1 - Atomic structure and the Periodic Table

Question 1

atomic number

Answer 1

the number of protons in the nucleus of an atom

Question 2

mass number

Answer 2

the sum of the number of protons and the number of neutrons in the nucleus of an atom.

Question 3

fact about electrons and protons in an atom

Answer 3

the number of electrons in an atom is equal to the number of protons in the nucleus because an atom is neutral

Question 4

isotope

Answer 4

Isotopes are atoms of the same element that contain the same number of protons and electrons, but a different number of neutrons.

Question 5

ions

Answer 5

Ions are charged particles formed when atoms lose or gain electrons to gain stability.

Question 6

Relative atomic mass

Answer 6

the weighted average mass of an atom of an element compared with 1/12th the mass of the atom of carbon-12.

Question 7

Relative isotopic mass

Answer 7

The mass of an atom of an isotope compared to 1/12th the mass of an atom of carbon-12.

Question 8

How do you calculate relative atomic mass when given masses of isotopes?

Answer 8

• Find the total mass of all the atoms (say there are 100 atoms).
• Find he weighted average mass by dividing the total mass by 100.

Question 9

What is a mass spectrometer used for?

Answer 9

A mass spectrometer can be used to find the relative abundances of isotopes in a sample of an element.

Question 10

What are the different parts of the mass spectrometer?

Answer 10

• ionisation area
• acceleration area
• drift region
• ion detector.

Question 11

Function of the ionisation area?

Answer 11

Vapourises and ionises the sample, to form positive ions.

Question 12

Function of the acceleration area?

Answer 12

Accelerates the ions by an electric/magnetic field, towards a negatively charged plate.

Question 13

Function of the drift region?

Answer 13

Allows the ions to drift through the flight tube to become separated according to mass.

Question 14

Function of the ion detector?

Answer 14

Detects when ions arrive at the detector.

Question 15

What affects how much an ion is deflected in MS?

Answer 15

• mass of the ion (lighter means deflected more).

- charge of the ion (higher positive charge means deflected more).

Question 16

First ionisation energy

Answer 16

the energy required when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge.

Question 17

Second ionisation energy

Answer 17

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 18

What are the 3 things that affect ionisation energy?

Answer 18

• the attraction of the nucleus (more protons, greater attraction).
• distance of the electrons from the nucleus.
• shielding of the attraction of the nucleus

Question 19

Why are successive ionisation energies always larger?

Answer 19

• second ionisation energy is always bigger than the first ionisation energy.
• once the first electron is removed, the force of attraction of the nucleus on each of the remaining electrons is greater, therefore the energy required to remove the next electron is greater.

Question 20

In a graph for successive ionisation energy of an element, why are there sometimes large jumps?

Answer 20

• say there is a large jump between the 2nd and 3rd electron being removed.
• the 3rd electron is in an inner shell closer and attracted much more by the nucleus than the 2nd electron.
• This electron experiences less shielding than the 2nd due to being in an inner shell, so force of attraction is greater.
• more energy is required to overcome the force of attraction and remove the 3rd electron.

Question 21

Why does He have such a high first ionisation energy and why is it bigger than H?

Answer 21

• first electron is in the first shell closest to the nucleus.
• experiences no shielding effect.
• Larger than H as it has one more proton.

Question 22

Why do first ionisation energies decrease down a group?

Answer 22

• as you go down the group
• although proton number increases
• the number of electron shells increases
• distance from the nucleus increase
• shielding effect increases.
• force of attraction of the nucleus on the outer electrons decreases.
• first ionisation energy decreases.

Question 23

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

Answer 23

• as you go across a period
• number of protons in the nucleus increases
• number of electrons increases on the same shell.
• shielding effect is the same.
• electrons are pulled in closer to the nucleus (due to increasing force of attraction of nucleus)
• force of attraction of nucleus on the outer electrons increases.
• first ionisation energy increases.

Question 24

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

Answer 24

• Na will have its outer electron in a 3s subshell, which is further away from the nucleus.
• also experiences a greater shielding effect.
• force of attraction of nucleus on outer electrons decreases.
• easier to remove, lower ionisation energy.

Question 25

Why is there a small drop from Mg to Al?

Answer 25

• first outer electron of Al is in the 3p subshell.
• first outer electron of Mg is in the 3s subshell.
• electrons in 3p subshell are higher in energy and also slightly shielded by the 3s subshell, so they are slightly easier to remove than electrons in the 3s subshell.
• hence lower first ionisation energy.

Question 26

Why is there a small drop from P to S?

Answer 26

• In sulphur, there are 4 electrons in the 3p subshell.
• the 4th electron causes the first 3p orbital to be completely filled (pair of electrons with opposite spin).
• there is slight repulsion between the two negatively charged electrons due to the opposite spin and the fact that they are both negative.
• this makes it easier to remove the first electron from S than P.
• therefore S has a lower first ionisation energy than P.

Question 27

What can an orbital hold?

Answer 27

An orbital can hold up to two electrons with opposite spin.

Question 28

how many electrons can s subshell hold?

Answer 28

2

Question 29

how many electrons can p subshell hold?

Answer 29

6

Question 30

how many electrons can d subshell hold?

Answer 30

10

Question 31

how many electrons can f subshell hold?

Answer 31

14

Question 32

what subshells in 1st shell

Answer 32

s

Question 33

what subshells in 2nd shell

Answer 33

s, p

Question 34

what subshells in 3rd shell

Answer 34

s, p, d

Question 35

what subshells in 4th shell

Answer 35

s, p, d, f

Question 36

Why does 4s get filled in first before 3d?

Answer 36

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

Question 37

What does it mean if an element is in the s-block of the periodic table?

Answer 37

Its outer electron is filling a s sub shell.

Question 38

shape of s orbital?

Answer 38

spherical

Question 39

shape of p orbital?

Answer 39

dumbbell shaped

Question 40

Define the term periodicity?

Answer 40

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

Question 41

Classification of the s, p, d blocks?

Answer 41

Elements are classified as s, p, d block, according to which orbitals the highest energy electrons are in.

Question 42

What happens to atomic radius as you go across a period?

Answer 42

• decreases
• number of protons in nucleus increases
• number of electrons increases on the same shell
• shielding effect stays very similar
• force of attraction of nucleus on the electrons increases.
• electrons are pulled in closer towards the nucleus.

Question 43

what happens to 1st ionisation energy as you go across a period?

Answer 43

• generally tends to increase.
• due to increasing number of protons in nucleus as the electrons are being added to the same shell.
• sometimes there are a few dips.
• sometimes the outermost electron may be in a sub shell that has higher energy than the one before. (e.g electrons in 3p are higher in energy than 3s, so they are easier to remove).
• sometimes it may be because the outermost electron has filled in an orbital of a sub shell. When the outermost electron is present, there is slight repulsion between the two negatively charged electrons, making it easier to remove the first electron, hence dip in first ionisation energy.

Question 44

What happens to melting and boiling points as you go across a period (in this case period 1)

Answer 44

• Na, Mg, Al have metallic bonding. Metallic bonding is strong. The more electrons in the outer shell, the more that are released into the sea of delocalised electrons, stronger bonding. Also a smaller sized ion with a greater positive charge also makes the metallic bonding stronger. A lot of energy is required to overcome the strong metallic bonds.
• Si is a giant covalent structure. There are many covalent bonds between the Si atoms. Covalent bonds are strong and require a lot of energy to overcome, therefore high mp and bp.
• Cl2, S8, P4 are simple molecules. There are weak London forces between the molecules which require little energy to overcome, therefore low mp and bp.
• Ar (argon) is monoatomic and there are only weak London forces between the atoms, which require very little energy to overcome.

Question 45

Why does S8 have a higher boiling point than P4?

Answer 45

S8 molecules have more electrons than P4 molecules, therefore there are stronger London forces between S8 molecules than P4, which require more energy to overcome.