Topic 2 - Chemical Bonds Flashcards
(126 cards)
1
Q
What is the wave function?
A
Wavefunction: The probability of finding the electron in a particular area.
- The reason why the wave function is used is because electrons show wave-like properties. This means they don’t always have a fixed position (Wave is not a single particle). Hence, the wave function informs us of the probable locations of the electrons.
Note –> diagram shows different shades –> representative of different probabilities.
2
Q
How does one represent the wave function?
A
Mathematical description –> Ψ (x,y,z) –>Three axis as we are dealing with three dimensions.
3
Q
How does the graph of the wave function look like?
A
4
Q
What is the electron density?
A
Electron density (Ψ2)
- Obtained by squaring the wave function
- Gives the number of electrons in a given area.
5
Q
What is the graphical representation of the electron density?
A
6
Q
What is the radial electron distribution?
A
Radial electron distribution (4πr2)(Ψ2)
- Multiply the volume of a sphere with electron density gives the e- distribution in a given volume. Basically, gives the probability of finding an electron in a given ‘layer’ relative to the nucleus.
- Equation provides us with the most likely radius where e- are found –> represented by maximum on graph.
7
Q
What does the graphical representation of the radial wave function look like?
A
8
Q
What is the wave function + electron distribution of the 2s and 2p orbitals?
A
Note –> shading/ ± is used to represent wave function direction, not charge.
9
Q
What is the equation for electrostatic potential energy?
A
Electrostatic potential energy –> Coulobmb’s Law
E α (q1 x q2)/(r)
which is equivalent to….
E = K (q1 x q2)/(r)
Where…
q = charge
r = seperation distance
k = constant
10
Q
What can coulomb’s law (Electrostatic potential energy) be used for?
A
When two objects with charge q1 and q2 respectively are located r distance away from each other we can calculate the electrostatic potential energy using the constant K.
- Greater charges/smaller distance –> greater E (visa versa)
Results:
- Positive E –> forces are repulsive
- Negative E –> forces are attracting
- Neutral object –> 0 interaction.
11
Q
What is the Aufbau principal?
A
Electrons are placed in orbitals starting with the lowest energy, working up.
12
Q
What is the Pauli exclusion principle?
A
It states that there are a maximum of 2 electrons per orbital and they must spin in opposite directions.
13
Q
What is Hund’s rule?
A
It states that when multiple orbitals of the same energy are available, electrons are distributed among them and spin parallel (before pairing electrons).
14
Q
What are molecular orbitals?
A
Molecular orbitals explain what happens to the orbitals when a molecule is formed.
They are helpful in explaining a specific type of bonding (delocalization) and properties (magnetism) of molecules which the Lewis structure doesn’t explain.
- Atomic orbitals are associated with atoms
- Molecular orbitals are associated with molecules –> are spread across the entire molecule.
15
Q
In what situations are molecular orbitals formed?
A
Atoms only bond to form a molecule when there are more favourable interactions than unfavourable interactions.
For example: H2 –> 4 favourable and 2 unfavourable –> H2 molecule is formed.
16
Q
What determines whether a successful bond is formed?
A
When talking about the molecular orbitals - a successful bond depends on whether the wave functions of the two atoms are in-phase or out of phase.
In-Phase –> Succesful bond
Out of phase –> No successful bond
Note:
In-Phase –> Waves combine together –> amplitude of waves is added.
Out of phase –> Waves cancel each other out –> destruction of the wave.
17
Q
What do the A.O wave function graph, M.O wave function graph, radial distribution and electron density look like for H2.
A.O –> Atomic orbital (before combined)
M.O –> After orbitals are combined.
A
In phase
- Bond is formed –> cylindrical symmetry
- Bond is always at a fixed distance - energetically favourable.
- Known as bonding molecular orbitals.
Out of phase
- No bond formed (also cylindrical symmetry).
- Known as an anti-bonding orbital
18
Q
Do the number of atomic orbitals and molecular orbitals need to equal each other?
A
Yes, the number of atom orbitals and molecular orbitals must equal each other.
19
Q
What do the energy level diagrams of H2 look like? What can they tell us about the bonding taking place?
A
The molecular bonding orbital is at a lower energy than the individual atomic orbitals –> hence, the formation of the M.O is energetically favourable –> provides stability.
20
Q
What do the energy level diagrams of He look like? What can they tell us about the bonding taking place?
A
In He’s case both of the bonding and anti-bonding orbitals get filled.
Stabilisation gained by the bonding orbitals is counter-balanced by the destabilisation of the anti-bonding orbital. No energetic reason to He2 –> hence it is not found in nature in this form.
21
Q
Can the molecular orbital theory be extended to π orbitals?
A
Yes, the M.O theory can also be extended to π bonds.
22
Q
Draw the bonding and anti-bonding orbitals for the sigma bonds formed by pi bond head-on overlap.
A
23
Q
Draw the bonding and anti-bonding orbitals for the formation of π bonds.
A
Important to realize that the Pi bond isn’t as strong as the sigma bond. This results in destabilisation of the of π* orbital having a lower energy than σ* .
24
Q
Draw the three different bonds that can be formed in an O2 molecule.
Hint - Think about different Pi overlaps.
A
The geometry of molecular orbitals in O2.
- Head on collision of 2p orbitals –> sigma bond
- Vertical side-ways overlap of 2p orbitals –> Pi bond
- Horizontal side-ways overlap of 2p orbitals –> Pi bond.
25
What does the orbitals size correspond to?
The orbital size correspond to the point where there is the highest probability of finding electrons.
26
When 2p orbitals are involved in bonding what happens to the 2s and 1s Orbitals?
When 2p orbitals are bonding the 2s and 1s orbital can't participate in bonding as they lack the sufficient size --\> they can't physically reach.
27
Draw the energy level diagram for O2 .
28
The formula for calculating bond order?
Bond order = (e- in bonding orbitals - e- in anti-bonding orbitals)/2
Bond order --\> gives the net number of shared pairs of electrons. i.e. A double bond corresponds to a bond order of 2.
29
Explain the bonding in methane (reference hybridisation)
According to the atomic structure of carbon, carbon wouldn't be able to form the tetrahedral shape of methane.
This is overcome by the hybridisation of orbitals to form Sp3 hybrid orbitals.
1. One electron in 2s orbital promoted to 2p
2. 2s and 2p orbitals hybridize to form Sp3 orbitals.
Each orbital has one electron so it can now form 4 sigma bonds with hydrogen --\> this has a tetrahedral arrangement (109.5o).
30
What is a Fisher projection?
A diagram that shows 3D nature of the molecule.
- Atoms coming out from the side --\> go out of the plane of the paper.
- Atoms at the top and bottom go into the plane of the paper.
31
Explain the bonding in ethene (reference hybridisation).
The carbons in ethene need to form 3 sigma bonds and one double bond. Orbitals need to be hybridized to Sp2 hybrid orbitals in order to make this possible.
1. The electron is promoted to P orbital
2. Orbitals hybridize to form Sp2
- Sp2 form sigma bonds and the extra P orbital forms Pi bond.
- Note we would expect the electron in the P orbital to move down to the Sp2 orbital but the orbitals are close in terms of energy that pairing electrons is more energetically unfavourable.
32
Explain the energy level diagram for the bonds in methane.
33
Explain the energy level diagram for the double bond in ethene.
34
What is the average bond energy + bond length of a single and double bond?
Single bond (σ) --\> 350 KJ/Mole --\> 1.54 Å (0.154nm)
Double bond (σ + π) --\> 600 KJ/mole --\> 1.34 Å (0.134nm)
As we can see from this a Pi bond is roughly 50 kJ/mole weaker than a sigma bond.
35
Definition of an isomer?
Isomer --\> same molecular formula but a different structural formula (different arrangement of atoms)
36
Structural isomer definition? Characteristics?
Same molecular formula but a completely different arrangement of atoms.
1. Chemically different
2. Same number of bonds
3. Composition is the same --\> different arrangement
37
What is geometric isomerism?
This type of isomerism involves the restricted rotation of a bond due to the presence of a double bond or ring structure. Results in the formation of Cis (same) and trans (opposite) isomers.
The only way to interconvert would be to break the bond and rotate the bond by 180 degrees --\> This process requires energy as electrons need to be returned to their atomic orbitals (higher energy level).
38
What are different types of stereoisomers?
Stereoisomerism: Molecules that have the same molecular formula but have a different arrangement of the atoms in space.
Types
1. Configurational --\> interconverted by breaking bonds
a) Optical Isomerism
b) Geometric isomers
C) Diastereomers
2. Conformational --\> no bonds need to break
39
When are two chiral molecules enantiomers?
They are enantiomers (isomers of each other) if they are mirror images and non-superimposable.
Note - Chiral centre must have 4 different groups attached.
40
What are diastereomers?
These are stereoisomers that are not mirror images and non-superimposable.
Both molecules have different chemical and physical properties.
41
What's the difference between absolute and relative stereochemistry?
**Absolute stereochemistry**: is the precise arrangement of atoms in space --\> use the stereochemical description of R and S.
**Relative stereochemistry:** Compares the arrangement of atoms in space of one compound with those of another. Use the stereochemical description of L and D.
Basically comparing and matching similar functional groups to figure out whether something is L or D.
42
Explain the stereoisomerism of glucose.
All carbons on the glucose molecules are diastereomers apart from one which is an enantiomer.
- Change in the position of the -OH only takes place on one carbon which is carbon 4 --\> involves breaking bonds to interconvert.
CHECK SLIDE!!!
43
What is conformational isomerism?
Conformational isomers (or conformers or rotational isomers or rotamers) are stereoisomers produced by rotation (twisting) about σ bonds.
44
What are the different conformations?
There are a total of 4 different conformations that one must consider.
1. One must examine the molecule in order to see whether it is eclipsed or staggered.
2. If it is staggered one must examine the bond angles between the two groups (usually examine the atom with the most protons --\> repel the most --\> further away --\> more stable).
Furthest angle --\> Called Anti
Anything else --\> Gauche
45
Differences in in eclipsed and staggered on an energy level diagram?
46
Differences between anti and gauche on an energy level diagram?
47
What is a Newman projection?
Diagram used to show different conformations.
Note the angle between the groups on the different carbons --\> known as a torsion angle.
48
Explain the dissociation of H2O and how the value of Kw is derived.
Dissociation of H2O
H2O \<---\> H+ + OH-
Equilibrium expression:
Kw = Keq = ([H+][OH-])/(1) = 10-14
Note that H2O is ingored --\> concentration is so large that change have minimal impact on Equilibrium constant.
Hence, in a neutral solution....
[H+][OH-] = 10-7
49
What's the definition of pH?
pH = -log[H+]
50
Write out the equation for the dissociation of a proton from acetic acid and create an expression for Ka.
CH3COOH \<-----\> CH3COO- + H+
Acid \<-----\> Conjugate base (acetate) + proton
Ka = ([CH3COO-][H+])/[CH3COOH]
51
From a general expression of Ka, derive the Henderson Hassel bach equation.
52
When will the pH = Pka in the Henderson Hasselbach equation?
When there are equal concentrations of Acid and conjugate base.
Log[A-]/[HA] = 0
Which means that...
pH = Pka
53
Using the Henderson-Hasselbach equation figure out whether the protonated or deprotonated form of aspartic acid is preferred at physiological pH.
Physiological pH = 7.2
pKa of aspartic acid = 4.2
7.2 = 4.2 + log [A-]/[HA]
3 = log [A-]/[HA]
103 = [A-]/[HA]
Ratio is 1000:1
For every 1000 A- there is 1 HA.
Hence, the most common form of the acid is the aspartate ion.
54
Explain using orbitals why aspartate is preferred instead of aspartic acid.
The carboxylate ion forms a resonance structure --\> Charged is distributed over a larger area --\> dispersed charge --\> increase stability --\> explains why aspartate is preferred.
55
Using the Henderson-Hasselbach equation figure out whether the protonated or deprotonated form of lysine side chain is preferred at physiological pH.
pH = 7.2
pKa = 10.2
7. 2 = 10.2 + log[B]/[BH+]
- 3 = log[B]/[BH+]
10-3 = [B]/[BH+]
103 = [BH+]/ [B]
For every 1 unprotonated lysine there are 1000 protonated lysine molecules.
Amino acid side chain will be positively charged at physiological pH.
56
Explain the charge distribution in the arginine side chain.
Basically what happens is that the C=NH group bonds to H+ to form C=NH2+ --\> subsequently the electrons from a lone pair from another nitrogen move towards positive charge --\> breaks double bond --\> but nitrogen that has now lost a lone pair needs to form a double bond --\> Double bond moves between amine groups.
57
Explain why arginine is more basic than lysine.
It is more basic as it more readily accepts the proton because it can form a resonance structure --\> distributes the charge --\> more stable.
58
Explain how phosphate is able to form 5 bonds in phosphoric acid.
1. Because the orbitals in the third energy level are similar in terms of energy + electrons paired in 3s orbital experience repulsion --\> the electron from the 3s promotes to the 3d orbital (as shown in picture).
2. The 3s and 3p orbitals hybridize --\> sp3 hybrid orbital
Now the phosphate can form 4 sigma bonds and 1 pi bond.
59
Explain the resonance system in phosphate ions.
Once phosphoric acid has become completely deprotonated it can form resonance structure --\> the double bond moves between all the P --- O bonds --\> disperses the negative charge --\> increases stability.
Hence, all oxygens have a partial negative charge.
IMPORTANT --\> favours the hydrolysis of ATP to ADP and Pi can form stable resonance.
60
What does a titration graph between a strong acid and base look like?
61
What does a titration graph of a weak acid and strong base look like?
62
Where is the buffer region of a weak acid?
The buffer is represented by the plateau on the graph before the sharp increase.
It occurs when the pH = pKa --\> good buffer region (plus minus 1 pH).
63
Why is a weak acid a good buffer when the pH = pKa?
This is because the concentration of both acid and conjugate base are equal. Hence, these molecules will be able to absorb any H+ or OH- added to the solution.
64
What are two important points to remember about a weak acid buffer?
2 main points
1. Weak acid can act as a goof buffer between plus minus 1 pH of pKa.
2. pH of the buffer is independent of dilution --\> this is because when diluted the ratio of protonated to unprotonated remains the same --\> no change in pH.
However, there is a decrease in buffering capacity.
65
What does the titration graph of phosphoric acid look like?
66
How does the fraction of the different protonated/deprotonated phosphoric acid molecules change when pH changes?
1. pH = pKa = 2.1
50: 50 ---\> H3PO4 and H2PO4-
2. pH = pKa = 6.9
50: 50 ---\> H2PO4- and HPO42-
3. pH = pKa = 12.4
50: 50 ---\> HPO42- and PO43-
67
What is the dominant species of Phosphoric acid when pH is equal to 7.2.
Examining the ratio of
H2PO4- and HPO42- at a pH = 7.2
7.2 = 6.9 + log [HPO42-]/[H2PO4-]
100.3 = [HPO42-]/[H2PO4-]
2 = [HPO42-]/[H2PO4-]
ratio is 2:1
For every H2PO4- you will have two HPO42-
68
What are the pKa values for the oxygens on phosphate?
Pka for the bottom -OH --\> 1.0 --\> 100% of the time deprotonated.
Pka for side -OH --\> 6.1 --\> 90% of the time will be found deprotonated
69
How can amino acids act as a buffer?
Amino acids are Zwitterions --\> NH2 forms NH3+ and COOH forms COO-
Hence, amino acids can absorb/neutralize both H+ ions and OH- ions --\> important for amino acids floating freely.
70
How can a weak acid buffer be created?
Weak acid buffer can be created by mixing equal weak acid and its salt --\> weak acid high concentrated --\> salt fully dissociates --\> conjugate base highly concentrated.
Weak acid + C.B --\> Good buffer.
71
What is the relationship between frequency and wavelength?
λ = C / ν
Wavelength = Speed of light / Frequency
This shows us that the frequency and wavelength are inversely proportionate.
72
What the relationship between energy and wavelength?
Because we know now that Energy = h/ν
We can derive the following equation:
E = (hc)/λ --\> energy is inversely proportionate to the wavelength but proportionate to the frequency.
73
What can happen to electrons in bonding orbitals when provided with energy?
The electrons can promote and move into the anti-bonding orbitals.
74
How can one calculate the absorbance? What are the different results?
Absorbance = log (IInput/Itransmitance)
This is the log of the ratio of light that entered versus the light that actually passed through.
The log compresses the scale to small positive numbers. Thus...
Transmittance = 100% = 1.0 --\> Absorbance = 0
Transmittance = 10% = 0.1 --\> Absorbance = 1
Transmittance = 1% = 0.01 --\> Absorbance = 2
Absorbance of 2 is usually the limit as the spectrophotometer can't detect smaller quantities.
75
In reality, do orbitals have many different energy levels?
Yes, this is because the orbitals have different vibrational energy levels.
Consequently, when electrons are promoted, a whole range of wavelengths are absorbed not just a specific wavelength --\> results in a curved shaped graph.
However, there is a peak absorbance --\> show wavelength that corresponds to the most frequent energy transition.
76
What is the relationship between double bonds and absorption?
The more double bonds present --\> the more absorption takes place in the light spectrum.
77
How can absorbance be used to identify specific ring structures?
Ring structures absorb a particular wavelength which can be identified.
78
What is the Beer-Lambert law state?
It states that....
A = E x c x I
A = Abosrbance
E = Molar extinction coefficient (value derived from graph --\> as it changes depending on wavelength).
(M-1 cm-1)
c = concentration (Molar)
I = Path length (cm)
79
How does the property of fluorescence work?
80
Examples of heterocyclic aromatic systems?
Hetero - meaning containing something else than carbon (N/O).
1. Pyridine
2. Pyrimidine
3. Pyrole
81
Explain the ring structure in pyridine?
Each atom in the ring provides 1 electron to the delocalised system --\> results in aromatic ring formation.
82
Explain the ring structure in Pyrimidine?
Each atom provides one electron to ring structure --\> allows for aromatic ring formation --\> Nitrogen still have two lone pairs.
Note that all the atoms are Sp2 hybridised which matches the 120 bond angle of the ring.
Thymine + uracil both have this ring structure --\> can be detected using UV.
83
Explain the ring structure in Pyrole?
Each atom provides one electron to ring structure except for nitrogen which provides 2 electrons (gets it from the lone pair) --\> allows for aromatic ring formation.
Note that all the atoms are Sp2 hybridised which doesn't match the 108 bond angle of the ring --\> makes the ring structure less stable.
84
What is the ratio of protonated to unprotonated histidine at physiological pH?
pH of 7.2 and pKa of 6.6.
80% unprotonated
20% protonated
85
Draw the different resonance and tautomers of the histidine ring.
86
Why do atoms bond together in the first place?
Atoms bond with one another in order to gain stability --\> electrons distribute themselves within the atomic orbitals in arrangements that have the lowest possible energy --\> to give the atom the greatest stability.
When chemical bonds form between atoms to generate compounds --\> valence electrons redistribute themselves so that the atoms gain full valence shells.
87
Do you consider all orbitals when talking about full valency\>
No --\> refer to the S and P orbitals --\> only concerned with those 8 electrons --\> this is where the Octet rule was derived from (caution --\> not an absolute rule --\> applies to atoms up to period 2.
88
Definition of electronegativity?
It indicates how strongly an atom of that element can attract an electron.
89
The link between electronegativity and bonding?
High electronegativity --\> complete transfer of electrons from one atom to another --\> one atom has the 'strength' to completely tear an electron away from the neighbouring atom --\> Ionic bonding
Lower electronegativity --\> atoms exert a similar pull --\> complete transfer is unlikely --\> neither atom has the strength to tear of electrons from the other --\> covalent bonding.
The difference in electronegativity greater or equal to 1.7 --\> ionic bonding --\> difference less --\> Covalent.
90
How do bonding and anti-bonding orbitals differ when it comes to covalent bond formation?
Bonding and anti-bonding orbitals have opposing effects on the formation of a covalent bond.
A pair of electrons occupying the bonding orbitals facilitates covalent bonding.
A pair of electrons occupying anti-bonding orbitals inhibits covalent bond formation.
91
What is a hypervalent atom? What determines whether an atom can expand its octet?
An atom with more than eight valence electrons
The answer may lie in the presence or absence of d orbitals in the valence shell of the element being considered.
Good biological example --\> Phosphorus --\> uses a vacant 3d orbital to hold some of its valence electrons.
92
Do atoms of the same element show different valencies?
Yes, some elements can exhibit different valencies and therefore contribute to different numbers of covalent bonds.
93
What is a conjugate system?
A conjugate system is one which contains electrons that are delocalised.
It arises from the overlap of adjacent P orbitals in any molecule featuring a sequence of alternating single and double bonds.
94
Can delocalisation occur in non-conjugate system?
Yes, delocalization of electrons in compounds that do not possess a conjugate system of bond does occur. It happens when there are two or more different arrangements of bond --\> ozone.
95
What are the 6 key types of non-covalent interactions?
1. Dispersion forces
2. Permanent dipolar interactions
3. Steric repulsion
4. Hydrogen bonds
5. Ionic interactions
6. Hydrophobic forces.
96
Are non-covalent interactions significant?
Not significant when examining the interactions as single events/in isolation but overall the net effect of many non-covalent interactions become noticeable and significant.
97
Difference between inter and intramolecular forces
1. Intramolecular interactions operate between separate parts of the same molecule --\> within the molecule.
2. Intermolecular interactions --\> operate between different molecules.
98
What is a permanent dipole?
Permanent dipole --\> a molecule that has a permanently partially positively charged region and a permanently partially negatively charged region.
Note --\> polar bonds can cancel each other out if the molecule is symmetrical
99
What are dispersion forces?
A dispersion force is a force of attraction that exists between the two areas of opposite charge, which form the induced dipole.
Occurs due to the random movement of electrons --\> creates temporary dipole --\> induces dipoles in the surrounding molecule --\> attraction.
100
What two factors influence the prevalence of dispersion forces?
1. Shape --\> Planar flat molecules are able to associate more than molecules with irregular shapes --\> as dispersion forces are stronger in close proximity --\> when the molecules are able to closely associate --\> stronger interactions.
2. Size --\> Larger molecule --\> more atoms --\> more electrons --\> more/larger induced dipoles may exist.
101
What are permanent dipolar interactions?
Permanent dipolar interactions are the attractive forces that exist between opposite partial charges on polar molecules.
- They are permanent/last longer
102
What are induced dipolar interactions?
Electrostatic interactions that arises when a permanent dipole on a polar molecule induces a temporary dipole on a non-polar molecule --\> creates an electrostatic force of attraction between both molecules
103
What is steric repulsion?
Steric repulsion --\> When two molecules approach each other --\> the molecules on their surface first interact --\> both groups of electrons will both carry a negative charge --\> thus these groups repel.
104
Definition of van der Waals forces?
A term used to describe the overall interaction between two molecular species once both the attractive dispersion forces and dipolar interactions and the repulsive forces of steric repulsion are taken into account.
105
What are hydrogen bonds?
Hydrogen bonds represent a special class of dipolar interactions --\> strong interaction between the hydrogen on one molecule to an electronegative atom either on another molecule or a spatially distinct part of the same molecule.
106
What are two critical factors needed for hydrogen bond formation?
1. Hydrogen atom must be bonded to an electronegative atom (Oxy, Nirto, fluorine) --\> compound containing this arrangement is known as the **hydrogen bond donor.**
2. The hydrogen atom can only form a hydrogen bond with a non-bonding pair of electrons on another electronegative atom --\> compound containing this arrangement is known as the **hydrogen bond acceptor**.
107
What are ionic forces?
Ionic forces are the attractive forces that exist between ionic species --\> species carrying a fully positive and negative charge.
This is not limited to ionic compounds --\> also important for covalent molecules with a full positive or negative charge.
108
What is a salt-bridge?
Ionic forces that operate between two oppositely charged amino acid side chains in a protein is called a salt bridge.
An important force that stabilizes 3D structures.
109
What are hydrophobic forces?
Hydrophobic interaction occurs when structural rearrangements move the hydrophobic portions of a molecule away from their aqueous surroundings.
Once clustered together the hydrophobic portions are stabilized through a network of dispersion forces that operate between them.
Basically --\> hydrophobic forces drive the shielding of hydrophobic molecules from their aqueous surroundings.
110
What three key factors determine the shape and structure of a molecule?
1. Bond Lengths
2. Bond Angles
3. Bond rotation
111
Definition of bond length?
The distance between the nuclei of two covalently bonded atoms.
Dictated by:
1. Atomic radii of the atoms joined together
2. Type of covalent bond
112
Definition of the atomic radius?
Usually defined as half the distance between two covalently bonded atoms when the two atoms are identical.
113
What two main factors influence atomic radii?
1. Valence shell number increases --\> atomic radius increases.
2. An increasing number of valence electrons enter the same valence shell --\> atomic radius decreases slightly --\> more valence electrons --\> greater attraction between nucleus and valence electrons.
114
What does the VSEPR theory state?
The valence pairs surrounding a central atom arrange themselves in such a way that they minimize repulsive forces that operate between them --\> basically they position themselves as far away as possible --\> decreases interaction --\> increased stability.
Applies to non-bonding and bonding pairs.
115
Why do non-bonding electrons distort the expected geometry?
The distortion in geometry arises because non-bonding pairs occupy more space than bonding pairs --\> so surrounding electrons pairs have to move slightly further away in order to evade repulsive forces from the non-bonding pair --\> this relatively larger push exerted by non-bonding pairs results in a slight distortion of the bond angles.
116
What is the hybridisation of atomic orbitals?
The hybridisation of atomic orbitals --\> a change in shape of the orbitals which allow the electrons they contain to adopt new geometries.
Helps the explain the geometries seen using the VSEPR theory.
117
Difference between configuration and conformation?
Configuration --\> describes how its composite atoms are joined together in three-dimensional space --\> the configuration is fixed and can only be changed by breaking bonds.
Conformation --\> Describes the overall three-dimensional structure at a particular moment in time --\> this is variable --\> applies to single bonds --\> double/triple bonds can't rotate.
118
What is steric hindrance?
There are circumstances where bond rotation around a single sigma bond is not possible --\> rotation is blocked because different groups of atoms simply get in the way of each other.
119
Can delocalised systems of electrons rotate?
When delocalization occurs, the electrons are Pi-bonding electrons that are shared between more than two atoms --\> shares both single and double bond character.
But they are UNABLE to rotate.
120
Can bonds in cyclic molecules rotate?
No, they are unable to rotate --\> they lack the flexibility to allow it to move into and adopt different conformations.
121
What do the values of Ka and Kb tell you about acid/base strength?
Large Ka --\> Strong acid
Small Ka --\> Weak acid
Large Kb --\> Strong Base
Small Kb --\> Weak base
Note --\> there is an inverse relationship between the strength of a base and the value of Ka
- Weak base --\> large Ka and small Kb
- Strong base --\> small Ka and large Kb
122
The important relationship between Ka, Kb and Kw?
Multiplying Ka and Kb together will always equal Kw (1x10-14).
This relationship tell us that if we know the value for Kb of a weak base, we can calculate Ka, the acid dissociation constant for the conjugate acid (Visa Versa).
This is true for all weak acids and bases.
123
Impact of changing pH on the dissociation of a weak acid?
- A decrease in pH leads to a decrease in the extent of dissociation of a weak acid.
- An increase in pH leads to an increase in the extent of dissociation of a weak acid.
Due to Le Chatelier's principle --\>, for example, a decrease in pH means there are more H+ in solution --\> drives the reaction away from the dissociation of the weak acid.
124
An important point to remember when pH=pka
When pH equals the Pka --\> acid and conjugate base are found in equal concentrations in solution.
pH \> Pka --\> there will be more dissociated acid in solution.
pH \< pKa --\> there will be more undissociated acid in solution.
125
Important points to remember when making a buffer.
Two types of buffer
1. Weak acid and the salt of its conjugate base
2. Weak base and the salt of its conjugate acid
Requirements
- We must use an acid-base conjugate pair
- Must use a weak acid or base
- Both compounds must be present at sufficiently high concentrations.
126
How can we calculate the pH of the buffer solution?
Using the Henderson-Hasselbach equation.
