* Topic 1 - Atomic structure and the periodic table

Question 1

what is the position, relative mass and relative charge of a proton

Answer 1

Position - nucleus
Relative mass - 1
Relative charge - +1

Question 2

What is the position, relative mass and relative charge of a neutron

Answer 2

Position - Nucleus
Relative mass - 1
Relative charge - 0

Question 3

What is the position, relative mass and relative charge of an electron

Answer 3

Position - Orbitals
Relative mass - 1/1840
Relative charge - -1

Question 4

What is the atomic number

Answer 4

The atomic number, Z, is the number of protons in the nucleus.

Question 5

What is the mass number

Answer 5

The mass number ,A, is the total number of protons and neutrons in the atom.

Question 6

What is an isotope

Answer 6

Isotopes are atoms with the same number of protons, but different numbers of neutrons

Question 7

Define relative isotopic mss

Answer 7

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

Question 8

Comment on the similarity in chemical and physical properties of isotopes

Answer 8

Isotopes have similar chemical properties because they have the same electronic structure. They may have slightly varying physical properties because they have different masses.

Question 9

Define relative atomic mass

Answer 9

Relative atomic mass is the average mass of one atom compared to one twelfth of the mass of one atom of carbon-12

Question 10

Define relative molecular mass

Answer 10

Relative molecular mass is the average mass of a molecule compared to one twelfth of the mass of one atom of carbon-12

Question 11

What can a mass spectrometer be used to identify

Answer 11

The mass spectrometer can be used to determine all the isotopes present in a sample of an element and to therefore identify elements.

Question 12

The relative atomic mass quoted on the periodic table is a weighted average of all the (1)

Answer 12

1 - isotopes

Question 13

For each isotope, the mass spectrometer can measure… (1) and (2)

Answer 13

a m/z (mass/charge ratio) and an abundance

Question 14

If asked to give the species for a peak in a mass spectrum, what do you give?

Answer 14

the charge and mass number eg. ^24Mg^+

Question 15

What is the equation for R.A.M

Answer 15

Σ(isotopic mass x % abundance) / 100

Question 16

Sometimes two electrons may be removed from a particle forming a (1) ions

Answer 16

1 - 2+

Question 17

^24Mg^2+ with a 2+ charge would have a m/z of (1)

Answer 17

1 - 12

Question 18

How do you measure the Mr of a molecule from a mass spectrometer

Answer 18

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.

Question 19

What is the peak from which the Mr value is calculated called

Answer 19

parent ion or molecular ion

Question 20

Give 4 uses of mass spectrometers

Answer 20

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

Question 21

Define first ionisation energy

Answer 21

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

Question 22

What is the equation for first ionisation energies

Answer 22

H(g) -> H^+(g) + e^-

Question 23

Define second ionisation energy

Answer 23

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

Question 24

What are 3 factors factors that affect ionisation energy

Answer 24

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
   weaker the attraction to the nucleus)
3. 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)

Question 25

What does patterns in successive ionisation energies of an element give us important information about

Answer 25

the electronic structure for that element

Question 26

Why are successive energies always larger

Answer 26

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.

Question 27

How are ionisation energies linked to electronic structure (eg explain a big jump between the 4th and 5th ionisation energy diagram)

Answer 27

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

Question 28

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

Answer 28

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

Question 29

What is periodicity in a ionisation energy diagram

Answer 29

The shape of the graph for periods two and three is similar. A repeating pattern across a period is called periodicity.

Question 30

What does the pattern in the first ionisation energy give us useful information about

Answer 30

electronic structure

Question 31

Why has helium the largest first ionisation energy?

Answer 31

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

Question 32

Why do first ionisation energies decrease down a group?

Answer 32

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

Question 33

Why is there a general increase in first ionisation energy across a period?

Answer 33

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.

Question 34

Why has Na a much lower first ionisation energy than Neon?

Answer 34

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.

Question 35

*Why is there a small drop from Mg to Al?

Answer 35

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

Question 36

*Why is there a small drop from P to S?

Answer 36

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.

Question 37

What is the Bohr electron model

Answer 37

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.

Question 38

Describe the A level electron model

Answer 38

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

Question 39

What are the sub levels in the principle levels 1,2,3 and 4

Answer 39

1 - 1s 2 - 2s,2p 3 - 3s,3p,3d 4 - 4s,4p,4d,4d

Question 40

What does the shape of orbitals represent

Answer 40

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.

Question 41

Why does the 3d shell get filled after the 4s shell

Answer 41

3d is higher in energy than 4s and so gets filled after the 4s

Question 42

What shape are s sub levels

Answer 42

spherical

Question 43

What shape are p sub levels

Answer 43

dumbbell shaped

Question 44

Describe how the blocks represent the outer shells in an atom

Answer 44

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

Question 45

What is periodicity

Answer 45

Periodicity is the repeating pattern of physical or chemical properties going across the periods

Question 46

What happens to atomic radii across a period

Answer 46

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

Question 47

Explain the trend in 1st ionisation energy across a period

Answer 47

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!

Question 48

Explain the melting and boiling points in Na, Mg and Al

Answer 48

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.

Question 49

Explain the melting and boiling points in Si

Answer 49

Macromolecular: many strong covalent bonds between atoms high energy needed to break covalent bonds– very high mp + bp

Question 50

Explain the melting and boiling points in Cl2, S8, P4

Answer 50

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

Question 51

Explain the melting and boiling points in Ar

Answer 51

monoatomic: weak London Forces between atoms

Question 52

How do the explanations in boiling and melting points differ from period 1 to period 2

Answer 52

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)