Atomic Structure

Question 1

Relative charges of proton neutrons and electrons:

Answer 1

+1, 0, -1

Question 2

Relative masses of protons, neutrons and electrons:

Answer 2

1, 1, negligible

Question 3

Behaviour of protons, neutrons and electrons in an electric field (direction and angle of deflection):

Answer 3

Protons: towards the negatively charged plate, y˚ where y˚y˚

Question 4

What is direction of deflection dependent on?

Answer 4

Charge of the particle.

Question 5

What is the angle of deflection proportional to?

Answer 5

charge/mass ratio of the particle.The larger the charge q, the stronger the attraction towards the oppositely charged plate. The larger the mass m, the more difficult it is to cause the particle to deviate towards the oppositely charged plate.

Question 6

What is the distribution of mass in an atom?

Answer 6

Bulk of the mass is in the nucleus.

Question 7

What is the distribution of charge in an atom?

Answer 7

Positive charges are located in the nucleus, negative charges are located in the electron cloud around the nucleus.

Question 8

What is nucleon number?

Answer 8

Total number of protons and neutrons.

Question 9

What is an isotope?

Answer 9

Atoms with the same number of protons but different number of neutrons.

Question 10

How are electrons contained in an atom?

Answer 10

Around the nucleus are electronic shells, each shell contains subshells, each subshell contains orbitals, each orbital contains electrons.

Question 11

What happens when the principal quantum value of a shell, n, increases?

Answer 11

The further the shell is from the nucleus, the higher the energy level of the shell, the weaker the electrostatic attraction between the nucleus and electrons, the more diffused the orbital.

Question 12

What is an orbital?

Answer 12

A region in space where there is a high probability of finding an electron. Each can accommodate 2 electrons, has a distinctive geometrical shape, and its energy is the energy of the electron occupying it.

Question 13

Traits of the s orbital:

Answer 13

• Spherical shape
• Non-directional
• As n increases, s orbital becomes more diffuse

Question 14

Traits of the p orbitals:

Answer 14

• Dumbbell shape
• Three p orbitals (Px, Py, Pz)
• Directional as electron density is concentrated in certain directions along the x, y and z axes
• The three orbitals are degenerate
• As n increases, p orbital becomes more diffuse

Question 15

Traits of the d orbitals:

Answer 15

For Dxz, Dxy, Dyz:
- 4 lobed shape
- Lobes pointing between the axes
For Dx^2-y^2:
- 4 lobed shape
- Lobes aligned along the x and y axes
For Dz^2:
- Dumbbell surrounded by a small ring at its waist
- Aligned along the z axis

Question 16

Why is the 4s orbital filled before the 3d orbital?

Answer 16

Due to its proximity to the nucleus, s orbitals are affected to a greater degree than p and d orbitals and hence their energies decrease more rapidly than the other orbitals. As atomic number increases, there is increased nuclear charge, and the energies of orbitals decrease.

Question 17

What is the Aufbau Principle?

Answer 17

Electrons fill orbitals from the lowest energy orbitals upwards (4s before 3d). This is because lowest energy states are most stable.

Question 18

What is Hund’s Rule?

Answer 18

Orbitals of a subshell must be occupied singly before pairing to minimise inter-electronic repulsion.

Question 19

What is the Pauli Exclusion Principle?

Answer 19

Each orbital can hold up to a maximum of 2 electrons and they must be of opposite spins. This ensures the magnetic attraction can counterbalance the electronic repulsion from their identical charges, allowing them to be stable.

Question 20

Anomalous electronic configuration of Cr:

Answer 20

Instead of [Ar] 3d^4 4s^2, electronic configuration is [Ar] 3d^5 4s^1. 3d and 4s orbitals are similar in energy levels for Cr, and this minimises inter-electronic repulsion.

Question 21

Anomalous electronic configuration of Cu:

Answer 21

Instead of [Ar] 3d^9 4s^2, electronic configuration is [Ar] 3d^10 4s^1. This is because a fully filled 3d subshell is unusually stable due to symmetrical charge distribution around the nucleus.

Question 22

What is ground state?

Answer 22

Electrons are in the orbitals of the lowest available energy level.

Question 23

What is excited state?

Answer 23

Electrons absorb energy and are promoted to a higher energy level.

Question 24

What is an isoelectronic species?

Answer 24

Species with the same total number of electrons.

Question 25

What affects electrostatic attraction?

Answer 25

Number of electronic shells, nuclear charge, shielding effect.

Question 26

How does number of shells affect electrostatic attraction?

Answer 26

As number of shells increases, n increases, distance between nucleus and valence electrons increases, electrostatic attraction decreases.

Question 27

How does nuclear charge affect electrostatic attraction?

Answer 27

As number of protons increases, nuclear charge increases, electrostatic attraction increases.

Question 28

How does shielding effect affect electrostatic attraction?

Answer 28

As number of inner electrons increases, shielding effect experienced by valence electrons increases, electrostatic attraction decreases.

Question 29

What is shielding?

Answer 29

Repulsion of negatively charged electrons in inner shells on those in outer shells to prevent them from experiencing the full effect of the actual nuclear charge.

Question 30

Why does shielding ability of electrons decrease from s>p>d>f?

Answer 30

d and f orbitals are more diffuse.

Question 31

What is effective nuclear charge?

Answer 31

Resultant positive charge experienced by valence electrons taking into account shielding effect (nuclear charge - shielding effect).

Question 32

How does atomic radii change across a period?

Answer 32

It decreases across a period. Across a period, number of shells remain the same, number of protons and hence nuclear charge increases, number of electrons increases as well but are added to the valence shell, hence shielding effect is constant. Effective nuclear charge increases, electrostatic attraction between nucleus and valence electrons increases, thus electron cloud size decreases.

Question 33

How does atomic radii change down a group?

Answer 33

Down a group, number of electronic shells increases, distance between nucleus and valence electrons increases, shielding effect experienced by valence electrons increases. Despite increasing nuclear charge, electrostatic attraction decreases, resulting in increase in electron cloud size.

Question 34

How does cationic radius compare to atomic radius?

Answer 34

Cationic radius is smaller. Both have the same number of protons and hence same nuclear charge, however cation has one less electronic shell. Electrostatic attraction increases, decreasing electron cloud size of cation.

Question 35

How does anionic radius compare to atomic radius?

Answer 35

Anionic radius is larger. Both have the same number of protons and hence same nuclear charge. However, anions have more electrons, thus electron-electron repulsion increases, so electrostatic attraction decreases, increasing electron cloud size of anion.

Question 36

How does ionic radii of isoelectronic species change across a period?

Answer 36

It decreases. Ions have the same shielding effect, however number of protons and hence nuclear charge increases, thus effective nuclear charge increases. Electrostatic attraction also increases, decreasing electron cloud size.

Question 37

What is first ionisation energy?

Answer 37

Energy required to remove 1 mol of electrons from 1 mole of M to form 1 mole of M+.

Question 38

What is n ionisation energy?

Answer 38

Energy required to remove 1 mol of electrons from 1 mole of M(n-1)+ to form 1 mole of M(n)+.

Question 39

How does first ionisation energies change across a period?

Answer 39

It generally increases, Number of shells remains the same, number of protons and hence nuclear charge increases, number of electrons increase but electrons are added to valence shell, so shielding effect is the same. Effective nuclear charge increases, electrostatic attraction increases and hence more energy is required to remove the electron.

Question 40

Why does first ionisation energy decrease from Group 2 to Group 13?

Answer 40

The np electron to be removed from the Group 13 atom is at a higher energy level than the ns electron to be removed from the Group 2 atom. less energy is required to remove the former than the latter.

Question 41

Why does first ionisation energy decrease from Group 15 to Group 16?

Answer 41

The np electron to be removed from the Group 16 atom is paired while that from the Group 15 electron is not. Due to inter-electronic repulsion, less energy is required to remove the former.

Question 42

How does first ionisation energies change down a group?

Answer 42

It decreases. Down a group, number of shells increases, distance between nucleus and valence electrons increases, shielding experienced by valence electrons increases. Despite increasing nuclear charge, electrostatic attraction decreases, hence energy required to remove electrons decreases as well.

Question 43

How do successive ionisation energies change?

Answer 43

They increase, as with each electron removed, the atom/ion becomes increasingly positive, which attracts electrons more strongly.

Question 44

What is electronegativity?

Answer 44

A relative measure of its ability to attract bonding electrons.

Question 45

How does electronegativity change across a period?

Answer 45

It increases. Across a period, number of shells remains the same, number of protons and hence nuclear charge increases. Number of electrons increases but electrons are added to the valence shell and hence shielding effect remains constant. Effective nuclear charge increases, electrostatic attraction between nucleus and bonding electrons increases.

Question 46

How does electronegativity change down a group?

Answer 46

It decreases. Down a group, number of shells increases, distance between bonding electrons and nucleus increases, shielding effect also increases. Despite increasing nuclear charge, electrostatic attraction between nucleus and bonding electrons increases.