Chapter 7: Periodicity

Question 1

Describe the arrangement of the periodic table in terms of atomic number, groups and periods. How does this correlate to an element’s physical and chemical properties?

Answer 1

• Reading from left to right, the elements are arranged in order of increasing atomic number. Each successive element has atoms with one extra proton.
• The elements are arranged in vertical columns called groups. Each element in a group has atoms with the same number of outer-shell electrons and similar properties.
• The elements are arranged in horizontal rows called periods. The number of the period gives the number of the highest energy electron shell in an element’s atoms.

Question 2

What is periodicity?

Answer 2

Periodicity is the repeating trend in properties of the elements across a period. The most obvious periodicity in properties is the trend from metals to non-metals. But others include trends in electron configuration, ionisation energy, structure

Question 3

1. What is the trend across a period for electron configuration?
2. What is the trend down a group for electron configuration? What effect does this have?
3. How is the periodic table divided into blocks?

Answer 3

1. Each period starts with an electron in a new highest energy shell. For each period, the s- and p-sub-shells are filled in the same way – a periodic pattern.
2. Elements in each group have the same number of electrons in their outer shell. Hence, they also have the same number of electrons in each sub-shell in the outer shell. This similarity in electron configuration gives elements in the same group their similar chemistry.
3. The elements in the periodic table can be divided into blocks corresponding to their highest energy sub-shell. This gives four blocks (s, p, d and f).

Question 4

How do blocks show the order of sub-shell filling?

Answer 4

Question 5

What are the names and numbers of all the groups in the perodic table?

Answer 5

Question 6

1. What is an ionisation energy?
2. Define the first ionisation energy?

Answer 6

1. Ionisation energy measures how easily an atom loses electrons to form positive ions.
2. The first ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.

Question 7

Explain the three factors affecting ionisation energy.

Answer 7

Three factors affect the attraction between the nucleus and the outer electrons of an atom, and therefore, the ionisation energy.

• Atomic radius. The greater the distance between the nucleus and the outer electrons, the less the nuclear attraction. The force of attraction falls of sharply with increasing distance, so atomic radius has a large effect.
• Nuclear charge. The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons.
• Electron shielding. Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons. This repulsion, called the shielding effect, reduces the attraction between the nucleus and the outer electrons.

Question 8

1. Define the second ionisation energy.
2. Why is the second ionisation energy for helium higher than the first ionisation energy?

Answer 8

1. The second ionisation energy is the energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.
2. In a helium atom, there are two protons attracting two electrons in the 1s sub-shell. After the first electron is lost, the single electron is pulled closer to the helium nucleus. The nuclear attraction on the remaining electron increases and more ionisation energy will be needed to remove this second electron.

Question 9

How do successive ionisation energies provide evidence for the different electron energy levels in a given atom?

Answer 9

Large increases between ionisation energies suggest that there is a change in the shell from which the electron is being removed.

This is because as the shell number decreases, the electrons are closer to the nucleus and with less shielding. Therefore, significantly higher energy is needed to remove them.

Question 10

Using the data shown, identify:

• the number of electrons in the outer shell
• the group of the element in the periodic table
• the identity of the element.

Explain your answer.

Answer 10

• The ionisation energies steadily increase but there is a large increase between the third and fourth ionisation energies. This shows that the fourth electron is being removed from an inner shell.
• Therefore there are three electrons in the outer shell and the element must be in Group 3.
• Since it is in Period 3, the element must be aluminium.

Question 11

What is the trend in first ionisation energy down a group? Explain this trend

Answer 11

First ionisation energies decrease down a group. You can see this trend by comparing the first ionisation energies of the three noble gases helium, neon and argon.

This happens because, as you go down a group,

• atomic radius increases
• more inner shells so shielding increases
• nuclear attraction on outer electrons decreases
• first ionisation energy decreases.

Question 12

What is the trend in first ionisation energy across a period? Explain this trend.

Answer 12

There is a general increase in first ionisation energy across the first three periods.

This happens because, as you go across a period,

• the nuclear charge increases
• there is similar shielding as the electrons are in the same shell
• nuclear attraction increases
• atomic radius decreases
• first ionisation energy increases.

Question 13

Describe and explain sub-shell trends in first ionisation energy across Period 2.

Answer 13

Although first ionisation energy shows a general increase across both Period 2, the first ionisation energy graph shows three rises and two falls:

• a rise from lithium to beryllium
• a fall to boron followed by a rise to carbon and nitrogen
• a fall to oxygen followed by a rise fluorine and neon.

This is linked to the pattern of fillinf of the s - and p-sub-shells.

Question 14

Compare the first ionisation energy of beryllium to boron. Explain.

Answer 14

The fall in first ionisation energy from beryllium to boron marks the start of filling the 2p sub-shell.

The 2p sub-shell in boron has a higher energy that the 2s sub-shell in beryllium. Therefore, in boron the 2p electron is easier to remove than one of the 2s electrons in beryllium. The first ionisation energy of boron is less than the first ionisation energy of beryllium.

Question 15

Compare the first ionisation energy of nitrogen to oxygen. Explain.

Answer 15

The fall in first ionisation energy from nitrogen to oxygen marks the start of electron pairing in the p-orbitals of the 2p sub-shell.

• In nitrogen and oxygen, the highest energy electrons are in a 2p sub-shell.
• In oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom.
• Therefore, the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen.

Question 16

What is metallic bonding? Explain.

Answer 16

Metallic bonding is the strong electrostatic attraction between cations (positive ions) and delocalised electrons.

• In a solid metal structure, each atom has donated its negative outer-shell elecrons to a delocalised sea of electrons. These electrons are mobile and are able to move throughout the structure.
• The positive cations left behind consist of the nucleus and the inner electron shells of the metal atoms. The cations ar fixed in position maintaining the structure and shape of the metal.

Question 17

Describe the structure of a solid metal.

Answer 17

In a metal structure, billions of metal atoms are held together by metallic bonding in a giant metallic lattice.

Question 18

Describe the four main properties of metals.

Answer 18

Most metals have:

• strong metallic bonds – attraction between positive ions and delocalised electrons
• high electrical conductivity
• high melting and boiling points
• very low solubility.

Question 19

Describe and explain how metals can conduct electricity. Compare this with the conductivity of ionic compounds.

Answer 19

Metals conduct electricity in solid and liquid states. When a volatge is applied across a metal, the delocalised electrons can move through the structure, carrying charge.

In comparison, ionic compounds can conduct electricity in the liquid state (when the ions are free to move) but have no mobile charge carriers in the solid state.

Question 20

Explain why metals have high melting and boiling points.

Answer 20

For most metals, high temperatures are necessary to provide the large amount of energy needed to overcome the strong electrostatic attraction between the cations and electrons in metallic bonding.

Question 21

Describe and explain the solubility of metals.

Answer 21

Metals do not dissolve as any interactions between polar solvents and the charges in a metallic lattice would lead to a reaction rather than dissolving, as with sodium and water.

Question 22

Compare the structure of simple covalent and giant covalent molecules. Explain.

Answer 22

• Most non-metallic elements exist as simple covalently bonded molecules. In the solid state, these molecules form a simple molecular lattice structure held together by weak intermolecular forces. These structure therefore have low melting and boiling points.
• The non-metals boron, carbon and silicon have different lattice structure: a very large number of atoms are held together by a network of strong covalent bonds to form a giant covalent lattice.

Question 23

Describe and explain the tetrahedral arrangement of carbon and silicion.

Answer 23

• Carbon and silicon atoms have four electrons in the outer shells. Carbon (in its diamond form) and silicon use these four electrons to form covalent bonds to other carbon or silicon atoms.
• The result is a tetrahedral structure with bond angles of 109.5˚ by electron-pair repulsion.

Question 24

Describe and explain the properties of giant covalent lattice structures.

Answer 24

• Melting and boiling points. Giant covalent lattices have high melting and boiling points. This is because covalent bonds are strong. High temperatures are necessary to provide the large quantity of energy needed to break the strong covalent bonds.
• Solubility. Giant covalent lattices are insoluble in almost all solvents. The covalent bonds holding together the atoms in the lattice are too strong to be broken by interaction with solvents.
• Electrical conductivity. Giant covalent lattices are non-conductors of electricity. The only exceptions are graphene and graphite, which are forms of carbon.
  
  • In graphene and graphite, carbon forms giant covalent structures based on planar hexagonal layers. Only three electrons of the four outer-shell electrons are used in covalent bonding. The remaining electron is released into a pool of delocalised electrons. Hence, graphene and graphite are good electrical conductors.

Question 25

Describe and explain the structure of graphene and graphite.

Answer 25

• Graphene and graphite are giant covalent structures of carbon based on planar hexagonal layers with bond angles of 120˚ by electron-pair repulsion.
  
  • Graphite is composed of parallel layers of hexagonally arranged carbon atoms. The layers are bonded by weak London forces.
  • Graphene is a single layer of graphite composed of hexagonally arranged carbon atoms linked by strong covalent bonds.
• In graphene and graphite, only three electrons of the four outer-shell electrons are used in covalent bonding. The remaining electron is released into a pool of delocalised electrons. They are therefore good electrical conductors.

Question 26

Describe and explain the trend in melting point of the elements in Period 2 and Period 3.

Answer 26

Across Period 2 and Period 3,

• the melting point increases from Group 1 to Group 14 (4)
• there is a sharp decrease in melting point between Group 14 (4) and Group 15 (5)
• the melting points are comparatively low from Group 15 (5) to Group 18 (0).

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The increase followed by the sharp decrease in melting points marks the change from giant metallic to giant covalent to simple molecular structures.

This trend in melting points continues across the s- and –blocks from Period 4 downwards as well.