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A-LEVEL CHEMISTRY > A1 Atomic Structure + Bonding > Flashcards

A1 Atomic Structure + Bonding Flashcards

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1
Q

Define an isotope

A
  • Atoms of the same element
  • With different numbers of neutrons and different masses
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2
Q

Define Relative Atomic Mass

A
  • The weighted mean mass of an atom of an element
  • Compared to 1/12th the mass of an atom of Carbon-12
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3
Q

Define Relative Isotopic Mass

A
  • The mass of an atom of an isotope
  • Compared to 1/12th the mass of an atom of Carbon-12
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4
Q

Define Weighted Mean Mass

A
  • The mean mass
  • Taking into account the relative abundancies of the isotopes
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5
Q

What does the atomic number of an atom tell you

A

Number of protons

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6
Q

What does the mass number of an atom tell you

A

The number of neutrons and protons

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7
Q

How do you calculate the RFM given abundances and relative atomic masses

A
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8
Q

On a mass spectrometer, what do the peaks indicate

A

Isotopes

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9
Q

What happens to all elements when in a mass spectrometer

A

Develop a 1+ charge

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10
Q

How to calculate the RFM from a mass spectrometer, where the relative abundances don’t equal 100

A
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11
Q

What peaks should be shown on mass spectrometers with diatomic elements

A

All possible combinations of diatomic molecules, and all monotomic atoms

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12
Q

How many electrons occupy the first shell

A

2

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13
Q

How many electrons occupy the 2nd shell

A

8

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14
Q

How many electrons occupy the 3rd shell

A

18

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15
Q

How many electrons occupy the 4th shell

A

32

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16
Q

How many orbitals in s subshell

A

1

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17
Q

How many orbitals in p subshell

A

3

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18
Q

How many orbitals in d subshell

A

5

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19
Q

How many electrons in ANY orbital

A

2

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20
Q

Define an orbital

A
  • A 3D reigon of space around the nucleus, that can hold up to 2 electrons
  • With opposite spins
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21
Q

What is the shape of an s orbital

A
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22
Q

What is the shape of a p orbital

A
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23
Q

What is the exception to the rule of subshells filling with increasing energy

A

4s fills before 3d

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24
Q

How do electrons fill sub-shell orbitals as you go along the period

A

Singularly, before doubling up in each orbital

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25
Q

Why is an element in the x block

A

As it’s highest energy electrons occupies the x subshell

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26
Q

When writing electronic configurations, how should you write it when the 3d subshell is full

A

Write it before 4s

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27
Q

What electrons are removed when positive ions form from d block elements

A

4s are removed before 3d

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28
Q

Electronic configuration for chronium

A
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29
Q

Electronic configuration for copper

A
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30
Q

Define Ionic Bonding

A
  • The strong electrostatic attraction
  • Between positive and negative ions
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31
Q

Define an Ionic Lattice

A

A repeating pattern of oppositely charged ions

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32
Q

When drawing ionic lattices, what must be included in each circle

A

Both the formula and the charge

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33
Q

Why do ionic compounds have high melting/boiling points

A
  • Giant ionic lattice structure
  • Has lots of very strong ionic bonds
  • Lots of energy is required to break all the strong ionic bonds
34
Q

What ions are always soluble

A
  • Na+
  • K+
  • NH4+
  • NO3-
35
Q

State and explain the conductivity of solid ionic compounds

A
  • Does not conduct
  • Ions are fixed in position
  • Not mobile and free to move
  • Around giant ionic grid-like lattice structure
36
Q

State and explain the conductivity of aqueous and molten ionic compounds

A
  • Does conduct
  • Ions aren’t fixed in position
  • Ions are mobile and free to move
  • Around the giant ionic grid-like lattice stucture
37
Q

Define covalent bonding

A
  • The strong electrostatic attraction
  • Between a shared pair of electrons and the nuclei of bonded atoms
38
Q

Define a Dative Covalent Bond

A

A covalent bond where both electrons are donated by one atom

39
Q

What is used to measure the strength of a covalent bond

A

Average bond enthalpy

40
Q

Why does diamond/silicon have a very high melting/boiling point

A
  • Each carbon atom makes 4 strong covalent bonds with other carbon/silicon atoms
  • Lots of energy is needed to break all the strong covalent bonds
  • In the giant tetrahedral covalent lattice
41
Q

Why does graphite/graphene have a very high melting/boiling point

A
  • Each carbon atom makes 3 strong covalent bonds with other carbon atoms
  • Lots of energy is needed to break all the strong covalent bonds
  • In the giant covalent lattice structure
42
Q

State and explain the conductivity of diamond/silicon

A
  • Cannot conduct
  • Each carbon/silicon atom makes 4 strong covalent bonds with other carbon/silicon atoms
  • There are no delocalised electrons, or ions mobile and free to move
  • Around the giant tetrahedral covalent lattice strucutre
43
Q

State and explain the conductivity of graphite/graphene

A
  • Does conduct
  • Each carbon/silicon atoms makes three strong covalent bonds with other carbon atoms
  • There is one delocalised electron per atom mobile and free to move
  • Around the giant covalent lattice structure
44
Q

Why is graphite soft

A
  • Layers of carbon atoms are held together by weak forces of attraction
  • Layers are able to slide over each other
45
Q

Which giant covalent structures are soluble

A

None of them

46
Q

Why do simple molecular substances have a low boiling point

A
  • Molecules held together by weak intermolecular forces
  • Not a lot of energy is needed to overcome the weak intermolecular forces that act between molecules
47
Q

State and explain the conductivity of simple molecular substances

A
  • None conduct
  • As all the molecules are neutral
  • As theres no delocalised electrons or ions present
48
Q

Define Metallic Bonding

A
  • Strong electrostatic attraction
  • Between a lattice of cations and a sea of negatively charged delocalised electrons
49
Q

Draw a metallic lattice

A
50
Q

Why do metallic compounds have high melting/boiling points

A
  • There are lots of strong metallic bonds
  • In the giant metallic lattice structure
  • Lots of energy is required to break all the strong metallic bonds
51
Q

State and explain the conductivity of metals in all states

A
  • Conducts in all states
  • As the sea of delocalised electrons are mobile and free to move
  • Around the giant metallic lattice structure
52
Q

When are metallic substances soluble

A

Never

53
Q

Draw a linear molecule

A
54
Q

Draw a trigonal planar molecule

A
55
Q

Draw a tetrahedral molecule

A
56
Q

Draw a trigonal bipyramid molecule

A
57
Q

Draw a octahedral molecule

A
58
Q

Draw a trigonal pyrimid molecule (NH3)

A
59
Q

Draw a bent molecule (H2O)

A
60
Q

What are the relative repulsions of lone electron pairs and bonded electron pairs

A

Lone pairs repel more than bonded pairs

61
Q

How to write an answer, comparing bond angles

A
  • State the number of bonded and lone pairs of each molecule
  • Lone pairs repel more strongly than bonded pairs
  • Bonded electron pairs repel equally
62
Q

Define electronegativity

A
  • The ability of an atom to attract the bonding electrons
  • In a covalent bond
63
Q

What is the most electronegative element

A

Flourine

64
Q

How to determine what element is more electronegative

A

Whatever is closest to flourine in the periodic table

65
Q

In a covalent bond, which atom do the electrons move closer to

A

The more electronegative atom

66
Q

How to explain if this molecule is or isn’t polar

A
  • Has polar bonds
  • Molecule is symmetrical
  • Dipoles cancel out
  • Non-polar
67
Q

How to explain if this molecule is or isn’t polar

A
  • Has polar bonds
  • Molecule isn’t symmetrical
  • Dipoles don’t cancel out
  • Polar
68
Q

Where do all intermolecular forces act

A

Between molecules

69
Q

In what molecules do London forces occur

A

In every molecule (NOT giant covalent)

70
Q

Why as molecules get larger, do their boiling points increase

A
  • More electrons
  • Stronger London forces
  • More energy required to break London forces
71
Q

How are London forces induced

A
  • Electrons move randomly in molecule
  • Creates a temporary dipole in the molecule
  • Induces temporary dipoles in neighbouring molecules
72
Q

Describe a permanent dipole force

A
  • Polar molecules have dipoles
  • Dipoles interact to form the dipole-dipole force
73
Q

Between what molecules does hydrogen bonding occur

A
  • One lone pair of flourine, oxygen or nitrogen
  • Hydrogen of another molecule
74
Q

What is the relative strength of the different intermolecular forces

A
  • Hydrogen bonding strongest
  • Permenant dipole-dipole forces
  • London forces
75
Q

Draw a hydrogen bonding diagram for two water molecules

A
76
Q

What must be included in all hydrogen bonding diagrams

A
  • Lone pairs
  • Dipoles
  • Hydrogen bond
77
Q

Describe and explain the anomolous properties of ice - relativly high melting point

A
  • Hydrogen bonding is very strong
  • So lots of energy needed to overcome it
78
Q

Describe and explain the anomolous properties of ice - ice is less dense than water

A
  • Water molecules held apart in an open lattice structure
  • By hydrogen bonds
79
Q

State and explain the solubility of non-polar substances in non-polar solvents

A
  • Soluble
  • IMF’s form between molecules in solvent and molecules in solute
  • Weakens the IMF’s in the simple molecular solvent
80
Q

State and explain the solubility of polar substances in non-polar solvents

A
  • Insoluble
  • Attraction between molecules in solvent and ions in ionic lattice
  • Not strong enough to break ionic bonds in ionic lattice
81
Q

State and explain the solubility of polar substances in polar solvents

A
  • Soluble
  • Polar bonds in solute
  • Attract polar bonds in solvent

A-LEVEL CHEMISTRY (11 decks)

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