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Atomic Structure And Periodic Structure Flashcards

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1
Q

What are atoms

A

All the stuff elements and compounds are made of

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2
Q

What are the subatomic particles that make up a atom

A

Electron neutron and protons

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3
Q

What charge do electrons have

A

-1

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4
Q

What takes up the most volume in a atom

A

The orbitals

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5
Q

Where is most of the mass concentrated in the atom

A

The nucleus

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6
Q

Is the diameter of the nucleus big or small compared to the diameter of the atom

A

Small

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7
Q

What’s in the nucleus

A

Protons neutrons

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8
Q

What is the relative mass of the proton

A

1

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9
Q

What is the relative mass of the neutron

A

1

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10
Q

What is the relative mass of the electron

A

0.0005

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11
Q

What is relative charge of proton

A

+1

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12
Q

What is the relative charge of the neutron

A

0

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13
Q

What is the relative charge of a electron

A

-1

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14
Q

What is mass number

A

Total number of protons and neutrons in the nucleus

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15
Q

What is atomic mass

A

Number of protons in the nucleus

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16
Q

What is true about protons of an element

A

All atoms of the same element have the same number of protons

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17
Q

If a atom is neutral what is true about the protons and electrons

A

The are equal

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18
Q

How do you work out the number of neutrons

A

Mass number minus atomic number

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19
Q

What does a negative charge on atom mean

A

More electrons than protons

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20
Q

What does positive charge on atom mean

A

More protons than electrons

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21
Q

What are isotopes

A

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

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22
Q

What decides the chemical properties of an element

A

The number and arrangement of electrons

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23
Q

why do isotopes of the same element have the same chemical properties have the same chemical properties

A

because they have the same arrangement and number of electrons which determines the chemical properties

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24
Q

why do isotopes of an element have different physical properties

A

because physical properties tend to depend on the mass of an atom

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25
Q

what do physical properties depend on

A

mass of the atoms that make it up

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26
Q

what is the relative atomic mass

A

weighted mean mass of an atom compared to 1/12th of the mass of an atom of carbon 12

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27
Q

what is the relative isotopic mass

A

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

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28
Q

what is the relative molecular mass or relative formula mass

A

the average mass of a molecule or formula unit compare to 1/12th of the mass of an atom of carbon 12

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29
Q

how do you work out the relative molecular mass

A

add up the relative atomic mass values

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30
Q

what is relative molecular mass used for

A

simple molecules

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31
Q

what is relative formula mass used for

A

used for compounds that are ionic or giant covalent

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32
Q

how do you work out the relative formula mass

A

add the relative atomic masses of the ions or atoms

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33
Q

how to work out the relative atomic mass of an element from its isotopic abundance

A

step 1: multiply each relative isotopic mass by its % relative isotopic abundance and add up the results
step 2: divide by 100

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34
Q

where are mass spectra produced from

A

mass spectrometers

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35
Q

what are mass spectrometers used for

A

devices used to find out what samples are made up of by measuring the masses of their components

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36
Q

what can mass spectra tell us

A

relative isotopic masses and abundances of different elements

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37
Q

what is on the y axis and x axis on a mass spectra

A

abundance of ions on the y axis
and m/z values on the x axis

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38
Q

what is a m/z values

A

mass/charge ratio

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39
Q

what can you assume about the m/z value

A

charge of ions is +1 so we can assume x axis is the relative isotopic mass

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40
Q

how to work out relative atomic mass from a graph

A

step 1:multiply each relative isotopic mass by its relative isotopic abundance and add up the results
step 2:divide by the sum of the isotopic abundances

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41
Q

Silicon can exist in three isotopes. 92.23% of silicon is 28Si and 4.67% of silicon is 29Si. Given that the Ar of silicon is 28.1, calculate the abundance and isotopic mass of the third isotope.

A

100% – 92.23% – 4.67% = 3.10%
28.1 = ((28 × 92.23) + (29 × 4.67) + (X × 3.10)) ÷ 100
28.1 = (2717.87 + (X × 3.10)) ÷ 100
2810 – 2717.87 = X × 3.10
29.719 = X
So the isotopic mass of the third isotope is 30

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42
Q

Chlorine has two isotopes. 35Cl has an abundance of 75%
and 37Cl has an abundance of 25%. Predict the mass spectrum of Cl2.

A

step 1:
35Cl – 35Cl: 0.75 × 0.75= 0.5625
35Cl – 37Cl: 0.25 × 0.75= 0.1875
37Cl – 35Cl: 0.25 × 0.75= 0.1875
37Cl – 37Cl: 0.25 × 0.25= 0.0625
step 2: 0.1875 + 0.1875 = 0.375.
step 3:Divide all the relative abundances by the smallest relative abundance to get the smallest whole number ratio.
[(35Cl – 35Cl) (35 + 35 = 70) (0.5625 ÷ 0.0625 = 9)]
[(35Cl – 37Cl) (35 + 37 = 72) (0.375 ÷ 0.0625 = 6)]
[(37Cl – 37Cl) (37 + 37 = 74) (0.0625 ÷ 0.0625 = 1)]

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43
Q

what happens when molecules are bombarded with electrons in mass spectrometry

A

electron removed from molecule to form molecular ion m+

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44
Q

how do you molecular mass of a compound when looking at mass spectra

A

look at molecular ion peak

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45
Q

on a mass spectra which peak is the molecular ion peak

A

the one with the biggest m/z value

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46
Q

what do electrons move around the nucleus on

A

quantum shells/energy levels

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47
Q

what do shells contain

A

subshells

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48
Q

what do subshells contain

A

orbitals

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49
Q

how many electrons can orbitals hold

A

2

50
Q

how many electrons can s subshell hold

A

2

51
Q

how many electrons can p subshell hold

A

6

52
Q

how many electrons can d subshell hold

A

10

53
Q

how many electrons can f subshells hold

A

14

54
Q

what is an orbital

A

a bit of space an electron moves in

55
Q

what is true about the energy of orbitals within the same subshells

A

same energy

56
Q

what is spin pairing

A

when electrons within the same orbital have to spin in the opposite direction

57
Q

what shape do s orbitals have

A

spherical

58
Q

what shape do p orbitals have

A

dumbell shape

59
Q

how many p orbitals are there and what direction are they to each other

A

3 p orbitals
and 90 degrees to each other

60
Q

how can you represent electrons in orbitals

A

arrows in in boxes

61
Q

how do electrons fill up subshells

A

fill up the lowest energies first

62
Q

why does the 4s subshell fill up before the 3d subshell

A

becuase the 4s subshell has a lower energy level than the 3d subshell despite having a bigger quantumn number

63
Q

how do electrons fill up orbitals

A

fill up orbitals singly before they pair up

64
Q

electron configuration of argon /[Ar]

A

[Ar]1s²2s²2p⁶3s²3p⁶

65
Q

what is the outer shell configuration of an s block element

A

s¹ or s²

66
Q

what is the outer shell configuration of a p block element

A

s²p¹ to s²p⁶

67
Q

what do chromium and copper do that other d Block elements don’t do

A

donate one of their 4s electrons to the 3d sub shell

68
Q

why do chromium and copper donate one of their 4s electrons to the 3d sub shell

A

because they’re mores stable with a full or half full subshell

69
Q

what is electromagnetic radiation

A

energy transmitted as waves with a spectrum of different frequencies

70
Q

what is the order of the electromagnetic spectrum in increasing energy/frequency and decreasing wavelength

A

radio waves, microwaves, infrared, visible light, ultraviolet, x rays, gamma rays

71
Q

from radio waves to gamma rays what happens in terms of frequency / energy

A

increasing frequency / energy

72
Q

what happens in terms of wavelength from radio waves to gamma rays

A

decreasing wavelength

73
Q

mnemonic for em spectrum for increasing energy

A

rich men in vegas use expensive girls

74
Q

what happens when an electron is in ground state

A

electrons are on the lowest possible energy level

75
Q

what can an electrons do if they take in energy

A

move to higher energy levels

76
Q

what is it called when a electron moves to higher energy level

A

excited state

77
Q

how do electrons release energy

A

by dropping down from higher to lower energy level

78
Q

energy levels are discrete because

A

the energy levels have fixed values

79
Q

what do emission spectrum show

A

shows the frequencies of light emitted when electrons drop down from higher energy level to a lower one

80
Q

what do emission spectra look like

A

frequencies Emitted appear as coloured lines on a dark background

81
Q

why are line spectra unique for each element

A

each element has different electron arrangement so frequencies of radiation absorbed and emitted are different

82
Q

why is there sets of lines in emission spectra

A

because each set represents electrons moving to different energy levels

83
Q

in emission spectra why do they lines get closer together

A

because frequency increases

84
Q

what are the four principles when it coms to electron shells

A

*electrons can only exist in fixed orbitals or shells not anywhere between
*when an electron moves between shells electromagnetic radiation is emitted or absorbed
*each shell has fixed energy
*because energy of shell is fixed the radiation will have fixed frequency

85
Q

what supports the idea that energy levels are discrete

A

atom has clear lines for different energy levels

86
Q

describe how the electron gets to one energy level to the other

A

jumps from one to another with no in-between stage

87
Q

what provide evidence that electron exist in quantum shells

A

*emission spectra

88
Q

what is ionisation

A

the removal of one or more electrons

89
Q

what is the First ionisation energy

A

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

90
Q

what symbol must you always use when writing ionisation equations

A

gas (g)

91
Q

is ionisation an endothermic process or a exothermic process and why

A

endothermic process because energy is put in to ionise

92
Q

what are the three factors that affect ionisation energy

A

nuclear charge
electron shell
shielding

93
Q

how does nuclear charge affect ionisation energy

A

the more protons there are in the nucleus the more positively charged the nucleus and the stronger the attraction for the electrons

94
Q

how does electron shell affect the ionisation energy

A

Attraction falls off very rapidly with distance. An electron in an electron shell close to the nucleus will be much more strongly attracted than one in a shell further away

95
Q

how does shielding affect the ionisation energy

A

Attraction falls off very rapidly with distance. An electron in an electron shell close to the nucleus will be much more strongly attracted than one in a shell further away.

96
Q

what does a high ionisation energy mean

A

means there’s a strong attraction between the electron and the nucleus, so more energy is needed to overcome the attraction and remove the electron.

97
Q

what happens to ionisation energy as you go down the group

A

As you go down a group in the periodic table, ionisation energies generally fall, i.e. it gets easier to remove outer electrons.

98
Q

why does ionisation energy decrease down a group

A

*Elements further down a group have extra electron shells compared to ones above. The extra shells mean that the atomic radius is larger, so the outer electrons are further away from the nucleus, which greatly reduces their attraction to the nucleus.
*The extra inner shells shield the outer electrons from the attraction of the nucleus.

99
Q

what is the second ionisation energy

A

The second ionisation energy is the energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions.

100
Q

what does a graph of successive ionisation energy provide evidence for

A

A graph of successive ionisation energies provides evidence for the shell structure of atoms.

101
Q

what does decrease in ionisation energy provide evidence for

A

A decrease in ionisation energy going down a group provides evidence that electron shells really exist.

102
Q

why do successive ionisation energies increase

A

This is because electrons are being removed from an increasingly positive ion — there’s less repulsion amongst the remaining electrons, so they’re held more strongly by the nucleus.

103
Q

what are the big jumps on a successive ionisation energy graph

A

The big jumps in ionisation energy happen when a new shell is broken into — an electron is being removed from a shell closer to the nucleus.
2 electrons x from 1st shell.

104
Q

who was the first scientist to put the elements in a meaningful order

A

Dmitri Mendeleev

105
Q

how is the periodic table arranged

A

periods (rows) and groups (columns).

106
Q

what is true about the elements in a period

A

All the elements within a period have the same number of electron shells

107
Q

what is periodicity

A

repeating trends in the physical and chemical properties of the elements
across each period

108
Q

All the elements within a group have the same number of electrons in their outer shell what does this mean

A

This means they have similar chemical properties.

109
Q

how many outer shell electrons do s block elements have

A

The s-block elements (Groups 1 and 2) have 1 or 2 1s outer shell electrons

110
Q

what happens to atomic radius across a period

A

Atomic Radius Decreases across a Period

111
Q

why does atomic radius decrease across a period

A

As the number of protons increases, the positive charge of the nucleus increases. This means electrons are pulled closer to the nucleus, making the atomic radius smaller.

112
Q

what happens to ionisation energy across a period

A

Ionisation Energy Increases Across a Period

113
Q

why does ionisation energy increase across a period

A

because the number of protons is increasing, which means a stronger nuclear attraction extra electrons are roughly at the same level and so there’s little extra shielding affects or extra distance to lessen the attraction

114
Q

what is true about removing electrons from higher energy sub shells compared to lower energy sub shells

A

Generally, it requires more energy to remove an electron from a higher energy subshell than a lower energy
subshell

115
Q

why are there do drops between group 2 and 3 on first ionisation energy graphs

A

sub shell structure

116
Q

why is there drops between group 5 and 6 on first ionisation graph

A

electron repulsion

117
Q

what is true about singly filled full and partially filled sub shells

A

In general, elements with singly filled or full subshells are more stable than
those with partially filled subshells, so have higher first ionisation energies.

118
Q

what happens to the boiling points in metals

A

For the metals (Li, Be, Na, Mg and Al), melting and boiling points increase across the period because the metallic bonds get stronger. The bonds get stronger because the metal ions have an increasing number of delocalised electrons and a decreasing radius
(i.e. the metal ions have a higher charge density). This means there’s a stronger attraction between the metal ions and delocalised electrons, so stronger metallic bonding.

119
Q

what happens to boiling points of giant covalent lattices

A

The elements with giant covalent lattice structures (C and Si) have strong covalent bonds linking all their atoms together.
A lot of energy is needed to break all of these bonds. So, for example, carbon (as graphite or diamond) and silicon have the highest boiling points in their periods.

120
Q

what happens to the boiling points of simple molecular structures

A

Next come the simple molecular structures (N2, O2 and F2, P4, S8 and Cl2). Their melting points depend upon the strength of the London forces between their molecules. London forces are weak and easily overcome, so these elements have low melting and boiling points.More electrons in a molecule mean stronger London forces . For example, in Period 3 a molecule of sulfur (S8) has the most electrons, so it’s got higher melting and boiling points than phosphorus and chlorine.

121
Q

what happens to boiling points of noble gases

A

The noble gases (Ne and Ar) have the lowest melting and boiling points in their periods because they exist as individual atoms (they’re monatomic) resulting in very weak London forces.

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