TOPIC 1 - Atomic structure and the Periodic Table

Question 1

What is the relative mass of an electron?

Answer 1

1/2000

Question 2

What are isotopes?

Answer 2

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

Question 3

What is the Relative Atomic Mass?

Answer 3

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

Question 4

What is the Relative Isotopic Mass?

Answer 4

Mass of an atom of an isotope compared to 1/12th of the mass of an atom of carbon-12.

Question 5

What is the Relative Molecular Mass?

Answer 5

Mean mass of a molecule compared to 1/12th of the mass of a carbon-12 atom.

Question 6

What is the M+ peak?

Answer 6

Molecular ion peak. Shows the relative molecular mass.

Question 7

What is the M+1 peak?

Answer 7

The peak that has +1 m/z than the molecular ion peak due to the existence of a carbon 13 isotope.

Question 8

How many orbitals has the ‘s’ subshell got? How many electrons?

Answer 8

1 orbital. 2 electrons.

Question 9

How many orbitals has the ‘p’ subshell got? How many electrons?

Answer 9

3 orbitals. 6 electrons.

Question 10

How many orbitals has the ‘d’ subshell got? How many electrons?

Answer 10

5 orbitals. 10 electrons.

Question 11

How many orbitals has the ‘f’ subshell got? How many electrons?

Answer 11

7 orbitals. 14 electrons.

Question 12

Which subshells has the first quantum shell got? How many electrons?

Answer 12

1s. 2 electrons.

Question 13

Which subshells has the second quantum shell got? How many electrons?

Answer 13

2s, 2p. 8 electrons.

Question 14

Which subshells has the third quantum shell got? How many electrons?

Answer 14

3s,3p,3d. 18 electrons.

Question 15

Which shape does the s orbital have?

Answer 15

Spherical.

Question 16

Which shape does the p orbitals have?

Answer 16

Dumbell shape. 3 orbitals. Px, Py and Pz.

Question 17

What does spin pairing mean?

Answer 17

Electrons occupying the same orbital have different spins.

Question 18

What are the different subshells from lowest energy to highest?

Answer 18

1s, 2s, 2p, 3s, 3p, 4s, 3d

Question 19

What is the electron configuration of Chromium?

Answer 19

1s², 2s², 2p⁶, 3s², 3p⁶, 4s¹, 3d⁵

An electron is excited from the 4s to fill the 3d orbitals.

Question 20

What is the electron configuration of Copper?

Answer 20

1s², 2s², 2p⁶, 3s², 3p⁶, 4s¹, 3d¹⁰

an electron is excited from 4s to fill the 3d subshell.

Question 21

Which block is the one in red?

Answer 21

s block

Question 22

Which block is the one in purple?

Answer 22

d block

Question 23

Which block is the one in green?

Answer 23

p block

Question 24

Which block is the one in purple?

Answer 24

f block.

Question 25

How are line spectra formed? How are they used to identify an element?

Answer 25

1. Electron is in ground state.
2. Electron absorbs energy and is excited (moves to a higher-energy shell)
3. Electron moves back to ground state and releases energy in the form of electromagnetic radiation with a specific frequency.
4. Emission spectra shows the frequency of light emitted.

This is unique for every element.

Question 26

How do line spectra prove that electrons exist in quantum shells?

Answer 26

Defined lines in the emission spectra shows that there are specific frequencies and therefore there are discrete energy values in difference between quantum shells. This proves that electrons exist in shells only.

Question 27

What is ionisation energy?

Answer 27

Minimum amount of energy required to remove 1 mole of electrons fro 1 mole of atoms in gaseous state.

i.e. Na(g) > Na+(g) + e-

Question 28

How does shielding affect ionisation energies?

Answer 28

The more electron shells between the positive nucleus and the negative outer electron that is being removed, the weaker the attraction and so less energy is needed.

Question 29

How does nuclear charge affect ionisation energies?

Answer 29

The more protons in the nucleus, the bigger the positive charge and so the bigger the attraction between nucleus and electrons therefore more energy is needed to remove an electron.

Question 30

How does atomic size affect ionisation energies?

Answer 30

The bigger the atom, the further away the outer electrons are from the nucleus and so the attractive force between them is less so its easier to remove electrons.

Question 31

Why does ionisation energies decrease as we go down the group?

Answer 31

• Atomic radius increases, so outer electrons are further away from the nucleus, so attractive force is weaker, so energy required decreases.
• Shielding increases since there are more shells between the nucleus and the electron and so the attractive force is less and so less energy is required.
• This offsets the fact that there’s an increasing charge.

Question 32

Why does atomic radius decrease across a period?

Answer 32

• Increased nuclear charge pulls electrons towards nucleus.
• Although more electrons, all are in the same shell so shielding is similar.

Question 33

Why does atomic radius increase down a group?

Answer 33

• More quantum shells, so more shielding.
• This offsets the fact that there’s an increasing nuclear charge.

Question 34

Why does ionisation energy increase as we go across a period?

Answer 34

• Increasing number of protons increases attractive forces towards the nucleus.
• Shielding and size of the atom are similar.

Question 35

Why is there a dip in Al?

Answer 35

Aluminium has its outer shell electron in a higher energy subshell, further away from the nucleus than magnesium.

Al - …3s2, 3p1

Mg - …3s2

Therefore the force of attraction of the electron with the nucleus is weaker and so less energy is needed.

Question 36

Why is there a dip in S?

Answer 36

Both P and S have lectrons in 3p sub shell so shielding is the same.

However, sulfur has an orbital with two electrons while phosphorus doesn’t.

There’s repulsion between electrons in the same orbitals so less energy is needed to remove the electron

Question 37

Why is there an increase from Na to Al?

Answer 37

• Increasing positive charge of ions
• Increaseing number of delocalised electrons.
• Therefore bonds are stronger and so more energy is needed to overcome them.

Question 38

Why is there a peak in Si?

Answer 38

Giant covalent (macromolecular) structure. Many strong covalent bonds hold silicon together so a large amount of energy is required.

Question 39

Why is there a decrease from Si to P?

Answer 39

P₄ is a simple molecular formula so less energy is needed to overcome the intermolecular forces than in the giant covalent in Si.

Question 40

Why is there an increase from P to S?

Answer 40

S₈ larger simple molecular structure so london forces are stronger than in P₄ so more energy is needed to overcome intermolecular forces.

Question 41

Why is there a decrease from S to Cl?

Answer 41

Cl₂ smaller molecular structure so decrease london forces so weaker intermolecular forces and so less energy is needed.

Question 42

Why is there a decrease from Cl to Ar?

Answer 42

Ar exists as individual atoms so weaker forces of attraction.