Topic 2 Bonding & Structure Flashcards
1
Q
Ice & I2 have open lattice structures. What is that?
A
Simple molecules held together in a lattice by intermolecular forces of attraction.
2
Q
Ionic bonding
A
Strong electrostatic attraction between oppositely-charged ions held in a giant lattice.
3
Q
Why do ionic radii decrease down a group?
A
More shells of electrons.
4
Q
What experiment can be used to show a compound is ionic?
A
Electrodes in a circuit are placed in a solution of the ionic compound. The bulb lights up.
5
Q
Isoelectronic
A
Ions with the same electron configuration as a Noble gas.
6
Q
What is the trend in bond strength/enthalpy and bond length down group 7?
A
Bond lengths increases and bond strength/enthalpy decreases.
Exception: F-F, which has the shortest bond length yet the lowest bond enthalpy.
7
Q
Why is F-F an exception to the bond length & strength trend down group 7?
A
The lone pairs around each F atom repel, which weakens the covalent bond.
8
Q
Covalent bond
A
Strong electrostatic attraction between 2 nuclei and the shared pair of electrons between them.
9
Q
Sigma bond
A
Covalent bond where the p-orbitals overlap end-to-end.
10
Q
Pi bond
A
Covalent bond that forms only after a sigma bond has formed, usually by the sideways overlap of p-orbitals.
11
Q
Dative/coordinate covalent bond
A
Formed when one atom donates both bonding electrons.
12
Q
Bond length
A
Sum of the atomic radii.
13
Q
What determines the strength of a covalent bond?
A
Bond length. Longer bonds are weaker, hence the hydrolysis of the C-X bond in halogenoalkanes.
14
Q
P & S in period 3
A
Can hold more than 8 e- in the outer shell due to the availability of energetically favourable orbitals.
15
Q
Describe the bonding in graphene.
A
Same as in graphite.
Single layer of hexagonal rings.
Delocalised e- are free to move to conduct electricity.
16
Q
Describe the bonding in graphite.
A
Each C has 3 sigma bonds & 1 delocalised e-.
Planar hexagonal rings held together by London forces.
Layers can slide over each other, and e- are free to move to conduct electricity.
17
Q
Where can the delocalised e- move in graphite?
A
Electrons flow parallel to the layers. The energy gap is too big to allow perpendicular flow.
18
Q
Describe the bonding in diamond.
A
Each C forms 4 sigma bonds. Tetrahedral. Numerous, strong covalent bonds.
19
Q
Explain why some ionic compounds have higher melting points than others.
A
Compare charge and ionic radius. Higher charge density of the cation = greater attraction between the cation & anion.
20
Q
Migration of ions experiment
A
Electrolysis of a green copper (II) chromate (VI) solution. Cu2+ ions migrate to the cathode and appear yellow. Chromate ions migration to the anode and appear blue.
21
Q
In a negative ion, which atom gains the e-?
A
The more electronegative one.
22
Q
Why do alcohols have higher boiling points than alkanes of the same chain length?
A
Alcohols have hydrogen bonding, but alkanes have only London forces.
23
Q
Why does HF have hydrogen bonding but HBr does not?
A
F is small & electronegative, so creates a greater dipole moment in HF, meaning it can attract the lone pair on the F of another molecule.
24
Q
Why do branched isomers have lower boiling points than straight-chain isomers?
A
Branched molecules have fewer points of contact, so weaker London forces.
25
Upon what does the strength of London forces depend?
- The number of e-.
- Branching/the number of points of contact.
26
Describe the trend in boiling points of hydrogen halides.
HF has the highest due to hydrogen bonding. Decrease down to HCl then increase thereafter due to an increasing number of e-.
27
How do London forces arise?
Oscillations of electrons in a molecule results in a temporary dipole, which induces a dipole in a neighbouring molecule.
London forces are always present!
28
Electronegativity
The ability of an atom to attract bonding pairs of e- in a covalent bond.
29
Simple test: is a liquid polar?
A charged rod will bend a stream of liquid, if the molecules are polar.
30
Why do some molecules have polar bonds but not an overall dipole?
The dipoles in CH4 of CCl4 are symmetrically arranged, so they cancel out.
If the polar bonds/dipoles are not symmetrically arranged, the molecule has an overall dipole.
31
Why does water have an unusually high boiling point for a small, simple molecule?
Hydrogen bonding.
32
When does water have its maximum density?
4°C
33
Why is ice less dense than liquid water?
Molecules form the maximum number of hydrogen bonds in ice to form an open lattice structure.
34
Why are some organic molecules soluble in ethanol but not water?
There are fewer H bonds to overcome in ethanol compared to water.
35
What is typically true when a solid dissolves?
The sum of the hydration enthalpies > the lattice enthalpy.
36
What are formed when dissolving occurs?
Ion-dipole interactions. The dipole in the water surrounds the ions.
37
What does the solubility of a solute in a solvent depend upon?
The compatibility of the intermolecular forces, e.g., polar molecules dissolve more readily in polar solvents.
38
When are London forces greater?
In molecules with more electrons and a larger surface area/or points of contact.
39
Lone pairs of electrons
Repel each other more than bonding pairs do.
40
How can you answer a question on shape and bond angles?
- State the shape and bond angle.
- Explain in terms of the number of bonding pairs, the number of lone pairs and that lone pairs repel more.
- Maximum separation, minimum repulsion.
41
2 electron-dense areas & no lone pairs: shape, bond angle; example.
Linear, 180° and CO2.
42
3 electron-dense areas & no lone pairs: shape, bond angle; example.
Trigonal planar, 120° and BF3.
43
3 electron-dense areas including 1 lone pair: shape, bond angle; example.
Bent, less than 120° and NO2^-.
44
4 electron-dense areas & no lone pairs: shape, bond angle; example.
Tetrahedral, 109.5° and ammonium ion.
45
4 electron-dense areas including 1 lone pair: shape, bond angle; example.
Trigonal pyramidal, 107° and NH3.
46
4 electron-dense areas including 2 lone pairs: shape, bond angle; example.
Bent, 104.5° and H2O.
47
5 electron-dense areas & no lone pairs: shape, bond angle; example.
Trigonal bipyramidal, 90° & 120° and PCl5.
48
5 electron-dense areas including 1 lone pair: shape, bond angle; example.
Seesaw SF4.
49
5 electron-dense areas including 2 lone pairs: shape, bond angle; example.
T shaped ICl3.
50
5 electron-dense areas including 3 lone pairs: shape, bond angle; example.
Linear I3^-.
51
6 electron-dense areas & no lone pairs: shape, bond angle; example.
Octahedral, 90° and SF6 (g).
52
6 electron-dense areas including 1 lone pair: shape, bond angle; example.
Square pyramidal IF5.
53
6 electron-dense areas including 2 lone pairs: shape, bond angle; example.
Square planar ICl4^-.
54
6 electron-dense areas including 3 lone pairs.
T-shaped.
55
6 electron-dense areas including 2 lone pairs.
Linear.
56
Why cannot lattice enthalpy be measured directly in a single experiment?
Gaseous ions cannot be isolated; they would attract instantly.
57
A Born-Haber cycle is a vector diagram. If more than one mole is involved in a reaction, what happens?
Multiply the enthalpy change accordingly.
58
Explaining the difference between experimental and theoretical lattice energies.
No significant difference between the two value: 100% ionic bonding.
Experimental is more exothermic than theoretical: the cation is small and highly-charged and/or the anion is large, which means the anion will be polarised by the cation and bonding will have covalent character.
59
Of what is lattice enthalpy a measure?
The strength of ionic bonding. The more exothermic the value, the stronger the bonding.
60
Why does P form PCl5, but N does not form NCl3?
P can expand its octet. N doesn’t have d-orbitals, so can only accommodate 8 electrons in its outer shell.
61
Delocalised electrons
Electrons not associated with any particular atom or covalent bond.
62
Electrical conductivity of metals
Generally increases as the number of outer-shell electrons increases.
63
Metallic bonding
The electrostatic force of attraction between the nuclei of metal cations & delocalised electrons.
64
Why are melting points of metals typically high?
Metals have giant lattice structures with many electrostatic forces between the cation nuclei & the delocalised electrons, which require lots of energy to overcome.
65
What factors affect the melting points of metals?
- The number of delocalised electrons per cation.
- The size of the cation. The smaller the cation, the closer the delocalised electrons are to the nucleus of the cation, so the electrostatic forces of attraction are stronger.
66
Group 1 metals
Lower melting temperatures than group 2 metals.
67
Metals in the d-block
Typically have high melting points as they have more delocalised electrons per ion.
68
What 2 factors affect the thermal conductivity of metals?
- Free-moving delocalised electrons pass kinetic energy along the metal.
- The cations are closely-packed and pass kinetic energy from one cation to another.
69
Malleability
Hammered or pressed into different shapes.
70
Ductility
Drawn into a wire.
71
Why are metals malleable & ductile?
When stress is applied to a metal, the layers of cations slide over each other. As the delocalised electrons are free-moving, they move with the cations and prevent strong forces of repulsion forming between cations in different layers.
72
Electrostatic interaction between ions in ionic compounds
Not directional: only the distance between the ions, not their orientations, matters.
73
Stronger ionic bonding
Smaller ions (, so more closely-packed lattices) with higher charges.
74
Ionic radius decreases as the number of protons increases.
As the positive charge of the nucleus increases, the electrons are attracted more strongly and pulled closer to the nucleus.
75
Why do ionic compounds have high melting points?
Ionic solids are giant lattice networks of oppositely-charged ions. The combined electrostatic forces among all the ions is large, so lots of energy is required to overcome the forces of attraction for the ions to break free from the lattice and be able to slide past one another.
76
Why are ionic solids brittle?
When stress is applied to an ionic solid, the layers of ions slide over one another. Ions of the same charge are side by side, so repel each other.
77
Why do molten (or aqueous) ionic compounds conduct electricity?
Ions are free to move and will migrate to electrodes of opposite charges when a potential difference is applied.
78
The ability of an ionic compound to conduct electricity when molten/aqueous as evidence for the existence of ions
When a direct current is passed through molten NaCl:
At the negative electrode: 2Na+ + 2e- --> 2NA
At the positive electrode: 2Cl- --> Cl2 + 2e-
79
The effect of passing a direct current through a solution of copper (II) chromate (VI) solution
Cu2+(aq) = blue at the negative electrode
CrO4^2- (aq) = yellow at the positive electrode
(Green in the middle.)
80
pi bond
Sideways overlap of 2 p-orbitals.
Cannot form until a sigma bond has been formed.
Results in high electron density above & below the molecule.
81
sigma bond
End-on overlap of either 2 s-orbitals or 2 p-orbitals.
82
When is covalent bond formed?
When an atomic orbital containing a single electron from one atom overlaps with another atomic orbital, also containing a single electron, from another atom.
83
Why are alkenes more reactive than alkanes?
The pi bond is weaker than the sigma bond, hence alkanes readily undergo addition reactions.
84
Bond length
The distance between the nuclei of 2 atoms that are covalently bonded together.
85
The shorter the bond, the greater the bond strength.
This is due to the increased electrostatic attraction between the 2 nuclei and the electrons in the overlapping atomic orbitals.
86
General trend in electronegativity
Decreases down a group.
Increases from left to right across a period.
87
Electronegativity
The ability of an atom to attract a bonding pair of electrons in a covalent bond.
88
Polar covalent bond
A type of covalent bond between 2 atoms where the bonding electrons are unequally distributed. This means one atom carries a slight positive charge & the other atom carries a slight negative charge.
89
How can we show a polar covalent bond?
Electron density maps. The atom with the higher electronegativity will have more electron density, so the electron density will not be symmetrical about the whole molecule. The contour lines are closer together near the more electronegative atom.
90
Discrete/simple molecule
Electrically neutral group of 2 or more atoms held together by chemical bonds.
91
Examples of the expanded octet
SF6 & PCl5 (elements in period 3 or beyond where the 3d subshell is available).
92
Fewer than 8 electrons in its outer shell
BCl3 & BCl2
93
When does a dative covalent bond form?
When an empty orbital of 1 atom overlaps with an orbital containing a lone pair of electrons of another atom. Represented by an arrow. E.g., Al2Cl6.
94
Molecules with multiple bonds in VSEPR
Treat each multiple bond as a single pair of electrons.
95
2 bond pairs 0 lone pairs
Linear, 180°
96
3 bond pairs 0 lone pairs
Trigonal planar, 120°
97
4 bond pairs 0 lone pairs
Tetrahedral, 109.5°
98
5 bond pairs 0 lone pairs
Trigonal bipyramidal, 120, 90 & 180 degrees.
99
6 bond pairs 0 lone pairs
Octahedral, 90° & 180°
100
3 bond pairs 1 lone pair
Trigonal pyramidal, 107 degrees.
101
2 bond pairs 2 lone pairs
V-shaped/bent, 104.5°
102
2 key points for VSEPR
- Electron pairs arrange themselves around the central atom, so that repulsion around them is at a minimum.
Lone pair-lone pair repulsion > lone pair-bond pair repulsion > bond pair-bond pair repulsion.
103
Shape: NH4^+ vs. NH3
NH3 is trigonal pyramidal (107°) & NH4^+ is tetrahedral (109.5°).
104
Dipole
Exists when 2 charges of equal magnitude & opposite signs are separated by a small distance.
105
When is a molecule with polar bonds non-polar?
When the dipoles cancel out due to symmetry (note: trigonal pyramidal & V-shaped aren't symmetrical).
106
How do London dispersion forces arise?
Electron density fluctuates over time. If at any instant the electron density becomes unsymmetrical in the molecule, a dipole is generated, which induces a dipole in a nearby molecule.
107
Features of London forces:
- Always present.
- The more points of contact between molecules, the greater the overall London force.
- The greater the number of electrons, the greater the fluctuations in electron density & the larger the instantaneous dipoles created.
108
Permanent dipole-dipole forces
2 molecules with dipoles will attract one another. (Or a molecule with a permanent dipole induces a dipole in another molecule.)
109
Why are London forces often the most significant interaction when just London forces & permanent dipole-dipole forces are present?
London forces are always aligned to produce a favourable interaction. Permanent dipole-dipole forces are not always favourably aligned (due to the random movement of electrons).
110
Hydrogen bond
An intermolecular interaction between a H atom of a molecule bonded to an atom, which is more electronegative than H and another atom in the same or a different molecule.
111
Boiling point of unbranched alkanes increases with increasing molecular mass.
- The number of electrons per molecules increases, so instantaneous & induced dipoles increase in strength.
- Instantaneous dipole-induced dipole forces exist at each point of contact between molecules, so the more points of contact, the greater the overall intermolecular force of attraction.
112
Why do branched alkanes have lower boiling points than their unbranched isomers?
The more branching, the fewer points of contact, so they don't pack together as well, so intermolecular forces are weaker.
113
Boiling point of alcohols
Typically higher than similar alkanes due to hydrogen bonding, however, as chain length increases, the London forces increase until eventually they predominate.
114
Boiling points of hydrogen halides
Steady increase from HCl to HI due to the increasing no. of electrons, so increasing London forces, but HF is much higher due to hydrogen bonding.
115
2 anomalous properties of H2O
- Relatively high melting & boiling points despite so few electrons.
- The density of ice at 0°C is less than that of water at 0°C.
116
Why does HF have a lower boiling point than H2O?
HF forms an average of 1 H bond per molecule, water forms an average of 2 H-bonds per molecule, so hydrogen bonding is much more extensive in water. (Not all the hydrogen bonds in HF are broken upon vaporisation as HF is substantially polymerised, even in the gas phase.)
117
The density of ice at 0°C is less than that of water at 0°C. Why?
Molecules in ice are arrange in rings of 6 held together by H bonds, which creates large areas of open space inside the rings in ice. The ring structure is destroyed when it melts.
118
Choosing suitable solvents
- The solute particles must be separated from each other & surrounded by solvent particles.
- Forces of attraction between the solvent & solute must be strong enough to overcome the solvent-solvent forces & solute-solute forces.
119
Solubility of alcohols
Soluble in water as it can form H bonds to the water, but solubility decreases with increasing hydrocarbon chain length as London forces predominate between the alcohol molecules.
120
Bromine in hexane
Often used instead of bromine water to test for unsaturation as both are non-polar, so miscible & the compound being tested is also likely to be non-polar.
121
Some polar compounds are insoluble in water.
The solute-solvent attractions are not enough to replace the strong hydrogen bonding in water.
122
Metallic lattice
A regular arrangement of cations surrounded by delocalised electrons.
123
Ionic lattice
A regular arrangement of positive & negative ions.
124
Giant covalent lattices
A giant network of atoms linked to each other by covalent bonds.
Diamond, graphite, graphene & Silicon (IV) oxide.
125
Diamond
Each C atom forms 4 sigma bonds to 4 other C atoms in a giant 3D tetrahedral arrangement. All bond angles are 109.5°.
Hard due to very strong C-C bonding.
High melting point due to lots of strong C-C bonds that require lots of energy to break.
126
Graphite
Layered structure. Each C atom is bonded to 3 others via sigma bonds. the 4th electron on each C is in a p-orbital. The C atoms are close enough for the p-orbitals to overlap with one another to produce a cloud of delocalised electrons above & below the plane of the ring.
127
Graphite as a lubricant
Layers slide easily over each other as there are only weak London forces between the layers & gases adsorb to the surface.
128
Graphite has a high melting temperature.
Lots of covalent bonds to break.
129
Graphite conducts electricity.
The delocalised electrons are free to move under a potential difference. It can only conduct electricity parallel to its layers as delocalised electrons cannot move between the layers.
130
Graphene
One-atom thick layer of graphite. 200x stronger than steel, absorbs light, thermal conductor and can be shaped into fullerenes & carbon nanotubes.
131
Iodine structure
Molecular lattice (when solid). The iodine diatomic molecules are held together in a regular pattern by London forces. Crystalline.
132
Examples of molecular solids
Iodine, S8, white phosphorous (P4), Buckminster fullerene (C60), dry ice (solid CO2) & sucrose (C12H22O11).
133
Molecular solids have low melting & boiling points
Little energy is required to overcome the weak intermolecular (London forces). London forces increase with number of electrons, so a macromolecular solid, e.g., poly(ethene) has a much higher melting point than ethene.
134
Molecular: properties
Generally insoluble, unless H-Bonding is possible (as in sucrose) or it reacts with H2O (e.g., Cl2). Non-conductors. Generally low melting & boiling points.
135
Giant metallic & giant covalent: solubility
Insoluble.
136
If asked for a dot & cross diagram..
Check whether it's asking for all the electrons or just the outer shell.
137
Why are metals malleable/ductile?
The layers of ions can easily slide over each other.
138
Carbonates require high decomposition temperatures. Why?
Strong bonds within the carbonate ion.
139
Covalent bond.
Strong electrostatic attraction between 2 nuclei & the shared pair of electrons.
140
Nitrogen **cannot** expand...
... its octet!
141
AlF3 vs. AlCl3. Why does AlF3 have a higher sublimation point?
Al-F is more polar than Al-Cl.
AlF3 has a giant structure with strong electrostatic forces of attraction between the ions.
AlCl3 is a mostly covalent small molecule, so only the weak intermolecular London forces need to be broken.
142
When comparing the density of two elements, consider...
The RAMs of each element & the form in which the element exists, e.g., metals have layers of cations close together.
143
There can be very different values for the compressibility of graphite. Why?
Lower values refer to the weak London forces between the layers. Higher values refer to the strong covalent bonds within each layer.
144
Dative covalent =
Coordinate bonds.
145
How can Boron form a dative covalent bond with a species with a lone pair?
The boron atom is electron-deficient, so can accept 2 electrons to complete the octet.
146
How does hydrogen bonding cause ice to be less dense than water?
Ice is more open due to the 3D lattice structure. Hydrogen bonds are longer than covalent bonds.
147
Why can phosphorous expand its octet, but nitrogen cannot?
P can accommodate 10 electrons as it has 3d-orbitals available for the promotion of electrons. Nitrogen has no 2d-orbitals, so can only accommodate 8 electrons in its outer shell.
148
Describe how London forces arise between molecules.
Uneven distribution of electrons due to their random movement results in a temporary dipole in the first molecule. This induces a dipole on another molecule.
