this learn about periodicity

1. Chapter 10.
Periodicity
by Joko Susilo

2. Learning objectives

3. 10.1 Structure of the Periodic Table
We now know that the chemical elements
are arranged in the Periodic Table in order
of atomic number, not atomic mass as first
thought. This explains why Mendeleev had
to re-order some elements in his table
(which was developed before scientists
knew about the structure of the atom). The
modern Periodic Table is shown in Figure
10.3 in Section 10.2. There are 18 groups
(vertical columns) in the Periodic Table.
The rows across the table are called
periods. In this chapter we will be looking
for patterns going across the third period,
from sodium (Na) to argon (Ar). The
patterns seen across Period 3 are seen
across other periods, too. This recurrence
of the same pattern is called periodicity.

6. 10.2 Periodicity of physical properties
Periodic patterns of atomic radii
Figure 10.2: The distance between the
two nuclei of the same type of atom can
be determined and then divided by two to
arrive at the atomic (single covalent)
radius.
There are other measures of atomic radii,
such as metallic radii and van der Waals’
radii. However, covalent radii can be
obtained for most elements, so these
provide the best data for comparison
purposes across a period.

7. Periodic patterns of atomic radii
The atoms of the noble gases in Group 18, such as argon in Period 3, do not
have a covalent radius, as they do not form bonds with each other. Their
atomic radii can be determined from their van der Waals’ radius. This is found
by measuring the distance between the nuclei of two neighbouring, touching
atoms which are not chemically bonded together. This distance is divided by
two to give the van der Waals’ radius. This figure will be higher than the single
covalent radius of any given element, as there is no overlap of electron clouds
involved in the van der Waals’ radius.

8. Periodic patterns of atomic radii
Figure 10.4: Plotting the atomic radii (single covalent radii) against
atomic number for the elements in Period 3 (argon not included).
The atomic radius decreases across Period 3,
as shown in Figure 10.4. The same pattern is
also found in other periods. Across a period,
the number of protons (and so the nuclear
charge) and the number of electrons
increases by one with each successive
element. The extra electron added to the
atoms of each successive element occupies
the same principal quantum shell (energy
level). This means that the shielding effect
remains roughly constant (see Section 2.2),
so the greater attractive force exerted by the
increasing positive nuclear charge on the
outer (valence) shell electrons pulls them in
closer to the nucleus. For this reason, the
atomic radius decreases across the period.

9. Periodic patterns of ionic radii
From your work in Chapter 4, you will know that the atoms of metallic elements produce positively
charged ions (called cations), such as Na+. By contrast, the atoms of non-metallic elements form
negatively charged ions (called anions), such as Cl−. What pattern in ionic radii do we see going
across Period 3? The data are shown in Table 10.2 and are displayed graphically in Figure 10.5.

10. Periodic patterns of ionic radii
Figure 10.5: Plotting the ionic radii against atomic number
for the elements in Period 3 (argon not included).
The positively charged ions have effectively lost their
outer shell of electrons (the third principal quantum shell
or energy level) from their original atoms. For this
reason, the cations are much smaller than their atoms.
To add to this effect, there is also less shielding of the
outer electrons in these cations compared with their
original atoms.
Going across the period, from Na+ to Si4+, the ions get
smaller for reasons similar to those for the decreasing
atomic radii across a period. The increasing nuclear
charge attracts the outermost (valence- shell) electrons
in the second principal quantum shell (energy level)
closer to the nucleus with increasing atomic number.
The negatively charged ions are larger than their
original atoms. This is because each atom will have
gained one or more extra electrons into their third
principal quantum shell. This increases the repulsion
between its electrons, while the nuclear charge remains
constant. This increases the size of any anion
compared with its atom.

13. Periodic patterns of melting points and electrical conductivity

14. Periodic patterns of melting points and electrical conductivity
Sodium, magnesium and aluminium, at the start of Period 3, are all metallic elements. As you saw in Section 4.6,
their metallic bonding can be described as positive ions arranged in a giant lattice held together by a ‘sea’ of
delocalised electrons. The delocalised electrons are those from the outermost (valence) shell. These delocalised
electrons are free to move around within the structure of the metal. When a potential difference is applied, the
delocalised electrons drift through the metal towards the positive terminal. Both the melting point and the electrical
conductivity increase from sodium to magnesium to aluminium. This can be explained by the number of electrons
each metal donates into the ‘sea’ of delocalised electrons and the increasing charge on the metal ions in the giant
metallic lattice. Each sodium atom donates just one electron, forming Na+ ions in the lattice, whereas each
aluminium atom donates three electrons, forming Al3+ ions. This makes the metallic bonding in aluminium
stronger, as the electrostatic forces of attraction between its 3+ ions and the larger number of negatively charged
delocalised electrons holding the giant structure together are stronger. There are also more delocalised electrons
available to drift through the structure when aluminium metal conducts an electric current, making aluminium a
better electrical conductor than sodium.

15. Periodic patterns of melting points and electrical conductivity
The element in the centre of Period 3, silicon, has the highest melting
point because of its giant molecular structure (also called a giant
covalent structure). Every silicon atom is held to its neighbouring
silicon atoms by strong covalent bonds. However, its electrical
conductivity is much lower than the metals at the start of the period
because there are no delocalised electrons free to move around
within its structure. Silicon is classed as a semimetal, or metalloid.
The elements to the right of silicon are all non-metallic elements.
They exist as relatively small molecules. Sulfur exists as S8
molecules, phosphorus as P4 molecules and chlorine as Cl2
molecules. Although the covalent bonds within each molecule are
strong, there are only relatively weak instantaneous dipole– induced
dipole forces between their molecules (see Section 4.7). Therefore, it
does not take much energy to break these weak intermolecular forces
and melt the elements. At room temperature, phosphorus and sulfur
are solids with low melting points and chlorine is a gas. Argon gas
exists as single atoms with very weak instantaneous dipole–induced
dipole forces between these atoms.
Figure 10.6: Plotting the melting point against
atomic number for the elements in Period 3.

16. Periodic patterns of melting points and electrical conductivity

18. Periodic patterns of first ionisation energies

19. Periodic patterns of first ionisation energies
In general, the first ionisation energy increases
across Period 3 as the positive nuclear charge
increases and electrons successively fill the
third quantum shell. As electrons are in the
same shell, the shielding effect is similar in
atoms of each element. There are small ‘dips’ in
the general trend across the period between Mg
and Al, and between P and S. The same
pattern appears in Period 2 for Be and B, and N
and O. The explanation given in Section 2.6
also applies here in Period 3.
Figure 10.7: Plotting the first ionisation energy (X(g) → X+(g) + e−)
against atomic number for the elements in Period 3.

23. 10.3 Periodicity of chemical properties
Reactions of Period 3 elements with oxygen
Some of these experiments may be demonstrated by your teacher. Watch carefully to see
what happens and note down your observations.
1. Sodium reacts vigorously when heated and placed in a gas jar of oxygen. The sodium
burns with a bright yellow flame (Figure 10.8). The main product when sodium burns in a
limited amount of oxygen is a white solid, sodium oxide:
Na(s) + O2(g) → 2Na2O(s)
2. Magnesium also reacts vigorously when heated in oxygen, forming magnesium oxide.
Aluminium metal is protected by a layer of aluminium oxide, but powdered aluminium
does react well with oxygen. Both metals burn with bright white flames.
Mg(s) + O2(g) → 2MgO(s)
4Al(s) + 3O2(g) → 2Al2O3(s)
3. Silicon reacts slowly with oxygen to form silicon(IV) oxide (silicon dioxide):
Si(s) + O2(g) → SiO2(s)

24. 4. Phosphorus reacts vigorously with oxygen. A yellow or white flame is seen, and clouds
of white phosphorus(V) oxide are produced:
4P(s) + 5O2(g) → P4O10(s)
Reactions of Period 3 elements with oxygen
5. Sulfur powder, once ignited, burns gently with a blue flame in a gas jar of oxygen gas.
Toxic fumes of sulfur dioxide gas are produced (Figure 10.9):
S(s) + O2(g) → SO2(g)
Further oxidation of sulfur dioxide gives sulfur trioxide. Their systematic names are
sulfur(IV) oxide and sulfur(VI) oxide, respectively.
6. Chlorine and argon do not react with oxygen.

25. Reactions of Period 3 elements with chlorine
1. When sodium metal is heated then plunged into a gas jar of chlorine there is a
vigorous reaction, forming sodium chloride:
2Na(s) + Cl2(g) → 2NaCl(s)
2. Magnesium and aluminium also react vigorously with chlorine gas:
Mg(s) + Cl2(g) → MgCl2(s)
2Al(s) + 3Cl2(g) → Al2Cl6(s)
3. Silicon reacts slowly with chlorine, as it does with oxygen, giving silicon(IV)
chloride:
Si(s) + 2Cl2(g) → SiCl4(l)
4. Phosphorus also reacts slowly with excess chlorine gas:
2P(s) + 5Cl2(g) → 2PCl5(l)
5. Sulfur does form chlorides, such as SCl2 and S2Cl2, but you do not need to
cover these for your examination.
6. Argon does not form a chloride.

26. Reactions of sodium and magnesium with water
1. Sodium reacts vigorously with cold water, melting into a ball of molten metal (Figure
10.10). It moves across the surface of the water, giving off hydrogen gas. It quickly gets
smaller and smaller until it disappears, leaving a strongly alkaline solution (e.g. pH 14)
of sodium hydroxide behind:
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
2. By contrast, fresh magnesium reacts extremely slowly with cold water, taking several
days to produce a test-tube of hydrogen gas. The solution formed is very weakly
alkaline (e.g. pH 11), as any magnesium hydroxide formed is only slightly soluble.
Therefore, a lower concentration of OH−(aq) ions enters the solution compared with
the result when sodium is added to water. This is because sodium hydroxide is much
more soluble in water than magnesium hydroxide.
When heated, magnesium does react vigorously with water in the form of steam to
make magnesium oxide and hydrogen gas:
Mg(s) + H2O(g) → MgO(s) + H2(g)

29. 10.4 Oxides of Period 3 elements
Oxidation numbers of oxides
The maximum oxidation number of each element rises as we cross the
period. This happens because the Period 3 element in each oxide can use
all the electrons in its outermost shell in bonding to oxygen (ox. no. = −2).
They all exist in positive oxidation states because oxygen has a higher
electronegativity than any of the Period 3 elements. See Section 12.2 for
more about electronegativity.

30. 10.5 Effect of water on oxides and hydroxides of Period 3 elements
The oxides of sodium and magnesium react with water to form hydroxides. The presence of
excess aqueous hydroxide ions, OH−(aq), makes these solutions alkaline:
Magnesium oxide and magnesium hydroxide are commonly used in indigestion remedies
(Figure 10.11).
These basic compounds neutralise excess acid in the stomach, relieving the pain:
MgO(s) + 2HCl(aq) → MgCl2(aq) + H2O(l)
Mg(OH)2(s) + 2HCl(aq) → MgCl2(aq) + 2H2O(l)

31. Aluminium oxide does not react or dissolve in water, which is why an oxide layer can protect
aluminium metal from corrosion. However, it does react and dissolve when added to acidic
or alkaline solutions.
• With acid:
Al2O3(s) + 3H2SO4(aq) → Al2(SO4)3(aq) + 3H2O(l)
• With hot, concentrated alkali:
Al2O3(s) + 2NaOH(aq) + 3H2O(l) → 2NaAl(OH)4(aq)
When aluminium oxide reacts with an acid it behaves like a base: it forms a salt
(aluminium sulfate in the example with dilute sulfuric acid above) plus water.
When it reacts with an alkali it behaves like an acid: reacting to form a salt (sodium
tetrahydroxoaluminate in the example with sodium hydroxide above). Compounds that
can act as both acids and bases, such as aluminium oxide, are called amphoteric.

32. Silicon dioxide is also insoluble in water. Water cannot break down its giant molecular
structure. However, it will react with and dissolve in hot, concentrated alkali:
SiO2(s) + 2NaOH(aq) → Na2SiO3(aq) + H2O(l)
Silicon dioxide acts as an acid when it reacts with sodium hydroxide, forming a salt
(sodium silicate) plus water. It does not react with acids, so it is classed as an acidic
oxide.
Phosphorus(V) oxide reacts vigorously and dissolves in water to form an acidic solution
of phosphoric(V) acid (pH 2):

33. The oxides of sulfur, SO2 and SO3, both react and dissolve in water, forming acidic
solutions (pH 1):

34. The effect of bonding and electronegativity on the structure and acidic / basic nature of
Period 3 oxides
Table 10.9 shows a summary of the acidic / basic nature of the Period 3 oxides. You need to
know this summary for your examination. We can explain the behaviour of the oxides by
looking at their structure and bonding (Table 10.10 and Figure 10.12).

36. Going across a period, the elements get more electronegative as electrons are more
strongly attracted by the increasing positive nuclear charge (see Section 4.7). The
electronegativity values, which indicate the strength of the attraction of an atom for
the electrons in a bond, are shown in Table 10.11.

37. The electronegativity of oxygen is 3.5. The greater the difference in electronegativity
between the Period 3 element and oxygen, the more likely it is that the oxide will have
ionic bonding. Electrons will be transferred from sodium, magnesium and aluminium
atoms (forming positively charged ions) to oxygen atoms (forming O2− ions) when their
oxides are formed. The other Period 3 elements will form covalently bonded oxides in
which bonding electrons are shared.
Notice the high melting points of the giant ionic and giant covalent structures, leading to
the use of:
• magnesium oxide to line the inside of furnaces
• aluminium oxide and silicon dioxide to make ceramics, with giant covalent structures
designed to withstand high temperatures and provide electrical insulation.

38. The oxides of the metals sodium and magnesium, with purely ionic bonding, produce
alkaline solutions with water as their oxide ions, O2−(aq), become hydroxide ions, OH−(aq).
The oxide ions behave as bases by accepting H+ ions from water molecules:
O2−(aq) + H2O(l) → 2OH−(aq)
By contrast, the covalently bonded non-metal oxides of phosphorus and sulfur dissolve
and react in water to form acidic solutions. The acid molecules formed donate H+ ions to
water molecules, behaving as typical acids. For example, sulfuric(VI) acid:
H2SO4(aq) + H2O(l) → H3O+(aq) + 𝐻𝑆𝑂4
−
(aq)
The insoluble oxides of aluminium and silicon show their acidic nature by reacting and
dissolving in an alkaline solution, such as hot, concentrated sodium hydroxide solution,
forming a soluble salt. This behaviour is typical of a covalently bonded oxide. However,
aluminium oxide also reacts and dissolves in acidic solutions, forming a soluble salt:
behaviour typical of a basic metal oxide with ionic bonding. Because aluminium oxide
behaves in these two ways, this provides evidence that the chemical bonding in aluminium
oxide is not purely ionic nor purely covalent. It is amphoteric.

40. 10.6 Chlorides of Period 3 elements
Oxidation numbers of the Period 3 chlorides
The oxidation numbers rise as we cross Period 3, until we reach sulfur in Group 16. This
happens because the Period 3 elements from sodium to phosphorus use all the electrons in
their outermost shell, their valence electrons, in bonding to chlorine (ox. no. = −1). They all
exist in positive oxidation states because chlorine has a higher electronegativity than any of
the other Period 3 elements (see Table 10.12).

41. 10.7 Effect of water on chlorides of Period 3 elements
As with the oxides of Period 3 elements, the chlorides also show characteristic behaviour
when we add them to water. Once again, this is linked to their structure and bonding (Table
10.13).

42. At the start of Period 3, the ionic chlorides of sodium (NaCl) and magnesium (MgCl2) do not
react with water. The polar water molecules are attracted to the ions, dissolving the
chlorides by breaking down the giant ionic structures. The solutions formed contain the
positive metal ions and the negative chloride ions surrounded by water molecules. The
metal ions and the chloride ions are called hydrated ions:

43. Aluminium chloride is sometimes
represented as AlCl3, which suggests that
its chemical bonding is likely to be ionic:
with Al3+ ions and Cl− ions in a giant lattice.
In solid hydrated aluminium chloride
crystals, this is the case. However, without
water it exists as Al2Cl6. This can be
thought of as a dimer (two molecules
joined together) of AlCl3. Al2Cl6 is a
covalently bonded molecule (see Figure
10.13).
Figure 10.13: The chemical bonding in Al2Cl6.

44. • Each relatively small and highly charged Al3+ ion is hydrated and causes a water
molecule bonded to it to lose an H+ ion. This turns the solution acidic. We can show this
in an equation as follows:
[Al(H2O)6]3+(aq) → [Al(H2O)5OH]2+(aq) + H+(aq)
• The non-metal chlorides SiCl4 and PCl5 are hydrolysed in water, releasing white fumes
of hydrogen chloride gas in a rapid reaction (Figure 10.14).
SiCl4(l) + 2H2O(l) → SiO2(s) + 4HCl(g)
• The SiO2 is seen as an off-white precipitate. Some of the hydrogen chloride gas
produced dissolves in the water, leaving an acidic solution (hydrochloric acid).
Phosphorus(V) chloride also undergoes hydrolysis when added to water:
PCl5(s) + 4H2O(l) → H3PO4(aq) + 5HCl(g)
Both products are soluble in water and are highly acidic.

46. 10.8 Deducing the position of an element in the Periodic Table
A. Element G forms a chloride, which reacts with water to form a solution of pH 1. It forms an
oxide which has a melting point of 1610 °C. The oxide does not dissolve in or react with
aqueous sodium hydroxide. Deduce the possible position of G in the Periodic Table.
Solution
1. Use the pH data to decide whether G is likely to be a metal or non-metal.
Low pH suggests that it is a metal chloride which undergoes hydrolysis, so is likely to be in
Group 13 to 17.
2. Use the solubility data.
The oxide doesn’t dissolve in sodium hydroxide. This rules out aluminium oxide because
this reacts with sodium hydroxide. So G could be a giant covalent structure or simple
molecular structure.
3. Use melting point data.
The melting point is high so it is likely to be a giant covalent structure in Group 14.
4. It cannot be in Period 2 because CO2 is a gas. So it could be in Period 3 or lower.

47. 10.8 Deducing the position of an element in the Periodic Table
B. Selenium is in Group 16 and Period 4 of the Periodic Table. Predict some physical
and chemical properties of selenium.
1. Identify the likely structure of selenium from its position in the Periodic Table and comparison with
other Group 16 elements.
It’s a non-metal in Group 16 so comparing with sulfur it should have a simple molecular structure.
2. Identify the physical properties related to the structure.
Simple molecular structures have (relatively) low melting points, do not conduct electricity and are
insoluble in water.
3. Identify the chemical properties related to reaction with water, reaction with chlorine and reaction
with oxygen.
Simple molecules do not react with water.
Reacts with chlorine to form simple molecule SeCl4 (which reacts with water vapour in the air to
produce hydrogen chloride).
Reacts with oxygen to form an oxide of possible formula SeO2 (by comparison with sulfur).