Atomic structure and the periodic table

Question 1

What is the relative isotopic mass?

Answer 1

the mass of an isotope of an atom relative to the mass of 1/12 the atom of carbon-12

Question 2

What is the relative atomic mass?

Answer 2

The average mass of an atom relative to the mass of 1/12 of an atom of carbon-12

Question 3

what is the principal quantum number (n) ?

Answer 3

The number given to each shell e.g. the third shell has n=3

Question 4

What happens when we have 2 electrons in the same orbital?

Answer 4

They must have opposite spins

Question 4

What is an atomic orbital?

Answer 4

A region around the nucleus that can hold up to 2 electrons with opposite spins

Question 5

What are the different types of atomic orbitals?

Answer 5

S, P, D, F

Question 6

Define details of S orbitals

Answer 6

• It has a spherical shape and every electron shell contains a single s orbital
• The size of the orbitals increase with increasing shell number

Question 7

Define details of p orbitals

Answer 7

• There are 3 p orbitals
• They are shaped like a dumb bell
• They are orientated perpendicular to one another
• With increasing shell number, the lobes become longer and larger

Question 8

How many d orbitals are there?

Answer 8

5

Question 9

How many f orbitals are there?

Answer 9

7

Question 10

How are electrons imagined?

Answer 10

As small spinning charges which rotate around their own axis in either clockwise or anticlockwise direction. The spin creates a tiny magnetic field with each pole facing up or down

Question 11

What happens to electrons with the same spin?

Answer 11

they repel each other, known as spin-spin repulsion

Question 12

What do sub shells contain?

Answer 12

one or more atomic orbitals

Question 13

How much space does each electron occupy?

Answer 13

Even though there is repulsion between them, they occupy the same region of space in orbitals

Question 14

What is the order the electrons fill the orbitals in?

Answer 14

1s - 2, 2s - 2, 2p - 6, 3s - 2, 3p - 6, 4s - 2, 3d - 10, 4p - 6

Question 15

Why does 4s come before 3d in the orbital order?

Answer 15

It has a lower energy level than 3d

Question 16

Which elements do not have the expected electron structure and why?

Answer 16

Chromium and Copper
To decrease the electron repulsion and get a lower overall energy level, the electrons occupy the 3d orbital first. (for copper its 3d10, 4s1 and not 3d9 and 4s2)

Question 17

What is ionisation energy?

Answer 17

A measure of the amount of energy needed to remove electrons from atoms

Question 18

What is the first ionisation energy?

Answer 18

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous positive ions

Question 19

Give an example of the first ionisation energy?

Answer 19

Na(g) = Na+(g) +e-

Question 20

What does the value of the 1st I.E. depend on?

Answer 20

The electronic structure

Question 21

Why is the ionisation energy value for helium higher than for hydrogen?

Answer 21

• There are 2 protons in the nucleus instead of 1, meaning that the nuclear charge is greater
  -The distance of the outer electron to the nucleus decreases (decreased atomic radius)
• The outer electron is held more tightly to the nucleus so more energy is required to pull it out of the atom

Question 22

Why is the Ionisation energy for lithium lower than for helium?

Answer 22

• Although it has 3 protons instead of 2, the valence electron is in another shell
• Filled inner shells exert a shielding effect which lowers the effective nuclear pull
• It is further away from the nucleus so the electron has a lower nuclear attraction and less energy is required to pull it away

Question 23

Why does the 1st ionisation energy show a general increase across a given period?

Answer 23

There is an increased nuclear charge, but no extra electron shielding so the distance from the nucleus stays the same but there is a greater nuclear charge

Question 24

Why does the ionisation energy decrease down a group?

Answer 24

- All though there is an increase in nuclear charge, there is an increase in shells so the distance of the outer electron to the nucleus increases - Shielding effect increases so the attraction of valence electrons to the nucleus decreases (increased atomic radius) - The outer electron is held more loosely to the nucleus so becomes easier to remove

Question 25

Why is the ionisation energy for oxygen lower than for nitrogen?

Answer 25

All though there is no new electron shell here, the extra electron has paired up with one of the electrons already in one of the 2p orbitals. The repulsive force between the 2 paired-up electrons means that less energy is required to remove one of them

Question 26

What are successive ionisation energies of an element?

Answer 26

More than one electron can be removed from an atom and each time you remove an electron there is a successive ionisation energy

Question 27

What is the second ionisation energy?

Answer 27

The energy required to remove 1 mole of electrons from 1 mole of gaseous 1+ positive ions to form one mole of 2+ positive ions

Question 28

What is an example of the second ionisation energy?

Answer 28

Na+(g) = Na2+(g) + e-

Question 29

Why are successive ionisation energies always greater than the previous one?

Answer 29

The electron is being pulled away from a more positive species

Question 30

When do large increase occur?

Answer 30

When there is a change of shell as there is a big decrease in shielding

Question 31

What can large increases be used for?

Answer 31

To predict the group of an unknown element

Question 32

What happens wherever there has been a large increase in ionisation energy?

Answer 32

There has been a change in energy level from which the electron has been removed

Question 33

Why does the ionisation energy across a period increase?

Answer 33

- The nuclear charge increases - The atomic radius of the atoms decrease so outer shell is pulled closer to nucleus - Shielding remains constant - Becomes harder to remove an electron as you move across so more energy is needed

Question 34

What do elements across a period show?

Answer 34

Repeating patterns in physical and chemical properties

Question 35

Why is there a decreased atomic radius across a period?

Answer 35

There is an increased nuclear charge, so a higher pull from the nucleus

Question 36

Why is there a small drop in the 1st ionisation energy between Mg and Al?

Answer 36

Mg has outer electrons in the 3s sub shell whereas Al starts to fill the 3p sub shell. Al's electron is slightly easier to move as 3p electrons are higher in energy

Question 37

What is electronegativity?

Answer 37

The power of an atom to attract the 2 electrons in a covalent bond

Question 38

What is the trend of electronegativity across a period?

Answer 38

Across a period, there are more protons in the nucleus so a smaller atomic radius so there is a stronger attraction between the nucleus and 2 electrons in a covalent bond

Question 39

How is melting and boiling points affected by the size of an atom in metallic bonding?

Answer 39

- The bigger the atom, the more electrons there are, so the greater the forces - The greater the number of electrons holding them together, the greater the B.p and M.p

Question 40

What also makes the bonding stronger in metallic bonding?

Answer 40

A smaller sized ion with a greater positive charge

Question 41

Do metallic bonds need higher or lower energy in general?

Answer 41

Higher

Question 42

What molecular term is used to describe silicone?

Answer 42

Macromolecular

Question 43

Why does silicone have a very high boiling point and melting point?

Answer 43

It has many strong covalent bonds between atoms so a high energy is needed to break them

Question 44

What are the examples of simple molecular?

Answer 44

cl2(g), s8(s), P4(s) - (this is an example for one period but it counts for all the other periods as well)

Question 45

What are the forces between the simple molecular?

Answer 45

There are weak London forces between molecules so little energy is needed to break them - it has a low melting and boiling point

Question 46

Why does S8 have a higher melting point than P4?

Answer 46

It has more electrons (s8=128, p4=60) so it has stronger London forces between molecules

Question 47

What is Argon classified as?

Answer 47

Monoatomic - there are weak London forces between the atoms