IB Chemistry Review. Unit I. Topics 2, 3, and 12.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
In the Bohr model of the atom, electrons are assigned __________.
According to the Heisenberg Uncertainty Principle, it is impossible to know precisely both the position and the
__________ of an electron.
All of the orbitals in a given electron shell have the same value of the __________ quantum number
In a px orbital, the subscript x denotes the __________ of the electron.
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.
Which one of the following orbitals can hold two electrons?
Which electron configuration represents a violation of the Pauli Exclusion Principle? This means would you know
what to look for?
Which one of the following is the correct electron configuration for a ground-state nitrogen atom?
The ground state electron configuration of Fe is __________.
The ground-state electron configuration of the element __________ is [Kr]5s24d5
The ground-state electron configuration of __________ is [Ar] 4s 14d5 (yes, 4s1 is correct, this atom is excited).
Which one of the following configurations depicts an excited oxygen atom?
Which electron configuration represents a violation of Hund's rule for an atom in its ground state?
The ground state configuration of fluorine is __________.
The ground state configuration of tungsten is __________
1
The element that has a valence configuration of 4s is _________.
Which of the following elements has a ground-state electron configuration different from the predicted one?
How many different principal quantum numbers can be found in the ground state electron configuration of nickel?
The valence shell (level) of the element X contains 2 electrons in a 5s subshell (sublevel). Below that shell,
element X has a partially filled 4d subshell. What type of element is X?
How many quantum numbers are necessary to designate a particular electron in an atom?
Each p-subshell can accommodate a maximum of __________ electrons.
__________-orbitals are spherically symmetrical.
The total number of orbitals in a shell is given by __________.
The principal quantum number of the first d subshell is __________.
The n = 1 level contains __________ p orbitals. All the other shells contain __________ p orbitals.
There are __________ orbitals in the second level
The second shell in the ground state of atomic argon contains __________ electrons.
In a ground-state manganese atoms, the __________ sublevel is half filled.
What is the correct ground-state electron configuration for copper?
All of the __________ have a valence shell electron configuration ns1.
The largest principal quantum number in the ground state electron configuration of iodine is
Elements in the modern version of the periodic table are arranged in order of increasing
In which orbital does an electron in a phosphorus atom experience the greatest shielding?
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.
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 __________.
Element M reacts with chlorine to form a compound with the formula MCl 2. Element M is more reactive than
magnesium and has a smaller radius than strontium. This element is __________.
Oxides of most nonmetals combine with water to form __________.
What is the atomic number of a neutral atom which has 51 neutrons and 40 electrons?
31
3
39. How many electrons does the ion 15 P contain?
40. The only noble gas that does not have the ns2np6 valence electron configuration is __________.
41. Electrons in the 1s subshell are much closer to the nucleus in Ar than in He due to the larger __________ in Ar.
42. Which species have the same number of electrons?
2–
I.
S
–
II. Cl
III. Ne
43. Give the correct order for atomic radius for Mg, Na, P, Si and Ar.
2+
44. What is the electron arrangement of the Mg ion?
45. The atomic radii of the alkali metals increase down a group because __________.
63 2+
65 +
46. Which statement about the species Cu and Cu is correct?
47. Of the following elements, which has the largest first ionization energy?
48. Which property decreases down group 7 in the periodic table?
49. Which of the following is an isotope of 24
12 Mg?
50. Which properties of the alkali metals decrease going down group 1?
51. An element is in group 4 and period 3 of the periodic table. How many electrons are in the highest occupied
energy level of an atom of this element?
52. Which oxides produce an acidic solution when added to water?
I.
P4O10
II.
MgO
SO3
III.
53. Which group on the periodic table has the lowest first ionization energies?
54. Which of the following has the largest second ionization energy?
31
55. How many protons, neutrons and electrons are present in each atom of P?
56. Chlorine is much more apt to exist as an anion than is sodium. This is because __________.
57. Sodium is much more apt to exist as a cation than is chlorine. This is because __________.
58. Which one of the following is a metalloid?
59. This element is more reactive than lithium and magnesium but less reactive than potassium. This element is…
60. What is the order of increasing energy of the orbitals within a single energy level?
61. Which property generally decreases across period 3?
62. Which property increases down group 1?
63. Which change explains why the boiling points of the halogens increase as their molecular masses increase?
64. 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.
65. Hydrogen is unique among the elements because __________.
66. Which ion below has the largest radius?
2+
67. What is the electron arrangement (this is simpler than the electron configuration) of the Mg ion?
68. 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.
+
69. Deduce the numbers of protons and electrons in the K ion.
+
70. Deduce the electron arrangement for the K ion.
71. Values for the successive ionization energies for an unknown element are given in the table below.
 In which group of the periodic table would the unknown element be found?
First ionization
Second ionization
Third ionization
–1
–1
–1
energy / kJ mol
energy / kJ mol
energy / kJ mol
420
3600
4400
72. What happens when sodium is added to water?
I. A gas is evolved
II.
The temperature of the water increases
III.
A clear, colourless solution is formed
Fourth ionization
–1
energy / kJ mol
5900
“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, clearly
labeling each series. (5)
3. Deduce the electron arrangement (this is simpler than the electron configuration) for the K+ ion. (1)
4. 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
5. The graph below represents the successive ionization energies of sodium. The vertical axis plots log
ionization energy instead of ionization energy to allow the data to be represented without using an
unreasonably long vertical axis.
State the full electron configuration of
sodium and explain how the successive
ionization energy data for sodium are related
to its electron configuration. (4)
6. Magnesium is the eighth most abundant element in the earth’s crust. The successive ionization
energies of the element are shown below.
(i) Define the term first ionization energy and state the equation for the first ionization of
magnesium. (3)
(ii) Explain the general increase in successive ionization energies of the element. (2)
(iii) Explain the large increase between the tenth and eleventh ionization energies. (3)
7. 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) Explain how information from this graph provides evidence for the existence of main
energy levels and sub-levels within atoms. (4)
(iii) State what is meant by the term second ionization energy. (1)
(iv) Sketch and explain the shape of the graph obtained for the successive ionization energies
of potassium using a logarithmic scale for ionization energy on the y-axis against
number of electrons removed on the x-axis. (4)
8. Describe the emission spectrum of hydrogen. Outline how this spectrum is related to the energy
levels in the hydrogen atom.
9. Define the term first ionization energy. (2)
10. Silicon has three stable isotopes, 28Si, 29Si and 30Si. The heaviest isotope, 30Si, has a percentage
abundance of 3.1 %. Calculate the percentage abundance of the lightest isotope to one decimal
place. The relative atomic mass of silicon is 28.09.
11. State three factors that affect the degree of deflection of ions in a mass spectrometer.
12. Distinguish between a continuous spectrum and a line spectrum.
13. Explain why the first ionization energy of magnesium is higher than that of sodium. (2)
14. 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)
15. With reference to electronic arrangements, suggest why the reaction between rubidium and water
is more vigorous than that between sodium and water. (2)
16. The relative atomic mass of naturally occurring copper is 63.55. Calculate the abundances of 63Cu
and 65Cu in naturally occurring copper.
17. Explain why the melting points of the elements decrease down group 1 and increase down group
7. (2)
18. 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)
19. 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)
(iii) Explain why sodium conducts electricity but phosphorus does
not. (2)
20. 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)
21. Identify a radioactive isotope of carbon and state one of its uses.
22. Which quantities are the same for all neutral atoms of chlorine?
23. (a) Define the term relative atomic mass (Ar). (1)
(b) Relative atomic masses are obtained using a mass spectrometer. Draw a simple annotated
diagram of the mass spectrometer. (5)
24. Explain why the first ionization energy of aluminium is lower than the first ionization energy of
magnesium. (2)
25. The graph of the first ionization energy plotted against atomic number for the first twenty
elements shows periodicity. However, it is really not necessary to answer the questions below but
it was on the IB question.
(i) Define the term first ionization energy and state what is meant by the term periodicity. (2)
(ii) State the electron arrangement 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 why a sulfide ion, S2–, is larger than a chloride ion, Cl–. (1)
(v) Explain why the melting points of the Group 1 metals (Li → Cs) decrease down the group
whereas the melting points of the Group 7 elements (F → I) increase down the group.(2)