Atomic Structure Flashcards
Stage one - electron impact
Electron impact - electrons are fired at the sample using a high energy electron gun. This causes the ejection of the electron from the sample forming an ion of 1+ charge.
Eg Mg —> Mg+ + e-
The electron gun is a hot wire filament that acts the source of electrons. The ions are always positively charged.
Electron impact is used for elements and substances with low formula mass. Electron impact can cause larger organic molecules to fragmen
Stage 1 - electrospray ionisation
Electrospray ionisation - methord used for high mr compounds such as protiens and other macro molecules.
The sample X is dissolved in a volatile solvent (eg water or methanol) and injected through a fine hypodermic needle to give a fine mist (aerosol).
The tip of the needle is attached to the positive terminal of a high-voltage power supply.
The particles are ionised by gaining a proton (ie an H+ ion which is simply one proton) from the solvent as they leave the needle producing XH+ ions (ions with a single positive charge and a mass of Mr + 1).
X(g) + H+ —> XH+(g)
The solvent evaporates away while the XH+ ions are attracted towards a negative plate where they are accelerated.
Stage 2 acceleration
Positive ions are accelerated by an electric field
To a constant kinetic energy.
Given that all the particles have the same kinetic energy, the velocity of each particle depends on its mass. Lighter particles
have a faster velocity, and heavier particles have a slower
KE=½ mass x velocity.
Stage 3 flight tube
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
Stage 4 detection
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
For each isotope the mass spectrometer can measure a m/z (mass/charge ratio) and an removed from a particle forming a 2+ ion. abundance
When do subshell energy increase?
As we move away from the nucleus
Rules for writing electron configurations
• Electrons enter the lowest energy orbital first.
• Electrons enter the orbitals singularly, they only pair when no empty orbitals are available.
• Pair of electros in orbitals must have opposite spin - this minimises the repulsion between two electrons in the same orbital.
Why does the electron configuration not show individual orbitals in the 2p shell
The eelctron configuration only shows the subshells and not the individual orbitals
The energy of the 3d shell
The energy of the 4s subshell is less than the energy in the 3d subshell. This means we fill the 4s subshell before we fill the 3d subshell
What elements are exeptions to writing the elctron configuration
Chromium and copper
Why is chromium an exception of the electron configuration
The 3d subshell is more stable when it is either half full or completely full.
So in the case of chromium, by having only one electron in the 4s subshell, it can have a half full 3d subshell making it more stable.
Why is chromium an exception of the electron configuration
The 3d subshell is more stable when it is either half full or completely full.
And in the case of copper, by having only one electron in the 4s subshell, it can have a full 3d subshell.
Why is chromium an exception of the electron configuration
The 3d subshell is more stable when it is either half full or completely full.
And in the case of copper, by having only one electron in the 4s subshell, it can have a full 3d subshell.
Easiest way to find the electon configuration
Easiest way to figure out the electron configuration, count the periods downwards, this determines amount of energy levels for example: 3 periods down means that the outer electron is in the third energy level. Next count how many across the their subshell group laterally. This tells us how many electrons are in the subshell of the highest energy level.
Electron configuration of ions
When ions form the subshell with the highest energy gains or loses an electron. In this case the 3s subshell has the highest energy so two electrons are lost from there.
Once the 4s subshell contains electrons, it has now higher energy than the 3d subshell. So when forming ions, the electrons are always lost from the 4s subshell before the 3d subshell. There are no exceptions to this rule.
Electrons absorbing energy
Electrons can absorb energy from heat, electricity or electromagnetic radiation. When it does this, it moves up an energy level.The energy absorbed is then transferred back to the electron once it has fallen to ground state.
Evidence for the energy levels
• Neil Bohr proposed the idea of electrons being found in fixed energy
levels.
• The emission spectrum for the hydrogen atom shows discrete bands-spectral lines.
• These lines have a characteristic frequency and therefore energy value, corresponding to the energy levels of the H atom.
• The discrete energy levels are sometimes referred to as “quanta”-packets.
• The spacing between energy levels is not equal- the energy levels converge as they get further from the nucleus.
How many energy levels does an atom have
Each atom has infinity number of shells however not all of them are occupied with electrons. The lowest energy occupied in an atom is called ground state.
Definition of first ionisation energy
the first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of +1 ions.
X —> X+ + e-
What type of reaction is ionisation
When atoms lose electrons they are positive. This is endothermic and requires energy.
What causes dips in graphs
Subshells
Orbitals
Orbitals:
• each energy level is divided into subshells
• Each subshell is made up of orbitals
• An orbital is three a three dimentional volume of space where this is a probability of finding an electron
• Each orbital can hold up to a maximum of 2 electrons
• There are four types of orbitals, S, P, D and F
Subshell S
Number of orbitals - 1
Maximum number of electrons in shells - 2
Subshell S
Number of orbitals - 1
Maximum number of electrons in shells - 2
Subshell p
Number of orbitals - 3
Maximum number of electrons in shells - 6
Subshell d
Number of orbitals - 5
Maximum number of electrons in shells - 10
Subshell f
Number of orbitals - 7
Maximum number of electrons in shells - 14
Factors that affect ionisation energy:
Nuclear charge
the greatest numbers of protons in the nucleus. The greater the attraction between nucleus and the outer electron. Therefore more energy is needed to knock it off.
Factors that affect ionisation energy - Distance from the nucleus
- if the outer electron is further from the nucleus, the nuclear attraction is weaker meaning less energy is required to remove the electron.
Factors that affect ionisation energy - Shielding by inner electrons
- the greater the number of inner electrons between the outer electron and the nucleus, the weaker the nuclear attraction . Less energy is needed to remove electrons, the lower the ionisation energy.
What does a rapid increase of ionsation energy mean
Changing in energy levels
What happens to first ionisation energy as we go down the group
First ionisation energy decreases down the group because the outer electrons in the elements are further from the nucleus. Because there is more shielding, there is a weaker attraction between the outer electrons and the nucleus so the electron is more easily lost. This means less energy is required to remove an electron.
What happens to first ionisation energy as we move across a period
Across a period -
• 1st ionisation energy increases across a period because there are more protons in the nucleus, the shielding is about the same.
• The distance between the nucleus and outer electrons decreases.
• Therefore the attraction between the outer electrons and nucleus is stronger so more energy is needed to remove an electron.
Why do noble gases not form ions
Their ionisation energy is too high
second ionisation energy
the second ionisation energy is the enrgy needed to remove one mole of electrons from one mole of 1+ ions in their gaseous state to form one mole of 2+ ions (also in their gaseous state)
second ionisation energy
the second ionisation energy is the enrgy needed to remove one mole of electrons from one mole of 1+ ions in their gaseous state to form one mole of 2+ ions (also in their gaseous state)
Why do successive ionisation energies increase?
Compared to the second electron shell, the first shell is closer to the nucleus and electrons in the first shell experience much less shielding. This means the electrons in the first shell have a greater attraction to the nucleus compared to the electrons in the outer shell.
Kinetic energy
0.5xmxv2
Velocity
Distance over time
Avagadros
6.02x10-23
Avagadros
6.02x10-23
Avagadros
6.02x10-23
Time of flight
T = D X (square root) mass over 2KE
How to calculate relative atomic mass
(isotopic mass x % abundance)
Over
100
Measuring the Mr of a molecule
If a molecule is put through a mass spectrometer with an Electron impact ionisation stage 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 Molecular ion equal to the relative molecular mass , molecule.
This peak is called the parent ion or molecular ion
What happens if a molecule is put through a mass spectrometer with electro spray ionisation
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.
Describing the alevel model of the atom
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 dtwo electrons have opposite spin
Definition :First ionisation energy
The first ionisation energy is the enthalpy change when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge
Definition :Second ionisation energy
The second ionisation energy is the enthalpy change when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge
Factors that affect ionisation energy
There are three main factors
1.The attraction of the nucleus
(The more protons in the nucleus the greater the attraction)
2. The distance of the electrons from the nucleus
(The bigger the atom the further the outer electrons are from the nucleus and the
of weaker the attraction to the nucleus)
- Shielding of the attraction of the nucleus
(An electron in an outer shell is repelled by electrons in complete inner shells,
weakening the attraction of the nucleus)
Why are successive ionisation energies 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.
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
Why has helium the largest first ionisation energy?
ts 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
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
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.
Why has Na a much lower first ionisation energy than neon?
This is because 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.
Why is there a small drop 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 sheilded by 3s electron
W hy is there a small drop from P to S?
With sulfur there are 4 electrons in the 3p sub shell and the 4th is starting to doubly fill drops as they are the first 3p orbital.
When the 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
