Download topic 3: periodicity

TOPIC 3: PERIODICITY
Periodicity refers to repeated patterns of physical and chemical properties. The study of periodicity is
made easier through the use of graphs of atomic number against the different properties we will study.
Such graphs show repeated patterns of minima and peaks which is referred to as periodicity.
3. 1. THE PERIODIC TABLE
3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number.
3.1.2 Distinguish between the terms group and period.
3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z = 20.
3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its
positionperiodic
in the periodic
 Modern
table:table.
elements are in order of increasing atomic number.
 group = vertical column of elements; within a group there are distinct chemical similarities but also
gradual changes in properties.
 period = horizontal row of elements from an alkali metal to noble gas.
 Relationship between electronic configuration of elements and their position in the Periodic Table:
 group number indicates the number valence electrons in the last main energy level or the
highest occupied energy level (these valence electrons decide the chemical and physical
properties)
 period number indicates the main energy level that is being filled or highest occupied energy
level.
3. 2. PHYSICAL PROPERTIES
3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and
melting points for the alkali metals ( Li -Cs ) and the halogens ( F - I ).
3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for
elements across period 3.
a) atomic radius
How do we measure the atomic radius?
 in non-metallic elements: the atomic radius is half the distance between the centres of two covalently
bonded atoms (= covalent radius). In table 8 of the data booklet, there is no atomic radius for inert
gases as they do not make covalent bonds in diatomic molecules.
 in metallic elements: the atomic radius is half the internuclear distance in the metallic crystal.
These radii can only be measured as distances between two nucleii as only the positions of the nucleii
can be established accurately enough.
The atomic radius is limited to the point at which the nuclear attraction is cancelled out by the electron
repulsion. In a simplified way, the atomic radius is similar to the distance between the nucleus and the
most outer electron.
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Trends in the Period Table
across period 3 from left to right
Atomic size/radius decreases as you move across because:


the nuclear charge increases;
electrons are added to same main energy level /similar shielding effect as the number of complete
inner energy levels is the same
As a result of the above the effective nuclear charge increases (effective nuclear charge = the proton
number – the number of electrons on filled main energy levels e.g. in the case of sodium this is +11 – 10
= +1) which attracts the outer electrons more strongly.
The increased repulsion between the electrons which would cause the radius to increase is cancelled
out by the increased nuclear charge.
alkali metals
Increases as you go down because:


the number of main energy levels increases (more electrons) which cancels out the increase in
nuclear charge; outermost electron is placed in a higher energy level which is further away from the
nucleus.
the shielding effect increases as you go down as there are more complete energy levels.
shielding effect
The screening or shielding effect from the nuclear charge by the inner electrons; as a result the valence
electrons do not feel the full effect of that nuclear charge e.g. the 1s and 2s electrons shield the 2p
electrons from the nuclear charge. The greater the shielding effect, the less an electron is attracted by
the nucleus and the lower the ionisation energy.
The shielding effect depends on:
 the number of inner energy levels; shielding effect is always greater when an inner energy level is full;
 the closer the inner electrons are to the nucleus, the greater their shielding;
 a full inner main energy level has a greater shielding effect than an incompletely filled energy level;
electrons on the same energy level do NOT shield each other.
halogens
Increases as you go down because:


the number of main energy levels increases (more electrons) which cancels out the increase in
nuclear charge; outermost electron is placed in a higher energy level which is further away from the
nucleus.
the shielding effect increases as you go down.
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b) ionization energy
3.2.1 Define the terms first ionization energy and electronegativity.
Ionization energy refers to the minimum amount of energy required to remove an electron from one
mole of gaseous atoms (or ions); it is measured in kilojoules per 1 mole and is defined by the following
equation:
atom in ground state (g) +
ionisation energy
 atom + (g) + e -
Ionization energy has a positive value as it is an endothermic process – energy is needed.
The ionization of an atom depends on:
 the distance of the electron from the nucleus (for the outer most electron this will be the atomic
radius) that will be removed from the nucleus/the energy level it is on/size of attraction;
 the size of the positive nuclear charge (=number of protons);
 the screening or shielding effect from the nuclear charge by the inner electrons.
The greater the ionisation energies in an atom/ion, the more stable the electron configuration of that
atom/ion.
Trends in the Periodic Table
We could consider the first ionisation energies only (is the energy needed to remove the most outer
electron) or we could consider all ionisation energies in each atom, the same trends will be observed.
 first ionisation energy:
 Minimum energy needed to remove first electron from 1 mole of gaseous or free atom to
make a gaseous ion;
 X+ (g)
X (g)
+ e-
 the more strongly the electron is attracted to the nucleus, the greater the amount of energy
needed (measured in Joules per mole);
 successive ionisation energy:
 energy needed to remove second, third,... electrons from 1 mole of gaseous ions to form
gaseous ions; eg
second ionization energy
X+ (g)
 X2+ (g)
+ e-
 the lower the energy level an electron is on, the more energy is needed;
successive ionization energies increase as the electron-electron repulsion decreases (this
repulsion puts electrons onto higher energy levels) + electrons are removed from positive
ions.
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across period 3
 ionisation energy increases when moving to the right because more energy is needed because:
 the increased nuclear charge;
 smaller atomic radius (outermost electron closer to nucleus);
 electrons go in the same energy level (similar shielding effect).
The result is a stronger attraction which pulls the valence electrons closer to the nucleus/stronger
attraction.
(HL only) The dips in first ionisation energy can be explained in the following way:
 dip from Be to B: boron’s 2p electron feels extra shielding of 2s electrons and has therefore a lower
ionisation energy;
 dip from N to O: oxygen fourth 2p electron goes into the first 2p orbital on which there is already an
electron with opposite spin; the resulting repulsion between these two electrons - which is minimised
because of their opposite spin - makes it easier to remove that electron.
alkali metal group (Li-Cs)
Ionisation energy decreases as you go down the group because:



atomic radius increases/outer electron to be removed is further away from nucleus, reducing the
attraction between the valence electrons and the nucleus (this offsets the increased nuclear charge).
Increased shielding as there are more energy levels so there is less effective nuclear charge
More inner electrons also means greater repulsion.
halogens
Ionisation energy decreases as you go down the group because:



atomic radius increases/outer electron to be removed is further away from nucleus, reducing the
attraction between the valence electrons and the nucleus (this offsets the increased nuclear charge).
Increased shielding as there are more energy levels so there is less effective nuclear charge
More inner electrons also means greater repulsion.
c) electronegativity
3.2.1 Define the terms first ionization energy and electronegativity.
3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table.
Electronegativity is the ability for an atom to attract a bonding pair of electrons (a shared pair in a
diatomic covalent bond). The more commonly used scale is the Pauling scale in which all values are
measured relative to fluorine which has the maximum electronegativity of 4.0.
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Trends in the Periodic Table:
across period 3
Increases when moving to the right because:



the positive nuclear charge increases;
the atomic radius decreases so the nucleus can attract other electrons better;
the difference in shielding effect is minimal as all have the same number of filled inner shells.
alkali metals
Decreases as you go down:



the number of main energy levels increases, increasing the shielding effect on free electrons
(=electrons part of another atom),
increased the atomic radius;
the two above cancel out the increased nuclear charge.
halogens
Decreases as you go down:



the number of main energy levels increases, increasing the shielding effect on free electrons
(=electrons part of another atom),
increased the atomic radius;
the two above cancel out the increased nuclear charge.
d) ionic radius: cations
Cations are always smaller than the atom from which they are made. This is because when an electron
is removed the other electrons experience less repulsion between them and they contract towards the
nucleus; the contraction is the greatest when the removed electron (s) is/are from a higher energy level.
Also in some ions, the remaining electrons are also at lower energy levels. When electrons are
removed, there is also a greater effective nuclear charge (i.e. same number of protons for a lower
number of electrons) on the remaining electrons.
Trends in the Period Table
across period 3
Ionic radius decreases as the cations have the same number of electrons (= isoelectronic) but the
positive charge increases pulling more on the electrons.
alkali metals
Increases as you go down as the number of main energy levels increases which offsets the increase in
nuclear charge.
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halogens Do not form cations.
d) ionic radius: anions
Anions are always larger than the atoms they are derived from as the added electrons increase the
repulsion between the valence electrons pushing them further away from the nucleus. The electrons
are held less tightly.
Trends in the Period Table
across period 3
Ionic radius decreases as the anions have the same number of electrons (= isoelectronic) but the
positive charge increases pulling more on the electrons.
alkali metals
Do not form anions.
halogens
Increases as you go down as the number of main energy levels increases which offsets the increase in
nuclear charge.
For all isoelectronic ions: the greater the atomic number, the smaller the
radius.
e. melting point
Melting point depends on the type of structure the substance has (more on this in the next topic).
Trends in the Period Table
across period 3
Melting point increases in the order Na  Mg  Al  Si and then decreases sharply as a result of the
different types of bonding as you move across the period starting from metallic (Na, Mg and Al) to giant
covalent (Si) to intermolecular (P, S, Cl, Ar).
Metallic bonds become stronger as the number of delocalised/valence electrons per atom increases and
the size of the ions decreases which explains Na  Mg  Al (in Na there is only 1 delocalized electron
per 1 Na atom; also the greater the number of delocalized electrons, the greater the charge on the
positive metal ion).
Si is covalent solid which means its lattice is held together by strong covalent bonding, which is stronger
than metallic bonding which is why Si has the highest.
From phosphorus to argon the mp decreases as they are simple molecular substances with weak Van
der Waals’ forces. These forces get weaker as the mass of the molecules overall decreases, P4, S8, Cl2
and Ar.
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alkali metals
Decreases as you go down as the metallic bond becomes weaker; as the radius increases, the
delocalised/valence electrons (the same number for each atom in group 1) are further away from the
nucleus so there is a weaker attraction between the metal ion and the free electrons. These delocalized
electrons are also more shielded by more electron shells.
halogens
Increases as you go down as the strength of the non-polar intermolecular forces (Van Der Waal’s)
increases due to greater instantaneous polarity within the molecules.
Summary
physical property
across period 3
down group 1
down group 7
atomic radius
decreases
increases
increases
ionization energy
increases
decreases
decreases
electronegativity
Increases
decreases
decreases
cation radius
decreases
increases
do not form cations
anion radius
decreases
do not form anions
increases
melting point
increases to Si and then
decreases
decreases
increases
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3. 3. CHEMICAL PROPERTIES
3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group.
Alkali metals (Li to Cs)
Similarities:
 they are all in the same group because they all have 1 valence electron.
 all alkali metals are very reactive (low ionisation energy); because of their low ionisation energies
(their only valence electron is well shielded and is further away from the nucleus ) they are easily
oxidised, they are very good reducing agents:
M (g)  M+ (g) + e they have low first ionization energies but high second ionization energies because of similar electron
configuration.
 they reduce water to hydrogen and hydroxide forming an alkali solution which is why they are called
alkali metals: (observations: the metals move around and fizz). Also a gas is produced.
2Na (s) + 2H2O (l)  2NaOH (aq) + H2 (g)
 they readily react with air producing distinctly coloured flames as they absorb and emit visible light;
 they have low melting and boiling points (compared to other metals), low density (they all float on
water) and are soft because of weaker - relative - metallic bonds;
 they react easily with halogens to form white ionic compounds called halides; they form compounds
with formula XY with halogens. These reactions are highly exothermic
2Na (s) + Cl2 (g  2NaCl (s)
Differences: eg observations: Li just floats on water, Na might catch fire, K always catches fire!!!
More metallic as you go down/more reactive/more ionic:



more reactive; as you go down the first ionisation energy decreases as outer electrons are further
from the nucleus. The outer electrons are also more shielded. Fr is the most reactive with the
lowest first ionisation energy, Fr is also the best reducing agent.
lower m.p/b.p;
higher density: Li, Na and K float on water whilst the Rb, Cs and Fr sink.
Halogens (Cl, Br, I)
Similarities
 7 electrons in outer shell;
 all reactive/good oxidising agents (*remove electrons)/become easily reduced;
Hal + e-  HalPage 8 of 16
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 form halide ions: Hal- /oxidation state is mostly -1 (only fluorine has -1 as its only oxidation state; the
other have variable oxidation states)
 diatomic molecules/single covalent bonds;
 more soluble in organic solvents (eg hexane) than in water;
 react with each other; more reactive halogen displaces less reactive one (you need to be able to
explain the experimental procedure of these displacement reactions and all the observations eg
colour changes that occur which are summarised below) (See Practical on Halogens).
Solution (aq)
Chlorine water
(Cl2 (aq))
Potassium
chloride
(Cl- (aq) =
ions)
Potassium
bromide
(Br- (aq) =
ions)
Potassium
iodide
(I- (aq) =
ions)
Initial colour
Colour after
shaking
colourless
Products
Equation
Colour after
shaking
Products
Bromine water
(Br2(aq))
orange
Iodine water
No change
No change
No reaction
No reaction
From colourless to
yellow/red
Bromine + potassium
chloride
(I2(aq))
brown
No change
No reaction
Equation
Cl2 +2Br- Br2 + 2Cl-
Colour after
shaking
Products
From colourless to
red/brown
Iodine + potassium
chloride
From orange to
brown colour
Iodine + potassium
bromide
Equation
Cl2 +2I- I2 + 2Cl-
Br2 +2I-  I2 + 2Br-
The element not included in this table is fluorine because it is far too reactive. Flourine displaces all
other halide ions according to the equation:
F2 +
2Cl-/Br-/I-/At- 
Cl/Br/I/At2 +
2F-
 react with water: fluorine and chlorine to form acidic solutions; eg
Cl2 (g) + H2O (l)  HCl (aq) + HClO (aq)
2F2 (g) + 2H2O (l)  4HF (aq) + O2 (aq)
 react with silver ions to form solid silver halides some of which decompose in sunlight as shown by
the table below.
test
Add AgNO3 (aq)
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chloride (Cl- (aq) =
ions)
bromide (Br- (aq) =
ions)
iodide (I- (aq) = ions)
white ppt
pale yellow/cream ppt
yellow ppt
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Periodicity
Effect of sunlight
ppt turns purple grey
ppt turns grey to yellow
no change to ppt
The above table shows:


darkening of silver halide as you go down;
that silver iodide is not sensitive to light but the other two silver halides are (which is why they are
used in photography).
Differences:
As you go down:
 reactivity decreases; fluorine is the most reactive halogen and best oxidising agent (why does the
oxidising ability decrease as you go down?) (fluorine has the highest electronegativity because of its
smallest radius and little shielding);
F2 oxidises Cl-, Br- and I- to Cl2, Br2 and I2 because it is a better oxidizing agent.
Cl2 oxidises Br- and I- to Br2 and I2 because it is a better oxidizing agent.
Br2 oxidises I- (but not Cl- or F-) to I2 as it is a better oxidizing agent than I2.
I2 will not react with F-,Cl-, and Br- as it is the weakest oxidizing agent of the four.
 astatine ions are oxidized the easiest of all halide ions because of lowest electronegativity of all
halogens;
 less soluble in water;
 higher melting and boiling points increase;
 colours get darker.
Period 3
3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.
As you go across Period 3 in the Periodic Table there is a change from metallic to non-metallic
character; this is shown by:

increase in ionisation energy as you go across; this means the elements are less likely to give up
electrons and form positive ions and ionic bonds which is typical metallic behaviour;

properties of the oxides also show a change in properties from metallic to non-metallic:

metals form oxides which are ionic and can act like bases (in the case of sodium oxide it
dissolves easily in water and forms an alkali) and can neutralise acids;
Na2O


+
H2O  2 NaOH
Al2O3 is an amphoteric oxide which means it can neutralise both acids and bases;
Non-metals form oxides which dissolve in water – except silicon dioxide – to form acids; they
are called acidic oxides and can neutralise bases or alkalis.
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Equation acidic oxide dissolving:
SO3
H2O  H2SO4
+

high electronegativity towards the right making the atoms better at gaining electrons.

argon does not form any compounds as it is a noble gas with full outer shell.
1
2
3
4
Na
Mg
Al
5
6
7
0
P
S
Cl
Ar
Group
Element
Structure of
element
Formula of
oxide
Structure of
oxide
Acid-base
character of
oxide
 giant metallic
Na2O

MgO
giant ionic
 basic 

Al2O3

Si
giant
covalent
SiO2
giant
covalent
amphoteric


simple molecular
P4O10
SO3
Cl2O7
P4O6
SO2
Cl2O
 simple molecular 

acidic  no oxide
IB past paper questions
PAPER 1
1. (M06) What is the electron arrangement of silicon?
A. 2.4
B. 2.8
C. 2.8.4
D. 2.8.8
2. (M06) Which reaction results in the formation of a coloured substance?
3. (N06) Why do the boiling points of the halogens increase down the group?
A. There is an increase in bond enthalpy.
B. There is an increase in bond polarity.
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C. There is an increase in the strength of temporary dipoles.
D. There is a decrease in electronegativity
4. (N05) Which properties are typical of most non-metals in period 3 (Na to Ar)?
I. They form ions by gaining one or more electrons.
II. They are poor conductors of heat and electricity.
III. They have high melting points.
A. I and II only
B. I and III only
C. II and III only
D. I, II and III
5. (N05) A potassium atom has a larger atomic radius than a sodium atom. Which statement
about potassium correctly explains this difference?
A. It has a larger nuclear charge.
B. It has a lower electronegativity.
C. It has more energy levels occupied by electrons.
D. It has a lower ionization energy.
6.
Which pair of elements reacts most readily?
A.
7.
Li + Br2
A.
K + Cl2
I only
Atomic radius
II.
Melting point
B. I and II only
C.
III. Electronegativity
I and III only
D.
I, II and III
Li
B. Be
C. B
D.
Mg
Which of the reactions below occur as written?
A.
10.
D.
For which element are the group number and the period number the same?
A.
9.
C. K + Br2
Which of the following properties of the halogens increase from F to I?
I.
8.
B. Li + Cl2
I.
Br2 + 2I– → 2Br– + I2
II.
Br2 + 2Cl– → 2Br– + Cl2
I only
B.
II only
C.
Both I and II
D. Neither I nor II
Rubidium is an element in the same group of the periodic table as lithium and sodium.
It is likely to be a metal which has a
A.
high melting point and reacts slowly with water.
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Periodicity
11.
B.
high melting point and reacts vigorously with water.
C.
low melting point and reacts vigorously with water.
D.
low melting point and reacts slowly with water.
When the following species are arranged in order of increasing radius, what is the correct
order?
A.
12.
13.
B.
K+, Ar , Cl–
C.
Cl–, K+, Ar
D. Ar, Cl–, K+
What increases in equal steps of one from left to right in the periodic table for the elements
lithium to neon?
A.
the number of occupied electron energy levels
B.
the number of neutrons in the most common isotope
C.
the number of electrons in the atom
D.
the atomic mass
Which property decreases down group 7 in the periodic table?
A.
14.
Cl–, Ar, K+
atomic radius
B. electronegativity
C. ionic radius
D. melting point
Which two elements react most vigorously with each other?
A.
chlorine and lithium
B.
chlorine and potassium
C.
iodine and lithium
D.
iodine and potassium
15. (M06) Which statement is correct for a periodic trend?
A. Ionization energy increases from Li to Cs.
B. Melting point increases from Li to Cs.
C. Ionization energy increases from F to I.
D. Melting point increases from F to I.
16. (N06) Which equation represents the first ionization energy of fluorine?
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PAPER 2
1. (M07)
(a) State the meaning of the term electronegativity.
[1]
(b) State and explain the trend in electronegativity across period 3 from Na to Cl.
[2]
(c) Explain why Cl2 rather than Br2 would react more vigorously with a solution of I- (aq) [2]
2. (M06)
(a) Explain why
(i) the first ionization energy of magnesium is lower than that of fluorine.
[2]
(ii) magnesium has a higher melting point than sodium.
[3]
(b) Discuss the acid-base nature of the period 3 oxides. Write an equation to illustrate
the reaction of one of these oxides to produce an acid, and another equation of
another of these oxides to produce a hydroxide.
[5]
3. (N06) Information about the halogens appears in the Data Booklet.
(a) (i) Explain why the ionic radius of chlorine is less than that of sulfur.
[2]
(ii) Explain what is meant by the term electronegativity and explain why the electronegativity of
chlorine is greater than that of bromine.
[3]
(b) For each of the following reactions in aqueous solution, state one observation that would be
made, and deduce the equation.
(i) The reaction between chlorine and sodium iodide.
(ii) The reaction between silver ions and chloride ions.
[2]
[2]
(c) Deduce whether or not each of the reactions in (b) is a redox reaction, giving a reason in each
case.
[4]
4. (N05)
(a) (i) Define the term ionization energy.
[2]
(ii) Write an equation, including state symbols, for the process occurring when measuring
the first ionization energy of aluminium.
[1]
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(b) The first ionization energies of the elements are shown in Table 7 of the Data Booklet.
Explain why the first ionization energy of magnesium is greater than that of sodium.
[2]
(c) Lithium reacts with water. Write an equation for the reaction and state two observations that
could be made during the reaction. [3]
5. (Specpap)
(a) Many of the properties of the elements follow trends which can be related to patterns in
the Periodic Table. Describe and explain the trends in atomic radii and ionisation
energies:
(i)
upon descending Group 1
(ii)
across Period 3
(b) Discuss the (redox) reactions of the halogens (Cl2, Br2, I2) with halide ions (Cl-, Br-, I-).
include ionic equations as appropriate. Describe and account for any colour change that
takes place.
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[12]
[8]
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