Chapter 7

Question 1

Describe the arrangement of elements in the periodic table

Answer 1

By increasing atomic ( proton ) number

Question 2

Describe the arrangement of elements ( in periods ) in the Periodic Table

Answer 2

show repeating trends in physical and chemical properties (periodicity)

Question 3

What is periodicity ?

Answer 3

The repetition of properties of the chemical elements across periods in the periodic table

Question 4

Describe the arrangement of elements ( in groups ) in the Periodic table

Answer 4

They have similar chemical properties

Question 5

Explain the Periodic trend in electron configurations across Periods 2 and 3

Answer 5

Across period 2, the 2s sub shell fills with 2 electrons and then the 2p subs shells fill with 6 electrons. In period 2, the highest energy electrons are in the second shell.
In period 3 the pattern repeats itself with the 3s and 3p sub-shells. In period 3, the highest energy electrons are in the third shell.

Question 6

How do you classify elements into s-, p- and d-blocks?

Answer 6

s block element = highest energy electron is a s electron
p block element = highest energy electron is a p electron
d block element = highest energy electron is a d electron

Question 7

Define first ionisation energy

Answer 7

The energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to make a mole of gaseous 1+ ions ( in kJmol-1 )

Question 8

Which element has the highest first ionisation energy ?

Answer 8

Helium

Question 9

What is meant by successive ionisation energies ?

Answer 9

Successive ionisation energies are the amounts of energy needed to remove electrons one after the other from an atom

Question 10

What 3 factors affect the first ionisation energies of elements ?

Answer 10

1) Distance from the nucleus - the further the electron is from the nucleus, the lower the
nuclear attraction and the I.E. will be decreased
2) Nuclear charge - the more protons in the nucleus, the greater the nuclear
attraction and the I.E. will be increased
3) Shielding - filled inner shells exert a shielding effect lowering the nuclear attraction so the I.E. will be decreased

Question 11

Explain the trend in first ionisation energies across Periods 2 and 3, in terms of attraction, nuclear charge and atomic radius?

Answer 11

Nuclear charge increases
Same shell – same shielding
Nuclear attraction increases
Atomic radius decreases
First ionisation energy increases

Question 12

What 2 elements are exceptions to the trend in first ionisation energies across Periods 2 and 3 and why ?

Answer 12

Boron and Oxygen

eg Beryllium vs Boron : the 2p subshell has higher energy than 2s sub shell therefore the 2p electron is easier to remove and less energy is needed so Boron has a lower 1st ionisation energy than Beryllium
eg Oxygen vs Nitrogen : the 2 paired 2p electrons in Oxygen repel each other, these electrons are at higher energies, which makes it easier to remove an electron from an Oxygen atom therefore O has lower 1st IE than N

Question 13

Explain the trend in first ionisation energies down a group, in terms of attraction, nuclear charge and atomic radius?

Answer 13

The atomic radius increases
Nuclear charge increased but cancelled out by other factors
Shielding increases (more inner shells)
Overall, nuclear attraction on outer electrons decreases
First ionisation energy decreases

Question 14

How to predict from successive ionisation energies of the number of electrons in each shell of an atom and the group of an element?

Answer 14

A large jump in the ionisation energy graph indicates that an electron is being removed from a more stable shell. The number of electrons in the outer shell of the atom determines the group of the element. For example, if the fifth electron experiences a large jump in ionisation energy, the element is in group 5.
The first ionisation energy is usually lower than the second because the attraction between the nucleus and the outer electrons increases each time an electron is removed. This makes it harder to remove the next electron, requiring more energy.
Large increase in ionisation energy when an electron comes from a new shell.
Slight dip in ionisation energy if an electron is removed from a p-orbital compared to an s-
orbital because the 3p subshell has higher energy.
Slight dip in energy if removal of an electron reduces electron pairing.
Removing an electron from an atom will make it smaller.
Adding an electron to an atom will make it bigger.
So anions are usually large, and cations are usually small.

Question 15

What is the trend in successive ionisation energies ?

Answer 15

They increase in value
Once an electron has been removed, the remaining electrons are held more tightly ( outer electron is closer to nucleus, experiences greater attraction )
Electron has to be removed from a + charged ion rather than a neutral atom so more energy required to remove an electron

Question 16

Define metallic bonding and explain why this is the case

Answer 16

Metallic bonding is the strong electrostatic attraction between cations (positive ions) and delocalised electrons. Each metallic atom has donated its negative outer shell electron/s to a shared pool of electrons, which are delocalised through the whole structure ( they are mobile ). The + cations left behind are fixed in position

Question 17

Explain a giant metallic lattice structure

Answer 17

There are no molecules so no inter molecular forces
Each metallic atom has donated its negative outer shell electron/s to a shared pool of electrons, which are delocalised through the whole structure ( they are mobile ). The + cations left behind are fixed in position.

Question 18

Explain solid giant covalent lattices of carbon (diamond, graphite and graphene) and silicon as networks of atoms bonded by strong covalent bonds?

Answer 18

Diamond : strong, no layers, 4 covalent bonds per atom, high m.p, non conductor, insoluble in all solvents ( no possible interactions which could occur with solvent molecules )
Graphite : soft and weak structure, layer move over each other easily ( bonded by weak london forces ), 3 covalent bonds per atom, conductor due to delocalised electron between layers ( only 3 electrons used in bonding ), trigonal planar shape, 120° bond angle
Graphene : thick sheet of carbon atoms, strong material, conductor of heat and electricity due to delocalised electrons
Silicon : similar structure to diamond, forms four bonds, giving a tetrahedral arrangement with a bond angle of 109.5°

Question 19

Silicon dioxide vs Carbon dioxide

Answer 19

CO2 : sublimes at low temp, gas at room temp, simple covalent molecular structure with weak intermolecular forces
SiO2 : high mp and bp, solid at room temp, giant covalent network structure

Carbon atoms are small enough to form double bonds to oxygen atoms
Silicon atoms are too large to form double bonds to oxygen atoms

SiO2 : tetrahedral arrangement, each silicon covalently bonded to 4 oxygens, no delocalised electrons so doesn’t conduct electricity, insoluble in water and organic solvents as there are no possible attraction that could occur between solvent molecules and Si/O atoms which could overcome covalent bonds in giant structure

Question 20

Describe the physical properties of giant metallic lattices, including melting and boiling points, solubility and electrical conductivity in terms of structure and bonding?

Answer 20

High electrical conductivity.
Metals conduct when solid and in molten (liquid) form. The delocalised electrons can move throughout the structure, carrying charge.
Apart from mercury and some metals not in the d-block, a high temperature is required to provide the quantity of energy needed to overcome the electrostatic attraction between
cations and delocalised electrons
Metals are not soluble in water. Some, however, react with water
Metals are insoluble and do not dissolve in solvents.

Question 21

Describe the physical properties of giant covalent lattices, including melting and boiling points, solubility and electrical conductivity in terms of structure and bonding?

Answer 21

The strong covalent bonds make these structures very stable.
They have high melting and boiling points. High temperatures are needed to provide enough energy to break the strong covalent bonds.
They are insoluble in almost all solvents.
They are mainly non-conductors, with a few exceptions like graphite and graphene.
the structures of carbon dioxide and silicon dioxide are quite different.
But carbon itself (the element) can form giant covalent structures.
Not only can carbon form giant covalent structures, it can form more than one type.
Many elements can exist in different forms depending on how their atoms are
arranged/bonded.
Different forms of an element are called allotropes.

Question 22

Explain the variation in melting points and boiling points across Periods 2 and 3 in terms of structure and bonding?

Answer 22

Melting point increases from Group 1 to Group 4 : metallic to giant covalent
Group 4 element has highest m.p and strongest structure
From Group 5 onwards, m.p is low as these are simple/discrete molecules with simple molecular structures
Metals ( Group 1 to 3 ) : giant metallic lattice held together by strong attractive forces
Giant covalent network ( Group 4) : held together by strong covalent bonds
Simple molecules ( Group 5 to 8 ) : held together by weak intermolecular forces