Periodic table

Class 11 Chemistry, unit 10, book pages 212–228

Practise 67 MCQs Take the unit quiz Review 24 cards

The periodic table arranges the elements in order of increasing atomic number, so that their properties repeat in periods and groups. This unit shows how the electronic configuration of an element fixes its period, group and block, and how the elements are classified as metals, non-metals or metalloids. It then explains the trends in electron affinity, ionization energy and electronegativity, and uses them to predict the reactions of sodium and magnesium, the acid and base character of oxides, the pH of chlorides in water, and the identity of unknown elements.

Learning outcomes (18)

Explain the arrangement of elements in the periodic table.
Explain that the periodic table is arranged into four blocks associated with the four sub-levels s, p, d and f.
Recognize that the period number (n) is the outer energy level that is occupied by electrons.
State that the number of the principal energy level and the number of valence electrons in an atom can be deduced from its position on the periodic table.
Identify the position of metals, non-metals and metalloids in the periodic table.
Explain that vertical and horizontal trends in the periodic table exist for atomic radius, ionic radius, ionization energy, electron affinity and electronegativity.
Recognize that trends in metallic and non-metallic behaviour are due to the trends in valence electrons.
Deduce the electronic configuration of an atom from the element's position in the periodic table, and vice versa (based on s, p, d and f sub-shells).
Write equations for the reactions of Na and Mg with oxygen, chlorine, and water.
Explain the variation in the oxidation number of the oxides and chlorides (NaCl, MgCl2) in terms of their outer shell (valence shell) electrons.
Describe (including writing equations for) the reactions, if any, of the oxides (acidic and basic) with water (including the likely pHs of the solutions obtained).
Explain with the help of equations for the acid / base behaviour of the oxides and hydroxides NaOH and Mg(OH)2, including amphoteric behaviour where relevant, in reactions with acids and bases (NaOH only).
Explain with equations for the reactions of the chlorides with water, including the likely pHs of the solution obtained.
Explain the variations and trends in terms of bonding and electronegativity.
Suggest the type of chemical bonding present in the chlorides and oxides from observation of their physical and chemical properties.
Predict the characteristic properties of an element in a given group by using knowledge of chemical periodicity.
Deduce the nature, possible position in the periodic table and identity of unknown elements from given information about physical and chemical properties.
Explain the trends in the ionization energy and electron affinity of the Group 1 and Group 17 elements.

The periodic law

Chemists have always searched for order among the elements. The most important result of that search is the systematic arrangement of the elements, which is considered the most important achievement in chemistry.

When Moseley discovered atomic numbers in 1913, it became clear that atomic numbers could provide the basis for a systematic arrangement of the elements. The modern periodic table therefore shows the elements in order of increasing atomic number, together with their repeating properties.

A table that shows the systematic arrangement of the elements is called a periodic table. It is based on the periodic law.

Periodic law: When elements are arranged in order of increasing atomic number, their properties repeat periodically.
Periodic table: A table that shows the systematic arrangement of the elements, based on the periodic law.

10.1Periods and groups of elementsOften examined

Based on the periodic law, the elements are arranged in seven horizontal rows called periods. Elements with similar chemical properties lie in the same vertical columns, which are called groups (or families).

The elements are listed in ascending atomic number, from left to right and from top to bottom. Hydrogen (H, atomic number 1) is in the upper left corner. Helium (He, atomic number 2) is in the upper right corner. Lithium (Li, atomic number 3) is at the left end of the second row.

Figure 10.1: The periodic table of elements. Periods are the horizontal rows, numbered 1 to 7. Groups are the vertical columns, numbered 1 to 18 (IUPAC). The two rows at the bottom are the f-block elements.

Periods

The number of elements per period varies from 2 in period 1 to 32 in period 7. The first three periods are called short periods and the rest are long periods. The properties of the elements change gradually from left to right across a period. From one period to the next, the pattern of properties repeats itself, in accordance with the periodic law.

Numbering the groups

Two numbering systems are often used. In the traditional system and the old IUPAC system the letters A and B are used. The first two groups are IA and IIA, the last six groups are IIIA to VIIIA, and the middle groups are group B. In 1988 IUPAC decided that the groups would be numbered 1 to 18 from left to right.

Group numbering and names
Old system	IUPAC (1988)	Name given in the book
IA	1	Alkali metals
IIA	2	Alkaline earth metals
Group B (middle)	3 to 12	Transition elements
IIIA	13	Boron family
IVA	14	-
VA	15	-
VIA	16	Oxygen family (chalcogens)
VIIA	17	Halogens
VIIIA	18	Noble gases

Elements in the same group have the same number of valence electrons, so the group number shows the number of valence electrons. Groups 1 and 2 have 1 and 2 valence electrons. Group 13 elements have 3 valence electrons, group 14 elements have 4, group 15 elements have 5, and so on. For groups 13 to 18, the group number equals the number of valence electrons plus 10.

Groups are named after the topmost element of the group, and some groups have special names. The metallic elements of group 1 are the alkali metals. Group 2 elements are the alkaline earth metals. Group 13 is the boron family. Group 16 is the oxygen family, also called the chalcogens. Group 17 elements are the halogens. Group 18 elements are the noble gases, because they do not readily undergo chemical reactions.

Group A elements are called normal or representative elements, and are also called main group elements. Group B elements are called transition elements.

10.2Blocks of elements in the periodic table

The block of an element is its position in the periodic table, based on the type of sub-shell that holds its valence electrons. The valence sub-shell may be s, p, d or f, so the table has four blocks: the s-block, the p-block, the d-block and the f-block.

s-block

Elements of groups 1 and 2 (except He) are the s-block elements. They hold their valence electrons in s sub-shells. The alkali and alkaline earth metals are in the s-block.

p-block

Elements of groups 13 to 18 are the p-block elements. They hold their valence electrons in p sub-shells. This block contains both metals and non-metals.

d-block

Elements of groups 3 to 12 (sub-group B) are the d-block elements. They hold their valence electrons in d sub-shells. These elements are called transition metals.

f-block

The two rows at the bottom of the periodic table are the f-block elements. They hold their valence electrons in f sub-shells. They are also called the lanthanides and actinides, or inner transition metals.

Figure 10.2: The four blocks of the periodic table. The s-block is groups 1 and 2, the d-block is groups 3 to 12, the p-block is groups 13 to 18, and the f-block is the two rows at the bottom.

Block	Groups	Valence sub-shell	Description in the book
s	1 and 2 (except He)	s	Alkali and alkaline earth metals
p	13 to 18	p	Contains both metals and non-metals
d	3 to 12	d	Transition metals (sub-group B)
f	Two rows at the bottom	f	Lanthanides and actinides (inner transition metals)

10.3Position of an element in the periodic tableOften examined

The electronic configuration of an element gives basic information about its position in the periodic table. By examining the valence shell configuration, the n value and the sub-shell type, you can identify the element's group, period and block.

To find the group, count the electrons in the valence shell. For an s-block element, the group number is the number of valence electrons. For a p-block element, the group number is the number of valence electrons plus 10.

The period is the horizontal row. It equals the largest n value of the valence shell. If the valence electrons are in the L shell (n = 2), the element belongs to period 2. If n = 3, it belongs to period 3.

Example 10.1: Position of nitrogen and oxygen

Find out the position of the following elements in the periodic table: (a) nitrogen (atomic number 7), (b) oxygen (atomic number 8).

(a) Nitrogen

Write the electron configuration of N (Z = 7).
$1 s^{2} 2 s^{2} 2 p^{3}$
The valence shell is $2 s^{2} 2 p^{3}$ . Its n value is 2, so the period number is 2.
Count the valence electrons: 2 + 3 = 5. The group number is 5 + 10 = 15 (group VA).
$Group = 2 + 3 + 10 = 15$

(b) Oxygen

Write the electron configuration of O (Z = 8).
$1 s^{2} 2 s^{2} 2 p^{4}$
The valence shell is $2 s^{2} 2 p^{4}$ . Its n value is 2, so the period number is 2.
Count the valence electrons: 2 + 4 = 6. The group number is 6 + 10 = 16 (group VIA).
$Group = 2 + 4 + 10 = 16$

N: period 2, group 15 (VA); O: period 2, group 16 (VIA)

Worked example: position of aluminium and calcium from atomic number

Find the period, group and block of (a) aluminium (atomic number 13) and (b) calcium (atomic number 20).

Given

$Z (Al) = 13$
$Z (Ca) = 20$

To find

\text{period, group and block of Al and Ca}

Write the electronic configuration of aluminium, $Z = 13$ .
$1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{1}$
The valence shell has the largest $n$ , which is 3, so aluminium is in period 3. The valence electrons are $3 s^{2} 3 p^{1}$ : three in all, and the last electron goes into a p sub-shell, so it is a p-block element. For the p-block the group number is the number of valence electrons plus 10.
$group = 3 + 10 = 13$
Write the electronic configuration of calcium, $Z = 20$ .
$1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2}$
The largest $n$ is 4, so calcium is in period 4. The valence shell is $4 s^{2}$ : two electrons in an s sub-shell, so it is an s-block element and the group number is just the number of valence electrons.
$group = 2$
Check against the common mistake. Aluminium has three valence electrons but it is in group 13, not group 3, because the p-block rule adds 10. Group 3 belongs to the d-block.

Al : period 3, group 13, p-block; Ca : period 4, group 2, s-block

The valence shell configuration can also be found from an element's position. The period number gives the n value of the valence shell, and the group number gives the total number of valence electrons (for the p-block, after subtracting 10).

Example 10.2: Valence shell configuration of phosphorus and neon

Obtain the valence shell electronic configuration of (a) phosphorus and (b) neon.

Use the rules: the period number is the n value of the valence shell, and the group number is the number of valence electrons. Distribute the electrons in the sub-shells of the valence shell.

(a) Phosphorus is in period 3, so n = 3 and the valence shell is M, with 3s and 3p sub-shells. Its group number is V, so there are 5 valence electrons.
2 electrons fill the 3s sub-shell and the remaining 3 fill the 3p sub-shell.
$3 s^{2} 3 p^{3}$
(b) Neon is in period 2, so n = 2 and the valence shell is L, with 2s and 2p sub-shells. Its group number is VIII, so there are 8 valence electrons.
2 electrons fill the 2s sub-shell and the remaining 6 fill the 2p sub-shell.
$2 s^{2} 2 p^{6}$

(a) 3 s^{2} 3 p^{3}; (b) 2 s^{2} 2 p^{6}

Concept assessment exercise 10.1 (1): Valence shell of Mg and Cl

Find the valence shell electronic configuration of Mg and Cl from their positions in the periodic table.

Mg (Z = 12): the configuration is $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2}$ . The valence shell is $3 s^{2}$ , so it is in period 3 and group 2 (2 valence electrons).
Cl (Z = 17): the configuration is $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{5}$ . The valence shell is $3 s^{2} 3 p^{5}$ , so it is in period 3. It has 2 + 5 = 7 valence electrons, so the group number is 7 + 10 = 17.

Mg: 3 s^{2}; Cl: 3 s^{2} 3 p^{5}

Concept assessment exercise 10.1 (2): Position of K and S

Find the position of K (atomic number 19) and S (atomic number 16) in the periodic table.

K (Z = 19): the configuration is $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{1}$ . The valence shell is $4 s^{1}$ , so it is in period 4 with 1 valence electron, in group 1.
S (Z = 16): the configuration is $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{4}$ . The valence shell is $3 s^{2} 3 p^{4}$ , so it is in period 3. It has 2 + 4 = 6 valence electrons, so the group number is 6 + 10 = 16 (group VIA).

K: period 4, group 1; S: period 3, group 16 (VIA)

10.4Metals, non-metals and metalloidsOften examined

The periodic table gives a general framework for organizing the elements. On the basis of their properties and their position in the table, the elements are generally classified in three categories: metals, non-metals and metalloids.

Metals appear on the left side of the periodic table. Metals are good conductors, malleable and ductile.

Non-metals are on the right side of the periodic table. They are bad conductors and are non-malleable and non-ductile.

A leader line separates metals from non-metals. Metals lie on the left of the line and non-metals on the right. Hydrogen is the exception: it is a non-metal, but it lies on the left. The elements next to the leader line are the metalloids. They possess the properties of both metals and non-metals, and they are also called semi-metals.

Table 10.3: Position of metals, non-metals and metalloids in the periodic table. Metals are shaded, non-metals are outlined, and the metalloids lie along the boundary between metals and non-metals.

Feature	Metals	Non-metals	Metalloids
Position	Left side of the table	Right side of the table (hydrogen is also a non-metal)	Next to the leader line
Conductivity	Good conductors	Bad conductors	Properties of both metals and non-metals
Malleable and ductile	Yes	No (non-malleable and non-ductile)	Properties of both metals and non-metals
Also called	-	-	Semi-metals

10.5Periodicity of properties

The electronic configuration of the elements shows a periodic fluctuation with increasing atomic number. Therefore, the physical and chemical properties of the elements vary in a periodic manner.

Elements with similar valence shell electronic configurations are placed in the same group, one below the other. Chemical properties depend on the electronic configuration of the valence shell. Because all the elements of a group have a similar valence shell configuration, they have similar chemical properties.

Physical properties depend on the size of the atoms. Atomic size changes gradually from top to bottom in a group (as the book shows in its earlier discussion of atomic size). The elements of a group therefore show a gradation of physical properties.

In a period, the number of electrons in the valence shell increases gradually from left to right. The chemical and physical properties of the elements also differ in the same gradual way across the period.

10.5.1Electron affinityOften examined

Electron affinity explains the formation of anions. It is defined as the amount of energy released when an electron is added to the valence shell of an isolated gaseous atom to form a uni-negative gaseous ion. It is also called the first electron affinity.

Electron affinity: The amount of energy released when an electron is added to the valence shell of an isolated gaseous atom to form a uni-negative gaseous ion (also called the first electron affinity).
Shielding effect: The reduction in the force of attraction between the nucleus and the valence electrons, caused by the electrons present in the inner sub-shells.

X X_{(g)} + e X^{-} X X^{-} X_{(g)} (energy is released)

Figure 10.7 shows the electron affinities of the main group elements. Three factors affect electron affinity.

1. Nuclear charge

Electron affinity tends to increase from left to right in the periodic table. This is mainly due to the increase in nuclear charge. As the number of protons in the nucleus increases, the attraction for electrons increases, so the atom accepts an additional electron more easily.

2. Atomic size

Electron affinity generally decreases from top to bottom in a group, because atomic size increases. Larger atoms have their electrons farther from the nucleus, so the attraction between the nucleus and the outer electrons is weaker. It is therefore more difficult for larger atoms to accept extra electrons.

3. Shielding effect

The electrons in the inner shells partly shield the outer electrons from the attraction of the nucleus. The more inner shells there are, the more the outer electrons are shielded. This shielding can lead to a decrease in electron affinity.

Across a period, electron affinity increases. The nuclear charge increases and the atomic size decreases, which binds the extra electron more tightly to the nucleus, while the shielding effect stays about the same. Down a group, electron affinity decreases. The increase in shielding and in atomic size makes the added electron bind less tightly, so less energy is released.

Figure 10.7: Electron affinities of the main group elements, as printed in the book. Most values are negative, so a more negative value means more energy released.
Period	Group 1	Group 2	Group 13	Group 14	Group 15	Group 16	Group 17	Group 18
1	H −73	-	-	-	-	-	-	He 0
2	Li −60	Be 0	B −27	C −122	N +7	O −141	F −328	Ne 0
3	Na −53	Mg 0	Al −44	Si −134	P −71.7	S −200	Cl −349	Ar 0
4	K −48	Ca 0	Ga −29	Ge −120	As −77	Se −195	Br −325	Kr 0
5	Rb −47	Sr 0	In −29	Sn −121	Sb −101	Te −190	I −295	Xe 0
6	Cs −45	Ba 0	Tl −30	Pb −110	Bi −110	Po −180	At −270	Rn 0

10.6.1Metallic behaviour

The valence electrons play a key role in determining the chemical and physical behaviour of the elements. The next sections use this idea to explain the trends in metallic and non-metallic behaviour.

Metallic behaviour

Metals are usually found on the left side of the periodic table. They tend to lose valence electrons easily and form cations, so that they reach a stable electron configuration.

This behaviour is due to three reasons:

Metals usually have 1 to 3 valence electrons.
Their ionization energy is usually low, so they can lose these electrons easily.
They have low electronegativity and do not tend to attract electrons.

10.6.2Non-metallic behaviour

Non-metals are usually found on the right side of the periodic table. Their non-metallic behaviour is due to their tendency to gain or share valence electrons in order to achieve a stable structure.

Non-metals usually have 5 to 7 valence electrons. They need only a few more electrons to fill their valence shell.

Non-metals have a high ionization energy. This makes them less likely to lose electrons and more likely to share or gain electrons. They also have high electronegativity, which allows them to gain electrons easily, or to share electrons to form covalent bonds.

10.6.3Trends in groups

Group 1 alkali metals have a strong metallic character. Group 17 (VIIA) halogens have a strong non-metallic character.

The metallic character of the group 1 elements increases as you move down the group. This is due to the increase in atomic radii and their ability to lose electrons easily.

Below group 17, non-metallic behaviour decreases. This is due to the increase in atomic radii and the decrease in their ability to attract electrons.

Group 14 (IVA) consists of carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb). Going down the group, the non-metallic character decreases and the metallic character increases.

Group 14: the change from non-metal to metal going down the group
Element (group 14)	Classification in the book
C (carbon)	Typical non-metal
Si (silicon)	Metalloid
Ge (germanium)	Metalloid
Sn (tin)	Typical metal
Pb (lead)	Typical metal

10.6.4Electronegativity and type of bondOften examined

Electronegativity: The ability of an atom to attract the electrons towards itself in a chemical bond.
Electronegativity difference (ΔEN): The difference in electronegativity between two bonded atoms. It gives a rough indication of the expected nature of the bond.

The electronegativity difference (ΔEN) of the bond between two bonded atoms gives a rough indication of the expected nature of the bond, and hence the bond type. The ionic character of a bond increases as the difference in electronegativity of the bonded atoms increases.

Δ E N = E N_{A} - E N_{B}

ΔEN	Bond type (rough indication)
Greater than 1.8	Ionic
0.4 to 1.8	Polar covalent
Less than 0.4	Covalent (non-polar)

Bond type of

NaCl

and

MgO

(ionic bonds)

Use the electronegativity difference to find the bond type in

NaCl

and

MgO

Given

$E N (Na) = 0.93$
$E N (Cl) = 3.16$
$E N (Mg) = 1.31$
$E N (O) = 3.44$

To find

\text{bond type in $NaCl$ and $MgO$ }

$NaCl$ : the difference is 3.16 minus 0.93. This is 2.23, which is greater than 1.8, so the bond in $NaCl$ is ionic.
$Δ E N = 3.16 - 0.93 = 2.23$
$MgO$ : the difference is 3.44 minus 1.31. This is 2.13, which is greater than 1.8, so the bond in $MgO$ is ionic.
$Δ E N = 3.44 - 1.31 = 2.13$

NaCl and MgO both have ionic bonds

Bond type of C-H and H-F (covalent bonds)

Use the electronegativity difference to find the bond type in C-H and H-F.

Given

$E N (C) = 2.55$
$E N (H) = 2.2$
$E N (F) = 3.98$

To find

\text{bond type in C-H and H-F}

C-H: the difference is 2.55 minus 2.2 = 0.35. This is less than 0.4, so the C-H bond is non-polar covalent.
$Δ E N = 2.55 - 2.2 = 0.35$
H-F: the difference is 3.98 minus 2.2 = 1.78. This lies between 0.4 and 1.8, so the H-F bond is polar covalent.
$Δ E N = 3.98 - 2.2 = 1.78$

C-H: non-polar covalent; H-F: polar covalent

Worked example: ranking bonds by ionic character

Use the electronegativity differences to classify the bonds Cl-Cl, H-Cl, O-H, Mg-Cl and Na-F, and arrange them in order of increasing ionic character.

Given

$E N (H) = 2.2$
$E N (O) = 3.44$
$E N (Cl) = 3.16$
$E N (Mg) = 1.31$
$E N (Na) = 0.93$
$E N (F) = 3.98$

To find

\text{bond type of each bond and the order of ionic character}

Find ΔEN for each bond by subtracting the smaller electronegativity from the larger.
$Cl-Cl H-Cl O-H Mg-Cl Na-F = 3.16 - 3.16 = 0 = 3.16 - 2.2 = 0.96 = 3.44 - 2.2 = 1.24 = 3.16 - 1.31 = 1.85 = 3.98 - 0.93 = 3.05$
Compare each value with the limits: below 0.4 is covalent, 0.4 to 1.8 is polar covalent, above 1.8 is ionic. Cl-Cl is covalent (the atoms are identical, so ΔEN is 0). H-Cl and O-H are polar covalent. Mg-Cl (1.85) and Na-F (3.05) are ionic.
Mg-Cl is only just above 1.8, so it is a borderline case. ΔEN gives a rough indication only, which is why the figure shows it close to the polar covalent region.
Ionic character increases as ΔEN increases, so the order is the order of the numbers.

Bonds of this example on the ΔEN scale. Mg–Cl (accent) lies just inside the ionic region.

Cl-Cl < H-Cl < O-H < Mg-Cl < Na-F

The bond type scale from ΔEN. The bonds worked above fall at 0.35 (C-H, non-polar covalent), 1.78 (H-F, polar covalent), 2.13 (

MgO

, ionic) and 2.23 (

NaCl

, ionic).

10.6.5Trends in chemical propertiesOften examined

This section looks at the reactions of two typical metals, sodium (group 1) and magnesium (group 2), with oxygen, chlorine and water.

Reaction with oxygen

Sodium is a silvery white soft metal and an extremely reactive element. It readily combines with oxygen. In a limited supply of oxygen, sodium burns to form sodium oxide. In excess oxygen, it forms pale yellow sodium peroxide.

4 Na + O X_{2} 2 Na X_{2} O

2 Na + O X_{2} Na X_{2} O X_{2}

Magnesium is relatively hard and is also an extremely reactive element. When ignited, magnesium burns in oxygen with an intense white flame to give solid magnesium oxide.

2 Mg + O X_{2} 2 MgO

Reaction with chlorine

Sodium burns in chlorine with a bright yellow flame, producing a white solid, sodium chloride.

2 Na + Cl X_{2} 2 NaCl

Magnesium also burns in chlorine with an intense white flame, producing a white solid, magnesium chloride.

Mg + Cl X_{2} MgCl X_{2}

Reaction with water

Sodium reacts violently with cold water, producing hydrogen with a light explosion.

2 Na + 2 H X_{2} O 2 NaOH + H X_{2}

Magnesium reacts slowly with cold water to form a thin layer of magnesium hydroxide on its surface. This protective coating stops further reaction. In steam, however, magnesium burns with a white flame, forming magnesium oxide and hydrogen.

Mg + H X_{2} O (g) MgO + H X_{2}

10.6.6Oxidation number of oxides and chloridesOften examined

The oxidation number of an element in its oxides and chlorides corresponds to the number of electrons it uses for bonding. The maximum oxidation number of an element is the same as its total number of valence electrons. It is always positive, since oxygen and chlorine are more electronegative than the other elements.

Take period 3 as an example. In the oxides of Na, Mg, Al, Si, P, S and Cl, the maximum oxidation number increases across the period, from +1 in Na to +7 in Cl.

Maximum oxidation numbers of the period 3 elements in their oxides
Element (period 3)	Oxide	Maximum oxidation number
Na	\ce{Na2O}	+1
Mg	\ce{MgO}	+2
Al	\ce{Al2O3}	+3
Si	\ce{SiO2}	+4
P	\ce{P2O5}	+5
S	\ce{SO3}	+6
Cl	\ce{Cl2O7}	+7

Phosphorus and sulphur show variable oxidation numbers, because they can expand their octet. For example, in \ce{SO2} sulphur has oxidation number +4, but in \ce{SO3} it has +6.

10.6.6Acid and base character of oxidesOften examined

Metal oxides are generally basic in character, because they produce bases in water (pH above 7). Non-metal oxides are generally acidic, because they produce acids in water (pH below 7).

Reactions of oxides with water

\ce{Na2O + H2O -> 2NaOH}

\ce{MgO + H2O -> Mg(OH)2}

\ce{CO2 + H2O -> H2CO3}

\ce{P4O10 + 6H2O -> 4H3PO4}

The basic character of oxides decreases as you move along a period, which shows a decrease in metallic character. The electronegativity difference between each element and oxygen also decreases across the period, so the bonding changes from ionic to covalent.

The oxides of sodium and magnesium are ionic compounds. They have high melting and boiling points, are soluble in water, and their aqueous solutions conduct electricity. The other oxides are covalent. The oxides of P, S and Cl exist as gaseous molecules with low melting and boiling points.

The oxides of the period 3 elements change from strongly basic, through weakly basic and amphoteric, to weakly acidic and strongly acidic. Table 10.7 summarizes this trend.

Table 10.7: Nature of the oxides of the elements in period 3
Group	IA	IIA	IIIA	IVA	VA	VIA	VIIA
Period 3 element	Na	Mg	Al	Si	P	S	Cl
Oxide	\ce{Na2O}	\ce{MgO}	\ce{Al2O3}	\ce{SiO2}	\ce{P4O10}	\ce{SO2}	\ce{Cl2O7}
Nature of oxide	Strongly basic	Basic	Amphoteric (both acidic and basic)	Weakly acidic	Acidic	Strongly acidic	Very strongly acidic

Reactions of the oxides with acids and bases

Sodium oxide and sodium hydroxide are strongly basic. They react with acids to form a salt and water.

\ce{Na2O + 2HCl -> 2NaCl + H2O}

\ce{NaOH + HCl -> NaCl + H2O}

Magnesium oxide and magnesium hydroxide are also basic, but they are less basic than those of sodium.

\ce{MgO + H2SO4 -> MgSO4 + H2O}

\ce{Mg(OH)2 + H2SO4 -> MgSO4 + 2H2O}

Aluminium oxide and aluminium hydroxide are amphoteric. They react with both acids and bases.

\ce{Al2O3 + 2NaOH + 3H2O -> 2NaAl(OH)4}\quad\text{(sodium aluminate)}

\ce{Al2O3 + 6HCl -> 2AlCl3 + 3H2O}

Worked example: what amphoteric means for aluminium oxide

Aluminium oxide reacts with hydrochloric acid and also with sodium hydroxide. (a) Check that the equation with HCl is balanced. (b) State the role of \ce{Al2O3} in each reaction and explain why it is called amphoteric.

Count the atoms in \ce{Al2O3 + 6HCl -> 2AlCl3 + 3H2O}. Left: Al 2, O 3, H 6, Cl 6. Right: Al 2, Cl 6, H 6, O 3. Every element matches, so the equation is balanced.
\ce{Al2O3 + 6HCl -> 2AlCl3 + 3H2O}
With the acid, \ce{Al2O3} forms a salt and water, which is how a basic oxide behaves (like \ce{Na2O + 2HCl -> 2NaCl + H2O}). So here it acts as a base.
With sodium hydroxide, \ce{Al2O3} reacts with the base and forms sodium aluminate, so here it acts as an acid.
\ce{Al2O3 + 2NaOH + 3H2O -> 2NaAl(OH)4}
An oxide that reacts with both acids and bases is amphoteric. Aluminium lies between the strongly basic oxides on the left of period 3 and the acidic oxides on the right, so its oxide shows both behaviours.

The two roles of aluminium oxide. The accent marks the oxide that behaves both ways.

\ce{Al2O3}\ \text{is a base towards acids and an acid towards bases: amphoteric}

10.6.7Reactions of chlorides with water

Metal chlorides, such as sodium chloride and magnesium chloride, are ionic in nature. They dissolve in water to form neutral solutions with pH 7.

Aluminium chloride (\ce{AlCl3}), silicon tetrachloride (\ce{SiCl4}) and phosphorus pentachloride (\ce{PCl5}) are covalent compounds. They are also soluble in water, but they react more vigorously with water to form acidic solutions with low pH.

pH of the solutions formed by chlorides in water, as printed in the book
Chloride	Behaviour in water	pH of the solution
\ce{NaCl}	Dissolves to give a neutral solution	7
\ce{MgCl2}	Dissolves to give a neutral solution	7
\ce{AlCl3}	\ce{AlCl3 + 3H2O -> Al(OH)3 + 3HCl}	3
\ce{SiCl4}	\ce{SiCl4 + 4H2O -> Si(OH)4 + 4HCl}	0
\ce{PCl5}	\ce{PCl5 + 4H2O -> H3PO4 + 5HCl}	0

The equations show that the acidity comes from the hydrogen chloride (\ce{HCl}) formed in each reaction with water.

Worked example: chlorides of period 3 in water

(a) Show that \ce{PCl5 + 4H2O -> H3PO4 + 5HCl} is balanced. (b) How many moles of \ce{HCl} form when 0.25\,\mathrm{mol} of \ce{SiCl4} reacts completely with water? (c) Why is a solution of \ce{NaCl} neutral while that of \ce{SiCl4} is strongly acidic?

Given

n(\ce{SiCl4}) = 0.25\,\mathrm{mol}

To find

n(\ce{HCl})

(a) Count atoms. Left: P 1, Cl 5, H 8, O 4. Right: P 1, H 3+5=8, Cl 5, O 4. All match.
\ce{PCl5 + 4H2O -> H3PO4 + 5HCl}
(b) From \ce{SiCl4 + 4H2O -> Si(OH)4 + 4HCl}, 1 mol of \ce{SiCl4} gives 4 mol of \ce{HCl}. Multiply the amount of \ce{SiCl4} by 4.
n(\ce{HCl})=4\times0.25\,\mathrm{mol}=1.0\,\mathrm{mol}
(c) \ce{NaCl} is ionic. It only dissolves, no acid is made, and the pH stays 7. \ce{SiCl4} is covalent and reacts with water, producing \ce{HCl}, which is a strong acid, so the pH falls to 0.

n(\ce{HCl})=1.0\,\mathrm{mol}

10.6.8Type of chemical bonding in chlorides and oxides

By looking at their physical and chemical properties, we can infer the type of chemical bonds present in a compound. Ionic compounds usually have higher melting and boiling points, dissolve in water, and conduct electricity when molten. Covalent compounds often have lower freezing and boiling points, are insoluble in water, and do not conduct electricity.

Property	Ionic compound	Covalent compound
Melting and boiling points	Usually higher	Often lower
Solubility in water	Usually dissolves	Usually insoluble
Electrical conduction	Conducts when molten	Does not conduct

Specific properties and reactivity help to distinguish between ionic and covalent bonding. In some cases, a compound has a combination of both bond types, with both ionic and covalent character.

10.6.9Trends in ionization energy and electron affinity of groups 1 and 17Often examined

As you go down group 1, the ionization energy decreases. This is due to two factors:

An increase in the shielding effect.
An increase in atomic radii.

Both factors reduce the attraction between the nucleus and the valence electrons. Therefore, removing an electron requires less energy.

Worked example: why ionization energy falls from sodium to potassium

A student says: potassium has more protons than sodium, so it should hold its outer electron more strongly and have a higher ionization energy. Explain the mistake.

Both elements are in group 1 and have one valence electron. Sodium has three shells and potassium has four, as the figure shows.
Atomic radius: the extra shell makes the potassium atom larger, so its valence electron is farther from the nucleus.
Shielding effect: potassium has three inner shells between the nucleus and the valence electron, sodium only two, so the attraction is reduced more.
The book gives these two factors, increase in shielding and in atomic radius, as the reason ionization energy falls down group 1. They outweigh the extra nuclear charge, so less energy is needed to remove the outer electron of potassium.

Outer electron (accent) of Na and K. The line shows the larger distance from the nucleus in potassium.

\text{IE}(\ce{K})<\text{IE}(\ce{Na})

Ionization energy: The minimum amount of energy required to remove the outermost electron from an isolated gas atom.

As you move down group 17, the electron affinity decreases. This is due to the same factors described above. They reduce the attraction between the nucleus and the incoming electron.

Although fluorine is the most electronegative element, its electron affinity is lower than that of chlorine. This is because the atomic size of F is much smaller than that of Cl. The incoming electron is repelled by the electrons of the F atom.

10.6.10Identification of an unknown elementOften examined

The electronic configuration, the position in the periodic table and the properties of an element can be used together to identify an unknown element. The following examples show how.

Example 1: Unknown element with atomic number 19

An unknown element has atomic number 19. It is a highly reactive metal that reacts violently with water to produce hydrogen gas. Identify its group, period and identity.

The atomic number 19 gives the electron configuration.
1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^1
The valence shell is 4s1. The n value is 4, so the element is in period 4. It has one valence electron, so it is in group 1.
Its reaction with water confirms the position. Alkali metals react violently with water, so this element is an alkali metal.
The element is potassium (K).

\text{Potassium (K): period 4, group 1 (alkali metal)}

The next unknown element has atomic number 17. It exists as diatomic molecules in its elemental form and reacts with metals to form a salt.

Example 2: Unknown element with atomic number 17

An unknown element has atomic number 17. It exists as diatomic molecules and reacts with metals to form salts. Identify its group, period and identity.

The atomic number 17 gives the electron configuration.
1s^2\,2s^2\,2p^6\,3s^2\,3p^5
The valence shell is 3s2 3p5, so it has 7 valence electrons and is in group 17. The n value is 3, so it is in period 3.
A group 17 element that exists as diatomic molecules and reacts with alkali metals to form salts such as \ce{NaCl} is chlorine (Cl).
\text{Cl}_2 \text{ and } \text{NaCl}

\text{Chlorine (Cl): period 3, group 17 (halogen)}

Example 10.5: Identifying the position of an unknown element

An unknown element has atomic number 35. Determine its position in the periodic table and predict its properties.

Write the electronic configuration.
1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^{10}\,4p^5
The valence shell configuration is 4s^2\,4p^5, which shows 7 valence electrons. The element is in group VIIA or 17.
The n value of the valence shell is 4, so the element is in period 4. (The element with atomic number 35 is bromine, Br.)

From its position, its properties can be predicted:

It is a non-metal.
It must be a diatomic molecule.
It forms an acidic oxide, which is more acidic than that of chlorine.
Its maximum oxidation state would be 7.
It combines with metals to form salts.
It forms an acid with hydrogen, and so on.

\text{Period 4, group 17 (VIIA): a non-metal}

Formula sheet

Group number from valence electrons (groups 1 and 2)

G=N_\text{val}

N_\text{val} is the number of valence electrons. For the s-block (groups 1 and 2) the group number equals the number of valence electrons.

no unit

Group number for groups 13 to 18

G=N_\text{val}+10

N_\text{val} is the number of valence electrons. For the p-block (groups 13 to 18) the group number is the valence electron count plus 10.

no unit

Period number

P=n

n is the n value of the valence shell (the largest principal energy level occupied). It gives the period, the horizontal row.

no unit

Electronegativity difference

\Delta EN = EN_{A}-EN_{B}

EN_{A} and EN_{B} are the electronegativities of the two bonded atoms, with the difference taken as positive. \Delta EN above 1.8 gives an ionic bond, 0.4 to 1.8 a polar covalent bond and below 0.4 a covalent bond.

no unit

Key terms

Periodic law: When elements are arranged in order of increasing atomic number, their properties repeat periodically.
Periodic table: A table that shows the systematic arrangement of the elements, based on the periodic law.
Period: A horizontal row of elements in the periodic table. The period number is the n value of the valence shell.
Group (family): A vertical column of elements in the periodic table. Elements in a group have the same number of valence electrons and similar chemical properties.
Alkali metals: The metallic elements of group 1.
Alkaline earth metals: The elements of group 2.
Boron family: The elements of group 13 (IIIA).
Oxygen family (chalcogens): The elements of group 16 (VIA).
Halogens: The elements of group 17 (VIIA).
Noble gases: The elements of group 18 (VIIIA), which do not readily undergo chemical reactions.
Normal (representative) elements: Group A elements, also called main group elements.
Transition elements: Group B elements, in groups 3 to 12.
Block: The part of the periodic table (s, p, d or f) that corresponds to the sub-shell holding the valence electrons.
Metalloids: Elements next to the leader line between metals and non-metals, with properties of both. Also called semi-metals.
Shielding effect: The reduction in the force of attraction between the nucleus and the valence electrons, caused by the electrons in the inner sub-shells.
Ionization energy: The minimum amount of energy required to remove the outermost electron from an isolated gas atom.
Electron affinity: The amount of energy released when an electron is added to the valence shell of an isolated gaseous atom to form a uni-negative gaseous ion.
Electronegativity: The ability of an atom to attract the electrons towards itself in a chemical bond.
Electronegativity difference (ΔEN): The difference in electronegativity between two bonded atoms, used as a rough indication of the bond type.
Amphoteric oxide: An oxide that reacts with both acids and bases, such as aluminium oxide.

Multiple choice questionsp. 226–227

The book prints no answer key for these multiple choice questions; each key was worked out from the text and Figure 10.7.

In the periodic table, the number of the period of an element is the same as:

Aatomic number
Batomic mass
Cgroup number
Dprincipal quantum numberAnswer

Why

The period number is the principal quantum number n of the outer (valence) shell, so the answer is (d). The atomic number counts protons, the atomic mass is the average mass of the atoms and the group number counts the valence electrons, so (a), (b) and (c) describe other properties.

Which of the following possesses the highest electronegativity?

A\ce{F}Answer
B\ce{Cl}
C\ce{Br}
D\ce{I}

Why

Electronegativity falls down group 17 because each atom is larger and its bonding electrons are held less strongly. Fluorine, at the top of the group, is therefore the most electronegative. The book gives 3.98 for F and 3.16 for Cl in its bond calculations, and the elements below chlorine are lower still.

iii

Which of the following will produce a base in water?

A\ce{Na2O}Answer
B\ce{CO2}
C\ce{SO3}
D\ce{SiO2}

Why

Sodium oxide is a metal oxide. It dissolves in water to give sodium hydroxide, a base: \ce{Na2O + H2O -> 2NaOH}. \ce{CO2} and \ce{SO3} are non-metal oxides and give the acids \ce{H2CO3} and \ce{H2SO4}. \ce{SiO2} is only weakly acidic (Table 10.7 of the book), so it does not give a base.

Which is correct for magnesium, \ce{Mg} (atomic number 12)?

Aperiod 1 element
Bgroup 1 element
Cperiod 2 element
Dalkaline earth metalAnswer

Why

Magnesium has the configuration 1s^2\,2s^2\,2p^6\,3s^2. Its highest shell is n=3, so it is in period 3, and its two valence electrons put it in group 2. Group 2 metals are the alkaline earth metals, so (d) is correct. Options (a) and (c) give the wrong period and option (b) the wrong group.

Note. The book prints the atomic number as a subscript, _{12}Mg; it is read here as atomic number 12.

v2025 Annual I, Set B, MCQ 8

An element has electron configuration 1s^2\,2s^2\,2p^6\,3s^2\,3p^2. It belongs to:

Aperiod 2
Bgroup 14Answer
Cgroup IIIA
Dperiod 1

Why

The highest shell is n=3, so the element is in period 3 and options (a) and (d) are wrong. The valence electrons are 3s^2\,3p^2, which is 4 electrons. For a p-block element the group number is the valence electrons plus 10, so 4+10=14: group 14 (IVA). Group IIIA would need 3 valence electrons.

Which of the following elements has the highest ionization energy?

A\ce{Li}Answer
B\ce{Na}
C\ce{K}
D\ce{Rb}

Why

Ionization energy falls down a group (section 10.6.9 of the book). Lithium is at the top of group 1, so its atoms are the smallest, its outer electron is closest to the nucleus and the inner shells shield it least. Na, K and Rb are lower in the group, with larger atoms and more shielding, so each needs less energy to remove the outer electron.

vii

Which of the following elements has the lowest electron affinity?

AF
BCl
CBr
DIAnswer

Why

Electron affinity is the energy released when a gaseous atom gains an electron. The values in Figure 10.7 are -328 for F, -349 for Cl, -325 for Br and -295 for I. Iodine releases the least energy, so it has the lowest electron affinity. Chlorine releases the most. Fluorine releases less than chlorine because its small size makes the added electron repel the electrons already in its 2p sub-shell (section 10.6.9).

Note. The key is read as the least energy released. If the question meant the most negative value in Figure 10.7, the answer would be (b) Cl (-349).

viii

Which of the following oxides is amphoteric?

A\ce{MgO}
B\ce{CO2}
C\ce{Al2O3}Answer
D\ce{SO3}

Why

Aluminium oxide reacts with both acids and bases, which is what makes it amphoteric (section 10.6.5 of the book): \ce{Al2O3 + 6HCl -> 2AlCl3 + 3H2O} and \ce{Al2O3 + 2NaOH + 3H2O -> 2NaAl(OH)4}. \ce{MgO} is basic, while \ce{CO2} and \ce{SO3} are acidic, so they do not behave amphoterically.

Elements of which group have strong non-metallic character?

AG-IA
BG-IIIA
CG-IVA
DG-VIIAAnswer

Why

The halogens of group VIIA (group 17) are strong non-metals (sections 10.6.2 and 10.6.3). Group IA contains the strongly metallic alkali metals. Group IIIA has mixed character, and in group IVA non-metallic character falls down the group, with carbon a non-metal and tin and lead metals. Hence G-VIIA is the answer.

Which atom has the smaller ionization energy?

A\ce{B}
B\ce{Al}Answer
C\ce{C}
D\ce{N}

Why

Boron and aluminium are both in group 13, and ionization energy falls down a group, so \ce{Al}, which is further down, has the smaller value. Carbon and nitrogen lie to the right of boron in period 2, and ionization energy rises across a period, so both are higher than boron.

Short questionsp. 227

i3 marks

What is a period?

Solution

A period is a horizontal row of the periodic table. Elements are placed in ascending order of atomic number from left to right, so their properties change gradually along the row and then repeat in the next period.

The period number is the principal energy level n of the valence shell. An element whose valence electrons are in the n=3 shell is in period 3. The number of elements per period varies from 2 in period 1 to 32 in period 7; periods 1 to 3 are the short periods and the rest are long periods.

ii3 marks

Differentiate between s and p-block elements.

Solution

The block of an element is set by the sub-shell that its valence electrons fill. Elements of the s-block have their valence electrons in s sub-shells, and those of the p-block have their valence electrons in p sub-shells.

Feature	s-block	p-block
Groups	Groups 1 and 2 (except He)	Groups 13 to 18
Valence electrons	In s sub-shells	In p sub-shells
Elements	Alkali and alkaline earth metals	Metals, non-metals and metalloids
Examples	\ce{Na}, \ce{Mg}	\ce{Cl}, \ce{Ar}

In short, the s-block elements are the alkali and alkaline earth metals, while the p-block mixes metals, metalloids and non-metals.

Note. The text places He outside the s-block, but Figure 10.2 draws the s-block bracket beside He in the first row. The answer follows the text.

iii3 marks

Find out the position of sulphur (atomic number 16) in the periodic table.

Solution

The electron configuration of sulphur is:

1s^2\,2s^2\,2p^6\,3s^2\,3p^4

The highest shell is n=3, so sulphur is in period 3. Its valence electrons are 3s^2\,3p^4, which is 6 electrons. Sulphur is a p-block element, so its group number is the valence electrons plus 10: 6+10=16. It is therefore in group 16 (VIA), the oxygen family or chalcogens.

\text{Period } 3,\ \text{group } 16\ (\text{VIA})

iv3 marks

Find out the position of an element in the periodic table which has electronic configuration 1s^2,2s^2,2p^6,3s^2,3p^3.

Solution

The total number of electrons is 2+2+6+2+3=15, so the element is phosphorus, \ce{P} (atomic number 15). Its valence shell is 3s^2\,3p^3, with n=3, so it is in period 3.

The valence shell has 5 electrons (2+3). For a p-block element the group number is 5+10=15, so the element is in group 15 (VA). The last electron enters a p sub-shell, so it is a p-block element.

\text{Period } 3,\ \text{group } 15\ (\text{VA}),\ \text{p-block}

v3 marks

In what respect do metals differ from metalloids?

Solution

Metals lie to the left of the blue leader line in the periodic table. Metalloids are the elements next to that line, which is shown in blue (section 10.4 of the book).

Property	Metals	Metalloids
Position	Left of the blue leader line	Along the blue leader line
Electrical conduction	Good conductors	Show properties of both metals and non-metals
Malleability and ductility	Malleable and ductile	Show properties of both metals and non-metals

The main difference is that metals conduct electricity well and can be hammered into sheets or drawn into wires, whereas metalloids have a mixture of metallic and non-metallic properties. This is why metalloids are also called semi-metals.

vi3 marks

A cation is smaller than its atom. Justify.

Solution

A cation forms when an atom loses one or more valence electrons. The number of protons, and so the nuclear charge, stays the same, but the number of electrons falls. Take sodium as an example:

\text{Na}:\ 1s^2\,2s^2\,2p^6\,3s^1 \qquad \text{Na}^+:\ 1s^2\,2s^2\,2p^6

The outer 3s electron is removed, so the n=3 shell is lost and the ion has the size of the second shell only. The fewer electrons that remain also repel each other less, so the nucleus pulls the remaining electrons closer. Both effects make the cation smaller than the atom.

Long questionsp. 227–228

Describe the variation of electron affinity in the periodic table.

Solution

Definition

Electron affinity is the energy released when an electron is added to the valence shell of an isolated gaseous atom to form a uni-negative gaseous ion. It is also called the first electron affinity.

\ce{X_{(g)} + e^- -> X^-_{(g)}}\qquad \text{energy released} = \text{electron affinity}

Variation across a period

Electron affinity generally increases from left to right across a period, so more energy is released when the atoms of later elements gain an electron. In period 2 of Figure 10.7 the values run from -60 for Li to -328 for F. Two factors cause this: the nuclear charge increases, and the atomic size decreases while the shielding effect stays about the same, so the added electron is pulled more tightly.

The values in Figure 10.7 have exceptions to the general rise. Beryllium (0) and nitrogen (+7) release no energy, and the noble gas neon (0) does not accept an electron readily. The full 2s^2 sub-shell of Be and the half-filled 2p^3 sub-shell of N are stable, so the added electron is not favoured.

Variation down a group

Electron affinity generally decreases from top to bottom in a group. The atoms get larger and the inner shells shield the outer electrons more, so the added electron binds less tightly and less energy is released. In group 1 the values in Figure 10.7 fall in size from -60 (Li) to -45 (Cs).

Chlorine has a higher electron affinity than fluorine (-349 against -328), even though fluorine is the most electronegative element. Fluorine's atoms are so small that the incoming electron is repelled by the electrons already in its 2p sub-shell.

Electron affinity trends from the values in Figure 10.7: across period 2 (top) and down group 17 (bottom).

Direction	Change in electron affinity	Main reason
Across a period (left to right)	Increases (more energy released)	Nuclear charge increases and atomic size decreases; shielding stays about the same
Down a group (top to bottom)	Decreases (less energy released)	Atomic size increases and shielding increases, so the added electron binds less tightly

Explain the reactions of sodium and magnesium with: (i) oxygen, (ii) chlorine, (iii) water.

Solution

(i) Reactions with oxygen

Sodium burns in a limited supply of oxygen to form white sodium oxide. In excess oxygen it forms pale yellow sodium peroxide. Magnesium burns in oxygen with an intense white flame to form magnesium oxide. In each case the metal loses its valence electrons to oxygen.

\ce{4Na + O2 -> 2Na2O}\qquad \ce{2Na + O2 -> Na2O2}\ (\text{excess oxygen})

\ce{2Mg + O2 -> 2MgO}

(ii) Reactions with chlorine

Sodium burns with a bright yellow flame to form white sodium chloride. Magnesium burns with an intense white flame to form white magnesium chloride.

\ce{2Na + Cl2 -> 2NaCl}\qquad \ce{Mg + Cl2 -> MgCl2}

(iii) Reactions with water

Sodium reacts violently with cold water, with a light explosion, to form sodium hydroxide and hydrogen gas.

\ce{2Na + 2H2O -> 2NaOH + H2}

Magnesium reacts only slowly with cold water. A thin layer of magnesium hydroxide forms on the metal and protects it from further reaction. In steam, magnesium burns with a white flame to form magnesium oxide and hydrogen.

\begin{gathered}\ce{Mg + 2H2O -> Mg(OH)2 + H2}\ (\text{cold water})\\ \ce{Mg + H2O -> MgO + H2}\ (\text{steam})\end{gathered}

Note. The book prints \ce{2Na + O2 -> Na2O}, which is not balanced; the balanced equation is \ce{4Na + O2 -> 2Na2O}. It also prints \ce{Mg + 2H2O -> MgO + H2} for steam, which is not balanced (oxygen and hydrogen do not match). The balanced steam equation is \ce{Mg + H2O -> MgO + H2}.

Discuss the variation in the nature of oxides of period 3 elements.

Solution

Trend across the period

The oxides of the period 3 elements change from strongly basic (\ce{Na2O}) to very strongly acidic (\ce{Cl2O7}). Metal oxides dissolve to give bases, non-metal oxides give acids, and the oxide of aluminium is amphoteric.

Oxide	Nature (Table 10.7)	Reaction shown or stated in the book
\ce{Na2O}	Strongly basic	\ce{Na2O + 2HCl -> 2NaCl + H2O}; \ce{Na2O + H2O -> 2NaOH}
\ce{MgO}	Basic	\ce{MgO + H2SO4 -> MgSO4 + H2O}; \ce{MgO + H2O -> Mg(OH)2}
\ce{Al2O3}	Amphoteric	\ce{Al2O3 + 6HCl -> 2AlCl3 + 3H2O}; \ce{Al2O3 + 2NaOH + 3H2O -> 2NaAl(OH)4}
\ce{SiO2}	Weakly acidic	No reaction equation is given in the book
\ce{P4O10}	Acidic	\ce{P4O10 + 6H2O -> 4H3PO4}
\ce{SO2}	Strongly acidic	Gives an acidic solution in water
\ce{Cl2O7}	Very strongly acidic	Gives an acidic solution in water

Nature of the period 3 oxides (Table 10.7), from most basic to most acidic.

Why the nature changes

Across the period, metallic character falls and the electronegativity difference between each element and oxygen falls, so the bonding changes from ionic (Na2O and MgO) to covalent (the oxides of P, S and Cl). Ionic oxides of metals give basic solutions. Covalent oxides of non-metals give acidic solutions.

Aluminium lies at the boundary between the two groups, which is why its oxide reacts with both acids and bases.

Note. The book prints \ce{Na2O + H2O -> NaOH}, which is not balanced (the correct equation gives 2\ce{NaOH}); it prints 2\ce{MaAl(OH)4} for sodium aluminate, which should read 2\ce{NaAl(OH)4}; and it prints \ce{Al2O3 + 6HCl -> 3AlCl3 + 3H2O}, which should read 2\ce{AlCl3}.

Explain the variation in metallic and non-metallic character in the periodic table.

Solution

Definitions

Metallic character is the tendency of an element to lose valence electrons and form cations. Non-metallic character is the tendency to gain or share electrons to complete the valence shell.

Across a period (left to right)

The number of valence electrons rises from 1 to 8 across periods 2 and 3. Metals have 1 to 3 valence electrons, low ionization energy and low electronegativity, so they lose electrons easily. Non-metals have 5 to 7 valence electrons, high ionization energy and high electronegativity, so they gain or share electrons. Metallic character therefore decreases, and non-metallic character increases, from left to right.

Down a group (top to bottom)

In group 1, metallic character increases down the group. The atoms get larger, so their valence electron is lost more easily. In group 17, non-metallic character decreases down the group, because the larger atoms attract electrons less strongly. In group 14, carbon is a non-metal, silicon and germanium are metalloids, and tin and lead are metals, so metallic character increases down the group.

Trend	Across a period (left to right)	Down a group (top to bottom)
Valence electrons	Increase from 1 to 8	Same within a group
Metallic character	Decreases	Increases (group 1; group 14, from C to Pb)
Non-metallic character	Increases	Decreases (group 17, from F to I)
Main reason	Nuclear charge rises, so the valence electrons are held more strongly	Atoms grow larger, so the valence electrons are held less strongly

The blue leader line separates the metals on the left from the non-metals on the right, and the metalloids lie along it. Metalloids show both kinds of character.

Evaluate on the following giving reasons: (i) atomic radii decrease in a period; (ii) atomic radii increase in a group.

Solution

(i) Atomic radii decrease in a period

The statement is true. Across a period the electrons are added to the same valence shell, so the atoms do not gain a new energy level that would make them larger.

The nuclear charge increases by one proton at each element. The shielding of the valence electrons by the inner shells stays about the same, because the added electrons enter the same shell. The valence electrons are therefore pulled more strongly towards the nucleus, and the atom becomes smaller (the book explains the same factors in section 10.5.1).

(ii) Atomic radii increase in a group

The statement is true. Going down a group, each element has its valence electrons in a higher energy level (a larger n), so the outer electrons lie further from the nucleus.

Each step down also adds inner shells, which increase the shielding of the outer electrons. The nucleus pulls the valence electrons less strongly, so the atom is larger. The book states that atomic size gradually increases from top to bottom in a group (section 10.3).

How does the periodic table help in understanding the properties of elements?

Solution

Arrangement

The periodic table lists the elements in order of atomic number. Elements with similar valence electron configurations fall in the same group, and because chemical properties depend on the valence electrons, the elements of one group behave similarly. The properties repeat from one period to the next, as the periodic law states.

Predicting properties from position

The period gives the principal energy level n of the valence shell.
The group gives the number of valence electrons: for groups 1 and 2 the group number is the number of valence electrons, and for groups 13 to 18 the valence electrons are the group number minus 10.
The block (s, p, d or f) shows which sub-shell is being filled.
The trends follow from position. Atomic size increases down a group and decreases across a period; ionization energy and electronegativity fall down a group; metallic character increases down a group and decreases across a period.

For example, potassium (group 1, period 4) is a highly reactive alkali metal that reacts violently with water. An unknown element in group 1 can therefore be predicted to behave in the same way.

Explaining chemical behaviour and bonding

The electronegativity difference between two bonded atoms predicts the bond type: a difference above 1.8 gives an ionic bond (as in NaCl, with a difference of 2.23), a difference from 0.4 to 1.8 gives a polar covalent bond (as in H-F), and a difference below 0.4 gives a covalent bond (as in \ce{C-H}).

The group number also gives the maximum oxidation number of an element in its oxides and chlorides, and the trends across period 3 show whether those oxides are basic or acidic. The table therefore lets you predict the properties of an element before you test it.

Evaluate the significance of the periodic table.

Solution

The periodic table is significant because it turns a long list of elements into an ordered system. Arranging the elements by atomic number, in periods and groups, shows that their properties repeat.

Strengths

It gives the position of an element from its electronic configuration: period from n, group from the valence electrons and block from the sub-shell.
It predicts properties from the group or period, for example the reaction of an alkali metal with water.
It separates metals, non-metals and metalloids, and links these to conduction and malleability.
It explains the trends in atomic size, ionization energy, electron affinity and electronegativity, and through these the type of bonding and the nature of oxides and chlorides.
It helps identify an unknown element from its atomic number and properties, as in the book's identification examples.

Limitations

The trends are general, not exact. Electron affinity does not rise steadily across period 2 (Be 0, N +7 and Ne 0 in Figure 10.7), and fluorine has a lower electron affinity than chlorine.
The table shows patterns, not every value. The exact values of properties such as electron affinity must still be measured.

Overall, the periodic table is highly significant because it organises the elements into a system that predicts their properties. Its predictions are trends, however, and should be checked against the data for the element concerned.

How does the periodic table aid in understanding the characteristics of oxides across the period?

Solution

Across period 3 the oxides change from strongly basic (\ce{Na2O}) to very strongly acidic (\ce{Cl2O7}). The position of each element in the table explains this in three steps.

Group number: the maximum oxidation number of the element in its oxide equals its number of valence electrons, so it is +1 for Na, +2 for Mg, +3 for Al, +4 for Si, +5 for P, +6 for S and +7 for Cl.
Bonding: metallic character decreases across the period and the electronegativity difference between the element and oxygen falls. For MgO the difference is 3.44-1.31=2.13, which is ionic; for the oxides of non-metals on the right the bond is covalent.
Acid or base: ionic oxides of metals give basic solutions, for example \ce{Na2O + H2O -> 2NaOH}; covalent oxides of non-metals give acidic solutions, for example \ce{P4O10 + 6H2O -> 4H3PO4}.

Oxide	Group of the element	Nature (Table 10.7)
\ce{Na2O}	Group 1 (IA)	Strongly basic
\ce{MgO}	Group 2 (IIA)	Basic
\ce{Al2O3}	Group 13 (IIIA)	Amphoteric
\ce{SiO2}	Group 14 (IVA)	Weakly acidic
\ce{P4O10}	Group 15 (VA)	Acidic
\ce{SO2}	Group 16 (VIA)	Strongly acidic
\ce{Cl2O7}	Group 17 (VIIA)	Very strongly acidic

Aluminium oxide lies between the two groups and is amphoteric, so it reacts with both acids and bases. The table therefore lets you predict the nature of an oxide from the element's position: basic on the left and acidic on the right.

FBISE 2025 and 2026 questions

Questions from the FBISE HSSC-I 2026 (1st Annual) and FBISE HSSC-I 2025 (1st Annual) papers that belong to this unit, answered the way an examiner expects.

2026 Annual I, Set B, MCQ 41 mark

Br has higher electron affinity than K due to:

AShielding effect
BAtomic size
CNuclear chargeAnswer
DDensity

Why

Bromine (Z=35) and potassium (Z=19) are in the same period (period 4). Going from K to Br the nuclear charge rises from +19 to +35, while the extra electrons go into the same shell. The attraction between the nucleus and an incoming electron therefore becomes much stronger, so more energy is released when Br gains an electron. In the book's Figure 10.7 the values are K -48 and Br -325. The book gives nuclear charge as the main reason why electron affinity increases from left to right.

Shielding (option A) is wrong because the extra electrons are added to the same shell, so the shielding by inner shells stays about the same across the period and cannot explain the large difference. Atomic size (option B) is the tempting choice: Br is smaller than K, but the smaller size is itself a result of the greater nuclear charge, which is the main cause. Density (option D) plays no part in electron affinity.

2025 Annual I, Set B, MCQ 81 mark

Position of an element in periodic table with electronic configuration 1s^{2}, 2s^{2}2p^{6}, 3s^{2}3p^{2} is:

AIn Group-1 (I-A)
BIn Group-2 (II-A)
CIn Group-13 (III-A)
DIn Group-14 (IV-A)Answer

Why

The highest shell is n=3, so the element is in period 3. The valence shell is 3s^2\,3p^2, which holds 2+2=4 electrons. The last electron enters a p sub-shell, so it is a p-block element and its group number is the number of valence electrons plus 10, that is 4+10=14 (group IVA). The element is silicon.

Group 13 (option C) is the tempting wrong choice: it would need three valence electrons (3s^2\,3p^1). Groups 1 and 2 (options A and B) are s-block groups and would have only s electrons in the valence shell.

2026 Annual I, Set B, Q.2 (i)1x3 = 3 marks

Write three equations showing reaction of chlorides with water.

Solution

Metal chlorides such as \ce{NaCl} and \ce{MgCl2} are ionic and simply dissolve in water to give neutral solutions (pH 7). The covalent chlorides react with water (hydrolysis) and form hydrogen chloride, which makes the solution acidic. Three such reactions (1 mark each):

Aluminium chloride (1 mark)

\ce{AlCl3 + 3H2O -> Al(OH)3 + 3HCl}

Silicon tetrachloride (1 mark)

\ce{SiCl4 + 4H2O -> Si(OH)4 + 4HCl}

Phosphorus pentachloride (1 mark)

\ce{PCl5 + 4H2O -> H3PO4 + 5HCl}

In each reaction the acidity comes from the \ce{HCl} formed. The solutions have pH about 3 for \ce{AlCl3} and about 0 for \ce{SiCl4} and \ce{PCl5}.

2026 Annual I, Set B, Q.2 (xii)3 marks

What is meant by electronegativity? Explain its trend along the period of periodic table.

Solution

Definition (1 mark)

Electronegativity is the ability of an atom to attract the electrons towards itself in a chemical bond.

Trend along a period (2 marks)

Electronegativity increases from left to right across a period. Moving along the period the nuclear charge increases, the electrons are added to the same valence shell, so the shielding stays about the same and the atomic size decreases. The nucleus therefore pulls the shared pair of electrons more strongly. Group 1 elements have the lowest values and the halogens the highest.

Period 3	Na	Mg	Cl
Electronegativity	0.93	1.31	3.16

Period 2	C	O	F
Electronegativity	2.55	3.44	3.98

2026 Annual I, Set D, MCQ 41 mark

Which element shows highest electron affinity?

ANa
BMg
CAl
DClAnswer

Why

Electron affinity generally increases from left to right along a period because the nuclear charge increases and the atomic size decreases. Chlorine, a halogen at the right of period 3, has the largest value, -349 in the book's Figure 10.7, against -53 for Na and -44 for Al.

Magnesium (option B) has a full 3s^2 sub-shell, and its electron affinity is 0 in the book's table, the lowest of the four. Sodium and aluminium have a much smaller nuclear charge than chlorine, so they release far less energy when they gain an electron.

2025 Annual I, Set D, MCQ 91 mark

Which of the following decreases across the period?

AAtomic radiusAnswer
BIonization potential
CElectron affinity
DElectronegativity

Why

Across a period the nuclear charge increases while the electrons are added to the same valence shell. The nucleus pulls the electrons in more strongly, so the atomic radius decreases from left to right.

The other three properties all increase across the period: ionization energy increases because the outer electron is held more tightly, electron affinity generally increases (the book gives -53 for Na and -349 for Cl) and electronegativity increases. Electron affinity (option C) is the tempting choice because its values are printed with negative signs, but a more negative value means more energy released, so it increases across the period.

2025 Annual I, Set D, MCQ 101 mark

An element M reacts with \ce{O2}, \ce{H2O} and \ce{Cl2} to form the products \ce{M2O}, \ce{MOH} and \ce{MCl} respectively. The element is most likely:

ALiAnswer
BAl
CMg
DCa

Why

The formulas \ce{M2O}, \ce{MOH} and \ce{MCl} show that M forms the \ce{M+} ion, so it has one valence electron and belongs to group 1. Of the elements listed only lithium is in group 1. Like the book's example of sodium, a group 1 metal burns in oxygen to give an oxide (\ce{M2O}), reacts with water to give a hydroxide (\ce{MOH}) and burns in chlorine to give a chloride (\ce{MCl}).

Mg and Ca (options C and D) are in group 2 and form \ce{MO}, \ce{M(OH)2} and \ce{MCl2} (compare \ce{MgO}, \ce{Mg(OH)2} and \ce{MgCl2} in the book). Al (option B) is in group 13 and forms \ce{Al2O3}, \ce{Al(OH)3} and \ce{AlCl3}.

2026 Annual I, Set D, Q.2 (i)1x3 = 3 marks

Briefly explain effects of hydrolysis of any three chlorides on pH with equations.

Solution

Ionic chlorides do not change the pH of water, but covalent chlorides react with water and give hydrochloric acid, which lowers the pH. Three examples (1 mark each):

Chloride	Reaction with water	Effect on pH
\ce{NaCl} (ionic)	Dissolves without reaction: \ce{NaCl -> Na+ + Cl-}	Neutral, pH 7
\ce{AlCl3} (covalent)	\ce{AlCl3 + 3H2O -> Al(OH)3 + 3HCl}	Acidic, pH about 3
\ce{PCl5} (covalent)	\ce{PCl5 + 4H2O -> H3PO4 + 5HCl}	Strongly acidic, pH about 0

The acidity of the covalent chlorides is due to the \ce{HCl} formed (and also \ce{H3PO4} in the case of \ce{PCl5}). \ce{SiCl4} also hydrolyses completely (\ce{SiCl4 + 4H2O -> Si(OH)4 + 4HCl}, pH about 0), and \ce{MgCl2} is neutral like \ce{NaCl}.

2025 Annual I, Set D, Q.2 (iii) OR3 marks

How are the elements arranged in different blocks of periodic table?

Solution

Meaning of a block (1 mark)

The elements are arranged into blocks according to the type of sub-shell (s, p, d or f) that holds their valence electrons. The periodic table therefore has four blocks.

Schematic of the four blocks of the periodic table (after Figure 10.2).

The four blocks (2 marks)

Block	Groups	Valence electrons in	Description
s-block	1 and 2	s sub-shell	Alkali and alkaline earth metals
p-block	13 to 18	p sub-shell	Both metals and non-metals
d-block	3 to 12 (sub-group B)	d sub-shell	Transition metals
f-block	Two rows at the bottom	f sub-shell	Lanthanides and actinides (inner transition metals)

2025 Annual I, Set D, Q.2 (xi)3 marks

Differentiate between metallic and non-metallic behavior of elements.

Solution

Metallic behaviour is the tendency of an atom to lose its valence electrons and form cations. Non-metallic behaviour is the tendency to gain or share valence electrons to reach a stable electron configuration.

Property	Metallic behaviour	Non-metallic behaviour
Position in the periodic table	Left side	Right side
Valence electrons	Usually 1 to 3	Usually 5 to 7
Ionization energy	Low	High
Electronegativity	Low	High
Tendency	Lose electrons easily and form cations	Gain electrons (form anions) or share them to form covalent bonds
Nature of oxides	Basic (metal oxides give bases in water)	Acidic (non-metal oxides give acids in water)
Trend in the table	Decreases along a period; increases down a group (group 1, group 14)	Increases along a period; decreases down a group (group 17)

For example, sodium loses one electron to form \ce{Na+} (metallic), whereas chlorine gains one electron to form \ce{Cl-} (non-metallic).

2026 Annual I, Set D, Q.2 (xii)3 marks

What is meant by electronegativity? Explain its trend along the period of periodic table.

Solution

Definition (1 mark)

Electronegativity is the ability of an atom to attract the electrons towards itself in a chemical bond.

Trend along a period (2 marks)

Period 3	Na	Mg	Cl
Electronegativity	0.93	1.31	3.16

Period 2	C	O	F
Electronegativity	2.55	3.44	3.98

Exam focus

What is electronegativity? Explain its trend along a period. (short)
Asked in both 2026 sets (Set B Q.2 (xii) and Set D Q.2 (xii)), 3 marks each; give the definition and the reason (nuclear charge up, size down).
Go to answer
Write equations for the reactions of chlorides with water and state the effect on pH (AlCl3, SiCl4, PCl5, and NaCl as neutral). (short)
Asked in both 2026 sets (Set B Q.2 (i) equations; Set D Q.2 (i) effect on pH).
Go to answer
Trend of electron affinity across a period and down a group, and why Cl is higher than F. (long)
2026 Set B MCQ 4 (Br against K) and Set D MCQ 4 (highest electron affinity: Cl), and the book's long question 1.
Go to answer
Find the period, group and block of an element from its electron configuration. (mcq)
2025 Set B MCQ 8 (3s^2\,3p^2 gives group 14), the same as book exercise MCQ v; also book short questions 3 and 4.
Go to answer
Differentiate between metallic and non-metallic behaviour, and explain the trends in the table. (short)
2025 Set D Q.2 (xi); the book's long question 4 asks the trend across a period and down a group.
Go to answer
Properties that change across a period (atomic radius, ionization energy, electron affinity, electronegativity) and identifying a group from the formulas of its compounds. (mcq)
2025 Set D MCQ 9 (atomic radius decreases) and MCQ 10 (M2O, MOH, MCl means group 1).
Go to answer
How are the elements arranged in the s, p, d and f blocks? (short)
2025 Set D Q.2 (iii) OR; book short question 2 differentiates s and p blocks.
Go to answer
Reactions of sodium and magnesium with oxygen, chlorine and water. (long)
Book long question 2; not asked directly in the 2025 or 2026 papers, but 2025 Set D MCQ 10 uses these reactions.
Go to answer
Nature of the period 3 oxides (strongly basic to strongly acidic) with equations. (long)
Book long questions 3 and 8; this topic was not in the 2025 or 2026 paper.
Go to answer