Structure and Bonding Flashcards
1
Q
What are the four main types of bonding?
A
Ionic, Covalent, Metallic, and Hydrogen bonding.
2
Q
What is ionic bonding? Give some examples
A
The bonding occurs in compounds in which the constituent elements have a relatively large difference in electronegativity
Examples are NaCl and CsF
3
Q
What is covalent bonding? Give some examples
A
Involves the formal sharing of electrons between two atoms
They are found in elements with high electronegativity e.g. Cl₂, F₂, Br₂
4
Q
When are covalent bond formed in heteroatomic molecules?
A
Occurs when the different atoms have a small difference in electronegativity
5
Q
What is metallic bonding?
A
Found in elements that have a low electronegativity, with the electrons delocalised into a ‘sea of electrons’
6
Q
What is hydrogen bonding? Give an example of a molecule where hydrogen bonding occurs
A
The primary electrostatic interaction between a hydrogen atom attached to an electronegative atom and a second electronegative atom
It is an example of a non-covalent interaction
Water is the main example
7
Q
What is electronegativity?
A
The ability of an atom to attract an electron to itself
It is considered a slightly artificial concept as there is no way to directly measure it
8
Q
What are two common scales that electronegativity is measured on?
A
The Pauling scale and the Allen scale
9
Q
Where in the periodic table is electronegativity at its highest?
A
The top right (F)
10
Q
Where in the periodic table is electronegativity at its lowest?
A
The bottom left (Cs)
11
Q
What is the general trend in electronegativity as you go down the periodic table?
A
It decreases as the atoms get larger and so are further away from the nucleus
12
Q
What does ‘electropositive’ mean?
A
An element with a low electronegativity (the inverse of electronegative)
13
Q
What does the van Arkel-Ketelaar triangle allow us to predict?
A
The dominant type of bonding in a compound based on electronegativity and the difference in electronegativity
14
Q
What type of bonding is present if the electronegativity difference is greater than 2?
A
Ionic bonding
15
Q
What type of bonding is present if the electronegativity difference is greater than 0.5 but less than 2?
A
Polar covalent bonding
16
Q
What type of bonding is present if the electronegativity difference is less than 0.5?
A
Covalent bonding
17
Q
Why do bonds form?
A
To lower the energy of a system
18
Q
What does an exothermic reaction tell you about the energy level of the reactants and the products?
A
The bonds in the product are stronger than those in the starting material - it is said to be thermodynamically favourable
19
Q
Why is NaCl₂ only a hypothetical compound?
A
The ionic bonds are not strong enough to compensate for the very high second ionisation enthalpy of Na
The bonds are not strong enough to ‘repay’ the energy needed
20
Q
What is the general trend of the ionisation energy across a period?
A
It generally increases
21
Q
Why does the ionisation energy increase across a period?
A
The number of protons increases, increasing the nuclear charge
The additional electrons added are in the same area of space (have the same quantum principle number) so do not shield each other as effectively
22
Q
Why is there a dip in ionisation energy at Boron?
A
The 2p orbital does not penetrate the 2s orbital very well so the effective nuclear charge is lower
It is not as close to the nucleus leading to less attraction
23
Q
Why is there a dip in ionisation energy at Oxygen?
A
There are two electrons in the same p orbital and so there is electron-electron repulsion
24
Q
Which element has the maximum ionisation energy in period 2, and why does this make it relatively inert?
A
Neon has the highest ionisation energy, and it cannot form bonds strong enough to compensate for this, making it very hard to react with
25
What is the trend in ionisation energy as you go down the periodic table?
It decreases
As the principal quantum number increases, electrons spend time further away from the nucleus, decreasing the force of attraction
26
Why is there a smaller difference in ionisation energy between K and Rb compared to Li to Na?
Electrons have been added in the d orbital between K and Rb
This is due to the 10 electrons d orbital not effectively shielding the 10 added protons, therefore effective nuclear charge is higher
27
Do successive ionisation energies require more or less energy?
It always takes more energy - this is because it is harder to remove an electron from a cation due to electrostatic forces of attraction
28
Why is there a much higher increase in the 3rd ionisation energy of Be compared to the 2nd ionisation energy?
There are more protons than electrons, and the electrons are much closer to the nucleus as they have a higher principal quantum number
29
Why is it impossible for reactions involving ionisation energies of core electrons to occur?
The ionisation energy is so high that it is not possible to recoup the energy for bond formation
30
Usually, what type of reaction is electron gain enthalpy? Why?
Exothermic (or very close to it)
There is normally sufficient attraction between the nucleus and the 'new' electron to overcome electron-electron repulsion
31
What is the general trend in electron gain enthalpy across a period?
It generally decreases
32
Why is there an increase in electron gain enthalpy at Be?
The 'new' electron is being added into the 2p¹ orbital, and as this is further away from the nucleus, more energy is needed to accept this electron
33
Why is there an increase in electron gain enthalpy at N?
The 'new' electron is being added in the 2p orbital (which is the first that is doubly filled), leading to more electron-electron repulsion, meaning more energy is required to add the electron
34
What is lattice enthalpy directly proportional to?
( |z⁺| x |z-| ) / (r₊+ r₋)
|z⁺| = Charge on cation
|z-| = Charge on anion
r₊ = Radius of cation
r₋ = Radius of anion
| | = Modulus (absolute value)
35
What is the trend in lattice enthalpy when the anion or cation is larger?
It decreases (lattice enthalpy gets smaller)
36
What is the trend in lattice enthalpy when the ions have a greater charge? (e.g. Al³⁺ compared to Li⁺)
It increases as the bonds are stronger
However, more energy is needed to complete multiple, successive ionisations
37
Why are there no known stable compounds of Helium?
The ionisation energy is too high and the making of bonds cannot repay this debt
38
Why can Xenon and Krypton form stable compounds?
Their ionisation energies are low enough so that the bonds created can repay the debt
39
Why is the electron gain enthalpy for Noble Gases positive?
The extra electron resides in the next shell much further away from the nucleus, meaning that energy is needed to form the attraction between the nucleus and the outer electron
40
Why are half-filled shells (p⁴ to p³) more stable than shells that are not half-filled?
Going from p⁴ to p³, there are two electrons in one orbital, and this repulsion means that less energy is needed to form bonds
41
Why are half-filled shells (p² to p³) more stable than shells that are not half-filled?
It is gaining an electron with all other electrons having the same spin, imparting stability through 'exchange energy'
42
What is 'exchange energy'?
A quantum mechanical effort which means the electrons are as far apart from each other, leading to a greater attraction to the nucleus
43
How does bond length usually affect bond energy?
The shorter the bond, the greater the overlap of orbitals and therefore the bond energy increases
44
How does having more electrons in the covalent bond (single, double, triple) affect bond energy?
The bond electrons there are, the stronger the bond is
45
Why are Bond Dissociation Enthalpies averages?
Bond Dissociation Enthalpies for a given element combination can vary widely, and therefore it is best to take any average
46
What is the general trend in bond energy as the atom or both atoms get larger?
The bond length increases as there is less overlap of orbitals - due to the orbitals getting larger and more diffuse (spread out) - leading to a decrease in bond energy
47
What is the general trend in bond energy when atoms are bonded to a more electronegative atom?
The bond length usually decreases, therefore increasing the bond energy
Electronegative atoms have smaller orbitals, and the electrons are attracted more strongly to the nucleus, creating a stronger orbital overlap
48
Why is the C-N bond in H₃C-NH₂ weaker than the C-C bond in H₃C-CH₃, despite H₃C-NH₂ having a shorter bond length?
There is a repulsion between the lone pair of electrons on the Nitrogen atom and the methyl group, destabilising the structure
There is therefore a balance between the increased electronegative atom increasing bond strength and lone pair repulsion having the opposite effect
49
What are the three forces happening in a covalent bond?
Both electrons are attracted to both nuclei
The two electrons are repelling each other
The two nuclei are repelling each other
50
What happens when the internuclear distance is very short?
The energy is very high as the repulsion between the two nuclei is very large
51
What happens when the internuclear distance is very long?
Each electron is only attracted to its own nucleus
52
What happens at the 'equilibrium bond length'? (This is what happens when a stable covalent bond is formed)
The electron-nuclei attraction outweighs the two repulsive terms (electron-electron and nuclei-nuclei) by the most, creating a stable covalent bond
53
What is the maximum amount of electrons an element in period 2 can have in its outer shell?
8
These elements can form compounds with fewer than 8 electrons (BF₃ and [C(C₆H₅)₃]⁺ are 6 valence electron species
54
What is the term when there are more than 8 electrons in their outer shell?
'Expanding the octet' to form 'hypervalent' species
For example, PF₅ has 10 valence electrons, SF₆ has 12 valence electrons and XeF₆ has 14 valence electrons
55
To which sections of the periodic table do the Lewis Model not apply to?
Transition Metals, Lanthanides and Actinides
56
What are the four steps on how to draw a Lewis structure?
1) Find out the number of electrons in the valance shell (usually the group number)
2) Add one electron for each single, two for a double bond and three for a triple bond
3) Add one electron for each unit of negative charge formally on the atom and subtract one for a positive charge
4) The formal charge on an atom may be calculated from steps 1 - 2, then by subtracting the number of unshared electrons
57
What does VSEPR stand for?
Valence Shell Electron Pair Repulsion Theory
58
What is the purpose of VSEPR?
It is a simple method for determining the geometry of a given atom within a molecule
59
What are the three main assumptions in VSEPR?
1) Electrons form pairs (either bonding or lone pairs)
2) Electron pairs repel each other and molecules adopt a structure in which the electron pairs position themselves as far apart as possible
3) Lone pairs occupy more space than bonding pairs
60
Why do lone pairs occupy more space than bonding pairs?
Lone pairs are closer to the nucleus
61
What type of repulsion is the strongest (bonding pair-bonding pair or lone pair-lone pair)?
Lone pair-lone pair
62
What type of repulsion is the weakest (bonding pair-bonding pair or lone pair-lone pair)?
Bonding pair-bonding pair
63
How can we determine the geometry of an atom?
Using the number of electron pairs, bonding pairs and lone pairs
64
If an atom has no double or triple bonds, how can the number of lone pairs and bonding pairs be calculated?
1) Find out how many valence electrons there are in the atom
2) Add electrons for substituents (one for single, two for double and three for triple)
3) Add / remove electrons for charge
4) Divide by two to give the number of Valance Electron Pairs (VEP)
5) Deduct the number of substituents from VEP to give the number of lone pairs
65
If an atom has double and/or triple bonds, how can the number of lone pairs and bonding pairs be calculated?
1) Find out how many valence electrons there are in the atom
2) Add electrons for substituents (one for single, two for double and three for triple)
3) Add / remove electrons for charge
4) Divide by two to give the number of Valance Electron Pairs (VEP)
5) Deduct one for each attached atom (regardless of whether the bond is multiple or single - this is because we are looking for areas of electron density)
66
What Electron Pair Geometry (EPG) does an atom with two VEP have?
Linear
67
What EPG does an atom with three VEP have?
Trigonal Planar
68
What EPG does an atom with four VEP have?
Tetrahedral
69
What EPG does an atom with five VEP have?
Trigonal Bipyramid
70
What EPG does an atom with six VEP have?
Octahedral
71
How is the molecular shape determined?
From the electron pair geometry and the number of lone pairs
72
What molecular shape does an atom with an EPG of tetrahedral with one lone pair have?
Pyramidal
73
What molecular shape does an atom with an EPG of tetrahedral with two lone pairs have?
Bent
74
What molecular shape does an atom with an EPG of octahedral with two lone pairs have?
Square Planar
75
What molecular shape does an atom with an EPG of octahedral with one lone pair have?
Square based pyramid
76
Why does methane have bond angles of 109.5 whereas ammonia has bond angles of 106.7?
Lone pairs occupy more space than bonding pairs, and so the angles between bonding pairs become compressed as lp-bp repulsion is greater than bp-bp
77
How do electronegative elements impact the shape of molecules? Use the example of CH₃Cl
They often pull electrons towards themselves, meaning they change the shape of molecules
In CH₃Cl, the chlorine pulls the electrons in the C-Cl bond towards itself, creating a bond angle of 108.5 between the Cl and H, and 110.5 between the H-H
78
How do multiple bonds (double/triple) impact the shapes of molecules?
In general, multiple bonds occupy more space than single bonds and therefore induce further distortions in the structure
This makes sense as there are more electrons in the bond
79
How does the 'trigonal bipyramid' shape differ from shapes like tetrahedrons and octahedrons?
Not all of the positions are identical
There are two distinct environments - two axial positions (a) and three equatorial positions (e)
80
Does a trigonal bipyramid structure prefer to have the lone pair in an axial or equatorial position?
Equatorial - this is because it only has two 90-degree lone pair repulsions compared to three, creating a more stable environment
81
Does a trigonal bipyramid structure with two lone pairs prefer to have them in both axial positions, one equatorial and one axial, or both equatorial positions? Rank them in order of favourability
The most favoured is for both of them to be equatorial - this creates just four 90-degree bonding pair-lone pair repulsions
The least favoured is to have one equatorial and one axial lone pair - this creates a lone pair-lone pair repulsion which is very unfavourable
Both lone pairs in axial positions is between the two in terms of favourability as it forms six 90-degree bonding pair-lone pair repulsions
82
Why does it not matter where the lone pair is placed in an octahedral complex?
All sites in an octahedron are identical, and therefore it does not matter
However, the introduction of a lone pair does cause a significant deviation in the geometry of the compound, for example, the different atoms might move out of the plane to accommodate the lone pair
83
What is the bond angle between the lone pairs in an octahedral complex with two lone pairs?
180 degrees, making the lone pairs as far apart as possible, creating a square planar complex
They are not 90 degrees apart as the repulsion would be a lot greater
84
Why can VSEPR not correctly predict the structure of some molecules e.g. [TeBr₆]²⁻?
The lone pairs can be stereochemically inactive so there it has no impact on the shape
This occurs towards the bottom of the periodic table as the lone pairs occupy spherical s orbitals
85
Why can VSEPR not be used on radicals?
They do not have an integer value of valance electron pairs
86
What sections of the periodic table does VSEPR not work well for?
The 'd' and 'f' block elements
87
What is sp³ hybridisation?
Where the s orbital and 3 p orbitals combine to form four of sp³ orbitals
This corresponds to the four corners of the tetrahedron
88
Where does the nucleus lie in sp³ hybridisation?
Not on the node - the node is a plane where there is no electron density
89
How is the C-H bond in methane formed?
There is an overlap between the sp³ orbital on carbon and the 1s orbital on hydrogen, resulting in a two-electron covalent bond
90
What is sp² hybridisation?
One s orbital and two p orbitals combine to form three sp² hybrid orbitals
This leaves one p orbital which lies in the plane that impacts all of the bonds equally
91
How is a double bond formed when the orbitals are hybridized?
The double bond is formed by the p orbital that is not hybridized, and is formed by the overlap of the p-orbitals (one from each atom)
92
Why can't atoms next to double bonds rotate?
It would force the p orbitals to break, making it very unfavourable
93
What is sp hybridisation?
One s orbital and one p orbital combine to form two sp hybrid orbitals
The two p orbitals that have not been hybridized lie in the planes that affect all of the bonds equally
94
How are triple bonds formed?
The two p orbitals that have not been hybridized overlap on two planes with the same happening on both atoms - this creates two pi bonds that are 90 degrees to each other
95
What type of hybridisation occurs in a molecule with the EPG of linear?
sp
96
What type of hybridisation occurs in a molecule with the EPG of trigonal planar?
sp²
97
What type of hybridisation occurs in a molecule with the EPG of tetrahedral?
sp³
98
What type of hybridisation occurs in a molecule with the EPG of trigonal bipyramidal?
sp³d
99
What type of hybridisation occurs in a molecule with the EPG of octahedral?
sp³d²
100
What orbitals are the lone pairs on water part of when hybridisation occurs?
There are sp³ hybrids
101
How is bond length affected by the amount of s-character?
The more s-character, the shorter the bond will be
This is because s-electrons are held closer to the nucleus
102
What will form shorter bond lengths: sp³ hybridisation or sp hybridisation?
sp hybridisation as there is more s-character
103
What is resonance?
When two forms do not have an independent existence and are not an equilibrium
104
What is the main limitation of the Valence Bond Theory?
It is not always correct e.g. we would predict sp² hybridisation with an O=O bond and two lone pairs on each oxygen, but we know this is wrong as O₂ is paramagnetic (it has two unpaired electrons)
