

Determine the correct formula of these complex salts based on the following information:

a) CoCl₃(NH₃)₆ + 3 Ag⁺ ———> 3 AgCl + complex ion

b) Ni(H₂O)₆Cl₂ + 2 Ag⁺ ———> 2 AgCl + complex ion

19. What is the oxidation state of the transition metal ion, R, in each of the complex salts?

a) [R(NH₃)₄Cl₂ ][Ag(CN)₂] b) [Pt(H₂O)₆Cl₂] [R(CN)₄]₂

20. 4.736 g of BaCl₂•xH₂O is heated and the hydrate becomes the anhydrous BaCl₂, which weighs 4.038 g. Calculate the relative moles of H₂O (the x above). (ANS: 2)

23.5 The Bronsted-Lowry Acid Base Theory Definition

23.9 Naming Bronsted Acids

21. Draw the electron-dot structures for the species in these acid-base reactions and use arrows to show the proton transfer.

a) HI + H₂O ———> H₃O⁺ + I^-

b) NH₄⁺ + OH^- ———> NH₃ + H₂O

22. HPO₄^2- + OH^- ———>

HPO₄^2- + H₃O⁺ ———>

Complete the above reactions verifying the amphiprotic nature of HPO₄^2-. Is PO₄^3- amphiprotic?

23. Write formulas that would go in the blanks for the following acid-base neutralization reactions:

a) _________+_________ ———> H₂O + KNO₃

b) Ba(OH)₂ + _________ ———> BaSO₄ + _________

24. Tests indicate that formic acid, HCO₂H, is monoprotic, or donates only one proton. How do you explain this fact when there are two H's in the formula?

25. Name these using their acid names. The concept of periodicity can help you.

H₂Se H₃AsO₄ HBrO₃ H₂SeO₃

26. Write the formula for the following acids:

a) iodous b) telleric c) perbromic d) hydroastinic

27. Complete the following acid-base reactions assuming the first reactant is the acid:

a) HCO₃^- + OH^- b) NH₄⁺ + PO₄^3-

23.10 Bronsted acid-base reactions are equilibrium reactions

23.11 Relative Bronsted Acid-Base Strength

23.12 Strong Acids Are Leveled In Water

28. (a) How does a strong acid differ from a weak acid?

(b) How does a person determine which side is favored in an acid-base reaction? Why will every acid-base reaction favor the formation of the weak acid and base?

29 What is the correlation between the strength of an acid and the strength of its conjugate base?

30. Specify the conjugate acid-base pairs for the following equations:

a) NH₃ + HBr ———> NH₄⁺ + Br^-

b) NH₄⁺ + OH^- ———> NH₃ + H₂O

c) H₃O⁺ + PO₄^3- ———> HPO₄^2- + H₂O

d) HSO₃^- + CN^- ———> HCN + SO₃^2-

31. Write the formula (and charge, if necessary) for the conjugate base for each of these acids:

a) H₂O b) NH₄⁺ c) CH₄

32. Write the formula of the conjugate acid for each of these bases:

a) HSO₃^- b) C₆H₅N

33. Use the Relative Strength of Acids and Bases Table to predict (i) which reactant is the acid, (ii) the product formed, and (iii) whether the products or reactants are favored at equilibrium using appropriate-sized equilibrium arrows.

a) HI + OH^- b) H₂O + H₂SO₄ c) HSO₄^- + H₃O⁺

d) HSO₃^- + NH₃ e) CO₃^2- + HS^- f) HCO₃^- + F^-

34. NH₄⁺ + OH^- ———> NH₃ + H₂O

H₂O + C₂H₅O^- ———> C₂H₅OH + OH^-

Each of the reactions above proceeds essentially to completion to the right. Which is the stronger base, ammonia, NH₃, or ethoxide ion, C₂H₅O^-? Explain.

35. For the reaction

HX + B^- ———> HB + X^-

K = 100. Specify the strongest acid in the system.

36. Are any of the following species

a) leveled in water? Which ones?

HNO₃ NH₄Cl NaC₂H₃O₂ NH₄NH₂ Na₂O

b) For those that are leveled, write the equation that shows what it is leveled to. Why are they leveled?

37. Certain reactions require the use of extremely strong acids. Why are these reactions not carried out in an aqueous environment?

38. Weak bases are tested with 100% H₂SO₄, not H₂O. Why?

23.14 Salt Dissociation and Ionization Can Affect Acidity & Alkalinity

23.15 Molecular Dissolution With No Ionization

39. List three salts that, when added to water, will not create an acidic nor a basic solution.

40. Given the following salts:

Na₂CO₃ (NH₄)₂SO₄ CrCl₃ K₂S NaClO₄ KHCO₃

a) Circle any spectator ions.

b) List anions that are capable of combining with H⁺ to form strong acids.

c) List anions that are capable of combining with H⁺ to form weak acids

d) Write equations for those salts, (without the spectator ions) that create a basic solution when dissolved in water. Show appropriate-sized equilibrium arrows.

e) Write equations for those salts (without the spectator ions) that create an acidic solution when dissolved in water. Show appropriate-sized equilibrium arrows.

41A. All of the following salts dissolve in water. Some react further to create either an alkaline environment, some react further to create an acidic environment.

Directions: first write the dissociation in water for each salt. Second, decide if either the cation or anion reacts further with water. If it does write this second equation with appropriate sized equilibrium arrows. If no further reaction occurs write NR.

a) KClO₄ b) NaC₂H₃O₂ c) NaCN d) Rb₂SO₄ e) MgSO₃

f) NaF g) K₂S h) Ca(NO₃)₂ i) MgI₂ j) NaNO₃

41B. All of the following salts dissolve in water. Some react further to create either an alkaline environment, some react further to create an acidic environment.

Directions: first write the dissociation in water for each salt. Second, decide if either the cation or anion reacts further with water. If it does write this second equation with appropriate sized equilibrium arrows. If no further reaction occurs write NR.

a) Na₂CO₃ b) K₂SO₃ c) NH₄Br d) Mg(NO₂)₂ e) Al(NO₃)₃

f) Ca(C₂H₃O₂)₂ g) CsF h) CsI i) (NH₄)₂SO₄ j) Al(ClO₄)₃

42. NaF and NaC₂H₃O₂ are salts capable of creating a basic environment when dissolved in water.

a) Write the equation that shows this.

b) Check to see which acid, HF or HC₂H₃O₂, is stronger.

c) Which salt is likely to produce the more basic solution when dissolved in water, NaF or NaC₂H₃O₂? Explain.

43. A solution of NaCN is more basic than a NaNO₂ solution. Which acid is stronger, HCN or HNO₂? Explain, using equations.

44. Salt AB₃ creates an acidic environment when dissolved in water.

a) Write two equations to show how this happens.

b) Salt CD creates a basic environment when dissolved in water. Write an equation to show how this happens.

45. The salt NaC₂H₃O₂ dissociates completely in water to make a 0.1 M solution. A further reaction occurs which makes the [OH^-] = 7.5 x 10^-6-M.

a) Write the reaction occurring after dissociation and

b) Calculate the K value for this equation.

46. A 0.3 M solution of the acid HC₂H₃O₂ ionizes 1.0% in water at a particular

HC₂H₃O₂(aq) + H₂O ———> C₂H₃O₂^- (aq) + H₃O⁺(aq)

Calculate the K value for this ionization.

23.16 Acid Ionization

23.17 Polyprotic Acids

47. How is the strength of the bond holding hydrogen to a molecule related to that molecule's acid strength?

48. For each of the following pairs donates a proton more successfully?

a) H₃PO₄ or H₃AsO₄ b) H₂SO₄ or H₂SO₃ c) H₂O or H₂S

Account for your selection in each case, a - c, using Lewis electron-dot structures.

49. Draw the Lewis electron-dot structure for the following acids: HBrO H₂SO₄

50. How does the concept of electronegativity explain why HBr is a stronger acid than HCl?

51. Which acid produces the weaker conjugate base in each pair?

a) HCl or HI b) HClO₄ or HClO₃ c) H₂SO₄ or H₂SO₃

52. Place the species in each of the following groups in order of increasing base strength.

a) BrO^-, BrO₂^-, BrO₃^- b) H₂PO₄^-, HPO₄^2-, PO₄^3-

53. Arrange the following in order of acid strength from weakest to strongest:

a) CH₃CO₂H, FCH₂CO₂H, F₂CHCO₂H, F₃CCO₂H

b) NH₄⁺, CH₃NH₃⁺, HONH₃⁺

54. HX is a stronger acid than HY.

a) Draw oval shapes for both molecules. Label one side H, the other X (or Y). Draw a net polarity (vector) in each, of respective size.

b) Which molecule has the higher dipole moment?

c) Which molecule has stronger intermolecular forces?

d) Which compound has the higher boiling point?

55. Write the

a) 3 ionization equations for the acid H₃AsO₄.

b) Which ionization (1, 2, or 3) produces the largest quantity of H₃O⁺. Explain.

23.18 A Specific Look at Substances With an O-H Combination and The Degree of Ionization

56. For the following compounds write equations to show how they dissociate into ions. Write ND if no dissociation and ionization occurs. Briefly account for your answer in each case.

a) RbOH b ) BrOH c) Cu(OH)₂ d) CH₃OH

57. Which ionization in water is more successful in each pair?

a) HCl or HI b) NaOH or Ca(OH)₂

c) H₂SO₄ or H₂SO₃ d) HClO or HBrO

e) Ca(OH)₂ or C₂H₅OH

“If I have seen farther than others, it is because I have stood on the shoulders of giants.”

Sir Isaac Newton (1642-1727)

Chapter 24 -Chemical Reactions

The Solubility Rules For Salts In Water

Substance	Solubility in water	Exception
1. The compounds formed with cations from Group Ia element (Li^+, Na^+, K^+, Rb^+, Cs^+)	dissociates, dissolves, forms aqueous (aq) solutions
2. The compounds formed with the ammonium ion, NH4^+	dissociates, dissolves, forms aqueous (aq) solutions
3. The compounds formed with the nitrate, NO3^-, ion	dissociates, dissolves, forms aqueous (aq) solutions
4. The compounds formed with Cl^–, Br^–, or I^–	dissociates, dissolves, forms aqueous (aq) solutions	Hg^+, Pb^2+, or Ag^+ are insoluble
5. The compounds formed with sulfate, SO4^2-	dissociates, dissolves, forms aqueous (aq) solutions	Ba^2+, Pb^2+, Sr^2+, or Ca^2+ are insoluble
6. The compounds formed with carbonate, CO3^2-, phosphate, PO4^3-, sulfide, S^2–, oxide, O^2–	Insoluble	Group Ia metals and ammonium are soluble
7. The compounds formed with hydroxide, OH^–	Insoluble	Group Ia metals and barium are soluble

CATEGORY OF STABLE SPECIES EXAMPLE

Gases CO₂, H₂, SO₂, NH₃, H₂S

Water H₂O

Weak Acids H₂S, HC₂H₃O₂ , HCN, H₂O

Weak bases NH₃, Cl^-, H₂O

Water-insoluble salts CaCO₃, AgCl

Complex ions Ag(NH₃)₂⁺, Cu(CN)₄^2-

Certain Metals Au, Pt, Ag

Network Substances C(diamond)

24.1 (Chemical Reactions) Review of Principles

1. Write equations for the following but do not balance. Use (aq), (s), (g), (l) for each substance.

a) A strip of magnesium is added to a solution of silver nitrate.

b) Excess hydrochloric acid is added to a solution of potassium sulfite.

c) Ethene, C₂H₄, is burned in air.

d) Solutions of zinc sulfate and sodium phosphate are mixed.

e) A stream of chlorine gas is bubbled through a solution of strontium iodide.

f) A piece of sodium is added to water.

g) Solutions of lead(II)nitrate and calcium chloride are mixed.

h) Chlorine gas is bubbled through a solution of potassium bromide.

i) Butanol, C₄H₉OH, is heated in air.

j) Nitrogen gas is directed over calcium metal.

k) A piece of zinc is dropped in a solution of copper(II)bromide.

l) Solutions of barium chloride and sodium sulfate are mixed.

m) Hot oxygen gas is directed over solid copper(II)sulfide.

n) A 0.1 M solution of hydrochloric acid is mixed with a 0.1 M solution of sodium phosphate.

o) Hydrogen peroxide is heated.

p) Magnesium sulfite is heated until it decomposes.

q) A solution of dilute nitric acid is poured over chunks of magnesium carbonate.

r) A solution of sodium hydroxide is mixed with a solution containing ammonium chloride.

s) Excess chlorine gas is passed over iron filings.

24.2 Chemical Processes That Are Spontaneous

24.3 Reactivity That Produces Stable Species That Are Gaseous, Aqueous Molecular, or Soluble Complexes

24.4 Adding a Strong Aqueous Acid To Produce a Weaker Acid

24.5 Addition of Aqueous Hydroxide Compounds

24.6 Complexation

2. For the following solid substances add either strong acid [HCl] or strong alkali [NaOH] or a complexing agent [NH_3, CN^-, OH^-]. Then complete the equation.

a) Na₂CO₃ b) AgBr c) K₂SO₃ d) CuSO₄ e) NaF f) Zn(OH)₂

g) NaC₂H₃O₂ h) Rb₂S i) NH₄I j) Al(OH)₃

3. For the following solid salts (i) tell what aqueous solution must be added to dissolve the salt and (ii) write the equation. A stable species is produced in each case.

a) magnesium carbonate b) calcium hydroxide

c) silver bromide d) ammonium sulfate

e) chromium(III)hydroxide (in this case don’t use acid)

f) sodium acetate g) iron(II)chloride

h) potassium sulfite i) copper(II)chloride

j) zinc chloride k) sodium sulfide

24.7 Writing Net-Ionic/Molecular Equations

4. For the following specify what you would add to produce a stable species. Write the products with (aq), (g), etc. Then write the net-ionic/molecular equation. Do not use water as a reagent. Be sure that ionic species have the correct charge. Remember to not dissociate stable species.

a) copper(II)sulfate(s) b) calcium hydroxide(s)

c) silver chloride(s) d) magnesium sulfite(s)

e) ammonium bromide(aq) f) sodium sulfide(aq)

g) potassium carbonate (aq) h) zinc hydroxide(s)—don't use acid here

5. What liquid aqueous reagent would you add to each of the following mixtures that would dissolve one compound into a new stable compound but not the other? Write the equation.

a) NaCl(s) and CaSO₄(s) b) Al(OH)₃(s) and Ca(OH)₂(s)

c) AgCl(s) and PbCl₂(s) d) CaCO₃(s) and CaSO₄(s)

e) KCl(s) and NH₄Cl(s)

6. You choose to investigate some of the solubility rules for two ions, the chromate ion, CrO₄^2-, and the oxalate ion, C₂O₄^2-. You are given solutions (A, B, C, D) of four water-soluble salts:

Solution	Salt solution	Color of solution
A	Na2CrO4	Yellow
B	(NH4)2C2O4	Colorless
C	AgNO3	Colorless
D	CaCl2	Colorless

When these solutions are mixed, the following observations are made:

Expt. number	Solns mixed	Result
1	A + B	No precipitate, yellow solution
2	A + C	Red precipitate forms
3	A + D	No precipitate, yellow solution
4	B + C	White precipitate forms
5	B + D	White precipitate forms
6	C + D	White precipitate forms

a) Write the net ionic equation when precipitation occurs.

b) Circle the precipitate formed, if any, in each of the experiments.

c) Based on these limited observations, which ion tends to form the more soluble salts, chromate or oxalate?

7. What chemical (write its complete formula) would you add to make the substance in column A become that in column B?

A B

a) BaCl₂(aq) BaSO₄(s)

b) CuCO₃(s) Cu(NO₃)₂(aq)

c) CaBr₂(aq) CaCl₂(aq) + Br₂(l)

8. If AlCl₃ is mixed with an aqueous solution of ammonia, NH₃, a precipitate of Al(OH)₃ is formed. Why?

9. Is an aqueous solution of NaHSO₄ acidic, neutral, or basic? What reaction with water occurs? If solid Na₂CO₃ is added to a solution of NaHSO₄, what reaction can occur between the CO₃^2- and HSO₄^- ions?

24.8 Qualitative Analysis

10. A solution containing an unknown number of metal ions is treated with dilute HCl; no precipitate forms. The pH is adjusted to about 1, and H₂S is bubbled through. Again no precipitate forms. The pH of the solution is then adjusted to about 8. Again H₂S is bubbled through. This time a precipitate forms. The filtrate from this solution is treated with sodium carbonate. No precipitate forms. Which metal ions from the flowchart are possibly present? Which are definitely absent within the limits of these tests?

11. A student who is in a great hurry to finish his laboratory work decides that his qualitative analysis unknown contains a metal ion from the insoluble carbonate group, group 4 from the flowchart. He therefore tests his sample directly with Na₂CO₃, skipping earlier tests for the metal ions in groups 1 through from the flowchart. He observes a precipitate and concludes that a metal ion from group 4 is indeed present. Why is this possibly an erroneous conclusion?

12. You are given unlabeled samples of NaCl and NH₄Cl (both are white) and asked to distinguish them from each other. You do it in a single step, adding only one reagent. What is the reagent?

13. Cu²⁺ and Zn²⁺ are present in a solution. Describe a short procedure separating these ions from solution.

14. A Group I unknown from the flowchart gives a white precipitate with HCl which is completely soluble in excess NH₃. Identify the unknown.

15. A solution upon treatment with HCl or with H₂S at pH 9 gives no precipitate. What, if anything, does this indicate about the cations that might be present (from the flowchart) in this solution?

16. A Group III unknown contains only Ni²⁺ and Al³⁺. It is treated with aqueous ammonia to give a colored precipitate. As more NH₃ is added part of the precipitate dissolves to form a deep blue solution. The precipitate remaining goes into solution when treated with excess NaOH. If acid is slowly added to this solution, a white precipitate forms, which dissolves as more acid is added. Write balanced equations for each reaction that took place.

17. Hydrochloric acid does not dissolve glass. By contrast, hydrofluoric acid, HF(aq), is a weak acid that dissolves glass and must therefore be stored in plastic containers.

a) Do you think that the ability of HF to dissolved glass has anything to do with its ionization in water and production of hydronium ions?

b) Which reaction will be more successful for 0.5 M solution of the acids:

Mg(s) + HCl(aq) à MgCl₂(aq) + H₂(aq)

Mg(s) + HF(aq) à MgF₂(aq) + H₂(aq)

Chapter 25

Effect Of Dissolved Reagents on pH

Numerical Evaluation of The Chemical Equilibrium of Ions and Molecules In Aqueous Solution

DISSOCIATION CONSTANTS FOR SELECTED WEAK ACIDS (25°C)

Acid	Reaction		Conjugate base	Ka
Iodic	HIO3 + H2O —>	H3O^+ + IO3^-	iodate	1.6 x 10^‑1
Benzoic	HC6H5CO2 + H2O —>	H3O^+ + C6H5CO2^-	benzoate	6.6 x 10^‑3
Hydrofluoric	HF + H2O —>	H3O^+ + F^-	fluoride	6.8 x 10^‑4
Nitrous	HNO2 + H2O —>	H3O^+ + NO2^-	nitrite	5.1 x 10^‑4
Formic	HCOOH + H2O —>	H3O^+ + COOH^-	formate	2.0 x 10^‑4
Acetic	HC2H3O2 + H2O —>	H3O^+ + C2H3O2^-	acetate	1.8 x 10^‑5
Hypochlorous	HOCl + H2O —>	H3O^+ + OCl^-	hypochlorite	3.0 x 10^‑8
Boric	H3BO3 + H2O —>	H3O^+ + H2BO3^-	dihydrogen borate	5.8 x 10^‑10
Ammonium ion	NH4^+ + H2O —>	H3O^+ + NH3	ammonia (molecules)	5.6 x 10^‑10
Hydrocyanic	HCN + H2O —>	H3O^+ + CN^-	cyanide	4.0 x 10^‑10
Water	H2O + H2O —>	H3O^+ + OH^-	hydroxide	1.0 x 10^‑14

DISSOCIATION CONSTANTS FOR SELECTED WEAK BASES

Name	Reaction	Dissociation Constant, Kb, 25°C
Trimethylamine	(CH3)3N + H2O <——> (CH3)3NH^+ + OH^-	6.5 x 10^‑5
Ethanolamine	HOC2H4NH2 + H2O<——> HOC2H4NH3^+ + OH^-	3.2 x 10^‑5
Ammonia	NH3 + H2O <——> NH4^+ + OH^-	1.8 x 10^‑ 5
Hydrazine	N2H4 + H2O <——> N2H5^+ + OH^-	1.7 x 10^‑6(20°C)
Hydroxylamine	HONH2 + H2O <——> HONH3^+ + OH^-	1.1 x 10^‑8(20°C)
Pyridine	C5H5N + H2O <——> C5H5NH^+ + OH^-	1.8 x 10^‑9
Aniline	C6H5NH2 + H2O <——> C6H5NH3^+ + OH^-	4.3 x 10^‑10

TRANSITION RANGE FOR VARIOUS INDICATORS

Indicator pH Transition Range Acid Color Base Color

Methyl violet 0.5-1.6 yellow blue

Thymol blue 1.2-2.8 red yellow

Thymol blue 8.0-9.6 yellow blue

Methyl orange 3.2-4.4 red yellow

Bromcresol green 3.8-5.4 yellow blue

Methyl red 4.8-6.0 red yellow

Chlorophenol red 5.2-6.8 yellow red

Bromthymol blue 6.0-7.6 yellow blue

Phenol red 6.6-8.0 yellow red

Neutral red 6.8-8.0 red yellow-orange

Phenolphthalein 8.2-10.0 colorless magenta

Thymolphthalein 9.4-10.6 colorless blue

Alizarin yellow 10.1-12.0 yellow red

25.2 Preliminary Definition and Calculations

1. Perform the following calculations:

a) [H₃O⁺] = 7.39 x 10^-2 M find pH. b) [OH^-] = 2.14 x 10^-9 M find pOH.

c) [H₃O⁺] = 6.32 x 10^-6 M find [OH^-]. d) [OH^-] = 4.78 x 10^-1 M find [H₃O⁺]

e) pH = 7.22 find [H₃O⁺]. f) pH = 4.31 find pOH.

g) pOH = 7.37 find pH. h) pH = 9.23 find [OH^-] .

i) [H₃O⁺] = 2.66 find pOH.

25.3 The Equilibrium Created By Dissolving Strong Acids and Bases In Water

25.4 The Equilibrium Created By Dissolving Weak Acids or Bases In Water

2. Comment on the relative size of K for the ionization of each of these acids in water. Write the equation, too.

a) HNO₃ b) HC₂H₃O₂

3. In an equilibrium mixture of each of the above acids in water, comment on the relative amount of molecules versus ions in each case.

4. What assumption can be made about the [H₃O⁺] of a strong acid in water?

5. Calculate the pH of the following solutions:

a) 0.0020 M HClO₄ b) 1.5 x 10^-4 M KOH

c) 0.040 g NaOH dissolved in 2.0-liters of water

d) 74.0 mg of calcium hydroxide dissolved in 750.0 ml of solution (assuming 100% dissociation)

e) Dilution of 1.0 ml of 0.20 M HCl to a total volume of 5.0 L.

f) The result of mixing 50.0 ml of 0.500 M HCl and 50.0 ml of 0.750 M NaOH.

g) The result of bubbling 6.37 liters of HCl gas at 25.0°C and 1.21 atm into 0.750-liter of water. Assume all the HCl ionizes.

6. The DH for the dissociation of water has a positive value. At 25°C, the K_w for water is 1 x 10^-14.

a) Is water dissociated to a greater or less degree at 100°C than at 25°C?

b) Is K_w higher or lower at 100°C?

c) Is the pH of boiling water greater or smaller than 7?

d) Is boiling water acid, basic, or neutral?

7. A 0.100 M solution of the weak acid HC₂H₃O₂ is 1.30% ionized in water. What is the pH of this solution?

8. Why can’t a solution of HCl be made up to have a pH = 10.0? .

9 . Calculate the mass of solid NaOH required to prepare 0.500 liter of NaOH with a pH = 12.0.

10. For the acid HF,

a) Write the ionization equation and comment on the degree of ionization in water.

b) Does a 0.100 M solution of HF have a pH = 1 (i.e. use the formula

pH = -log(0.1)?

c) Comment on the amount of HF molecules compared to H₃O⁺ in a 0.100 M solution.

d) What is the ionization constant for HF at 25°C?

11. Calculate the pH of the following solutions (Use tables in this unit):

a) 1.0 M HF b) 0.10 M C₆H₅NH₂, aniline c) 0.001 M HNO₃

12. A 1.0 x 10^-3 M solution of a weak acid, HX, is ionized 2.10% in water at 25°C. Calculate

a) [H₃O⁺] and [X^-] of this solution.

b) the ionization constant for this acid.

c) pH of the solution.

13. A solution of hypochlorous acid, HClO, has a pH equal to 4.2 at equilibrium.

a) Calculate the [H₃O⁺].

b) Calculate the [ClO^-].

c) Calculate the [HClO].

d) Calculate the ratio of conjugate base to acid in the solution.

e) Determine the percent ionization of this acid in the solution.

14. A solution of HF has a pH = 2.2. Calculate the grams of HF in 250.0 ml of this solution.

15. A 0.10 M solution of benzoic acid, HC₆H₅CO₂, is prepared. Calculate the

a) H₃O⁺ concentration,

b) percentage ionization of this acid in water.

c) ratio of conjugate base to acid in this solution?

16. KCN is a salt that dissolves in water.

a) Comment on whether the solution is acidic, neutral, or basic.

b) Write an equation for the reaction of CN^- ion in water.

c) The K value for this equation is 2.5 x 10^-5. Calculate the pH of a 0.10 M KCN solution.

25.5 The Relationship Between K_b and K_a

17. The K_a for HNO₂ is 5.1 x 10^-4

a) Write the ionization equation for HNO₂.

b) Write the ionization equation for the conjugate base and determine the K_b.

18. Use tables in this unit to determine the K as indicated for the reactions

a) Kb for IO₃^- + H₂O <———> HIO₃ + OH^-

b) Kb for CN^- + H₂O <———> HCN + OH^-

c) Ka for N₂H₅⁺ + H₂O <———> N₂H₄ + H₃O⁺

Calculate the equilibrium constants for these reactions at 25°C using the tables in this unit. The concept of additivity of reactions will be helpful.

d) N₂H₅⁺ + F^- <———> HF + N₂H₄

e) HOCl + C₅H₅N <———> C₅H₅NH⁺ + OCl^-

25.6 The Equilibrium Created By Salts That Produce H₃O⁺ or OH^- In Water

19. Calculate the pH for the solution that is

a) 1.0 M in NaF b) 1.0 M in HF.

20. Calculate the resulting pH after the following salts are dissolved in water to the specified concentration: (Assume 100% dissociation).

a) 0.50 M NaC₆H₅CO₂ (or NaC₆H₄COOH, sodium benzoate).

b) 0.10 M AlCl₃ Ka = 6.3 x 10^-6

c) 1.0 M NH₄NO₃

d) 1.0 M NaCl

21. It is found that 0.10 M solutions of the three salts — NaX, NaY, and NaZ have a pH of 7, 8, and 9, respectively.

a) Write the ionization equations for the reacting ion in each salt with appropriate-sized equilibrium arrows.

b) Which ion is the stronger base?

c) Arrange the acids — HX, HY, and HZ — in order of acid strength.

22. A 0.25 M solution of the salt NaX has a pH = 10.0.

a) Calculate Kb for X^-.

b) Calculate the K for the reaction: HX + H₂O ———> H₃O⁺ + X^-

23. Write the equation for the neutralization of

i) HCl & NaOH ii) HC₂H₃O₂ & NaOH

24. What is the pH of a solution after

i) 500.0 ml of 0.50 M HCl is neutralized with 500.0 ml of 0.50 M NaOH?

ii) 500.0 ml of 0.50 M HC₂H₃O₂ is neutralized with 500.0 ml of 0.50 M NaOH?

25. A 0.50 M KNO₂ solution is prepared. Calculate the ratio of molecular acid to its conjugate base in the solution.

25.7 Solutions With Acids and Their Conjugates

26. A bottle contains concentrated formic acid which is 88% (i.e. this means it is 88 g HCOOH/100 g solution). Its specific gravity is 1.29. Its formula weight is 46.03 g/mole.

a) Calculate the molarity of the concentrated acid.

b) A chemist wishes to make up 1.0 liter solution of 0.10 M formic acid. What volume of the concentrate must be used?

c) Calculate the HCOOH, COOH^-, H₃O⁺ concentrations and the pH of the 0.10 M HCOOH solution.

d) 8.5 g of sodium formate, NaCOOH, crystals are added to 250.0 ml of the above solution without affecting its volume. Calculate the parameters asked for in (c).

25.8 Buffer Solutions — Definition and Reactions

25.9 Buffer Problems

27. 1.0 liter of a 0.50 M HNO₂ solution was made.

a) Write the ionization equation for HNO₃.

b) What must be added to make this solution a buffer?

c) Write the reaction that occurs should HCl be added to the buffer above the ionization equation from (a).

d) Write the reaction that occurs should NaOH be added to the buffer above the ionization equation from (a).

e) Do the same as in a-d for 1.0 liter of 0.40 M HONH₂, a weak base.

28. What is the function

a) of a buffer?

b) What effect on pH does the addition of H₂O have on a buffer?

29. Classify each of the following solutions as i) a good buffer, ii) a weak buffer, iii) not a buffer at all:

a) 10^-5 M HC₂H₃O₂ + 10^-5 M NaC₂H₃O₂ b) 1.0 M HCl + 1.0 M NaCl

30. A 1.0-liter solution containing 0.50 M HC₂H₃O₂ and 0.50 M C₂H₃O₂^-.

a) Calculate its pH.

b) Calculate its pH after 0.10 mole of HCl is added to it and mixed. Assume the reaction with HCl is 100% successful.

31. A buffer is prepared where formic acid, HCOOH, is 0.50 M as is its conjugate base, COOH^-. Calculate the pH

a) of the buffer.

b) after 0.05 mole of HCl is added to 1 liter of the buffer.

c) after 0.09 mole of NaOH is added to 1 liter of the original buffer.

32. A 2.0 M solution of the weak base hydrazine, N₂H₄, is prepared. 500.0 ml of it is mixed with 500.0 ml of 1.0 M HCl.

a) Calculate the pOH of the resulting solution.

b) To 500.0 ml of the resulting solution 500.0 ml of 0.20 M N₂H₅⁺ is added. Calculate the pH after mixing.

33. The equations and constants for the dissociation of three different acids are given below.

HCO₃^- <———> H⁺ + CO₃^2- Ka = 4.2 x 10^-7

H₂PO₄^- <———> H⁺ + HPO₄^2- Ka = 6.2 x 10^-8

HSO₄^- <———> H⁺ + SO₄^2- Ka = 1.3 x 10^-2

a. From the systems above, identify the conjugate pair that is best for preparing an ideal buffer with a pH of 7.2. Explain your choice.

b. Explain briefly how you would prepare the buffer solution described in (a) with the conjugate pair you have chosen.

c. If the concentrations of both the acid and the conjugate base you have chosen were doubled, how would the pH be affected? Explain how the capacity of the buffer is affected by this change in concentrations of acid and base.

d. Explain briefly how you could prepare the buffer solution in (a) if you had available the solid salt of only one member of the conjugate pair and solutions of a strong acid and a strong base.

e. For any ideal buffer

i) why is the [H₃O⁺] equal to the K_a of the weak acid.

ii) If the pH of a buffer is 4.83, what is the K_a of the weak acid, HA? Assume it is an ideal buffer.

34. We need a buffer with pH 9.00. It is to be prepared from NH₃ and NH₄Cl. What should be the ratio of [NH₄⁺] to [NH₃] in the buffer?

35. What is the resulting pH when 4.00 grams of NaOH is dissolved in 1.0 liter of a solution containing both 0.50 M NH₃ and 0.50 M NH₄Cl?

36. Enough acid is added to a buffer to change the acid/base ratio by a factor of 2. How much does the pH change by?

37. 250.0 ml of a 1.0 M solution of NaNO₂ is mixed with 250.0 ml of 0.50 M HCl. The system is mixed. Then, 500.0 ml of 0.20 M NaOH is added and mixed. What is the resulting pH?

25.11 Titration — A Stoichiometric Evaluation

38. What is the purpose of a

a) titration?

b) What purpose does the indicator serve in a titration?

39. 30.00 ml of an HCl solution is titrated to an end point with 48.38 ml of 0.1013 M NaOH. Calculate the molarity of the acid solution.

40. How many milliliters of 0.100 M HCl are required to neutralize 25.0 ml of 0.100 M Ba(OH)₂?

41. A HCl solution is standardized using pure Na₂CO₃ as the primary standard. 0.5896 g of Na₂CO₃ is dissolved in water, indicator added, and then titrated with the HCl solution. 29.34 ml of the acid is required to reach the end point. Calculate the molarity of the HCl.

42. 0.8432 g of an impure sample of potassium biphthalate, HKC₈H₄O₄, is titrated to an end point with 21.06 ml of 0.0978 M NaOH.

HKC₈H₄O₄ + NaOH———> H₂O + NaKC₈H₄O₄

Calculate the fraction of the HKC₈H₄O₄ in the sample.

43. A 3.094-g sample of the monoprotic acid, HA, was dissolved in distilled water. 72.23 ml of 0.4862 M NaOH was required to neutralize the acid. Calculate the molecular weight of HA.

44. A 0.500‑gram sample of a weak, nonvolatile acid, HA, was dissolved in sufficient water to make 50.0 milliliters of solution. The solution was then titrated with a standard NaOH solution. Predict how the calculated molar mass of HA would be affected (too high, too low, or not affected) by the following laboratory procedures. Explain each of your answers.

a) After rinsing the buret with distilled water, the buret is filled with the standard NaOH solution; the weak acid HA is titrated to its equivalence point.

b) Extra water is added to the 0.500‑gram sample of HA.

c) An air bubble passes unnoticed through the tip of the buret during the titration.

45. Federal regulations set an upper limit of 50 parts per million (ppm) of NH₃ in the air in a work environment (that is, 50 mL NH₃ per 10⁶ mL of air). The density NH₃(g) at room temperature is 0.771 g/L. Air from a manufacturing operation was drawn through a solution containing 100 mL of 0.0105 M HCl. The NH₃ reacts with HCl as follows:

NH₃(aq) + HCl(aq) ——> NH₄Cl(aq)

After drawing air through the acid solution for 10.0 min at a rate of 10.0 L/min, the acid was titrated. The remaining acid required 13.1 mL of 0.0588 M NaOH to reach the equivalence point.

a) How many grams of NH₃ were drawn into the acid solution?

b) How many ppm of NH₃ were in the air?

c) Is this manufacturer in compliance with regulations?

46. An impure sample of Ba(OH)₂ weighing 0.500 g was placed in a flask. 50.00 ml of 0.100 M HCl was added to the flask to be sure that the HCl was in excess. To find out how much of the HCl was in excess, the flask was titrated to an end point with 7.50 ml of 0.200 M NaOH. Calculate the percentage of Ba(OH)₂ in the impure sample.

25.12 Titration — Monitoring With a pH Meter

47. The following are titration curves

a) Which curve represents a titration of a

•weak base with HCl titrant?

•strong acid with NaOH titrant?

•weak acid with NaOH titrant?

•strong base with HCl titrant?

b) How is a titration for a weak acid different from that of a strong acid in terms of the pH at the beginning of the titration?

c) Answer the same question for a weak base vs a strong base.

For the above curves

d) specify the approximate pH of the equivalence point.

e) specify which indicator would be suitable in each case. Use the table in this unit.

48. What data are needed to perform the following calculations associated with a titration curve?

a) Molarity of titrated substance.

b) Dissociation constant (K_a or K_b) of the titrated solution.

49. The following is the titration curve of 50.0 ml of NH₃, a weak base, by the strong acid, HCl.

a) Use the graph to calculate the molarity of the NH₃ solution.

b) Calculate the K_b for NH₃ using the graph.

c) Calculate the K_b for NH₃ instead at that pH point when 10.0 ml of HCl had been used.

d) What would happen if an indicator that changed color at pH 10 were used?

50. 20.0 ml of a 0.100 M N₂H₄, a weak base, is titrated to the equivalence point with 20.0 ml of 0.100 M HCl. Calculate the pH at this point.

51. 40.0 ml of HA, a weak acid, is titrated against a 0.400 M NaOH solution. The original HA solution has a pH of 2.3. The addition of 6.0 ml of NaOH raises the pH to 3.5. The equivalence point of the titration occurs when 20.0 ml of the NaOH is added.

a) Calculate the molarity of the acid.

b) Calculate the Ka for the acid.

i) before the titration started.

ii) when 6.0 ml of NaOH had been added.

c) Calculate the pH at the equivalence point.

52. The following observations are made about acids. Discuss the chemical processes involved in each case. Use chemical equations to elucidate.

a) When 10 milliliters of 0.10-molar H₂SO₄ is drained from a buret into an Erlenmeyer flask containing 40 milliliters of 0.10-molar NaOH, the pH changes only by about 0.5 unit. After 10 more milliliters of 0.10-molar H₂SO₄ is added, the pH changes by about 6 units upon addition of the last few drops.

b) The pH of 0.10 M HNO₃ is 1 but that of 0.10 M HNO₂ is higher than 1.

c) The first dissociation constant for H₂S in water is larger than the second.

53. The DG° for the reaction is 49.1 kJ.

HA + H₂O ———> H₃O⁺ + A^-

a) Calculate the pH for a 1.0 M solution of HA. Is it weak or strong?

b) How does the DG° for the ionization of HCl compare to the 49.1 kJ value for HA?

54. The molar solubility of PbCl₂ at 25°C is 1.59 x 10^-2-M.

a) Calculate the DG° for the dissociation of this salt in 25°C water.

b) In what way would the DG° for the dissolution of NaCl be different from PbCl₂’s?

55. Which of the following bar graphs (i, ii, iii, or iv) is representative of a

a) buffer preparation? b) strong acid equilibrium?

c) titration? d) weak acid equilibrium?

56. The following plots pH versus volume of base NaOH for the titration of various acids against NaOH.

Order the acids (HG to HA) from strongest to weakest.

Answers to Numerical Questions

1a) 1.13 b) 8.67 c) 1.58 x 10^-9 M d) 2.09 x 10^-14 M 1e) 6.03 x 10^-8 M f) 9.69

g) 6.63 h) 1.70 x 10^-5 M 1i) 14.425 5a) 2.7 b) 10.2 c) 10.70

d) 11.43 5e) 4.40 f) 13.1 g) 0.38 7) 2.89 9) 0.20 g

11a) 1.58 b) 8.82 c) 3

12a) 2.1 x 10^-5 M b) 4.5 x 10^-7 M c) 4.68 13a) 6.3 x 10^-5 M b) 6.3 x 10^-5 M c) 1.33 x 10^-1 M d) 4.7 x 10^-4/1 e) 4.7 x 10^-2% 14) 0.29 g

15a) 2.57 x 10^-2 M b) 25.7% c) 0.257

17c) 11.2 17b) 2 x 10^-11 18a) 6.3 x 10^-14 b) 2.5 x 10^-5 c) 5.9 x 10^-9

18d) 8.7 x 10^-6 e) 5.4 x 10^-3 19a) 8.6 b) 1.6 20a) 7.94 b) 3.1 c) 4.63 d) 7.00

22a) 4.0 x 10^-8 b) 2.5 x 10^-7 24i) 7.00 ii) 9.07 25) 6.3 x 10^-6/1 26a) 24.7 M

b) 4.05 ml c) 4.47 x 10^-3 M, pH 2.35 26d) 3.96 x 10^-5 M, pH 4.4 e) 3.3 x 10^-5 M, pH4.5

30a) 4.74 b) 4.57 31a) 3.7 b) 3.6 c) 3.9 32a) pOH 5.77 b) pH = 8.15

33e) 1.48 x 10^-5

34) 1.8/1

35) 9.43 36) -0.30 37) 4.25 39) 0.163 M 40) 0.050 L

41) 0.3792 M

42) 0.500 43) 88.1 g/n 46) 60.0% 49a) 0.10 M b) 1 to 2 x 10^-5 c) 1 to 2 x 10^-5

50) 4.77 51a) 0.200 M b) i) 1.28 x 10^-4 ii) 1.36 x 10^-4 c) 8.5 52a) 4.3

53a) 27.5 kJ

“Science is not everything, but science is very beautiful”

J. Robert Oppenheimer (1904-1967)

QCP Equation Types

Aa	alpha	Nn	nu
Bb	beta	Xx	xi
Gg	gamma	Oo	omicron
Dd	delta	Pp	pi
Ee	epsilon	Rr	rho
Zz	zeta	Ss	sigma
Ee	eta	Tt	tau
Qq	theta	Uu	upsilon
Ii	iota	Ff	phi
Kk	kappa	Cc	chi
Ll	lambda	Yy	psi
Mm	mu	Ww	omega

Greek Alphabet

Fundamental Physical Constants

Standard gravitational acceleration, g	9.8 m/s^2
Standard pressure	1 atmosphere 101,325 N/m^2 101,325 Pascals 760 mm Hg 760 torr 14.7 pounds/in^2
Speed of light in a vacuum, c	2.997 x 10^8 m/s
Charge of an electron	1.602 x 10^-19 Coulomb
Avogadro’s constant	6.0225 x 10^23
Faraday’s constant, F	9.6487 x 10^4 Coulombs/mole
Planck’s constant, h	6.6256 x 10^34 Joule•s
Ideal Gas constant, R	0.0821 L•atm/K•mol 0.008314 kJ/K•mol 0.00199 kcal/K•mol
Volume per mole of ideal gas at 1 atm and 0°C	22.414 liters
Electron•volt	1.602 x 10^-19 Joule
rest mass of electron	9.11 x 10^-31 kg
rest mass of proton	1.67 x 10^-27 kg
rest mass of neutron	1.67 x 10^-27 kg
mass of earth	5.98 x 10^24 kg

Measurement	Standard SI Unit	Common Lab Unit	Equivalency to SI
mass	kilogram (kg)	gram (g)	1 g = 0.001 kg
length	meter (m)	centimeter (cm) millimeter (mm)	1 cm = 0.01 m 1 mm = 0.001 m
volume	cubic meter (m^3)	liter (L) milliliter (mL) cubic centimeter (cm^3)	1 L = 1000 mL 1 mL = 1 cm^3
time	second (s)
area	meter^2
velocity	meter/second
pressure	newton/meter^2	pascal (Pa)
power	kg•m^2/s^2	watt (W)
electric current	coulombs/second	ampere (I)

Important Metric Prefixes

Prefix	Abbreviation	Examples
tera	T	10^12 meters/Tm
giga	G	10^9 meters/Gm
mega	M	10^6 meters/Mm
kilo	k	10^3meters/km
deci	d	10^-1 meter/dm
centi	c	10^-2 meter/cm
milli	m	10^-3meter/mm
micro	µ	10^-6 meter/µm
nano	n	10^-9 meter/nm
pico	p	10^-12 meter/pm
femto	f	10^-15 meter/fm
atto	a	10^-18 meter/am

Conversion Factors

Length 1 meter = 39.370 inches = 3.281 feet = 1.094 yard = 6.2137 x 10^-4 mile 1 inch = 2.540 centimeters 1 light•year = 5.88 x 10^12 miles = 9.4606 x 10^15 meters 1 mile = 5,280 feet 1 Angstrom = 1 x 10^-10meter 1 micron (µ) = 1 x 10^-6 meter 1 nanometer = 1 x 10^-9 meter	Mass 1 kilogram = 2.2046 pounds 1 pound = 453.59 grams 1 pound = 16 ounces 1 atomic mass unit (amu) = 1.6604 x 10^-24 gram 1 metric ton = 1000 kilograms
Volume 1 liter = 1.0567 quarts 1 gallon = 3.7854 liters 1 gallon = 4 quarts 1 liter = 1000 ml 1 m^3 = 1000 liters 1 foot^3 = 7.481 gallons = 28.32 liters 1 inch^3 = 16.387 ml

Temperature, °C	Pressure, mm Hg	Temperature, °C	Pressure, mm Hg
0	4.6	26	25.2
1	4.9	27	26.7
2	5.3	28	28.3
3	5.7	29	30.0
4	6.1	30	31.8
5	6.5	31	33.7
6	7.0	32	35.7
7	7.5	33	37.7
8	8.0	34	39.9
9	8.6	35	42.2
10	9.2	40	55.3
11	9.8	45	71.9
12	10.5	50	92.5
13	11.2	55	118.0
14	12.0	60	149.4
15	12.8	65	187.5
16	13.6	70	233.7
17	14.5	75	289.1
18	15.5	80	355.1
19	16.5	85	433.6
20	17.5	90	525.8
21	18.7	95	633.9
22	19.8	97	682.1
23	21.1	99	733.2
24	22.4	100	760.0
25	23.8	101	787.6

Vapor Pressure of Water

Common Ions

IA

IIA

IIIA

IVA

VA

VIA

VIIA

0

Li^+

N^3-

O^2-

F^-

Na^+

Mg^2+

<———————————Transition Metals———————————>

Al^3+

P^3-

S^2-

Cl^-

K^+

Ca^2+

Cr^2+ Cr^3+

Mn^2+ Mn^3+

Fe^2+ Fe^3+

Co^2+ Co^3+

Cu^+ Cu^2+

Zn^2+

Br^-

Rb^+

Sr^2+

Ag^+

Cd^2+

Sn^2+ Sn^4+

I^-

Cs^+

Ba^2+

Hg2^2+ Hg^2+

Pb^2+ Pb^4+

Polyatomic Ions or Radicals

-1	-2	-3
acetate, C2H3O2^-	carbonate, CO3^2-	phosphate, PO4^3-
hydrogen carbonate, HCO3^-	chromate, CrO4^2-
hydrogen sulfate, HSO4^-	sulfate, SO4^2-
cyanide, CN^-	sulfite, SO3^2-
hydroxide, OH^-	dichromate, Cr2O7^2-
nitrate, NO3^-	thiosulfate, S2O3^2-
nitrite, NO2^-
permanganate, MnO4^-
dihydrogen phosphate, H2PO4^-
perchlorate, ClO4^-		+1
chlorate, ClO3^-		ammonium, NH4^+
chlorite, ClO2^-
hypochlorite, ClO^-