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Year End Review

CHEMISTRY 40S Teacher: Mr. Andrew Hiebert

Once this review is done well in advance …

it is also highly recommended to: a) Read through your notes, and make condensed

notes of the key concepts on cue cards. b) Practice “redoing” the archetypal questions given

as examples in the notes. c) Go through your tests and make sure you could

redo each of the questions with the correct work. d) Create your own practice exams, and then do

them. You may wish to exchange practice exams with a peer for extra practice.

e) REMEMBER: Come for help as soon as you have trouble.

Getting help ‘last minute’ is often not as effective as getting help ‘well in advance’ of the final exam.

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Chemistry 40S Year End Review KINETICS – REACTION RATES

1. What is the basic definition of a reaction rate? 2. Sketch the result of the decomposition of a reactant [A], on a Concentration versus

Time graph with a solid line, then sketch the formation of a product [B] that results from the decomposition of A, using a dotted line.

[mol•L-1] Time [s]

3. a. What is the difference between endothermic and exothermic reactions? b. What is the difference between enthalpy and entropy? c. What makes a chemical reaction ‘spontaneous’? d. Most spontaneous reactions are exothermic, give an example of a spontaneous

reaction that is endothermic.

4. Sketch a graph of Potential Energy versus Reaction Progress for the potential energy changes during an ENDOTHERMIC reaction involving THREE ELEMENTARY STEPS. Identify the rate determining step.

5. What are two common ways in which reaction rates are empirically measured?

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KINETICS – REACTION RATES continued…

6. What are the 5 factors that affect a reaction’s rate? Provide a specific chemical reaction that aptly illustrates each factor. Remember the ‘pop your lid’ lab…

7. How is the instantaneous reaction rate found from a Concentration versus Time

graph?

7. Nitrogen dioxide can combine with oxygen gas to form a single product N2O5 gas. a. Write a balanced chemical equation for the reaction.

b. Suppose at a particular moment during the reaction, molecular oxygen is reacting

at the rate of 0.024 mol•L-1/s, at what rate is nitrogen dioxide reacting? [2 marks]

8. Given the following reaction and experimental data, determine the rate law equation for the reaction.

S2O 28(aq) − + 3I-(aq) → 2SO 2

4−

(aq) + I 3−

(aq)

Trial # [S2O 2

8(aq) − ] (mol•L-1) [I-(aq)] (mol•L-1) Initial Rate (mol•L-1/s)

1 0.080 0.034 4.4 x10-4

2 0.080 0.017 1.1 x10-4 3 0.16 0.017 2.2 x10-4

{The data above has been fabricated and does not represent the real reaction rate pattern}

9. What does it mean if a given reactant has a zero order of reaction for a given reaction? 10. What is ‘half-life’? What makes the half-life of a first order reaction distinct from the half-

life of other orders of reaction? 11. The decomposition of ethane (C2H6) into methyl radicals is a first order reaction with a rate

constant of 5.36x10-4 s-1 at 700oC. Calculate the half-life of the reaction in MINUTES.

12. The rate law for the reaction 2NOBr(g) → 2NO(g) + Br2(g)

at a given temperature is : r = k [NOBr]2 How much time is required to reduce the concentration of NOBr

from 0.900 M to 0.100 M, given that the half-life of this reaction is 2.00s when

[NOBr]o=0.900M?

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EQUILIBRIUM 1. Derive the ‘equilibrium law expression’ from the concepts of reaction rate formulas from

chemical kinetics, given the chemical equilibrium described by aA + bB ⇌ cC + dD. 2. What is the difference between a homogeneous equilibrium and a heterogeneous

equilibrium? Provide a specific example relating to the ‘equilibrium law expression’ using a balanced chemical equation.

3. What is the definition, as stated in the textbook, of Le Châtelier’s Principle? 4. Given 2CO2(g) ⇌ 2CO(g) + O2(g)

Sketch a well labeled graph to illustrate what would happen to the above equilibrium, if some oxygen was removed.

5. Identify three different ways a system could be ‘disturbed’/ ‘stressed’ to cause the equilibrium system to respond. Provide and example for each.

6. In a closed container at 55OC dinitrogen tetroxide gas, N2O4(g), decomposes to nitrogen

dioxide, NO2(g), with an equilibrium constant of 0.87. A sample of dinitrogen tetroxide is released into a rigid container under those conditions and is analyzed several minutes later. The concentrations at that time were found to be [N2O4(g)] = 0.90 mol/L and [NO2(g)]=0.90 mol/L. Determine whether the system is in equilibrium, and if not predict the direction in which the reaction will proceed to achieve equilibrium.

8. If 0.50 mol of iodine, 1.0 mol of chlorine, and 1.5 mol of the

Iodine monochloride are initially placed into a 2.00 L reaction vessel at 25oC, find the concentration of all entities at equilibrium?

I2(g) + Cl2(g) ⇌ 2ICl(g) K=81.9 at 25oC

9. What are the three main types of dynamic equilibrium in chemistry?

10. What are homogenous equilibria? Heterogeneous equilibria?

11. How does the equilibrium expression for heterogeneous equilibria differ from that of homogeneous equilibria?

12. Distinguish between the term equilibrium constant and reaction quotient. What are the

symbols used for each?

13. How can the reaction quotient be used to predict the flow of a reaction? Provide a specific example.

14. How can the reaction quotient be used to predict precipitation? Provide a specific

example.

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EQUILIBRIUM continued…

15. In general, when solving equilibrium problems an ICE chart is needed. What is an ICE chart? Given Q >K create an ICE chart that could be used to solve for the equilibrium concentrations if K was know. The theoretical reaction is as follows:

aA(aq) + bB(aq) ⇌ cC(aq) + dD(aq)

16. At times in the process of solving for “x” to determine the equilibrium concentrations a short cut called “the hundred rule” or the “5% rule” is used. Explain when the short cut can be used, and how the validity of the short cut is double checked.

17. What type(s) of algebra may need to be employed if the “hundred rule” or “5% rule”

can not be used? Provide an example or two.

18. How can solubility measured in mol/L (Molar) be converted into a solubility measured in g/L?

19. How is solubility different than the solubility product? Explain.

20. How can Ksp be used to find the solubility of an ionic compound? Provide a detailed

calculation example to illustrate.

21. Describe the significance of Q being >, or = ,or <, Ksp.

ACID & BASES 1. What is the Boyle definition of acids & bases? Why is the Boyle definition of acids and

bases considered too limited? [include at least 2 reasons]

2. What is the Arrhenius definition of acids & bases? Why is the Arrhenius definition of acids

and bases considered too limited? [include at least 2 reasons]

5. Determine the pH of a 0.100 M solution of ammonium chloride. 6. A sample of 0.0001 M HCl has very nearly the same pH as a 0.10 M solution of

CH3COOH. Are hydrochloric acid and acetic acid equally strong in these samples?

Explain.

7. Explain how water can auto-ionize. Show the equilibrium that results. Include the value of Kw at SATP

8. Find the concentration of H3O+ in each of the following solutions, then find the pH of each.

a) 0.0023 M HCl b) 0.05 M NaOH c) 0.145 M Ba(OH)2

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9. The pH of a solution is given as 2.8. What is the hydronium ion concentration of this

solution?

10. Taking neutralization into account, what is the pH of a solution prepared by adding 25.0

mL of 0.10 M KOH to 50 .0 mL of 0.06 M HNO3? (These are strong acids and bases.)

11. In a titration 20.00 mL of 0.4000 mol/L of HCl is titrated with standardized 0.4000 mol/L of NaOH. The student doing the titration missed the equivalence point by 0.01 mL. What would the pH of the combined solution be after 20.01 mL of sodium hydroxide is added? Note this is past the equivalence point.

12. In a titration of 30.00 mL of 0.3000 mol/L of the weak acid, ethanoic acid (HC2H3O2) with

a standardized solution of 0.3000 mol/L of the strong base, sodium hydroxide (NaOH), determine the pH of the solution during the titration just before the equivalence point, after 29.99mL of the sodium hydroxide had been titrated.

13. Calculate the pH that is formed when 0.100 mol of hydrochloric gas HCl(g) is dissolved

into a 1.00L solution of ammonia-ammonium chloride buffer containing 0.110 mol/L of NH3

and 0.110 mol/L of NH 4+ at equilibrium. Assume no change in volume of the buffer. 14. Show a derivation that proves that pH + pOH =14

ATOMIC THEORIES

1. Describe Ernest Rutherford’s atomic theory and provide a picture to illustrate his model, and provide a brief explanation of the experiments he performed that lead him to his new theory.

2. Use the Bohr model to explain how an atom can produce light.

Include a well labeled diagram in your explanation.

3. How were each of the four main quantum numbers (n, l, ml, ms) discovered? Name the scientist(s) and year involved in the discovery.

4. What are three (there are more than three…) key aspects of the quantum mechanics that can be best understood only when thinking about electrons as waves? Include at least two labeled diagrams in your answer.

5. What are the Aufbau principle, Hund’s Rule, & Pauli Exclusion Principle?

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ATOMIC THEORIES continued…

6. Explain how the Schrödinger’s cat analogy illustrates the distinctive attributes of quantum mechanics.

7. Explain how the periodic table is arranged according to the quantum theory of energy

levels, and sub-levels. 8. a) Draw the energy level diagram [with the energy levels , sub-

levels, and ↑↓ arrows] for the 33 electrons in Arsenic (As) b) Write out the electron configuration for Arsenic (As).

6. Draw the orbital pictures for the 3s, and 3p [if there are several

orbitals within the sub-level draw all of the possible configurations].

BIG PROJECTS 1. Explain the critical factors involved in producing a successful 'hot shot' that is able to produce heat, in a safe way, on command. If you could re-do your project, what would you have used, or made differently, to produce an even more successful product? 2. Explain the critical factors involved in producing a voltaic battery involving only two different metals, and some fruit or vegetable juice. If you could re-do your project, what would you have used, or made differently, to produce an even more successful product?

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Appendix Zero First Second Rate = k Rate = k[A] Rate =k[A]2 [A] = -kt + [A]0 In[A] = -kt + In[A]0

0

1 1[ ] [ ]

k tA A= ⋅ +

[A] versus t In[A] versus t 1[ ]A

versus t

Slope = -k Slope = -k Slope = k

012

[ ]2Atk

= 12

0.6931tk

= 102

1[ ]

tk A

=

log[ ]pH H += − log[ ]pOH OH −= − pH + pOH = 14

[ ] [ ] wH OH K+ −⋅ = w a bK K K= ⋅ a

acbbx2

42 −±−=

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