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Summary of know/edge
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SUMMARY OF
KNOWLEDGE
C Atoms consist Of a nucleus surrounded by one or more electron shells.
[2 Atoms contain three sub-atomic particles: electrons, protons and neutrons.
[J Protons and neutrons occupy the nucleus of the atom. The electrons move round the
nucleus in orbits.
9 Protons are- positively charged and have-a mass almost identical to neutrons. Neutrons
are not electrically charged. Electrons carry a negative charge equal in size to that of the
proton.
D The sub-atomic particles can be distinguished by their behaviour in electric and magnetic
fields.
w Atoms are electrically neutral due to the presence of equal numbers of protons and
electrons.
67
68
A TOMle STRuaURE
-. The atomic number (Z) is the number of protons 'in an atom of a chemical element The
mass number (A) is the number of protons and neutrons in the nucleus in an atom of a
chemical element.
~ A nuclide is an atom with a specific number of protons and neutrons. Nuclides are
described by the notation 1X, where Z represents the atomic number, A represents the
mass number and X represents the symbol of the chemical element.
Positive ions are formed when atoms lose electrons; negative ions are formed when
atoms gain electrons.
Isotopes of a chemical element have the same atomic number but different mass
numbers. Isotopes of a chemical element have the same chemical properties, but slightly
different physical properties.
A mass spectrometer is used to show the isotopic composition of a chemical element and
determine its relative atomic mass.
The relative atomic mass is the mass of a weighted average of a sample of atoms
compared to one twelfth of the mass of a carbon-12 atom. The relative isotopic mass is
the mass of a nuclide compared to one twelfth of the mass of a carbon-12 atom.
A mass spectrometer ionizes gaseous atoms and then deflects a beam of its positive ions.
A variable magnetic field is used to deflect the ions and bring them to a detector. Varying
the magnetic field strength brings ions of higher mass to the detector, creating a mass
spectrum.
..
The isotopes of some chemical elements are radioactive and release ionizing radiation.,>
These isotopes are known as radioisotopes.
The three types of ionizing radiation are alpha radiation (helium nuclei), beta radiation
(electrons) and gamma radiation (gamma rays).
Radioactive decay is an exponential process with a characteristic half-life, which is the
time for the rate of radioactive decay to decrease by half. The half-life is independent of
the amount of the isotope.
i.;
Radioisotopes have many uses in medicine (as tracers and radiotherapy) and dating
techniques, for example, radiocarbon dating. The uses of radioisotopes depend on their
half-life and the radiation they release.
Light is a form of energy and can be regarded as having the properties of an
electromagnetic wave and a particle. Waves are described by their wavelength, frequency
and speed.
A continuous spectrum contains all wavelengths from a band of the electromagnetic
spectrum.
Each line in an emission spectrum of an atom or chemical element corresponds to
electrons transitioninq from one energy level to another. During electron transitions
electromagnetic radiation is emitted or absorbed.
The lines in an emission or line spectrum become closer together as wavelength
decreases (or frequency increases).
The Bohr model of the atom assumes that electrons rotate in circular orbits around the
nucleus. The electrons in each orbit have a fixed amount of potential energy.
The frequency of the electromagnetic radiation and the gap between two energy levels
is given by Planck's equation: !'!f == hi, where f represents the frequency and h represents
Planck's constant.
The ground state is the state of an atom in which all electrons have their lowest energies.
If electrons are given additional energy and move from a lower to a higher state, then
the atom is said to be excited.
The Balmer series is caused by excited electrons falling back to the second energy level.
The Lyman series is caused by excited electrons falling back to the first energy level.
The ionization energy for the hydrogen atom is the energy needed to remove the
electron from the ground state of the gaseous atom.
·'1
L.
-------··.11
BD
I
Examination questions
69
Examination questions - a selection
•
Paper 1 IB questions and IB style questions
01 Which statement is correct about the isotopes of
an element?
A They have the same mass number.
B They have the same numbers of protons and
neutrons in the nucleus.
e They have more protons than neutrons.
D They have the same electron arrangementor
confiquration.
02
A chemical element with the symbol X has the
electron arrangement 2,8,6. Which chemical
species is this chemical element most likely to form?
A the ion X3+
e the compound HzX [2Wxz-J
B the ion X6+
D the compound XFa[X8+8F-]
Q3 Which of the following particles contains more
electrons than neutrons?
I
1H
A I only
B II only
II
T~CI-
III
f§K+
e
I and II only
D II and III only
Standard Level Paper 1, May 00, 06
Q4 What information about the structure of a helium
atom can be gained from its emission spectrum?
A Most of the mass of the atom is in its nucleus.
B A helium atom contains two electrons and two
protons.
e The electrons in the helium atom are held near
the nucleus.
D The electrons may exist in any of several energy
levels.
Q5 An element has the electron arrangement 2,8,6.
What is the element?
ACe
S
B P
D Ar
Q6 Which is an incorrect statement about the atomic
emission spectrum of hydrogen?
A The frequency of each line depends on the
difference in energy between the higher and
lower energy levels.
B The spectrum consists of several series of lines.
e Electronic transitions to the leve~·n = 2 give rise
to lines in the visible region.
D It is a continuous spectrum.
Q7 What is the correct number of each particle in a
fluoride ion, 19F-?
Protons Neutrons
A
9
10
B
9
10
e
9
10
D
9
Electrons.
8
9
10
19
10
Standard Level Paper 1, Nov 03, 05
Q8 What fraction of a radioisotope will remain after
three half-lives?
A 1/16
B 1/8
e
1/3
D 3/4
Q9 Why was the Bohr theory of the atom developed?
A To account for changes in gas volumes with
temperature.
B To account for the ratios by mass of elements
in compounds.
To account for the emission or line spectrum of
hydrogen atoms.
D To account for chemical formulas.
e
Q10 A particular element consists of two isotopes: 72%
of mass number 85 and 28% of mass number 87.
What is the expected range of the relative atomic
mass?
A less than 85
B between 86 and 87
betvveen 85 and 86
D more than 88
e
Q11 How many valence electrons (electrons in the
outermost shell) are present in the element of
atomic number 147
A4
c z
83
D1
Q12 Which one of the following atoms will have the
same number of neutrons as an atom of ~~Sr?
A §~y
e ~§Sr
8 ~fRb
D ~Kr
Q13 Which statement is correct for the emission
spectrum of the hydrogen atom?
A The lines converge at lower energies.
B The lines are produced when electrons move
from lower to higher energy levels.
e The lines in the visible region involve electron
transitions into the energy level closest to the
nucleus.
o The line corresponding to the qreatesternission
of energy is in the ultraviolet region.
Standard Level Paper 1, Nov 03, 06
Q14 Naturally occurring chlorine consists of the
isotopes chlorine-35 and chlorine-37. The relative
atomic mass of chlorine is 35.5. Which one of the
following statements is true?
A The chlorine-35 and chlorine-37 atoms are
present in equal amounts.
8 The ratio of chlorine-37 atoms to chlorine-35
atoms is 2 : 1.
e The ratio of chlorine-37 to chlorine-35 atoms is
37/35.
D There are three times as many as chlorine-35
atoms as chlorine-37 atoms.
70
A TOMle STRUaURE
Q15 Which statement is correct about a line emission
spectrum?
A Electrons neither absorb nor release energy as
they move from low to high energy levels.
B Electrons absorb energy as they move from
high to low energy levels.
C Electrons release energy as they move from low
to high energy levels.
DElectrons release energy as they move from
high to low energy levels.
Standard Level Paper 1, Nov OS, 06
Q16 Which electronic transition within a hydrogen
atom requires the greatest energy?
A n= 1 ~n=2
B n=3~n=5
Q18 The atomic numbers and mass numbers for four
different nuclei are given in the table below.
Which two are isotopes?
Atomic number
II
III
IV
A I and II
B II and III
Mass number
101
102
102
103
258
258
260
259
.C III and IV
0 I and IV
Standard Level Paper 1, Nov 98, 06
Q19 All isotopes of uranium have the same:
I
number of protons
1/ number of neutrons
III mass number
A I only
C III only
B II only
D I and III only
Q20 Which is the correct sequence for some of the
various stages that typically occur in the analysis of
an element during mass spectrometry?
A vaporization, electron bombardment,
acceleration, deflection, detection
B electron bombardment, vaporization,
acceleration,deflection, detection
C vaporization, electron bombardment,
deflection, acceleration, detection
o deflection, acceleration, electron
bombardment, vaporization, detection
and IB style questions
Q1 The element bromine exists as the isotopes 79Br
. and 81Br,and has a relative atomic mass of 79.90.
a Copy and complete the following table t9
show the numbers of sub-atomic particles in
the species shown.
[3]
An atom of 79Br
An ion of 81Br
Protons
Neutrons
Electrons
b
C n=2~n=3
D n=5~n=oo
Q17 Which of the following radioisotopes is used in
nuclear medicine to image the thyroid gland?
A iodine-131
C fluorine-18
B carbon-14
D uranium-235
I
Paper 2'18 questions
State and explain which of the two isotopes
79Brand 81Bris more common in the element
bromine.
[1]
Standard Level Paper 2, Nov OS, 03
Q2 The element silver has two isotopes, l~~Agand
lS1Ag,and a relative atomic mass-of 107.87.
a Define the term isotope.
[1]
b State the number of protons, electrons arid
neutrons in 12tAg+.
[2]
c State the name and the mass number of the
isotope relative to which all atomic masses are
measured.
[1]
Q3 A sample of iridium is analysed in a mass
spectrometer. The first and last processes in mass
spectrometry are vaporization and detection.
a i State the names of the second and third
processes in the order in which they occur
in a mass spectrometer.
[2J
ii Outline what occurs during the second
process.
[2]
iii State and explain which one of the
following ions will undergo the greatest
deflection (under the same conditions in a
mass spectrometer):
1911r+
or
1931r+
[2]
b The sample of iridium is found to have the
following composition of stable isotopes:
Isotope
Relative abundance/%
i
Ir-191 Ir-193
37.1
62.9
Define the term relative atomic mass.
[2]
ii Calculate the relative atomic mass of this
c
sample of iridium, giving your answer to
two decimal places.
[2]
Iridium-192 isa short-lived radioisotope used
to treat cancer.Define the term radioisotope
and name another radioisotope used in nuclear
medicine.
[2]
Q4 Describe the emission or line spectrum of gaseous
hydrogen atoms and explain how this is related to
the energy levels in the atom.'
[3]
•• II II III Ii II II •..•••••.•.••••••••••
~
•••••.••••••••••••
'-~.....•...•
~.,......:-::'"~-,...,....:_--..:::-::~,.-~ ..~~-.-,.~.
Atomic structure
1023
_.::_
.._-::..~:'~
- .•..
-··~.....,.-'.:':"':':::"-·
__
-~~_:::-;--~.·7.:::~·--_;:_-_:_-_:::::_:_:_P7_---_:::~~-··-7~-~,,~~_,
------------
2 Atomic structure
Paper 1 IB questions and IB style questions
Q1 0
Isotopes are two or more atoms of the same element with
different numbers of neutrons (and therefore different
relative isotopic masses). They have the same numbers of
protons and hence the same number of electrons.
02 C
HzS: [2H+ XZ-j
Q3 A
1P
ii
Hi
Q4 D
17p
19 P
1e
18 e
18e
On
18 n
20 n
Emission spectra provide evidence that electrons exist in
fixed energy levels within an atom.
Q5 C The atomic number of sulfur is 16; it is in group 6 and
therefore has six valence electrons.
06 0 A discontinuous spectrum is produced.
Q7 C
Q8 B
Q9 C
9 protons (from periodic table); 10 electrons in atom (due
to the single negative charge) and 10 neutrons (19 - 9).
1~.l~.!.~.!.
2
4
8
To account for the emission or line spectrum of hydrogen
atoms.
010 C
Relative atomic mass is a weighted average. The more
abundant isotope has a mass number of 85.
011 A
The element is silicon with the electron arrangement 2,8,4.
012 B
~Y: 52 neutrons ~~Rb:52 neutrons ~Sr: 51 neutrons
~Kr: 48 neutrons.
013 0 They converge at high energies; lines are produced when
electrons move from higher to lower energies.
Q14 0
.
.
f chlori
ReIauve
atomic mass 0
onne
= (35x3)+(37x1)
3+ 1
=142=355
4
015 0
.
Electrons release energy as they move from high to low
energy levels: light is released.
016 A -Energy levels rapidly converge with an increase in n.
017 A
The thyroxine secreted by the thyroid is an iodinecontaining compound.
018 B
I. 157 neutrons; II. 156 neutrons; III. 158 neutrons and IV.
156 neutrons.
019 A
Atoms of isotopes have the same number of protons have
the same number of protons but different numbers of
neutrons. The mass number is the number of protons and
neutrons.
020 A
The sample is first vaporized by lowering the surrounding
pressure: a mass spectrometer can only accept gaseous
atoms or molecules. The molecules or atoms are then
subjected to bombardment by high energy electrons
which remove a valence electron to form unipositive ions.
The ions are then accelerated by an electric field before
entering a powerful magnetic field where they experience
deflection according to their mass-to-charge ratio.
1024 ANSWERS TO EXAMINA TlON QUESTIONS
--------
Paper 2 IS questions and IS style questions
An atom of 79Br An ion of8'Br
Q1 a
b
Q2 a
b
c
Q3 a
Protons
35
35
Neutrons
44
46
Electrons
35
36
[31
nBr because its relative isotopic is closer to 79.90
[1]
Atoms of the. same.element but with different mass
numbers and hence different number of neutrons.
47 protons, 46 electrons and 60 neutrons
Carbon-12 (12C).
!1J
III
[11
i
Ionization, acceleration
12J
The gaseous atoms formed from the sample are
bombarded with electrons traveliing at high speed with
large kinetic energies.
!2J
iii '"'Ir+
[1]
It has the lowest mass-to-charge ratio (or lowest
mass since the Charges are identical).
[1J
The weighted average of relative isotopic masses of all
the stable isotopes of an element relative to the mass
of one atom of carbon-12.
12J
ii (191 x0.371)+(l93
x 0.629) = {70.861 + 121397)
= 192.258 = 192.26 (2 d.p.)
[2J
iii An unstable isotope of an element that decays or
disintegrates spontaneously, emitting nudear or
ionizing radiation,
[1]
Cobalt-60, molybdenum-99, chromium-51, iron-59
and iodil1e-125
[1]
ii
b
Q4 The emission spectrum consist of a series of sharp or discrete
coloured lines on a black background.
[1 J
The lines converge together at high energy (high frequency).
The lines are generated when excited electrons move from
high energy levels to lower energy levels.
[1 J
light of a particular frequency is released during this process. [1J
---------------- ---~----
340 A TOMle STRUCTURE
TAL·Bol
CH--Pf
SUMMARY OF
KNOWLEDGE
11..
"._ The first ionization energy is the minimum energy required to remove a mole-of
electrons from a mole of gaseous atoms to form a moie of gaseous unipositive ions (under
standard thermodynamic conditions).
The ionization energy of an atom or ion is determined by the following factors: nuclear
charge, shielding effect and atomic radius (distance between nucleus and outer electrons).
On moving across periods 1, 2 and 3 there is a general increase in first ionization energy.
This is due to a large increase in nuclear charge which is accompanied by a small increase in
shielding.
The nuclear charge is the force that attracts all the electrons to the nucleus. The shielding
effect is the effect of shielding the outer electrons from the attraction of the nucleus by the
repelling effect of the inner electrons .
.-' There-are-two small decreases in first ionization energy in periods 2 and 3. The first is due
to a change in sub-shell (sub-level) and loss of shielding and the second is due to enhanced
-electron-electron repulsion. These factorsoutweighi:he
increase in nuclear charge with
increasing atomic number.
Electrons can be progressively removed from gaseous atoms and successive ionization
energies measured and plotted against number of electrons removed. The graphs offer
strong experimental evidence for shells (main energy levels) and sub-shells (sub-levels). The
relatively large increases in successive ionization energies correspond to changes in shells;
smaller changes correspond to changes in sub-shells or electron pairing.
-. Each shell can hold a maximum number of electrons: first shell (two electrons), second shell
(eight electrons), third shell (18 electrons) and fourth shell (32 electrons). (The number of
electrons in each shell is given by 2n2, where n represents the shell number.)
=.::. Each shell (energy level) is composed of one or more sub-shells (sub-levels), The first shell is
composed of one sub-shell, the second shell is composed of two sub-shells, the third shell is
composed of three sub-shells and the fourth shell is composed of four sub-shells.
"_ Electrons occupy regions or volumes of space called orbitals. Each atomic orbital can
accommodate a maximum of two electrons.
Electrons can occupy four types of orbitals: s, p, d and f. s orbitals are spherical in shape; p
orbitals are 'dumb-bell' shaped and arranged mutually perpendicularly to each other. p, d
and f orbitals all have the same energy within the same shell or sub-shell.
Orbitals (of the same type) retain the same shape but become larger (and their electron
density more diffuse) with an increase in shell number.
Summary of know/edge
341
,~.
C The first shell has an s sub-shell (one orbital); the second shell has an s and a psub-shell
::J
r:
L..
':J
c:
;]
(three orbitals); the third shell has an s sub-shell, a p sub-shell and a d sub-shell (five orbitals);
and the fourth shell has an s sub-shell, a p sub-shell, a d sub-shell and an f sub-shell (seven
orbitals).
Electrons are assigned to atomic orbitals of atoms (in their ground state) according to the
Aufbau principle, the Pauli exclusion principle and Hund's rule.
The Aufbau principle states that electrons occupy atomic orbitals in the order of the energy
levels of the orbitals. .
Copper and chromium atoms have anomalous electron configurations which violate the
Aufbau principle. The chromium atom is 3d54s1 and the copper atom is 3d104si because d
sub-shells that are half-filled or fully filled are particularly stable.
The Pauli exclusion principle states that only two electrons may occupy the same orbital and
that these two electrons must have opposite spins.
Hund's rule states when electrons are placed in a set of atomic orbitals with equal energies,
the electrons must occupy them singly with parallel spins before they occupy the orbitals
in pairs. As a consequence of Hund's rule atoms tend to maximize the number of unpaired
electrons.
The electron configuration of an atom describes how electrons are distributed among the
various orbitals in the various shells and sub-shells and is denoted by an orbital diagram or
spdf notation.
In orbital notation each orbital is represented by a box and each electron by an arrow. The
arrow head indicates the direction of electron spin.
In spdf notation the electron configuration is denoted by writing the symbol for the
occupied sub-shell and adding a superscript to indicate the number of electrons in that subshell. This can be condensed by describing the core electrons with the nearest noble gas
electron configuration.
Simple positive ions are formed by removing the appropriate number of outer electrons;
simple negative ions are formed by adding the appropriate number of electrons to the outer
sub-shell. Excited atoms have one or more electrons promoted to higher energy levels.
Many ions formed by non-transition metals obey the octet rule. These ions have a full outer
shell of eight electrons.
Excited atoms are formed when one or more electrons are promoted to higher energy
orbitals.
Examination questions - a selection
Paper 1 IB questions and IB style questions
Q4 An atom of chlorine has the electron configuration
[Ne]3s23pS What is the number of orbitals
occupied by at least one electron?
A 7
B 9
C 13
D 17
Ql What is the electron configuration for an atom
with Z = 22?
A 1s22s22p63s23p63d4
B 1s22s22p63s23p64s24p2
c 1s22s22p63s23p63d24p2
D 1s22s22p63s23p64s23d2
Higher Level Paper l;"May 03, Q5
Q5 For which of the following
Q2 For the species below, which one would require
Q6 Which of the following
the most energy for the removal
A Na+
B F
C F-
of an electron?
D Ar
Q3 How many unpaired electrons would the FeZ.•.ion
be expected to have?
A1
B3
(4
D6
pairs of species are the
chemical properties most similar?
A tH and tW
C ilNa and l§K
B l~C and I~N
D ~Li'"and ~Be
first ionization energy?
A 19K
B 6C
(12Mg
D 4He
Q7 What is the total number of electrons in d orbitals
in a tin atom 50Sn?
A 5
B 10
(20
D 0
Q8 Which of the following
ionization energy?
A Na
B F
-.:....-,...,.,-...,,---_.- -------.
elements has the smallest
atoms has the smallest first
C Be
D CI
342 A TOMle STRuaURE
Q9 What is the electron configuration
[Ar]
(
[Ar]3d5
B
[ArJ4s23d4
D
[Ar]3d5
B
(
[ArJ4s23d3
D
energies for two
X and Y, are given below.
elements,
Q10 What is the ground state electron configuration
the Fe3+ion?
A [ArJ4s13d5
Q20 The first three ionization
of C03+?
A
Element First ionization
energy/kJ mol-'
of
--------
x
[Ar]3d5
[Ar]3d3
y
Q11 Which set of chemical species is arranged
in order
of decreasing ionic radius?
A Ca2+ > 5r2+ > Ba2+
(K+
52-
B A13+> Mg2+ > Na+
Na+
> (1- >
D 02- > F- >
Which
A
following
A 12Mg
B 2sMn
(3SBr
D 36Kr
for the
atoms is incorrect?
1s22s22p63s2
7300
11800
1086
2350
4620
pair of elements
Na+
B
Na
3Li and 6C
(
B 4Be and lSP
(Mg2+
D A13+
Q15 How many unpaired electrons would the cobalt
27) in its ground state
be expected to have?
(4
D3
Q16 Which property
changes along with increasing
ionization energy?
A increasing non-metallic properties
increasing atomic radii
decreasing electron affinity
decreasing nuclear charge
A
B
4Be
(6C
orbitals?
D 7N
Q18 What is the maximum
number of electrons that
can occupy the 5d sub energy level?
A 20
B 10
(32
D 25
Q19 Which of the following
first ionization
A
lBArB
elements
chemical
C
Mg2+
species has the
D 52-
F-
(6C
(Cu2+
B Ni
D Ni2+
series differ from
each other mainly in the number
electrons?
A p electrons
B sand p electrons
( p and d and f electrons
D d electrons
of which
Q24 Which is the best periodic correlation
atomic number and another
A atomic masses
B ionic and atomic radii
( first ionization energies
D electron configuration
atomic
type
of
between
property?
D
55(S
series are the atoms
in order of increasing first ionization
A
Be, Mg, Ca
C
Be, B, C
B
0, F, Ne
D
Ne, 0, F
Q26 What increases in equal steps of one from left to
right in the periodic table for the elements
to neon?
A
B
has the largest
energy?
lSP
BO and 165
D 2He and 4Be
Q23 Elements in the first transition
arranged
energy?
state have two half-filled
lSP
and Y?
Q25 In which of the following
Q17 For which element would neutral isolated atoms in
the ground
X
configuration?
species would you expect
to have the largest radius?
A Na+
B N3(Ne
D F-
B
(
D
largest radius?
A A13+
B
A Mn
Q14 Which of the following
atom (element number
electronic configuration
AS
B2
could represent
,-,
Q22 Which atom or ion could have a 3dB electronic
[ArJ4s23d4
[Ar]3d104s24p5
[Ar]4s23d104p6
Q13 Which chemical species has the smallest radius?
A
----
520
Q21 Which of the following
Q12 Which ground state configuration
Second
Third ionization
ionization
energy/kJ mol-'
energy/kJ mol-1
the number
the number
- isotope
of occupied
of neutrons
electron energy levels
in the most common
( the number of electrons in the atom
D the atomic mass
Higher Level Paper 1, May
Q27 In which of the following
configurations
j
{He12s22p2
II
[He]2s22p3
I only
I and II only
as, Q6
ground-state
electron
are unpaired electrons present?
III [He]2s2p4
A
B
lithium
c
D
I, II and III
II and III only
Examination questions 343
Q28 What is the number of unpaired electrons in the
Cr3+ion?
A
B 2
C 3
0 5
d
State which period element C is in and explain
your reasoning.
[3]
e Why is the first ionization energy of element G
lower than that of element F?
[2]
a
Q29 A transition metal ion has the electronic
configuration X3+= [Ar]3d4. What is the atomic
number of element X?
A 24
B 22
C 25
0 26
Q30 What is the total number of p orbitals containing
one or more electrons in germanium (atomic
.
number 32)?
A 2
B 3
C 5
0 8
Higher Level Paper 1, May 04, Q5
Paper 2 IB questions and IB style questions
Q1 The graph below
ionization energy
Figure 1 refers to
Figure 2 refers to
group as element
shows the variation in first
of some chemical elements.
the chemical elements and
chemical elements in the same
C.
F
Q2 The successive ionization energies of germanium
are shown in the following table:
---Ionization energy/
kJ mol-1
1st
2nd
3rd
4th
5th
760
1540
3300
4390
8950
a
Identify the sub-level from which the electron
is removed when the first ionization energy of
germanium is measured.
[1]
b Write an equation, including state symbols,
for the process occurring when measuring the
second ionization energy of germanium.
[1]
c Explain why the difference between the 4th
and 5th ionization energies is much greater
than the difference between any two other
successive values.
[2]
Higher Level Paper 2, Nov 05, 02
Q3 The graph shows the variation in second ionization
energy of the elements silicon to argon.
L-
0
E
J
.Y
<,
>-
e'
CD
c
2900
2700
/
/~
2500
CD
c
2300
~
2100
0
x
x+1
x+2
x+3
x+4
N
'c
.Q
Atomic number
Figure 1
-0
C
0
o
1900
1700
<D
(j)
./
1500
/
Si
/
~
P
S
CI
Ar
Element
>-
e'<D
a
Explain the general increase in second
ionization energy from silicon to argon.
[3]
b Give the detailed electron configurations
of the S-;-(g)and CI+(g) ions.
[2]
C Explain why the increase from sulfur to chlorine
is significantly less than the increase to sulfur
from phosphorus.
[1]
c
(J)
c
o
~
.~
c
.Q
x-8
x
x+8
Atomic number
Figure 2
a
-:
Define the term first ionization energy of an
element.
[2]
b Element C, with atomic number x, is in group 2
of the periodic table. Justify this using all the
information from Figure 1.
[4]
C Explain the trend in the first ionization energy
as shown in Figure 2.
[3]
Atomic structure 103s---------------------------------------------------------
12 Atomic structure
Q12 B
The electron configuration for a manganese atom is 4s23d5•
Q13 0
Na 2,8,1 (11 protons), Na+2,8 (11 protons), Mg'+ 2,8 (12
protons) and AP+ 2, 8 (13 protons). The radii decrease due
to the increasing effectivenudear charge.
Q14
B
Na+2,8 (11 protons), N3--2,8(7 protons), F- 2,8 (9 protons),
Ne 2,8 (10 protons). The electrons in the nitride ion are
experiencing the lowest effective nuclear charge.
Q15
D
A cobalt atom has the following electron configuration:
1 s'2s'2p63s'3p63d74s'. The 3d sub-shell has two spin pairs
and three unpaired electrons.
Q16 A
Ionization energies generally increase across periods.
This results in a decrease in metallic behaviour, that is, a
decrease in tendency to form cations (positive ions).
Q17 C
A carbon atom has the ronowing electron configuration
1522£221". The p sub-shell has two unpaired electrons in
accordance with Hund's rule.
Q18
All d sub-energy levels have 10 orbitals and hence can hold
to a maximum of 10 electrons.
B
Q19 A
Ionization energies decrease down a group and generally
increase across a period. Ihe electrons in argon are
experiencing the highest effective nuclear charge.
Q20 A
X is in group 1 and Y is in group 2 or 4.
Q21 0
Al3+2,8 (13 protons), Mg'+ 2,8 (12 protons), F- 2,8 (9
protons), 52- 2,8,8 (16 protons). The electrons in the sulfide
ion are experiencing the lowest effective nudear charge.
Paper 1 IB questions and IB style questions
Q1 D
Q2
Titanium has an atomic number of 22. Its electron
configuration is 1s'2s'2p63s'3p64s'3cf2.
A The removal of an electron from a sodium ion would
involve the removal of a core electron from the second
shell close to the nucleus. This electron wouid be strongly
held and hence would have a large ionization energy. This
electron would have the highest effective nuclear charge.
Q3 C
The Fel+ion has the configuration [Ar]3d6. Hence there is
one spin pair and four unpaired electrons.
04 B 1s sub-shell: 1 occupied orbital; 2s sub-shell: 1 occupied
,.
orbital; 2p sub-shell: 3 occupied orbitals; 3s sub-shell: 1
occupied orbital; 3p sub-shell: 3 occupied orbitals: two
filled and one half-filled.
05
C
The two elements are both in the same group (group 1).
Elements in groups, especially those at the extremes of the
periodic table, have very similar chemical properties.
Q6 A
Ionization energies decrease down a group and generalty
increase across a period.
Q7 C
A tin atom has the following electron configuration:
1s22s'2p63s'3p63dl04s24p64dl°5s'5p>.
Q8 A
Ionization energies decrease down a group and generally
increase across a period. The valence electron in sodium is
experiencing the lowest effective nuclear charge.
Q9 D
Tne electron configuration of the cobalt(IIf}'ion is:
W2s'2p63s23p63d6. The 4s electrons are lost from a
transition metal atom before the 3d electrons.
Ql0
C
The-electron configuration of the iron(m)ion is
1s22s22pG3s'3p63d5.
The 4s electrons are lost from a
transition metal atom before the 3d electrons.
Q11
D
The electron arrangements of the species are all 2,8, that
is. they are ISO electronic. However, the atomic (proton)
numbers are 8, 9 and 11. The effective nuclear charge
increases left to right and hence radii decrease.
Q22
0
Ni 3dE4s2;Mn 3d54s2; CUlT 3d9; Ni2+3dS
Q23
0
Elements in the first transition series (with the exception of
copper and chromium) have a 3&4s2 configuration.
Q24
D
The atomic number is the number of protons in the
nucleus. The proton number determines the number of
electrons and hence the electron configuration.
Q25
B
Oxygen 2,6 (8 protons), fluorine 2.7 (9 protons), neon 2,8
(10 protons). Effective nuclear charge is increasing from left
to right
Q26 C
From lithium to neon the atomic number (proton number)
increases by one. This is accompanied by the addition of an
extra electron.
Q27
C
I has two unpaired 2p electrons; II has three unpaired 2p
electrons and III has two unpaired 2p electrons.
Q28
C
The chromium(m) ion has the electron configuration 3d3.
The chromium atom has the electron configuration 3d54s1.
Q29 C
X will have the electron configuration is [Ar]3d54s'.
Transition metal ions ionise via loss of 45 and then 3d
electrons. The total number of electrons in the atom is 18
+ 5 + 2 = 25.
Q30
A germanium atom has the following configuration:
Js22s22p€3s23p54sz
3d104p2.Three 2p orbitals, three 3p
orbitals and two unfilled 4p orbitals.
0
Paper 2 IB questions and IB style questions
Q1 a
The first ionization energy is the minimum energy required
to remove one mole of electrons from one mole of gaseous
atoms to form one mole of unipositive gaseous ions (under
standard thermodynamic conditions).
[2]
1036 ANSWERS TO EXAMINA nON QUESTIONS
7_-=-:",:"""'-'~_-+7:;-:~
~~~-_-.
-:------:-_-~----~~_~.-_--_-.
~--.~---__"-=----_--. ---~~~--.,-~~"""---- .------------
-----_--:--:-:--: -~--
-~--~.-
---- ---'-- ~----..----~----------
b
c
In a group 2 element (C) the electron (for the first
ionization energy) is removed from a spin pair in a s subshell. For D, a group 3 element, the electron (for the first
ionization energy) is removed from a 3p sub-shell further
away from the nucleus. The 3p sub-shell also experiences
more shielding. Hence there is a decrease in first ionization
energy when moving from C to D. The increase in nuclear
charge accounts for the increase from E (np2 to F np3). In
addition C cannot be in group 5 (the other group after
which there is a slight decrease) because there is a steady
rise for the next three elements (0 to E to F) indicating the
filling of a p sub-shelll.
[4}
As you move down group 2 the first ionization energies
decrease. As you descend the group the nudear charge
increases due to the presence of additional protons, but
the shielding effect progressively increases due to the
presence of an extra electron shell as you move from
one period to the next. The effect of the extra protons is
compensated for by the effect of the extra electrons. The
only factor left is the additional distance between the outer
electron and the nucleus: That lowers the first ionization
d
e
Q2 i
ii
Hi
Q3 a
energy.
13J
(is in period 3 since in period 2 the group 2 element
would have the higher ionization energy while in period 4
the group 2 element would have a lower ionization energy.
This can be accounted for by the progressively increasing
average distance the valence electron is from the nucleus.
C cannat be in the first period (He) as p-orbrtals are being
filled, it cannot be in the fourth period (Ca) as (x + 8)
would- not bring you back to group 2 again.
[3]
There is a slight decrease from F to G due to the presence
of a spin pair in the np sub-shell for G. The resulting
electron-electron repulsion is greater than the effect of
the increase in nuclear charge and reduces the ionization
energy of G+.
[2}
4p
(1]
Ge+(g) ~ Ge'+(gl+ e[1]
5th electron removed from 3rd energy level and 4th
electron from 4th energy level. The attraction by the
protons in the nucleus is greater since the electrons are
closer to the nucleus.
[2}
Across a period the atomic number increases progressively
by one unit, thus resulting in an additional proton being
added to the nucleus and thus an increase in the attraction
for all the electrons, including the outermost electron.
11]
Since the number of electron shells for all elements in
period 3 is constant, namely, three, there is very little
increase in shielding.
[1 J
Hence, the increase in nuclear charge across the period will
bring all the electrons progressively closer to the nucleus,
causing the removal of the second electron from the atom
to require progressively more energy.
{1 J
b
The electron configurations of the S+and (1+ ions are
1SZ2S22p63s23p3
and 1s22s22p63s23p'.
[2J
c
The (1+ ion has a spin pair of electrons in the 3p sub-shell.
This electron-electron repulsion counteracts, to some
extent, the increase in nuclear charge. Hence, the increase
in ionization energy is relatively small.
{1J