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a-level oxford chem > 3.1.1 Atomic Structure > Flashcards

3.1.1 Atomic Structure Flashcards

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1
Q

_____ take up most of the volume in atoms

A

Orbtitals take up most of the volume in atoms

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2
Q

Relative mass for an electron

A

1/1840

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3
Q

What letter represents the mass number?

A

A

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4
Q

What letter represents the atomic number?

A

Z

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5
Q

What type of ions have…

No. of electrons < No. of protons

A

Postive Ions

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6
Q

What type of ions have…

No. of electrons > No. of protons

A

Negative Ions

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7
Q

What holds the protons and neutrons?

A

Strong nuclear force

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8
Q

What holds electrons and protons together in atom?

A

Electrostatic forces of attraction

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9
Q

Why is the strong nuclear force stronger than electrostatic forces?

A

It overcomes repulsion between protons in nucleus

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10
Q

Strong nuclear force acts only over very ___ distances

A

SHORT distances (within nucleus)

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11
Q

What decides the chemical properties of an element?

A

No. & arrangement of electrons decides

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12
Q

Why do isotopes have the same chemical properties? (2)

A
  • ∵ they have same electron configuration
  • chemical properties depend on electrons
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13
Q

Isotopes have ___ _____ physical properties

A

slight different physical properties

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14
Q

Why do isotopes have slight different physical properties?

A

∵ physical properties depend on mass of atom

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15
Q

19th century: What did John Dalton say atoms were?

A
  • Solid spheres
  • Different spheres made different elements
  • (All atoms of an element = same mass)
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16
Q

1897: What did J.J. Thomson discover and what did it show?

A
  • Discovered the electron
  • Showed atoms weren’t solid and indivisible
  • (Model known as ‘plum pudding model’)
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17
Q

1909 - Ernest Rutherford: What did he find out?

A

Conducted the golden foil experiment:

  • Fired positively charged alpha particles at a very thin sheet of gold
  • Particles passed straight through gold & only small no. of particles were deflected backwards (pulm pudding model said = alpha particles would be deflected by the positive ‘pudding’ in atom)
  • = developed into nuclear model of atom
    • Tiny positive nucleus surrounded by ‘cloud’ of negative electrons - most of atom is empty space
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18
Q

What was Niels Bohr’s model & discovery?

A
  • Model: where electrons exist in shells or orbits of fixed energy
  • Discovered: When electrons move between shells, electromagnetic radiation (with fixed energy or frequency) is emitted/absorbed
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19
Q

What have modern day scientists discovered & so what did they do?

A
  • Electrons in same shell ≠ same energy
  • Bohr model = wrong ∴ they refined it & added sub-shells
  • (Isn’t perfect model but it’s simple and explains many experimental observations e.g. bonding & ionisation energy trends)
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20
Q

What are relative masses essentially?

A

Masses of atoms compared to carbon-12

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21
Q

Define Relative Atomic Mass (Ar) of an element (1x)

A

Average mass of an atom of an element on a scale where an atom of carbon-12 is 12

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22
Q

Define Relative Isotopic Mass

A

Mass of an atom of an isotope of an element on a scale where an atom of carbon-12 is 12

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23
Q

Define Relative Molecular Mass (Mr)

A

Average mass of a molecule on a scale where an atom of carbon-12 is 12

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24
Q

What does a mass spectrometer do and how?

A
  • It determines the mass of separate atoms (or molecules)
  • Works by forming ions from sample and then separating them according to the ratio of their charge to their mass
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25
Q

Name the 6 things that happen when a sample is squirted into time of flight (TOF) mass spectrometer

A
  1. Vacuum
  2. Ionisation
  3. Acceleration
  4. Ion Drift
  5. Detection
  6. Data Analysis
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26
Q

Describe the step vacuum in mass spectrometry (TOF)

A

Whole apparatus is kept under a vacuum to prevent ions produced from colliding with air molecules

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27
Q

Name the two ways you can ionise your sample in mass spectrometry (TOF)

A

2 methods:

  1. Electrospray ionisation
  2. Electron impact ionisation
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28
Q

Describe the method electrospray ionisation

A
  1. A high voltage is applied to a sample in a polar solvent
  2. Sample molecule, M, gains a proton forming MH+
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29
Q

Describe the method electron impact ionisation

A
  1. Sample is bombarded by high energy electrons
  2. Sample molecule loses an electron = become +1 ions (M+)
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30
Q

Describe the step acceleration in mass spectrometry (TOF)

A

Positively charged ions are accelerated by an electric field (attracted to negatively charged plate) so = they all have same kinetic energy

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31
Q

Describe the step ion drift in mass spectrometry (TOF)

A
  • Ions enter region with no electric field so they just drift through it
  • Lighter ions will drift faster than heavier ions
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32
Q

Describe the step detection in mass spectrometry (TOF) & state how abundance is measured

A
  • Lighter ions travel at higher speeds = reach detector in less time than heavier ions
  • Positive ions collected at detector
  • Causing current to flow / detected electrically
  • Abundance measured: idea that current depends on number of ions hitting detector
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33
Q

Describe the step data analysis in mass spectrometry (TOF)

A

Signal from detector is sent to a computer which generates a mass spectrum

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34
Q

What does the y-axis of mass spectrum represent?

A

Abundance of ions

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35
Q

What does the height of each peak give on the mass spectrum?

A

Relative isotopic abundance

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36
Q

If the sample is an element, what does each line represent on the mass spectrum?

A

A different isotope of the element

37
Q

What does the x-axis on the mass spectrum represent?

A

‘mass/charge’ ratio (m/z)

38
Q

Describe how to work out the relative atomic mass from mass spectrum (4)

A
  1. Spectrum gives relative abundance (of isotopes) & m/z (mass/charge ratio)
  2. Multiply m/z by relative abundance for each isotope
  3. Sum these values
  4. Divide by the sum of the relative iostopic abundances
39
Q

Why do elements with isotopes produce more than one line in a mass spectrum?

A

∵ isotopes = different masses

40
Q

Describe how you can use mass spectrometry to identify elements

A

You can see if the sample being analysed has the same relative abundances of isotopes

41
Q

Explain how you use mass spectrometry to identify molecules

A

mass/charge ratio (of peak) = relative molecular mass of molecule

42
Q

Electrons have ____ ______ & move around nucleus in certain regions of atom called ____________

A

Electrons have fixed energies & move around nucleus in certain regions of atom called shells/energy levels

43
Q

Each shell is given a number called ____ ____ _____

A

principal quantum number

44
Q

What is the principal quantum number?

A

2(n2)

45
Q

The further away a shell is from nucleus, the _____ its energy & the ____ its principal quantum number

A

higher its energy + larger its principal quantum number

46
Q

Electrons in same the shell ___ have same energy

A

DON’T

47
Q

Shells divided up into sub-shells which have ____ ______ energies

A

slightly different

48
Q

Sub-shells have different no. of orbitals which can hold up to ___ electrons

A

2

49
Q
A
50
Q

2 electrons in each orbital…

A

spin in opposite directions

51
Q

Electrons fill up ___ energy sub-shell 1st

A

Electrons fill up lowest energy sub-shell 1st

52
Q

Why do electrons fill orbitals singly before they start sharing?

A

∵ electrons repel each other

53
Q

Electron configuration: Give 2 examples of transition metals behaving unusually

A

Chromium (Cr) & copper (Cu) = donate 1 of their 4s electrons to 3d sub-shell

54
Q

Write the electron configuration for chromium

A
55
Q

Write the electron configuration for copper

A
56
Q

Electron configuration: what happens when transition metals become an ion?

A

They lose 4s electrons before their 3d electrons

57
Q

Groups 4-7 can _______ electrons when they form _____ ____

A

Groups 4-7 can share electrons when they form covalent bonds

58
Q

Why are the gases in Group 0 inert?

A

∵ completely filled s & p sub-shells

59
Q

Why does Chromium (Cr) & copper (Cu) donate 1 of their 4s electrons to 3d sub-shell?

A

∵ they’re happier with a more stable full or half-full d sub-shell

60
Q

Define first ionisation energy

A

Enthalpy change when 1 mole of gaseous 1+ ions is formed from 1 mole of gaseous atoms

61
Q

Ionisation is a ________ process ∵ you have to put energy in to ionise atom/molecule

A

endothermic

62
Q

Write an equation for the first ionisation of oxygen

A
63
Q

Name 3 rules about ionisation energies

A
  1. Must use gas state symbol (g) ∵ ionisation energies are measured for gaseous atoms
  2. Always refer to 1 mole of atoms
  3. Lower ionisation energy = easier it is to remove from ion
64
Q

Name 3 factors that affect ionisation energy

A
  1. Nuclear Charge
  2. Shielding
  3. Distance from Nucleus
65
Q

Describe how nuclear charge affects ionisation energy

A

More protons in nucleus = more positively charged nucleus is & stronger the attraction for electrons

66
Q

Describe how shielding affects ionisation energy

A

As no. of electrons between outer electrons & nucleus increases = outer electrons feel less attraction towards nuclear charge

Lessening of pull of nucleus by inner shells is called shielding (or screening)

67
Q

Describe how distance from nucleus affects ionisation energy

A

Attraction decreases rapidly with distance

(i.e. electron close to nucleus = much more strongly attracted than one further away)

68
Q

What is meant by high ionisation energy?

A

High ionisation energy = high attraction between electrons & nucleus = more energy needed to remove electron

69
Q

What provides evidence for shells structure of atoms?

A

Graph of successive ionisation energies

70
Q

Within each shell, successive ionisation energies ______

A

increase

71
Q

Why does successive ionisation energies increase within each shell?

A

∵ electrons are being removed from an increasingly positive ion = less repulsion amongst remaining electrons ∴ they’re held more strongly by nucleus

72
Q

When does big jumps in ionisation energy happen?

A

When a new shell is broken into = an electron is being removed from shell closer to nucleus

73
Q

Define second ionisation energy

A

Enthalpy change when 1 mole of gaseous 2+ ions is formed from 1 mole of gaseous 1+ ions

74
Q

State the equation for the second ionisation of oxygen

A
75
Q

State the equation for the nth ionisation energy

A
76
Q

Describe how you can use a successive ionisation energies graph to figure out which group an element belongs to

A

Count how many electrons are removed before the 1st big jump to find the group number

77
Q

Describe how you can use a successive ionisation energies graph to predict the electronic structure of elements

A

Working from right to left, count no. of points there are before each big jump to find how many electrons there are in each shell, starting with the first

78
Q

Name 2 trends in first ionisation energy

A
  • 1st ionisation energies of elements down a group of periodic table decrease
  • 1st ionisation energies of elements across a period generally increase
79
Q

Explain why ionisation energy decreases down Group 2

A
  1. Atomic radius increases/electron removed further from nucleus
  2. As group is descended more shielding = nucleus’ attraction reduces

Both of these factors = make it easier to remove outer electrons = lower ionisation energy

80
Q

Explain why ionisation energy increases across a Period (2x)

A
  1. Increased nuclear charge (no. of protons is increases = stronger nuclear attraction)
  2. Extra electrons enter roughly same energy level or similar shielding
81
Q

Drops between Groups __ and ___ show _____ Structure

A

Drops between Groups 2 and 3 show Sub-Shell Structure

82
Q

Describe and explain how aluminium provides evidence for the theory of electron sub-shells

A
  1. Aluminium’s outer electron = in 3p orbital rather than 3s
    • ∵ 3p orbital = slightly higher energy than 3s orbital
      • ∴ electron is found f_urther from nucleus_
  2. Additonal electron shielding - 3p orbital has additional shielding provided by 3s2 electrons
  3. Both these factors strong enough to override effect of increased nuclear charge = ionisation energy drops slightly, easier to remove electron
83
Q

Drops between Groups __ and __ Is due to Electron ______

A

Drops between Groups 5 and 6 Is due to Electron Repulsion

84
Q

Describe and explain how phosphorus & sulfur provides more evidence for the eletronic structure model

A
  1. (Shielding identical in phosphorus & sulfur atoms + electron is being removed from an identical orbital)
  2. In phosphorus’s case: electron being removed from singly-occupied orbital
  3. But in sulfur: electron removed from paired electrons in 3p orbital
  4. Electron repulsion between 2 electrons = electron easier to remove from pair
85
Q

Explain why the value of the first ionisation energy of neon is higher than that of sodium (2x)

A
  • Electron removed from a level of lower energy or e– removed from 2p rather than from 3s
  • Less shielding
86
Q

Write the electron configuration for calcium using noble gas symbols

A
87
Q

Which of Na+ and Mg2+ is the smaller ion? Explain why (2)

A

Mg2+

Has more protons with same sheilding

88
Q

Magnesium exists as three isotopes: 24Mg, 25Mg & 26Mg

A

24Mg percentage = 80%

26Mg percentage = 10%

a-level oxford chem (34 decks)

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