Chapter 8: Periodic Table and PeriodicityLong Questions with Answers

Information needed: electronic configuration

The position of an element depends on its electronic configuration, which comes from its atomic number. If we know the configuration (or the number of shells and of outer electrons), we can place it:

Chemical properties help too

The charge on its ion, the nature of its oxide (basic or acidic), whether it is a metal or non-metal and its reactivity with water all point to its group.

Is atomic mass helpful?

Not reliably. Mendeleev used atomic mass, but in some pairs the heavier element must come first because of its properties:

Isotopes also have different masses but the same position. So the atomic number, not the mass, decides the position; atomic mass gives only a rough idea.

Four blocks

Elements in the periodic table are classified into blocks depending upon the type of the subshell which is being filled: s, p, d and f.

Yes, they are helpful

So knowing the block tells us at once the general behaviour of an element.

(a) Atomic radius

Atomic radius is half the distance between the nuclei of two identical bonded atoms. In a period it decreases from left to right.

Reason: as we go from lithium to neon, electrons are added to the same outermost shell, while the charge on the nucleus increases from +3 to +10. This pulls the electrons closer to the nucleus, so the atoms get smaller.

(b) Ionisation energy

Ionisation energy is the energy required to remove the most loosely bound electron from an isolated gaseous atom. In a period it increases from left to right.

Reason: the atoms get smaller and the nuclear charge larger, so the attraction between the nucleus and the outer electrons is stronger and more energy is needed to remove one.

Exceptions

B is lower than Be (a new 2p subshell starts) and O lower than N (O has its first paired 2p electron), but the overall trend is upward.

1. Ionisation energy

Elements with low ionisation energy (metals of groups 1 and 2) lose electrons easily and form positive ions: they form ionic bonds with non-metals, and metallic bonds with each other.

2. Electron affinity

Elements with high electron affinity (halogens, oxygen) gain electrons easily and form negative ions with metals.

3. Electronegativity

The difference in electronegativity between two atoms decides the bond:

4. Atomic size

Large atoms on the left lose electrons easily (ionic, metallic); small atoms on the right attract electrons (covalent or anions).

5. Metallic character

Metal + metal → metallic bond; metal + non-metal → ionic; non-metal + non-metal → covalent.

Boron and silicon are also solids, but they are metalloids. Bromine is the only non-metal that is a liquid at room temperature; the rest (H₂, N₂, O₂, F₂, Cl₂, noble gases) are gases.

A period fills the outermost s and p

Each period represents the completion of a shell, and the number of elements in a period is fixed by the number of electrons accommodated. The elements of a period run from ns¹ to ns²np⁶.

Periods 2 and 3

Period 2 fills 2s and 2p: 2 + 6 = 8. Period 3 fills 3s and 3p: 2 + 6 = 8. The third shell also has a 3d subshell, but 3d is higher in energy than 4s, so it is filled only in period 4 (from scandium). So periods 2 and 3 both have 8 elements.

Other periods

Period 1 has only the 1s subshell (2 elements); periods 4 and 5 add ten d-block elements (18); periods 6 and 7 also add fourteen f-block elements (32).

Benefits

So the periodic table is the most useful single summary of chemistry.

1. Very small size

Li and Be are the first members of their groups, with only two shells; they are the smallest atoms of their groups (Li 152 pm, Be 113 pm).

2. High charge density of the ions

Li⁺ and especially Be²⁺ are tiny with a concentrated charge, so they pull the electron clouds of neighbouring anions and give their compounds covalent character (BeCl₂ is largely covalent).

3. High ionisation energy and electronegativity

They hold their electrons more strongly (Li 520, Be 899 kJ mol⁻¹), so they are less electropositive and less reactive than the members below them.

4. No d-orbitals in the outer shell

Their second shell has only s and p subshells, so Be can hold at most four bonds.

Examples of different behaviour

Mendeleev's merits

He grouped elements with similar properties, left gaps and predicted the properties of undiscovered elements (gallium, scandium, germanium), and corrected some atomic masses.

Why the modern table is better

Using atomic number removes the anomalies, explains isotopes, and links the position of every element directly to its electronic configuration.

Basis

The modern periodic table is based upon the arrangement of elements according to increasing atomic number. Properties of elements repeat after regular intervals, so elements of similar properties are placed in the same vertical columns.

Periods

There are seven periods. Each period except the first starts with an alkali metal and ends at a noble gas. Each period represents the completion of a shell; the number of elements in a period is fixed.

Lanthanides and actinides

In the 6th and 7th periods two series of fourteen elements each are placed at the bottom: lanthanides (after La = 57) and actinides (after Ac = 89). Together they are called rare earths or f-block elements.

Groups

There are eighteen groups. Elements in a group resemble one another in chemical properties since they contain the same number of electrons in their outermost shells. A group is also called a family.

Normal and transition elements

Groups 1, 2 and 13 to 17 contain normal elements: inner shells complete, outermost shell incomplete. Groups 3 to 12 are transition elements: inner sub-shells are being completed.

Blocks

s-block (groups 1, 2), p-block (13-18), d-block (between s and p) and f-block (two rows at the bottom), according to the subshell being filled.

Reason

The chemical properties of elements depend largely upon the number of electrons in their outermost shells; all elements of a group have the same number, so they show similar properties.

Group 1 (ns¹)

Strong tendency to lose their one electron, forming cations: electropositive metals. They react vigorously with water producing hydrogen and alkalis. Their oxides are strongly basic; they form halides with halogens. Reactivity increases down the group.

Group 2 (ns²)

They lose both outer electrons, forming dipositive ions; the tendency increases down the group. Their oxides are basic and form alkalis in water.

Group 16 (ns²np⁴)

They accept two electrons to form dinegative ions; oxygen is strongly electronegative. O + 2e⁻ → O²⁻; they form oxides and sulphides.

Group 17 (ns²np⁵)

A strong tendency to gain one electron to become an anion; very reactive non-metals that form salts with metals (halogens). Reactivity decreases down the group.

Charge on ions

Ionisation energy

The energy required to remove the most loosely bound electron from the valence shell of an isolated gaseous atom. Na → Na⁺ + e⁻, +496 kJ mol⁻¹. It increases from left to right in a period and decreases from top to bottom in a group, because smaller atoms hold their electrons more strongly.

Electron affinity

The energy change when an electron is added to the outermost shell of an isolated gaseous atom. F + e⁻ → F⁻, −328 kJ mol⁻¹. Smaller atoms attract the entering electron more strongly, so it increases across a period and decreases down a group (Cl −349 is an exception, above F).

Electronegativity

The force with which an atom attracts the shared pair of electrons towards itself in a bond. It increases from left to right in a period and decreases down a group. Most electronegative atoms are at the top right (F, O, N, Cl).

Difference between EA and EN

Electron affinity is an energy change of a free gaseous atom gaining an electron; electronegativity is the pull of a bonded atom on a shared pair and has no unit.

Metallic character

The tendency of an element to lose electrons and form positive ions (cations).

In a group

Since ionisation energy decreases down a group, elements have an increased ability to lose electrons, so both metallic character and reactivity increase down the group. Group 2: Ba > Sr > Ca > Mg. Group 14 goes from non-metal carbon to metals tin and lead.

In a period

Moving left to right the nuclear charge increases and the valence electrons are pulled strongly, making it difficult to lose them, so metallic character decreases. Period 3: Na, Mg, Al are metals; Si a metalloid; P, S, Cl non-metals.

Reactivity in a period

Reactivity decreases from the very reactive sodium to magnesium, aluminium and silicon; moving further right towards the non-metals, reactivity increases again (chlorine is very reactive).

Density

Density is the mass of a substance in unit volume. It generally increases from top to bottom in a group but varies less across a period.

Rules

Example 1: sodium, Z = 11

1s² 2s² 2p⁶ 3s¹ → period 3, s-block, group 1.

Example 2: chlorine, Z = 17

1s² 2s² 2p⁶ 3s² 3p⁵ → period 3, p-block, group 10 + 7 = 17.

Example 3: calcium, Z = 20

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² → period 4, s-block, group 2.

Example 4: iron, Z = 26

… 3s² 3p⁶ 3d⁶ 4s² → period 4, d-block, group 6 + 2 = 8.

Example 5: neon, Z = 10

1s² 2s² 2p⁶ → period 2, p-block, group 18.