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CHEM 30
1.
FINAL EXAM REVIEW ( 2012-13)
THERMOCHEMISTRY (CH 15)
- recognize endothermic and exothermic changes in energy; recognize endo/exo reactions
- write equations for endothermic and exothermic processes
- solve thermochemical problems using specific heats, molar heat capacities, q = mcΔt,
ΔH = q/n, etc.
- molecular kinetic energies; Ke = ½ mv2
- solving calorimetry problems
- use the variety of methods to determine ΔH: Hess’ law and heats of formation, bond energies
- recognize positive and negative changes in entropy in a reaction
- predicting changes in entropy, ΔS in a reaction
- determine spontaneity of reactions using the Gibbs free energy equation
- relationship between Gibbs free energy, entropy and enthalpy: ΔG = ΔH – TΔS
-comparing energy changes for physical, chemical and nuclear changes
2.
REACTION RATES (CH 16)
-calculate rate of reactions using: r = Δc/Δt
- explain the factors which effect reaction rates; predicting rate of reactions
- concentration vs time graphs – instantaneous rate of reaction
- use activation energy diagrams and kinetic energy diagrams to show effect of
temperature and catalysts on reaction rate
- explain reaction rates in terms of the collision theory - effect of changes in temperature,
concentrations, nature of the reactants, etc.
- identify reaction mechanisms: catalysis, reaction intermediates, net overall equations, ratedetermining steps
- explain how catalysts speed up reaction rates
- be able to explain one example of catalysis
- describe enzyme activity as an example of a biological catalyst
- use the Arrhenius equation to compare rates at different temperatures;
- use the Arrhenius equation to solve for activation energy and Temperature
3.
EQUILIBRIUM
(CH 17)
- recognizing equilibrium conditions, reversibility of reactions, meaning of equilibrium
- writing equilibrium expressions and solving problems ( Keq expressions)
- use Le Chatelier’s principle to predict how equilibria will shift if changes are imposed on systems
at equilibrium
- use ICE box methods to calculate equilibrium concentrations of substances
- understand the Haber principle as an example of applied equilibrium theory
- Le Chatelier’s principle, working with equilibrium expressions qualitatively and quantitatively.
- Common ion effect on equilibria in aqueous systems
4.
SOLUTIONS/SOLUBILITY (CH 14, CH.17.3)
- explain miscibility of liquids and solubility of solids and gases in liquids on the basis of their
intermolecular forces
- explain the effect of temperature and pressure on the solubility of gases and solids in liquids
- express the concentrations of solutions using: Molarity (M, mol/L) , %weight by volume, ppm and
ppb; use c = n/v to solve problems regarding solutions
- using the dilution formula C1V1 = C2V2; apply to serial dilutions
- use solubility table to predict the solubility of various ion combinations in aqueous solutions
- writing molecular and complete ionic equations for solution processes
- writing precipitate reactions, predicting precipitates
- use the solubility table to devise qualitative analysis separation of ions in solution
- apply techniques of qualitative analysis to identify presence of ions in solutions
- calculate maximum ion concentrations possible from Ksp values
- calculate a solubility in g/L or mol/L from Ksp values
- calculate a Ksp value from a given solubility
- quantitative prediction of precipitates using Qsp methods
- solve common ion problems with regards to aqueous solutions and solubility
5.
ACIDS AND BASES
(CH 18)
- theories of acids and bases: Arrhenius, Bronsted-Lowry, Lewis
- applying all three theories of acids and bases: completing neutralization reactions, Bronsted-lowry - equations and recognizing conjugate acids and bases
- predicting the pH of salt solutions
- using the water constant, Kw – calculations of [H3O+] [OH-] for strong acids and bases
- using the pH scale; performing calculations based on pH equations and titrations
- using ionization constants, Ka and Kb values to do calculations for weak acids and bases
- problems using common ion effect for weak acids/bases
- salt hydrolysis; predicting the pH of salt solutions
- buffering of solutions, acid-base indicators, and calculations using the common ion effect
- recognizing titration curves; solving titration problems
- using the Henderson – Hasselbach equation to solve buffer pH questions
6.
OXIDATION-REDUCTION
(CH 19, 20)
- assigning oxidation states to elements in compounds and equations;
- balancing redox equations using oxidation states;
- recognizing oxidized, reduced species, oxidizing agents and reducing agents;
- predicting spontaneity of reactions and E0 values for redox reactions
- understanding how an electrochemical cell works: half-reactions, salt bridge, direction of
electron flow, E0 values, shorthand notation, stoichiometry problems
- corrosion of metals: oxidation and reduction reactions involved; how to prevent or slow down the
rate of corrosion of iron
- application of electrochemistry: electroplating, batteries, electrolytic cells
- quantitative electrochemistry: faraday’s equation and quantitative electrolysis
Some Practice Questions:
THERMOCHEMISTRY
1.
2.
3.
Differentiate between exothermic and endothermic reactions.
o
What amount of energy is required to raise the temperature of 10.0 g of aluminum metal from 25.0 C
o
to 35.0 C?
Given the reaction: N2(g) + 3 H2(g)

2 NH3(g) + 83.6 kJ
a) is this reaction exothermic or endothermic?
b) how much heat will be released when 5.10 g of ammonia form?
4.
In an experiment using a simple styrofoam calorimeter, a student dissolves 4.62 g of KOH in 250
o
o
mL of distilled water. The temperature of the water increases from 21.6 C to 28.4 C Calculate the
molar enthalpy of solution for KOH in water.
5.
Calculate the change in enthalpy, ΔH, and predict the change in entropy, ∆S
for the following reaction: C2H6(g) + 7/2 O2(g) 
2 CO2(g) + 3 H2O(l)
6.
Predict whether ΔS is negative or positive for the following chemical reactions:
a) 2 HgO(s) 
2 Hg(l) + O2(g)
b) 2 SO2(g) + O2(g) 
2 SO3(g)
7.
Use heats of formation to determine the enthalpy change for the following reaction under standard
conditions:
CH4(g) + 2 O2(g) 
CO2(g) + 2 H2O(g)
8.
Use bond energies to calculate the approximate change in enthalpy for the reaction:
2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(g)
o
Bond
C–C
C–H
O=O
C=O
O–H
Bond Energy (kJ mol-1)
347
414
499
799
460
9.
While calculating a ΔH value experimentally in the lab, a student determined the value to be
-535 kJ mol-1. After consulting a chemistry handbook, the student discovered the accepted value to be
-547 kJ mol-1. The percent error in the experiment was:
A. 2.2%
B. 12.0%
C. 65.2%
D. 97.8%
10.
Which of the following reactions is endothermic?
A. H2(g) + ½ O2(g) → H2O(g)
B. C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g)
C. Al(s) + 3/2 F2(g) → AlF3(s) + 1306.2 kJ
D. H2O(g) + C(s) → CO(g) + H2(g)
ΔH = -242.76 kJ
ΔH = -2046.2 kJ
ΔH = +131.86 kJ
11.
Which of the following is an endothermic change?
A. Hot coffee cooling to room temperature
B. Steam condensing to water on a cold window
C. Frost forming on a windshield
D. Ice melting
12.
The molar heat of formation for some sodium halides are given below.
Na(s) + ½ I2(g) → NaI(s)
ΔH = -287.9 kJ
Na(s) + ½ Br2(g) → NaBr(s)
ΔH = -361.1 kJ
Na(s) + ½ Cl2(g) → NaCl(s)
ΔH = -411.2 kJ
Na(s) + ½ F2(g) → NaF(s)
ΔH = -570.7 kJ
A generalization that can be made about these compounds is that
A. iodine forms stronger bonds than chlorine does.
B. bromine compounds are less stable than chlorine compounds.
C. fluorine forms the weakest bond of the halides listed.
D. sodium halides absorb heat when formed.
13.
Use the reactions below to determine the enthalpy for the combustion of methane:
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
CH4(g) + ½ O2(g) → CH3OH(l)
CH2O(g) + H2O(g) → CH3OH(l) + ½ O2(g)
CH2O(g) + ½ O2(g) → HCOOH(g)
HCOOH(g) + ½ O2(g) → CO2(g) + H2O(g)
ΔH = -164.2 kJ
ΔH = +118.7 kJ
ΔH = -308.8 kJ
ΔH = -210.6 kJ
REACTION RATES
1.
For the following reaction, would the rate of loss of oxygen be the same as the rate of loss of
propane?
C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(l)
2.
Which of the following reactions are likely to be very rapid reactions? Why?
2+
2a) Zn (aq) + S (aq)  ZnS(s)
b) H2SO4(aq) + 2 NaOH(aq)  Na2SO4(aq) + 2 H2O(l)
c) Zn(s) + S(s)  ZnS(s)
3.
Draw an energy diagram for the reaction: X 
using the following information:
- the mechanism for the reaction consists of two steps
- the reaction is exothermic
- the second step is slower than the first step
4.
In an exothermic reaction, ΔH was found to be -40kJ, and the activation energy, Eato be 30 kJ. What would
be the activation energy of the reverse reaction, Ear?
5.
The reaction of nitrogen monoxide with oxygen to form nitrogen dioxide is thought to proceed via
the following mechanism:
step(1)
2 NO 
N2O2
step(2)
N2O2 + O2 
2 NO2
a) identify any reaction intermediates b) what is the net reaction?
6.
Identify the catalyst in the reaction mechanism below:
step (1)
NO + Br

NOBr
step (2)
NOBr + Br2 
NOBr2 + Br
7.
The rate of a chemical reaction depends
A. only on the collision frequency.
B. only on the collision efficiency.
C. on both the collision frequency and the collision efficiency.
D. on neither the collision frequency nor the collision efficiency.
Use the diagram below to answer questions 8 and 9.
3
1
1
2
4
8.
The activation energy for the reverse reaction is represented by
A. 1
B. 2
C. 3
D. 4
9.
The ΔH for the forward reaction is represented by:
A. 1
B. 2
C. 3
D. 4
2Draw a graph of [Te2O8 ] vs time to answer the following questions:
10.
2-
.
[Te2O8 ] in mol L
0.400
0.240
0.144
0.0864
0.05184
-1
time ( minutes )
0
6
12
18
24
a) determine the [Te2O82-] at 15 minutes.
b) determine the time required for the [Te2O82-] to drop to 0.069 mol.L-1
c) draw a tangent line to the graph at time = 21 minutes. Find the slope of this tangent line to
determine the rate of reaction at 21 minutes.
11
Assume that the rate of reaction doubles for every 10 °C rise in temperature. A student determined
that a certain reaction reached completion in 20 minutes at 0 °C. If the same reaction was carried out
at 30 °C, it would be expected to reach completion in
A. 20 minutes.
B. 10 minutes.
C. 5 minutes.
D. 2.5 minutes.
12.
A certain reaction has an activation energy of 125 kJ mol-1.
Compare the rates of this reaction at 500C and 100C
13.
A reaction at 2000C a reaction has an Arrhenius constant of 5.00 and a rate constant
of 0.80
Determine the activation energy for the reaction.
EQUILIBRIUM
1.
Write the equilibrium expression constant expression for the following equation:
S(s) + 3/2 O2(g)  SO3 (g)
2.
Consider the following reaction:
4 HCl (g) + O2(g)  2 H2O(g) + 2 Cl2(g) + 113 kJ
What effect would the following changes have on the equilibrium concentration of Cl2?
i) increasing the temperature of the system ii) decreasing the total pressure on the system
iii) increasing the concentration of oxygen gas iv) adding a catalyst
3.
Which of the following remains constant at equilibrium?
a) the concentration of the products
b) the concentration of the reactants
c) the ratio of product concentration to reactant concentration
4.
What does the magnitude of the equilibrium constant, Keq, indicate about the extent of the reaction?
5.
Explain why the value of Keq for an exothermic reaction increases as the temperature decreases.
6.
Write the equilibrium expression for each of the following processes:
a) Br2(g) + 5 F2(g)  2 BrF5(g)
b) Ca3(PO4)2(s) 
3 Ca2+(aq) + 2 PO43-(aq)
7.
At 10000C methane gas reacts with steam as follows: CH4(g) + H2O(g)  CO(g) + 3 H2(g)
In one experiment the equilibrium concentrations of the gases were:
[CH4] = 2.97 x 10-3 atm
[H2O] = 7.94 x 10-3 atm
-3
[CO] = 5.45 x 10 atm
[H2] = 2.1 x 10-3 atm
Calculate the value of the Keq for the reaction at this temperature.
8.
Given the reaction: N2(g) + O2(g)  2 NO(g)
H = 179 kJ
In which direction will the equilibrium shift, with the following changes:
i) addition of more nitrogen gas
ii) volume of the reaction vessel is doubled
iii) addition of a catalyst
iv) removal of some oxygen gas
9.
The following reaction occurs readily at 4250C: 2 NO(g) + Cl2(g)  2 NOCl(g)
The equilibrium constant, Keq , is 14.9 at this temperature. Predict the shift that the reaction would
take to establish equilibrium for each of the following starting conditions:
i) all gases are at a concentration of 0.200 mol/L
ii) [NOCl] = 0.100 atm; [NO] = 0.0500 atm ; [Cl2] = 0.300 atm
10.
If an equilibrium constant for a system has a value of 5.62 x 1049, would you expect to
find mostly reactant or mostly product after the system has come to equilibrium?
11.
If you wished to maximize the products in the following equilibrium, would you run the reaction at a
high pressure or a low pressure? CaCO3(s)  CaO(s) + CO2(g)
12.
Ammonia NH3(g) is produced according to the following equation.
N2(g) + 3 H2(g)  2 NH3(g)
When the system has reached equilibrium, the value of the equilibrium constant,Keq,
can be changed by
A. increasing the pressure.
B. increasing the temperature.
C. adding more nitrogen.
D. decreasing the concentration of ammonia.
13.
Anhydrous ammonia, an important fertilizer used in Saskatchewan’s agriculture industry, is
created by the following process: N2(g) + 3 H2(g)  2 NH3(g)
An experiment was carried out with initial [N2] = 0.96 M and initial [H2] = 0.72 M.
The [NH3] at equilibrium was found to be 0.24 M
[Initial]
[Change]
[Equilibrium]
[N2]
0.96 M
[H2]
0.72 M
0.24 M
Based on the above data, the Keq value is:
A. 0.16
B. 1.5
C. 2.6
14.
[NH3]
0
D. 120
The graph of concentration versus time over three time intervals for the equilibrium:
X(g)  Y(g) + Z(g) + heat is given below
C
o
n
c
e
n
t
r
a
t
i
o
n
14
12
10
8
6
4
2
0
X(g)
Y(g)
Z(g)
0
6
12
18
Time
The stress applied at t = 6 was:
A. a decrease in temperature
C. the addition of a catalyst
B. a decrease in pressure
D. a decrease in [X]
15.
The Ksp for silver iodide, AgI is known to be 1.5 x 10-16. Determine the [Ag+] in a saturated
solution of silver iodide.
16.
The Ksp value for AgBr is known to be 6.5 x 10-13. What is the maximum [Br-] in a solution
which contains a [Ag+] of 2.5 x 10-8 mol/L?
17.
The Ksp for Mg(OH)2 at 450C in water is known to be: 4.0 x 10-5
Determine the solubility of this hydroxide in mol L-1 and g L-1 at this temperature.
18.
Will a precipitate form when 62.6 ml of 0.0322 M CaCl2 solutions is mixed with 31.3 ml of
0.0145 M NaOH solution? The Ksp for Ca(OH)2 is 5.0 x 10-6
19.
The Ksp for scandium fluoride (ScF3) at 298 K is 4.2 x 10-18
Write the chemical equation for the solubility equilibrium of scandium fluoride in water.
What concentration of Sc3+ ions is required to cause a precipitate to form if the fluoride
ion concentration in a solution is 0.076 M ?
OXIDATION-REDUCTION/ELECTROCHEMISTRY
1.
Use half reactions to complete and balance the following redox reaction:
Zn + NO3- + H+  Zn2+ + NO2 + H2O
2.
Use oxidation states to balance the following redox reaction:
H+ + Cr2O72- + SO2  Cr3+ + HSO4- + H2O
3.
Use standard reduction potentials (E0 values) to determine if the following reaction is spontaneous:
Cu+ + Ag  Cu + Ag+
4.
Given the redox reaction:
Zn + Br2  ZnBr2
a) which atom is oxidized? b) which atom is reduced?
c) what is the oxidizing agent?
5.
Determine the oxidation states for all the atoms in the following chemical species:
a) KMnO4 b) V2O3 c) Cr2O72- d) Fe3O4 e) FeS
6.
Identify the reactant that is oxidized and the reactant that is reduced in the reaction:
H2C2O4 + H2O2  2 CO2 + 2 H2O
7.
If you wish to store a solution containing Sn(NO3)2 without having a reaction between the
solution and the container, the container could be made of:
A. aluminum
B. nickel
C. copper
D. zinc
8.
The equation fragment that would be found in a reduction half-reaction is:
A. X2+ → X3+
B. 2R→ R2
3+
2+
C. Z
→ Z
D. R
→ R+
9.
In the equation, 2I-(aq) + F2(g) → 2F-(aq) + I2(s) , the substance that is oxidized is:
A. IB. F2
C. FD. I2
10.
A magnesium – chromium electrochemical (voltaic) cell is set up as below:
Mg
NaNO3
Cr
a) in which direction will electrons flow through the wire?
b) write the anode and cathode half-reactions which would occur.
c) write the balanced overall reaction for this electrochemical cell.
Use the diagram below to answer questions 11 and 12
eV
Mg(s)
Cu (s)
1.0 mol L-1
Mg(NO3)2(aq)
1.0 mol L-1
Cu(NO3)2(aq)
porous
cup
11.
12.
The anode is made of:
A. Cu
B. Cu(NO3)2 (aq)
C. Mg
The reduction half-reaction for the cell is:
A. Cu2+(aq) + 2e- → Cu(s)
C. Mg2+(aq) + 2e- → Mg(aq)
D. Mg(NO3)2 (aq)
B. Cu(s) → Cu2+(aq) + 2eD. Mg(s) → Mg2+(aq) + 2e-
13.
A zinc-lead electrochemical cell is allowed to run long enough so that the mass of the anode changes
by 3.00 g. what will be the corresponding change in mass at the cathode?
14.
Determine the voltage produced by the following voltaic cell: FeFe2+  Br2  Br- standard
conditions
Which of the following methods will slow the corrosion of iron?
A. connecting iron to a piece of lead
B. wrapping tin around the iron
C. connecting iron to the negative terminal of a power source
D. connecting iron to the positive terminal of a power source
15.
16.
Consider the following equation for the rusting of iron:
4 Fe(s) + 3 O2(g) → 2 Fe2O3(s)
Which of the following statements about this reaction is NOT true?
A. This equation is an example of an oxidation-reduction?
B. Iron is the reducing agent in this reaction
C. Oxygen is oxidized in this reaction
D. Iron changes in oxidation number from 0 to +3
17.
What current must be used to plate 1.75 mole of copper on an electrode in 6.24 minutes?
18.
How many grams of silver will be deposited by a current of 1.00 A flowing for 9650 seconds?
19.
Explain why magnesium rods, when inserted into hot water tanks, prevent formation of rusty water
in the tank.
20.
To electroplate a penny with nickel, the penny must be connected to the source of direct current
as the:
A. cathode, and must be placed in a solution containing Ni2+(aq)
B. anode, and must be placed in a solution containing Ni2+(aq)
C. cathode, and must be placed in a solution containing Cu2+(aq)
D. anode, and must be placed in a solution containing Cu2+(aq)
ACIDS/BASES
1.
The [OH-] in a solution is 4.5 x 10-9M. what is the pH of the solution? is it acidic or basic?
2.
Identify the conjugate acid-base pairs in the reaction: H2O + HNO3  NO3- + H3O+
3.
If 15.0 ml of a 0.0300M HNO3 solution is required to titrate 25.0 ml of a KOH solution,
determine the concentration of the KOH solution.
4.
A solution has pH = 4.65. What are the hydronium ion and hydroxide ion concentrations?
5.
The [H3O+] in a 0.400 M solution of a weak acid, HX, is found to be 0.0260M. Calculate the
KA for this weak acid.
6.
Use LeChatelier’s principle to explain why the ionization of nitrous acid, HNO2, in aqueous solution
is greatly suppressed by the addition of some nitric acid, HNO3.
7.
Complete and balance the neutralization reaction below:
H2SO4(aq) + Al(OH)3(aq) 
8.
Complete the following Bronsted-Lowry acid/base reactions:
a) HSO4- + OH- 
b) H2PO4- + HBr 
9.
Why do we call an aqueous solution of HCl strong and a solution of acetic acid weak?
10.
What is the hydronium ion concentration in a solution of 0.150M in HNO2 and
0.0300M NaNO2? the ionization constant for HNO2 is 7.24 x 10-4
11.
Ammonia, NH3, a weak base is titrated with a strong acid, HNO3. Describe the shape and
characteristics of the titration curve for this experiment. What would be a suitable indicator
for this titration?
12.
Explain why an effective buffer system requires two components; give an example of a buffer
system.
13.
Using a Ka table, determine the strongest acid from the list below:
A. H2SO3(aq)
B. H2CO3(aq) C. HCl(aq)
D. HNO3(aq)
14.
The Kb for the weak base, hydrazine, N2H4 is known to be 1.50 x 10-6
Determine the Ka for its conjugate acid, N2H5+
What is the pH of a 0.02 M solution of hydrazine in water?
15.
HClO(aq), is a weak acid, which has a Ka = 2.9 x 10-8. The [H+] of a 0.85 M solution of this acid is:
A. 0.85 mol L-1 B. 1.6 x 10-4 mol L-1
C. 1.0 x 10-7 mol L-1 D. 2.5 x 10-6 mol L-1
16.
In a titration experiment, 15.0 ml of 0.250 M HCl was neutralized by 22.5 ml of a calcium
hydroxide , Ca(OH)2 , solution of unknown concentration. The concentration of the Ca(OH)2 was:
A. 0.266 mol L-1
B. 0.174 mol L-1
C. 0.133 mol L-1
D. 0.0833 mol L-1
17.
Determine the pH of a buffer prepared using equal volumes of 1.0 M CH3COONa and
0.4 M CH3COOH
18.
Which of the following weak acids would a student choose to prepare a buffer of
pH = 10.5 ?
a) hydrogen sulfide ion, HSc) dihydrogen phosphate ion, H2PO4-
The following graph shows the change in pH of a solution as a function of the volume of
hydrochloric acid added to 40.0 ml of 1.28 x 10-5 mol L-1 sodium hydroxide solution.
14
12
pH
19.
b) oxalic acid, HOOCCOOH
d) hydrogen carbonate ion, HCO3-
equivalence
point
10
8
6
4
2
0
0
20
40
60
80
100
volume HCl used (ml)
The pH of the acid used is
approximately
A. 2
B. 5
C. 7
D. 12
SOLUTIONS AND SOLUBILITY
1.
Write formulas for the following ionic compounds and predict their solubilities in water:
a) potassium phosphate b) copper II bromide c) nickel II sulfate
2.
What mass of Fe2(C2O4)3 is required to prepare 250.0 ml of a 0.0450M solution?
3.
Iodine crystals are relatively insoluble in water. However, they do dissolve in alcohol, benzene, and
several other organic solvents. Explain these solubilities.
4.
5.300 g of common salt, NaCl, are dissolved in 1 500 L of distilled water. Determine the
concentration of the resulting solution using:
a) Molarity, M b) parts per million, ppm c) parts per billion, ppb d) % weight by volume
5.
A precipitate forms in a sample of drinking water when SO42- (sulfate) ions are added, but not
when Cl- (chloride) ions are added. A cation which could cause this effect is:
A. Ca2+
B. Ag+
C. Na+
D. Mg2+
6.
The concentration of a solution made by dissolving 11.7 g of NaCl(s) in enough distilled water
to make 250 ml of solution is:
A. 0.10 mol L-1
B. 0.40 mol L-1
C. 0.80 mol L-1 D. 1.0 mol L-1
7
In a high school chemistry lab, a student accidentally mixed a solution of CaCl2(aq) with a
solution of Na2CO3(aq). A precipitate formed as a result. What is the overall ionic equation for
the reaction described above?
A. Ca2+(aq) + 2 Cl-(aq) + 2 Na+(aq) + CO32-(aq) → CaCO3(s) + 2 Na+(aq) + 2 Cl-(aq)
B. Ca2+(aq) + CO32- (aq) → CaCO3(s)
C. Ca2+(aq) + Cl-(aq) + Na+(aq) + CO32-(aq) → CaCO3(s) + NaCl(aq)
D. CaCl2(aq) + Na2CO3(aq) → CaCO3(s) + 2 NaCl(aq)
8.
In water analysis, it is sometimes necessary to separate the ions in solution using selective
precipitation. A solution is known to contain Be2+(aq) and Ra2+(aq) . Which pair of ions listed
below could be used in the order given to selectively precipitate the two cations from solution?
A.
S2-(aq) , CO32-(aq)
B. OH-(aq) , SO32-
C. Br-(aq), SO42-
D. NO3- , CH3COO-
9.
What volume of 12.0 mol L-1 HCl(aq) is needed to make 5.00 L of a 0.500 mol L-1 solution?
10.
When lead nitrate solution, Pb(NO3)2(aq) is added to a solution of sodium iodide, NaI(aq), a yellow
precipitate results. Write the net ionic equation representing this reaction.
11.
Effluent form a chemical factory is collected for analysis of Ag+. 2.00 L of effluent is
evaporated down to a 40.00 ml sample. The [Ag+] in the sample is determined to be
0.00400 M. Determine the [Ag+] in the effluent
12.
Determine the volume of 2.00 M nitric acid required to prepare 500.0 mL of a 0.800 M solution .
Good luck !!