* Topic 1 - Atomic structure and the periodic table Flashcards
what is the position, relative mass and relative charge of a proton
Position - nucleus
Relative mass - 1
Relative charge - +1
What is the position, relative mass and relative charge of a neutron
Position - Nucleus
Relative mass - 1
Relative charge - 0
What is the position, relative mass and relative charge of an electron
Position - Orbitals
Relative mass - 1/1840
Relative charge - -1
What is the atomic number
The atomic number, Z, is the number of protons in the nucleus.
What is the mass number
The mass number ,A, is the total number of protons and neutrons in the atom.
What is an isotope
Isotopes are atoms with the same number of protons, but different numbers of neutrons
Define relative isotopic mss
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
Comment on the similarity in chemical and physical properties of isotopes
Isotopes have similar chemical properties because they have the same electronic structure. They may have slightly varying physical properties because they have different masses.
Define relative atomic mass
Relative atomic mass is the average mass of one atom compared to one twelfth of the mass of one atom of carbon-12
Define relative molecular mass
Relative molecular mass is the average mass of a molecule compared to one twelfth of the mass of one atom of carbon-12
What can a mass spectrometer be used to identify
The mass spectrometer can be used to determine all the isotopes present in a sample of an element and to therefore identify elements.
The relative atomic mass quoted on the periodic table is a weighted average of all the (1)
1 - isotopes
For each isotope, the mass spectrometer can measure… (1) and (2)
a m/z (mass/charge ratio) and an abundance
If asked to give the species for a peak in a mass spectrum, what do you give?
the charge and mass number eg. ^24Mg^+
What is the equation for R.A.M
Σ(isotopic mass x % abundance) / 100
Sometimes two electrons may be removed from a particle forming a (1) ions
1 - 2+
^24Mg^2+ with a 2+ charge would have a m/z of (1)
1 - 12
How do you measure the Mr of a molecule from a mass spectrometer
If a molecule is put through a mass spectrometer 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 equal to the Mr of the molecule.
What is the peak from which the Mr value is calculated called
parent ion or molecular ion
Give 4 uses of mass spectrometers
planetary space probes; elements on other planets can be identified - elements on other planets can have a different composition of isotopes
drug testing in sport; identify chemicals in the blood and identify breakdown products from drugs in the body
Quality control in pharmaceutical industry; and to identify molecules from sample with potential biological activity
Radioactive dating; determine age of fossils or human remains
Define first ionisation energy
The first ionisation energy is the energy required when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge
What is the equation for first ionisation energies
H(g) -> H^+(g) + e^-
Define second ionisation energy
The second ionisation energy is the energy required when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge
What are 3 factors 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)
- The distance of the electrons from the nucleus
(The bigger the atom the further the outer electrons are from the nucleus and the
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)
What does patterns in successive ionisation energies of an element give us important information about
the electronic structure for that element
Why are successive 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 (eg explain a big jump between the 4th and 5th ionisation energy diagram)
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
Explain what group must the below element be in based off of its ionisation energies
1st - 590
2nd - 1150
3rd - 4940
4th - 6480
5th - 8120
there is a big jump between the 2nd and 3rd ionisations energies which means that this element must be in group 2 of the periodic table as the 3rd electron is removed from an electron shell closer to the nucleus with less shielding and so has a larger ionisation energy
What is periodicity in a ionisation energy diagram
The shape of the graph for periods two and three is similar. A repeating pattern across a period is called periodicity.
What does the pattern in the first ionisation energy give us useful information about
electronic structure
Why has helium the largest first ionisation energy?
Its 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 number of protons increases making the effective attraction of the nucleus greater. The electrons are being added to the same shell which has the same shielding effect and the electrons are pulled in closer to the nucleus.
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 shielded by the 3s electrons
*Why is there a small 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.
What is the Bohr electron model
An early model of the atom was the Bohr model (GCSE model) (2 electrons in first shell, 8 in second etc.) with electrons in spherical orbits. Early models of atomic structure predicted that atoms and ions with noble gas electron arrangements should be stable.
Describe the A level electron model
Principle energy levels numbered 1,2,3,4..
1 is closest to nucleus
Split into
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
Split into
Orbitals which hold up to 2 electrons of opposite spin
What are the sub levels in the principle levels 1,2,3 and 4
1 - 1s
2 - 2s,2p
3 - 3s,3p,3d
4 - 4s,4p,4d,4d
What does the shape of orbitals represent
Orbitals represent the mathematical probabilities of finding an electron at any point within certain spatial distributions around the nucleus.
Each orbital has its own approximate, three dimensional shape.
It is not possible to draw the shape of orbitals precisely.
Why does the 3d shell get filled after the 4s shell
3d is higher in energy than 4s and so gets filled after the 4s
What shape are s sub levels
spherical
What shape are p sub levels
dumbbell shaped
Describe how the blocks represent the outer shells in an atom
first two groups are s shell
transition metals are p shell
p block is group 3-8
f block is the two rows below the periodic table
What is periodicity
Periodicity is the repeating pattern of physical or chemical properties going across the periods
What happens to atomic radii across a period
Atomic radii decrease as you move from left to right across a period, because the increased number of protons create more positive charge attraction for electrons which are in the same shell with similar shielding.
Exactly the same trend in period 2
Explain the trend in 1st ionisation energy across a period
The general trend across is to increase. This is due to increasing number of protons as the electrons are being added to the same shell
There is a small drop between Mg + Al. Mg has its outer electrons in the 3s sub shell, whereas Al is starting to fill the 3p subshell. Al’s electron is slightly easier to remove because the 3p electrons are higher in energy.
There is a small drop between phosphorous and sulfur. Sulfur’s outer electron is being paired up with an another electron in the same 3p orbital.
When the second electron is added to an orbital there is a slight repulsion between the two negatively charged electrons which makes the second electron easier to remove.
Exactly the same trend in period 2 with drops between Be & B and N to O for same reasons- make sure change 3s and 3p to 2s and 2p in explanation!
Explain the melting and boiling points in Na, Mg and Al
Metallic bonding : strong bonding – gets stronger the more electrons there are in the outer shell that are released to the sea of electrons. A smaller sized ion with a greater positive charge also makes the bonding stronger. Higher energy is needed to break metallic bonds.
Explain the melting and boiling points in Si
Macromolecular: many strong covalent bonds between atoms high energy needed to break covalent bonds– very high mp + bp
Explain the melting and boiling points in Cl2, S8, P4
simple molecular : weak London forces between molecules, so little energy is needed to break them – low mp + bp
S8 has a higher mp than P4 because it has more electrons (S8 =128)(P4=60) so has stronger London forces between molecules
Explain the melting and boiling points in Ar
monoatomic: weak London Forces between atoms
How do the explanations in boiling and melting points differ from period 1 to period 2
Similar trend in period 2
Li,Be metallic bonding (high mp) B,C macromolecular (very high mp) N2,O2 molecular (gases! Low mp as small London Forces)
Ne monoatomic gas (very low mp)
