Atomic structure Flashcards
Proton
Relative mass: 1
Relative charge: +1
Neutron
Relative mass: 1
Relative charge:0
Electron
Relative mass: 1/1836
Relative charge: -1
Structure of an atom:
An atom has a positively charged nucleus ( containing protons and neutrons) and is surrounded by shells containing electrons
Atomic Number (Z)
Number of protons in the nucleus of an atom
Mass Number (A) (Nucleon Number)
Sum of the protons and neutron in the nucleus
Relative atomic mass
Relative atomic mass is the average mass of an atom of an element relative to one twelfth of the mass of an atom of carbon-12
Relative molecular mass
Relative molecular mass is the average mass of a molecule relative to one twelfth of the mass of an atom of carbon-12
Relative isotopic mass
Relative isotopic mass is the mass of one atom of an isotope compared to one twelfth of the mass of one atom of carbon-12
Isotopes
Atoms of the same element that have the same number of protons but different number of neutrons
Analysing mass spectrum
The abundance of the isotopes is always shown on the Y axis. It can be shown as a % or as a nominal value. If it is in %, all your isotopes must add to give 100%
M/z is always shown on the X as just the mass of the isotope divided by the charge. As most have just a +1 charge, this is the same as their isotopic mass
The highest peak is taken as the “base peak” and in the relative abundance method, it is assigned as 100% and all other peaks are assigned as a percentage of that.
No of peaks = number of isotopes
How to work out relative atomic mass?
To work out the relative atomic mass:
R.A.M = (isotopic mass of A x relative abundance of A) + (isotopic mass of B x relative abundance of B) / total relative abundance (usually 100)
How to predict mass spectra
To predict mass spectra:
- Write the percentages as decimals
- Create a table showing the isotope combinations
- Any molecules that are the same, add the abundances up
- Divide all the relative abundances worked out before by the smallest value.
This will give you a whole number ratio which can be used to predict your spectra
How can mass spectrometry be used to determine the relative molecular mass of a molecule:
In a mass spectrum for molecules the peaks show fragments of the original molecule
The last peak is the m+1 peak or the molecular ion peak
This is the same as the relative molecular mass of the molecule
What is ionisation energy?
Ionisation energy is the amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state
Successive ionisation energy
The energy required to remove than 1 electron (one by one) from the same atom is called successive ionisation
How does the number of protons affect ionisation energies?
Size of the positive nuclear charge:
As the nuclear charge increases, its attraction for the outermost electron increases and more energy is required to remove an electron.
This means that the ionisation energy increases.
How does the size of the atom affect ionisation energies?
Size of atom (distance of outermost electron from the nucleus):
As atomic size increases, the attraction of the positive nucleus for the negative electron decreases and less energy is required to remove an electron.
This means that the ionisation energy decreases.
How does shielding affect ionisation energies?
Shielding (screening) effect of inner shell electrons:
The outermost electron is screened (shielded) from the attraction of the nucleus by the repelling effect of the inner electrons.
As shielding increases, the attraction of the positive nucleus for the negative electron decreases and less energy is required to remove an electron.
This means that the ionisation energy decreases.
How does the fact that an electron is on its own in an orbital or paired with another electron (Spin to spin repulsion) affect ionisation energies?
Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.
How does the electron sub-shell from which the electron is removed affect the ionisation energy?
If an electron already has a high energy, then the energy it needs to gain in order to be removed will not be very large.
If, however, the electron is in an orbital of a low energy quantum shell, for example the 1s orbital, then it will need to gain considerably more energy to be removed
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.
Trend in 1st ionisation energy across period
The 1st ionisation energy increases across a period
Nuclear charge increases as you go across a period and there is no change in the amount of shielding or the distance of the outermost electron from the nucleus
This leads to the attraction between the outer electron and the nucleus increasing
Therefore, more energy is required to remove an electron ie; a higher ionisation energy is needed
Trends in ionisation energy across period 3:
In period 3 there is a general increase in the ionisation energies but there is a small drop from Mg to Al and from P to S
Why is there a 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 shielded by the 3s electrons
Why is there a drop from P to S?
With sulphur there are 4 electrons in the 3p sub shell and the 4th is starting to doubly fill 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
Trend in first ionisation energy down a group:
1st ionisation energy decreases down a group
As you go down a group, the nuclear charge increases. However, the amount of shielding and the distance from the nucleus also increases
This lowers the effective nuclear charge on the nucleus, leading to a decrease in the attraction between the nucleus and the outermost electron so less energy is required to remove electron from outermost shell
Hence ionisation energy down the group decreases
Understand how ideas about electronic configuration developed from:
I. The fact that atomic emission spectra provide evidence for the existence of quantum shells;
When atoms in the gaseous state are given energy, the electrons move to higher energy levels.
Eventually they return back to the lower energy levels and emit electromagnetic radiation as they do so
The radiation emitted will have a fixed frequency as the energy of the shells are fixed
The fact that only certain frequencies of electromagnetic radiation are emitted, rather than a continuous spectrum is evidence that the energy of electrons in atoms can only have certain, fixed values and not a continuous range of values
Understand how ideas about electronic configuration developed from:
II. The fact that successive ionisation energies provide evidence for the existence of quantum shells and the group to which the element belongs:
The patterns in successive ionisation energies for an element give us important information about the electronic structure for that element.
Successive ionisation energies increase as more and more electrons are removed.
Large jumps in the ionisation energy reveal where electrons are being removed from the next principle energy level.
How many electrons can fit in each quantum shells?
Quantam shell 1: 1s subshell and 2 electrons
Quantam shell 2: , 2s, 2p subshells and 8 electrons
Quantam shell 3: 3s, 3p,3d subshells and 18 electrons
Quantam shell 4: 4s,4p,4f,4d subshells and 32 electrons
Orbitals
Each type of sub-level contains one or more orbitals.
An orbital is the region where the electrons are most likely to be found.
Each orbital can hold up to 2 electrons, with opposite spins.
Each subshell has a different type of orbital, an s subshell has an s-orbital, a p subshell p-orbital, etc.
Shape of S subshell
A s-orbital has a spherical shape.
Shape of P subshell
A p-orbital has a 3-dimensional dumb-bell shape.
There are three p-orbitals, px, py, and pz at right angles to one another.
How many electrons occupy s, p and d-subshells:
The subshells s, p, d, and f contain the following number of orbitals respectively, where every orbital can hold up to two electrons maximum:
s: 1 orbital, 2 electrons.
p: 3 orbitals, 6 electrons.
d: 5 orbitals, 10 electrons
Pauli Exclusion Principle
Pauli Exclusion Principle: states that a maximum of two electrons can occupy a single atomic orbital and only if the electrons have opposite spins
Hund’s Rule:
Hund’s Rule: electrons occupy orbitals singly before pairing up
The Aufbau principle
The Aufbau principle states that electron’s occupy orbitals in order of increasing energy.
The order of occupation is as follows:
1s<2s<2p<3s<3p<4s<3d
How to write electron configuration of ions?
When writing the electron configuration of ions, it is important to add or subtract the appropriate number of electrons.
For negative ions add electrons.
For positive ions remove electrons.
Electronic configuration of chromium and copper:
The electron configurations of chromium and copper are exceptions to the normal rules of orbital filling:
Chromium has the electronic configuration: 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^1, 3d^5
Copper has the electronic configuration: 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^1, 3d^10
In each case the 4s orbital contains one electron.
This is because the 4s and 3d sub-levels lie very close together in energy, and the 3d being either half full or completely full is a lower and more stable energy arrangement.
How does the electronic configuration determines the chemical properties of an element:
All the elements in a main group contain the same outer electron configuration.
The elements within a group have similar chemical properties because they have the same outer electronic configuration
Periodicity
Periodicity is the general trends that can be seen in the periodic table.
Periodicity:
Atomic radius
Atomic radii decrease as you move from left to right across a period.
This is because number of protons increases, so does the nuclear charge.
This results in an increase in the attractive force between the nucleus and the outer electrons
This increase in attractive force offsets the increase in electron – electron repulsion as the number of electrons in the outer quantum shell increases
Periodicity:
Ionisation energy
There is a general trend across is to increase.
This is due to increasing number of protons as the electrons are being added to the same shell
Exception of electron configuration of ions
The exception is transition metals
4s IN before 3d and 4s OUT before 3d
Trends in successive ionisation energies
There is a general increase in successive ionisation energy due to the increasing effective charge in the nucleus ( as electrons are removed)
Large ‘jumps’ are caused by electrons being taken from an energy level that is closer to the nucleus (ie, decreased distance and shielding, so stronger attraction)
Pattern - In group 1 - large jump between successive IE 1 & 2
In group 2 - large jump between successive IE 2 & 3 etc
Features of ionisation energy graph
Group 1 has the lowest
Group 8 has highest
Large drops due to new energy level
Trend in bonding of period 3 elements
Na - Al - metallic bonding
Silicon - giant covalent
P - Cl - simple covalent
Ar - monoatomic
Trend in melting point of period 3 elements
The first three elements are all metallic. There is an increase in the melting points from Na to Al ,due to the increase in ionic charge, which results in a stronger electrostatic attraction with delocalised electrons. Therefore, more energy is needed to break these forces of attraction
Silicon has a very high melting point as is has strong covalent bonds that require a lot of energy to break
The last four elements (Si to Ar) is linker to the mr of the molecule. The higher the mr, the greater the strength of the London forces, so more energy is required and there is a higher melting point.
