topic 1 - atomic structure and the periodic table Flashcards
1
Q
What do mass spectrometers produce?
A
Mass spectra
2
Q
What can mass spectrometers identify and compare?
A
The relative isotopic masses and relative abundances of different elements by producing mass spectra
3
Q
Why are mass spectrometers used to analyze atoms?
A
Because individual atoms are too small to be counted and weighed
4
Q
What is the first step in using a mass spectrometer?
A
A gaseous sample is placed into the device
5
Q
How is the sample ionized in mass spectrometry?
A
By bombardment with electrons
6
Q
Why are atoms/molecules converted to positive ions in mass spectrometry?
A
For detection in the gaseous/vapour state
7
Q
Why is there a high vacuum inside the mass spectrometer?
A
To prevent interference from atoms and molecules in air
8
Q
How are positive ions accelerated in a mass spectrometer?
A
They are accelerated towards a negatively charged detection plate due to an electric field
9
Q
How does the mass analyser separate ions in mass spectrometry?
A
By mass:charge ratio (m/z)
10
Q
How does a magnetic field affect ions in a mass spectrometer?
A
Ions are deflected into a curved path, with heavier ions experiencing less deflection and ions with more positive charge experiencing more deflection
11
Q
What does the ion detector in mass spectrometry do?
A
It detects and counts the number of ions for each different m/z value
12
Q
What happens when positive ions hit the negatively charged detection plate?
A
They gain an electron, producing a flow of charge
13
Q
What determines the current produced by the ion detector?
A
The greater the abundance, the greater the current produced
14
Q
What information can be obtained from the mass spectrum graph?
A
Number of isotopes, most abundant and least abundant isotopes, isotopic composition, RAM, and relative molecular masses of compounds
15
Q
What is assumed about the charge of the ion in mass spectrometry?
A
It is assumed that the ion produced has a 1+ charge
16
Q
What happens if an ion with greater charge (2+) is produced?
A
The ion is affected more by the magnetic field, producing a curved path of smaller radius
17
Q
How does a 2+ charge affect the mass:charge ratio on the mass spectrum?
A
The mass:charge ratio is halved and appears as half the expected m/z value
18
Q
What does the small peak at half the expected m/z value represent?
A
A 2+ charged ion
19
Q
What causes the existence of a small peak in mass spectrometry?
A
- More (specific number) electrons knocked out at ionisation stage, so same mass to charge ratio as charge increases/decreases (relate to context of question)
20
Q
How do you calculate the RAM from a mass spectrum?
A
Multiply each relative isotopic mass by its relative isotopic abundance, add the results, and divide by the sum of isotopic abundances
21
Q
What might cause the RAM of a sample to be different from the RAM given in periodic table?
A
- Different abundances of isotopes
- Other isotopes could be present
22
Q
Where is most of the mass of an atom concentrated?
A
Most of the mass of the atom is concentrated in the nucleus as it consists of protons and neutrons.
23
Q
What is the atomic number?
A
The number of protons in the nucleus of an atom – this is also the number of electrons.
24
Q
What is the mass number?
A
The number of protons and neutrons in the nucleus of an atom.
25
What are isotopes?
Atoms with the same number of protons in the nucleus but a different number of neutrons (same atomic number, different mass number).
26
What chemical properties do isotopes have?
Isotopes have the same chemical properties as these are decided by the number and arrangement of electrons.
27
What physical properties do isotopes have?
Isotopes have different physical properties as these depend on the mass of the atom.
28
What is relative isotopic mass?
The mass of one atom of an isotope relative to 1/12th the mass of a carbon-12 atom. The values are relative and do not have units.
29
Why are relative isotopic masses not always whole numbers?
They are not always whole numbers as the mass of a proton is slightly less than the mass of a neutron, though they usually are whole numbers.
30
What is relative atomic mass?
The average mass of an atom of an element (containing a mixture of isotopes) relative to 1/12th the mass of a carbon-12 atom. The values are relative and do not have units.
31
Why are relative atomic masses not whole numbers?
They are the average value for the mixture of isotopes found naturally, so they are not in whole numbers.
32
What is relative molecular mass?
The relative molecular mass of an element or compound is the sum of the relative atomic masses (RAM) of all the atoms in its molecular formula.
33
What is relative formula mass?
The relative formula mass of a compound is the sum of the relative atomic masses of all the atoms or ions in its formula. This applies to ionic or giant covalent compounds, where ions do not affect the mass.
34
How do you calculate relative atomic mass?
∑(isotopic abundance * isotopic mass number)/100
35
What happens to energy levels of shells as they move further from the nucleus?
Shells further from the nucleus have a greater energy level than those close to the nucleus.
36
What do shells contain?
Shells contain different types of subshells which have different numbers of orbitals.
37
What is an orbital?
An orbital is a region of space where electrons move randomly, and it can hold up to *two electrons* with opposite spins, and there is *>95% probability of finding an electron*
38
How is each orbital defined?
Each orbital is defined by its energy level, shape and direction in space
39
What are the shapes of s and p orbitals?
s-orbitals are spherical, whereas p-orbitals have dumbbell shapes. There are three p-orbitals at right angles to one another.
40
Do orbitals in the same subshell have different energy levels?
No, orbitals within the same subshell have the same energy; electrons occupy the lowest energy orbitals before filling higher energy orbitals.
41
Why must electrons spin in opposite directions within an orbital?
Electrons must spin in opposite directions so the magnetic attraction from their opposite spins can counteract the electrical repulsion of their negative charges, a process called spin-pairing.
42
How do electrons occupy orbitals in relation to pairing?
Electrons occupy orbitals singly before pairing up to minimize repulsion, known as the Aufbau principle.
43
What happens if electrons are unpaired?
If electrons' spins are unpaired and unbalanced, it creates repulsion between them, making the atom unstable. Electrons may rearrange to improve stability
44
What happens to the energy levels as shells get further from the nucleus?
As shells get further from the nucleus they become closer in energy. The difference in energy between the second and third shells is smaller than between the first and second shells.
45
What occurs between the third and fourth shells?
When the fourth shell is reached, there is an overlap between the 3d orbital of the third shell and the 4s orbital of the fourth shell. This means 4s orbital is filled before 3d
46
Which electrons are lost first when a d-block element ionises?
The 4s orbital as it is the outer orbital
47
What evidence from an ionisation diagram shows that _?
* _3p subshell is higher in energy than the 3s
* Ionisation energy of Al < than that of Mg
* _Maximum number of electrons that can be occupied in a subshell
* Number of elements before a drop in energy
* _No more than three unpaired electrons can be accommodated in the 3p subshell
* Fall in energy from P to S
* Idea that there is a gradual increase from Al to P, followed by a drop when electron pairing is necessary
48
What shows what block an element belongs to?
* Depends which orbital highest energy electron is in
49
What is electromagnetic radiation?
Electromagnetic radiation is energy transmitted as waves, with a spectrum of different frequencies.
50
What happens to frequency and wavelength along the electromagnetic spectrum?
Along the spectrum, radiation increases in frequency and decreases in wavelength.
51
Where do electrons exist in an atom?
Electrons exist in fixed energy levels, with each level having a fixed value of energy.
52
What is the ground state of an atom?
In their ground state, atoms have their electrons in the lowest possible energy levels.
53
What happens if electrons absorb energy?
If electrons absorb energy, they move to higher energy levels further from the nucleus, becoming excited.
54
What happens if electrons emit energy?
If electrons emit energy, they drop down to a lower energy level.
55
Why does radiation have a fixed frequency?
As the energy of shells is fixed and the distance between them, the radiation emitted or absorbed has a fixed frequency.
56
Can electrons exist between energy levels?
No, electrons can only exist in fixed orbitals and not anywhere in between.
57
What is a line spectrum (emission spectrum)?
A line spectrum shows the frequencies of light emitted as a line when electrons drop from a higher to a lower energy level after being excited and moving up, supporting the idea of discrete energy levels. (include idea from flashcard 58)
58
Why do different elements emit different colors of light?
Each element has different energy gaps between subshells, so they emit different colors as *specific* frequencies of radiation absorbed and emitted vary as different electronic arrangement
59
When is the wavelength of light visible?
The wavelength of light is only visible if emitted within the visible region of the electromagnetic spectrum.
60
What does each set of lines in a line spectrum represent?
Each set of lines represents electrons moving to different energy levels, supporting the idea that energy levels have fixed values.
61
Why do lines in a line spectrum get closer together?
The lines get closer together as the frequency increases, meaning higher energy corresponds to higher frequency.
62
What would happen if shells/subshells didn’t exist?
If shells/subshells didn’t exist, electrons could take any energy on a continuum, and any wavelength of light could be emitted and white light observed
63
Why does an ion of an element may be smaller than atom of same element?
* Number of shells may vary, *electrons may be held more tightly*
* Same number of protons attracting fewer electrons
* Less repulsion between (remaining electrons)
64
What is the first ionisation energy?
The first ionisation energy is the energy needed to remove one electron from each atom in a mole of gaseous atoms to form one mole of gaseous 1+ ions.
65
What is first ionisation energy of successive elements evidence of?
Electron subshells.
66
Why is ionisation an endothermic process?
Ionisation is an endothermic process because energy is absorbed to ionise an atom or molecule and overcome the attraction between an electron and the nucleus.
67
What does a high ionisation energy indicate?
A high ionisation energy indicates a strong attraction between the electron and the nucleus, requiring more energy to overcome this attraction and remove the electron.
68
What are the three factors that affect ionisation energy?
The three factors are nuclear charge, electron shell (distance from nucleus), and shielding.
69
How does nuclear charge affect ionisation energy?
As the number of protons increases, the nucleus becomes more positively charged, strengthening the attraction for the outermost electrons.
70
How does distance from the nucleus affect ionisation energy?
As the distance between the nucleus and the outermost electron increases, the attraction becomes weaker.
71
How does shielding affect ionisation energy?
Inner electron shells create a 'barrier' that reduces the attraction between the nucleus and the outer electrons, known as effective nuclear charge.
72
What does a graph of successive ionisation energies suggest?
A graph of successive ionisation energies suggests that electrons exist in energy levels within atoms and can indicate the group an element belongs to.
73
What does a sudden large increase in ionisation energy indicate?
A sudden large increase indicates a change in energy level, as the electron is being removed from an orbital closer to the nucleus, requiring more energy.
74
Why does ionisation energy increase with successive ionisations?
With each successive ionisation, electrons are being removed from an increasingly positive ion with less repulsion among the remaining electrons, so they are held more strongly by the nucleus.
75
Why is the outer electron relatively easy to remove?
The outer electron is relatively easy to remove due to shielding, which reduces the attraction between the nucleus and the outer electron.
76
What do all elements in a period (row) have in common?
All elements in a period have the same number of electron shells, leading to repeating trends in physical and chemical properties.
77
What do all elements in a group (column) have in common?
All elements in a group have the same number of electrons in their outer shell *electronic configuration*, leading to a trend of similar chemical properties.
78
What happens to ionisation energy along a period?
Ionisation energy increases due to a decreasing atomic radius and greater forces of attraction.
79
What causes the increase in ionisation energy as you move along a period?
The atomic number increases, leading to a greater positive charge in the nucleus, which pulls electrons closer and decreases atomic radius.
80
How do extra electrons added across a period affect shielding?
The extra electrons are added to the outer energy level, making no difference to shielding.
81
What are the exceptions to the trend in ionisation energy in Period 2?
In Period 2, boron (B) and oxygen (O) are exceptions because they contain unpaired electrons that require less energy to remove, resulting in lower first ionisation energies.
82
What are the exceptions to the trend in ionisation energy in Period 3?
In Period 3, aluminium (Al) and silicon (Si) are exceptions for the same reason as in Period 2; they have unpaired electrons that require less energy to remove, resulting in lower first ionisation energies.
83
Additionally, what causes the increase in ionisation energy from Al to Si in Period 3?
* Si has *greater nuclear charge* (not atomic number) than Al, and electron removed from the same subshell (3p)
* Greater force of attraction between the nucleus and outer electrons/outer electrons (or highest energy) held more strongly by nucleus
84
What happens to the number of shells as you move down a group?
The number of shells increases, resulting in a greater distance between the nucleus and the outermost shell.
85
How does the increasing atomic radius down a group affect ionisation energy?
Electrons are located further away from the nucleus due to the increase in number of shells, leading to reduced electrostatic attraction and a larger atomic radius.
86
What is the effect of increased nuclear charge on ionisation energy down a group?
Despite an increase in nuclear charge, the effects of electron shielding and increased atomic radius outweigh this increase, resulting in lower ionisation energy.
87
What happens to ionisation energy down a group?
Ionisation energy decreases down a group due to an increasing atomic radius and electron shielding which reduces the effect of the electrostatic forces of attraction.
88
Why is there a drop in ionisation energy from Group 2 to 3?
Electron removed from the group has to come from a 3p orbital, which is further from the nucleus and has additional shielding provided by the electrons in 3s orbital - override the effect of increased nuclear charge
89
Why is there a drop in ionisation energy from Group 5 to 6?
Due to electron pairing, as elements in Group 5 have stable electrons configuration with unpaired electrons whereas in Group 6 elements have to begin pairing up electrons causing repulsion so less energy is needed to remove outer electron in Group 6 than 5
90
What factors determine the melting temperature of an element?
The melting temperature depends on both the element's structure and the type of bonding between its atoms.
91
How do metallic bonds influence melting temperatures in Groups 1-3?
In metals, a higher number of delocalised electrons and a decreasing atomic radius strengthen the bonding between ions and electrons, resulting in higher melting temperatures that increase from Group 1 to Group 3.
92
What is unique about the melting behavior of carbon and silicon in Group 4?
Carbon and silicon have giant covalent structures with strong, highly directional covalent bonds, meaning that many bonds must break before the solid can melt.
93
What type of structures do non-metal elements in Groups 5, 6, 7, and 0 form?
These elements form simple molecular structures with weak intermolecular forces, resulting in low melting temperatures.
94
Why do noble gases have the lowest melting temperatures?
Noble gases exist as monatomic elements with full outer shell electrons, leading to weak forces between the individual atoms, resulting in the lowest melting temperatures.
