Chapter 23

Coordinate Covalent Bonding and

Lewis Acid Base Reactivity

 

Table of comparable charge density

Li+

1.7

Be2+

6.5

 

 

 

 

 

 

 

 

 

 

 

Na+

1.1

Mg2+

3.1

 

 

 

 

 

 

 

 

 

 

Al3+

6.0

K+

1.75

Ca2+

2.0

 

Ti4+

5.9

 

Cr3+

4.6

Mn2+

2.5

Fe3+

4.5

Co2+

2.4

Ni2+

2.9

Cu2+

2.9

Zn2+

2.7

Ga3+

4.8

Rb+

.68

Sr2+

1.8

Y3+

3.2

Zr4+

5.0

 

 

 

 

 

 

Ag+

.79

Cd2+

2.1

In3+

3.7

Cs+

.59

Ba2+

1.5

La3+

2.6

 

 

 

 

 

 

Pt4+

6.2

 

Hg2+

1.8

Tl3+

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-

*Al3+

sp3

4

tetrahedral

OH-

Al3+

sp3d2

6

octahedral

H2O

*Cr3+

sp3

4

tetrahedral

CN-  NH3  Cl-  OH-

Cr3+

sp3d2

6

octahedral

CN-  NH3  Cl-

*Cu2+

dsp2

4

square planar

CN-  NH3  Cl-

*Fe3+

sp3d2

6

octahedral

SCN-   CN-  H2O

*Zn2+

sp3

4

tetrahedral

CN-  NH3  OH-

Pt4+

dsp2

4

square planar

CN-  NH3  Cl-

Pt4+

sp3d2

6

octahedral

CN-  NH3  Cl-

 

 

COMMON LEWIS BASES (LIGANDS)

Ligand

Name of Ligand

H2O

aquo

NH3

ammine

CN-

cyano

Cl-

chloro

OH-

hydroxo

SO42-

sulfato

NO3-

nitrato

C2O42-

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

SO42-

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

SO32-

Ammonium ions

NH4+

ammonia 

NH3

Hydrogen carbonate ions

HCO3-

carbonate ions

CO32-

Hydrogen sulfide ions

HS-­

sulfide ions

S2-­

Water

H2O

hydroxide ions

OH-

Hydroxide ions              Very

OH­-

oxide ions                       Very

O2­-

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- ———> BeF42-             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)42-

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) CuCl42-                      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) ZnCl42-                          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 BaCl2xH2O 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.  HPO42-   +   OH-  ———>   

       HPO42-   +   H3O+ ———>

       Complete the above reactions verifying the amphiprotic nature of HPO42-. Is PO43- 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+ + PO43-

 

 

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+   + PO43- ———>  HPO42-  +  H2O

d)   HSO3- + CN- ———> HCN + SO32-

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) CO32- + 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-, HPO42-, PO43-

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+, Pb2+, or Ag+ are insoluble

5. The compounds formed with sulfate, SO42-

dissociates, dissolves, forms aqueous (aq) solutions

Ba2+, Pb2+, Sr2+, or Ca2+ are insoluble

6. The compounds formed with carbonate, CO32-, phosphate, PO43-, sulfide, S2–, oxide, O2–

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)42-

                        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 [NH3, 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,  CrO42-, and the oxalate ion, C2O42-. 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 CO32- 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) ŕ MgCl2(aq) + H2(aq)

Mg(s) + HF(aq) ŕ 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 Kw 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 Kw 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 Kb and Ka

17.  The Ka 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 Kb.     

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+ + CO32-                      Ka = 4.2 x 10-7

                        H2PO4-   <———> H+ + HPO42-                     Ka = 6.2 x 10-8

                        HSO4-    <———>      H+ + SO42-                   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 Ka of the weak acid.

ii)   If the pH of a buffer is 4.83, what is the Ka 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 (Ka or Kb) 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 Kb for NH3 using the graph.

c)   Calculate the Kb 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

Text Box: The Dissociation of a Salt In Water
NaCl(s)  > Na+(aq)  +  Cl-(aq)

Ionic or Double Replacement Reactions That Produce Precipitates
NaCl(aq)  +  AgNO3(aq) --> NaNO3(aq)   +   AgCl(s)
Double Replacement Reactions That Produce Gases And Water
very unstable	     very stable
NH4OH  -->   NH3(g)   +   H2O(l)
H2CO3    -->   CO2(g)    +   H2O(l)
H2SO3     -->  SO2(g)    +   H2O(l)
Na2CO3(aq)  + HCl(aq)  --> NaCl(aq) + H2O + CO2(g)

Double Replacement Reactions That Produce Weak Acids
NaF(aq)    +    HCl(aq) -->  NaCl(aq)  +  HF(aq)

Composition
H2(g)   +   O2(g)  --> H2O(l)

Decomposition
NH3(g)   --> N2(g)  +  H2(g)
KClO3(s) --> KCl(s)   +    O2(g)
Combustion Of A Hydrocarbon
C4H10(g)   +   O2(g) --> CO2(g)    +      H2O

Metal Replacement
Zn(s)    +    CuSO4(aq)  --> ZnSO4(aq)    +     Cu(s)
 Mg(s)     +     HCl(aq)	--> MgCl2(aq)      +      H2(g)

Active Metals In Water
Na(s)    +    H2O --> NaOH(aq)    +    H2(g)

Nonmetal Replacement
KI(aq)    +    Br2(l)  -->  KBr(aq)  +  I2(s)

 

 

 

 

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/s2

 

 

Standard pressure

1 atmosphere

101,325 N/m2

101,325 Pascals

760 mm Hg

760 torr

14.7 pounds/in2

Speed of light in a vacuum, c

2.997 x 108 m/s

Charge of an electron

1.602 x 10-19 Coulomb

Avogadro’s constant

6.0225 x 1023

Faraday’s constant, F

9.6487 x 104 Coulombs/mole

Planck’s constant, h

6.6256 x 1034 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 1024 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 (m3)

liter (L)

milliliter (mL)

cubic centimeter (cm3)

1 L = 1000 mL

1 mL = 1 cm3

 

time

second (s)

 

 

area

meter2

 

 

velocity

meter/second

 

 

pressure

newton/meter2

pascal (Pa)

 

power

kg•m2/s2

watt (W)

 

electric current

coulombs/second

ampere (I)

 

 

 

 

 

 

 

 

 

 

 

Important Metric Prefixes

Prefix

Abbreviation

Examples

tera

T

1012 meters/Tm

giga

G

109 meters/Gm

mega

M

106 meters/Mm

kilo

k

103 meters/km

deci

d

10-1 meter/dm

centi

c

10-2 meter/cm

milli

m

10-3 meter/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 1012 miles = 9.4606 x 1015 meters

1 mile = 5,280 feet

1 Angstrom = 1 x 10-10 meter

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 m3 = 1000 liters

1 foot3 = 7.481 gallons = 28.32 liters

1 inch3 = 16.387 ml

 

 

 

Text Box: Work and Energy
1 electron volt = 1.6021 x 10-12 erg = 23.061 kcal/mole
1 Joule = 107 ergs
1 calorie = 4.1840 Joules
1 kilocalorie = 4.1840 kiloJoules

Temperature					Time
°C = 5/9(°F - 32)				1 hour = 60 minute = 3,600 seconds
°F = 9/5°C + 32					1 minute = 60 seconds 
°K = °C + 273					1 year = 365.25 days

Pressure
1 atmosphere = 760 mm Hg or torr
1 atmosphere = 101,325 Pascals = 14.7 pounds/in2 = 1.013 bar

Specific Heat
1 calorie/g•°C = 4.184 Joule/g•°C

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

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+

 

 

 

 

 

 

N3-

 

 

O2-

 

 

F-

 

 

 

Na+

 

 

Mg2+

 

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

 

 

Al3+

 

 

 

P3-

 

 

S2-

 

 

Cl-

 

 

 

K+

 

 

Ca2+

 

 

 

 

 

Cr2+

Cr3+

 

Mn2+

Mn3+

 

Fe2+ Fe3+

 

Co2+

Co3+

 

 

Cu+

Cu2+

 

Zn2+

 

 

 

 

 

 

Br-

 

 

 

Rb+

 

 

Sr2+

 

 

 

 

 

 

 

 

 

 

Ag+

 

 

Cd2+

 

 

 

Sn2+

Sn4+

 

 

 

 

I-

 

 

 

Cs+

 

 

Ba2+

 

 

 

 

 

 

 

 

 

 

 

Hg22+

Hg2+ 

 

 

Pb2+

Pb4+

 

 

 

 

 

 

 

 

 

Polyatomic Ions or Radicals

 

-1

-2

-3

acetate, C2H3O2-

       carbonate, CO32-

       phosphate, PO43-

   hydrogen carbonate,

HCO3-

       chromate, CrO42-

 

hydrogen sulfate, 

HSO4-

       sulfate, SO42-

 

     cyanide, CN-

       sulfite, SO32-

 

     hydroxide, OH-

       dichromate, Cr2O72-

 

     nitrate, NO3-

       thiosulfate, S2O32-

 

     nitrite, NO2-

 

 

permanganate,  MnO4-

 

 

dihydrogen phosphate,                     H2PO4-

 

 

    perchlorate, ClO4-

 

+1

     chlorate,  ClO3-

 

ammonium, NH4+

     chlorite, ClO2-

 

 

     hypochlorite, ClO-