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Chemistry > Ch2.2 Electrons, Bonding and Structure > Flashcards

Ch2.2 Electrons, Bonding and Structure Flashcards

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1
Q

How are electrons arranged in atoms?

A

In energy levels/shells

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2
Q

Where is the lowest energy level located in an atom?

A

Closest to the nucleus

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3
Q

In what order do shells become filled in an atom?

A

Shells fill outwards from the nucleus

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4
Q

How many shells are there in each atom?

A

An infinite amount (despite them possibly being empty they are still there)

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5
Q

What is a primary quantum number of a shell?

A

Each shell has a primary quantum number (n)
1st shell: n=1, 2nd shell: n=2

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6
Q

What is the maximum number of electrons that can be held in each shell?

A

2n²
(n being the primary quantum number of the shell)

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7
Q

Where do subshells exist?

A

Within the shells

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8
Q

What are the different types of subshell?

A

4 types: s,p,d,f

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9
Q

How many electrons can each type of subshell hold?

A

s-subshells hold up to 2 e-
p-subshells hold up to 6 e-
d-subshells hold up to 10 e-
f-sebshells hold up to 14 e-

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10
Q

How many of each subshell do the first four shells in an atom hold?

A

1st: one s-subshell
2nd: one s-subshell and one p-subshell
3rd: one s-subshell, one p-subshell and one d-subshell
4th: one of each subshell (spdf)

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11
Q

What does the Aufbau Principle help us to remember?

A

The order in which the subshells of an atom fill

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12
Q

What does Aufbau mean?

A

Building

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13
Q

Describe how the Aufbau principle should be used

A

Write out the primary quantum numbers in a vertical line and write the possible subshells next to them in increasing e- capacity. Draw diagonal lines going towards the bottom left of the page and follow the arrows to see which order the subshells fill in

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14
Q

What does core notation help with?

A

Saves time with writing out spdf structures

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15
Q

How does core notation work?

A

When writing out spdf structure if a shell becomes completely full you can write the subshells involved as the symbol of a noble gas eg. iron= (Ar)4s²3d⁶

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16
Q

What order do subshells fill in? (generally)

A

In order of increasing energy level

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17
Q

Are all of the energy levels of different shells linear or do they overlap?

A

They overlap

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18
Q

Where are orbitals located?

A

Within subshells which are within shells

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19
Q

How many electrons can orbitals hold?

A

maximum of 2 in each

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20
Q

How many orbitals are in an s-subshell and of what kind?

A

One s-orbital

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21
Q

How many orbitals are in an p-subshell and of what kind?

A

Three p-orbitals

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22
Q

How many orbitals are in an d-subshell and of what kind?

A

Five d-orbitals

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23
Q

How many orbitals are in an f-subshell and of what kind

A

Seven f-orbitals

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24
Q

What shape is an s-orbital?

A

Spherical

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25
Q

What shape is a p-orbital?

A

Dumbbell/hourglass shape

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26
Q

In a p-subshell, how are the three p-orbitals arranged?

A

At right angles, one on each axis (x,y and z)

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27
Q

What order do electrons fill orbitals in?

A

In order of increasing energy level

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28
Q

Which orbital would be filled first out of a vacant one and one with an electron in already?

A

Vacant orbitals fill before electrons start to pair up in orbitals as it is a lower energy option for the electron

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29
Q

What must happen when there is two electrons in the same orbital?

A

They must have opposite spins

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30
Q

What is box notation?

A

A way of showing shells, subshells and orbitals in an atom

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31
Q

How can box notation be shortened?

A

By using the previous noble gas’ symbol and finishing the box notation for all the remaining subshells
(basically write the symbol of the most recent completely full shell then add the subshells and the orbitals in afterwards)

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32
Q

What is the flame colour of burning lithium?

A

Brick red

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33
Q

What is the flame colour of burning potassium?

A

Lilac

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34
Q

What is the flame colour of burning calcium?

A

Orange/red

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35
Q

What is the flame colour of burning sodium?

A

Yellow

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36
Q

What is the flame colour of burning copper?

A

Green/blue

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37
Q

What is the flame colour of burning strontium?

A

Scarlet red

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38
Q

What is the flame colour of burning barium?

A

Bright green

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39
Q

What is the flame colour of burning magnesium?

A

Bright white

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40
Q

What was the bunsen burner invented for?

A

For flame tests

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41
Q

Do any metals burn as the same colour?

A

No there is 80 known metals and they all burn as unique colours in a bunsen burner

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42
Q

How do you prepare for a flame test?

A
  1. Take an inert substance (eg. Platinum/Pt) and dip it in HCl
  2. Put the metal into a roaring blue flame until there is no colour change (this cleans the metal and eliminates contamination)
  3. Dip the loop into the solid sample and put into the flame to observe colour
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43
Q

What did Niels Bohr use for evidence of ionisation energies?

A

Emission spectra

44
Q

What model did Niels Bohr come up with?

A

A model of the atom with a positively-charged nucleus surrounded by a series of shells in which electrons could be found

45
Q

How many shells did Bohr think there was in each atom?

A

An infinite amount

46
Q

How are spectral lines produced to be viewed as emission spectra?

A
  • if electrons absorb enough energy they could jump into a higher energy level shell
  • when they fall back to their original position (or a position with lower energy than where they jumped to) they release this energy as light of a particular frequency
  • these colours of light produce the emission spectra
47
Q

What piece of equipment is needed to view an emission spectra?

A

A spectroscope

48
Q

What is ionisation?

A

When electrons gain enough energy to leave an atom completely rather than just move up an energy level

49
Q

What are anions and cations?

A

Anions: negatively charged, gained electrons
Cations: positively charged, lost electrons

50
Q

What is the first ionisation energy? (definition)

A

The amount of energy (kJ) needed for one mole of gaseous atoms to lose one mole of electrons to form one mole of gaseous 1+ ions

51
Q

How does ionisation energy increase across a period?

A

There is a net increase of energy required to remove an electron

52
Q

What is shielding in terms of electrons in atoms?

A

Anything blocking the pull of the nucleus from something being pulled away eg. an electron

53
Q

What are successive ionisation energies?

A

The energy required to remove each electron from an atom until all of its electrons have been stripped away.
eg. 1IE is energy to remove first e-
2IE is energy to remove second e-
etc.
(Have to be in gaseous state)

54
Q

What is the common factor with successive ionisation energies of any element and why?

A

They always increase. This is because with each electron you remove, the next electron will be pulled away from a more positive species. Therefore the pull is stronger and it’s harder to remove.

55
Q

What does a big jump in successive ionisation energies indicate?

A

When there is a change in shells (move in towards next closest shell to nucleus)

56
Q

If you are given a table of successive ionisation energies and asked what group the element is in how do you work out the answer?

A

However many plots there are before the first big jump = group number

57
Q

What are the 3 types of chemical bonding in broad terms?

A
  • ionic bonding: metal and non-metal losing and gaining electrons
  • covalent bonding: non-metals sharing electrons
  • metallic bonding: metals bonding together
58
Q

What are the names of the two types of bonding which are between covalent and ionic bonding?

A
  • polarisation of the anion (swapping but with some sharing)
  • polar covalent bonding (sharing but charges develop)
59
Q

What is the definition of covalent bonding?

A

A bond formed by a shared pair of electrons between nuclei

60
Q

Examples of substances with single covalent bonds

A
  • methane: CH4
  • ammonia: NH3
61
Q

Examples of substances with double covalent bonds

A
  • oxygen: O2
  • carbon dioxide: CO2
  • alkenes (CnH2n): eg. ethene: C2H4
62
Q

Examples of substances with triple covalent bonds

A
  • nitrogen: N2
  • alkynes (CnH2n-2): eg. ethyne: C2H2
63
Q

What are dative covalent bonds also known as?

A

Co-ordinate bonds

64
Q

What is a dative covalent bond?

A

Where both bonding electrons come from only one of the atoms

65
Q

What is the checklist for a question about the shape of a molecule?

A
  1. Count the number of electron pairs (including lone pairs) around the outer shell of the central atom
  2. How many are bonding pairs vs. lone pairs?
  3. Will repulsion be equal or not? If not, why?
  4. Come up with associated shape and angle
66
Q

Which pair of electrons repels more? Bonding or lone?

A

Lone pairs repel more, they reduce other bond angles by 2.5°

67
Q

What are intermolecular forces?

A

Forces of attraction acting between molecules

68
Q

Are covalent bonds or intermolecular forces stronger?

A

Covalent bonds are stronger. Intermolecular forces are easily overcome by temperature

69
Q

What are the different types of intermolecular forces?

A
  • induced dipole-dipole interactions (London forces / dispersion forces)
  • permanent dipole-dipole interactions
  • hydrogen bonds
70
Q

What is the shape of a molecule determined by?

A

The electron pairs surrounding the central atom, generally only the ones in the outermost shell.
Electron pairs repel each other so pairs move as far away as possible from each other to minimise repulsion

71
Q

What shape is a carbon dioxide molecule?

A

Linear (straight line)

72
Q

What shape would a molecule with 3 bonding areas be?

A

‘trigonal planar’ (trigonal shape)

73
Q

What shape would a molecule with 4 bonding areas be?

A

Tetrahedral

74
Q

What shape would a molecule with 5 bonding areas be?

A

Trigonal bipyramidal

75
Q

What shape would a molecule with 6 bonding areas be?

A

Octahedral

76
Q

Give two examples of where dative covalent bonds occur

A

ammonium ion and oxonium ion

77
Q

What is the difference between inter and intra molecular forces?

A

INTERmolecular forces act between molecules and INTRAmolecular forces act within molecules

78
Q

Which is stronger, inter or intra molecular forces? Give examples

A

INTRAmolecular forces are stronger eg. ionic, covalent and metallic bonding
INTERmolecular forces are weaker eg. hydrogen bonds

79
Q

Rank the 3 types of intermolecular forces in order from strongest to weakest

A
  1. hydrogen bonds
  2. permanent dipole-dipole forces
  3. induced dipole-dipole forces
80
Q

Where do hydrogen bonds form?

A

Between molecules in which either fluorine, oxygen of nitrogen (FON) is directly attached to hydrogen atoms

81
Q

Why are hydrogen bonds the strongest type of intermolecular force?

A

They involve fluorine, oxygen or nitrogen which are the most electronegative elements

82
Q

Define electronegativity

A

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

83
Q

How does hydrogen bonding affect the boiling point of substances it occurs in? Give examples

A

It makes them much higher than other molecules with a similar molecular mass
eg. water Mr = 18 and liquid at room temp, methane Mr = 16 and gas at room temp

84
Q

Are substances with hydrogen bonding miscible?

A

2 substances with hydrogen bonding are miscible (they will mix) whereas non-polar substances don’t mix with polar substances eg. water and oil

85
Q

What are the anomalous properties of water?

A
  • high Mp and BP for its Mr
  • its solid form (ice) is less dense than its liquid form (water)
86
Q

Why is ice less dense than water and explain how this benefits our planet?

A

When ice freezes the hydrogen bonds lock the H2O molecules into an open framework so molecules are further apart than they are in the liquid state.
If ice sank, our oceans would freeze from the bottom upwards and our planet would have become a giant ice ball long ago

87
Q

What is a permanent dipole attraction?

A

A weak intermolecular force that arises between 2 permanently polar molecules

88
Q

Which 2 elements are most likely to form permanent dipole interactions?

A

Chlorine and bromine.
eg. In HCl, a dipole is formed on each molecule as the chlorine is electronegative and pulls the shared electrons closer to itself so the positive H and negative Cl from different molecules attract.

89
Q

Why are permanent dipole attractions not the same as hydrogen bonds?

A

Hydrogen bonds require F, O or N to be directly bonded to hydrogen for hydrogen bonds to form

90
Q

How are induced dipoles formed?

A

If a molecule with a permanent dipole approaches a molecule with no dipoles then it can induce a dipole to form temporarily

91
Q

Explain how an induced dipole can be formed between HF and HH

A

If two molecules get close enough (one polar and one not) the 𝛿⁻ end of the HF forces the electron in one of the HH atoms away. This causes an induced dipole to form

92
Q

When do induced dipoles disappear?

A

When the molecules move apart again after coming in close contact with one another

93
Q

What is the proper name for London forces/ Van Der Waal’s forces

A

Instantaneous induced dipole pairs

94
Q

What is the process of an instantaneous induced dipole bond forming?

A
  1. atoms/molecules are surrounded by clouds of electrons
  2. electrons are constantly moving
  3. at an instant there will be more electrons on one side of the molecule than the other
  4. at that point an instantaneous dipole forms with a 𝛿⁻ end and a 𝛿⁺ end
  5. in the brief time it exists, the dipole can induce a dipole in any close neighbouring atoms/molecules
95
Q

Where are instantaneous induced dipole attractions present?

A

In every substance

96
Q

What is the mnemonic to remember the order of substances produced in fractional distillation from light to heavy and the actual names?

A

Rich Girls Never Kiss Dirty Boys
- refinery gases
- gasoline
- naphtha
- kerosene
- diesel
- bitumen

97
Q

What is an isomer?

A

Compounds that have the same molecular formula but different structural formula

98
Q

How is the shape of a molecule determined?

A

VSEPR: valence shell electron pair repulsion theory
(simply, regions of negative charge in the outer shells of atoms repel eachother so they are as far apart as possible)

99
Q

Give an example of a molecule with 2 regions of negative charge and state its shape and bond angle

A

Carbon dioxide (CO2), 180°, linear shape

100
Q

Give an example of a molecule with 3 regions of negative charge and state its shape and bond angle

A

Boron trifluoride (BF3), 120°, trigonal planar shape

101
Q

Give an example of a molecule with 4 bonding pairs and state its shape and bond angle

A

Methane (CH4), 109.5°, tetrahedral shape

102
Q

Give an example of a molecule with 3 bonding pairs and 1 lone pair and state its shape and bond angle

A

Ammonia (NH3), 107°, pyramidal shape

103
Q

Give an example of a molecule with 2 bonding pairs and 2 lone pairs and state its shape and bond angle

A

Water (H2O), 104.5°, non-linear shape

104
Q

Give an example of a molecule with 5 bonding pairs and state its shape and bond angle

A

Phosphorus pentachloride (PCl5), 90° and 120°, trigonal bipyramidal shape

105
Q

Give an example of a molecule with 6 bonding pairs and state its shape and bond angle

A

Sulphur hexafluoride (SF6), 90°, octahedral shape

106
Q

Define ionic bonding

A

The bond formed between a positively charged and negatively charged ion due to the electrostatic forces of attraction between them

107
Q

Define metallic bonding

A

The attraction between metal ions and delocalised electrons

Chemistry (17 decks)

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