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IB chemistry > topic 2/12- atomic structure/ionisation > Flashcards

topic 2/12- atomic structure/ionisation Flashcards

1
Q

define relative atomic mass, Ar

A

the average mass of the isotopes in a naturally occurring sample of an element, compared to 1/12th the mass of a carbon-12 atom

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2
Q

what do we use to determine the Ar of an element from its isotopic composition?

A

mass spectrometer

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3
Q

what is another name for protons and neutrons?

A

nucleons

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4
Q

what do A and Z stand for in:

A
X (element)
Z

A

A= mass number (protons + neutrons)
Z= atomic number (protons)

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5
Q

Describe how colour varies along the EM spectrum

A

High energy/frequency = PURPLE end of spectrum
Low energy/frequency= RED end of spectrum

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6
Q

Describe the relationship between frequency/energy and wavelength

A

As frequency/energy increases, wavelength decreases

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7
Q

Describe the difference between a line and a continuous spectrum (2)

A

a continuous spectrum contains all frequencies/energies
a line spectrum contains only discrete frequencies/energies

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8
Q

When are emission spectra produced?

A

When photons are emitted by atoms as excited electrons return to a lower energy level.

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9
Q

When are absorption spectra produced?

A

When photons are absorbed by atoms as excited electrons move to a higher energy level.

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10
Q

How does the line emission spectrum of hydrogen provide evidence for the existence of electrons in discrete energy levels?

A

The lines on the emission spectrum are separate up to the convergence point, suggesting electrons can only exist at specific energy levels.

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11
Q

Describe the first three spectral series.

A

Lyman- down to/up from n=1- UV
Balmer- down to/up from n=2- Visible Light
Paschen- down to/up from n=3- Near Infrared

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12
Q

Describe the relationship between energy of transition and colour for the hydrogen emission spectrum

A

The energy of the transition is proportional to the frequency of the emitted light.

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13
Q

State the 4 main points of Bohr’s theory.

A
  1. Electrons can only exist at specific energy levels (shells)
  2. A photon of light is emitted/absorbed when electrons move between energy levels
  3. The energy of the photon equals the difference between the energy levels
  4. E=hxf so the frequency of the emitted light is proportional to its energy.
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14
Q

What happens to the discrete energy levels of electrons at higher energies?

A

They converge

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15
Q

How does convergence affect the emission spectrum?

A

It leads to the lines in the emission spectrum becoming continuous

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16
Q

State the sub-levels of the main energy levels in order of increasing energy

A

s, p, d, and f

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17
Q

What is the number of orbitals and electrons in the s sub-level?

A

1 orbital = max 2 electrons

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18
Q

What is the number of orbitals and electrons in the p sub-level?

A

3 orbitals = max. 6 electrons

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19
Q

What is the number of orbitals and electrons in the d sub-level?

A

5 orbitals= max.10 electrons

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20
Q

What is the number of orbitals and electrons in the f sub-level?

A

7 orbitals= max 14 electrons

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21
Q

What is the shape of an s orbital?

A

a sphere

22
Q

What is the shape of a p orbital?

A

Dumbbell shaped

23
Q

What is an orbital?

A

A region of space where there is a high probability of finding up to 2 electrons.

24
Q

Describe the sub levels of n=1

A

1s

25
Q

Describe the sub levels of n=2

A

2s, 2p

26
Q

Describe the sub levels of n=3

A

3s, 3p, 3d

27
Q

Describe the sub levels of n=4

A

4s, 4p, 4d, 4f

28
Q

How many electrons can each orbital hold?

A

2 electrons with opposite spins

29
Q

What defines the energy of an electron in an orbital?

A

The electron configuration of the atom and its chemical environment

30
Q

State the Aufbau principle and give an example

A

Orbitals fill up in order of increasing energy (4s filled before 3d as it has less energy)

31
Q

State Hund’s rule

A

Orbitals in a sub-level are first occupied singly by electrons to keep them as far apart in the atom as possible. Only when every orbital is singly occupied do the electrons pair up in orbitals

32
Q

State the Pauli exclusion principle (3)

A
  • each atomic orbital can hold max.2e-
  • e- have a spin, or spinning motion, in one of two directions (at same rate), clockwise/anticlockwise
  • e- can only occupy the same orbital if they have opposite, or paired spins.
33
Q

Describe the configuration of the last 2 energy sub-levels in Chromium, Cr

A

Cr 4s13d5

34
Q

Describe the configuration of the last 2 energy sub-levels in Copper, Cu

A

Cu 4s13d10

35
Q

What is the name given to the point at which a line emission spectrum becomes continuous?

A

The convergence limit

36
Q

What does the limit of convergence correspond to?

A

First ionisation energy

37
Q

State the equation linking energy and frequency.

A

Energy = frequency x Planck’s constant

38
Q

What is the value for Planck’s constant?

A

6.63 x 10-34 Js

39
Q

Define first ionisation energy

A

The minimum energy required to remove one electron from each atom in a mole of gaseous atoms.

40
Q

Define second ionisation energy

A

The minimum energy required to remove one electron from each ion in a mole of gaseous singly charged positive ions.

41
Q

What is the significance for trends in first ionisation energy across periods?

A

It accounts for the existence of main energy levels and sub-levels in atoms.

42
Q

Explain the trend in first ionisation energy down a group.

A

IE1 decreases- there is more shielding and more distance as you go down a group

43
Q

Explain the trend in first ionisation energy down a period

A

IE1 increases- there is an increase in nuclear charge, shielding stays the same, distance from the nucleus only decreases slightly.

44
Q

Explain the change in first ionisation energy between group 18 and group 1

A

Decreases- shielding increases, distance from nucleus increases

45
Q

Explain the change in first ionisation energy between group 2 and 3

A

Decrease:
- the valent electron electron in group 3 is in a p orbital and is shielded by the s-electrons in the outer shell.
- less energy needed to remove the outer electron in group 3

46
Q

Explain the change in first ionisation energy between group 5 and 6

A

decrease:
- in g6 the extra electrons goes into a half-full p orbital
- the 2 spin paired electrons repel each other slightly
- the repulsion slightly outweighs the nuclear charge.

47
Q

how does a mass spectrometer work?

A

higher mass=deflected less
lower mass=deflected more
higher charge=deflected more
lower charge= deflected less

48
Q

describe how a mass spectrometer works

A
  1. vaporisation- gaseous atoms/molecules produced
  2. ionisation- atoms bombarded by high-energy electrons, forming a positively charged species
    X (g) + e- -> X+ (g) + 2e-
  3. acceleration- positive ions attracted to negatively charged plates ad accelerated in the electric field
  4. deflection- positive ions deflected by a magnetic field perpendicular to their path. The degree of deflection depends on the mass-to-charge ratio
  5. detection- detector detects species of a particular ratio. the ions hit the counter and an electrical signal is generated
49
Q

factors affecting 1st ionisation energy

A
  1. nuclear charge
  2. distance from nucleus
  3. extent of shielding
50
Q

differences and similarities between emission and absorption spectrum of hydrogen

A

same- both due to electrons changing energy levels and moving through shells
different- absorption has continuous spectrum background and is up the shells, emission has black background and is down the shells

51
Q

what are degenerate orbitals?

A

orbitals in the same sub-shell that have the same energy

IB chemistry (29 decks)

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