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A-Level Chemistry Edexcel > 1: Atomic Structure and the Periodic Table > Flashcards

1: Atomic Structure and the Periodic Table Flashcards

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1
Q

Explain the structure of an atom in terms of electrons, protons and neutrons.

A

Protons and neutrons contained in the nucleus.
Electrons in shells around the nucleus.

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2
Q

Give the relative mass of protons, neutrons and electrons

A

Proton: 1
Neutron: 1
Electron: 1/1840

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3
Q

Give the relative charge of protons, neutrons and electrons

A

Proton: +1
Neutron: None
Electron: -1

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4
Q

If an atom has an atomic number = Z and a mass number = M, determine the number of each type of sub-atomic particle.

A

Proton: Z
Neutron: M - Z
Electron: Z

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5
Q

Explain the term ‘isotopes’.

A

Atoms with the same atomic number, but different mass numbers.

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6
Q

Define atomic number

A

The number of protons in the nucleus of an atom

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7
Q

Define mass number

A

the total number of protons and neutrons in the nucleus of an atom

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8
Q

Define relative isotopic mass

A

The mass of an atom of a particular isotope of an element, relative to 1/12th the mass of a carbon-12 atom.

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9
Q

Define relative atomic mass

A

The weighted average mass of all isotopes of an element, relative to 1/12th the mass of a carbon-12 atom.

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10
Q

What is relative molecular mass?

A

The average mass of a molecule of a substance, relative to 1/12th the mass of a carbon-12 atom.

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11
Q

What is relative formula mass?

A

The sum of the atomic masses of all the atoms in a particular formula.
- Used for compounds with giant structures.

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12
Q

Mass spectrometry

What happens in the ionisation stage?

A

Electrons from the electron gun collide with the particles and form charged ions.

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13
Q

Mass spectrometry

What happens in the acceleration stage?

A

An electric field accelerates the particles to a uniform speed.

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14
Q

Mass spectrometry

What happens in the deflection stage?

A

A magnetic field deflect the particles.
Particles with less mass and more charge are deflected more.
An electromagnet varies the strength of the field.

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15
Q

Mass spectrometry

What happens in the detection stage?

A

The ions create an electric current which is detected.

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16
Q

Define 1st ionisation energy.

A

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

17
Q

Define successive ionisation energies.

A

The energy required to
remove 1 electron from each ion
in 1 mole of gaseous ions
to form 1 mole of gaseous (2+, 3+, 4+, …) ions

18
Q

How does the number of protons influence ionisation energy?

A

More protons -> stronger attraction -> higher IE

19
Q

How does electron shielding influence ionisation energy?

A

More shielding -> weaker attraction -> lower IE

20
Q

How does the electron subshell from which the electron is removed influence ionisation energy?

A

Greater distance -> weaker attraction -> lower IE

21
Q

State and explain the general trend in first ionisation energy as you go across a period.

A

Nuclear charge increases,
Shielding increases slightly,
Atomic radius decreases,
Strength of attraction between nucleus and outer electron increases,
so first ionisation energy increases.

22
Q

State and explain the trend in first ionisation energy as you go down a group.

A

Nuclear charge increases,
Shielding increases,
Atomic radius increases,
Strength of attraction between nucleus and outer electron decreases,
so first ionisation energy decreases.

23
Q

What ideas provide evidence for the existence of quantum shells?

A
  • Atomic emission spectra
  • Successive ionisation energies
  • First ionisation energies of successive elements
24
Q

What do the atomic emission spectra provide evidence for?

A

The existence of quantum shells.

25
Q

What do successive ionisation energies provide evidence for?

A

The existence of quantum shells and the group to which the element belongs.

26
Q

What do the first ionisation energies of successive elements provide evidence for?

A

Electron sub-shells.

27
Q

How many electrons can fill the first four quantum shells?

A

1: 2
2: 8
3: 18
4: 32

28
Q

What is an orbital?

A

A region within an atom that can hold up to 2 electrons with opposite spins.

29
Q

What is the shape of an s orbital?

A

2 electrons in an s subshell, therefore 1 orbital in the shape of a sphere.

30
Q

What is the shape of a p orbital?

A

6 electrons in a p subshell, therefore 3 orbitals

31
Q

What is the order that the subshells are filled and the maximum number of electrons in each subshell?

A 
1s - 2
2s - 2
2p - 6
3s - 2
3p - 6
4s - 2
3d - 10
4p - 6
32
Q

What rules describe how electrons fill subshells and orbitals?

A

Electrons enter the lowest energy orbital available.
Electrons fill subshells singly, before pairing up.
Two electrons in the same orbital must have opposite spins.

33
Q

What is the electronic configuration of copper in 1s notation? (group 11 period 4)

A

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

34
Q

What is the electronic configuration of chromium in 1s notation? (group 6 period 4)

A

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

35
Q

What are the four blocks of the periodic table?

A
36
Q

What determines the chemical properties of an element?

A

Electronic configuration

37
Q

Define periodicity.

A

A repeating trend in across different periods.

38
Q

State and explain the trend in the melting and boiling points from sodium to aluminium.

A

Ioinic radius decreases,
Charge of ions increases,
So number of delocalised electrons increases,
So attraction between ions and electrons increases,
So melting and boiling point increases.

39
Q

Explain why silicon has a high melting and boiling point compared to other period 3 atoms.

A

Silicon forms a giant covalent structure,
Held together by strong covalent bonds.

Other non-metals form simple molecular structures held together by weak intermolecular forces.

A-Level Chemistry Edexcel (34 decks)

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