FIC May 23 exam Flashcards
(131 cards)
1
Q
Identify how responses to electric fields differ between alpha, beta and Gamma radiation ?
A
Alpha - Helium Nucleus - 2+ charge - attracted to -ve plate - small deflection
Beta - High speed electrons - light - attracted to +ve plate - large deflection
Gamma radiation - high-energy radiation similar to X-rays - no deflection
2
Q
Which is the incorrect answer ?
(A) An alpha particle shows a small deflection towards the negative plate of an electric field
(B) The charged particles detected in cathode rays by JJ Thomson were electrons
(C) Milliken determined the charge on the electron by observing the rate of fall of charged oil droplets
(D) A gamma ray splits and shows an equal deflection toward the negative and positive plate of an electric field.
A
(D) A gamma ray is not charged
3
Q
What is the order of the electromagnetic spectrum in terms of decreasing wavelength and increasing frequency ?
A
Gamma rays - X-rays - UV - Visible - IR - Microwaves - Radio
Raging Martians invaded Venus using X-ray guns
4
Q
Calculate the energy (in J) of light with a frequency of 9 x 10*15 s-1
A
5.96 x 10*18
E = hv
5
Q
Calculate the energy of a wave of wavelength 520 nm
A
3.82 x 10*-19J
E = hc / wavelength
6
Q
Describe the visible changes when T increases on black body objects
A
As T increases glows red then Blue / White
7
Q
How is wavelength affected by changes in T in black body radiation ?
A
Wavelength (red / lower T) > Wavelength (blue / higher T)
As T increases the peak in intensity of radiation will shift to a shorter wavelength
8
Q
Outline the term ‘Ultraviolet catastrophe’
A
Classical physics suggests a black body should emit an infinite amount of energy as the frequency of the radiation approaches infinity. However this would breakdown the laws of physics.
Disproved because experiments showed energy did not increase indefinitely, reached maximum and decreased at higher frequencies.
8
Q
Outline the term ‘Ultraviolet catastrophe’
A
Classical physics suggests a black body should emit an infinite amount of energy as the frequency of the radiation approaches infinity. However this would breakdown the laws of physics.
Disproved because experiments showed energy did not increase indefinitely, reached maximum and decreased at higher frequencies.
9
Q
What is the relationship between energy and wavelength
A
Energy is inversely proportional to wavelength
10
Q
Suppose that yellow visible light can be used to eject electrons from a certain metal surface. What would happen if ultraviolet light was used instead and why ?
A
UV Light has a shorter wavelength, higher frequency and therefor higher energy - so the electron would be ejected with a greater kinetic energy than those ejected by the yellow light.
11
Q
Explain the difference between absorption and emission
spectra
A
Émission spectra
- electron drops from higher energy level to lower
- photon of energy is emitted
- detect wavelength of this line
Absorption spectra
- photon of energy is absorbed.
- Electron is excited from lower energy level to higher.
- Detect wavelength of this line.
12
Q
Explain how Bohr explained line spectra
A
Electron in an atom can only occupy certain circular orbits (corresponding to certain energies and specific radius)
13
Q
How do you calculate energy using plank’s constant and frequency
A
E = hv
14
Q
Identify the key features of the Bohr model which hold true.
A
Electrons exist only in discrete energy levels - described by quantum numbers
Energy is involved in moving an electron from one state to another.
15
Q
As the electron in a hydrogen atom jumps from n = 3 orbit to the n = 7 orbit, does it absorb or emit energy?
A
It absorbs energy.
The electron moves from a lower-energy to a higher-energy state (absorption).
16
Q
How do you calculate wavelength using planck constant, mass and frequency ?
A
Wavelength = Planck constant / mass * frequency
17
Q
Does a moving tennis ball generate matter waves?
If so, can we observe them?
A
Yes!! The mass of the electron is 9.109 x 10-31 kg is much smaller than the mass of a tennis ball 0.05kg. Therefore, the wavelength associated with a tennis ball is too small to observe.
18
Q
Calculate the uncertainty in the position of an electron moving with an uncertainty in the velocity of 5 ms-1 ?
Electron with a mass of me = 9.109 x 10*-31
A
1.16 x 10*-5
Delta x * Delta P = h / 4pi
Delta p = Mass * uncertainty in the velocity
Delta p = (9.109 x 10*-31) x 5
Delta x = (h / 4pi) / Delta P
19
Q
What does wave-function show ?
A
Wave-function gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time.
20
Q
What is and what are the allowed values of the principle quantum number (n) ?
A
1, 2, 3, 4… (integers)
Shows the spatial orientation
21
Q
What type of orbital does the azimuthal quantum number (l) correlate too ?
A
Value of l (type of orbital)
0(s) 1(p) 2(d) 3(f)
22
Q
What is and what are the allowed values of the Azimuthal quantum number (l) ?
A
Defines the shape and energy of the orbital.
Allowed values range from 0 to n-1.
23
Q
What is and what are the allowed values of the Magnetic quantum number (ml) ?
A
Described the three-dimensional orientation of the orbital.
Values are integers ranging from -l to l
24
What are the allowed values of the spin quantum number (ms) ?
+ or = 1/2.
25
Sketch a Radial wavefunction graph for a 1s orbital
Check answer
26
Sketch a radial wavefunction graph for a 2s orbital
Check answer
27
Sketch the radial wavefunction graphs for 3s, 3p and 3d orbitals
Check answer
28
What is the equation for No. of radial nodes
n - l - 1
29
How does the n value correlate to the distance between energy levels ?
Energy levels become closer together as n value increases
30
How does number of electrons affect repulsion ?
As the number of e-s increases, so does the repulsion between them.
31
In quantum mechanics, what does it mean for subshells to be 'degenerate' ?
An energy level is degenerate if it corresponds to two or more different measurable states of a quantum system.
31
In quantum mechanics, what does it mean for subshells to be 'degenerate' ?
An energy level is degenerate if it corresponds to two or more different measurable states of a quantum system.
32
Outline Pauli's exclusion principle
No two fermions (particles with half integer spin) can occupy the same quantum state simultaneously. Which is why the spin quantum numbers of a quantum state must be +1/2 and -1/2
33
Outline the Aufbau Principle
Electrons in atoms build up from the lowest energy levels.
Order of filling up orbitals - 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p
Exceptions to the rule due to extra stability associated with half-filled and filled sets of d-orbitals which can occur since the 4s and 3d orbitals are very close in energy.
34
Outline Hunds rule
For degenerate orbitals (ie p or d/0/ the lowest energy is attained when the number of electrons with the same spin is maximised.
35
Briefly describe the the components of an electrochemical cell
Comprised of two half cells, a electrode in contact with a redox solution.
36
Define electrode potential
Energy when the redox states within the redox couple at equilibrium.
37
What are voltaic cells ?
Two connected half cells also known as a battery. E released in redox reaction may be used to perform electrical work.
38
Define and outline the equation for electromotive force (emf)
Joining two electrodes to form a cell creates a potential difference between them called the electromotive force. At equilibrium the emf is the difference between the two electrode potentials.
emf = E(+) - E(-)
emf sometimes written as Delta E
39
Calculate emf for this daniel cell
Zn I Zn2+ (aq) I I Cu2+(aq) I Cu
-0.76v 0.34v
0.34v - (-0.76 volts) = 1.10 volts
40
Define Gibbs function (delta G) ?
Measure of useful work done in a chemical reaction occurring in the cell which may be equated with the electrical output of the cell.
41
What is the equation for Gibbs free energy ?
Delta G = Number of electrons in redox equation x Faraday constant (charge of 1 mole of electrons) x electromotive force
42
Define radial distribution function
RDF = Probability that electron lies in a spherical cell of radius, r, thickness dr - useful to find most probable distance of finding an electron from the nucleus.
43
What types of bonds are contained within single and multiple bonds ?
In single bonds the bonds are always sigma bonds.
In multiple bonds (double, triple etc) one bond is sigma the rest are pi.
44
How are sigma bonds characterised ?
Head to head overlap
Shared electron density with internuclear axis.
45
How are Pi bonds characterised ?
Side to side overlap
Electron density above and below the internuclear axis.
46
What is the difference between in phase vs out of phase ?
In phase refers to constructive overlap of atomic orbitals. Results in molecular orbital with lower energy and greater stability than the original atomic orbitals.
Out of phase refers to the destructive overlap of atomic orbitals. Results in molecular orbital with higher energy and lower stability than the original atomic orbitals.
47
What happens as overlap of orbitals increases ?
As overlap of orbitals increases
- bonding MO e lowered (stabilised)
- Anti-bonding MO e raised (destabilised) more compared to the atomic orbitals from which they are formed.
48
What characteristics do atomic orbitals need to have for significant overlap to occur ?
- Similar size
- Similar energies
- Same symmetry
49
What is the calculation for bond order ?
Bond order = 1/2 ((no. of bonding e-)-(no. of anti-bonding e-))
50
What does 'end on' overlap of p orbitals form ?
End on overlap forms sigma and sigma* Molecular orbitals.
51
What are the two types of interactions between atoms with p orbitals ?
The Pz orbitals can overlap head-on or end-on to form a sigma molecular orbitals.
The other two sets of p orbitals overlap side-on to form a pi molecular orbitals.
52
What are the two types of orbital overlap ?
Head on overlap
Present in sigma bonds - if rotate internuclear axis there is no phase change
Side on overlap
Present in Pi bonds - if rotate internuclear axis there is a phase change
53
What signs indicate symmetry in orbitals ?
Is the phase the samw @ x,y,z as @ -x,-y,-z
yes - gerade (g)
no - ungerade (u)
54
How do you identify something as paramagnetic or diamagnetic ?
Paramagnetic materials have unpaired electrons, while diamagnetic materials have all their electrons paired. This is because unpaired electrons generate magnetic fields themselves, which contribute to the overall magnetisation of the material.
55
Describe how bonding and antibonding MOs are formed between 1s and 2p orbitals.
1s - A bonding molecular orbital is formed when the two 1s orbitals combine in phase, antibonding formed when the two combine out of phase
2p - Bonding - When 2p orbitals combine in phase, a bonding molecular orbital is formed. this has a node at the midpoint between the two nuclei. Which reflects the cancelation of the two waves. Antibonding - when the two p orbitals combine out of phase an antibonding mo is formed. This has a node at the midpoint where the waves add up to zero.
56
Draw the MO diagram for O2, O2+ and O2-
Check.
57
Which is more electro-ve, a bonding MO or an Anti-bonding MO ?
Bonding MO - Bigger contribution - more electro-ve atom.
Antibonding MO - Bigger contribution - less electro-ve atom
58
Define metallic character
Preference to lose electrons metals tend to easily oxidise and form + ions.
59
Define ionisation energy
Ionisation energy is defined as the energy change when an electron is removed from an atom of the element in the gas phase.
60
Define Ionic radius
Internuclear distance in ionic solids.
61
Define electron gain energy
Energy change that occurs when an electron is attached to an atom in the gas phase.
62
Describe the trend in ionisation energy across the periodic table
Left to right - increase
Top to bottom - decrease
63
What are the two anomalies in the trends in ionisation energy in the periodic table
Be - B = ionisation of B easier due to favourable 2s2 configuration
N - O = stability due to electron spin of 4th electron.
64
Why is ionisation energy of sodium smaller than that of magnesium ?
Magnesium has a smaller radius and higher nuclear charge than sodium. Meaning more energy will be required to remove the electron from the same orbital (3s)
65
Which has a higher first ionisation energy, Be or B ?
Be has higher IE removing e- that is further away plus stability from full 2s.
66
How are hydroxides formed ?
Hydroxides are formed on the reaction of group 1 metals (or their oxides) with water.
67
The reaction of group 1 metals with water creates vigorous reactions. explain why rubidium and caesium reactions are dangerous
Because reaction is exothermic which ignites H2 -> risk of explosion
68
Briefly compare molecular orbital theory and valence bond theory
Molecular orbital theory
Ground and excited state.
Quantitative picture of bonding.
Delocalised MOs for electrons (can be spread over whole molecule)
Valence bond theory
Ground state
Qualitative picture of bonding
Localised bond + e- pairs.
69
What is the difference between ionic and covalent bonding ?
Covalent = shared pair of electrons (eg H-H)
Ionic = Handed over electrons (eg Na+Cl-)
70
What is the difference between a regular covalent bond and a polar covalent bond ?
In covalent bonding electron pairs are shared equally between atoms, in polar covalent electron pairs are shared unequally giving a polar covalent bond.
71
Draw the lewis structure for GeCl3-. Hint the -ve charge means there is an extra electron associated with one of the chlorides to form a dative covalent bond.
Check.
72
What does a higher value of electronegativity mean ?
Higher electronegativity = more likely to attract electrons towards itself.
73
Define the Octet rule
An octet of electrons consists of full s- and p- sub-shells on an atom, and gives rise to a stable, inert electronic configuration.
If there are no vacant orbitals, there is no reactivity.
74
Briefly outline the exceptions to the OCTET rule
Molecules with odd numbers of electrons. Reduced octets (ie less than 8 electrons) and expanded octets.
75
Explain why group 2 elements are strong reducing agents
G2 elements are strong reducing agents because they gave a relatively low ionisation energy and high electropositivity which means they have a strong tendency to lose electrons and form cations in reactions.
76
What is a reducing agent ?
In a redox reaction a reducing agent donates electrons to another species, causing that species to be reduced while the species itself is oxidised.
77
Why are group 2 peroxides generally more stable down the group ?
Because of the decrease in ionisation energy and increase in size of the atoms as you move down the group.
78
Briefly outline what valence bond theory is
Valence bond theory is the idea that valence atomic orbitals combine or interact to form chemical bonds.
79
Write the resonance structures for HF & ClI
H-F => H+F- => H-F+
Cl - I => Cl+I- => Cl-I+
80
What are the characteristics of sp hybrids ?
One s, one p orbital mixed
Linear shape (180 degrees)
Form triple bonds normally
One sigma two pi bonds.
81
What are the characteristics of sp2 hybrids ?
one s, two p orbitals mixed
Trigonal shape (120 degrees)
Form double bonds
One sigma one pi bond
82
What are the characteristics of sp3 hybrids ?
One s, three p orbitals mixed
Tetrahedral shape (109 degrees)
Form single bonds
Sigma only.
83
What are the characteristics of sp3 hybrids ?
One s, three p orbitals mixed
Tetrahedral shape (109 degrees)
Form single bonds
Sigma only.
84
Using a sketch of the Lewis structure of chlorine, explain the Lewis treatment of electrons in bonding.
Check correctly drawn lewis structure
The lewis theory suggests that atoms in a molecule share electrons to achieve a more stable electron configuration.
85
Briefly outline the assumptions of the Valence Shell Electron Repulsion Theory
1. Atoms in a molecule held together by pairs of electrons known as bonding pairs (Lewis model)
2. Some atoms have pairs of electrons that are not involved in bonding (lone pairs)
3. As electron pairs are negatively charged, they repel one another such that all electron pairs are as equally spread across the molecule as possible.
86
Outline the approach for determining a molecules domain and structure
1. Count the number of bonding environments (bonds)
2. Count the number of non-bonding environments (lone pairs)
3. Add numbers together to determine electron domain shape.
4. Select molecular geometry from possibilities within of electron domain, based on bonds and lone pairs.
87
Identify the electron domain of methane (CH4)
4 Bonding environments
0 Non-bonding environments
4 + 0 = 4
4 total environments = Tetrahedral domain
88
Identify the electron domain and molecular geometry of Ammonia (NH3)
3 Bonding environments
1 Non-bonding environments
3 + 1 = 4 total environments = Tetrahedral domain
3 bonding environments + 1 non-bonding = Trigonal Pyramidal molecular geometry
89
What is the electron pair arrangement and approximate bond angle of a molecule with 2 electron groups ?
0LP 2BP = Linear (180 degrees)
90
What are the possible electron pair arrangements and approximate bond angles of a molecule with 3 electron groups ?
0LP 3 BP = Trigonal Planar (120 degrees)
1LP 2BP = Bent (<120 degrees)
91
What are the possible electron pair arrangements and approximate bond angles of a molecule with 4 electron groups ?
0LP 4BP = Tetrahedral (109.5 degrees)
1LP 3BP = Trigonal Pyramid (<109.5 degrees (~107 degrees)
2LP 2BP = Bent (<109.5 degrees (~105 degrees)
92
What are the possible electron pair arrangements and approximate bond angles of a molecule with 5 electron groups ?
0LP 5BP = Trigonal Bipyramidal (90 & 120 degrees)
1LP 4BP = See-saw (<90 & <120 degrees)
2LP 3BP = T-structure (<90 degrees)
3LP 2BP = Linear (180 degrees)
93
What are the possible electron pair arrangements and approximate bond angles of a molecule with 6 electron groups ?
0LP 6BP = Octahedral (90 & 90 degrees)
1LP 5BP =Square Pyramidal (90 & <90 degrees)
2LP 4BP = Square Planar (90 degrees)
94
How do charged species differ when drawing their structures
Negative
Structure has an extra electron so add electron to central atom
Positive
Structure Has one less electron so remove electron from central atom.
95
What are the limitations of the Valence Electron Shell Repulsion Theory ?
Uneven numbered molecules cannot be predicted because VSEPR is based on the Lewis model.
Heavy atoms have their shape incorrectly predicted by VSPER because of the inert-pair effect.
96
Write the half equations for Magnesium in air
Mg (s) + 1/2O2 (g) => MgO (s)
Mg => Mg2+ + 2e-
1.2O2 + 2e- => O2-
97
Write the half equations for
2NaCl(l) => 2Na (s) + Cl2 (g)
2Na+ (l) + 2e- => 2Na (s)
2Cl- (l) => Cl2 (g) + 2e-
98
What are the 4 primary rules when assigning oxidation states ?
1. Elements in pure states always zero
2. Sum of all oxidation states = 0 in uncharged compound
3. Oxidation states in simple ions = charge
4. Oxidation states in compounds with charges = charge
99
What are the oxidation states for Fluorine , Oxygen, Hydrogen, Chlorine and group 1 & 2 metals
Fluorine = -1 apart from in F2
Oxygen = -2 unless fluorinated
Hydrogen = +1 unless in hydrides
Chlorine = -1 except with O or F
Group 1 and 2 metals follow their respective numbers.
100
Identify what has been oxidised and what has been reduced in this reaction
2Na (s) + 2H2O (l) => 2NaOH (aq) + H2 (g)
Na changes from 0 > +1
H changes from +1 > 0
O no change
Therefor sodium has been oxidised and hydrogen is reduced
101
What is the trend in metallic character across the p-block ?
Metallic character decrease along the period resulting in formation of covalent bonds rather than ionic for oxides and halides. This happens because the atom gets bigger so the shells gets bigger meaning the electrons are further away from the nucleus.
102
Define the 'inert-pair effect ?
Inert pair effect = tendency of outershell electrons to remain unshared in compounds of post-transition metals.
103
What happens to bond enthalpy down periods of the p block and why does this happen ? And define bond enthalpy
Bond enthalpy decreases down the p block due to an increase in atomic size. The increase in atomic size leads to an increase in the distance between the nuclei of the bonded atoms, resulting in a decrease in the strength of the bond.
Bond enthalpy = energy required to break one mole of covalent bonds between atoms in a gaseous molecule.
104
Discuss the trend in oxidation state as you proceed down group 13.
All group 3 elements from 0 and +3 oxidation states due to having three valence electrons in its outer shell. As you go down +1 oxidation state becomes increasingly prevalent. +1 oxidation state allowed due to inert pair effect. -s electrons shielded from bonding.
105
Why is BF3 used as an electron acceptor in many synthesis
- Boron's ability to accept electrons makes it a good Lewis acid.
- This property allows BF3 to form coordinate covalent bonds with other molecules by accepting an electron pair. Leading to formation of a stable complex.
- BF3 is often used as a catalyst in organic synthesis reactions due to its Lewis acid behaviour.
106
Explain why Boron halides are excellent Lewis acids and rationalise why Bl3 is a better Lewis acid than BF3.
Boron Halides are excellent Lewis acids because their electronic structure (6 valence e-) allows for empty p-orbitals.
BI3 is a better Lewis acid because : Iodine atoms are larger than fluorine atoms, meaning the electron cloud around the iodine atom in BI3 is more easily distorted. Iodine has more electrons than fluorine which means that the BI3 has a greater electron density. Iodine has a lower electronegativity than fluorine, which means that is is less able to attract electrons towards itself. Meaning iodine atoms more likely to accept electron pairs from other molecules or ions.
107
Define the term allotrope and give three allotropes of carbon
An allotrope is when a chemical element can exist in multiple forms in the same physical state or phase.
Diamond, Graphite and fullerenes.
108
How is an Octet achieved in oxygen ?
Octet is achieved by sharing 2 e-, usually through covalent bonds. Second most electronegative element.
109
What are the two allotropes of oxygen ?
O2 - colourless, odourless gas, blue liquid.
O3 - Pale blue gas
110
How is pure oxygen formed ?
Pure oxygen is formed from the decomposition of chlorate or peroxide.
111
Write the equation for oxygen in the atmosphere (ozone)
3O2 (g) <=> 2O3 (g)
112
What type of reaction is ozone formation ?
An endothermic reaction - preferable in higher temperatures.
During the reaction ozone undergoes photochemical reactions where O-O bonds break.
O3 (g) => O2(g) + O* - oxygen radical formed.
113
How is the oxygen radical formed ?
O3 consists of O=O-O
O-O bonds are weaker than O=O bonds and absorb UV-B and UV-C radiation.
114
What is the reaction between O3 and chlorine radicals and describe what this process means ?
Cl* (g) + O3 (g) => ClO* (g)+ O2 (g)
Chlorinates species in the atmosphere form radicals and attack ozone, converting it to oxygen. This is advantageous in the troposphere, but disadvantageous in the stratosphere.
115
What is formed with the combustion of group 1 metals in oxygen ?
Combustion of group 1 metals in oxygen forms different oxygenated species.
Lithium forms oxides, sodium forms peroxides and potassium and anything heavier forms superoxides.
116
Explain how ozone is formed in the upper stratosphere. Use chemical equations to show how chlorinates species attack the ozone layer.
In upper stratosphere - temperature is high. O3 formation is endothermic so high chance of O3 formation.
O3 - O=O*-O- => O2 + O-
Cl* + O3 => ClO- + O2
ClO* + O- => O2 + Cl*
Auto catalysing so reaction will continue until termination.
Cl* + Cl* => Cl2
117
Why is Sulfur unique in terms of its allotropes ?
Sulfur exhibits one of the widest range of allotropes known because it can catenate (form chains/rings with itself). S2 - S8
118
What are the two oxides that sulfur forms ?
Sulfur Dioxide (SO2) and Sulfur trioxide (SO3).
119
Describe the solubility of sulfur dioxide
SO2 is soluble in water and will form sulfurous acid.
120
What is formed when fluorine reacts with oxygen ?
Oxygen forms toxic gases with fluorine.
2OH - (aq) + 2F2 (g) => OF2 (g) + 2F- (aq) + H2O (l)
This is the only time oxygen forms positive oxidation states.
Oxygen dihalide stability decreases as halide gets bigger - I2O is unknown.
121
Write chemical equations to show the synthesis of sulfuric acid from elemental sulfur, S8. State safety considerations for the reactions and how unsafe conditions are avoided practically.
SO (g) + 8O2 (g) => 8SO2 (g)
SO2 (g) =(V2O3)> SO3 (g)
SO3 (g) + H2O => H2SO4 (g) (corrosive mist)
Usually SO3 pumped into OLEUM, H2S207 to reduce heat given OH. SO3 dissolving in water is dangerously exothermic.
122
What is the standard configuration of group 17 elements ?
Ns3Np5 - They are each one electrons short of a noble-gas configuration and so are often found in the -1 oxidation state.
123
What is the trend in bond dissociation enthalpies in group 17 ? and identify the anomaly.
General trend is that bond dissociation enthalpies of dihalogens decrease down the group because : Larger halogens lead to longer, weaker covalent bonds, Reduced electrostatic interaction between nuclei and valence electrons, easier to break bonds.
Fluorine is anomalously weak.
124
What is formed when chlorine dissolves in water ?
Cl2 (g) + 2H2O (l) => HOCl (aq) + H3O (aq) + Cl- (aq)
Chlorine dissolves in water to form Hypchlorous acid (HOCl) and hypochlorite ion (OCl-) is formed in basic conditions
125
What is a polyiodide and what is the most common ?
Polyiodides are widely known and useful - as the name suggests, they comprise entirely of iodine only.
I3- is the most common.
126
Write balanced equations to predict the products of reactions between Xe and F2 if they are in :
A 1.5 ratio at 5atm
Xe (g) + 2F2 (g) => XeF4 (g)
127
Briefly explain what is meant by the “photoelectric effect” and what key concept about
the nature of light was developed as a result of observing this phenomenon
The photoelectric effect refers to the phenomenon where electrons are emitted from a material when light of a certain frequency or higher is shone on it.
The key concept about the nature of light developed as a result of observing this phenomenon is that light can behave as both a wave and a particle
128
Write the electron configuration of Rb
1s2 2s2 2p6 3s2 3p6 4s2 4p6 5s1
129
In the Bohr model of the hydrogen atom, when the electron is in its ground state, it orbits the nucleus at a specific radius of 0.53 Å.
In the quantum mechanical description of the hydrogen atom, the most probable distance of the electrons from the nucleus is 0.53 Å.
Why are these two different statements?
Bohr model = 100% certainty of finding electron in that area
Quantum mechanical description = Less than 100% certainty.
