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Y11 Chem 2024 > Atomic Structure > Flashcards

Atomic Structure Flashcards

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1
Q

Subshell

A

A specific energy level within an electron shell. Designated s, p, d or f in order of increasing energy

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2
Q

Order of subshells in increasing energy

A

s, p, d, f

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3
Q

Orbital

A

A region of space in which up to two electrons may be located

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4
Q

How many orbitals in a p subshell?

A

3

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5
Q

How many electrons can fit in a d subshell?

A

10

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6
Q

How many orbitals in an s subshell?

A

1

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7
Q

How many electrons would be in an f subshell if all the orbitals were half-filled?

A

7

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8
Q

Order of subshell filling up to 6p

A

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p

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9
Q

Pauli Exclusion Principle

A

A maximum of two electrons can fit into any orbital, providing that those electrons have opposite spin

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10
Q

Hunds Rule

A

Electrons will arrange themselves into a subshell in such a way as to maximise the number of half-filled orbitals. OR Electrons will half-fill all orbitals in a subshell before any are filled completely.

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11
Q

Aufbau principle

A

Electrons will move into subshells in order from lowest energy up to highest energy

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12
Q

Ground state

A

A situation where all electrons in an atom or ion are in the lowest possible energy levels (subshells)

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13
Q

Excited state

A

A situation where not all electrons in an atom or ion are in the lowest possible energy levels (subshells)

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14
Q

Which two elements in the first row of the d block have only one 4s electron?

A

Chromium and Copper

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15
Q

What causes a spectral line on an emission spectrum?

A

An excited electron relapsing into a lower energy subshell

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16
Q

What shape is an s orbital?

A

Sphere

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17
Q

How many electrons can fit into the 3rd shell, and how many of these electrons fit into each subshell?

A

A total of 18 - 2 in the s subshell, 6 in the p subshell and 10 in the d subshell

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18
Q

What is meant by the term ‘isotopes’?

A

Two atoms that have the same number of protons (atomic number) but a different number of neutrons (or different mass number)

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19
Q

From which subshell would an iron atom lose its first electron?

A

The 4s subshell

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20
Q

Trend in atomic radius across a period (L to R), and why

A

Atomic radius decreases because core charge increases but the number of shells remains the same

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21
Q

Trend in atomic radius down a group, and why

A

Atomic radius increases because the number of shells increases but core charge remains the same

22
Q

Trend in core charge across a period (L to R), and why

A

Core charge increases because the atomic number increases but the number of inner shell electrons remains the same

23
Q

Trend in core charge down a group, and why

A

Core charge remains the same because the atomic number and the number of inner shell electrons increase by the same amount

24
Q

How to calculate core charge

A

Number of protons in the nucleus minus the number of inner shell electrons

25
Q

Core charge of a Silicon atom

A

4+ (14 protons minus 10 inner shell electrons)

26
Q

Core charge of a Magnesium 2+ ion

A

10+ (12 protons minus 2 inner shell electrons)

27
Q

Core charge of a Sulfide ion

A

6+ (16 protons minus 10 inner shell electrons)

28
Q

Trend in electronegativity across a period (L to R), and why

A

Electronegativity increases because core charge increases and atomic radius decreases slightly (note that noble gases are an exception as they do not easily participate in bonding)

29
Q

Trend in electronegativity down a group, and why

A

Electronegativity decreases because atomic radius increases (due to increasing number of electron shells) but core charge remains constant

30
Q

Definition of ‘first ionisation energy’

A

The amount of energy required to remove an electron from each of a mole of gaseous atoms (of that element) to produce a mole of 1+ ions

31
Q

Trend in first ionisation energy across a period (L to R) and why

A

First Ionisation Energy increases due to increasing core charge and slight decrease in atomic radius

32
Q

Trend in first ionisation energy down a group and why

A

First Ionisation Energy decreases due to increasing atomic radius (due to increasing number of electron shells) but core charge remains constant

33
Q

Define Electronegativity

A

The force of attraction experienced by bonding electrons towards the nucleus of that element

34
Q

Which group has the highest electronegativity in any period after the first?

A

Group 7 (the halogens)

35
Q

Which group has the lowest electronegativity in any period after the first?

A

Group 1 (the alkali metals)

36
Q

Which group has the highest first ionisation energy in any period after the first?

A

Group 8/0 (the noble gases)

37
Q

Which group has the lowest first ionisation energy in any period after the first?

A

Group 1 (the alkali metals)

38
Q

How can you work out the group number of an element using successive ionisation energies?

A

The number of electrons removed before the first big ‘jump’ in ionisation energies is the number of valence electrons (and therefore group number)

39
Q

Why is the third ionisation energy of calcium much greater than the second?

A

The third electron is removed from a lower-energy shell than the first two. This electron will be closer to the nucleus and experience much less electron shielding.

40
Q

What is electron shielding?

A

A repulsive force exerted by electrons that are closer to the nucleus (ie in lower energy shells) on electrons that are further from the nucleus (ie in higher-energy shells), pushing them away from the nucleus

41
Q

Trend in the reactivity of metals down groups 1 and 2, and why

A

Reactivity increases because electrons are lost more easily, because the atomic radius increases but core charge remains the same

42
Q

What is common about metals that are found in the native state?

A

They are very unreactive. Examples include gold, silver and platinum.

43
Q

Trend in the reactivity of non-metals down group 7, and why

A

Reactivity decreases because it becomes harder for the halogen to attract an electron due to increasing atomic radius (more electron shielding) while core charge remains the same

44
Q

Definition of relative isotopic mass

A

Mass of that isotope relative to 1/12 the mass of a carbon-12 isotope

45
Q

Definition of relative atomic mass

A

The weighted average mass of all the isotopes of that element, weighted by their relative abundance

46
Q

Formula to calculate moles from mass and molar mass

A

n = m/Mr

47
Q

Formula to calculate moles from concentration and volume of a solution

A

n = CV

48
Q

Formula to calculate moles from volume and molar volume of a gas (at SLC)

A

n = V/Vm (or V/24.8)

49
Q

Formula to calculate number of particles (N) from moles and Avogadro’s number (6.02 x 10^23)

A

N = n x 6.02 x 10^23

50
Q

How many moles of hydrogen atoms are present in 2.5 moles of water?

A

5 moles of hydrogen atoms

51
Q

How many moles of atoms are present in 2.5 moles of water?

A

7.5 moles of atoms (2.5 moles of oxygen atoms plus 5 moles of hydrogen atoms)

Y11 Chem 2024 (13 decks)

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