1. Atomic Structure Flashcards
1
Q
What model used to display atoms was initially thought to be scientifically correct?
A
- Plum pudding model
2
Q
Describe the plum pudding model
A
- Atoms consisted of a sphere of positive charge
- With small negative charges distributed evenly within in
3
Q
What model is used to display atoms in modern day?
A
- Electron shell model
An early model of the atom was the Bohr model (GCSE model) (2 electrons in first shell, 8 in second etc.) with electrons in spherical orbits. Early models of atomic structure predicted that atoms and ions with noble gas electron arrangements should be stable.
4
Q
Why is the nuclear model still used by scientists when more accurate models have since been developed?
A
- The model is easy to draw and understand
- Fits well with most observations e.g. bonding, ionisation, energy trends
5
Q
State the force that holds protons and electrons together
A
- Electrostatic force
6
Q
State the force that holds protons and neutrons together
A
- Strong nuclear force
7
Q
Describe the electron shell model
A
- Atom consists of a small, dense central nucleus
- Surrounded by orbiting electrons in electron shells
8
Q
When was the electron shell model developed?
A
- During the discovery of Rutherford ‘s scattering experiment in 1911
9
Q
What does the nucleus consist of?
A
- Protons and neutrons
- Giving it an overall positive charge
- It contains almost the entire mass of the atom
10
Q
What characteristics does a neutral atom have in terms of its subatomic particles?
A
- The number of electrons is equal to the number of protons due to the relative charges
11
Q
State the relative charge, mass, and position of a proton
A
- Charge: +1
- Mass: 1
- Position: nucleus
12
Q
How would you calculate the number of neutrons?
A
- Mass number - atomic number
13
Q
State the relative charge, mass, and position of a neutron
A
- Charge: 0
- Mass: 1
- Position: nucleus
14
Q
State the relative charge, mass, and position of an electron
A
- Charge: -1
- Mass: 1/1840 (negligible)
- Position: Orbitals
15
Q
What does the maximum number of orbiting electrons that can be held by any single shell depend on?
A
- The number of shells
16
Q
How can you calculate the max number of orbiting electrons that can be held by any single shell?
A
- 2n²
e.g. electrons in shell 2 = 2(2²) = 8 electrons
17
Q
What must happen before the next electron shell is filled?
A
- Each electron shell must fill before the next one can hold any electrons
18
Q
State the symbol that represents mass number
A
- A
19
Q
How can mass number be calculated?
A
- Sum of protons and neutrons in an atom
20
Q
State the symbol that represents atomic mass
A
- Z
21
Q
What is the atomic number equal to?
A
- Equal to the number of protons in an atom
22
Q
Define relative atomic mass
A
- The mean mass of an atom taking into account all of its isotopes compared to the mass of carbon-12 on a scale which an atom of C-12 has a mass of 12
23
Q
What does the relative atomic mass take into account?
A
- The relative abundances of the different isotopes of an element
24
Q
State the symbol that represents relative atomic mass
A
- Ar
25
Describe the R.A.M quoted on the periodic table
* Weighted average of all the isotopes
26
Fig: spectra for magnesium from mass spectrometer
% abundance
20
24Mg+
60 40
If asked to give the species for a peak in a mass spectrum then give charge and mass number e.g. 24Mg+
27
State the equation to find the R.A.M of an element using % abundance
R.A.M = sigma isotopic mass x % abundance) x100
For above example of Mg
R.A.M = [(78.7 x 24) + (10.13 x 25) + (11.17 x 26)] /100 = 24.3
28
State the equation to find the R.A.M of an element using relative abundance
R.A.M = sigma (isotopic mass x relative abundance) total relative abundance
29
How would you calculate the R.A.M of Tellurium?
- 124-Te relative abundance 2; 126-Te relative abundance 4; 128-Te relative abundance 7; 130-Te relative abundance 6
* R.A.M = [(124x2) + (126x4) + (128x7) + (130x6)]
19 = 127.8
30
Explain how you would: Copper has two isotopes 63-Cu and 65-Cu. The relative atomic mass of copper is 63.5. Calculate the percentage abundances of these two isotopes.
Example: Copper has two isotopes 63-Cu and 65-Cu. The relative atomic mass of copper is 63.5. Calculate the percentage abundances of these two isotopes.
63.55 = yx63 + (1-y)x65 63.55 = 63y +65 -65y 63.55 = 65 -2y
2y = 1.45 y = 0.725
%abundance 63-Cu =72.5%
%abundance 65-Cu = 27.5%
31
Mass spectra for Cl2 and Br2
Cl has two isotopes Cl35 (75%) and Cl37(25%) Br has two isotopes Br79 (50%) and Br81(50%)
These lead to the following spectra caused by the diatomic molecules
The 160 peak has double the abundance of the other two peaks because there is double the probability of 160 Br79-Br81 + as can be both Br79-Br81 and Br81-Br79
32
Name ways mass spectrometers have been used in astrology
* Included in planetary space probes so that elements on other planets can be identified.
* Elements on other planets can have a different composition of isotopes.
33
How does electron impact in the ionisation stage affect a molecules spectra / measuring the Mr?
* It will often break up and give a series of peaks caused by the fragments.
* The peak with the largest m/z, however, will be due to the complete molecule and will be equal to the relative molecular mass , Mr ,of the molecule.
* This peak is called the parent ion or molecular ion
34
How does electro spray ionisation in the ionisation stage affect a molecules spectra / measuring the Mr?
* Fragmentation will not occur.
** There will be one peak that will equal the mass of the MH+ ion. It will therefore be necessary to subtract 1 to get the Mr of the molecule.
* So if a peak at 521.1 is for MH+, the relative molecular mass of the molecule is 520.1.
35
State the equation to find Mr (C-12)
* Mr = mean mass of an atom of an element / 1/12 x mean mass of C-12 isotope
36
Define isotopes
* Atoms of the same element with the same atomic number
* but with a different number of neutrons
* resulting in a different mass number
37
Describe the chemical properties of isotopes
* (Neutral atoms of) isotopes have similar chemical properties because they have the same electronic structure (proton number & electron configuration is the same)
* The sharing and transfer of electrons is unaffected
38
Describe the physical properties of isotopes
* However the different mass number means they have different physical properties
39
Describe how ions are formed
* Ions are formed when an atom loses or gains electrons
*Meaning it is no longer neutral and will have an overall charge
40
Define mass spectrometry
* An analytical technique used to identify different isotopes and find the overall relative atomic mass of an element
41
Describe Time of Flight Mass Spectrometry
* A form of mass spectrometry
* Records the time it takes for ions of each isotope to reach a detector
42
What can the mass spectrometer be used to determine?
* The mass spectrometer can be used to determine all the isotopes present in a sample of an element and to therefore identify elements
43
What is produced during TOF mass spectrometry?
* Spectra
44
What does a spectra show?
* Each isotope present
45
State the conditions a mass spectrometer must be under and why
* It needs to be under a vacuum otherwise air particles would ionise and register on the detector
46
State the 5 stages of TOF mass spectrometry
* Ionisation
* Acceleration
* Ion Drift
* Detection
* Analysis
47
Describe what happens during the ionisation stage of TOF mass spectrometry
* A sample of an element is vaporised and injected into the mass spectrometer where a high voltage is passed over the chamber
* Causes electrons to be removed from the atoms (it is ionised) leaving +1 charged ions in the chamber
48
Describe what happens during electron impact (TOF)
* A vaporised sample is injected at low pressure
* An electron gun fires high energy electrons at the sample
* This knocks out an outer electron
* Forming positive ions with different charges e.g. Ti (g)Ti+ (g)+ e–
49
When is electron impact used during TOF?
* For elements and substances with low formula mass.
* Electron impact can cause larger organic molecules to fragment.
50
Describe what happens during electron spray ionisation (TOF)
* The sample is dissolved in a volatile, polar solvent
* Injected through a fine needle giving a fine mist or aerosol
* The tip of needle has high voltage
* at the tip of the needle the sample molecule, M, gains a proton, H+, from the
solvent forming MH+
* M(g) + H+ MH+(g)
* The solvent evaporates away while the MH+ ions move towards a negative plate
51
When is electro spray ionisation used in TOF?
* Preferably for larger organic molecules.
* The ‘softer’ conditions of this technique mean fragmentation does not occur.
52
Describe what happens during the acceleration stage (TOF)
* These positively charged ions are then accelerated by an electric field
* Towards a negatively charged detection plate
* With a constant kinetic energy
53
Describe what happens during the ion drift / flight tube stage (TOF)
* The ions are then deflected by a magnetic field into a curved path
* The radius of their path is dependent on the charge and mass of the ion
*The positive ions with smaller m/z values will have the same kinetic energy as those with larger m/z and will move faster.
*The heavier particles take longer to move through the drift area.
*The ions are distinguished by different flight times
=/
t = time of flight (s)
d = length of flight tube (m)
= velocity of the particle (m s–1
54
Describe what happens during the detection stage (TOF)
* When the positive ions hit the negatively charged detection plate, they gain an electron producing a flow of charge
* The greater the abundance, the greater the current produced
The ions reach the detector and generate a small current, which is fed to a computer for analysis. The current is produced by electrons transferring from the detector to the positive ions. The size of the current is proportional to the abundance of the species
Sometimes two electrons may be removed from a particle forming a 2+ ion. 24Mg2+ with a 2+ charge would have a m/z of 12
For each isotope the mass spectrometer can measure a m/z (mass/charge ratio) and an abundance
55
Describe what happens during the analysis stage of TOF mass spectrometry
* The current values are then used in combination with the flight times to produced a spectra print-out with the relative abundance of each isotope displayed
56
What does the velocity of each particle in the ion drift stage depend on?
* The velocity of each particle depends on its mass.
* Lighter particles have a faster velocity, and heavier particles have a slower velocity.
* (Given that all the particles have the SAME kinetic energy)
57
State the equation to find kinetic energy
* KE = 1⁄2 mv^2
* KE = kinetic energy of particle (J)
* m = mass of the particle (kg)
* v^2 = velocity of the particle (ms–1)
58
A sample of nickel was analysed and one of the isotopes found was 59Ni. The ions were accelerated to have 1.000 x 10-16 J of kinetic energy and travelled through a flight tube that was 0.8000 m long.
How long would one ion of 59Ni+ take to travel along the flight tube?
The Avogadro constant L = 6.022 × 1023 mol–1
A sample of nickel was analysed and one of the isotopes found was 59Ni. The ions were accelerated to have 1.000 x 10-16 J of kinetic energy and travelled through a flight tube that was 0.8000 m long.
How long would one ion of 59Ni+ take to travel along the flight tube?
The Avogadro constant L = 6.022 × 1023 mol–1
59
What may be produced during the ionisation process?
* A 2+ charged ion may be produced
60
Explain how a 2+ charged ion affects the mass to charge ratio
* This means it will be affected more by the magnetic field producing a curved path of a smaller radius
* As a result, its mass to charge ratio (m/z) is halved and this can be seen on spectra as a trace at half the expected m/z value
61
How can Ar be calculated using spectra?
* Ar = m/z x abundance / total abundance
62
What does spectra produced by the mass spectrometry of chlorine display?
* A characteristic pattern in a 3:1 ratio for Cl+ ions
* 3:6:9 ratio for Cl2+ ions
* This is because one isotope is more common than the other and the chlorine molecule can form in different combinations
63
Describe the structure of electrons in an atom
* Held in clouds of negative charge called orbitals
64
What is the difference between orbitals, shells and sub shells?
* orbital is a region, only has 2 electrons
* shells can have numerous electrons
* sub shells
64
orbitals
65
Principle energy levels numbered 1,2,3,4..
1 is closest to nucleus
Sub energy levels labelled s , p, d and f
s holds up to 2 electrons
p holds up to 6 electrons
d holds up to 10 electrons f holds up to 14 electrons
Orbitals which hold up to 2 electrons of opposite spin
66
What are the different types of orbitals?
* s, p, d, f
67
Describe the shape of each orbital
* s-orbital: spherical
* p-orbital: dumbell / infinity
* d-orbital: four-leaf clover
68
Describe the relationship with orbitals and the Periodic Table
* Each orbital correspond with blocks on the Periodic Table
* Each element in the bock has outer electrons in that orbital
69
How many electrons can each orbital fill?
* s-orbital: 2 electrons
* p-orbital: 6 electrons
* d-orbital: 10 electrons
70
Sub-level
1 1s
2 2s, 2p
3 3s, 3p, 3d
4 4s, 4p, 4d, 4f
71
Shapes of orbitals
Orbitals represent the mathematical probabilities of finding an electron at any point within certain spatial distributions around the nucleus.
Each orbital has its own approximate, three dimensional shape.
It is not possible to draw the shape of orbitals precisely.
72
An atom fills up the sub shells in order of increasing energy (note 3d is higher in energy than 4s and so gets filled after the 4s) 1s2s2p3s3p 4s3d4p5s4d5p
73
Describe the order in which orbitals are filled
* The energy of orbitals increases from s to d meaning the orbitals are filled in this order
* Each orbital is filled before the next one is used to hold electrons
74
Writing electronic structure using letters and numbers
*s sublevels are spherical
* p sublevels are shaped like dumbbells
For oxygen 1s2 2s2 2p4
Number of main Name of energy level type of
sub-level
For calcium 1s2 2s2 2p6 3s2 3p6 4s2
75
Explain how you would draw a spin diagram
* An arrow is one electron
* Box represents one orbital
* The arrows going in the opposite direction represents the different spins of the electrons in the orbital
* When filling up sub levels with several orbitals, fill each orbital singly before starting to pair up the electrons
76
The periodic table is split into blocks.
A s block element is one whose outer electron is filling a s-sub shell e.g. sodium 1s2 2s2 2p6 3s1
A p block element is one whose outer electron is filling a p-sub shell e.g. chlorine 1s2 2s2 2p6 3s2 3p5
A d block element is one whose outer electron is filling a d-sub shell e.g. vanadium 1s22s22p63s23p6 4s23d3
77
Show how you would write sodium in terms of its orbitals
* 11 electrons: Na = 1s2 2s3 2p6 3s1
* It has 3 energy levels and 3 orbitals holding 11 electrons
78
Describe spin within an orbital
* Within an orbital, electrons pair up with opposite spin so that the atom is as stable as possible
* Electrons in the same orbital must have opposite spin
79
How is spin represented in an orbital diagram?
* With arrows
80
What are the three rules for writing out electron configurations?
* The lowest energy orbital is filled first
* Electrons with the same spin fill upon orbital first before pairing begins
* No single orbital holds more than 2 electrons
81
What are the exceptions to the rules when writing out electron configurations?
* If electron spins are unpaired and therefore unbalanced, it produces a natural repulsion between the electrons making the atom very unstable
* If this is the case, the electrons may take on a different arrangement to improve stability
82
Give an example of an exception to the rules when writing out electron configurations
* 3p4 contains a single pair of electron with opposite spin making it unstable
* Therefore the electron configuration changes to become 3p3 4s1 which is a much more stable arrangement
83
Electronic structure for ions
When a positive ion is formed electrons are lost from the outermost shell.
Mg is 1s2 2s2 2p6 3s2 becomes Mg2+ is 1s2 2s2 2p6
Electronic structure of d-block elements
When a negative ion is formed electrons are gained O is 1s2 2s2 2p4 becomes O2- is 1s2 2s2 2p6
`* 4s sub shell fills and empties its electrons first
84
Electronic structure of d-block elements
When a negative ion is formed electrons are gained O is 1s2 2s2 2p4 becomes O2- is 1s2 2s2 2p6
The electronic structure of the d-block has some complications. As mentioned earlier, conventionally we say that 4s fills before 3d and so we write them in that order. There is, however, disagreement in the scientific community about whether this is true.
If you look at the electronic structures below you will see both chromium and copper have an unusual arrangement in having a half filled 4s sub shell.
You will also see that when d-block elements form ions they lose the 4s electrons first.
You may find if you research different reasons for these observations. It may well be many of the reasons are false and we have to accept that some things in chemistry don’t neatly follow patterns we can explain. You do need to learn these electronic structure though!
85
Sc 1s22s22p63s23p6 4s23d1 Ti 1s22s22p63s23p6 4s23d2 V 1s22s22p63s23p6 4s23d3 Cr 1s22s22p63s23p6 4s13d5 Mn 1s22s22p63s23p6 4s23d5 Fe 1s22s22p63s23p6 4s23d6 Co 1s22s22p63s23p6 4s23d7 Ni 1s22s22p63s23p6 4s23d8 Cu 1s22s22p63s23p6 4s13d10 Zn 1s22s22p63s23p6 4s23d10
Sc3+ [Ar] 4s03d0 Ti3+ [Ar] 4s03d1 V3+ [Ar] 4s03d2 Cr3+ [Ar] 4s03d3 Mn2+ [Ar] 4s03d5 Fe3+ [Ar] 4s03d5 Co2+ [Ar] 4s03d7 Ni2+ [Ar] 4s03d8 Cu2+ [Ar] 4s03d9 Zn2+ [Ar] 4s03d10
86
Define ionisation energy
* The minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state
* OR the enthalpy change when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge
87
State the equation that represents first ionisation energy
* H(g) ---> H+(g) + e-
* The equation for 1st ionisation energy always follows the same pattern (it does not matter if the atom does not normally form a +1 ion or is not gaseous)
88
Define second ionisation energy
* Enthalpy change when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge
89
State the equation the represents second ionisation energy
* Ti+ (g) --> Ti2+(g) + e-
90
What is ionisation energy measured in?
* kJmol-1
91
State the three main factors that affect ionisation energy
* The attraction of the nucleus (nuclear charge)
* The distance of the electrons from the nucleus (atomic radius)
* Shielding of the attraction of the nucleus
92
Explain how shielding affects ionisation energy
* An electron in an outer shell is repelled by electrons in complete inner shells, weakening the attraction of the nucleus
93
Explain how nuclear charge affects ionisation energy
* The more protons in the nucleus the greater the attraction)
94
Explain how atomic radius affects ionisation energy
* The bigger the atom the further the outer electrons are from the nucleus and the weaker the attraction to the nucleus
95
Describe the importance of patterns in successive ionisation energies
* The patterns in successive ionisation energies for an element give us important information about the electronic structure for that element.
96
When do successive ionisation energies occur?
* When further electrons are removed
* This usually requires more energy because as electrons are removed the electrostatic force of attraction between the positive nucleus and the negative outer electron increases
* More energy is therefore needed to overcome this attraction so ionisation energy increases
97
Explain why successive ionisation energies are always larger
* The second ionisation energy of an element is always bigger than the first ionisation energy. When the first electron is removed a positive ion is formed.
* The ion increases the attraction on the remaining electrons and so the energy required to remove the next electron is larger.
98
How are ionisation energies linked to electronic structure?
* The fifth electron is in a inner shell closer to the nucleus and therefore attracted much more strongly by the nucleus than the fourth electron.
* It also does not have any shielding by inner complete shells of electron
99
What trend does first ionisation energy follow along a period (in the periodic table)?
* First ionisation energy increases due to a decreasing atomic radius and a greater electrostatic forces of attraction
100
1
2
3
4
5
Ionisation energy kJ mol-1
590
1150
4940
6480
8120
Here there is a big jump between the 2nd and 3rd ionisations energies which means that this element must be in group 2 of the periodic table as the 3rd electron is removed from an electron shell closer to the nucleus with less shielding and so has a larger ionisation energy
101
What trend does first ionisation energy follow down a group (in the periodic table)?
* First ionisation energy decreases due to an increasing atomic radius and shielding which reduces the effect of the electrostatic forces of attraction
102
What trend does first ionisation energy follow across a period?
* Nuclear charge increases
* atomic radius is smaller
* higher ionisation energy
* However electrons also experience shielding (same so has no effect)
103
he shape of the graph for periods two and three is similar. A repeating pattern across a period is called periodicity.
The pattern in the first ionisation energy gives us useful information about electronic structure
104
* When successive ionisation energies are plotted on a graph, a sudden large increase indicates a change in energy level
* This is because the electron is being removed from an orbital closer to the nucleus so more energy is required to do so
105
* This large energy increase provides supporting evidence for the atomic orbital theory
106
Why is the first ionisation of Al lower than expected ? check this
* The first ionisation energy of Aluminium is lower than expected due to a single pair of electrons with opposite spin
* As a result there is a natural repulsion which reduces the amount of energy needed to be put in to remove the outer electron
107
patterns in second ionisation energy
If the graph of second ionisation or each successive element is plotted then a similar pattern to the first ionisation energy is observed but all the elements will have shifted one to the left.
The group 1 elements are now at the peaks of the graph
Lithium would now have the second largest ionisation of all elements as its second electron would be removed from the first 1s shell closest to the nucleus and has no shielding effects from inner shells. Li has a bigger second ionisation energy than He as it has more protons.
108
Why has helium the largest first ionisation energy?
* Its first electron is in the first shell closest to the nucleus and has no shielding effects from inner shells. He has a bigger first ionisation energy than H as it has one more proton
109
Why do first ionisation energies decrease down a group?
* As one goes down a group, the outer electrons are found in shells further from the nucleus and are more shielded so the attraction of the nucleus becomes smaller
110
Why is there a general increase in first ionisation energy across a period?
* As one goes across a period the electrons are being added to the same shell which has the same distance from the nucleus and same shielding effect.
* The number of protons increases, however, making the effective attraction of the nucleus greater.
111
Why has Na a much lower first ionisation energy than neon?
* Na will have its outer electron in a 3s shell further from the nucleus and is more shielded.
* So Na’s outer electron is easier to remove and has a lower ionisation energy.
112
Why is there a small drop in ionisation energy from Mg to Al?
* Al is starting to fill a 3p sub shell
* Whereas Mg has its outer electrons in the 3s sub shell.
* The electrons in the 3p subshell are slightly easier to remove because the 3p electrons are higher in energy and are also slightly shielded by the 3s electrons
113
Why is there a small drop in ionisation energy from P to S?
* Sulfur - 4 electrons in the 3p sub shell & 4th is starting to doubly fill the first 3p orbital.
* When second electron is added to a 3p orbital there is a slight repulsion between the two negatively charged electrons which makes the second electron easier to remove.
* Therefore sulphur requires lower ionisation energy
* Phosphorus: 1s2 2s2 2p6 3s2 3p3
* Two electrons of opposite spin in the same orbital
* Sulfur: 1s2 2s2 2p6 3s2 3p4
