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A Level Chemistry Year 1 > Electronic Structure > Flashcards

Electronic Structure Flashcards

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1
Q

What is the number given to each shell called?

A

Principal quantum number

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2
Q

What can be said about the energy and principal quantum number when a shell gets further away from the nucleus?

A

The higher it’s energy and the larger it’s principal quantum number is

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3
Q

What are shells divided into?

A

Sub-shells

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4
Q

What are the 4 types of sub-shell?

A

S
P
D
F

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5
Q

What is an orbital?

A

A region of a sub-shell that contains a maximum of 2 electrons with opposite spins

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6
Q

What is spin?

A

A type of momentum possessed by an electron which can either be up or down

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7
Q

How many electrons can s sub-shells hold?

A

2

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8
Q

How many electrons can p sub-shells hold?

A

6

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9
Q

How many electrons can d sub-shells hold?

A

10

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10
Q

How many electrons can f sub-shells hold?

A

14

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11
Q

What can be said about the energy of orbitals in the same sub-shell?

A

They have the same energy

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12
Q

What is spin-pairing?

A

When there are 2 electrons in an orbital which spin in opposite directions

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13
Q

What shape are s-orbitals?

A

Spherical

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14
Q

What shape are p-orbitals?

A

Dumbbell shape

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15
Q

What is the formula for the number of electrons a shell can hold?

A

2n^2 (n being shell number)

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16
Q

What is electron configuration?

A

The number of electrons and how they’re arranged

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17
Q

Why do 4s sub-shells fill first before 3d sub-shells?

A

4s sub-shells have a lower energy level

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18
Q

What did Werner Heisenberg came up with?

A

The theory that you cannot know the speed of an electron or it’s exact position

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19
Q

What was Heisenberg’s theory called?

A

Uncertainty principle

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20
Q

How do ions form?

A

When atoms lose or gain electrons to become positively or negatively charged

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21
Q

What is electrostatic attraction?

A

A strong force between positive and negative ions

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22
Q

What is an ionic bond?

A

An electrostatic attraction between 2 oppositely charged ions

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23
Q

What is an ionic compound?

A

When oppositely charged ions form an ionic bond

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24
Q

What are ionic crystals?

A

Giant lattice (regular structure) of ions

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25
Q

How does a lattice form?

A

When each ion is electrostatically attracted in all directions to ions of the opposite charge

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26
Q

Describe the electrical conductivity of ionic compounds

A

Conduct electricity when molten or dissolved, but not when solid

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27
Q

Why do ionic compounds not conduct electricity when solid?

A

There are no mobile ions as they’re fixed in position by strong metallic bonds

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28
Q

Why do ionic compounds conduct electricity when molten or dissolved?

A

They have ions that are free to move (mobile ions) so can carry charge

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29
Q

Describe the melting and boiling points of ionic compounds?

A

Have high melting and boiling points

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30
Q

Why do ionic compounds have high melting and boiling points?

A

The giant ionic lattices are held together by strong electrostatic forces (which takes lots of energy to overcome)

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31
Q

Describe the solubility of ionic compounds

A

Tend to dissolve in water

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32
Q

Why do ionic compounds tend to dissolve in water?

A

Water molecules are polar, so the water molecules pull the ions away from the lattice and cause it to dissolve

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33
Q

What does polar molecule mean?

A

Part of the molecule has a negative charge, the other bits have a positive charge

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34
Q

How are molecules formed?

A

When 2 or more atoms bond together

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35
Q

What are molecules held together by?

A

Strong covalent bonds

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36
Q

What is a covalent bond?

A

The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

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37
Q

Give an example of a molecule with triple bonds

A

Nitrogen (N2)

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38
Q

Give an example of a molecule with double bonds

A

Oxygen (O2)

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39
Q

What is different between a single covalent bond and a dative covalent bond?

A

In a single covalent bond, atoms share electrons with the electrons coming from each atom. Dative covalent bonds shares electrons but both electrons come from the same atom

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40
Q

How can you show a dative bond on a diagram?

A

By an arrow, pointing away from the donor atom

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41
Q

What is different about boron trifluoride?

A

The boron only has 6 electrons on its outer shell

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42
Q

What does expanding the octet mean?

A

Having an outer shell with more than 8 electrons

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43
Q

Give an example of a molecule that expands the octet

A

Sulfur Hexafluoride, has 12 electrons in it’s outer shell

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44
Q

What does enthalpy mean?

A

Another word for energy

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45
Q

What does the average bond enthalpy tell you?

A

The strength of a covalent bond

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46
Q

What does average bond enthalpy measure?

A

The average energy needed to break a bond

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47
Q

What does the shape of a molecule depend on?

A

The number of pairs of electrons in the outer shell of the central atom

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48
Q

What are shared electrons called?

A

Bonding pairs

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49
Q

What are unshared electrons called?

A

Lone pairs

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50
Q

What do pairs of electrons do?

A

Repel each other (as they’re both negatively charged)

51
Q

Which repels more, lone pairs or bonding pairs?

A

Lone pairs

52
Q

Why do lone pairs repel more?

A

Being closer to the central atom causes lone-pairs to take up more of the available ‘bonding space’

53
Q

What is electron pair repulsion theory?

A

That the greatest bond angles are between lone pairs of electrons, and bond angles between bonding pairs are reduced because they’re pushed together by lone pair repulsion

54
Q

When drawing shapes of molecules, what does a broken line (e.g. |||||) show?

A

A bond pointing away from you

55
Q

When drawing shapes of molecules, what does a wedge line show?

A

A bond pointing towards you

56
Q

When drawing shapes of molecules, what does a line show?

A

A bond that isn’t pointing towards or away from you

57
Q

What is the bond angle of a linear shape?

A

180°

58
Q

What electron pairs does a linear shape have?

A

2 bonding pairs, 0 lone pairs

59
Q

Why do linear molecules have a bond angle of 180°?

A

The pairs of bonding electrons repel each other, so are positioned as far away as possible from each other

60
Q

What electron pairs does a trigonal planar have?

A

3 bonding pairs, 0 lone pairs

61
Q

What is the bond angle of a trigonal planar?

A

120°

62
Q

What electron pairs does a non-linear molecule (with 3 electron pairs) have?

A

2 bonding pairs, 1 lone pair

63
Q

What is the bond angle of a non-linear with 3 electron pairs?

A

A bit less than 120°

64
Q

What electron pairs does a tetrahedral have?

A

4 bonding pairs, 0 lone pairs

65
Q

What is the bond angle of a tetrahedral?

A

109.5°

66
Q

What electron pairs does a trigonal pyramidal have?

A

3 bonding pairs, 1 lone pair

67
Q

What is the bond angle of a trigonal pyramidal?

A

107°

68
Q

What electron pairs does a non-linear (with 4 electron pairs) have?

A

2 bonding pairs, 2 lone pairs

69
Q

What is the bond angle of a non-linear with 4 electron pairs?

A

104.5

70
Q

What is a molecule with 5 bonding pairs called?

A

Trigonal bipyramidal

71
Q

What are the bond angles in a trigonal bipyramidal?

A

90° and 120°

72
Q

Why do trigonal bipyramidal molecules have 2 bond angles?

A

Because 3 of the atoms will form a trigonal planar shape (120°) and the other 2 atoms will be at 90°

73
Q

What is a molecule with 6 bonding pairs called?

A

Octahedral

74
Q

What are all the bond angles in an octahedral?

A

90°

75
Q

What types of bonds can only be purely covalent?

A

Only bonds between atoms of a single element (e.g. diatonic gases)

76
Q

What is electronegativity?

A

The ability of an atom to attract the bonding electrons in a covalent bond

77
Q

What is the general trend in electronegativity?

A

Increases as you move towards Fluorine

78
Q

What is the trend of electronegativity down a group?

A

It decreases

79
Q

What is the trend of electronegativity across a period?

A

It increases

80
Q

What does a higher electronegativity number mean?

A

Means that element is better able to attract the bonding electrons

81
Q

What are the covalent bonds in diatonic gases?

A

Non-polar

82
Q

What does non-polar mean?

A

Where the atoms have equal electronegativities and so the electrons are equally attracted to both nuclei

83
Q

What is a polar bond?

A

A covalent bond where a difference in electronegativity has caused a shift in electron density in the bond

84
Q

In a covalent bond between 2 atoms with different electronegativities, where are the bonding electrons pulled to?

A

The most electronegative atom

85
Q

In a polar bond, what does the difference in electronegativity between the atoms cause?

A

A permanent dipole

86
Q

What is a dipole?

A

A difference in charge between the 2 atoms caused by a shift in electron density in the bond

87
Q

Between 2 atoms, the greater the difference in electronegativity…

A

… The more polar the bond

88
Q

What do polar molecules have?

A

An overall dipole

89
Q

What determines whether a polar molecule has an overall dipole?

A

The arrangement of polar bonds

90
Q

Why does CO2 have no overall dipole?

A

Because it is a linear molecule, the shape of the molecule is symmetrical. This causes the dipoles to cancel each other out, so it will be non-polar and have no overall dipole

91
Q

For a greater difference in electronegativity, what will the difference in type of bond be?

A

The less covalent and the more ionic it will be

92
Q

In general, what will the bonds be if the difference in electronegativity is less than 0.4?

A

Non-polar covalent bonds

93
Q

In general, what will the bonds be if the difference in electronegativity is between 0.4 and 2.0?

A

Polar covalent bonds

94
Q

In general, what will the bonds be if the difference in electronegativity is greater than 2.0?

A

Ionic

95
Q

What are intermolecular forces?

A

Forces between molecules

96
Q

How strong are intermolecular forces compared to covalent, ionic or metallic bonds?

A

Much weaker

97
Q

What are the 3 types of intermolecular bonding?

A

Permanent dipole-dipole interactions
Induced dipole-dipole interactions
Hydrogen bonding

98
Q

Which is the strongest of the 3 types of intermolecular bonding?

A

Hydrogen bonding

99
Q

What does induced dipole-dipole interactions cause?

A

All atoms and molecules to be attracted to each other

100
Q

What cause a temporary dipole?

A

When the electrons in an atom are to more side than the other (increasing the charge on that side)

101
Q

What does a temporary dipole cause?

A

Another temporary induced dipole in the opposite direction of a neighbouring atom (which causes them to attract)

102
Q

What trend in strength of induced dipole-dipole interactions occurs the larger the molecule?

A

Larger the molecule, the larger the electron clouds meaning stronger induced dipole-dipole interactions

103
Q

Why do molecules with a greater surface area have stronger induced dipole-dipole interactions?

A

The have more of the electron cloud exposed

104
Q

What can induced dipole-dipole interactions also be called?

A

London forces

105
Q

What are permanent dipole-dipole interactions?

A

Where the ẟ+ and the ẟ- charges on polar molecules cause weak electrostatic forces of attraction between molecules

106
Q

What is Hydrogen bonding?

A

A bond that only occurs when Hydrogen is covalently bonded to fluorine, nitrogen or oxygen

107
Q

What is charge density?

A

A measure of how much positive or negative charge there are in a certain volume

108
Q

What is the charge density of Hydrogen like?

A

Very high as it is small

109
Q

What do molecules with Hydrogen bonding normally contain?

A

-OH or -NH groups

110
Q

What are the properties of molecules with Hydrogen bonding?

A

Soluble in water

High melting and boiling points

111
Q

Why do water and ammonia have very high melting points compared to the other hydrides in their group?

A

Because more energy is needed to break the Hydrogen bonds

112
Q

Why do icebergs float on water?

A

Ice is less dense than water

113
Q

How are water molecules in ice held together?

A

By Hydrogen bonds

114
Q

Why is ice less dense than water?

A

Ice has Hydrogen bonds. Since Hydrogen bonds are long, the molecules in ice will be further apart than in liquid (making it less dense)

115
Q

When drawing Hydrogen bonds, where does the ẟ+ go?

A

Goes on the H atom

116
Q

When drawing Hydrogen bonds, where does the ẟ- go?

A

On the electronegative atom

117
Q

What are the bonds and forces like in simple covalent compounds?

A

Strong bonds within compounds, but weak forces between the molecules

118
Q

What is the electrical conductivity of simple covalent compounds?

A

They don’t conduct electricity (as they’re overall uncharged)

119
Q

What is the melting and boiling points of simple covalent compounds like?

A

Melting and boiling points are low, as the intermolecular forces that hold together the molecules are weak, so don’t need much energy to break

120
Q

When melting or boiling and simple covalent compounds, what bonds/forces do you have to break?

A

Have to break the intermolecular forces that hold together the molecules in a compound (But not the covalent bonds that hold the atoms together in a molecule)

121
Q

What is the trend in melting and boiling points as intermolecular forces get stronger?

A

Melting and boiling points increase

122
Q

What is the solubility of simple covalent compounds with Hydrogen bonds?

A

Compounds with Hydrogen bonds can form Hydrogen bonds with water molecules, so will be soluble

123
Q

What is the solubility of simple covalent compounds that are non-polar?

A

Insoluble

124
Q

What are the 2 factors, as you go down the Group 7 Hydrides, that affect the strength of the intermolecular forces?

A

The polarity of the molecules decrease, so the strength of the permanent dipole-dipole interactions decreases
The number of electrons in the molecule increases, so the strength of the permanent dipole-dipole interactions increase

A Level Chemistry Year 1 (9 decks)

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