Inorganic Chemistry > Topic 2 - Bonding and Structure > Flashcards

Topic 2 - Bonding and Structure Flashcards

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1
Q

What is an ion?

A

A positively or negatively charged atom (or group of atoms).

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2
Q

How are ions formed?

A

When electrons are transferred from one atom to another.

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3
Q

What are positive ions called?

A

Cations

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4
Q

What are negative ions called?

A

Anions

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5
Q

How do single atoms form ions?

A

They gain or lose 1, 2 or 3 electrons so that they have a full outer shell.

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6
Q

Is the charge on a metal ion positive or negative?

A

Positive

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7
Q

Is the charge on a non-metal ion positive or negative?

A

Negative

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8
Q

What is an ionic bond?

A

The strong electrostatic attraction between two oppositely charged ions.

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9
Q

What is the force responsible for ionic bonding?

A

Electrostatic attraction

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10
Q

What does the fomula of an ionic compound tell you?

A

The ratio of the ions in that compound.

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11
Q

How does the electrostatic attraction relate to the strength of an ionic bond?

A

The stronger the attraction, the strongere the bond.

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12
Q

What two factors affect the strength of an ionic bond?

A
  • Ionic charges

* Ionic radius

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13
Q

How does ionic charge affect the strength of an ionic bond?

A

The greater the charge, the stronger the bond.

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14
Q

How does ionic radius affect the strength of an ionic bond and why?

A
  • The smaller the radii, the stronger the bond.

* This is because smaller ions can pack more closely together and electrostatic attraction gets weaker with distance.

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15
Q

What is charge density of an ion?

A

The amount of charge per unit area or volume.

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16
Q

How does charge density affect the strength of an ionic bond?

A

Ions with high charge density form stronger bonds than those with low charge densities.

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17
Q

What is an ion with high charge density?

A

An ion with a large charge spread over a small area.

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18
Q

What is an ion with low charge density?

A

An ion with a small charge spread over a large area.

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19
Q

NaF and CaO.

Which has the higher melting point and why?

A
  • CaO
  • It is made up of Ca2+ and O2- ions, while NaF is made up of Na+ and F- ions.
  • So the charges in the CaO are greater, resulting in stronger ionic bonds.
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20
Q

NaF and CsF.

Which has the higher melting point and why?

A
  • NaF
  • The ionic radius of Ca+ is greater than that of Na+.
  • So Na+ and F- ions can pack more tightly than Ca+ and F- ions, resulting in stronger ionic bonds.
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21
Q

Describe how the size of an ion changes as you go down a group. Why?

A

It is increased, because extra electron shells are added.

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22
Q

What are isoelectronic ions?

A

Ions of different atoms with the same number of electrons.

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23
Q

Describe how the radius of isoelectronic ions changes as the atomic number increases.

A

It decreases, because there is greater attraction per electron from the protons, so they are pulled in closer.

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24
Q

What type of diagram is used to show ionic bonding?

A

Dot-and-cross diagrams

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25
Remember to practise drawing out dot-and-cross diagrams for ionic bonding.
See pg 20 of revision guide
26
Describe the structure that ionic compounds form.
Giant ionic lattice
27
In an ionic compound, why does a giant ionic lattice form?
Each Ion is electrostatically attracted in all directions to ions of the opposite charge.
28
What are the physical properties of ionic compounds?
* High melting point * Soluble in water, but not in non-polar solvents * Non-conductors when solid, but do conduct when molten or dissolved * Brittle
29
Describe how the physical properties of ionic compounds is evidence for the giant ionic lattice model.
* High melting point -> Tells you that the ions are held together by a strong attraction, like that between positive and negative ions * Soluble in water, but not in non-polar solvents -> Particles must be charged * Non-conductors when solid, but do conduct when molten or dissolved -> Ions are present, which are only free to move when molten or dissolved * Brittle -> Layers can’t be pulled over each other, or the electrostatic repulsion would be too great
30
Aside from the physical properties of ionic compounds, what else is evidence for the theory of ionic bonding? Explain how this works.
Migration of ions in electrolysis: • When you electrolyse a green solution of copper(II) chromate(VI) ions in some wet filter paper, the filter paper turns blue at the cathode and yellow at the anode. • This is due to blue copper(II) ions in solution and yellow chromate(VI) ions in solution, which move to the electrodes. • The solution is yellow to start off with because the ions mix.
31
What is the colour change observed in the migration of ions experiment observed? Why?
* Green colour (copper(II) chromate(VI) ions) splits into blue and yellow at the electrodes * Blue is due to copper(II) ions * Yellow is due to dichromate(VI) ions
32
What colour are dichromate(VI) ions?
Yellow
33
What types of bonds hold a molecule together?
Covalent
34
Describe how a covalent bond forms.
* Two atoms share electrons so they’ve both got full outer shells of electrons. * A covalent bond is the strong electrostatic attraction between the two positive nuclei and the shared electrons in the bond
35
What diagrams can be used to represent covalent bonding?
Dot-and-cross diagrams
36
Remember to practise drawing out dot-and-cross diagrams for covalent bonding.
Pg 22 of revision guide, including examples
37
Describe the dot-and-cross diagram for the bonding in CO.
* C has two non-bonding electrons * O had two non-bonding electrons * In the covalent bond -> 2 electrons from the C + 4 electrons from the O (this includes a dative covalent bond)
38
What things balance at the given bond length for a bond?
* Repulsion between positively charged nuclei | * Attraction between the positive nuclei and the area of electron density
39
What determines the enthalpy of a bond?
Its length.
40
How does the length of a bond affect its enthalpy?
The shorter the bond, the higher the bond enthalpy.
41
Compare the length and bond enthalpy of C-C, C=C and C triplet bond?
* C-C is the longest and has the smallest enthalpy | * C triplet bond is the shortest and has the highest enthalpy
42
What is a dative covalent bond?
A covalent bond where bond of the electrons are donated by one atom.
43
Describe how an ammonium Ion involves a dative covalent bond.
* Ammonia has a lone pair of electrons on the N | * This lone pair is donated to form a dative covalent bond with a H⁺ ion to give an ammonium ion
44
When a dative covalent bond is being drawn, which way does the arrow point?
The direction in which the electrons are being donated.
45
Describe the structure of AlCl₃.
* Each Al forms single covalent bonds with 3 Cl’s * However, in certain conditions, two AlCl₃ molecules can combine to form Al₂Cl₆. * This happens when one Cl in each molecule donates a lone pair to the Al in the other molecule, forming two dative covalent bonds * This gives aluminium a full outer shell (See diagram pg 23 of revision guide)
46
Remember to practise drawing out the structure of Al₂Cl₆.
See diagram pg 23 of revision guide
47
Give some examples of molecules where dative covalent bonding is seen.
* CO * NH₄⁺ * Al₂Cl₆
48
Order these in terms of the angle size: • Bonding pair/bonding pair • Lone pair/bonding pair • Lone pair/Lone pair
* Smallest angle: Bonding pair/bonding pair * Lone pair/bonding pair * Largest angle: Lone pair/lone pair
49
Do lone pair electrons or bonding pair electrons repel more?
Lone pair electrons
50
What is the name for the way of predicting molecular shapes according to how much electron pairs repel?
Electron pair repulsion theory
51
Describe how much an extra lone pair decreases the bonding angles by in a molecule with 4 electron pairs around the central atom.
2.5°
52
Describe a tetrahedral molecule’s bond angle and how they change with extra lone pairs.
* 109.5° * Decreases by 2.5° for every extra lone pair, so: * 109.5° -> 107° -> 104.5°
53
Describe how you can predict the shape of a molecule around a central atom.
1) Find the central atom. 2) Work out the number of bonding pairs and lone pairs of electrons around it. 3) Use this information to predict the shape of the molecule.
54
For a molecule with 2 bonding pairs around the central atom, give the: • Bond angles • Name
* Bond angles = 180° | * Name = Linear
55
For a molecule with 3 bonding pairs around the central atom, give the: • Bond angles • Name
* Bond angles = 120° | * Name = Trigonal planar
56
For a molecule with 2 bonding pairs and 1 lone pair around the central atom, give the: • Bond angles • Name
* Bond angles = 119° | * Name = Bent
57
For a molecule with 4 bonding pairs around the central atom, give the: • Bond angles • Name
* Bond angles = 109.5° | * Name = Tetrahedral
58
For a molecule with 3 bonding pairs and 1 lone pair around the central atom, give the: • Bond angles • Name
* Bond angles = 107° | * Name = Trigonal pyramidal
59
For a molecule with 2 bonding pairs and 2 lone pairs around the central atom, give the: • Bond angles • Name
* Bond angles = 104.5° | * Name = Bent
60
For a molecule with 5 bonding pairs around the central atom, give the: • Bond angles • Name
* Bond angles = 90° and 120° | * Name = Trigonal bipyramidal
61
For a molecule with 4 bonding pairs and 1 lone pair around the central atom, give the: • Bond angles • Name
* Bond angles = 87° and 102° | * Name = Seesaw
62
For a molecule with 3 bonding pairs and 2 lone pair around the central atom, give the: • Bond angles • Name
* Bond angles = 87.5° | * Name = Distorted T
63
For a molecule with 6 bonding pairs around the central atom, give the: • Bond angles • Name
* Bond angles = 90° | * Name = Octahedral
64
For a molecule with 5 bonding pairs and 1 lone pair around the central atom, give the: • Bond angles • Name
* Bond angles = 81.9° | * Name -> Square pyramidal
65
For a molecule with 4 bonding pairs and 2 lone pair around the central atom, give the: • Bond angles • Name
* Bond angles = 90° | * Name = Square planar
66
Remember to practise drawing out and naming different shapes of molecules.
Pg 25 of revision guide
67
What type of structure do covalently bonded atoms form?
* Simple covalent molecule | * Giant covalent structure
68
Give three examples of giant covalent structures.
* Carbon * Silicon * Silicon dioxide
69
Describe the structure of diamond.
Tetrahedral, with each carbon bonding with another four carbons. (See diagram pg 26 of revision guide)
70
Describe the structure of silicon dioxide.
* Silicon atoms form a tetrahedral shape, like in diamond * Oxygen atoms are between each of the two silicon atoms (See diagram pg 26 of revision guide)
71
Describe and explain the properties of giant covalent structures.
* High melting points -> Need to break a lot of very strong bonds * Hard -> Due to very strong bonds * Insoluble -> Atoms are more attracted to their neighbours in the lattice than to solvent molecules. Also, no ions. * Can’t conduct electricity -> No charged ions or free electrons.
72
What is the only exception to giant covalent structures not being able to conduct electricity?
Graphite
73
Describe the structure of graphite.
* Each carbon forms 3 single bonds with other carbons -> This forms a hexagonal structure * There are multiple sheets held together by weak forces * Each carbon has a free electron (to conduct electricity)
74
What allows graphite to conduct electricity?
Each carbon has a single free electron that is free to move around and conduct electricity.
75
What is one layer of graphite called?
Graphene
76
Describe the structure of graphene.
* It is a single sheet | * Each carbon forms 3 single bonds with other carbons -> This forms a hexagonal structure
77
What are the properties of graphene?
* Conductor of electricity * Transparent * Strong * Light
78
Describe the structure of a metal.
Exist as giant metallic lattices: • The outermost shell of the metal atoms are delocalised -> This leaves a positive metal ion • Positive metal ions are electrostatically attracted to the sea of delocalised electrons • The overall lattice structure is made up of layers of positive metal ions, separated by layers of electrons
79
Describe and explain the properties of metals.
* High melting point -> Strong attraction between the positive ions and sea of delocalised electrons * Malleable + Ductile -> Layers of metal ions are free to slide over each other, without disrupting the bonding * Good thermal conductors -> Electrons can pass kinetic energy to each other * Good electrical conductors -> Delocalised electrons are free too move and carry a current * Insoluble (except in liquid metals) -> Metallic bonds are very strong
80
What factors affect the melting point of a metal?
* No. of delocalised electrons per atom * Size of the metal ion * Lattice structure
81
Define electronegativity.
The ability of an atom to attract the binding electrons in a covalent bond.
82
On what scale is electronegativity measured?
Pauling
83
What value are very electronegative elements given on the Pauling scale?
High
84
What are the properties of very electronegative elements?
* High nuclear charges | * Small atomic radii
85
How does electronegativity change across periods and groups?
* Increases across periods | * Increases up groups
86
What causes a polar bond?
When the two elements on either side of the bond have very different electronegativities.
87
What is the name for a difference in charge between the atoms at each end of a polar bond?
Dipole
88
What are the only truly covalent bonds?
Those between two atoms of the same element.
89
How does the difference in electronegativity between two elements tell you about the bond between them?
The greater the difference, the more ionic the bond.
90
Do polar bonds always make polar molecules?
No, often the polar bonds may point in opposite directions and cancel out.
91
What are the 3 types of intermolecular force?
1) London forces 2) Permanent dipole-permanent dipole bonds 3) Hydrogen bonds
92
What is another name for London forces?
Instantaneous dipole-induced dipole bonds
93
What type of molecules form London forces?
All atoms and molecules do
94
Describe how London forces form.
* Electrons in charge clouds are always moving quickly * At any instant, the electrons in an atom are likely to be made to one side than another -> This is an instantaneous dipole * This dipole can induce another temporary dipole in the opposite direction on a neighbouring atom -> This is an induced dipole * This can induce another dipole in another atom * These dipoles are constantly being created and destroyed, but the overall effect is for the atoms to be attracted
95
What types of forces hold molecules in a lattice?
London forces
96
Describe the bonding in iodine.
* Iodine atoms within I₂ are held together by strong covalent bond. * The whole molecules are then held together in a molecular lattice arrangement by weak London forces
97
What is the effect of stronger intermolecular forces?
The melting and boiling posits are higher.
98
What factors affect the size of the intermolecular forces within an organic compound?
* Size | * Surface Area
99
Compare the strength of intermolecular forces in straight and branched-chain organic compounds.
Straight molecules have stronger intermolecular forces due to their increased surface area and better ability to pack closer together.
100
Compare the melting and boiling points of butane and methylpropane (both C₄H₁₀).
Butane has a higher melting and boiling point because it is straight, not branched-chain.
101
Describe how permanent dipole-permanent dipole bonds form.
δ⁺ and δ⁻ charges on polar molecules cause weak electrostatic forces of attraction between molecules.
102
Can a molecule have both London forces and permanent dipole-permanent dipole bonds?
Yes
103
Compare the melting and boiling point of molecules with just London forces and those with both London force and permanent dipole-permanent dipole bonds.
The ones with both tend to have higher melting and boiling points.
104
What is the strongest intermolecular force?
Hydrogen bonding
105
Describe the formation of a hydrogen bond.
• H atom bonded to N, O or F can form a hydrogen bond with the lone pair of electrons on N, O or F of another molecule This is because: • The H is given a strong positive by the very electronegative N, O or F • It is also very small, so it has a high charge density
106
How can you tell whether something is a hydrogen bond?
There is a H atom with an N, O or F either side of it.
107
What is the effect of hydrogen bonding?
It raises the melting and boiling point of a molecule hugely.
108
Describe and explain the graph of boiling points of hydrogen halides.
* HF has a very high boiling point -> Due to hydrogen bonding * Large drop to HCl -> Since it can’t form hydrogen bond * Steady increases to HBr and HI -> Increasing size of the molecule means that there are stronger London forces (which overrides the effect of the weaker dipoles)
109
Why does HBr have a higher boiling point than HCl, even though the strength of the permanent dipole-permanent dipole increases?
The increased strength of the London forces overcomes this decrease.
110
Remember to revise the graph of the boiling points of Group 6 hydrides (like H₂O).
Pg 32 of revision guide
111
What is the effect of hydrogen bonds on solubility?
It makes the substance soluble, because it can form hydrogen bonds with the water molecules, allowing them to mix and dissolve.
112
Describe the structure of ice.
* 6 H₂O molecules line up to form a hexagon | * There are hydrogen bonds holding the structure together
113
Remember to practise drawing out the structure of ice.
Pg 33 of revision guide
114
Why does ice float?
In ice: • 6 H₂O molecules line up to form a hexagon • There are hydrogen bonds holding the structure together • This is a structure that wastes a lot of space When the ice melts: • The water molecules lose their lattice structure and become more closely packed • This increases the density, so water has a lower density and floats
115
What is the effect of hydrogen bonding on volatility?
It reduces the volatility (i.e. it increases the boiling point)
116
Compare and explain the volatility of alcohols and equivalent alkanes.
* The alcohols have lower volatility (i.e. higher boiling points) * Because they can form hydrogen bonds due to the -OH group * So it takes more energy to overcome the intermolecular forces and boil them
117
What things have to happen in order for a substance to dissolve?
* Bonds in the substance have to break * Bonds in the solvent have to break * Bonds have to form between the substance and the solvent The strength of the new bonds has to be about the same as, or greater than, the bonds broken.
118
What are the two types of solvent?
* Polar | * Non-polar
119
Do all polar solvents form hydrogen bonds?
Most, but not all. For example, propanone only forms London forces and permanent dipole-permanent dipole bonds.
120
What intermolecular forces do non-polar solvents have?
London forces
121
What is the process of water molecules surrounding ions in dissolving called?
Hydration
122
In what solvents do ionic substances dissolve?
Polar solvents
123
Why don’t some ionic substances dissolve in water?
The bonding between their ions is too strong - so it is stronger than the bonds they would form with the water molecules.
124
How does the solubility of alcohols change with their size?
The larger the alcohol, the less soluble it is, because: • Small alcohols -> The OH groups can form H-bonds with the water molecules • Large alcohols -> The hydrocarbon chain can’t form H-bonds with the water molecules
125
Do all polar molecules dissolve in water?
No, usually it is just those that can form hydrogen bonds with the water.
126
Do halogenoalkanes dissolve in water and non-polar solvents?
* Water -> No | * Non-polar solvents -> Yes, those that form permanent dipole-permanent dipole bonds
127
Explain the solubility of halogenoalkanes in water.
They are not soluble because the hydrogen bonding between the water molecules is stronger than the bonds that would be formed with halogenoalkanes.
128
What do non-polar substances dissolve in best?
Non-polar solvents
129
What are the general rules for solubility?
Substances usually dissolve best in solvents with similar intermolecular forces.
130
``` Describe ionic compounds in terms of: • Melting/Boiling point • State at RTP • Electrical conductivity when solid • Electrical conductivity when liquid • Solubility in water ```
* Melting/Boiling point -> High * State at RTP -> Solid * Electrical conductivity when solid -> No * Electrical conductivity when liquid -> Yes * Solubility in water -> Yes
131
``` Describe simple molecular compounds in terms of: • Melting/Boiling point • State at RTP • Electrical conductivity when solid • Electrical conductivity when liquid • Solubility in water ```
* Melting/Boiling point -> Low * State at RTP -> Usually liquid or gas (but may be solid) * Electrical conductivity when solid -> No * Electrical conductivity when liquid -> No * Solubility in water -> Depends on whether it can hydrogen bond
132
``` Describe giant covalent compounds in terms of: • Melting/Boiling point • State at RTP • Electrical conductivity when solid • Electrical conductivity when liquid • Solubility in water ```
* Melting/Boiling point -> High * State at RTP -> Solid * Electrical conductivity when solid -> No * Electrical conductivity when liquid -> N/A since they usually sublime * Solubility in water -> No
133
``` Describe metallic substances in terms of: • Melting/Boiling point • State at RTP • Electrical conductivity when solid • Electrical conductivity when liquid • Solubility in water ```
* Melting/Boiling point -> High * State at RTP -> Solid * Electrical conductivity when solid -> No * Electrical conductivity when liquid -> N/A since they usually sublime * Solubility in water -> No
134
Substance X has a melting point of 1045 K. When solid, it is an insulator, but once melted it conducts electricity. Identify the type of structure present in substance X.
Ionic
135
Remember to practise predicting the properties of a material from its structure and vice versa.
Pgs 36 and 37 of revision guide
Inorganic Chemistry
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