Atomic Structure And Periodic Structure Flashcards
1
Q
What are atoms
A
All the stuff elements and compounds are made of
2
Q
What are the subatomic particles that make up a atom
A
Electron neutron and protons
3
Q
What charge do electrons have
A
-1
4
Q
What takes up the most volume in a atom
A
The orbitals
5
Q
Where is most of the mass concentrated in the atom
A
The nucleus
6
Q
Is the diameter of the nucleus big or small compared to the diameter of the atom
A
Small
7
Q
What’s in the nucleus
A
Protons neutrons
8
Q
What is the relative mass of the proton
A
1
9
Q
What is the relative mass of the neutron
A
1
10
Q
What is the relative mass of the electron
A
0.0005
11
Q
What is relative charge of proton
A
+1
12
Q
What is the relative charge of the neutron
A
0
13
Q
What is the relative charge of a electron
A
-1
14
Q
What is mass number
A
Total number of protons and neutrons in the nucleus
15
Q
What is atomic mass
A
Number of protons in the nucleus
16
Q
What is true about protons of an element
A
All atoms of the same element have the same number of protons
17
Q
If a atom is neutral what is true about the protons and electrons
A
The are equal
18
Q
How do you work out the number of neutrons
A
Mass number minus atomic number
19
Q
What does a negative charge on atom mean
A
More electrons than protons
20
Q
What does positive charge on atom mean
A
More protons than electrons
21
Q
What are isotopes
A
Atoms of the same element with same number of protons but different number of neutrons
22
Q
What decides the chemical properties of an element
A
The number and arrangement of electrons
23
Q
why do isotopes of the same element have the same chemical properties have the same chemical properties
A
because they have the same arrangement and number of electrons which determines the chemical properties
24
Q
why do isotopes of an element have different physical properties
A
because physical properties tend to depend on the mass of an atom
25
what do physical properties depend on
mass of the atoms that make it up
26
what is the relative atomic mass
weighted mean mass of an atom compared to 1/12th of the mass of an atom of carbon 12
27
what is the relative isotopic mass
mass of an isotope compared to 1/12th the mass of an atom of carbon 12
28
what is the relative molecular mass or relative formula mass
the average mass of a molecule or formula unit compare to 1/12th of the mass of an atom of carbon 12
29
how do you work out the relative molecular mass
add up the relative atomic mass values
30
what is relative molecular mass used for
simple molecules
31
what is relative formula mass used for
used for compounds that are ionic or giant covalent
32
how do you work out the relative formula mass
add the relative atomic masses of the ions or atoms
33
how to work out the relative atomic mass of an element from its isotopic abundance
step 1: multiply each relative isotopic mass by its % relative isotopic abundance and add up the results
step 2: divide by 100
34
where are mass spectra produced from
mass spectrometers
35
what are mass spectrometers used for
devices used to find out what samples are made up of by measuring the masses of their components
36
what can mass spectra tell us
relative isotopic masses and abundances of different elements
37
what is on the y axis and x axis on a mass spectra
abundance of ions on the y axis
and m/z values on the x axis
38
what is a m/z values
mass/charge ratio
39
what can you assume about the m/z value
charge of ions is +1 so we can assume x axis is the relative isotopic mass
40
how to work out relative atomic mass from a graph
step 1:multiply each relative isotopic mass by its relative isotopic abundance and add up the results
step 2:divide by the sum of the isotopic abundances
41
Silicon can exist in three isotopes. 92.23% of silicon is 28Si and 4.67% of silicon is 29Si. Given that the Ar of silicon is 28.1, calculate the abundance and isotopic mass of the third isotope.
100% – 92.23% – 4.67% = 3.10%
28.1 = ((28 × 92.23) + (29 × 4.67) + (X × 3.10)) ÷ 100
28.1 = (2717.87 + (X × 3.10)) ÷ 100
2810 – 2717.87 = X × 3.10
29.719 = X
So the isotopic mass of the third isotope is 30
42
Chlorine has two isotopes. 35Cl has an abundance of 75%
and 37Cl has an abundance of 25%. Predict the mass spectrum of Cl2.
step 1:
35Cl – 35Cl: 0.75 × 0.75= 0.5625
35Cl – 37Cl: 0.25 × 0.75= 0.1875
37Cl – 35Cl: 0.25 × 0.75= 0.1875
37Cl – 37Cl: 0.25 × 0.25= 0.0625
step 2: 0.1875 + 0.1875 = 0.375.
step 3:Divide all the relative abundances by the smallest relative abundance to get the smallest whole number ratio.
[(35Cl – 35Cl) (35 + 35 = 70) (0.5625 ÷ 0.0625 = 9)]
[(35Cl – 37Cl) (35 + 37 = 72) (0.375 ÷ 0.0625 = 6)]
[(37Cl – 37Cl) (37 + 37 = 74) (0.0625 ÷ 0.0625 = 1)]
43
what happens when molecules are bombarded with electrons in mass spectrometry
electron removed from molecule to form molecular ion m+
44
how do you molecular mass of a compound when looking at mass spectra
look at molecular ion peak
45
on a mass spectra which peak is the molecular ion peak
the one with the biggest m/z value
46
what do electrons move around the nucleus on
quantum shells/energy levels
47
what do shells contain
subshells
48
what do subshells contain
orbitals
49
how many electrons can orbitals hold
2
50
how many electrons can s subshell hold
2
51
how many electrons can p subshell hold
6
52
how many electrons can d subshell hold
10
53
how many electrons can f subshells hold
14
54
what is an orbital
a bit of space an electron moves in
55
what is true about the energy of orbitals within the same subshells
same energy
56
what is spin pairing
when electrons within the same orbital have to spin in the opposite direction
57
what shape do s orbitals have
spherical
58
what shape do p orbitals have
dumbell shape
59
how many p orbitals are there and what direction are they to each other
3 p orbitals
and 90 degrees to each other
60
how can you represent electrons in orbitals
arrows in in boxes
61
how do electrons fill up subshells
fill up the lowest energies first
62
why does the 4s subshell fill up before the 3d subshell
becuase the 4s subshell has a lower energy level than the 3d subshell despite having a bigger quantumn number
63
how do electrons fill up orbitals
fill up orbitals singly before they pair up
64
electron configuration of argon /[Ar]
[Ar]1s²2s²2p⁶3s²3p⁶
65
what is the outer shell configuration of an s block element
s¹ or s²
66
what is the outer shell configuration of a p block element
s²p¹ to s²p⁶
67
what do chromium and copper do that other d Block elements don't do
donate one of their 4s electrons to the 3d sub shell
68
why do chromium and copper donate one of their 4s electrons to the 3d sub shell
because they're mores stable with a full or half full subshell
69
what is electromagnetic radiation
energy transmitted as waves with a spectrum of different frequencies
70
what is the order of the electromagnetic spectrum in increasing energy/frequency and decreasing wavelength
radio waves, microwaves, infrared, visible light, ultraviolet, x rays, gamma rays
71
from radio waves to gamma rays what happens in terms of frequency / energy
increasing frequency / energy
72
what happens in terms of wavelength from radio waves to gamma rays
decreasing wavelength
73
mnemonic for em spectrum for increasing energy
rich men in vegas use expensive girls
74
what happens when an electron is in ground state
electrons are on the lowest possible energy level
75
what can an electrons do if they take in energy
move to higher energy levels
76
what is it called when a electron moves to higher energy level
excited state
77
how do electrons release energy
by dropping down from higher to lower energy level
78
energy levels are discrete because
the energy levels have fixed values
79
what do emission spectrum show
shows the frequencies of light emitted when electrons drop down from higher energy level to a lower one
80
what do emission spectra look like
frequencies Emitted appear as coloured lines on a dark background
81
why are line spectra unique for each element
each element has different electron arrangement so frequencies of radiation absorbed and emitted are different
82
why is there sets of lines in emission spectra
because each set represents electrons moving to different energy levels
83
in emission spectra why do they lines get closer together
because frequency increases
84
what are the four principles when it coms to electron shells
*electrons can only exist in fixed orbitals or shells not anywhere between
*when an electron moves between shells electromagnetic radiation is emitted or absorbed
*each shell has fixed energy
*because energy of shell is fixed the radiation will have fixed frequency
85
what supports the idea that energy levels are discrete
atom has clear lines for different energy levels
86
describe how the electron gets to one energy level to the other
jumps from one to another with no in-between stage
87
what provide evidence that electron exist in quantum shells
*emission spectra
88
what is ionisation
the removal of one or more electrons
89
what is the First ionisation energy
the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions
90
what symbol must you always use when writing ionisation equations
gas (g)
91
is ionisation an endothermic process or a exothermic process and why
endothermic process because energy is put in to ionise
92
what are the three factors that affect ionisation energy
nuclear charge
electron shell
shielding
93
how does nuclear charge affect ionisation energy
the more protons there are in the nucleus the more positively charged the nucleus and the stronger the attraction for the electrons
94
how does electron shell affect the ionisation energy
Attraction falls off very rapidly with distance. An electron in an electron shell close to the nucleus will be much more strongly attracted than one in a shell further away
95
how does shielding affect the ionisation energy
Attraction falls off very rapidly with distance. An electron in an electron shell close to the nucleus will be much more strongly attracted than one in a shell further away.
96
what does a high ionisation energy mean
means there’s a strong attraction between the electron and the nucleus, so more energy is needed to overcome the attraction and remove the electron.
97
what happens to ionisation energy as you go down the group
As you go down a group in the periodic table, ionisation energies generally fall, i.e. it gets easier to remove outer electrons.
98
why does ionisation energy decrease down a group
*Elements further down a group have extra electron shells compared to ones above. The extra shells mean that the atomic radius is larger, so the outer electrons are further away from the nucleus, which greatly reduces their attraction to the nucleus.
*The extra inner shells shield the outer electrons from the attraction of the nucleus.
99
what is the second ionisation energy
The second ionisation energy is the energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions.
100
what does a graph of successive ionisation energy provide evidence for
A graph of successive ionisation energies provides evidence for the shell structure of atoms.
101
what does decrease in ionisation energy provide evidence for
A decrease in ionisation energy going down a group provides evidence that electron shells really exist.
102
why do successive ionisation energies increase
This is because electrons are being removed from an increasingly positive ion — there’s less repulsion amongst the remaining electrons, so they’re held more strongly by the nucleus.
103
what are the big jumps on a successive ionisation energy graph
The big jumps in ionisation energy happen when a new shell is broken into — an electron is being removed from a shell closer to the nucleus.
2 electrons x from 1st shell.
104
who was the first scientist to put the elements in a meaningful order
Dmitri Mendeleev
105
how is the periodic table arranged
periods (rows) and groups (columns).
106
what is true about the elements in a period
All the elements within a period have the same number of electron shells
107
what is periodicity
repeating trends in the physical and chemical properties of the elements
across each period
108
All the elements within a group have the same number of electrons in their outer shell what does this mean
This means they have similar chemical properties.
109
how many outer shell electrons do s block elements have
The s-block elements (Groups 1 and 2) have 1 or 2 1s outer shell electrons
110
what happens to atomic radius across a period
Atomic Radius Decreases across a Period
111
why does atomic radius decrease across a period
As the number of protons increases, the positive charge of the nucleus increases. This means electrons are pulled closer to the nucleus, making the atomic radius smaller.
112
what happens to ionisation energy across a period
Ionisation Energy Increases Across a Period
113
why does ionisation energy increase across a period
because the number of protons is increasing, which means a stronger nuclear attraction extra electrons are roughly at the same level and so there's little extra shielding affects or extra distance to lessen the attraction
114
what is true about removing electrons from higher energy sub shells compared to lower energy sub shells
Generally, it requires more energy to remove an electron from a higher energy subshell than a lower energy
subshell
115
why are there do drops between group 2 and 3 on first ionisation energy graphs
sub shell structure
116
why is there drops between group 5 and 6 on first ionisation graph
electron repulsion
117
what is true about singly filled full and partially filled sub shells
In general, elements with singly filled or full subshells are more stable than
those with partially filled subshells, so have higher first ionisation energies.
118
what happens to the boiling points in metals
For the metals (Li, Be, Na, Mg and Al), melting and boiling points increase across the period because the metallic bonds get stronger. The bonds get stronger because the metal ions have an increasing number of delocalised electrons and a decreasing radius
(i.e. the metal ions have a higher charge density). This means there’s a stronger attraction between the metal ions and delocalised electrons, so stronger metallic bonding.
119
what happens to boiling points of giant covalent lattices
The elements with giant covalent lattice structures (C and Si) have strong covalent bonds linking all their atoms together.
A lot of energy is needed to break all of these bonds. So, for example, carbon (as graphite or diamond) and silicon have the highest boiling points in their periods.
120
what happens to the boiling points of simple molecular structures
Next come the simple molecular structures (N2, O2 and F2, P4, S8 and Cl2). Their melting points depend upon the strength of the London forces between their molecules. London forces are weak and easily overcome, so these elements have low melting and boiling points.More electrons in a molecule mean stronger London forces . For example, in Period 3 a molecule of sulfur (S8) has the most electrons, so it’s got higher melting and boiling points than phosphorus and chlorine.
121
what happens to boiling points of noble gases
The noble gases (Ne and Ar) have the lowest melting and boiling points in their periods because they exist as individual atoms (they’re monatomic) resulting in very weak London forces.
