Atomic Structure and the Periodic Table

Question 1

Put these in order of size - sub-shells, orbitals, shells

Answer 1

From largest to smallest - shells, sub-shells, orbitals

Question 2

What are orbitals?

Answer 2

An orbital is a region within an atom that can hold up to two electrons with opposite spins.

Question 3

What shape are s-orbitals?

Answer 3

Spheres

Question 4

What shape are p-orbitals?

Answer 4

‘Dumbbell’ shape. Can look like and 8 or an infinity depending on orientation. There are 3 types of p orbital (depending on if they are on the x, y or z axis).

Question 5

How many orbitals does an s-subshell have?

Answer 5

1

Question 6

How many orbitals does a p-subshell have?

Answer 6

3

Question 7

How many orbitals does a d-subshell have?

Answer 7

5

Question 8

How many orbitals does an f-subshell have?

Answer 8

7

Question 9

How many electrons can be found in each type of subshell?

Answer 9

s-subshells can hold up to 2 electrons
p-subshells can hold up to 6 electrons
d-subshells can hold up to 10 electrons
f-subshells can hold up to 14 electrons

Question 10

What is the combination of subshells for each shell?

Answer 10

1st - 1s
2nd - 2s2p
3rd - 3s3p3d
4th - 4s4p4d4f

Question 11

What is the maximum number of electrons which can be held in each shell?

Answer 11

1st - 2
2nd - 8
3rd - 18
4th - 32

Question 12

Which has higher energy, the first or the second quantum shell?

Answer 12

The second quantum shell has higher energy than the first.

Question 13

What is the order of energies of the orbitals in the 4th shell?

Answer 13

4s<4p<4d<4f

Question 14

How do you write the electronic configuration of atoms using 1s notation?

Answer 14

The shell number is written first, then the sub-shell type. The number of electrons occupying that subshell are written in superscript above the subshell type.
e.g. Magnesium atom would be written as
1s^2 2s^2 2p^6 3s^2

Question 15

How to you use electrons-in-boxes notation?

Answer 15

Draw a box to represent a sub-shell, divided into orbitals (e.g. p-subshells would have boxes divided into 3 as they have 3 orbitals). Draw arrows to represent the electrons (up to 2 in each orbital). If there are 2 electrons in an orbital, they must have opposite spins (so the arrows must point in opposite directions).

Question 16

What is Hund’s principle (filling order)?

Answer 16

Electrons fill subshells singly before pairing up due to the electrostatic repulsion between electrons.

Question 17

What is the Pauli Exclusion principal (orbitals)?

Answer 17

Two electrons cannot occupy the same orbital unless they have opposite spins.

Question 18

Which subshells are filled first?

Answer 18

Subshells with the lowest energy are filled first.
The filling order goes 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d and so on

Question 19

Will electrons always be found in the orbitals where they are supposed to be?

Answer 19

No. There is a 90% probability of finding the electron within the orbital’s boundaries.

Question 20

What is the difference between a 1s subshell and a 2s subshell?

Answer 20

2s subshells have the same shape as 1s subshells (spherical) but are larger

Question 21

What is the first ionisation energy?

Answer 21

The energy required to remove one electron from each atom in 1 mole of gaseous atoms to form one mole of gaseous 1+ ions.

Question 22

What is the second ionisation energy?

Answer 22

The energy required to remove one electron form every ion in a mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions.

Question 23

What is the general equation for ionisation?

Answer 23

X^(n-1) —> X^n + e-

Question 24

How does number of protons (nuclear charge) affect ionisation energy?

Answer 24

The greater the number of protons in an atom, the stronger the positive charge of the nucleus. The stronger the positive charge of the nucleus, the stronger the electrostatic attraction between the nucleus and the electrons. This means that, generally, more energy is required to remove the outer electron of an atom with a greater number of protons than an atom with a lesser number of protons (higher ionisation energy).

Question 25

How does electron shielding affect ionisation energy?

Answer 25

Electron shielding occurs when there are electrons between the outer electrons and the nucleus. These inner electrons repel the outer electrons (as they both have negative charges), meaning the outer electrons have a weaker attraction to the nucleus, so less energy is required to remove one of them, so greater shielding means lower ionisation energy.

Question 26

How does the shell or subshell from which an electron is removed affect ionisation energy?

Answer 26

Different shells and subshells lie at different distances from the nucleus. The further away the outer electron is from the nucleus, the weaker its attraction to the nucleus, and therefore the less energy is needed to remove it. Therefore, the higher the energy of the quantum shell form which an electron is removed (and the further the distance from the nucleus), the lower the ionisation energy. The subshell also has an impact. Generally, electrons being removed from s-subshells have higher ionisation energies than electrons being removed from p-subshells. This is because the shape of an s-subshell means that electrons found in them are likely to be closer to the nucleus than those found in p-subshells, so more energy is required to remove an electron from an s-subshell than a p-subshell. This is one reason why electrons in the same shell have slightly different ionisation energies.

Question 27

Why do successive ionisation energies increase in magnitude?

Answer 27

The first ionisation energy is the lowest because electrons are being removed from the outermost shell. This shell is the furthest distance from the nucleus, and so the electrons in it experience the least amount of attraction. It requires the least energy to remove them. If there is more than one electron in the outer shell, the next ionisation energies will increase gradually until the next shell is reached. This is because when electrons are removed from a shell, the electron-electron repulsion between the remaining electrons becomes less strong. This causes the remaining electrons to be more strongly attracted to the nucleus and so more energy is required to remove them. Also, ionisation energies increase when electrons are removed from s-subshells instead of p-subshells, as electrons in s-subshells are likely to be closer to the nucleus than those in p-subshells. There are large increases in ionisation energy when an electron is removed from the next shell down. This is because the next shell down is much closer to the nucleus, so is more strongly attracted to it. Also, there is less electron shielding in lower shells, so attraction to the nucleus is stronger. Overall, successive ionisation energies increase gradually within a shell, and jump when a new shell is reached.

Question 28

Explain the trend in ionisation energy across a period.

Answer 28

As you move from left to right across a period, the number of protons in each atom increases. This means that the electrons become more strongly attracted to the nucleus as you go along, so more energy is required to remove them. However, one extra electron is also added to the same quantum shell on each occasion, which causes greater electron-electron repulsion, decreasing the ionisation energy. However, the increasing nuclear charge has a greater effect than the shielding, so the general trend is that ionisation energies increase across a period.

Question 29

Why is there a decrease in ionisation energy between magnesium and aluminium?

Answer 29

In aluminium, the electron is being removed from a p-subshell, whereas in magnesium, the electron is being removed from the s-subshell. P-subshells are at higher energy than s-subshells, so it takes less energy to remove an electron from aluminium than magnesium. This outweighs the effect of increased nuclear charge.

Question 30

Why is there a decrease in ionisation energy between phosphorus and sulfur?

Answer 30

In P, all of the orbitals in its outer shell are half-filled, whereas in S, one of the outer shell’s orbitals is full. The two electrons in the full orbital repel each other, and so are less strongly attracted to the nucleus. Less energy is required to remove an electron from the full orbital in S than the half-filled orbitals in P, so the ionisation energy decreases.

Question 31

Explain the trend in ionisation energies down a group.

Answer 31

As you go down a group, the number of protons in each atom increases, so this factor alone would increase ionisation energy as attraction to the nucleus would be stronger. However, as you go down a group, a new quantum shell is added with each element. Each new shell is further from the nucleus, so experiences a weaker attraction to it. Furthermore, there is more electron shielding the further out the outer shell is, which repels the electrons in the outer shell, making them easier to remove. The effect of shielding and increasing numbers of shells outweighs the effect of increasing nuclear charge and so the general trend is that as you go down a group, ionisation energies decrease.

Question 32

What is periodicity?

Answer 32

Periodicity is a repeating pattern across a period of the periodic table.

Question 33

Explain the trend in atomic radii across a period.

Answer 33

As you go across a period, the atomic radius of the elements decreases. This is because the nuclear charge increases and so the electrostatic attraction between the electrons and the nucleus is stronger, meaning the electrons are pulled closer to the nucleus. While there is a slight increase in shielding and electron-electron repulsion as you go along the period (as the number of electrons increases), this is outweighed by the increase in nuclear charge, so the atomic radii decrease.

Question 34

What is the atomic radius?

Answer 34

It is a measure of the size of an atom. It is the distance from the centre of the nucleus to the boundary of the electron cloud. However, since this is difficult to measure, the atomic radius is usually found by determining the distance between two nuclei and dividing by two.

Question 35

What is the trend in melting and boiling temperatures across a period.

Answer 35

In group 2, the trend is that the mtp and btp start high then decrease rapidly as you enter the non-metals.
In group 3, the trend is that the mtp and bpt start fairly high, then go really high at silicon, then decrease rapidly when the other non-metals are reached.

Question 36

Explain the trend in melting and boiling points across a period

Answer 36

In period 2, they start high. This is because the metallic elements come first. They have metallic bonding, which consists of strong electrostatic forces of attraction which occur between positive ions and the sea of delocalised electrons. These require lots of energy to break, so the melting and boiling points are high. Then, there is a rapid increase between beryllium and boron as boron has a giant covalent structure. Covalent bonds are the extremely strong electrostatic forces of attraction which occur between positive nuclei and negative shared pairs of electrons. Giant covalent structures have very high melting points as in order for them to melt, all of their covalent bonds need to be broken, which requires significant amounts of energy. Then, the temperatures drop significantly between carbon and nitrogen as nitrogen has a simple molecular covalent structure. There are weak intermolecular forces between molecules, which don’t require much energy to break, so mtp and bpt are low. They generally decrease along the non-metals as the molecules get smaller. Smaller molecules have weaker intermolecular forces.

Question 37

What is relative isotopic mass?

Answer 37

The mass of an atom of an isotope compared to 1/12th the mass of an atom of carbon-12. (Always a whole number).

Question 38

What is relative atomic mass?

Answer 38

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

Question 39

What is relative molecular mass?

Answer 39

The average mass of a molecule on a scale where carbon-12 is 12. (Used for simple molecular covalent substances)

Question 40

What is relative formula mass?

Answer 40

The average mass of a formula unit on a scale where carbon-12 is 12. (Used for giant structures)

Question 41

What does a mass spectrum show?

Answer 41

A mass spectrum is a graph which shows the results of mass spectrometry. It plots % abundance against mass to charge ratio, so can be used to work out the mass of isotopes of an element, or their abundances. A mass spectrum can be used to calculate the relative atomic mass of a sample.

Question 42

What are the stages of mass spectrometry?

Answer 42

1. Vapourisation
2. Ionisation
3. Acceleration
4. Deflection
5. Detection
6. Spectrum plotting

Question 43

How are particles accelerated in a mass spectrometer?

Answer 43

An electric field interacts with the charged particles, causing them to accelerate.

Question 44

How are particles deflected in a mass spectrometer?

Answer 44

A magnetic field is at right angles to the electric field. The charged particles interact with the magnetic field and change direction. The amount that they are deflected by indicates their mass.

Question 45

Why are there multiple peaks on a mass spectrum of chlorine?

Answer 45

The two peaks at 35 and 37 represent the isotopic masses. Their height is in a 3:1 ratio. There are 3 more peaks at 70, 72 and 74, representing molecules of CL2 with different combinations of the isotopes. There is a 9/16 chance of getting 2 Cl-35 atoms, a 6/16 chance of getting a Cl-35 and a Cl-37 atom, and a 1/16 chance of getting two Cl-37 atoms. So the ratio of the peak heights is 9:6:1.

Question 46

When is the oxidation state of a substance zero?

Answer 46

When the substance is neutral (e.g. Na, MgCl2, H2)

Question 47

When is the oxidation state of a substance not zero?

Answer 47

When the substance has an overall charge (e.g. Li+, CO32-)

Question 48

What is the oxidation state of H+ and O2-?

Answer 48

+1 and -2

Question 49

What is the oxidation state of carbon in CO32-?

Answer 49

C + (3x-2) = -2
C + -6 = -2
C = 4