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Atomic Structure Flashcards

1
Q

What do atoms consist of?

A

A central nucleus

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2
Q

What does this central nucleus contain?

A

Protons and neutrons

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3
Q

What is the nucleus surrounded by?

A

Electrons in energy levels/shells

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4
Q

State what else you know about the atom

A

Atoms have no electric charge
They contain the same number of protons as electrons
The nucleus is tiny compared to the size of the whole atom

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5
Q

What is the atomic number?

A

The number of protons

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6
Q

What is the mass number?

A

The number of protons + the number of neutrons

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7
Q

Name the relative masses of all the subatomic particles

A

Proton= 1
Neutron = 1
Electron= 1/1836

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8
Q

State the relative charges of all the subatomic particles

A

Proton= +1
Neutron= 0
Electron= -1

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9
Q

How do you calculate the number of neutrons?

A

Mass number - atomic number

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10
Q

What do atoms of the same element have?

A

The same number of protons
It is the number of protons that determines what type of atom it is
Atoms of different elements have different numbers of protons

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11
Q

What is an isotope?

A

Atoms with the same number of protons but a different number of neutrons
Atoms of the same element with the same atomic number but a different mass number

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12
Q

How are electrons arranged in atoms and ions?

A

Electrons are arranged in electron shells/ energy levels- which have sub shells
Each sub level consists of electron orbitals
Each orbital can hold two electrons with opposite spins

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13
Q

What are electron orbitals?

A

A region of space in which the electron spends most of its time- most likely to be in

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14
Q

What are the different sub levels?

A

s
p
d
f

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15
Q

How many orbitals can be held in each sub level?

A

s= 1
p= 3
d= 5
Only need to know these three and not d

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16
Q

How are the electron orbitals filled up?

A

Electrons enter the lowest energy orbital available
Electrons prefer to occupy orbitals on their own and only pair up when no empty orbitals of the same energy are available

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17
Q

What is the aufbau principle?

A

When electrons enter the lowest energy orbital available
Electrons fill the lowest orbital of energy first starting with the lowest energy level and working its way up to the highest energy level
Fills the 1s before the 2s

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18
Q

What is hunds rule?

A

Where electrons prefer to occupy orbitals on their own and only pair up when no empty orbitals of the same energy are available
Electrons prefer to be in orbitals on their own before pairing up as this requires more energy

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19
Q

How do electron orbitals fill up in ions?

A

The electrons in the highest energy levels are lost first
When losing electrons- electrons are lost from 4s before 4d

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20
Q

What happens when electrons fill 4s and 3d?

A

4s goes above 3d in energy
These energy levels are very close

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21
Q

What are two exceptions to the expected pattern?

A

Chromium - Cr
Copper- Cu

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22
Q

What is the electron configuration for chromium?

A

1s2 2s2 2p6 3s2 3p6 4s1 3d5
NOT
1s2 2s2 2p6 3s2 3p6 4s2 3d4

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23
Q

What is the electron configuration for copper?

A

1s2 2s2 2p6 3s2 3p6 4s1 3d10
NOT
1s2 2s2 2p6 3s2 3p6 4s2 3d9
This is slightly lower in energy arrangement than Cr as the reduced electron repulsion makes up for the fact one electron is in a slightly higher energy level

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24
Q

What does ionisation energies give evidence for?

A

How electrons are arranged in atoms

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25
What is the first ionisation energy?
The energy required to remove one electron from each atom in a mole of gaseous atoms producing one mole of 1+ gaseous ions
26
What is the second ionisation energy?
The energy required to remove the second electron Not both electrons
27
What three factors affect ionisation energy?
Distance from the nucleus- atomic radius Nuclear charge Shielding
28
What happens to the first ionisation energies in group 2 ?
They decrease down the group
29
Why is there a dip in first ionisation energies from group 2 to group 3 ?
Group 2 = s orbital Group 3 = p orbital Electron is easier to remove from p orbital than s orbital- takes less energy( less is needed)
30
Why is there a dip in first ionisation energies from group 5 to group 6 ?
Electron pair repulsion
31
Why does argon have a higher first ionisation energy than potassium?
Argon has less shielding Electron from argon is from shell 3 but shell 4 in potassium
32
Why does phosphorous have a higher first ionisation energy than sulfur?
Less electron repulsion in phosphorus There is a lot of electron pair repulsion in sulphur Electron in phosphorous is from orbital with one electron Sulphur has this same orbital with two electrons
33
Why does magnesium have a higher first ionisation energy than calcium?
Less electron shielding in magnesium and so a smaller atomic radius(stronger attraction to nucleus) Calcium has one more energy level than magnesium- larger atomic radius(less attraction and so not as much energy is needed) Ionisation energy decreases from magnesium to calcium
34
Why does magnesium have a higher first ionisation energy than aluminium?
Electron in magnesium is from 3s but from 3p in aluminium 3p is higher in nervy than 3s
35
Why does fluorine have a higher first ionisation energy than oxygen?
Fluorine has a smaller atomic radius than oxygen Atomic radius decreases across the periodic table Fluorine has more protons than oxygen and so fluorine has a higher nuclear charge
36
Why does sodium have a higher second ionisation energy than magnesium?
Na+ has less electron shielding Na+ has a smaller radius Electron in Na+ is from second shell where as third shell for Mg+
37
Why does sodium have a higher second ionisation energy than neon?
Na+ has more protons than Ne+ Therefore Na+ has a higher nuclear charge and higher ionisation energy
38
Why does sulfur have a higher second ionisation energy than phosphorus?
S+ has more protons than P+ S+ has a higher nuclear charge and so a higher ionisation energy
39
Why is it harder to remove the second energy electron?
The second ionisation energy is higher than the first ionisation energy Removing it from a positive ion
40
What is the equation for kinetic energy?
KE= 1/2mv2 OR mv2/2
41
How is relative atomic mass calculated?
From the mass numbers of all the isotopes
42
State what you know about a mass spectrometer
They weigh atoms Lighter ions travel faster than heavier ones Atoms are turned into ions
43
What do you know about relative abundance?
Shows the abundance of each peak relative to each other
44
What is percentage abundance?
Compares what’s detected to 100
45
What is m/z?
Mass to charge ratio
46
Define relative atomic mass
Mean mass of an atom measured in relation to 1/12 th the mass of a carbon 12 atom
47
How is the periodic table ordered?
Elements based on atomic number
48
What does abundance mean?
The quantity or amount of an isotope present in a sample expressed as a percentage or ratio
49
What is a mass spectrometer?
Analytical machine used in laboratories Measures masses of different elements,isotopes or compounds
50
What 3 things does a mass spectrometer measure?
Relative atomic mass Relative molecular mass Abundance of isotopes
51
What are the four main steps in mass spectrometery?
Element is placed in mass spectrometer and atoms get ionised Ions are accelerate using an electric field and are sent through a flight tube or ion drift At the end of the ion drift abundance’s are measured - how long it took them in the flight tube Data is used to calculate the mass of the ion
52
What happens during ionisation?
An atom or molecule needs to gain a charge so it can be accelerated in the flight tube Two methods of doing this: Remove an electron - electron impact ionisation Add a proton- electro spray ionisation
53
What is electron impact ionisation?
The sample is bombarded with high energy electrons These knock electrons off and leave atoms with a positive charge Atom/ molecule loses an electron This takes place in s gaseous state
54
What is electro spray ionisation?
Sample is dissolved in a volatile solvent like methanol This is injected through a positively charged needle at a high voltage This causes each molecule to gain a proton from the solvent
55
What happens during acceleration?
Once ions are in form they are attracted to the negatively charged plate or electric field At this stage all the ions are given the same amount of kinetic energy Atoms or molecules are first ionised so they can be accelerated your the electric field
56
What is ion drift?
Ions travel before they reach the detector Each ion has the same kinetic energy As you increase the mass of the atom it’s velocity will decrease Atoms travel at different speeds
57
What is the difference between heavier and lighter ions?
Heavier ions take longer to reach the detector Lighter ions reach the detector first
58
What happens during detection?
Ions hit the detector plate after travelling down the flight tube This plate is negatively charged so when positive ions hit it a current flows
59
What is the abundance of the ion proportional to?
The size of the current that flows The flow of current is proportional to abundance of that ion
60
How can the mass of the ion be calculated?
From The time it took the ion to reach the detector
61
What is the equation for the mass of an ion?
m= 2KE/v2
62
What is constant in these calculations?
Distance if the ion drift Kinetic energy inside the mass spectrometer
63
What is the mass of one mole of atoms?
The relative atomic mass
64
How to calculate the mass of an ion?
RAM/La
65
What is La in equations?
Avogadro constant 6.022 x10^23 atoms
66
Equation for velocity?
V= d/t
67
How did schrodinger describe electrons?
As waves and clouds of probability Where you are most likely to find these electrons
68
What does each energy level contain?
Orbitals Each orbital has a unique shape
69
State what you know about the s orbital
Lowest energy of the orbitals 1 s orbital per energy level Can have 2 electrons Looks like a sphere
70
State what you know about the p orbital
Next in energy from s 3 p orbitals in each energy level Can have 2 electrons in each orbital Looks like two barons with their ends tied together
71
State what you know about the d orbital
Highest energy orbital Can have 5 d orbitals per energy level Looks like 4 balloons with their ends tied together
72
What do not all energy levels have?
All the orbitals First energy level doesn’t have p or d orbitals Second energy level doesn’t have d orbital
73
What does the first energy level contain?
Just the 1s orbital
74
What does the second energy level have?
2s sub shell and 2p sub shell
75
Why does the third energy level have?
3s sub shell 3p sub shell 3d subshell
76
What does the fourth energy level have?
4s sub shell 4p sub shell
77
What can each orbital hold?
Two electrons All orbitals hold different amounts of energy
78
What is special about the 3d and 4s orbitals?
3d orbital has more energy than the 4s orbital Important to remember
79
How is chromium an exception?
Fills normally until 4s orbital Instead of placing two in it there is one instead and then continues to the 3d orbitals
80
How is copper an exception?
Fills in a similar way to chromium but goes further and pairs up the electrons in the 3d orbital instead of going back and filling the 4s orbital This leaves a loan electron
81
What is short hand electronic structure?
The electron configuration starts with the symbol of the noble gas in the previous period followed by the additional configuration of the electrons for the given element
82
What are the basic fundamentals of electronic structure?
In each shell there are sub shells There are s,p and d sub shells Not all shells have the sub shells Each sub shell has orbitals
83
What groups are the s block on the periodic table?
Group 1 and 2 Highest energy electron level is in s orbital
84
What do you need to remember about the 4s orbital?
Empty it first and fill it first
85
What 4 things effect ionisation energy?
Nuclear charge- number of protons Atomic radius- distance from the nucleus Electron repulsion Electron shielding
86
What do bigger atoms have?
Electrons further away from the nucleus Force of attraction between the nucleus and electron is lower
87
What does electron shielding do?
Lowers the amount of energy needed to ionise atoms
88
What is different for each element?
Ionisation energy
89
State what you know about nuclear charge in relation to ionisation energy
Number of protons in nucleus- increasing number of protons in nucleus increases nuclear charge The more protons in the nucleus the stronger the attraction between the nucleus and electrons Nuclear charge increases across and down the periodic table Increase in nuclear charge increases ionisation energy
90
State what you know about atomic radius in relation to ionisation e
The further the outer electrons are from the nucleus the weaker the attraction Atomic radius decreases across the periodic table and up the group As atomic radius decreases ionisation energy goes up
91
State what you know about electron repulsion in relation to ionisation energy
All electrons are held in orbitals As electrons pair up inside orbitals there is a repulsion between the negatively charged particles The repulsion makes it easier to remove the highest energy electron Most notable from nitrogen to oxygen As electron repulsion increases ionisation energy decreases
92
State what you know about electron shielding in relation to ionisation energy
Electrons in orbitals between the nucleus and highest energy electrons lower the attraction on the nucleus- makes them easier to remove This increases as you go down the periodic table p orbitals are shielded more from the nucleus than s orbitals As electron shielding increases, ionisation energy decreases

Chemistry (7 decks)

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