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Chemistry A level (Edexcel) > TOPIC 1 - Atomic structure and the Periodic Table > Flashcards

TOPIC 1 - Atomic structure and the Periodic Table Flashcards

1
Q

What is the relative mass of an electron?

A

1/2000

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2
Q

What are isotopes?

A

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

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3
Q

What is the Relative Atomic Mass?

A

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

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4
Q

What is the Relative Isotopic Mass?

A

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

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5
Q

What is the Relative Molecular Mass?

A

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

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6
Q

What is the M+ peak?

A

Molecular ion peak. Shows the relative molecular mass.

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7
Q

What is the M+1 peak?

A

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

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8
Q

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

A

1 orbital. 2 electrons.

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9
Q

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

A

3 orbitals. 6 electrons.

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10
Q

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

A

5 orbitals. 10 electrons.

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11
Q

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

A

7 orbitals. 14 electrons.

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12
Q

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

A

1s. 2 electrons.

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13
Q

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

A

2s, 2p. 8 electrons.

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14
Q

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

A

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

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15
Q

Which shape does the s orbital have?

A

Spherical.

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16
Q

Which shape does the p orbitals have?

A

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

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17
Q

What does spin pairing mean?

A

Electrons occupying the same orbital have different spins.

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18
Q

What are the different subshells from lowest energy to highest?

A

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

19
Q

What is the electron configuration of Chromium?

A

1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5

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

20
Q

What is the electron configuration of Copper?

A

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

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

21
Q

Which block is the one in red?

A

s block

22
Q

Which block is the one in purple?

A

d block

23
Q

Which block is the one in green?

A

p block

24
Q

Which block is the one in purple?

A

f block.

25
Q

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

A
  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.

26
Q

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

A

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.

27
Q

What is ionisation energy?

A

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-

28
Q

How does shielding affect ionisation energies?

A

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.

29
Q

How does nuclear charge affect ionisation energies?

A

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.

30
Q

How does atomic size affect ionisation energies?

A

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.

31
Q

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

A
  • 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.
32
Q

Why does atomic radius decrease across a period?

A
  • Increased nuclear charge pulls electrons towards nucleus.
  • Although more electrons, all are in the same shell so shielding is similar.
33
Q

Why does atomic radius increase down a group?

A
  • More quantum shells, so more shielding.
  • This offsets the fact that there’s an increasing nuclear charge.
34
Q

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

A
  • Increasing number of protons increases attractive forces towards the nucleus.
  • Shielding and size of the atom are similar.
35
Q

Why is there a dip in Al?

A

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.

36
Q

Why is there a dip in S?

A

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

37
Q

Why is there an increase from Na to Al?

A
  • Increasing positive charge of ions
  • Increaseing number of delocalised electrons.
  • Therefore bonds are stronger and so more energy is needed to overcome them.
38
Q

Why is there a peak in Si?

A

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

39
Q

Why is there a decrease from Si to P?

A

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

40
Q

Why is there an increase from P to S?

A

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

41
Q

Why is there a decrease from S to Cl?

A

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

42
Q

Why is there a decrease from Cl to Ar?

A

Ar exists as individual atoms so weaker forces of attraction.

43
Q
A

Chemistry A level (Edexcel) (19 decks)

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