Atomic Structure

Question 1

Subshell

Answer 1

A specific energy level within an electron shell. Designated s, p, d or f in order of increasing energy

Question 2

Order of subshells in increasing energy

Answer 2

s, p, d, f

Question 3

Orbital

Answer 3

A region of space in which up to two electrons may be located

Question 4

How many orbitals in a p subshell?

Answer 4

3

Question 5

How many electrons can fit in a d subshell?

Answer 5

10

Question 6

How many orbitals in an s subshell?

Answer 6

1

Question 7

How many electrons would be in an f subshell if all the orbitals were half-filled?

Answer 7

7

Question 8

Order of subshell filling up to 6p

Answer 8

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p

Question 9

Pauli Exclusion Principle

Answer 9

A maximum of two electrons can fit into any orbital, providing that those electrons have opposite spin

Question 10

Hunds Rule

Answer 10

Electrons will arrange themselves into a subshell in such a way as to maximise the number of half-filled orbitals. OR Electrons will half-fill all orbitals in a subshell before any are filled completely.

Question 11

Aufbau principle

Answer 11

Electrons will move into subshells in order from lowest energy up to highest energy

Question 12

Ground state

Answer 12

A situation where all electrons in an atom or ion are in the lowest possible energy levels (subshells)

Question 13

Excited state

Answer 13

A situation where not all electrons in an atom or ion are in the lowest possible energy levels (subshells)

Question 14

Which two elements in the first row of the d block have only one 4s electron?

Answer 14

Chromium and Copper

Question 15

What causes a spectral line on an emission spectrum?

Answer 15

An excited electron relapsing into a lower energy subshell

Question 16

What shape is an s orbital?

Answer 16

Sphere

Question 17

How many electrons can fit into the 3rd shell, and how many of these electrons fit into each subshell?

Answer 17

A total of 18 - 2 in the s subshell, 6 in the p subshell and 10 in the d subshell

Question 18

What is meant by the term ‘isotopes’?

Answer 18

Two atoms that have the same number of protons (atomic number) but a different number of neutrons (or different mass number)

Question 19

From which subshell would an iron atom lose its first electron?

Answer 19

The 4s subshell

Question 20

Trend in atomic radius across a period (L to R), and why

Answer 20

Atomic radius decreases because core charge increases but the number of shells remains the same

Question 21

Trend in atomic radius down a group, and why

Answer 21

Atomic radius increases because the number of shells increases but core charge remains the same

Question 22

Trend in core charge across a period (L to R), and why

Answer 22

Core charge increases because the atomic number increases but the number of inner shell electrons remains the same

Question 23

Trend in core charge down a group, and why

Answer 23

Core charge remains the same because the atomic number and the number of inner shell electrons increase by the same amount

Question 24

How to calculate core charge

Answer 24

Number of protons in the nucleus minus the number of inner shell electrons

Question 25

Core charge of a Silicon atom

Answer 25

4+ (14 protons minus 10 inner shell electrons)

Question 26

Core charge of a Magnesium 2+ ion

Answer 26

10+ (12 protons minus 2 inner shell electrons)

Question 27

Core charge of a Sulfide ion

Answer 27

6+ (16 protons minus 10 inner shell electrons)

Question 28

Trend in electronegativity across a period (L to R), and why

Answer 28

Electronegativity increases because core charge increases and atomic radius decreases slightly (note that noble gases are an exception as they do not easily participate in bonding)

Question 29

Trend in electronegativity down a group, and why

Answer 29

Electronegativity decreases because atomic radius increases (due to increasing number of electron shells) but core charge remains constant

Question 30

Definition of 'first ionisation energy'

Answer 30

The amount of energy required to remove an electron from each of a mole of gaseous atoms (of that element) to produce a mole of 1+ ions

Question 31

Trend in first ionisation energy across a period (L to R) and why

Answer 31

First Ionisation Energy increases due to increasing core charge and slight decrease in atomic radius

Question 32

Trend in first ionisation energy down a group and why

Answer 32

First Ionisation Energy decreases due to increasing atomic radius (due to increasing number of electron shells) but core charge remains constant

Question 33

Define Electronegativity

Answer 33

The force of attraction experienced by bonding electrons towards the nucleus of that element

Question 34

Which group has the highest electronegativity in any period after the first?

Answer 34

Group 7 (the halogens)

Question 35

Which group has the lowest electronegativity in any period after the first?

Answer 35

Group 1 (the alkali metals)

Question 36

Which group has the highest first ionisation energy in any period after the first?

Answer 36

Group 8/0 (the noble gases)

Question 37

Which group has the lowest first ionisation energy in any period after the first?

Answer 37

Group 1 (the alkali metals)

Question 38

How can you work out the group number of an element using successive ionisation energies?

Answer 38

The number of electrons removed before the first big 'jump' in ionisation energies is the number of valence electrons (and therefore group number)

Question 39

Why is the third ionisation energy of calcium much greater than the second?

Answer 39

The third electron is removed from a lower-energy shell than the first two. This electron will be closer to the nucleus and experience much less electron shielding.

Question 40

What is electron shielding?

Answer 40

A repulsive force exerted by electrons that are closer to the nucleus (ie in lower energy shells) on electrons that are further from the nucleus (ie in higher-energy shells), pushing them away from the nucleus

Question 41

Trend in the reactivity of metals down groups 1 and 2, and why

Answer 41

Reactivity increases because electrons are lost more easily, because the atomic radius increases but core charge remains the same

Question 42

What is common about metals that are found in the native state?

Answer 42

They are very unreactive. Examples include gold, silver and platinum.

Question 43

Trend in the reactivity of non-metals down group 7, and why

Answer 43

Reactivity decreases because it becomes harder for the halogen to attract an electron due to increasing atomic radius (more electron shielding) while core charge remains the same

Question 44

Definition of relative isotopic mass

Answer 44

Mass of that isotope relative to 1/12 the mass of a carbon-12 isotope

Question 45

Definition of relative atomic mass

Answer 45

The weighted average mass of all the isotopes of that element, weighted by their relative abundance

Question 46

Formula to calculate moles from mass and molar mass

Answer 46

n = m/Mr

Question 47

Formula to calculate moles from concentration and volume of a solution

Answer 47

n = CV

Question 48

Formula to calculate moles from volume and molar volume of a gas (at SLC)

Answer 48

n = V/Vm (or V/24.8)

Question 49

Formula to calculate number of particles (N) from moles and Avogadro's number (6.02 x 10^23)

Answer 49

N = n x 6.02 x 10^23

Question 50

How many moles of hydrogen atoms are present in 2.5 moles of water?

Answer 50

5 moles of hydrogen atoms

Question 51

How many moles of atoms are present in 2.5 moles of water?

Answer 51

7.5 moles of atoms (2.5 moles of oxygen atoms plus 5 moles of hydrogen atoms)