Atomic theory 5 Flashcards
1
Q
What was thought about electrons prior to Bohr’s model?
A
Electrons were able to posses any amount of energy, so they could be found anywhere around the nucleus.
2
Q
What did Bohr show?
A
That the amount of energy an electron possessed was a definite quantifiable amount.
3
Q
1st shell n=1
A
2(1)² = 2
4
Q
What did Bohr show that ‘n’ increases?
A
The energy states become progressively closer to one another.
5
Q
4th shell n=4
A
2(4)² = 32
6
Q
2nd shell n=2
A
2(2)² = 8
7
Q
What did Bohr notice?
A
That when white light is passed through a prism the light is split into an array of colours called the spectrum.
8
Q
Give an example of a continuous spectrum
A
A rainbow
9
Q
When light is passed through a prism the light is split into an array of colours called a spectrum. What is this known as?
A
Continuous spectrum
10
Q
Why is the light splitting into an array of colours known as the continuous spectrum?
A
As it consists of a continuous range of wavelengths (colours)
11
Q
Bohr carried out experiments using….
A
Light coming from a hydrogen discharge tube.
12
Q
What happened when he passed the light coming out from the discharge tube through a prism?
A
Only lines of a few wavelengths were present in the resultant spectra.
13
Q
What did Bohr call it when he saw a series of lines?
A
An emission line spectrum.
14
Q
What did the fact that the spectrum consisted of a series of lines indicate?
A
That only certain energy emissions are possible.
15
Q
What did Bohr do after his experiment was finished?
A
He repeated the experiment with other elements.
16
Q
What did Bohr notice after he repeated the experiment with other elements?
A
Noticed that each element had its own characteristic line spectrum.
17
Q
What was concluded because each element had its own characteristic line spectrum?
A
They each had a ‘fingerprint’ = a means of identification
18
Q
What was concluded because each element had its own characteristic line spectrum?
A
They each had a ‘fingerprint’ = a means of identification.
19
Q
In an atom, where do electrons revolve around?
A
Revolve around the nucleus in certain allowed orbits or shells.
20
Q
How much energy does an electron have when in a shell?
A
A definite amount of energy
21
Q
What is an energy level?
A
One of the discrete amounts of energy that an electron has when it is in an atom.
22
Q
What is the energy level dependant on?
A
Dependant on the distance from the nucleus.
23
Q
When close to the nucleus, the electron has …… energy
A
Little
24
Q
When far from the nucleus the electron has ……. energy
A
More
25
What happens to the amount of energy an electron contains when it stays in a particular energy level?
It remains the same
26
What happens when an atom absorbs energy?
The electrons jump from a lower to a higher energy level.
27
What happens when an atom absorbs energy?
The electrons jump from a lower to a higher energy level.
28
What state is an atom in when the electrons jump from a lower to a higher energy level?
The 'excited' state
29
Two points on the 'excited' state.
- Temporary
| - Unstable
30
What happens after the 'excited' state is reached?
The excited electron will fall back to a lower energy level.
31
What happens when the excited electron falls back to a lower energy level?
Energy is given off
32
How much energy is given off as the electron falls back to a lower energy level and why?
As the electron can only fall back to a certain definite energy levels, only fixed amounts (quantifiable amounts) of energy can be given off.
33
What energy level does hydrogen's only electron occupy?
The lowest available energy level - ground state.
34
What happens when energy is given to the electron by heating it by example?
The electron is promoted to a higher energy level.
35
Give two points on the state when energy is given to the electron by heating it for example?
- Unstable
| - Temporary
36
What happens when the electron falls back to a lower level?
The energy difference is emitted as a photon of light. E2 - E1 = E
37
What is each line in the spectrum a result of?
The electron moving from one energy level to a lower one.
38
What does each transition element have?
A definite amount of energy and appears as a line of a particular colour in the line spectrum mission.
39
What does each line in the line emission spectrum have?
A definite frequency
40
What is indicated by the fact that each line has a definite frequency?
Only a limited number of energy changes are possible within the structure of an atom.
41
E =
hf
42
E = hf meaning
```
E = energy
h = Planck's constant
f = frequency
```
43
f =
1/λ
44
f = 1/λ meaning
```
f = frequency
λ = wavelength
```
45
What is the Lyman series?
When electrons fall from a higher energy to the n=1 energy level, we get a set of lines called the Lyman series.
46
What is the Balmer series?
When electrons fall from a higher energy level to the n=2 energy level, we get a set of lines called the Balmer series.
47
What series do we get in the atomic emission spectrum of hydrogen?
Balmer series
48
What is the Paschen series?
When electrons fall from a higher energy level to the n=3 energy level, we get a series of lines called the Paschen series.
49
Why do different elements have unique atomic spectra? (L.C)
Each element has a different arrangement of energy levels and a different electronic configuration, giving rise to different electron transitions (jumps) from higher to lower energy levels.
50
Sodium - colour of flame
Amber/yellow
51
Potassium - colour of flame
Lilac
52
Copper - colour of flame
Blue/green
53
Lithium - colour of flame
Crimson
54
Strontium - colour of flame
Red
55
Barium - colour of flame
Yellow/green
56
***Explain why different metals have different flame colours.
n
57
Can atoms absorb light?
Yes
58
What happens if white light is passed through an element in its *gaseous* form?
The light that comes out has wavelengths missing.
59
If white light is passed through an element in its *gaseous* form, the light that comes from it has missing wavelengths. Ste how they appear + energy amount.
The missing wavelengths appear as dark lines and have the same energy as would appear in the emission spectrum.
60
What is the amount of light absorbed proportional to?
The concentration of the element.
61
Why is AAS used to determine the quantity of 'heavy metals' present in water analysis and the amount of lead in a blood sample?
As the amount of light absorbed is proportional to the concentration of the element.
62
What is used to determine the quantity of 'heavy metals' present in water analysis and the amount of lead in a blood sample?
AAS
63
The amount of light absorbed is proportional to the concentration of the element. What does this tell us?
That atoms in the ground state can absorb the same amount of energy as they would emit in the excited state.
64
What is an absorption spectrum?
A series of dark lines against a coloured background.
65
What is an emission spectrum?
A series of coloured lines against a dark background.
66
AAS
Atomic absorption spectrometry
67
Use of AAS
Used to analyse the amount and type of elements present in a sample (quality control).
68
Street lights
Sodium streetlights give out light with a distinct colour (amber).
69
Fireworks
The elements used in fireworks give out their distinct colours when they are heated.
70
How does an electron move according to Bohr's theory?
An electron moves in a fixed path or orbit around the nucleus.
71
According to Bohr's theory an electron moves in a fixed path orbits around the nucleus. What element does this work with when dealing with and why?
Hydrogen as it only has one electron.
72
State four limitations of Bohr's theory. (L.C)
1. This theory failed when applied to atoms with more than one electron i.e all others!
2. Does not take wave-particle duality into account
3. Does not allow for uncertainty (probability)
4. Does not explain the discovery of sublevels.
73
What did Louis Victor de Broglie propose?
That electrons may behave as waves as well as particles, like light. (proven 4 yrs later)
74
Who proposed that electrons may behave as waves as well as particles, like light, and when?
Louis Victor de Broglie in 1923.
75
Who in 1926 published the first paper on wave mechanics in which the electron in atoms was modelled as a wave rather than as a particle?
Erwin Schrodingers
76
What was the new model that Erwin Schrodingers published in 1926?
The electrons in atoms was modelled as a wave rather than as a particle.
77
Define Heisenberg's uncertainty principle 1927
This states that it is impossible to know both the position and the speed (velocity) of an electron at the same time as electrons move in a wave motion.
78
What must be taken into account because the mass of an electron is so small?
They are disturbed by the measuring techniques.
79
The mass of an electron is disturbed by the measuring techniques. What does this mean?
The more precise one measurement the less precise the other.
80
What are we left to deal with because the mass of an electron is so small that they are disturbed by the measuring techniques?
Left with dealing with the probability of finding an electron at a particular position within an atom.
81
Define atomic orbital
An atomic orbital is a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
82
Define energy level
Is the discrete (exact) amount of energy an electron has when it is in an atom.
83
Define energy sublevel
A group of atomic orbitals within an atom, all of which have the same energy.
84
What can each shell/energy level be sub-divided into?
Sublevels
85
What does an electron have when it is in a particular sublevel?
A particular (quantifiable) amount of energy.
86
1s
s = 2 electrons
87
2s
s = 2 electrons
88
2px2py2pz
p = 6 electrons
89
3s
s = 2 electrons
90
3px3py3pz
p = 6 electrons
91
4s
s = 2 electrons
92
3d(5 orbital)
d = 10 electrons
93
4px4py4pz
p = 6 electrons
94
4d (5 orbitals)
d = 10 electrons
95
4f (7 orbitals)
f = 14 electrons
96
n = 1
1s
97
n = 2
1s, 2s, 2px2py2pz
98
n = 3
1s, 2s, 2px2py2pz, 3s, 3px3py3pz, 4s, 3d
99
n = 4
1s, 2s, 2px2py2pz, 3s, 3px3py3pz, 4s, 3d, 4px4py4pz, 4d, 4f
100
What shape are s orbitals?
Spherical
101
What shape are p orbitals?
Dumbbell
102
Capacity of s
2
103
Capacity of p
6
104
Capacity of d
10
105
Capacity of f
14
106
All orbitals
1s, 2s, 2p(x2py2pz), 3s, 3p(x3py3pz), 4s, 3d, 4p(x4py4pz), 4d, 4f
107
Why do the 4s and 3d sublevels overlap?
Because the 4s sublevel is of lower energy than the 3d and therefore must be filled first.
108
What is very important to remember when drawing s and p orbitals?
Always put in the labels for the axes
109
What are the 3 orbitals that the p sublevel are split into?
px, py, pz
110
Capacity of px, py, px
- Equal energy, all 2 (meaning p = 6)
111
What gets filled first the 3d sublevel or the 4s sublevel?
4s
112
Why does the 4s sublevel get filled first?
As it is of lower energy
113
Define Aufbau principle
Electrons occupy the lowest available energy level.
114
Define Hunds rule of maximum multiplicity
When two or more orbitals of equal energy are available (i.e 2px 2py 2pz), electrons fill them singly before filling them in pairs.
115
Define Pauli exclusion principle
No more than two electrons can occupy an orbital and this they can only do if they have opposite spin.
116
For Hunds rule of maximum multiplicity, what must you have?
Orbitals of equal energy
117
Write the full electronic configuration for Na atom. (first look up the atomic number to see how many electrons the Na atom has)>
11e : 1s₂, 2s₂, 2px₂, 2py₂, 2pz₂, 3s1
118
Write the full electronic configuration for Na ion. (look up atomic number and take 1e away as Na loses electron, brackets must be added!!)
11e - 1e = 10e : [ 1s₂2s₂2p₆]+
119
Is a half filled p-sublevel stable or unstable?
VERY STABLE
120
Is the atomic number always a whole number?
Yes
121
Which number is the atomic number?
The number above the symbol in the periodic table.
122
s-block elements
All the elements in groups I and II
123
p-block elements
All the elements in groups III and O
124
d-block elements
The elements in-between groups II and III
125
Why are all the elements in group I and II s-block elements?
As they have their highest energy electron in an s orbital.
126
Why are all the elements in groups III and O known as p-block elements?
As they have their highest energy electron in a p-orbital.
127
Why are all the elements in-between groups II and III known as d-block elements?
As their highest energy electron ENTER a d orbital.
128
What are the two exceptions to electronic configurations?
Chromium (Cr) and copper (Cu)
129
What is the electronic configuration of Cr - 24e?(exception learn off)
1s₂, 2s₂, 2p₆, 3s₂, 3p₆ 4s1, 3d₅ (instead of 4s₂, 3d₄)
130
What is the electronic configuration of Cu - 29e?(exception learn off)
1s₂, 2s₂, 2p₆, 3s₂, 3p₆ 4s1, 3d₁₀ (instead of 4s₂, 3d9)
131
Why are there two exceptions to electronic configuration?
As half-filled and full sublevels are more stable than partially filled sublevels.
132
What do transition elements have?
Partially filled d-sublevels.
133
Do transition elements form positive or negative ions?
Positive ions
134
Why do transition elements form positive ions?
Because they have only 1 or 2 electrons on their outer shells.
135
What are transition elements?
Metals
136
What two elements are actually d-block elements? (should be transition elements but AREN'T)
Zinc and scandium
137
Define a transition metal
A transition metal is one that forms at least one ion with a partially filled d sublevel.
138
List 3 characteristics of transition elements
- Show variable valency
- Form coloured ions/compounds
- Act as catalysts in many reactions
139
Does scandium act as a catalyst?
No
140
What colour compounds does scandium form?
White compounds
141
What colour compounds does zinc form?
White compounds
142
What ions does scandium form?
+3 ions
143
What ions does zinc form?
+2 ions
144
Does zinc act as a catalyst?
No
145
What theory is behind the flame tests?
Different metals emit different colour flames to the Bunsen burner due to each metal having different electron arrangements.
146
Apparatus for flame tests
Safety glasses, Bunsen burner, nichrome (or platinum) wire.
147
Materials needed for flame tests
Concentrated hydrochloric acid (two different batches), salts of sodium, lithium, copper, potassium, calcium.
148
What is the first step in the flame tests?
The nichrome wire was first cleaned by dipping it in concentrated hydrochloric acid (batch one).
149
What happens after the nichrome wire was first cleaned? (2nd step flame tests)
It was heated to red hot in the Bunsen flame until no flame colour was observed.
150
What happens after the nichrome wire was heated? (3rd step in flame tests)
The wire was then dipped into a clean sample of hydrochloric acid (batch two) and then into the salt to be tested.
151
What happens after the wire was dipped into the salt? (4th step in flame tests)
The wire was then held in the Bunsen flame.
152
What happens after the wire was held in the Bunsen flame? (5th step in flames test)
The colour of the flame was noted and recorded.
153
What happens after the colour of the flame was noted and recorded? (last step in flame test)
This procedure was repeated for the other salts, ensuring that the wire was cleaned thoroughly each time.
154
Write the electron configuration (s. p etc) for an iron atom. (L.C)
1s₂ 2s₂ 2p₆ 3s₂ 3p₆
155
What term is used to refer to the condition of the hydrogen atom when its electron occupies the E1 level? (L.C)
Ground state
156
What term is used for the condition of the hydrogen atom when its electron occupies any of the levels E2, E3 etc? (L.C)
Excited state
157
What causes the electron to leave the E1 level? (L.C)
- It acquires energy
| - It is heated
158
Why does the electron not remain in any of the levels E2, E3 etc? (L.C)
Higher energy states unstable
159
The visible lines in the atomic emission spectrum of a sample of hydrogen are produced when electrons fall to a particular energy level. Identify this energy level. (L.C)
E₂ / n = 2
160
Bohr's theory considered electrons as tiny particles restricted to orbits. How does modern atomic theory describe the behaviour of electrons? (L.C)
Electrons have both wave and particle properties.
161
What are orbitals? (L.C)
An atomic orbital is a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
162
Identify the main energy levels involved in the electron transition that gives rise to the first (red) line of the Balmer series in the emission spectrum of the hydrogen atom. (L.C)
2 and 3
163
Define an atomic orbital. (L.C)
An atomic orbital is a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
164
Write the ground state s, electron configuration for a carbon atom. How many orbitals are occupied? (L.C)
1s₂ 2s₂ 2p₂
| 4 orbitals occupied
165
Distinguish between the ground state and the excited states of the electron in a hydrogen atom. (L.C)
Ground - n = 1
| Excited - n = 2,3
166
How can the electron in a hydrogen atom become excited? (L.C)
Add heat
167
Explain the origin of the series of visible lines in the emission spectrum of hydrogen. What name is given to this series? (L.C)
- Excited electrons fall back from n = 3,4 to n = 2, the energy lost is emitted as light of different frequencies.
- Balmer series
168
Explain why there is no yellow line in the hydrogen emission spectrum. (L.C)
No corresponding excited state
169
Describe how to carry out a flame test to confirm the presence of lithium in a salt sample. (L.C)
Clean a platinum (nichrome) wire in concentrated hydrochloric acid
Dip rod in salt and hold salt in hot part of Bunsen flame
Red crimson colour is a positive result for lithium
170
Define an atomic orbital. (L.C)
An atomic orbital is a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
171
Distinguish between a 2p orbital and a 2p sublevel. (L.C)
2p sublevel consists of three 2p orbitals of equal energy
172
Write the s, p electron configuration for a calcium atom. (L.C)
1s₂ 2s₂ 2p₆ 3s₂ 3p₆ 4s₂
173
Explain in terms of energy sublevels why the arrangement of electrons I the main energy levels in a calcium atoms is 2, 8, 8, 2 and not 2, 8, 10. (L.C)
- 4s sublevel lower in energy than the 3d
| - Electrons fill the 4s sublevel before the 3d
174
Explain how the line emission spectrum of hydrogen arises and provides evidence for the existence of energy levels. (L.C)
- In the ground state the hydrogen electron occupies the lowest available energy level.
- The electron can jump to a higher energy level if it recieves a certain amount of energy
- Excited state unstable
Evidence - Energy emitted corresponds to difference between the two energy levels.
175
Suggest an element that gives a blue-green colour to a fireworks display. (L.C)
Copper
176
Write the s, p configuration of a calcium atom in its ground state. (L.C)
1s₂ 2s₂ 2p₆ 3s₂ 3p₆ 4s₂
177
Give one significant difference between an electron in the 2s orbital and an electron in the 3s orbital of a calcium atom. (L.C)
- Energy of 2s electron is less than 3s
178
State the famous principle published in 1927, which bears the name Werner Heisenberg. (L.C)
Position and momentum of an electron cannot be found simultaneously.
179
Name the scientist whose work on energy levels in the hydrogen atom is depicted in the Google doodle reproduced above. (L.C)
Bohr
180
Distinguish between the terms energy level and atomic orbital. (L.C)
Energy level - Is the discrete amount of enegry an electron has when it is in an atom.
An atomic orbital - is a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
181
Write the electron configuration (s, p) of an atom of silicon showing the distribution of electrons in atomic orbitals in the ground state. (L.C)
1s₂ 2s₂ 2p₆ 3s₂ 3p₂
| Distribution on marking scheme (2014)
182
State how many (i) main energy levels, (ii) atomic orbitals are occupied in the silicon atom in its ground state. (L.C)
(i) 3
| (ii) 8
183
Use Bohr's atomic theory of 1913 to account for the emission spectrum of the hydrogen atom. (L.C)
- The electron in a hydrogen atom occupies fixed energy levels.
- In the ground state electrons occupy the lowest available energy levels.
- The electron can move to a higher energy level if it receives a certain amount of heat
- Excited state unstable
- Emitting excess energy in the form of a photon of light
184
Explain in terms of atomic structure, why different flame colours are observed in flame tests using salts of different metals. (L.C)
Metal atoms of different elements have different sets of energy levels therefore they emit different frequencies.
185
What colour is observed in a flame test on lithium chloride? (L.C)
Red
186
Describe the testing procedure in a flame test on lithium chloride. (L.C)
- Salt on platinum probe
| - Hold in at top of flame
187
Explain the term uncertainty principle. (L.C)
This states that it is impossible to know both the position and the speed of an electron at the same time as electrons move in a wave motion.
188
Give one factor that contributed to the need for modification of Bohr's 1913 theory other than Heisenberg's uncertainty principle. (L.C)
- Wave nature of electron
189
What is an atomic orbital? (L.C)
An atomic orbital is a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
190
Define atomic orbital. (L.C)
An atomic orbital is a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
191
Write the electron configuration (s, p etc) of the element manganese (Mn). (L.C)
1s₂ 2s₂ 2p₆ 3s₂ 3p₆ 4s₂ 3d₅
192
What do the electron configurations of the series of elements from scandium to zinc have in common? (L.C)
Electrons entering 3d sublevel / All end in 3d(x)
193
State the number of (i) sub-levels, (ii) orbitals, occupied by electrons in an argon atom in its ground state. (L.C)
(i) 5
| (ii) 9
194
Write the electron configuration (s, p) of an oxygen atom showing the arrangement of electrons in atomic orbitals. (L.C)
1s₂ 2s₂
| Arrangements in marking scheme (2012)
195
Outline Bohr's atomic theory based on the hydrogen emission spectrum. (L.C)
- The electron in a hydrogen atom occupies fixed energy levels
- An electron in an energy level does not radiate energy
- Electron occupies lowest energy levels available
- The electron can move to a higher energy level if it receives an amount of energy
- The photon must be exactly equal to the energy difference between the ground state and a higher energy level.
196
State two limitations of Bohr's theory that led to its modification. (L.C)
- Didn’t work for higher elements
| - Did not allow for uncertainty
197
Draw the shape of the p-orbital. (L.C)
Dumbbell drawn
198
Draw the shape of the s-orbital. (L.C)
In notes
199
State the maximum number of electrons that can be accommodated in a p-orbital. (L.C)
2
200
Write the s, p electron configuration for the potassium atom. (L.C)
1s₂ 2s₂ 2p₆ 3s₂ 3p₆ 4s₁
201
How many (i) energy sub-levels, (ii) individual orbitals, are occupied by electrons in a potassium atoms. (L.C)
(i) 6 sublevels
| (ii) 10 orbitals
202
Explain why there are electrons in the fourth main energy level of potassium although the third main energy level is incomplete. (L.C)
4s sublevel lower in energy than 3d
203
Write the electron configuration (s, p etc.) of the aluminium ion (Al³+) (L.C)
1s₂ 2s₂ 2p₆
204
Name the type of spectroscopy, based on absorptions within a particular range of electromagnetic frequencies, and used as a 'fingerprinting' technique to identify organic and inorganic compounds. (L.C)
infra-red / IR
205
Define energy level. (L.C)
One of the discrete amounts of energy that an electron has when it is in an atom.
206
Distinguish between ground state and excited state for the electron in a hydrogen atom. (L.C)
Ground - in lowest energy state
| Excited - higher energy state
207
Name the series of lines in the visible part of the line spectrum of hydrogen. (L.C)
Balmer series
208
Explain how the expression E2 - E1 = hf links the occurrence of the visible lines in the hydrogen spectrum to energy levels in a hydrogen atom. (L.C)
- E2 - E1 : Energy difference between higher and lower level
- f = frequency of line in spectrum
- h is Plank's constant
- Energy difference divided by frequency equals a constant.
209
Define energy level (L.C)
One of the discrete amounts of energy that an electron has when it is in an atom.
210
Write the electron configuration (s, p) for the sulphur atom in its ground state, showing the arrangement in atomic orbitals of the highest energy electrons. (L.C)
1s₂ 2s₂ 2p₆ 3s₂ 3p₄
| Arrangement in marking scheme (2007)
211
State how many (i) energy levels, (ii) orbitals, are occupied in a sulphur atom in its ground state. (L.C)
(i) 3
| (ii) 9
212
Describe how you would carry out a flame test on a sample of potassium chloride. (L.C)
n
213
Why do different elements have unique atomic spectra? (L.C)
Each element has a different distribution of energy levels giving rise to different electron transitions
214
What instrumental technique is based on the fact that each element has a unique atomic spectra? (L.C)
Atomic absorption spectrometry (AAS)
215
Define atomic orbital. (L.C)
An atomic orbital is a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
216
What does Heisenberg's uncertainty principle say about an electron in an atom? (L.C)
It is not possible to measure the exact position and energy of an electron in an atom simultaneously.
217
Write the electron configuration (s, p etc.) of a chromium atom in its ground state. (L.C)
1s₂ 2s₂ 2p₆ 3s₂ 3p₆ 4s1 3d₅
218
Name the series of coloured lines in the line emission spectrum of hydrogen corresponding to transitions of electron from higher energy levels to the second (n=2) energy level. (L.C)
Balmer series
219
Distinguish between atomic orbital and a sublevel. (L.C)
Atomic orbital - a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
Sublevel - A group of atomic orbitals within an atom, all of which have the same energy.
220
Describe how Bohr used line emission spectra to explain the existence of energy levels in atoms. (L.C)
- electrons in ground state
- fixed energies absorbed
- excited state unstable
- energy difference between levels gives specific frequency of light in spectrum
221
Why does each element have a unique line emission spectrum? (L.C)
- each element has a different distribution of energy levels giving rise to different electron transitions.
222
Name the instrumental technique that can be used to detect heavy metals and to measure their concentrations in a soil or water sample. (L.C)
Atomic absorption spectrometry (AAS)
223
Bohr's atomic theory was later modified. Give one reason why this theory was updated. (L.C)
Only worked for simple atoms
224
Define energy level (L.C)
One of the discrete amounts of energy that an electron has when it is in an atom.
225
Define atomic orbital (L.C)
An atomic orbital is a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
226
Write the electronic configuration (s, p etc.) of nitrogen. (L.C)
1s₂ 2s₂ 2p₃
227
Describe how the electrons are arranged in the orbitals of the highest occupied sub-level of a nitrogen atom in its ground state. (L.C)
One electron in each of the three p orbitals
228
Write the electronic configuration of a neutral copper atom. (L.C)
1s₂ 2s₂ 2p₆ 3s₂ 3p₆ 4s1 3d₁₀
229
Define atomic orbital. (L.C)
An atomic orbital is a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
230
What spectroscopic technique is used to detect heavy metals, e.g. lead, in environmental analysis? (L.C)
atomic absorption spectrometry / AAS
231
What is the colour of the light associated with the line spectrum of sodium? (L.C)
yellow / orange
232
Explain how line emission spectra occur. (L.C)
- electrons restricted to energy levels
- jump to higher levels
- fall back emitting energy as light
- energy difference between levels gives specific frequency of light in spectrum.
233
What evidence do line emission spectra provide for the existence of energy levels in atoms? (L.C)
- only fixed frequencies are emitted
| - therefore electrons must be restricted to certain energy values
234
Why is it possible for line emission spectra to be used to distinguish between different elements? (L.C)
- different elements have different spectra
235
Explain briefly why different metals produce different flame colours. (L.C)
Metal atoms of different elements have different sets of energy levels therefore they emit different frequencies.
236
Write the electronic configuration of the Al³+ ion. What neutral atom has the same configuration? (L.C)
1s₂ 2s₂ 2p₆
