Download IB Chemistry Review. Unit I. Topics 2

Topic 1 Review
1. What is the sum of all coefficients when the following equation is balanced using
the smallest possible whole numbers? __ C2H2 + __O2 → __CO2 + __H2O
2. 1.7g of NaNO3 (Mr = 85) is dissolved in water to prepare 0.20 dm3 of solution.
What is the concentration of the resulting solution in mol dm–3?
3. Which sample has the greatest mass?
A. 1 mol of SO2
B. 2 mol of N2O
C. 2 mol of Ar
D. 4 mol of NH3
4. The relative molecular mass of a gas is 56 and its empirical formula is CH2. What
is the molecular formula of the gas?
5. What mass, in g, of hydrogen is formed when 3 mol of aluminium react with excess
hydrochloric acid according to the following equation?
2Al(s) + 6HCl(aq) → 2AlCl3(aq) + 3H2(g)
6. What is the total number of hydrogen atoms in 1.0 mol of benzamide,
C6H5CONH2?
7. What is the sum of the coefficients for the equation when balanced using the
smallest possible whole numbers?
__N2H4(g) + __O2(g) → __NO2(g) + __H2O(g)
8. Chloroethene, C2H3Cl, reacts with oxygen according to the equation below.
2C2H3Cl(g) + 5O2(g) → 4CO2(g) + 2H2O(g) + 2HCl(g)
What is the amount, in mol, of H2O produced when 10.0 mol of C2H3Cl and 10.0
mol of O2 are mixed together, and the above reaction goes to completion?
9. Which amount of the following compounds contains the least number of ions?
A. 2 mol of NaOH
B. 1 mol of NH4Cl
C. 2 mol of CaCl2
D. 1 mol of Al2O3
10. What is the approximate molar mass, in g mol–1, of MgSO4•7H2O?
11. Which is both an empirical and a molecular formula?
A. C5H12
B. C5H10
C. C4H8
D. C4H10
12. What is the coefficient of Fe3O4 when the following equation is balanced using the
lowest whole numbers?
__ Al(s) + __ Fe3O4(s) → __ Al2O3(s) + __ Fe(s)
13. 6.0 mol of aluminium reacts with oxygen to form aluminium oxide. What is the
amount of oxygen, in mol, needed for complete reaction?
4Al(s) + 3O2(g) → 2Al2O3(s)
14. What is the total number of nitrogen atoms in two mol of NH4NO3?
15. On analysis, a compound with molar mass 60 g mol-1 was found to contain 12 g of
carbon, 2 g of hydrogen and 16 g of oxygen. What is the molecular formula of the
compound?
16. 8.5 g of NH3 are dissolved in H2O to prepare a 500 cm3 solution. Which statements
are correct?
I. NH3 is the solute and H2O is the solution
II. The concentration of the solution is 17 g dm-3
III. [NH3] = 1.0 mol dm-3
A. I and II only
B. I and III only
C. II and III only
D. I, II and III
17. What volume of sulfur trioxide, in cm3, can be prepared using 40 cm3 sulfur
dioxide and 20 cm3 oxygen gas by the following reaction? Assume all volumes are
measured at the same temperature and pressure.
2SO2(g) + O2(g) → 2SO3(g)
18. 4.00 mol of a hydrocarbon with an empirical formula of CH2 has a mass of 280 g.
What is the molecular formula of this compound?
19. Which coefficients would balance this equation?
__ MnO2 + __ HCl → __ MnCl2 + __ Cl2 + __ H2O
20. What does (l), (g), (s), (aq) refer to in a chemical equation?
21. Understand definitions, similarity, and differences in compounds, elements, mixtures (2
types).
22. Understand definitions, similarity, and differences in solids, liquids, gases and how they
change from state to state.
IB Chemistry Review. Unit I. Topics 2 (Atomic
Structure) and 3 (Periodicity)
Potential Multiple Choice Questions
1. In the Bohr model of the atom, electrons are assigned __________.
2. According to the Heisenberg Uncertainty Principle, it is impossible to know precisely both the
position and the __________ of an electron.
3. All of the orbitals in a given electron shell have the same value of the __________ quantum
number
4. In a px orbital, the subscript x denotes the __________ of the electron.
5. At maximum, an f-subshell (or level) can hold __________ electrons, a d-subshell (or sublevel)
can hold __________ electrons, and a p-subshell (or sublevel) can hold __________ electrons.
6. Which one of the following sublevels can only hold two electrons?
7. Which electron configuration represents and/or orbital diagram would be a violation of the Pauli
Exclusion Principle? This means could you spot an incorrect answer?
8. Which one of the following is the correct electron configuration for a ground-state nitrogen atom?
9. The ground state electron configuration of Fe is __________.
10. The ground-state electron configuration of the element __________ is [Kr]5s24d5
11. Rutherford…
12. Which one of the following configurations depicts an excited oxygen atom?
13. Which electron configuration represents a violation of Hund's rule for an atom in its ground state?
14. The ground state configuration of fluorine is __________.
15. The ground state configuration of tungsten is __________
1
16. The element that has a valence configuration of 4s is _________.
17. Which of the following elements has a ground-state electron configuration different from the
predicted one?
18. How many different principal quantum numbers can be found in the ground state electron
configuration of nickel?
19. The valence shell (level) of the element X contains 2 electrons in a 5s subshell (sublevel). Below
that sublevel, element X has a partially filled 4d sublevel. What type of element is X?
20. How many quantum numbers are necessary to designate a particular electron in an atom?
21. Each p-sublevel can accommodate a maximum of __________ electrons.
22. __________-orbitals are spherically symmetrical.
23. The total number of orbitals in a shell is given by the formula __________.
24. The principal quantum number of the first d subshell is __________.
25. The n = 1 level contains __________ p orbitals. All the other shells contain __________ p
orbitals.
26. There are __________ orbitals in the second level.
27. The second level in the ground state of atomic argon contains __________ electrons.
28. In a ground-state manganese atom, the __________ sublevel is half filled.
29. What is the correct ground-state electron configuration for copper?
30. All of the __________ have a valence shell electron configuration ns1.
31. The largest principal quantum number in the ground state electron configuration of iodine is
32. Elements in the modern version of the periodic table are arranged in order of increasing
33. In which orbital does an electron in a phosphorus atom experience the greatest shielding?
34. The first ionization energies of the elements __________ as you go from left to right across a
period of the periodic table, and __________ as you go from the bottom to the top of a group in
the table.
35. In general, as you go across a period in the periodic table from left to right:
(1) the atomic radius __________;
(2) the electron affinity becomes __________ negative; and
(3) the first ionization energy __________.
36. Element M reacts with chlorine to form a compound with the formula MCl2. Element M is more
reactive than magnesium and has a smaller radius than strontium. This element is __________.
37. What is the atomic number of a neutral atom which has 51 neutrons and 40 electrons?
31 3
38. How many electrons does the ion 15
P contain?
39. The only noble gas that does not have the ns2np6 valence electron configuration is __________.
40. Electrons in the 1s sublevel are much closer to the nucleus in Ar than in He due to the larger
__________ in Ar.
41. Which species have the same number of electrons?
I.
S2–
II. Cl–
III. Ne
42. Give the correct order for atomic radius for Mg, Na, P, Si and Ar.
43. What is the electron arrangement of the Mg2+ ion?
44. The atomic radii of the alkali metals increase down a group because __________.
45. Which statement about the species 63Cu2+ and 65Cu+ is correct?
46.
47.
48.
49.
50.
Of the following elements, which has the largest first ionization energy?
Which property decreases down group 7 in the periodic table?
Which of the following is an isotope of 24
12 Mg?
Which properties of the alkali metals decrease going down group 1?
An element is in group 14 and period 3 of the periodic table. How many electrons are in the
highest occupied energy level of an atom of this element?
51. Which group on the periodic table has the lowest first ionization energies?
52. How many protons, neutrons and electrons are present in each atom of 31P?
53. Chlorine is much more apt to exist as an anion than is sodium. This is because __________.
54. Sodium is much more apt to exist as a cation than is chlorine. This is because __________.
55. Which one of the following is a metalloid?
56. This element is more reactive than lithium and magnesium but less reactive than potassium. This
element is…
57. What is the order of increasing energy of the orbitals within a single energy level?
58. Which property generally decreases across period 3?
59. Which property increases down group 1?
60. Which statements about period 3 are correct?
I. The electronegativity of the elements increases across period 3.
II.
The atomic radii of the elements decreases across period 3.
III.
The oxides of the elements change from acidic to basic across period
3.
61. Hydrogen is unique among the elements because __________.
62. Which ion below has the largest radius?
63. Which statement about the numbers of protons, electrons and neutrons in an atom is always
correct? Yes, a weird question but think of possible answers if a MC question.
64. Deduce the numbers of protons and electrons in the K+ ion.
“Written” Section
1. Deduce the numbers of protons and electrons in the K+ ion. (1)
2. Draw and label an energy level diagram for the hydrogen atom. In your diagram show how
the series of lines in the ultraviolet and visible regions of its emission spectrum are
produced.
3. The electron configuration of chromium can be expressed as [Ar]4sx3dy.
(i) Explain what the square brackets around argon, [Ar], represent. (1)
(ii) State the values of x and y. (1)
Annotate the diagram below showing the 4s and 3d orbitals for a chromium atom using
an arrow, and , to represent a spinning electron. (1)
4s
3d
4. Describe the trend of electron affinity in the periodic tab le.
5. Describe the emission spectrum of hydrogen. Outline how this spectrum is related to the
energy levels in the hydrogen atom.
6. Define the term first ionization energy. (2)
7. Distinguish between a continuous spectrum and a line spectrum.
8. Explain why the first ionization energy of magnesium is higher than that of sodium. (2)
9. Carbon and silicon belong to the same group of the periodic table.
(a) Distinguish between the terms group and period in terms of electron
arrangement. (2)
(b) State the period numbers of both carbon and silicon. (1)
10. The relative atomic mass of naturally occurring copper is 63.55. Calculate the abundances
of 63Cu and 65Cu in naturally occurring copper.
11. The periodic table shows the relationship between electron arrangement and the properties
of elements and is a valuable tool for making predictions in chemistry.
(i) Identify the property used to arrange the elements in the periodic table. (1)
(ii) Outline two reasons why electronegativity increases across period 3 in the periodic
table and one reason why noble gases are not assigned electronegativity
values. (3)
12. Answer the questions below
(i) Define the term first ionization energy of an atom. (2)
(ii) Explain the general increasing trend in the first ionization
energies of the period 3 elements, Na to Ar. (2)
13. Answer the questions below
(i)
Outline two reasons why a sodium ion has a smaller radius than a sodium atom. (2)
(ii) Explain why the ionic radius of P3– is greater than the ionic radius of Si4+. (2)
14. Define the term relative atomic mass (Ar). (1)
15. Explain why the electron affinity is positive energy (kJ/mol) for most of the noble gases.
16. Define the term mass number.
17. Explain why the first ionization energy of aluminium is lower than the first ionization
energy of magnesium. (2)
18. The graph of the first ionization energy plotted against atomic number for the first twenty
elements shows periodicity.
(i) Define the term first ionization energy and state what is meant by the term
periodicity. (2)
(ii) State the electron configuration of argon and explain why the noble gases, helium,
neon and argon show the highest first ionization energies for their respective
periods. (3)
(iii) A graph of atomic radius plotted against atomic number shows that the atomic radius
decreases across a period. Explain why chlorine has a smaller atomic radius than
sodium. (1)
(iv) Explain two reasons why a sulfide ion, S2–, is larger than a chloride ion, Cl–. (2)
(v) Explain how information from this graph provides evidence for the existence of
energy levels within atoms. (2)
IB Chemistry I. Topic 4 (Bonding) Review 2016
1. What is the correct formula of cerium(III) phosphate?
2. What is the formula of magnesium fluoride?
3. The electronegativities of four different elements are given below (the letters are not their
chemical symbols).
W
X
Y
Z
Element
0.9
1.2
3.4
4.0
Electronegativity
Based on this information which atoms could form an ionic bond?
4. What are the correct formulas of the following ions?
Ammonium
Hydrogencarbonate
Phosphate
+
2–
A.
NH4
HCO3
PO4–
B.
NH3+
HCO3–
PO43–
C.
NH4+
HCO32–
PO42–
D.
NH4+
HCO3–
PO43–
5. What happens when magnesium metal reacts with chlorine gas?
A. Each magnesium atom loses two electrons and each chlorine atom gains two electrons.
B. Each magnesium atom gains one electron and each chlorine atom loses one electron.
C. Each magnesium atom loses two electrons and each chlorine atom gains one electron.
D. Each magnesium atom gains one electron and each chlorine atom loses two electrons.
6. Which is the best description of ionic bonding?
A. The electrostatic attraction between positively charged nuclei and an electron pair
B. The electrostatic attraction between positive ions and delocalized negative ions
C. The electrostatic attraction between positive ions and delocalized electrons
D. The electrostatic attraction between oppositely charged ions
7. Which statements best describe the structure of sodium chloride, NaCl?
I. Each sodium ion is surrounded by six chloride ions.
II. The chloride ions are arranged octahedrally around each sodium ion.
III. The lattice forms a cubic structure.
8. Metal “M” has only one oxidation number (ionic charge) and forms a compound with the
formula MCO3.
Which formula is correct?
A. MNO3
B. MNH4
C. MSO4
D. MPO4
9. Describe the structure and bonding in silicon dioxide and carbon dioxide.
10. What compound is formed when lithium reacts with selenium?
11. What type of solid materials are typically hard, have high melting points and poor electrical
conductivities?
I.
Ionic
II.
Metallic
III.
Bucky balls, diamonds, silicon, silicon dioxide
12. Which particles are responsible for electrical conductivity in metals?
13. Which compound forms hydrogen bonds in the liquid state?
A. C2H5OH
B. CHCl3
C. CH3CHO
D. (CH3CH2)3N
14. Describe and compare four features of the structure and bonding in the three allotropes of
carbon: diamond, graphite, grapheme, and C60 fullerene.
15. Draw the Lewis structure of CO2 and predict its shape and bond angle.
16. Describe the structure and bonding in SiO2.
17. Explain why silicon dioxide is a solid and carbon dioxide is a gas at room temperature.
18. Describe the bonding within the carbon monoxide molecule.
19. The Lewis structure of SO2 is given below.
.. ..
: O ―S O :
..
..
What is the shape of the SO2 molecule? Bond angles?
20. Explain the electrical conductivity of molten sodium oxide.
21. Which molecule has a non-bonding pair of electrons on the central atom?
A. BF3
B. SO2
C. CO2
D. SiF4
22. When C2H2, C2H4 and C2H6 are arranged in order of increasing carbon-carbon bond
strength (weakest bond first), what is the correct order?
23. Ammonia, NH3, is a weak base. Draw the Lewis structure of ammonia and state the shape
of the molecule and its bond angles.
24. The conjugate acid of ammonia is the ammonium ion, NH4+. Draw the Lewis structure of
the ammonium ion and deduce its shape and bond angles.
25. Which molecule has the shortest carbon-oxygen bond length?
A. CH3COOH
B. CH3CH2OH
C. CO2
D. CO
26. Which statement about the bonding between carbon atoms is correct?
A. In C60 fullerene each carbon atom is covalently bonded to three other carbon
atoms.
B. In C60 fullerene each carbon atom is covalently bonded to four other carbon
atoms.
C. In graphite each carbon atom is covalently bonded to four other carbon atoms.
D. In graphite each carbon atom forms a double covalent bond with three other carbon
atoms.
27. Which intermolecular forces exist between molecules of carbon monoxide, CO?
28. Which molecule is polar/dipole?
A. CH2Cl2
B. BCl3
C. Cl2
D. CCl4
29. Which compound has a covalent macromolecular network structure?
A. MgO(s)
B. Al2O3(s)
C. P4O10(s)
D. SiO2(s)
30. Which species have a coordinate covalent bond?
I. CO
II. NH3
III. H3O+
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
31. Which order is correct when the following compounds are arranged in order of increasing
melting point?
A. CH4 < H2S < H2O
B. H2S < H2O < CH4
C. CH4 < H2O < H2S
D. H2S < CH4 < H2O
32. Ethene, C2H4, and hydrazine, N2H4, are hydrides of adjacent elements in the periodic
table.
(a)
(i)
Draw Lewis (electron dot) structures for C2H4 and N2H4 showing all valence
electrons.
(ii) State and explain the H–C–H bond angle in ethene and the H–N–H bond angle
in hydrazine.
(b)
The polarity of a molecule can be explained in terms of electronegativity.
(i) Define the term electronegativity.
(ii) Compare the relative polarities of the C–H bond in ethene and the N–H bond in
hydrazine.
(iii) Hydrazine is a polar molecule and ethene is non-polar. Explain why ethene is nonpolar.
(c)
The boiling point of hydrazine is much higher than that of ethene. Explain this
difference in terms
of the intermolecular forces in each compound.
33. Draw the Lewis structures for carbon monoxide, CO, carbon dioxide, CO2 and methanol,
CH3OH.
List, with an explanation, the three compounds in order of increasing carbon to oxygen
bond length
(shortest first).
34. Predict the shape and bond angles for the following species:
(i) CO2
(ii) CO32–
(iii) BF4–
35. State what best describes metallic bonding.
36. Explain, using diagrams, why CO and NO21- are polar molecules but CO2 is a non-polar
molecule.
37. Describe the structure and bonding in silicon dioxide.
38. PF3, SF2 and SiF4 have different shapes. Draw their Lewis structures and use the VSEPR
theory to predict the name of the shape of each molecule.
39. What is the correct order if the compounds are arranged in order of increasing boiling
point?
A. CH4 < CH3Cl < SiH4 < CH3OH
B. CH3OH < CH4 < CH3Cl < SiH4
C. CH3OH < CH3Cl < SiH4 < CH4
D. CH4 < SiH4 < CH3Cl < CH3OH
40. How many atoms is each carbon directly bonded to in each of its allotropes?
41. (i) Outline the principles of the valence shell electron pair repulsion (VSEPR) theory.
(ii) Use the VSEPR theory to deduce the shape of H3O+ and C2H4. For each species, draw
the Lewis
structure, name the shape, and state the value of the bond angle(s).
(iii) Predict and explain whether each species is polar.
(iv) Using Table 7 of the Data Booklet (electronegativity), predict and explain which of the
bonds O-H, O-N
or N-H would be most polar.
42. Predict and explain which of the following compounds consist of molecules (meaning
covalently bonded):
NaCl, BF3, CaCl2, N2O, P4O6, FeS and CBr4.
43. Diamond, graphite, grapheme, and C60 fullerene are three allotropes of carbon.
(i) Describe the structure of each allotrope.
(ii) Compare the bonding in diamond and graphite.
44. Predict the shape of a molecule of AsH3
45. If the electronegativity difference between two is greater than 1.8 or so, what type of bond will they
form?
46. What is the shape of a molecule of carbon tetrachloride, CCl4?
47. How many valence electrons will hydrogen share in a bond?
48. If a molecule has three electron domains around its center atom and two of those are bonds and pair is
non-bonded, what is the name of its shape?
49. What is a molecule called that is composed of two identical atoms covalently bonded?
a. Would it be polar?
i. Why or why not?
50. Do molecules (covalent) tend to have lower or higher melting points than ionic compounds? Why?
51. Will the molecule CF4 dissolve in water?
b. Why or why not?
52. Why do atoms share electrons in covalent bonds?
53. What causes water molecules to have a bent shape, according to VSEPR theory?
54. According to VSEPR theory, if there are three electron domains (negative charge centres) all
involved in bonds in the valence shell of an atom, they will be arranged in a(n) __________
geometry.
155. The molecular geometry of BrO2 is _________.
1-
56. The molecular geometry of the BrO3 ion is __________.
57. The molecular geometry of the left-most carbon atom in the molecule is ____. The second
carbon ______
58. An electron domain (electron centres) consists of __________
59. According to VSEPR theory, if there are three electron domains on a central atom, they will be
arranged such that the shape will be __________.
60. A molecule has the formula AB3 and the central atom is in a different plane from the surrounding
three atoms. Its molecular shape is __________.
61. Elemental iodine (I2) is actually a solid at room temperature. What is the major attractive force
that exists among different (I2) molecules in the solid?
62. Hydrogen bonding is a special case of __________.
63. Describe where the valence electrons in metals are. What properties does this give most metals?
64. How do the bond angles in CH4, NH3 and H2O compare?
65. Which combination of the characteristics of element X, a metal, and element Y, a nonmetal, is most likely to lead to ionic bonding?
A.
B.
C.
D.
66.
67.
68.
69.
70.
71.
72.
73.
74.
75.
76.
77.
78.
X
low ionization energy
low ionization energy
high ionization energy
high ionization energy
Y
high electronegativity value
low electronegativity value
high electronegativity value
low electronegativity value
Which particles are responsible for electrical conductivity in metals?
Explain why sodium oxide has a higher melting point than sulfur trioxide.
Do London forces increase or decrease going down a group on the periodic table? Explain.
As far as London forces are concerned, what shapes of molecules have a higher boiling
point?
What types of molecules or compounds can conduct electricity and why?
Explain the concept of “like dissolves like”.
Describe an alloy.
What type of molecules have the highest melting point?
Explain why the melting points of the elements decrease down group 1 and increase down
group 17.
Explain why sodium conducts electricity but phosphorus does not.
What is 2 + 2?
What is the average high school GPA of UW’s freshman this year?
What is the average high school GPA of WSU’s freshman this year?
Energetics/Thermochemistry (Topic 5) and Chemical
Kinetics (Topic 6) Review
Exothermic
Endothermic
Enthalpy
Temperature
Heat
calorimeter
1.
2.
3.
4.
Enthalpy change
Bond enthalpy
Hess’s law
Rate of reaction
Maxwell-Boltzmann Distribution
Collision theory
Activation energy
Catalyst
Average rate
Temperature
Surface area
Concentration
Be able to identify exothermic and endothermic reactions based on their sign.
a. Be able to draw a graph of each and identify all “areas” and energy.
Be able to explain the difference between heat and temperature
Have a basic understanding of coffee cup calorimeter and how it’s used.
a. Be aware of major sources of error in a calorimeter.
Be able to use the enthalpy change formula: ΔH = mcΔT (same as q = mcΔT)
5.
6.
7.
8.
9.
Know how to convert joules to kilojoules when asked to do so.
Use the bond enthalpy formula and bond enthalpy values from the table to calculate ΔH.
Be able to use Hess’s law to find ΔH of an unknown reaction when given a series of reactions..
Be able to calculate the average rate of reaction based on a change in time and change in concentration.
Be able to determine information about particles from a Maxwell-Boltzmann distribution.
a. Be able to draw a Maxwell-Boltzmann distribution curve.
i. Show how the curve “changes” when there is an increase in temperature.
ii. Show how the curve “changes” when a catalyst is used.
10. Know the basic factors involved in the collision theory.
11. Be able to explain the activation energy of a reaction and what factors affect it.
12. Be able to explain how temperature, pressure, surface area, and concentration affect reaction rate according
to the collision theory.
1. Which types of reaction are always exothermic?
2. Hydrochloric acid is reacted with large pieces of calcium carbonate, the reaction is then
repeated using calcium carbonate powder. How does this change affect the activation
energy? The collision frequency?
3. Which statement is true about using sulfuric acid (shown as H+ below) as a catalyst in the
following reaction?

H ( aq )
CH3–CO–CH3(aq) + I2(aq) 

 CH3–CO–CH2–I(aq) + HI(aq)
4. Which changes would increase the rate of the reaction below?
C4H10(g) + Cl2(g) → C4H9Cl(l) + HCl(g)
5. Which of the following can increase the rate of a chemical reaction?
6. The enthalpy change for the following reaction is -483.6 kJ:
2H2 (g)  O2 (g)  2H2O(g)
Therefore, the enthalpy change for the following reaction is __________ kJ:
- 967.2
4H2 (g)  2O2 (g)  4H2O(g)
answer =
7. A sample of aluminum metal absorbs 9.86 J of heat, upon which the temperature of the
sample increases from 23.2 °C to 30.5 °C. Since the specific heat capacity of aluminum
is 0.90 J/g-K, the mass of the sample is __________ g. answer = 1.5 g
8. Which statements describe the action of a catalyst?
9.
The specific heat capacity of lead is 0.13 J/g-K. How much heat (in J) is required to raise
the temperature of 15g of lead from 22 °C to 37 °C?
10. Using the equations below:
C(s) + O2(g) → CO2(g)
∆Hο = –394 kJ mol–1
Mn(s) + O2(g) → MnO2(s) ∆Hο = –520 kJ mol–1
What is ∆H, in kJ, for the following reaction? answer = 126 kJ mol–1
MnO2(s) + C(s) → Mn(s) + CO2(g)
11. The ΔH for the solution process when solid sodium hydroxide dissolves in water is 44.4
kJ/mol. When a 13.9-g sample of NaOH dissolves in 250.0 g of water in a coffee-cup
calorimeter, the temperature increases from 23.0 °C to __________ °C. Assume that the
solution has the same specific heat as liquid water, i.e., 4.18 J/g-K. answer = 37.0 °C
12. ΔH for the reaction: IF5 (g)  IF3 (g)  F2 (g) is __________ kJ, give the data below.
answer = 355
IF(g)  F2 (g)  IF3 (g)
ΔH = -390 kJ
IF(g)  2F2 (g)  IF5 (g)
ΔH = -745 kJ
13. Given the following reactions
ΔH = -28.0 kJ
Fe2O3 (s)  3CO(s)  2Fe(s)  3CO2 (g)
ΔH = +12.5 kJ
3Fe(s)  4CO2 (s)  4CO2 (g)  Fe3O4 (s)
The enthalpy of the reaction of Fe 2 O3 with CO is __________ kJ.
answer = - 59.0
3Fe2O3 (s)  CO(g)  CO2 (g)  2Fe3O4 (s)
14. Given the following reactions:
ΔH = -790 kJ
2S(s)  3O2 (g)  2SO3 (g)
ΔH = -297 kJ
S(s)  O2 (g)  SO2 (g)
The enthalpy of the reaction in which sulfur dioxide is oxidized to sulfur trioxide is
__________ kJ.
answer = -196
2SO2 (g)  O2 (g)  2SO3 (g)
15. Given the following reactions
ΔH = 178.1 kJ
CaCO3 (s)  CaO(s) +CO2 (g)
ΔH = -393.5 kJ
C(s, graphite)  O2 (g)  CO2 (g)
The enthalpy of the reaction is __________ kJ. answer = 571.6
CaCO3 (s)  CaO(s)  C(s, graphite)  O2 (g)
16. Given the data in the table below, H
for the reaction is __________ kJ. answer =
rxn
-130.4
Ca(OH)2  2H3AsO4  Ca(H2 AsO4 )2  2H2O
17. H for the reaction 2CO(g)  O2 (g)  2CO2 (g) is __________ kJ. answer = -566.4
18. ΔH for an endothermic process is __________ while ΔH for an exothermic process is
__________.
19. For a given process at constant pressure, ΔH is negative. This means that the process is
__________.
20. Consider the following reactions.
1
Cu2O(s) + O2(g) → 2CuO(s)
∆HO = –144 kJ
2
Cu2O(s) → Cu(s) + CuO(s)
∆HO = +11 kJ
What is the value of ∆HO, in kJ, for this reaction? answer = -155 kJ
1
Cu(s) + O2(g) → CuO(s)
2
21. Consider the two reactions involving iron and oxygen.
2Fe(s) + O2(g) → 2FeO(s)
∆HO = –544 kJ
4Fe(s) + 3O2(g) → 2Fe2O3(s)
∆HO = –1648 kJ
What is the enthalpy change, in kJ, for the reaction below? answer = - 560 kJ
4FeO(s) + O2(g) → 2Fe2O3(s)
22. What is the energy, in kJ, released when 1.00 mol of carbon monoxide is burned according
to the following equation? answer = -282 kJ
2CO(g) + O2(g) → 2CO2(g)
ΔHo = –564 kJ
23. A chemical reaction that absorbs heat from the surroundings is said to be __________
and has a __________ ΔH at constant pressure
24. The reaction 4Al(s)  3O2 (g)  2Al2O3 (s) ΔH° = -3351 kJ is __________, and
therefore heat is __________ by the reaction.
25. The units of specific heat are __________.
26. When 0.800 grams of NaOH is dissolved in 100.0 grams of water, the temperature of the
solution increases from 25.00 °C to 27.06 °C. The amount of heat absorbed by the water
is __________ J. (The specific heat of water is 4.18 J/g-°C.) Answer: 868
27. Which substance in the reaction below either appears or disappears the fastest?
4NH 3 +7O2  4NO2 +6H 2O
28. The peroxydisulfate ion (S2O82 ) reacts with the iodide ion in aqueous solution via the
reaction:
S2O82 (aq)  3I  2SO4 (aq)  I3  (aq)
An aqueous solution containing 0.050 M of S2O82- ion and 0.072 M of I  is prepared, and
the progress of the reaction followed by measuring [ I  ]. The data obtained is given in the
table below.
The average rate of disappearance of I  between 400.0 s and 800.0 s is __________ M/s. answer
= 2.8  105
The average rate of disappearance of I  in the initial 400.0 s is __________ M/s. answer = 3.8 x
10-5
29. The rate of a reaction depends on __________.
30. Methanol is made in large quantities as it is used in the production of polymers and in
fuels.
The enthalpy of combustion of methanol can be determined theoretically or
experimentally.
1
CH3OH(l) + 1 O2(g) → CO2(g) + 2H2O(g)
2
a. Using the information from Table 10 of the Data Booklet, determine the theoretical
enthalpy of combustion of methanol. answer => ∆H = –540 (kJ mol–1);
b. The enthalpy of combustion of methanol can also be determined experimentally in
a school laboratory. A burner containing methanol was weighed and used to heat
water in a test tube as illustrated below.
The following data were collected.
Initial mass of burner and methanol / g
80.557
Final mass of burner and methanol / g
80.034
Mass of water in test tube / g
20.000
Initial temperature of water / °C
21.5
Final temperature of water / °C
26.4
i. Calculate the amount, in mol, of methanol burned. answer = 0.0163 mol
ii. Calculate the heat absorbed, in kJ, by the water. answer = 0.41 kJ
iii. Determine the enthalpy change, in kJ mol–1, for the combustion of 1 mole
of methanol. answer = –25 kJ mol–1
The Data Booklet value for the enthalpy of combustion of methanol is –726 kJ mol–1.
 Suggest why this value differs from the values calculated in Parts (a) and (b).
31. Enthalpy changes can be determined using average bond enthalpies. Define the term
average bond enthalpy and discuss variations in bond enthalpies from table to table.
32. In an experiment to measure the enthalpy change of combustion of ethanol, a student
heated a copper calorimeter containing 100 cm3 of water with a spirit lamp and collected
the following data.
Initial temperature of water: 20.0 °C
Final temperature of water: 55.0 °C
Mass of ethanol burned:
1.78 g
Density of water:
1.00 g cm–3
o Use the data to calculate the heat evolved when the ethanol was combusted.
answer = 14.6 kJ
o Calculate the enthalpy change of combustion per mole of ethanol. answer = –
378 kJ mol–1
o Suggest two reasons why the result is not the same as the value in the Data
Booklet.
33. Explain and label the enthalpy level diagram below.
34. Consider the reaction between gaseous iodine and gaseous hydrogen.
I2(g) + H2(g)
2HI(g) ∆HO = –9 kJ
Why do some collisions between iodine and hydrogen not result in the formation of the
product?
35. Factors that affect the rate of a chemical reaction include particle size, concentration of
reactants and the temperature of the reaction.
 Define the term rate of a chemical reaction.
 Discuss the above three characteristic properties of reactant particles which
affect the rate of reaction as described by the collision theory.
36. Below, sketch two Maxwell-Boltzmann energy distribution curves (in the same graph) for
the same sample of gas, one at a temperature T and another at a higher temperature T  .
Label both axes. Explain why raising the temperature increases the rate of a chemical
reaction.
37. The graph below shows how the volume of carbon dioxide formed varies with time when a
hydrochloric acid solution is added to excess calcium carbonate in a flask.
Explain the shape of the curve to the right.
38. Define the term activation energy, Ea.
39. State two conditions necessary for a reaction to take place between two reactant particles.
40. Sketch an enthalpy level diagram to describe the effect of a catalyst on an exothermic
reaction.
41. Graphing is an important method in the study of the rates of chemical reaction. Sketch a
graph to show how the reactant and product concentration change with time in a typical
chemical reaction taking place in solution. Show how the rate of the reaction at a particular
time can be determined for either the reactant or product.
42. A solution of hydrogen peroxide, H2O2, is added to a solution of sodium iodide, NaI,
acidified with hydrochloric acid, HCl. The yellow colour of the iodine, I2, can be used to
determine the rate of reaction.
H2O2(aq) + 2NaI(aq) + 2HCl(aq) → 2NaCl(aq) + I2(aq) + 2H2O(l)
The experiment is repeated with some changes to the reaction conditions. For each of
the changes that follow, predict, stating a reason, its effect on the rate of reaction.
 The concentration of H2O2 is increased at constant temperature.
 The solution of NaI is prepared from a fine powder instead of large
crystals.
43. Explain why the rate of a reaction increases when the temperature of the system increases.
IB Chemistry Review Topics 7 & 8
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
Bronsted-Lowry acid/
Chemical equilibrium
Hydrogen ion
Molarity
base
Conjugate acid/ base
Equilibrium constant
Disassociate
Activation energy
Kc
Le Chatelier’s Principle
Donor
Acid deposition
Strong acids/ base
Catalyst
Acceptor
Acid rain
Weak acids/ base
Haber Process
Concentration
Arrhenius
pH
Contact process
Endothermic
alkaline
System
Proton
Exothermic
Tia
Know how to determine whether something is Bronsted-Lowry acid or base.
Know how to determine the conjugate base of an acid and the conjugate acid of a base.
Explain what makes something a strong acid or strong base.
Be able to recognize the assigned strong and weak acids and bases.
Explain what equilibrium has to do with weak acids and weak bases.
Explain why a weak acid or base will have a different pH then a strong acid or base.
Explain the relationship between forward and reverse reactions and a double-ended arrow.
Be able to write an equilibrium constant from a balanced equation.
Be able to explain endothermic and exothermic reactions in terms of energy as a reactant or product.
Be able to explain how Le Chatelier’s Principle is related to pressure/volume, temperature and amounts of products
and reactants in a reaction.
Be able to explain how a catalyst affects a reaction in equilibrium.
Be able to explain what factors are involved in the Haber process and their connection to Le Chatelier’s Principle.
Be able to explain what factors are involved in the Contact process and their connection to Le Chatelier’s Principle.
1. The equilibrium between nitrogen dioxide, NO2, and dinitrogen tetroxide, N2O4, is shown below.
2NO2(g)
N2O4(g)
Kc = 0.01
What happens when the volume of a mixture at equilibrium is decreased at a constant temperature?
2. Which does dynamic imply about chemical equilibria?
3. At equilibrium… __________.
4. Which one of the following will change the value of an equilibrium constant?
5. For the following reaction Kc = 1.0 × 10–5 at 30 °C. 2NOCl(g)
2NO(g) + Cl2(g)
Which relationship is correct at equilibrium at this temperature?
6. Consider the equilibrium between methanol, CH3OH(l), and methanol vapor, CH3OH(g).
CH3OH(l)
CH3OH(g)
What happens to the position of equilibrium and the value of Kc as the temperature decreases?
7. What is the equilibrium constant expression, Kc, for the following reaction? Remember, you don’t
actually need a number, just show how to get it.
N2O4(g)
2NO2(g)
8. Consider the endothermic reaction below.
5CO(g) + I2O5(g)
5CO2(g) + I2(g)
According to Le Chatelier’s principle, which change would result in an increase in the amount of
CO2?
9. What is the effect of an increase of temperature on the yield and the equilibrium constant for the
following reaction?
2H2(g) + CO(g)
CH3OH(l)
∆HO = –128 kJ
10. Consider the following equilibrium reaction.
2SO2(g) + O2(g)
2SO3(g)
∆Ho = –197 kJ
Which change in conditions will increase the amount of SO3 present when equilibrium is reestablished?
11. Consider the following reversible reaction.
Cr2O72–(aq) + H2O(l)
2CrO42–(aq) + 2H+(aq)
What will happen to the position of equilibrium and the value of Kc when more H+ ions are added at
constant temperature?
12. What effect will an increase in temperature have on the Kc value and the position of equilibrium in
the following reaction?
N2(g) + 3H2(g)
2NH3(g)
ΔH = –92 kJ
13. An increase in temperature increases the amount of chlorine present in the following equilibrium.
PCl5(s)
PCl3(l) + Cl2(g)
What is the best explanation for this (keep in mind kinetics as well)?
14. What will happen when at a constant temperature if more iodide ions, I–, are added to the
equilibrium below?
I2(s) + I–(aq)
I3–(aq)
15. What changes occur when the temperature is increased in the following reaction at equilibrium?
Include both the direction of the shift and how Kc is affected.
Br2(g) + Cl2(g)
2BrCl(g) ∆Hο = +14 kJ mol–1
16. Methanol may be produced by the exothermic reaction of carbon monoxide gas and hydrogen gas.
CO(g) + 2H2(g)
CH3OH(g) ∆HO = –103 kJ
 State the equilibrium constant expression, Kc, for the production of methanol.
State and explain the effect of changing the following conditions on the amount of methanol
present at equilibrium:
 increasing the temperature of the reaction at constant pressure.
 increasing the pressure of the reaction at constant temperature.
 The conditions used in industry during the production of methanol are a temperature of 450
°C and pressure of up to 220 atm. Explain why these conditions are used rather than those
that could give an even greater amount of methanol.
17. An example of a homogeneous reversible reaction is the reaction between hydrogen and iodine.
H2(g) + I2(g)
2HI(g)
 Outline the characteristics of a homogeneous chemical system that is in a state of
equilibrium.
 Deduce the expression for the equilibrium constant, Kc.
 Predict what would happen to the position of equilibrium and the value of Kc if the
pressure is increased from 1 atm to 2 atm.
18. A state of equilibrium can exist when a piece of copper metal is placed in a solution of copper(II)
sulfate. Outline the characteristics of a chemical system in dynamic equilibrium.
19. For an exothermic reaction state how an increase in temperature would affect both Kc and the
position of equilibrium.
20. In carbonated drinks containing dissolved carbon dioxide under high pressure, the following
dynamic equilibrium exists. CO2(aq)
CO2(g)
Describe the effect of opening a carbonated drink container and outline how this equilibrium is
affected.
21. Consider the following equilibrium.
2SO2(g) + O2(g)
2SO3(g)
ΔHo = –198 kJ mol–1
 Deduce the equilibrium constant expression, Kc, for the reaction.
 State and explain the effect of increasing the temperature on the yield of sulfur trioxide.
 State the effect of a catalyst on the value of Kc.
 State and explain the effect of a catalyst on the position of equilibrium.
22. Consider the following reaction taking place at 375 °C in a 1.00 dm3 closed container.
Cl2(g) + SO2(g)
SO2Cl2(g)
∆HO = –84.5 kJ
 Deduce the equilibrium constant expression, Kc, for the reaction.
 If the temperature of the reaction is changed to 300 °C, predict, stating a reason in each
case, whether the equilibrium concentration of SO2Cl2 and the value of Kc will increase or
decrease.
23. The equation for the main reaction in the Haber process is: N2(g) + 3H2(g)
2NH3(g)
∆H



State and explain the effect on the equilibrium yield of ammonia with increasing the pressure and the
temperature.
In practice, typical conditions used in the Haber process involve a temperature around 450 °C and a
pressure of 250 atm. Explain why these conditions are used rather than conditions that would give a
higher yield.
A chemist claims to have developed a new catalyst for the Haber process, which increases the yield of
ammonia. State the catalyst normally used for the Haber process, and comment on the claim made by
this chemist.
24. Consider the equilibrium below.
CH3CH2COOH(aq) + H2O(l)
CH3CH2COO–(aq) + H3O+(aq)
Which species represent a conjugate acid-base pair?
25. The pH of a solution changes from pH = 2 to pH = 5. What happens to the concentration of the
hydrogen ions during this pH change?
1. What is formed when a metal oxide reacts with an acid?
2. What is formed when a metal carbonate reacts with an acid?
3. What is formed when a metal reacts with and acid?
4. Which species behave as Brønsted-Lowry acids in the following reversible reaction?
H2PO4–(aq) + CN–(aq)
HCN(aq) + HPO42–(aq)
5. Which of the following are weak acids in aqueous solution?
I.
CH3COOH
II.
H2CO3
III. HCl
6. What is the conjugate base of H2CO3 according to the Brønsted-Lowry theory?
7. For equal volumes of 1.0 mol dm–3 solutions of hydrochloric acid, HCl(aq), and methanoic
acid, HCOOH(aq), which statements are correct?
I. HCl dissociates more than HCOOH
II. HCl is a better electrical conductor than HCOOH
III. HCl will neutralize more NaOH than HCOOH
8. A Br∅nsted-Lowry base is defined as a substance that __________
9. A substance that is capable of acting as both an acid and as a base is __________.
10. The molar concentration of hydronium ion (or H+) in pure water at 25°C is __________.
11. Water has a very low Kc (also referred to as Kw). This indicates that __________.
12. What would the pH be of a solution with a high concentration of hydroxide ions?
13. Nitric acid is a strong acid. This means that __________.
14. Of the following acids, __________ is not a strong acid.
15. What is the conjugate acid of NH3?
16. The conjugate base of HSO4- is __________.
17. The conjugate acid of HSO4- is __________.
18. Which list contains only strong acids?
19. What is the formula of the conjugate base of the hydrogenphosphate ion, HPO42–?
20. Define the terms acid and base according to the Brønsted-Lowry theory and state one example
of a weak acid and one example of a strong base.
21. Describe two different methods, one chemical and one physical, other than measuring the pH,
that could be used to distinguish between ethanoic acid and hydrochloric acid solutions of the
same concentration.
22. Black coffee has a pH of 5 and toothpaste has a pH of 8. Identify which is more acidic and
deduce how many times the [H+] is greater in the more acidic product.
23. Define the terms acid and base according to the Brønsted-Lowry theory. Distinguish between a
weak base and a strong base. State one example of a weak base.
24. Describe two different properties that could be used to distinguish between a 1.00 mol dm–3
solution of a strong monoprotic acid and a 1.00 mol dm–3 solution of a weak monoprotic acid.
25. Explain, using the Brønsted-Lowry theory, how water can act either as an acid or a base. In
each case identify the conjugate acid or base formed.
26. Which 0.10 mol dm-3 solution would have the highest conductivity- HCl or CH3COOH?
Explain.
27. What is the ionic product constant for pure water (Kw) at 298K?
28. Calculate the pH of at 4.2 X 10-3 M HCl solution?
29. A solution of milk has a pH of 6.70. Calculate [H+].
30. Calculate the pH of at 11.2 X 10-10 M HCl solution?
31. Calculate the hydrogen ion concentration of a solution with a pH of 2.82
32. Calculate the hydrogen ion concentration of a solution with a pH of 10.98
33. Calculate the pH of 1.0 x 10-2 mol dm-3 solution of sodium hydroxide.
34. Define acid deposition.
35. What two are the two major pollutants that cause acid deposition?
36. Acid rain is results primarily from what two acids being formed?
37. Show the equations for how an oxide of nitrogen can produce nitric acid (HNO3).
38. Show the equations for how the combustion of sulfur can produce sulfuric acid (H2SO4).
39. State the typical pH of rain water. Of acid rain?
IB Chemistry I. Topic 9.1 Review.
1.
2.
3.
4.
5.
What are transferred in an oxidation-reduction reaction?
In the reaction of sodium with oxygen, which atom is reduced?
In the reaction of sodium with oxygen, which atom is the reducing agent?
In the reaction of calcium with chlorine, which atom is oxidized?
In the reaction of calcium with chlorine, which atom is the oxidizing agent?
6. Which species could be reduced to form NO2?
7. If an atom is reduced in a redox reaction, what must happen to another atom in the system?
8. Which ion can be most easily reduced?
Cu
Ni
Zn
Ca
9. What is the reducing agent in the following reaction?
2Na + 2H O  2NaOH + H
10. What is the reducing agent in the following reaction?
2Na + S  Na S
11. What happens to the manganese in the following reaction?
2MnO4–(aq) + 5H2O2(aq) + 6H+(aq) → 2Mn2+(aq) + 8H2O(l) + 5O2(g)
12. What happens (oxidized/reduced, oxidation # increase/decrease) to iodine when
iodate ions, IO3–, are converted to iodine molecules, I2?
13. When iron oxide becomes iron, what type of reaction occurs?
14. Identify the pair of metals that lists the more easily oxidized metal on the left.
Ag, Na
Ca, Al
Fe, K
K, Li
15. Which definition of oxidation is correct?
16. What is the oxidation half-reaction for the following unbalanced redox equation?
Cr O
+ Fe  Cr + Fe
17. What is the reduction half-reaction for the following unbalanced redox equation?
Cr O
+ NH  Cr O + N
18. Which element increases its oxidation number in the following reaction?
2Na + 2H O  2NaOH + H
19. In the following unbalanced reaction, which atom is reduced?
H O + Cl + SO
HCl + H SO
20. Identify the atom that increases in oxidation number in the following redox reaction.
2MnO + 2K CO + O
2KMnO + 2CO
21. What is the IUPAC name of Fe2O3?
22. Balance and state the oxidizing agent and reducing agent
Mn 2+ + NaBiO3  Bi 3+ + MnO4 1- + Na 1+
23. Consider the following three redox reactions.
Cd(s) + Ni2+(aq) → Cd2+(aq) + Ni(s)
Ni(s) + 2Ag+(aq) → Ni2+(aq) + 2Ag(s)
Zn(s) + Cd2+(aq) → Zn2+(aq) + Cd(s)
 Deduce the order of reactivity of the four metals, cadmium, nickel, silver and
zinc and list in order of decreasing reactivity.
 Identify the best oxidizing agent and the best reducing agent.
24. In which species does sulfur have an oxidation number of 0?
25. What is the reducing agent in the reaction below?
2MnO4–(aq) + Br–(aq) + H2O(l) → 2MnO2(s) + BrO3–(aq) + 2OH–(aq)
26. Consider the following reaction.
MnO4–(aq) + 8H+(aq) + 5Fe2+(aq) → Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)
 What reactant is the oxidizing agent and gains electrons.
27. Fertilizers may cause health problems for babies because nitrates can change into
nitrites in water used for drinking.
 Define oxidation in terms of oxidation numbers.
 Deduce the oxidation states of nitrogen in the nitrate, NO3–, and nitrite, NO2–,
ions.
28. Nitric acid reacts with silver in a redox reaction.
__ Ag(s) + __ NO3–(aq) + ____ → __ Ag+(aq) + __ NO(g) + ____
 Using oxidation numbers, deduce the complete balanced equation for the
reaction showing all the reactants and products.
29. Define oxidation in terms of electron transfer.
30. Chlorine can be made by reacting concentrated hydrochloric acid with potassium
manganate(VII), KMnO4.
2KMnO4(aq) + 16HCl(aq) → 2MnCl2(aq) + 2KCl(aq) + 5Cl2(aq) + 8H2O(aq)
 State the oxidation number of manganese in KMnO4 and in MnCl2.
 Deduce which species has been oxidized in this reaction and state the change in
oxidation number that it has undergone.
31. Which species is oxidized in the following reaction?
2Ag+(aq) + Cu(s) → 2Ag(s) + Cu2+(aq)
32. Which are redox reactions?
I.
2FeCl2 + Cl2 → 2FeCl3
II.
Mg + 2HNO3 → Mg(NO3)2 + H2
III.
H2O + SO3 → H2SO4
33. What is the reducing agent in this reaction?
Cu(s) + 2 NO 3 (aq) + 4H+(aq) → Cu2+(aq) + 2NO2(g) + 2H2O(l)
34. Organic matter present in water can be decomposed under aerobic and anaerobic
conditions by bacteria. Identify the product, different in each case, when compounds
containing the following elements are subjected to aerobic and anaerobic conditions.
Element
Aerobic decomposition
Anaerobic decomposition
Carbon
Nitrogen
Sulfur
Phosphorus
35. Water purity is often assessed by reference to its oxygen content.
 Outline the meaning of the term biochemical oxygen demand (BOD).
 Describe how the dissolved oxygen concentration in a river would decrease if
o a car factory releases warm water into the river after using it for cooling
o a farmer puts large quantities of a fertilizer on a field next to the river.
36. An iodine/thiosulfate redox titration is carried out to measure the dissolved oxygen present
in a water sample:
 A 263 cm3 water sample is taken and saturated with oxygen for five days
 The following reactions (don’t need to know) are related to the Winkler method
o 2 Mn2+ (aq) + O2 (aq) + 4OH– → 2 MnO(OH)2(s)
o MnO(OH)2(s) + 2 I–(aq) + 4H+ → Mn2+(aq) + I2(aq) + 3H2O
o 2 S2O32-(aq) + I2 (aq) → S4O62-(aq) + 2 I– (aq)
 It was found that 2.7 cm3 of a 0.0584 mol dm-3 solution of sodium thiosulfate
(Na2S2O3) was required to react with the iodine produced.
o What is the concentration (molarity) of oxygen in the water sample (1.4 x 10-4)
o What is the concentration (g dm-3) in the water sample? (4.8 x 10-3)
o What is concentration (in ppm) of oxygen in the water sample? (4.79)
Topic 10.1 and 10.2 (alkanes and alkenes only) Review
1.
2.
3.
4.
5.
6.
7.
8.
9.
Give the structural formula and the condensed structural formula of 2,3-dibromo-2methylhexane.
Compare and contrast empirical, molecular, and structural formulas.
Which of the following are isomers of each other?
 CH3(CH2)3CH3
 (CH3)2CHCH3
 CH3CH(CH3)CH2CH3
Give the IUPAC name of CH3CH2CH(CH3)CH3
Give the structural formula and the condensed structural formula of pentan-1-ol?
How many isomers of pentane exist (no cyclic structures)?
Give the structural formula of propan-1,2,3- triol.
Give the structural formula of butandianal.
Is the following a homologous series? Explain your answer.
CH3CH2CH(OH)CH3,
CH3CH2CH2CH2OH,
(CH3)3COH
What is the best definition of structural isomers?
Give the functional group found in amines.
The general formula RCHO if for…
Give the structural formula and the condensed structural formula of pentan-3-one?
Distinguish between a primary, secondary, and tertiary carbons in halogenalkanes.
Draw the skeletal structure of propan-1-ol.
Describe the characteristics of a homologous series.
What is more soluble in water, alkanes or carboxylic acids and why?
A carbonyl functional group indicates what class?
A hydroxyl functional group indicates what class?
What has a higher boiling point, methane or hexane and why?
What is the general formula of an alkane? Alkene? Alkyne? Carboxylic acids?
Distinguish between saturated and unsaturated carbon bonds.
Write the balanced equation for the combustion of propane.
Write the balanced equation for the incomplete combustion of propane.
Explain the reaction of an alkane with halogens in terms of free-radical substitution
mechanism involving homolytical fusion.
What type of energy must be added to alkanes in order for them to react?
What is a free radical?
The addition of hydrogen to unsaturated hydrocarbons is called…
Explain physical and chemical evidence for the structure of benzene.
Write the balanced equation for the reaction of ethene with hydrogen gas.
Draw the structural formula for the reaction of ethene with chlorine gas to create a
dihalogenalkane.
Draw the structural formulas for the reaction of ethene with HBr to create a
halogenalkane.
Write the balanced equation for the reaction of ethene with water to create an alcohol.
What is used to test for the presence of alkenes and what does it do?
Draw the structural formulas to show the polymerization of two ethenes. State the
name of the new polymer.
Is an apple organic?
Give the structural formula and the condensed structural formula of 2,3-dibromo-2methylhexane.
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Topic 11.1 and 11.2 Review
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Qualitative vs. quantitative data
Uncertainty in analog vs. digital equipment
Precision vs. Accuracy
Sig figs, addition/subtraction vs. multiplication/division rules
Experimental errors. Systematic vs. random
Absolute vs. relative uncertainty. Propagation of uncertainty (addition/subtraction vs.
multiplication/division rules)
7. Graphing- variables, axis, positive vs. negative correlation, slope, scatter plot,
extrapolation, “best fit” line,