Periodic Properties

The Periodic Table

Define periodic properties (periodicity).

Regularly repeating patterns of atomic, physical and chemical properties that can be predicted using the Periodic Table and explained using electronic configurations.

Give examples of periodic properties.

Atomic radius, melting temperature, boiling temperature and first ionisation energy.

Define atomic radius.

A measure of the size of an atom.

Why is atomic radius difficult to measure directly?

The electron cloud does not have a well-defined boundary.

How can atomic radius be determined?

By measuring the distance between the nuclei of two atoms and dividing it by two.

Define covalent radius.

Half the distance between the nuclei of two covalently bonded atoms of the same element.

Define van der Waals radius.

Half the distance between the nuclei of two neighbouring non-bonded atoms.

Define metallic radius.

Half the distance between the nuclei of neighbouring metal atoms in a metallic lattice.

Which radius is normally used when comparing trends in atomic radius?

Covalent radius.

Why should the same type of atomic radius be used when comparing elements?

Different methods of measuring radius give different values for the same element.

What happens to atomic radius across a period?

It decreases.

Why does atomic radius decrease across a period?

Nuclear charge increases while electrons are added to the same quantum shell, so the attraction between the nucleus and outer electrons increases.

What happens to nuclear charge across a period?

It increases because the number of protons increases.

Why does increased electron-electron repulsion not prevent atomic radius decreasing across a period?

The increase in nuclear attraction outweighs the increase in electron-electron repulsion.

What happens to the number of occupied principal quantum shells across a period?

It stays the same.

What general relationship exists between structure and melting/boiling temperature?

Giant lattice structures generally have high melting and boiling temperatures, while simple molecular substances generally have much lower melting and boiling temperatures.

Which types of structure generally have high melting and boiling temperatures?

Giant metallic and giant covalent lattice structures.

Which type of structure generally has low melting and boiling temperatures?

Simple molecular structures.

Why do giant structures generally have high melting and boiling temperatures?

Strong bonds or attractions extend throughout the giant lattice and require a large amount of energy to overcome.

Why do simple molecular substances generally have low melting and boiling temperatures?

Only relatively weak intermolecular forces between molecules need to be overcome.

What type of bonding do lithium and beryllium have?

Metallic bonding.

What type of structure do lithium and beryllium have?

Giant metallic lattices.

What type of bonding do boron and carbon have?

Covalent bonding.

What type of structure do boron and carbon have?

Giant covalent lattices.

What type of structure do nitrogen, oxygen and fluorine have?

Simple molecular structures.

In what form does nitrogen exist?

N₂ molecules.

In what form does oxygen exist?

O₂ molecules.

In what form does fluorine exist?

F₂ molecules.

What type of bonding do sodium, magnesium and aluminium have?

Metallic bonding.

What type of structure do sodium, magnesium and aluminium have?

Giant metallic lattices.

What type of bonding and structure does silicon have?

Covalent bonding in a giant covalent lattice.

What type of structure do phosphorus, sulfur and chlorine have?

Simple molecular structures.

In what molecular form does phosphorus commonly exist?

P₄.

In what molecular form does sulfur commonly exist?

S₈.

In what molecular form does chlorine exist?

Cl₂.

Why does silicon have a high melting and boiling temperature?

It has a giant covalent lattice containing many strong covalent bonds that require a large amount of energy to break.

Why do phosphorus, sulfur and chlorine have much lower melting and boiling temperatures than silicon?

They are simple molecular substances, so only intermolecular forces between molecules need to be overcome.

What is the general trend in first ionisation energy across Period 2?

It generally increases from lithium to neon.

What is the general trend in first ionisation energy across Period 3?

It generally increases from sodium to argon.

Why does first ionisation energy generally increase across a period?

Nuclear charge increases while electrons are added to the same principal quantum shell, so attraction between the nucleus and outer electron increases.

Why is the increase in first ionisation energy across a period not completely regular?

There are anomalies caused by differences in sub-shell energy and electron-electron repulsion between paired electrons.

Why is helium's first ionisation energy higher than hydrogen's?

Helium has a greater nuclear charge, which more than offsets the increased electron-electron repulsion caused by its second 1s electron.

What are the electronic configurations of hydrogen and helium?

H: 1s¹; He: 1s².

What are the two main first-ionisation-energy anomalies in Period 2?

Beryllium to boron and nitrogen to oxygen.

What are the corresponding anomalies in Period 3?

Magnesium to aluminium and phosphorus to sulfur.

Why is the first ionisation energy of boron lower than that of beryllium?

Boron's outer electron is in a higher-energy 2p orbital, whereas beryllium's outer electron is in a lower-energy 2s orbital, so less energy is required to remove boron's electron.

What are the electronic configurations of beryllium and boron?

Be: 1s²2s²; B: 1s²2s²2p¹.

Why does boron's 2p electron experience greater shielding than beryllium's 2s electron?

Boron has two inner electron sub-shells whereas beryllium has only one inner sub-shell relative to the electron being removed.

Why is the first ionisation energy of aluminium lower than that of magnesium?

Aluminium loses a higher-energy 3p electron, whereas magnesium loses a lower-energy 3s electron, so aluminium's electron is easier to remove.

What causes the Group 2 → Group 3 anomaly in first ionisation energy?

The electron removed from the Group 3 element is in a higher-energy p orbital rather than an s orbital.

Why is the first ionisation energy of oxygen lower than that of nitrogen?

Oxygen has a pair of electrons in one 2p orbital, increasing electron-electron repulsion and making one electron easier to remove.

What are the electronic configurations of nitrogen and oxygen?

N: 1s²2s²2p³; O: 1s²2s²2p⁴.

How are the three 2p electrons arranged in nitrogen?

Each occupies a separate 2p orbital.

How are the four 2p electrons arranged in oxygen?

One 2p orbital contains a pair of electrons while the other two each contain one electron.

Why does pairing electrons in the same orbital make an electron easier to remove?

The paired electrons repel each other, increasing their energy and reducing the energy required to remove one.

Why is the first ionisation energy of sulfur lower than that of phosphorus?

Sulfur contains a pair of electrons in one 3p orbital, causing increased electron-electron repulsion and making one electron easier to remove.

What causes the Group 5 → Group 6 anomaly in first ionisation energy?

Electron pairing in a p orbital increases electron-electron repulsion, making an electron easier to remove.

Why is the explanation that oxygen has a lower first ionisation energy because nitrogen has a 'stable half-filled p sub-shell' not preferred?

The difference is explained by increased electron-electron repulsion between the paired electrons in oxygen.

Why is the explanation that boron has a lower first ionisation energy because losing an electron gives it a 'stable full 2s orbital' incorrect?

Energy must be supplied to remove the electron, so the resulting ion is at a higher energy than the original atom.

Why does neon have the highest first ionisation energy in Period 2?

It has the greatest nuclear charge in Period 2 while its outer electrons are still in the second principal quantum shell, giving a strong attraction between the nucleus and outer electrons.

Why does argon have the highest first ionisation energy in Period 3?

It has the greatest nuclear charge in Period 3 while its outer electrons are still in the third principal quantum shell.