Topic 1: Atomic structure and the periodic table Flashcards
What is the atomic number
(what can it also be called?)
this is the number of protons that are contained within the nucleus of an atom. (it is also called the proton number)
What is the mass number
the mass number of an element is the sum of the number of protons plus the number of neutrons, in one atom of that element
What is an isotope
Isotopes are atoms of the same atomic number but different mass numbers.
All atoms of an element have the same number of protons. Differences in mass are caused bu different numbers of neutrons
What is Relative Atomic Mass (Ar)
Relative atomic mass is the average mass of an atom of an element relative to 1/12th of the mass of an atom of the isotope carbon 12
What is relative isotopic mass
Relative isotopic mass is the average mass of one atom of an isotope relative to 1/12th of the mass of an atom of the isotope carbon 12
Lithium, in its naturally occurring compounds, has two isotopes of relative isotopic masses 6.015 and 7.016. The percentage abundance of each isotope is 7.59 and 92.41 respectively. Calculate the relative atomic mass of Lithium to 3sf
6.94
the proportion of each isotope present in a sample of an element can be measured using an instrument called a (—–)
mass spectrometer
MASS SPECTROMETRY
On a spectrum of an element, we get one peak for each individual (———)
isotope
What is the height of a peak on a mass spectrometer proportional to
The height of each peak (the area under each peak) is proportional to the number of atoms (relative abundance) of this isotope in the sample tested
What happens to sample M when it is passed through a mass spectrometer (13)
- inject into the machine
- sample heated in vacuum chamber and is vaporised
- gaseous sample of substance bombarded with high energy e-
- an e- is knocked off from the edge of an atom to create a positive ion
- these are separated according to charge
- the positive ions are accelerated by an electric field
a focused stream of ions is passed through a magnetic field - the magnetic field deflects the ions (high molecular weight ions are deflected too little and hit the wall; low molecular weight ions are deflected too much and hit the wall)
- the deflection depends on mass and charge of the ion (m/z ratio)
- heavier ions are not deflected much
- ions with a higher charge are deflected more
- a detector measures the deflection of ions
- mass and charge put together into a mass/charge ratio
draw the mass spectrometer diagram of a chlorine molecules
3:1 at Cl-35 and Cl-37; Chlorine is diatomic so some molecules will fragment upon ionisation so there will also be a 9:6:1 ratio at Cl-70, Cl-72 and Cl-74
Boron has isotopic masses of Br-79 and Br-81. Identify the particles responsible for the peaks at m/z 158, 160 and 162.
158 = Br-79 and Br-79
160 = Br-79 and Br-81, Br-81 and Br-79
162 = Br-81 and Br-81
Boron has isotopic masses of Br-79 and Br-81. Work out the relative heights at m/z 158, 160, 162
158 = 1/2 x 1/2 = 1/4
160 = (1/2 x 1/2) + (1/2 x 1/2) = 1/2
162 = 1/2 x 1/2 = 1/4
Therefore, the relative heights at 158:160:162 = 1:2:1
How do you determine relative molecular mass of a polyatomic molecule and what is the extra line on the mass spectrometer from?
For a polyatomic molecule you will be asked to work out the relative molecular mass of a compound by considering what is called the molecular ion peak.
You have to be careful when analysing organic compounds. This is because there is always a small percentage of the carbon-13 isotope present in the compound, which can lead to what is often referred to as an M + 1 peak. This peak can often be seen in molecules with large masses, where the percentage of carbon-13 becomes significant. The peak is often missing, or insignificant, in molecules of small mass.
Why are we able to analyse organic molecules? (5)
- organic molecules do not hold charge well
- the ions break up or fragment
- they break into smaller ions and neutral fragments
- they do this in a regular pattern
- i.e. the same molecules always fragments in the same way: this allows us to analyse them
Which bonds are most likely to break?
weakest bonds which result in more stable fragments being produced - ions with a positive charge on a tertiary carbon atom are more stable than ions with a charge on a primary carbon atom
Predict the fragments of pentane which produced the peaks at the following m/z values:
29
43
57
72
29 - (C2H5)+
43 - (C3H7)+
57 - (C4H9)+
72 - (C5H12)+
*(they are all gases)
Describe the difference between relative atomic mass and relative isotopic mass.
The relative atomic mass is the average mass of an atom of an element compared to 1/12th of the mass of an atom of carbon-12. The relative isotopic mass, however, is the mass of an individual atom of a particular isotope relative to 1/12th of the mass of an atom of carbon-12.
Explain why the relative atomic masses of many elements are not exact whole numbers.
Relative atomic masses of many elements are not exact whole numbers due to the slightly different exact masses of protons and neutrons.
Calculate the relative atomic mass of a sample of magnesium that has the following isotopic composition:
* magnesium-24: 78.6%
* magnesium-25: 10.1%
* magnesium-26: 11.3%
Give your answer to three significant figures.
Relative atomic mass = ((78.6 x 24) + (10.1 x 25) + (11.3 x 26))/100
= 24.327 (to 3 s.f. = 24.3)
Calculate the relative atomic mass of a sample of magnesium that has the following isotopic composition:
* magnesium-24: 78.6%
* magnesium-25: 10.1%
* magnesium-26: 11.3%
Give your answer to three significant figures.
Relative atomic mass = ((78.6 x 24) + (10.1 x 25) + (11.3 x 26))/100
= 24.327 (to 3 s.f. = 24.3)
Calculate the relative atomic mass of neon given that it has 3 isotopes neon-20, neon-21 and neon-22. The percentage abundances are 90.92, 0.26 and 8.82 respectively.
Relative atomic mass =
(90.92 X 20) + (0.26 X 21) + (8.82 X 22) = 2017.9
2017.9/100 = 20.179 (to 3 s.f = 20.2)
How are quantum energy shells labelled?
- energy levels are numbered starting from those closest to the nucleus
- the first quantum shell is found closest to the nucleus, and is the lowest energy
- the second quantum shell is found further away from the nucleus
How many sub shells are there per quantum energy level
Quantum shell one only has one sub-shell = 1s
Quantum shell two has two sub-shells = 2s and 2p
Quantum shell three has three sub-shells = 3s, 3p and 3d
Quantum shell four has three sub-shells = 4s, 4p, 4d and 4f
Each sub shell has a slightly different (—–)
energy level
How many electrons can one orbital hold
2
how many orbitals does each sub shell contain
‘s’ sub-shells contains 1 orbital = 2 electrons
‘p’ sub-shells contains 3 orbital = 6 electrons
‘d’ sub-shells contains 5 orbital = 10 electrons
‘f’ sub-shells contains 7 orbital = 14 electrons
shape of s orbitals
- spherical
- one occur in every principal energy level
shape of p orbitals
- dumb bell shaped
- three occur in energy levels except the first
shape of d orbitals
- various shapes (look it up) dependent on axis
- five occur in energy levels expect the first and second
What is the main exception to the rule in filling sub shells
you fill 4s in before 3d unless 4s can be half filled or fully filled in which case it is filled first as this will be more stable (i.e. exceptions are copper (4s1 3d10) and chromium (4s1 3d5)
what is the order of filling orbitals
Orbitals are not filled in numerical order because the principal energy levels get closer together as you get further from the nucleus. This results in overlap of sub levels. The first example occurs when the 4s orbital is filled before the 3d orbitals.
What is Pauli Exclusion Principle
An orbital can hold only 0, 1 or 2 electrons only, and if two electrons are present then their spins must be exactly opposite
What si Hund’s rule of maximum multiplicity
When filling sub levels other than s, electrons are placed in individual orbitals before being paired up
why does the next electron not pair up before filling each p orbital individually first (in Carbon for example)?
As this would give rise to repulsion between the similarly charged species. Instead, it goes into another p orbital which means less repulsion, lower energy and more stability
(in Nitrogen for example) Following Hund’s rule, the next electron will not pair up so goes into a vacant p orbital. All three electrons are now unpaired. This gives (——), (——) and therefore (——-)
less repulsion, ; lower energy ; more stability
why does the 4s orbital get filled before the 3d orbitals?
This is because the principal energy levels get closer together as you go further from the nucleus coupled with the splitting into sub energy levels, the 4s orbital is of a LOWER ENERGY than the 3d orbitals so gets filled first.
Why is chromium an exception
its configuration is 4s1, 3d5 when it should be 4s2, 3d4
This is because to achieve a more stable arrangement of lower energy, one fo the 4s electrons is promoted into the 3d to give 6 unpaired electrons with lower repulsion
What does isoelectronic mean and give an example of 3 isoelectronic ions
Two or more atoms or ions are isoelectronic if they have the same electron configuration
for example,
Na+
Mg2+
Al3+
What are positive ions called and how are they formed
Positive ions are called cations and are formed by removing electrons from atoms
what are negative ions called and how are they formed
Negative ions (anions) are formed by adding electrons to atoms
From which orbitals are electrons first removed (excluding transition metals)
Electrons are removed first from the highest occupied orbitals
Are electrons first lost from the 3d block or 4s block
When filling orbitals 4s has a lower energy than 3d and so it is filled first.
However, 4s still behaves as the outermost highest energy orbitals and so 4s electrons are always lost first before 3d
what is the first ionisation energy
the first ionisation entry of an element is the energy required to remove an electron from each atom in one mole of atoms in the gaseous state
we can represent this as:
M(g) -> M+(g) + e-
why is the first ionisation energy endothermic
energy needs ti be put in to overcome the attraction between the electron and nucleus
Show the first three ionisation energies for an element, M
M(g) -> M+(g) + e-
M+(g) -> M2+(g) + e-
M2+(g) -> M3+(g) + e-
if the first three ionisation energies of aluminium are as follows:
1st = 578kJ/mol
2nd = 1817kJ/mol
3rd = 2745kJ/mol
Work ou the total value of the first three ionisation energies of aluminium
Em1 + Em2 + Em3 = 5140kJ/mol
State one piece of evidence for the existence of quantum shells (successive ionisation energies)
When successive energies of an element are listed there are steady increases, and big jumps occur at defined places (quantum shells)
What does a big jump in successive ionisation energies suggest and what does it allow you to deduce?
A big jump between successive ionisation energies marks the removal of an electron rom the next shell at a lower energy level (closer to the nucleus).
This should allow you to deduce which group the element is in.
Look at how many electrons are removed before the big jump, this tell you how many electrons are in the other shell and so the group.
Work out what group this element is from based off the following successive ionisation energies:
789
1577
3232
4356
16091
19785
- 4 electrons are ion the outermost shell
- therefore this element must come from group 4
What three factors affect ionisation energy
- nuclear charge
- distance from the nucleus
- shielding from filled inner electron shells
How does nuclear charge affect ionisation energy
As the nuclear charge increases, the electrostatic force of attraction between the nucleus and electrons increases.
Therefore more energy is required to remove each electron.
Ionisation energy increases.
How does distance from the nucleus affect ionisation energy
As the distance from the nucleus increases the electrostatic force of attraction between the nucleus and outer electrons decrease.
Therefore less energy is required to remove electrons.
So ionisation energy decreases.
How does shielding from filled inner electron shells affect ionisation energy
Inner shells filled with electrons will shield outer shells from the nucleus.
This shielding increases with increasing number of inner shells.
Shielding reduces the forces of attraction between the nucleus and outer electron.
Therefore less energy is required to remove an electron and the ionisation energy decreases.
What is the trend in ionisation energy down a group and why
down a group, the first ionisation energy decreases
- Although nuclear charge is increasing
- Distance from the nucleus and shielding from filled inner electron shells increases therefore overall the forces of attraction between the nucleus and outer electrons decreases.
- Less energy is required to removed an electron from the outer shell.
What happens to atomic radius across a period
atomic radius decreases
What happens to the first ionisation energy across a period
Across the period – the atomic radius decreases
As nuclear charge increases.
Shielding from filled inner shells stays to the same as the number of inner shells does not increase.
This leads to a general increase in first ionisation energy as you go from left to right.
The first four ionisation energies in kJ mol-1 of calcium are 590, 1150, 4940, and 6480. Explain why the second ionisation energy of calcium is larger than the first
The second ionisation energy is larger as the electron is being removed from a positive ion, so there is more attraction between the outer electron and nucleus
The first four ionisation energies in kJ mol-1 of calcium are 590, 1150, 4940, and 6480. Explain why the third ionisation energy is much larger than the second
- Calcium is in group 2
- The third ionisation energy removes and electron 3p. This orbital is closer to the nucleus and also has less shielding than 4s, which the electrons were previously removed from.
- This results in there being greater attraction between the electrons and the nucleus and so more energy is needed to overcome this attraction resulting in a greater third ionisation energy.
Nuclear Charge
The nucleus contains 2 particles protons and (1). Protons have a (2) charge and neutrons have no charge; therefore overall the charge on an nucleus is (3). Thus, the more protons within the nucleus the (4) the nuclear charge.
1 - neutrons
2 - positive
3 - positive
4 - greater
Distance from the nucleus
Ionisation energy is the energy required to remove an electron. Electrons are (1) charged and are therefore attracted to a positive nucleus. This attraction falls off rapidly with distance.
An electron close to the nucleus will be (2) strongly attracted to the nucleus than one further away.
1 - negatively
2 - more
If there was no attraction to the nucleus what would happen to the outer electron in sodium and lithium?
Electrons would be very easily removed
Why is there a big drop in first ionisation energy from Neon to Argon (look at a diagram)
- Ne =2,8Na=2,8,1
- Sodium’s outer electron is in the 3rd energy level and is repelled by to full shells of electrons
between it and the nucleus - Neon’s outer electron is repelled by only one full shell of electrons
- The extra repulsion in sodium outweighs the increased nuclear charge
- Therefore less energy is needed to remove the outer electron in sodium
Why is there a slight drop from magnesium to aluminium in terms of first ionisation energies?
- Aluminium’s outer electron is in a 3p orbital which has a slightly higher energy than a 3s orbital
- Less energy is needed to remove the outer electron in aluminium
Why is there a slight drop from phosphorus to sulfur in terms of first ionisation energy?
- In sulphur, the extra electron goes into a p orbital which already contains one electron
- The 2 spin paired electrons repel each other slightly
- This repulsion slightly outweighs the increased nuclear charge
- Less energy is needed to remove the outer electron from sulphur
What is the atomic radius
The atomic radius of an element is a measure of the size of its atoms. It is the distance from the centre of the nucleus to the boundary of the electron cloud
What type of bonding can cause ‘bonded atoms’. Where the atoms are pulled together and so the measured radius is less than if they are just touching.
Metallic bonding or covalent bonding (what you would get if you had metal atoms in a metallic structure, or atoms covalently bonded to each other. The type of atomic radius being measured here is called the metallic radius or the covalent radius depending on the bonding.)
What happens to the atomic radius if the atoms are just touching vs metallic or covalent radii? What is this type of radius called?
The attractive forces are much less, and the atoms are essentially “unsquashed”. This measure of atomic radius is called the van der Waals radius after the weak attractions present in this situation.
