Topic 1 - Atomic Structure & The Periodic Table

Question 1

Meaning Relative Atomic Mass (Ar)

Answer 1

The weighted mean mass of an atom of an element, compared to 1/12th of the mass of an atom of carbon-12

Question 2

Meaning Relative Isotopic Mass

Answer 2

The mass of an atom of an isotope, compared to 1/12th of the mass of carbon-12

Question 3

Meaning Relative Molecular Mass (Mr)

Answer 3

The mean mass of a molecule, compared to 1/12th the mass of an atom of carbon-12

Question 4

Name all of the orbitals

Answer 4

s, p, d, f

Question 5

Describe the s orbital

Answer 5

Has 1 orbital, so can hold 2 electrons

Question 6

Describe the p orbital

Answer 6

Has 3 orbitals, so can hold 6 electrons (2x3)

Question 7

Describe the d orbital

Answer 7

Has 5 orbitals, so can hold 10 electrons (5x2)

Question 8

Describe the f orbital

Answer 8

Has 7 orbitals, so can hold 14 electrons (7x2)

Question 9

Which shape does an S orbital have

Answer 9

Spherical shape

Question 10

Which shape does a p orbital have

Answer 10

Dumb-bell shaped

Question 11

Give the electron configuration of iron (26)

Answer 11

1s2 2s2 2p6 3s2 3p6 4s2 3d6

Question 12

Give the electron configuration of calcium +2

Question 13

Give the electron configuration of iron +3

Question 14

Give the electron configuration of copper (29)

Answer 14

1s2 2s2 2p6 3s2 3p6 4s1 3d10

Question 15

State the different radiations in order of the electromagnetic spectrum

Answer 15

Radio waves, micro waves, infrared, visible light, ultraviolet, X-rays, gamma rays

Question 16

When do electrons become excited

Answer 16

When electrons absorb energy (endothermic) as it moves up an energy level

Question 17

In an atomic emission spectra, why do the lines get closer

Answer 17

Because energy and frequency increases

Question 18

When do electrons show ultraviolet radiation

Answer 18

When an electron falls to the ground state (n=1)

Question 19

When do electrons show visible radiation

Answer 19

When an electron falls to the second energy level (n=2)

Question 20

When do electron show infrared radiation

Answer 20

When electrons fall to the third energy level (n=3)

Question 21

Meaning ionisation

Answer 21

The minimum amount of energy required to remove 1 mole of electrons from 1 mole of atoms in a gaseous state

Question 22

What are 3 factors of ionisation

Answer 22

Shielding, nuclear charge, atomic size

Question 23

How does shielding affect ionisation

Answer 23

The more electron shells between the positive nuclei and the negative electron thats is being removed, the less energy required to ionise because theres a weaker attraction

Question 24

How does nuclear charge affect ionisation

Answer 24

The more protons in the nucleus, the larger the attraction between the nucleus and the outer electrons. This means more energy is needed for ionisation

Question 25

How does atomic size affect ionisation

Answer 25

The bigger the atom, the further away the outer electrons are from the nucleus. The attractive force between the nucleus and the outer electrons reduces, so less energy is needed to ionise an electron

Question 26

Why does the atomic radii decrease across a period

Answer 26

The further across a period, there higher number of protons in the nucleus. This pulls the outer electrons further towards the nucleus. Shielding has no effect

Question 27

Why does the melting points of metals increase across a period

Answer 27

As the metals have more protons in their nucleus, there’s an increase in positive charge. So, there’s an increased number of delocalised electrons and smaller atomic radii. This means stronger metallic bonds, so more energy is needed to bream the binds.

Question 28

Explain why silicon (S8) has the highest melting point in period 3

Answer 28

Silicon is a giant covalent structure. This means it has many strong covalent bonds that hold the silicon atoms together. Thus, a large amount of energy is needed to overcome these string covalent bonds

Question 29

Why does sulfur have a lower melting point than phosphorus, even though it’s further across period 3

Answer 29

Sulfur has a larger simple molecular structure and has larger london forces than phosphorus

Question 30

Why does argon have the lowest melting point in group 3

Answer 30

Argon only exists as individual atoms as its a noble gas, so it has small london forces

Question 31

describe the trend in reactivity in group 1 and 2 elements

Answer 31

the elements increase in reactivity as you go down the group

Question 32

describe the trend in reactivity in group 6 and 7 elements

Answer 32

the elements decrease in reactivity as you go down the group

Question 33

why do group 1 and 2 elements increase in reactivity as you go down the group

Answer 33

going down the group, increases the distance of outer shell electrons from the nucleus, plus the shielding effect of the inner shell electrons overcomes the increasing nuclear charge. Thus, the outer shell electrons are held less strongly and are easier to remove in a reaction.

Question 34

why do group 6 and 7 elements increase in reactivity as you go up the group

Answer 34

going up the group, reduces the distance of the outer shell from the nucleus and decreases shielding effects on the decreasing positive charge, meaning that the outer shell electrons are more strongly attracted to the nucleus. Thus, the outer shell is more readily accepts another electron in a reaction

Question 35

why does the 1st ionisation energy for the elements increase across a period

Answer 35

as the outer electron is removed from elements of increasing atomic number within a period, it is being removed from the same quantum shell, but the nuclear charge is steadily becoming more positive so holding the electron with a greater force, thus more energy is needed to remove it

Question 36

why does the 1st ionisation energy for the elements decrease down a group

Answer 36

as the electron is removed from an outer shell which is increasing in distance from the nucleus. This increasing distance and the increased shielding effect of the increasing number of inner shells present, overcomes the fact that the nuclear charge is increasing. Thus, the outer electron is held less strongly and less energy is needed to remove it

Question 37

why do the successive ionisation energies increase for electrons taken from the same quantum shell

Answer 37

As you take successive electrons from the same quantum shell, the electrons are a similar distance from the nucleus, with the same inner shell shielding, but the ionic charge increases each time, so the attractive force on the electron being removed increases so it requires more energy to take the away

Question 38

why are there large jumps in successive energies for an elements such as sodium between Em1 to Em2 and Em9 to EM10

Answer 38

As you successfully take electrons away from Na, the 1st electron is from the 3rd Quantum shell, which is further from the nucleus than the 2nd electron, which is from the 2nd Quantum shell, which is nearer to the nucleus so attracts the electron more strongly. After 9 electrons have been removed, the 10th electron is now taken from the first Quantum shell, which is again much closer to the nucleus, so attracts the electron more strongly

Question 39

why is there a drop in the 1st I.E. between the elements Mg and Al in the 3rd period

Answer 39

the outer electron is removed from a 3P orbital in Al, but from the 3s orbital in Mg. The energy of the electrons in the 3p subshell is greater than the energy of the 3s subshell. Thus, it takes less energy to remove a 3p electron in Al, than a 3s electron in Mg