Autumn semester Flashcards
1
Q
How is the periodic table arranged?
A
Elements are sequenced in increasing atomic number across the periods, and arranged so that elements with similar chemical properties fall in the same groups.
2
Q
What are the sap,d,f blocks of the periodic table?
A
s-block = groups 1-2
p-block = 13- 18
d-block = 3-12
f block = bottom part
3
Q
What is the composition of an atom?
A
The atom consists of negatively charged electrons which move around a central nucleus containing positively charged protons and neutrons.
4
Q
What is an isotope?
A
An isotope are elements with the same atomic number but have different mass numbers (differ with amount of neutrons)
5
Q
What is the octet rule in relation to the driving force for reactions?
A
The octet rule states that atoms will continue to share electrons until they have acquired an octet of valence electrons. The driving force of any reaction is the formation of the electronic structure of a noble gas.
Each orbital contains maximum 2 electrons where there is 1 s orbital, 3 p orbitals, five d orbitals and 7 f orbitals.
6
Q
What is the concept and importance of the valence shell?
A
The valence shell is the highest energy set of orbitals that contain the outer (valence) electrons. It is important as it is the orbital that is gaining or losing electrons.
7
Q
How do you identify valence electrons for an element and why are they important?
A
Valence electrons are the electrons listed after the last noble gas configuration (outer electrons)
Valence electrons are important is chemistry as in an ionic compound is tell us the oxidation state and number of bond formed.
8
Q
How is the structure of the atoms found in experiment?
A
Rutherford > conducted an experiment in which alpha (HE2+)particles where shot through a gold foil. He found that he vast majority of the ions went straight through with minimal scattering (proving high amount of space between electrons) however some where reflected back at large angles suggesting there presence of a larger charged mass (the nucleus).
Additional he tried to use classical mechanics to explain atom behaviour, this failed as the electrons gradually decayed and collided with nucleus.
Bohr > suggested that energy of an electron in a particular orbital was quantised, he proved this by exciting hydrogen particles which showed shared bands of light emitted showing well defined gaps.
9
Q
What is the concept of energy levels around an atom?
A
Electrons located around an atom are found in defined area called energy levels, they represent the 3d space surrounding an atom where electrons are most likely to be.
10
Q
What is the definition of ionisation energy?
A
Ionisation energy is the amount of energy needed to remove 1 electron from an atom. The ionisation energies where found to agree exactly with those measured experimental in the atomic spectrum of hydrogen.
11
Q
What is electromagnetic radiation?
A
A from of energy consisting of oscillating electric and magnetic fields which travel at the speed of light (c).
e.g visible light, microwaves, x-rays … ( they all have wavelengths)
12
Q
What is the equation relating to frequency, wavelength and velocity of light?
A
λv = c
λ = wavelength (m)
v = frequency (Hz/s^-1)
c = speed of light (2.998 x 10^8 m/s)
13
Q
How does flame test observations relate to atomic structure?
A
The flame test is used to visually determine the identity of an unknown metal or metalloid ion based on the characteristic color the salt turns the flame of a bunsen burner. The heat of the flame converts the metal ions into atoms which become excited and emit visible light. The colour emitted corresponds with the energy change of the gaps between energy levels.
14
Q
What is a photon?
A
In some situation it was found that the behaviour of light/radiation cannot always be thought of as waves but instead particles. Max Planck proposed that electromagnetic radiation could only be emitted and absorbed in quanta of radiation called photons.
The energy of a photon is proportional too its frequency.
15
Q
What is the equation relating the energy of a photon to frequency?
A
E = hv
E = energy (j)
h = planck’s constant (6.63 x 10^-34 js)
v = frequency (Hz or s^-1)
16
Q
What is the photoelectric effect?
A
The ejection of electrons from a material pm irradiation by light (Light hits the material and electrons are thus displaced).
The electrons are only ejected from the surface is the frequency of the radiation is above a threshold value called the work function.
17
Q
How is the photoelectric effect applicable in the phototube?
A
a vacuum tube containing a cathode made of a metal with a small work function so that electrons would be easily emitted. The current released by the plate would be gathered by an anode held at a large positive voltage relative to the cathode
18
Q
What is the de Broglie relationship and how is it used?
A
De Broglie pointed out that the energies calculated for a wave and for a particle must be equal for anything behaving as both showing wave-particle duality.
λ = h/p
λ = Wave length (m) (limiting factor)
h = Planck’s constant (6.63 x 10^-34 js)
p = momentum of particle (mass x velocity)
19
Q
What is the uncertainty principle?
A
As an electron has wave-like properties with wave length = order of magnitude of the size of an atom, you can determine the probability that an electron is in a certain place.
20
Q
What is the Heisenberg uncertainty principle?
A
An electron has a position in space defined by x,y,z coordinates and has a momentum (p=mv) parallel to each axis (px,py,pz)
thus :
ΔxΔpx ≥ h / 4π
Where Δx, Δpx are the uncertainties in measuring x and px, we can know the position of an electron thus the less we know about its momentum (and vis versa).
21
Q
Example of the Heisberg uncertainty: calculate the minimum uncertainty in the position of an electron (mass = 9 x 10^-31 kg) moving at 3 x 10^5 m/s with an uncertainty in the velocity of +/- 10^2 m/s?
A
Δx = h / 4π x Δpx
Δpx = mass x velocity
= (6.63 x 10^-34) / 4π x ((9x10^-31)x(2x10^2)) = 293 nm
22
Q
What is the concept of a wave function and why is it used?
A
The wave function tell us how likely an electron is in a particular location at a given time its denoted by the symbol 𝚿 (psi). It is calculated using Schrodinger equation.
23
Q
What is the physical significance of the wave function?
A
Max Born suggested that the probability of finding a particle in any region of space is proportional to 𝚿^2.
Thus in a region were 𝚿^2 is large, the probability of finding and electron is high and vis versa.
(𝚿 = 0 there is a node)
24
Q
What is the Schrodinger equation?
A
Schtodinger used wave function to produce a mathematical equations to calculate the behaviour of an atom however it can only be solved for two body problems ( a system containing only a nucleus and an electron).
H𝚿 = E𝚿
𝚿 = R (r)Y (θ, Ф)
n,l l,ml
R =
25
What are the answers that can be obtained from the Schrondinger equation?
- the allowed wavefuntions for the particle
- solutions are only possible for certain energies
- the probability density for a particle at any point is proportional to the 𝚿^2 at that point.
- each allowed coltuion for the hydrogen atoms defines an allowed atomic orbital
26
What is the component symbols of the wave function?
R = radial wave function
Y = angular wave function
n = principle quantum number
l = angular momentum
ml = quantum number
27
What are the definition of all the quantum numbers and their allowed values?
n = the principle quantum number it determines the size of the orbital and for a hydrogen atom the energy. Its allowed values are (1,2,3,4,5 ....)
l = orbital angular momentum quantum number it determines the shape of the orbital (any whole number from 0 to (n-1).
ml = magnetic quantum number determines the orientation in space of the orbital (excluding in a magnetic field). Its values re any whole number between -l to +l.
28
How does the allowed value of the quantum numbers to the structure of the atomic orbitals?
Together the quantum numbers provides information of the properties of an atomic orbitals (the number and type of orbitals possible)
The type of orbital:
l = 0 (s orbital) l = 1 (p orbital) l = 2 (d orbital) and l = 3 (f orbital)
l = n - 1
e.g.
for n = 1, l = 0 and ml = 0 meaning that the orbital is a singular 1s orbital
for n = 2, l = 1 and ml = 0 so the orbital is one 2s but l can also be l = 1 and ml = -1,0,1 so the orbital is three distinct 2p orbitals.
- if the orbitals have the same value of n = same shell, if both n and l are the same = same sub shell
29
How can radial and angular wave functions be used to plot three dimensional shapes of the orbitals?
The wave function 𝚿 varies with position, this variation means that 𝚿 is a function of x, y and z. As atoms are spherical so 𝚿 can be represented as a function of one distance r and two angles θ, Ф (spherical coordinates) .
𝚿 = R (r)Y (θ, Ф)
n,l l,ml
30
What is radial wave function?
Shows the probability of finding an electron in a spherical shell at a distance r from the nucleus, this is found by plotting 4pir^2R^2
31
What is angular wave function?
Shows the shape and orientation of orbitals. (overall shape depends entirely on l and ml). Angular nodes are presents no of nodes = l.
32
How do you interpret useful information from radial probability function plots?
- As n increases from a given value of l the size increases
- there is a very small probability of finding the electron long way from the nucleus so accurate size definition is not easy
- there are areas called nodes where the probability of finding an electron in 0.
number of radial nodes = n - l - 1b e.g. 0 for 1s, 2p and 3d, 1 for 2s and 3p and 2 fro 3s.
33
What is the shapes and orientation of the s, p,d orbital?
s orbitals = these are spherical with the same phase of wave function across the whole boundary surface.
p orbitals = these are dungbell shaped with px,py and pz all have the same shape but different orientations and contains both negative and positive phases and one angular node where 𝚿 = 0 (2p) as n increases the number of nodal planes also increase.
d orbitals = have a clover like shapes with four lobes and two angular nodal planes.
34
What are the main types of chemical bonds?
- covalent bonds (electron sharing)
- ionic bonds (electron transfer)
- metallic bonding (positive metal ion held together by a sea of delocalised electrons)
35
What are the characteristics of the main chemical bonds?
- ionic bonds are formed between arose with low ionisation energies and high electron affinities
- covalent bonds are formed through the attraction between the shared pair of electrons and both positive nuclei
36
What is the definition of electronegativity?
The ability of an atom in a molecule to attract electron density to itself (ability to attracted the shared pair of electrons in a covalent bond)
- in a homonuclear bond the electrons are equally shared (H2)
- in a heteronuclear bond the sharing will be unequal (HCl)
The Pauline scale shows electronegativity increasing through a diagonally from left to right up the periodic table?
37
How does the periodic table aid in predicting types of chemical bonds?
The periodic table allows fro predictions of electronegativity to take place (Pauline scale)/
- no electronegativity creates a purse covlents bond (X-X)
- small to moderate electronegativity differences create polar covalent bonds (C - O with delta + and -)
- very large electronegativity differences make ionic bonds (K+Br-)
38
what are lewis structures and rules?
Lewis strutters allows us to determines the connectivity and formal charges of the atoms in a structure (not the shape).
- every atoms wants to achieve an octet of electrons
- each pair of shred electrons gives one bond
- often no more than 4 bonds to an atom
39
How do you draw lewis structures?
1. Add up total number of valence electrons (determine how may pairs of electrons can be used)
2. decide on the central atoms and draw it in the middle and surrounding with other bonding atoms
3. join the atom to central atoms with single bonds
4. keep adding remaining pair of electrons to form multiple bonds if appropriate
5. add any remaining electrons as lone pairs
6. check number of electrons in immediate surrounding and assign formal charges (circle around atom if e- inside circle = valence means no charge, if < valence e- = + and > valence e- = -)
40
Lewis structure example CO3^2-?
1. C = 4 valence e-
O = 6 valence e-
+ 2e- (add for negative charge) = 24 e- thus 12 electron pairs
2.central atom = C (lowest electronegativity)
3. sings bonds
4. 1 double bond between O and C as carbon is unlikely to have more than four bonds (can't expand octet)
5. remains e- pairs are lone pairs
6. 2 oxygens have 7e- in circle so both have negative charges the C and Oxygen attached with double bonds are neutral.
41
What are the main exceptions in lewis structures?
1. Some elements cannot achieve a full octet e.g. boron has 3 valence e- but can only form three bonds (need 5 to fit octet)
2. Elements in the second and subsequent rows can accommodate more than 8 electrons by expanding their octet e.g. SF6
3. Quite often their are several lewis structures for a given molecule or ion
42
What is the concept of resonance?
Resonance involves the blending of two or more lewis structures, where the atoms have the same relative positions but different electronic arrangement. the electrons are delocalised resulting in a resonance hybrid which is lower in energy.
e.g. CO^2 = there are three different positions of the double C-O bond (one shorter and two longer bonds) however experiments shows that all the C-O bonds are the same length showing resonance.
43
How do you use lewis structures to rationalise properties?
1. bond strength
N2O4 the lewis structure shows tray repulsion between positive charges between the conjoined N atoms led to a very weak N-N
2. Bond lengths
O-O distance in H2O2 is longer than F2O2 as H-O-O-H but F is more electronegative (pulls electron density) and has resonance forms making the bond longer in H2O2
3. explains stability
NCO^- = CNO-, NCO^- has one formal - charge and two resonance forms and CNO- haas three formal charges making it more unstable.
44
How do you use lewis structures to predict the most stable arrangements of atoms in a structure?
The most stable lewis structures keep
- formal changes to a minimum
- when formal charges are necessary they should be kept consistent with relative electronegativities
- like charges should be kept as far as possible for each other
45
What are coordinate bonds and their characteristics?
A coordination bond is where one atom (donor) gives both electrons to another (acceptor) to form a normal covalent bond. It is symbolised by an arrow.
46
What is meant by the term VSEPR?
VSEPR stands for Valence Shell Electron Pair Repulsion. It can be used to determine the 3D shape of a covalent molecule.
47
How can VSEPR theory be used to determine the appropriate system for molecules and ions?
1. determine the lewis structure for the molecule
2. count valence electrons
48
What are the names and shapes of the basic 3d shapes (ABn) of covalent molecules?
1. AB2 = linear
2. AB3 = trigonal planar
e.g BF4 e- from B = 3, e- from B-F = 3 resulting in 3 e- pairs.
3. AB4 = tetrahedral
e.g. CH4 e- from C = 4, e- form C-H bond = 4 so four electron pairs.
4. AB5 = trigonal bipyramid/square based pyramid
e.g. PF5 e- from P = 5 and P-F bonds = 5 so 5 electron pairs giving trigonal bipyramidal
5. AB6 = octahedral
e.g. SF6 electrons from S is 6 and S-F bond = 6 so 6 electron pairs.
49
How to ions impact the shape of a molecule?
For molecules that have a permeant charge, the electrons from the central atom and the electron from each bond are still counted however for every + charge 1 electron is subtracted and for every negative charge 1 electron is added.
e.g. electrons from N = 5 and electrons from N-H bond = 4, - 1e- for positive charge on central atom making it AB4 and tetrahedral.
50
How do lone pairs rationalised the shapes of molecules and ions?
If some of the central atoms electron do not participate in binding they are called lone pairs, they affect the shape as repulsion from lone pairs are greater that the bonded pairs as they are located closer to the central atom so must be counted as a region of electron density.
51
What are the main shapes of molecules/ions with lone pairs?
AB2L = v -shaped, based on trigonal planar
AB3L = pyramidal, based on tetrahedral
AB2L2 = v shaped, based on tetrahedral
AB4L = seesaw, based on trigonal bipyramidal, with lone pair in equatorial position.
AB3L2 = T shaped with lone pairs at equatorial positions
AB2L3 = linear, repulsion cancel each other out
AB5L = square based pyramidal
AB4L2 = square planar
lone pair of electrons with occupy the position with the least amount of repulsion.
52
What are the main bond angles with each shape?
Trigonal planar = 120 degrees
Tetrahedral = 109.5 degrees
Trigonal bipyramidal = ax and eq is 90 degrees and eq and eq is 120 degrees
Octahedral = 90 degrees
53
How do lone pairs modify bond angles?
Repulsion from lone pairs are greater that the bonded pairs as they are located closer to the central atom and thus take up more space this pushes the bonded pair of electrons closer together and lowers the angle between them.
BP - BP < BP-LP < LP-LP
54
How can you apply VSEPR theory with molecules containing coordination bonds?
VSPER theory is applied as normal and the shape of the molecule is not affected by the presence of coordination bonds.
55
How can you apply VSEPR theory with molecules containing multiple bonds?
When applying VSPER theory only the first bond between two atoms significantly affects the the structure, so the shape should be deduced assuming the molecule only contains single bonds. The electron count is then
1. valence electrons at central atom
2. valence electron for every single bond present
3. -1 electron for every double bond and -2 electrons for every triple bond.
however like lone pairs multiple bonds take up more space around the atom so can make the bond angles smaller than expected.
56
How can you apply VSEPR theory with molecules containing resonance?
VSPER theory should be applied to one of the resonance forms as normal, the double bond is delocalised due to the resonance meaning all bond lengths are the same.
e.g. CO3^2-
3 electrons from C
3 electrons from C-O bonds
-1 electron for double bond
6 total electrons
3 electron pairs
so trigonal planar shape with bond angles off 120 degrees.
57
What are the shapes of extended structures of carbon?
The extended strutters of carbon are diamond and graphite. Diamond contains tetrahedral carbons atoms in an extended network of covalent bonds (one of the hardest natural substance).
Graphite contains trigonal planar carbon atoms arranged in layers held together by weak intermolecular forces (very soft and electrical conductor)
58
What are the three main states of matter?
Gases > molecules are in a constant radon motions relatively in free spaces, far apart, easily compressible and low amount of order
Liquids > molecules at constant random motion, molecules closer together and have more order
Solids > molecules are in a fixed position and are close together with small amount of vibration and highest level of order.
59
What are dispersion dipoles and their characteristics?
Instantaneous - induced dipole > molecules do not have a localised pair of electron and instead can be pictured as 'clouds' of electrons density which move and can create an uneven distribution of electron density creating an instantaneous dipole, which then induces dipoles of neighbouring molecules, the two molecules are then attracted together by electrostatic interactions.
Characteristics
- exist for all molecules
- always attracts independent of molecular orientation
- increase with molecular weight (higher amount of electron make stronger polarisation)
- more compact molecules are less easily polarised (unbranched alkanes have more surface points of contact and thus stronger forces)
- very variable in strength
60
What are permanent dipoles and their characteristics?
If two electrons are bonded together with different electronegativities, a polar bond is created giving a permanent dipole. Two molecules are then attracted together by electrostatic interactions.
Characteristics
- weaker then dispersion forces due to the orientation dependance of the permanent dipole
61
What are dipole-induced dipole forces?
A molecule with a permanent dipole can induce an instantaneous dipole in neighbouring molecules, the molecules are then attracted together by electrostatic interactions
62
What are hydrogen bonds?
Hydrogen bonding is a moderates attractive force that exist between a hydrogen atom (covalently bonded to a very electronegative atom) and a lone pair of electrons on another electronegative atom (N , O or F). They are the strongest type of intermolecular force.
63
How do you predict the type of intermolecular interactions presents in a compound?
Every compound with have dispersion forces, dipole-dipole forces will only be present in compound with atoms that differ greatly in electronegativity, however in large compound Cush as (CCl4) the polarity with cancel each other out.
64
What are the relative strengths of bonding and intermolecular interactions?
1. hydrogen bonding = 25 KJ/mol
2. dipole - dipole = 2-15 KJ/mol
3. dispersion = very variable in strength
65
How do you predict the formation of ionic systems?
An ionic system with for through the transfer of electron between a non-metal and metal through oxidation and reduction to form an ionic bond. An ionic bond is thus the strong electrostatic attraction between a cation and an anion.
66
What are the main properties of ionic compounds and why?
- high melting and boiling point
- soluble in water
- conductor of electricity as molten or liquid but not as a solid
67
How do you predict the solubility of compounds in solvents and why?
Solubility if a substance is dependent upon the intermolecular interactions that occur. The general rule is like dissolves in like (polar substance in polar solvent and non polar substance in non polar solvent).
e.g. organic molecules down dissolve in water
68
What does the symmetry of a molecule affect?
- affects how the molecules assemble and thus react
- affects how the molecules pack together in a crustal (melting point)
- affects whether the molecule has a dipole moment
- determines whether atoms in a molecules are equivalent
- determines chirality
69
What is the definition of symmetry operation?
A symmetry operation is an action that transforms a molecule so the resultant molecule is unchanged (can be superimposed exactly on the original)
70
What are the symmetry operation's elements and symbols?
rotation = axis of rotation, Cn
reflection = mirror plane, σ
Inversion = centre of inversion, i
rotation-reflection = axis of improper rotation, Sn
Identity - 'whole of space', E
71
What is rotation?
A rotation around an axis is given the symbol Cn for a rotation 1/n of a revolution it is called the n-fold axis,
C2 axis = rotates by 180 degrees(1/2 rotation) e.g. H2O
C3 axis = rotates by 120 degrees (1/3) e.g. CH3Cl
C4 axis = rotates by 90 degrees (1/4) e.g. XeOF4
C6 axis = rotates by 60 degrees (1/6) e.g. benzene
72
What is inversion?
An inversion through a centre of inversion means that if any point of the molecules passes through the centre of inversion it remains the same e.g. ethane.
73
What is reflection?
This is when reflection within a place leaves the molecule the exact same. The reflection an be vertical σv or horizontal σh.
e.g. H20 has two vertical mirror planes
74
What is rotation-reflection?
This consist of rotating the molecules by Cn around an axis and then reflecting it onto a plane perpendicular to the same axis.
75
What is identity?
This consists of doing nothing to change the object at all, or oration guy 360 degrees (C1).
76
What is a point group and how can it be determined?
The point group of a molecule is identified by noting its symmetry elements and comparing these elements that define each point group. It can be determined by flowing the flow chart given.
77
How does point groups classification affect polarity and chirality?-
- a molecules cannot be polar if it belongs to a point group that includes a centre of inversion, any of the D groups and their derivatives, the cubic groups (T,O) and the icosahedral group (I) and their modifications.
- a molecule cannot be chiral if it contains an improper rotations axis (rotation- reflection)
78
Where it the s, p,d,f blocks on the periodic table?
S block - consist of groups 1 and 2
P block - consists of groups 13 -18
D- block - consists of groups 3-12
f -block - the lower periods below the main periodic table
79
Where are the metals, metalloids and non-metals found on the periodic table?
Metals = all of the left side of the periodic table finishing as Al and 'zigzagging' down to Lv.
Metalloids (have semi metal like properties) = starts at B and 'zigzags' down to Te and Po
Non-metals = everything on the right of that line, plus hydrogen.
80
How is the period table based on physical properties/structures?
The periodic table can help predict molecular structures of elements by their position in the periodic table
group = structures
18 = mononuclear gas
17 = dinuclear gas (one single bond)
16 = dinuclear gas (1 double bond)
16 = polynuclear solids (2 single bonds)
15 = dinuclear gas (1 triple bond)
15 = tetranuclear (P4)/ polynuclear solids (3 single bonds)
14 = carbon allotropes (chaoite, graphite and diamond)
13 = metallic systems
2 = metallic systems
1 = metallic systems
81
How can the periodic table be constructed from a theoretical basis?
On a theoretical bases the periodic table can be build using multiples theory's such as the Aufbau principle (lower energy sub-shells fill with electrons first) without violating the rules of the Pauli principle (no same electrons in an atom can have the same set of quantum numbers)
82
Why is the octet rule the driving force for reactions?
The octet rule is that all atoms are driving to achieve their most stable lowest energy state which vernally when the octet rule is achieved and the outer sub shell is fully filled. (8 valence electrons)
83
How do you predict the number of required bonds and structures for an element?
the number of required bond can be predicted using
- the location of the element in the periodic table
- determining bond order from the drawing of the lewis structure
- using quantitative calculations in a thermodynamic cycles to determines enthalpy change and thus the most stable structure and thermodynamically stable. e.g sulphur is most stable at S8 8 single bonds is lower energy whereas O2 has double bonds as that's its most stable formation.
- consider the effectiveness of pi bonding
84
How do you evaluate orbital overlap for possible catenation of multiple bonding?
For a multiple bond to form and be stable there need to be sufficient overlap meaning shorter interatomic distances between the P-orbtials to great a strong bond. Oxygen double bonding involves the overlap of two 2p orbitals which have short interatomic distances greater a strong overlap and stronge bond, whereas if a double bond where to take place in sulphur it would be between two 3p ortbial that have a larger interatomic distance and thus poor overlap and a weak bond so sulphur forms single bonds rather than oxygen doubles as it is more stable.
85
What is meant by allotropy?
It is the property of elements being able to exist in two or more different forms, within the same state.
e.g. carbon as diamond or graphite or Ozone (O3)
86
How do you rationalise the structure/property relationship of allotropes?
Allotrope form due to several reasons
- Different arrangement of particles
- different amount of energy/ temperature present will determine how the particles arrange themselves
- different methods of formation
87
How is ozone formed and why?
Ozones absorbs UV radiation and decompose to form O2, providing a shield protecting the earths surfaces from damaging radiation and temperatures.
O2 --> 2O
O2 + O --> O3
It reacts with greenhouse gases such as nitrogen oxides and halogens from CFCs
NO2 + O3 --> NO3 + O2
NO3 --> NO. + O2
NO. + O3 --> NO2 +O2
Net : 2O3 --> 3O2
Cl. + O3 --> ClO. + O2
ClO. + O --> Cl. + O2
Net : O3 + O --> 202
88
What are the carbon allotropes?
- Graphite (trigonal planar renounce system giving perfect hexagonal structured sheets, which weak intermolecular forces between)
- diamond (4 covalent bond per carbon making a 3D structure, good heat conductor)
- C60 'Buckyball' (truncated iscohedron, a football made up of hexagons and pentagons)
89
What is an oxidation state free energy diagram?
It present a picture representation of the relative stabilities of different oxidation states of an element in aqueous solutions. Its also know as the Frost diagram.
90
What is a redox couple?
It is the relative stability of two oxidation states and is given by free energy changes (ΔG) of the half reactions connecting them.
e.g.
V3+ + e- --> V2+ (reduction)
- if the reduction is negative then V2+ is more stable
- if the reduction is positive then V(III) is more stable
91
What is the Nernst equations?
The ΔG of reduction of a half reaction is related to the corresponding reduction potential by the Nernst equation
ΔG (reduction) = -nFE (reduction)
n = number of electrons involved in the half reactions
F = Faraday constant (96,487 c/Mol)
92
How do you use the Nernst equation?
The Nernst equation can be used to construct Frost diagrams, by using it to find the relative free energy changes
ΔG (reduction) / F = -nE (reduction)
ΔG (reduction) = - ΔG (oxidation)
so
ΔG (reduction) / F = -ΔG (oxidation)/F
93
How do you determine the most stable oxidation state of a redox couple from standard reduction potential data?
If E (reduction) is negative:
Zn2+ + 2e- --> Zn E = -0.76
ΔG (reduction) / F = -nE (reduction)
ΔG (reduction) = 1.52 v so Zn2+ is more stable.
If E (reduction) is positive
Cu2+ + 2e- --> Cu E = 0.34v
ΔG (reduction) = -0.68
Cu is more stable
94
How do you explain the theoretical basis for constructing an oxidation state free energy diagram(frost diagram?
To plot a frost diagram you plot oxidation state vs nE/V
- Free energy for a zero oxidation stare is defined as 0.0 v
- plotted in increasing oxidation states (need to work out ΔG (oxidation)) (ΔG(ox)/F = nE/V)
95
How do you interpret a Latimer diagram?
A Latimer diagram compresses all the standard potentials for redox diagrams of an element. An unstable species will have a lower standard potential to the left hand side than the right
96
How do you construct a Frost diagram from reduction potential data?
If you are given E (reduction)/ v to plot a frost digram you need nE?v so just times the reduction data by number of electron involved the reduction. Then work out each nE/v from the starting 0.00 point
x (0) = 0.00
x (I) = x+ + e- --> x (A)
x(II) = X2+ + 2e- --> X+ (B) x(II) = B
x(III) = X3+ + e- --> x2+ (C) X(III) = B + C
97
How do you interpret a Frost diagram to predict stabilities and feasibility of redox reactions?
- Change down a slope is favourable as it involves negative free energy
- The lower the free energy of each oxidation state is the more stable it is
- spontaneous change will not occur (a corresponding oxidation/reduction must accompany the change)
- oxidising and reducing power increases with the gradient not with the difference in oxidation states (the steeper the line the stronger the oxidising agent)
- The reducing power of an oxidation state is dependent on the gradient of the line joining it to the lower oxidation state. (steeper the decline the stronger the reducing agent)
- a bowl like concave points more thermodynamically stable with respect to disproportionation, and hill like concave point is stable with respect to disproportionation to neighbouring oxidation state
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What is the general trend for OSFE diagrams for transition metal triads?
- higher oxidation state is more accessible for heavier transition metals (Fe, Ru and Os)
- oxidising power of higher oxidation states reduced down trad
- occurrences of lower oxidation sates reduces down triad
Group 8 triad = Fe, Ru, Os
Group 6 triad = Cr, Mo, W
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What is the general trend for OSFE diagrams for main group elements?
- higher oxidation states are accessible for lighter elements
- oxidising power of the higher oxidation state increases down the group
- reducing power of (-II) increase down the group
- S (III) is unstable with respect to disproportionation
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What is meant by disproportionation and comproportionation reactions?
- a disproportionation reaction is during a reaction when the same element is both oxidised and reduced
- a comproportionation reaction is a reaction where the two reactant are the same element with different oxidation states that react to form a compound with an intermediate oxidation number (opposite of disproportionation)
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What is the physical significance of wave function?
A wave function (Ψ) is a mathematical function that relates the location of an electron at a given point in space (identified by x, y, and z coordinates) to the amplitude of its wave, which corresponds to its energy. Thus each wave function is associated with a particular energy E.
For example hydrogen like atoms the energy of their orbitals is determined solely by the principle quantum number n.
For multi electron atoms the inter electrons interaction results in a loss of degeneracy of orbital and dependence of the subsidiary quantum number l.
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What are the order of energies within and orbital (with high electron counts) and why?
1s,2s,2p, 3s,3p,4s,3d,4p
This is because of atomic numbers, electronic shields and atomic radii.
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What is the flame test observations and how does it relate to atomic structure?
If and atom or ion is excited an electron within can be promoted to a higher energy level, as they fall ack down to a lower energy, the energy is released as light each corresponding to a different wavelength and thus a different colour. The colour and wavelength is depended on which subshell the valence electron is moving so atomic structure (valence electron is important)
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What is the concept of electron shielding?
Electron shielding refers to the blocking of valence shell electron attraction by the nucleus, due to the presence of inner-shell electrons.
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What is Slaters rule/equation?
Slaters rule gives an estimation of the effective nuclear charge (Jeff) taking into consideration electron shielding:
Zeff = Z - S
Z = atomic number
S = sheildign coefficent
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how are orbital divided in slaters rule?
(1s) (2s2p) (3s3p) (3d) (4s 4p) (4d) (4f) (5s 5p) (5d) (5f)
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How do you calculate the shielding coefficient (S) for slaters rule?
1. There is no contribution from electrons in the group to the right of the one being considered
2. 0.35 added for each electron in the same group as the one being considered (except 1s when it is 0.30)
3. If the electron is in the ns or np orbital then all electron in the shell below it (n-1) contribute 0.85 and shells n-2 and lower contribute 1.00
4. If the electron is in the nd or nf orbital then all electron below it contribute 1.00
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What are the key trends in atomic and ionic radii?
- atomic radii decreases along a period
- atomic radii increases down a group
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What are the reasons for the key trends in atomic and ionic radii?
- atomic radius increases down a group as although nuclear charge is increasing, the atomic radius is increasing and thus have a higher shielding effect meaning Zeff is decreasing and attraction to outer electron is decreasing. Differences in shielding ability of different subshells causes some discrepancies in this trend (first e- in d block has a contraction effect, same for f block).
- Atomic radii decreases along a period as nuclear charge is increasing, shielding is remaining fairly constant so the Zeff is increasing and electron cloud pulled closer to the nuclear lowering atomic radii.
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What is the d-block and f-block contraction effect?
- The Lanthanide Contraction is the result of a poor shielding effect of the 4f electrons increases attraction to outer orbitals.
- The d block contraction, in the transition metals with d electrons as we move from left to right across the periodic table, the element’s atomic radius only decreases slightly. This is because they have the same amount of s electrons, but are only differing in d electrons and another shell is not added. The d electrons are not good at shielding the nuclear charge, so the atomic radius does not change much as electrons are added.
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What are the main types of chemical bond and characteristics?
The main types of chemical bonding are
covalent = the electrostatic attraction between atomic radius and a shred pair of electrons
ionic = the electrostatic attraction between a negative and positive ion
metallic = the electrostatic attraction between a positive metal cation and a sea of delocalised electrons.
Most real bonds are intermediate meaning most ionic bonds have some covalent character and most covalent bonds have some ionic character.
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How can you use the periodic table to predict bonding?
Across the periodic table the general trend starts with metallic lattices, then giant covalent structure then simple molecules each with different characteristics.
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How do you construct a born-harber cycle for ionic compounds?
A born-harber cycle allows use to understand and determine lattice energies of ionic solids.
1. Element changed into gaseous Staes (atomisation)
2. Ionisation to form the positive cation
3. electron affinity to gain the negative anion
4. lattice entry to form the ionic solid
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How do you use born barber cycles to calculate unknown values and predict reaction feasibility?
Each value can be combined into a single equation equaling the overall enthalpy change allowing us to rearrange and calculate unknowns.
For the reaction to be feasible the electron affinity need to result in an ion having a lower energy than the anion.
X (g) + e- --> X- need to be -Ea
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What is the definition of ionisation energy?
This is the energy change associated with the removal of one electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.
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What is the definition of electron affinity?
Electron affinity is the energy change associated when an electron is added to a neutral atom to create a 1- ion.
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What is the definition of lattice energy?
Lattice enthalpy is the measure of the strength of the forces between the ions in an ionic solid. The greater the lattice enthalpy the stronger the forces (the forces are only broken when ions are gaseous).
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What are the main trends in ionisation energy and affect of electron configuration?
- Ionisation energy decreases down a group as the atomic radii is increasing, shielding is increasing thus the nuclear attraction is reducing and removing an electron is easier, exception occur when there is d-bloc contraction and lanthanoid contraction
- across a group ionisation energies increase as the atomic radii is getting smaller, shielding is constant so atomic attraction is increasing making it harder to remove the outer electron. (exception are moving from S to P subshell and the start of electron pairing)
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What are the main trends in electron affinity and affect of electron configuration?
- Down a group electron affinity is decreasing as the atomic radii is increasing, more shielding and thus less nuclear attraction making it harder to attract an electron. (exceptions are F as the top of the group/most electronegative so pulls electron density making a very small atom causing e- e- repulsion making harder to gain an electron making it less exothermic than expected)
- across a group electron affinity increases as atomic radii is getting smaller so high nuclear attraction making it easier to gain an electron (exception are at the start of electron pairing increasing electron repulsion).
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What is Kapustinkiis equation and how do you use it?
Kapustinkiis equation allows you to calculate lattice enthalpy (UL) for any compound if you know the radii of the constituent ions.
UL (KJ/mol) = (121400 x (z+)(z-)v / (r+ + r-)) x (1- 34.5 / (r+ + r-)
Z+ = cation charge
Z- = anion charge
v = number of ions per formula unit
r+ = radius of cation (pm)
r - = radius of anion (pm)
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How do you qualitatively predict the relative sizes of the lattice energies?
Ionisation energy increases when:
- increases ion charge (Z+ and Z-)
- increase in number of ions (v)
- decrease in ionic radii (r+ and r-)
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What is the typical close packing structure for an ionic system?
The term closest packed structure refers to the most tightly packed composition of crystal structures (lattices). Unfilled spaces with always exist in the packing of spheres, the aim is to minimises the volume of unfilled space to achieve the 'closest packed structure'.
the electronic configuration is [N.G.C]ns1
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How is the structure of an ionic lattice determined?
Ions are charged spheres, there aim is too try and surround themselves with as many oppositely charged ions as possible as closely as possible (increases lattice enthalpy). Structure takes into accountability the radius ration and the cation to anion ratio when determining structure.
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How does the radius ratio influence the type of solid state structure of different compounds?
The smaller the radius of the ions means that they can be packed closer together and thus increasing the size of attraction. Smaller cations occupy tetrahedral holes, larger ones in octahedral and If the cation is too large the anions will adopt a more open structure e.g. a cubic array.
e.g.
NaCl has 6:6 coordination number
ZnS 4:4
SiO2 4:2
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What are tetrahedral and octahedral sites?
In ionic structures the larger anion are arranged in a closed packed array, with the smaller cations occupying the empty 'holes' between. There are two types tetrahedral or octahedral holes.
Tetrahedral > smaller hole that is found between 3 anion in one plane and 1 adjacent anion, arranged like the corners of a tetrahedral shape
Octahedral > the larger hole is found at the centre of 6 anions (3 in one layer and 3 in an adjacent layer) located at the corners of a octahedral shape. e.g. NaCl
A single ionic lattice can have both type of hole.
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What are the different ions lattice shapes?
Hexagonal close packed (HCP) the third layer has the same arrangement of spheres as the first layer covering all tetrahedral holes a-b-a-b-a-b) and cubic close packed third layer stacked into the depressions of the second layer covering all the octahedral holes (a-b-c-a-b-c)
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What is meant by oxide and what is its typically reactions?
An oxide is a chemical compound with one or more oxygen atoms combined with another element.
Group 1 metals rapidly react with oxygen to produce several ionic oxides (M2O).
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What is meant by superoxide and what is its typically reactions?
A superoxide is when potassium, rubidium and caesium react with excess oxygen to produce MO2, (ox number of oxygen is -1/2) they are used to generate oxygen in self-contained breathing apparatus.
4KO2 + 2CO2 --> 2K2CO3 + 3O2
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What is meant by peroxide and what is its typically reactions?
Peroxides form when lithium and sodium react with excess oxygen to form M2O2 with the oxidation number of oxygen equal to -1.
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What are the typical organometallic complexes of group 1 elements?
An organometallic is a compound that contains at least one metal to carbon bond. Due to their relatively low charge densities group 1 form very few organometallic complexes, lithium forms the most as it is the smallest cation thus the highest charge density.
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Why are crown ether, cryptands and host- guest complexes are useful?
The most stable complexes are formed by multi dentate ligands such as crowns and croplands. They are useful as they have the ability to form complies with low charge density group 1 metals, with the size of the cavity in the middle able to be tailored to fit each cation. They are useful as they permit the rapid and easy separation of the alkali metals.
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How are transitions from covalent to ionic systems rationalised and how does it effect reactivity?
Ionic and covalent bonds are the two extremes of bonding, polar covalent is the intermediate between the two. Polarity (measure of separation if charge in a compound) is a spectrum where each compound lies the closer to each extreme the more reactive the molecules.
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What is the concept of charge density and what is the diagonal relationship of the chemistry of LI and Mg?
Mg2+ and Li+ have similar charge densities as the higher charge of Mg2+ is partially offset by its larger size. They thus react in a similar manner
- direct formation of nitrides (Li3N, Mg3N2)
- direct formation of oxides (Li2O, MgO) by combustion in dioxide
- there oxysalts (Li2CO3, MgCO3) readily decompose into oxides
- formation of covalent organometallics (MeLi, Me2Mg)
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How do you predict the occurrence of covalence due to polarisation effects?
The polarising 'power' and tendency of an anion to become polarised by the cation enhances the formation of covalent bonds
- Small cations have high polarising power due to greater concentration of positive charge in a smaller areas (elements at the top of the periodic table)
- large anion have high polarisability due to larger size making outer electrons more easily distorted by the cation
- large charges as electrostatic attraction of to the cation to electrons of anion increases.
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What is meant by the term Bond Dissociation enthalpy?
Enthalpy required dissociate a specific bond in a gaseous molecules (in its ground state) to form two gaseous fragments (in their ground states).
X2 (g) --> 2X (g)
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Why must you take care when using BDE data and why the values in related example system?
As values will differ considerably depending on the compound involved. With bond enthalpies varying due to the surrounding atoms.
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How do you construct Born-Haber cycles for reactions of covalent compounds?
Construct a enthalpy change of formation diagram with the elements in their natural state at the bottom and arrow from both product and reactant pointing downwards. Calculate each bond enthalpy, then equate the two routes and solve the equation.
ΔHr = enthalpy change for bonds broken + enthalpy change for bonds formed.
Another way is to draw out the full structural formula then figure out the net change.
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What are the key trends in bond strength and bond length?
- down a group the BDE reduces as covalent bond strength reduces (atomic radii increasing so valence electrons further away and reduction in electrostatic interaction and overlap becomes poor)
- exceptions are fluorine as the the small size of the atom causes inter electronic repulsion destabilising the bond.
- heteronuclear bonds are vernal;y stronger than homonuclear bonds as charge separations leads to electrostatic interactions additional to the covalent bond interaction.
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What is meant by the term electronegativity?
Electronegativity is the ability of an electron to attract electron density towards itself in a molecule.
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What is Pauling's equation concerning electronegativity?
Bond dissociation energy (A-B) - 0.5 X bond dissociation energy (A-A) + bond dissociation energy (B-B) = 96.49 (xA - xB)^2
It gives the difference in electronegativities, symbol is Xp
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How do you interpret the Mulliken and Allred- Rochow electronegativity?
Mulliken (Xm) and Allred-Rochow gives alternative electronegativity scales
Xm ∝ 0.5(first ionisation energy - first electron affinity)
X = 0.744 +35.9(Jeff/r^2)
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What is the Schomaker-Stevenson relationship relating bond lengths to electronegativity?
The ionicity (charge separation) associated with heteronuclear bonds results in a shorter bond length calculated by the Schomaker-Stevenson relationship
interatomic distance d(A-B) pm = ra (covalent radii of A) + rb (covalent radii of B) - 9(Xa - Xb)
X = electronegativities
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Calculation of electronegativites full calculations method
1. Calculate bond enthalpy using a Born-Haber cycle of N-H
ΔrH = (number of bonds x bond dissociation energy of reactant 1) + (number of bonds x bond dissociation energy of reactant 2) - (number of bonds x bond dissociation energy of products)
2. repeat but with N-N bond enthalpy
3. calculate the electronegativy of N (XN)
96.49(ΔXN) = (BDE of N-H) - 0.5(BDE N-N) + (BDE H-H)
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