Elements, Atomic Structure, Radioactivity, Electronic Structure of Atoms Flashcards
1
Q
What do atoms consist of?
A
Three sub atomic particles and a dense nucleus
2
Q
What are the three subatomic particles found in an atom?
A
Electron, neutron, proton
3
Q
What is the relative charge of -
Proton
Neutron
Electron
A
P- +1
N- neutral
E- -1
4
Q
What is the relative mass [amu] of-
Proton
Neutron
Electron
A
P-1
N-1
E-1/1838
5
Q
What is the location within an atom of
Proton
Neutron
Electron
A
P- Nucleus
N- Nucleus
E- In energy levels orbiting the nucleus
6
Q
Who discovered proton?
A
Rutherford
7
Q
Who discovered neutron?
A
Chadwick
8
Q
Who discovered the electron?
A
JJ Thomson
9
Q
What do electrons occupy in an atom?
A
Definite energy levels orbiting the nucleus
10
Q
Who discovered the existence of energy levels?
A
Niels Bohr
11
Q
How did Niels Bohr find the existence of energy levels
A
He studied the emission spectra of hydrogen atoms that had been given energy and found a series of lines of specific energy, indicating that only specific energy jumps [absorbance and emissions] were possible in an atom
12
Q
What provides strong evidence for the existence of energy levels within the atom?
A
Line spectra
13
Q
What is the name given to the spectra of Paschen, Balmer, Lyman?
A
Electromagnetic spectrum
14
Q
What region of the electromagnetic spectrum does Paschen lie?
A
Infrared [passion-red]
15
Q
What region of the electromagnetic spectrum does Balmer lie?
A
Visible
16
Q
What region of the electromagnetic spectrum does Lyman lie?
A
Ultraviolet
17
Q
What is an energy level?
A
It’s the fixed amount of energy that an electron can have in an atom
18
Q
Where do energy levels start in atom?
A
Closest to Nucleus
19
Q
What formula did Neil’s Bohr produce to work out the maximum number of electrons that would fit each orbital?
A
2n^2
20
Q
What are energy levels further divided into
A
Sublevels
21
Q
What are sublevels?
A
Group of atomic orbitals, all of which have the same energy
22
Q
What is an orbital?
A
A region in space where there is a high probability of finding electrons
23
Q
What shape is an s orbital?
A
Spherical
24
Q
What shape is the P orbital?
A
Dumbbell
25
Features of electrons
- occupy lowest available energy level
- fit max of two electrons in each sublevel
- wave-particle dualism
- not possible to determine the position and velocity of an electron at the same time
26
What causes the nuclear charge in an atom?
Protons
27
What is the atomic number?
It’s the number of protons in the nucleus of an atom
28
What does the atomic number tell you?
PEP
- position on periodic table
- number of electrons present
- number of protons
29
What is a mass number?
The sum of the number of protons and neutrons in the nucleus of an atom
30
What are isotopes?
Atoms of an element that contain the same number of protons but different number of neutrons and thus have different atomic masses
31
What is the relative atomic mass?
It’s the average mass of an atom of an element relative to one twelfth the mass of an atom of carbon-12
32
Why do we need a relative atomic mass?
In order to represent different varieties of isotopes and relative abundance’s
33
What is radioactivity?
It’s the spontaneous breaking up of unstable nuclei accompanied by the emission of radiation
34
What are the three different types of radiation?
Alpha particle, Beta particle, Gamma Ray
35
What is an alpha particle?
Helium nuclei with a positive charge and little penetrating ability
36
What is a beta particle?
Electrons with a negative charge and greater penetrating ability than alpha particles
37
What are gamma rays?
High energy electromagnetic radiation, with greater penetrating ability than beta particles
38
What stops penetrating power of an alpha particle?
A sheet of paper
39
What stops penetrating power of a beta particle?
5mm of aluminium
40
What stops penetrating power of an alpha particle?
Thick block of lead
41
Example of alpha particle and use
Americium-241
| Smoke detectors
42
Example of beta particle and use
Carbon-14
| Carbon dating [archaeology]
43
Example of gamma ray and use
Cobalt-60
| Radiotherapy
44
What is the reason for the difference in mass numbers of isotopes?
Different number of neutrons
45
What did the ancient greeks propose?
All substances are composed of four elements - earth, air, water and fire
46
What did Robert Boyle propose?
Element definition
| -an element is a substance that cannot be broken down into simpler substances by chemical means
47
What did John Dalton propose about the atom?
- atoms are tiny, indivisible, indestructible particles
| - that are unique to each element
48
What did John Dalton propose about the element?
Elements combine to form compounds, they do so in fixed whole number ratios
49
What is the law of conservation of mass?
It states that matter is neither created nor destroyed in the course of a chemical reaction
50
What did Crookes discover?
Cathode rays
| Maltese cross placed in path of rays, cast sharp shadow at end of discharge tube
51
What did Johnstone Stony propose?
Suggested a name for the negative particles [cathode rays] - electrons
52
How did JJ Thomson discover electron?
Showed cathode rays were negatively charged particles when attracted to positive plate in electric field
53
What did Thomson discover
Electron
| Their charge to mass ratio
54
What did Henry Mosley discover?
Atomic number
55
How did Henry Mosley discover the atomic number?
Using X-rays, measured the number of positive charges [later called protons] in atoms and found the number was different for every element - unique
56
What was the significance of the discovery of the atomic number?
Confirmed Mendeleev’s positioning of some elements in periodic table
57
What did Millikan discover?
Measured change of electron accurately, conducting his oil drop experiment
58
What did Rutherford discover?
Nucleus and Proton
59
How did Rutherford discover the nucleus?
Bombarded gold foil with alpha particles
60
What was the result of the Geiger-marsden experiment?
- few were repelled back
- most pass straight through
- some were deflected at large angles
61
What was concluded from Geiger marsden experiment?
Atom consists mostly of space, with a dense positive nucleus
62
How was the proton discovered?
Rutherford bombarded atom’s nuclei and eventually pushed out positive particles known as protons
63
What did James Chadwick discover?
Neutron
64
How was the Neutron discovered?
Bombarded beryllium with alpha particles which caused neutral particles with same mass as protons to be pushed out
65
What did Niels Bohr discover?
The existence of energy levels, that electrons have fixed energies
Defined energy levels and 2n^2
66
What does the Heisenberg uncertainty principle state?
It is impossible to measure the velocity and position of an electron at the same time
67
What did Humphrey Davy discover?
Isolated new elements such as calcium and magnesium using electrochemical techniques
68
What did Dobereiner propose?
Law of triad -
Chemicals with similar properties were to be grouped in threes, the middle one having a mass midway between the other two
Discovered, Li, Na, K
69
What did Newlands propose?
Law of octaves
| Chemical properties of every eight element were repeated
70
What did Mendeleev propose?
Periodic table
| Arranged elements known and unknown in periodic lines according to increasing atomic weight
71
What did Becquerel discover?
Spontaneous radioactive emissions
72
What did Marie Curie discover?
Isolated two highly radioactive elements -polonium and radium
73
Who coined the term “radioactivity”
Marie Curie
74
What substance was used to discover polonium and radium?
Purified Pitchblende
75
What substance was used to discover spontaneous radioactive emissions?
Uranium salts
76
What were the features of elements in a triad?
- similar chemical properties
| - atomic weight of the middle elements would be midway between the other two
77
What was the law of triads limitation?
Restricted to only a small number of elements
78
What link made by dobereiner was essentially important in drawing up periodic table?
Link between different elements depended upon atomic weights
79
Limitations of the law of octaves
- tried to force all known elements to fit pattern and fitted very reactive metals in the same group as unreactive metals
- only worked for a small number of elements
80
What relationship was made in the periodic table?
Properties of elements and their atomic weights
81
Differences between Mendeleev’s periodic table and the Modern Periodic table
```
Mendeleev -
gaps left for undiscovered elements
in order of increasing atomic weight
Modern -
gaps filled
in order of increasing atomic number
```
82
What property of cathode rays allowed Thomson to measure charge to mass ratio
They are deflected by magnetic fields
83
Describe Thompson’s plum pudding model
Atoms were positive spheres in which negatively charged electrons were embedded
84
Who disproved plum pudding model?
Rutherford
| instead - mostly made up of empty space with positive dense nucleus, electrons orbiting nucleus
85
What was the Bohr Model of the Atom?
Electrons were arranged in a series of concentric circular orbits at an increasing distance from the nucleus
86
Name the series of coloured lines in the line emission spectrum of hydrogen corresponding to transitions of electrons from higher energy levels to the second n=2 energy level
Balmer series
87
Name the series of coloured lines in the line emission spectrum of hydrogen corresponding to transitions of electrons from n=3 to n=1
Ultraviolet region
88
Give an example of a radioactive isotope and state one common use made of this isotope?
Americium -241 [smoke alarms]
Carbon -14 [archaeological dating]
Cobalt-60 [radiotherapy]
89
What is used to calculate the relative atomic mass of an element?
Mass spectrometer
90
What is the principle of the mass spectrometer?
Positive charged ions are separated on the basis of their relative masses moving in a magnetic field
91
What are the five processes in mass spectrometry?
- vaporisation
- ionisation
- acceleration
- separation
- detection
92
Give a use of mass spectrometry
Drug tests
93
Explain what occurs in vaporisation in mass spectrometry
Sample becomes heated and turns into a gas
94
Explain what occurs in ionisation in mass spectrometry
High energy electrons are shot at sample which knocks out an electron to form a positive ion
95
Explain what occurs in acceleration in mass spectrometry
Passes through an electric field to accelerate positive ions to high speed
96
Explain what occurs in separation in mass spectrometry
Pass through magnetic field and are deflected : lighter ions are deflected more than heavier ions. in this way, particles with different masses can be separates and identified
97
Explain what occurs in detection in mass spectrometry
ions are detected
98
What instrument was fundamental in detecting the existence of isotopes?
Mass spectrometer
99
What is the difference between a nuclear reaction and a chemical reaction?
Chemical - changes in distribution of electrons, cannot change one element into another element
Nuclear - changes in nucleus and can cause elements to change into other elements
100
What is a radioisotope?
Unstable, radioactive isotopes
101
What is the half-life of a radioactive isotope?
It’s the time taken for half of the atoms in a sample of the isotope to decay
102
What gas makes up more than half of the background radiation?
Radon gas
103
What are we told from carrying out the flame test
Each metal has their own characteristic colour, therefore the colours obtained can be used to identify the metals present in unknown compounds
104
Why is the line spectrum of each element unique?
each element has a different electron configuration
| giving rise to different electron transitions (jumps)
105
What colour does barium turn in flame tests?
Green
106
What colour does copper turn in flame tests?
Blue-green
107
What colour does lithium turn in flame tests?
Deep red
108
What colour does potassium turn in flame tests?
Lilac
109
What colour does sodium turn in flame tests?
Yellow
110
What colour does strontium turn in flame tests?
Red
111
Procedure for flame tests
- using a soak wooden splint
- crush a little bit of the salt to be tested using mortar and pestle
- dip wooden splint in salt and put over Bunsen flame and note the characteristic colour given off
- repeat the experiment with other salts
112
Why is the emission spectra useful in identifying an element?
The emission spectrum of an element is characteristic of that element due to the fact that atoms of different elements contain different sets of energy levels (due to different number of electrons in shell) and therefore emit characteristic colour
113
What is the ground state?
The lowest energy state [level]
114
What is the excited state?
A higher energy state [level]
115
What occurs when energy is emitted in an atom?
Electron moves from a higher energy level to a lower energy level
116
Under normal circumstances, where is the hydrogen electron found?
The ground stage
117
What occurs if an electron receives enough energy?
It will move to an excited state [here it’s relatively unstable]
118
Why does the electron in the hydrogen often atom eventually fall back down?
It is unstable in the excited state, therefore drops back down emitting energy
119
What is the aufbau principle?
States that electrons will occupy the lowest energy sublevel available
120
What is the atomic radius of an element?
It is half the distance between the nuclei of two atoms of the element that are joined together by a single covalent bond
121
Describe and account for the trend in atomic radii of the elements across the second period
Decreases
Increase in Nuclear charge - exerts a greater attractive force on the outer electrons
No screening effect - electrons are held closer and therefore radius is smaller
122
Describe and account for the trend in atomic radii of the elements going down a group
Increases
Addition of extra energy levels -increase in number of filled shells
Screening effect - electrons in inner shells partially neutralise the nuclear charge by repelling the outer electron
123
What is first ionisation energy
It’s the energy required to remove the most loosely bound electron from an isolated atom in its ground state
124
What is the equation of first ionisation energy?
X -> X^+ + e^-
125
What is the first ionisation energy measured in?
KJ per mole
126
Describe and account for the trend in first ionisation energy of the elements across period
General increase -
Increase in nuclear charge which holds the electrons more tightly making them difficult to remove
Decrease in atomic radius holds electrons in more tightly and also make them difficult to remove
No screening effect
127
Describe and account for the trend in first ionisation energy of the elements going down a group
Decrease
Increase in atomic radius makes it easier to remove an electron despite the increased nuclear charge
Screening effect of inner levels - electrons in inner shells partially neutralise the nuclear charge by repelling the outer electron
128
Account for the decrease in first ionisation energy between nitrogen and oxygen
[write out electronic configuration]
Nitrogen is more stable due to its half filled 3p orbital
| Oxygen has a pair of electrons in px sublevel. If electron was taken away it would be stable.
129
Explain why the second ionisation energy of sodium is significantly higher than the first while the increase in the second ionisation energy or Neon compared to its first is relatively small
[write out electronic configuration]
```
Na = 1st electron removed is from the third shell which gives sodium a high stability configuration, but loss of second is from high stability atom, higher
Ne = both electrons are removed from same sublevel
```
130
Peaks in trend table
Explain :
High peaks, low peaks and medium peaks
[draw electronic configuration]
The high peaks - represent already stable atoms with filled p sublevel
Medium peaks - half filled sub levels or filled sublevels
Low troughs - atoms happy to have electrons removed due to unstable electronic configuration
```
HIGH = STABLE
LOW = UNSTABLE
```
131
What is the second ionisation energy?
Minimum energy required to remove the most loosely bound electron from each monopositive ion in a mole of these ions
132
How is the second ionisation energy represented in an equation?
X^+ -> X2+ + e^-
133
How does the definition of second ionisation energy differ from that of first ionisation energy?
The electron is removed from each monopositive ion
134
Evidence supporting Bohr Theory
- ionisation energy always rises due to protons in the nucleus pulling on less and less electrons causing a decrease in the radius of an atom
- energy jumps = small differences between ionisation energies for electrons belonging to the same shell but large differences of electrons are removed from different shells
135
Bohrs limitations
- only worked for hydrogen
- did not take into account of wave-particle duality
- did not take into account Heisenberg’s uncertainty principle
- didn’t include the discovery of sublevels
136
Features of transition metals
High densities
| High boiling /meltingpoints
137
Calculate the mass of sodium chloride required to prepare 500cm^3 of a .9% w/v saline solution
```
G/500 =.9
G= 500(.9)
G= 450
450/100 - 4-5 g
[per 100cm^3]
```
138
M^2+ has 25 electrons and 32 neutrons
| What is the atomic number and the mass number?
Atomic number [protons] - 27
| Mass number [protons and neutrons] - 59
139
How many electrons and neutrons has the aluminium ion
27
13 Al ^3+
Electrons [-3] - 13-3 = 10
| Neutrons - 27-13 = 14
140
Explain why relative atomic masses are rarely whole numbers
Average of mass numbers of isotopes of an element
141
What changes take place in the nucleus of an atom when beta decay occurs?
Neutron changes into proton
| Electron emitted
142
When 19.05g of copper reacted with nitrogen, 20.45g of copper nitride.
Deduce the empirical formula of copper nitride
Cu3N
19. 05/64 - .3
20. 45-19.05 = 1.4/78 = .1
3: 1
143
Find the empirical formula of a compound containing 40% sulfur and 60% oxygen, by mass
40/32 - 1.25
60/16 - 3.75
1:3
SO3
144
Calculate the daly mass of potassium iodide needed to supply .15mg of iodide ion
.15/127 - .00118
| .00118 x 166 - .196
145
How many neutrons are there in
| .14g of carbon 14
.14/14 - .01
.01 x 6 x 10^23 x 8 NEUTRONS =
4.8 x10^22
146
When hydrogen gas was passed over 1.59g of copper oxide, 1.27g of metallic copper were produced. Find by calculation the empirical formula of the copper oxide
1.59 - 1.27 = .32
1.27/64 = .02
.32/16 = .02
147
Complete and balance the equation for the chemical reaction that occurs when sodium is added to ethanol
C2H5OH + Na ->
C2H5OH + Na -> C2H5ONa + 1/2 H2
148
State the number of sublevels and orbitals in an argon atom (atomic no : 18)
1s2,2s2,2p6,3s2,3p6
Sublevels - 5 (s,s,p,s,p)
Orbitals - 9 (s,s,px,py,pz,s,px,py,pz)
149
Write a balanced nuclear equation for the beta particle decay of iodine-131 (atomic no - 53)
53 -1
| 131 I -> 0 é + Xe
150
Give two properties of cathode rays
Positively charged
Negligible mass
Travel in straight lines
151
Write the electron configuration of the oxygen (oxide) ion O2-
1s2,2s2,2p6
152
How many atoms of iron are there in a 30g bowl of cornflakes that contains .0024g of iron per 30g serving
.0024/mr(56) = 4.29 x 10-5 x 6 x 10^23 = 2.6 x 10^19
153
How many iron atoms should be consumed daily to meet the recommended daily intake of iron in a diet of 0.014g
.014 / 56 = .00025 x 6 x10^23 = 1.5 x 10^20
154
A 500cm^3 can of beer contains 21.5cm^3 of ethanol. Calculate its % alcohol
21.5/500 x 100/1 = 4.3%
155
Name the scientist who identified cathode rays as subatomic particles
JJ. Thomson
156
Calculate the percentage carbon, by mass, in methylbenzene
Carbon (mr of 12) - 7 carbons x 12 = 84
Hydrogen (mr of 1) = 8 hydrogens x 1 = 8
84/84+8 x100/1 = 91.3%
157
What was the nature of the bombarding particles in rutherfords experiment and from what material was the foil made?
Alpha particles
| Gold foil
158
Draw the shape of the three p orbitals
(diagram of x and y axis and the dumbell on all three axis)
159
Describe how Bohr used line emission spectra to explain the existence of energy levels in an atom / Account for the emission spectrum of the hydrogen atom
Studied emission spectra of hydrogen and found only specific energy jumps occured as electron occupies fixed energy levels. Normally, the hydrogen atom is in its ground state which is the lowest energy state, however when it is supplied energy it moves to a higher energy level or excited state. However it is unstable and loses energy to fall back to a lower energy level. The energy difference between levels gives a spefic wavelength of light in spectrum.
160
Account for the sharp decrease in ionisation energy values between elements 18 and 19
18 - stable due to full p sublevel
| 19 - Electron is removed from fourth shell
161
Calculate the relative atomic mass
multiply percentage by mass and add the two isotopes and divide by 100
162
Outline the contribution of JJ.Thomson
Plump pudding model - pos spehere with neg electrons embedded
Charge to mass ratio - electron
163
Why is it difficult to specify the absolute boundary of an atom
Heisenberg uncertainty principle
164
The full set of successive ionisation energy values for the electrons in carbon [1086,2353,4620,6223,37831,47277]
How does this set provide evidence for
i) number of electrons in a carbon atom
ii) number of electrons in each main energy level in a carbon atom
i) each value is one electron
| ii) large increase after fourth value -> four electrons in outer shell
165
How can the electron in a hydrogen atom become excited?
Add heat [energy]
166
Explain the origin of the series of visible lines in the emission spectrum of hydrogen
When an excited electrons falls down from a higher energy level falls to the second energy level. The energy lost is emitted as light of different frequencies
167
Explain why there is no yellow line in the hydrogen emission spectrum
No corresponding energy transition [energy loss]
168
Distinguish between a 2p orbital and a 2p sublevel
2p sublevel consists of three 2p orbitals of equal energy
169
Explain in terms of energy sublevels why the arrangement of electrons is 2,8,8,2 and not 2,8,10 in calcium atom
4s sublevel is lower in energy than 3d
170
Give one way of detecting the presence of cathode rays
Shadow cast by anode
171
Give one significant difference between an electron in the 2s orbital and an electron in the 3s orbital of a calcium atom
- > Energy of 2s electron is less than 3s
| - > Nuclear attraction for 2s greater than 3s
172
How many main energy levels and atomic orbitals are occupied in the silicon atom in its ground state? (atomic number - 14)
E.L - 3
| Orbitals - 8
173
Explain why the first ionisation energy value of silicon (14) is
i) greater than that of aluminium (13)
ii) less than that of carbon (6)
i) Smaller atomic radius and greater nuclear charge
| ii) Greater atomic radius and lossely bound electron is more shielded from nucleus
174
Explain how the graph provides evidence for energy levels in the silicon atom
```
Sharp increase (jump) in ionisation energy from 4th to 5th showing that this is the first electron to be removed from 2nd main level
Sharp (bigger) increase (jump) in ionisation energy from (from 12th to 13th showing that this is the first electron to be removed from 1st main level.
```
175
Give one other factor that also contributed to the need for modification of Bohr's 1913 Theory (other than heisenbergs principle)
Wave-particle duality
| Didn't account for discovery of sublevels
176
What as the basis used by Mendeleev in arranging the elements in his periodic table?
When arranged elements in increasing atomic weight there is a periodic occurrence of similar elements (similar chem properties) which were arranged in groups
177
Give two reasons why Mendeleev left spaces in his periodic table
For undiscovered elements
| Similar elements with similar chemical properties fit better to groups
178
Why did Mendeleev place the element tellurium before iodine?
Fit chemical properties better to groups despite the fact tellurium has a higher atomic weight.
179
Explain why the alkali metals are reactive and why this reactivity increases down the group
i) Readily lose single electron in outer shell
ii) Going down, atomic radius increases due to number of shells added and therefore electrons are further away from nucleus and outer electron is readily lost
180
What resut was expected by Rutherford when thin gold foil was first bombarded with alpha particles?
Most pass straight though with only slight changes of direction
181
How was the model of the atomic structure changed as a result of Rutherfords experiment
Nucleus - positive charge is not spread out but concentrated into small dense nucleus in centre
Empty space - rest of atom is empty space,occupied only by electrons to balance positive charge
182
Equation for alpha decay
Element = new element by -4 mass number and minus 2 the atomic number + 4 2 He
183
Equaton for beta decay
Element = new element by -1 atomic number + 0 1 é
184
DECAY MATHS QUESTION - Radium undergoes alpha decay, starting with 10 x 10-4 moles of radium-223, how many of these atoms remain when 87.5% of the sample has decayed
1 ) how much is left in percentage (100-87.5 = 12.5)
2 ) multiply 12.5 x original moles of atoms and divide by 100 = .0000125
3 ) multiply by avogadros constant - 6 x 10^23 = 7.5 x 10^18
185
In what order were the subatomic particles discovered?
Electron (JJ Thomson)
Proton (Rutherford)
Neutron (Chadwick)
186
State two assumptions of Dalton's atomic theory of 1808
- Small
- Indivisible
- Identical atomic mass of particular element
187
Who measured the ratio of charge to mass of the electron e/m?
JJ Thomson
188
Who proved that the electrons in an atom reside in an electron cloud surrounding a small dense positive central nucleus
Rutherford
189
Who measured the charge on the electron?
Milikan
190
State the max number of electrons that can be accommodated in a p orbital
2
191
Mendeleev predicted the properties of the elements gallium and germanium years before either of them was discovered. Explain the basis for his predictions
Arranged elements in order of inc atomic weight and elements with similar chemical properties in groups with gaps for the undiscovered elements
192
Graph for ionisation energies of first 31 elements, name B and what is the numerical value of x (e.papers pg.64)
Helium
| look at periodic table to get ionisation energy
193
Explain why some alpha particles were deflected at large angles as they passed through the gold foil
Repelled when passing near positive nucleus
194
Why were some alpha particles reflected back along their original paths
Collided with dense nucleus
195
Why were only a very small number of particles reflected back along their original paths
Nucleus is very small and most of the atom is empty space
196
Draw a labelled diagram to show the new structure of the atom proposed by Rutherford
Circle with small nucleus in centre and the circle is labelled electron cloud
197
Explain why carbon 12 dating can be used
When alive, C12 and C14 ratio is the same however when it dies, the unstable C14 decays and stable c12 remains the same
198
Alpha particles are hazardous to health, state one risk associated with exposure to alpha radiation
Causes cancer
199
Explain why the occupants of a house fitted with smoke detectors containing americium-241 are not at risk from alpha radiation emitted by these devices
Not very penetrating
200
Householders are advised to replace the batteries in smoke detectors regularly. Explain whether or not americium-241 needs to be replaced regularly also
No, half life is over 400 years
201
Account for the sharp decrease in ionisation energies between 18 and 19
18- stable due to full 3p orbital, harder as the electron is being removed from 3p
19- outer electron further from nucleus
202
1s2, 2s2, 2p6, 3s2, 3p6, 4s1
Explain why there are electrons in the fourth main energy level of potassium although the third main energy level is incomplete
4s is lower in energy than 3d
203
Explain how the expression E2-E1 = hf links the occurence of the visible lines in the hydrogen spectrum to energy levels in a hydrogen atom
E2-E1 = energy difference between higher (e.g. E2) and lower (e.g. E1)
f : frequency of line in spectrum
each definite frequency produced are due to electrons falling from a particular higher level to particular lower level
the expression indicates that the energy difference (E2 – E1) is proportional to the frequency (f)
204
What instrumental technique is based on the fact that each element has unique atomic spectra
AAS
205
''The minimum energy required to completely remove the most loosely bound electron from a mole of gaseous atoms in their ground state'' defines an important property of every element
i) Identify the energy quantity defined above
ii) State the unit used to measure this quanity
i) first ionisation energy
| ii) KJ per mole
206
Would it take more or less energy to remove the most loosely bound electron from an atom if that electron were not in its ground state? Explain
Less - already gained energy and also further from nucleus
207
An element has a low first ionisation energy value and a low electronegativity value. What does this information tell you about how reactive the element is likely to be, and what is likely to happen to the atoms of the element when they react
Reactive
| Lose electrons and become positively charged
208
How did Moseley help clarify the perodic table?
He discoered atomic number and clarified the periodic table by increasing atomic number whereas before it was increasing atomic mass
