1. Atomic structure and the Periodic Table

Question 1

atomic number

Answer 1

proton number
all isotopes have the same number of protons
equal to no of electrons if uncharged

Question 2

mass number

Answer 2

total number of protons and neutrons in the nucleus

Question 3

mass and charge of subatomic particles

Answer 3

proton 1, +1
electron 1/2000, -1
neutron 1, 0

Question 4

What are isotopes?

Answer 4

Different atoms of the same element with the same proton number but different neutron number

Question 5

What is the Relative Atomic Mass?

Answer 5

The weighted mean mass of an atom of an element compared to 1/12th of the mass of an atom of carbon-12

Question 6

What is the Relative Isotopic Mass?

Answer 6

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

Question 7

What is the M+ peak?

Answer 7

Molecular ion peak. Shows the relative molecular mass.

Question 8

What is the M+1 peak?

Answer 8

The peak that has +1 m/z than the molecular ion peak due to the existence of a carbon 13 isotope.

Question 9

Name the subshells, their orbitals and how many electrons each can hold

Answer 9

S, p, d, f
1, 3, 5, 7
2, 6 , 10, 14

Question 10

Which shape does the s orbital have?

Answer 10

Spherical.

Question 11

Which shape does the p orbitals have?

Answer 11

Dumbell (8) shape. 3 orbitals. Px, Py and Pz.

Question 12

What does spin pairing mean?

Answer 12

Electrons occupying the same orbital have different spins.

Question 13

What are the different subshells from lowest energy to highest?

Answer 13

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

Question 14

What is the electron configuration of Chromium?

Answer 14

1s², 2s², 2p⁶, 3s², 3p⁶, 4s¹, 3d⁵

An electron is excited from the 4s to fill the 3d orbitals.

Question 15

What is the electron configuration of Copper?

Answer 15

1s², 2s², 2p⁶, 3s², 3p⁶, 4s¹, 3d¹⁰
an electron is excited from 4s to fill the 3d subshell.

Question 16

Which block is the one in red?

Answer 16

s block

Question 17

Which block is the one in purple?

Answer 17

d block

Question 18

Which block is the one in green?

Answer 18

p block

Question 19

Which block is the one in purple?

Answer 19

f block.

Question 20

How are line spectra formed? How are they used to identify an element?

Answer 20

1. Electron is in ground state.
2. Electron absorbs energy from surroundings and is excited (moves to a higher-energy shell)
3. Electron moves back to ground state and releases energy in the form of electromagnetic radiation with a specific frequency.
4. Emission spectra shows the frequency of light emitted.

This is unique for every element.

Question 21

How do line spectra prove that electrons exist in quantum shells?

Answer 21

Defined lines in the emission spectra show energy levels are discrete, electrons jump without an inbetween stage. This proves electrons exist in shells and each shell has a fixed energy.

Question 22

What is ionisation energy?

Answer 22

Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous 1+ions

i.e. Na(g) > Na+(g) + e-

Question 23

How does shielding affect ionisation energies?

Answer 23

The more electron shells between the positive nucleus and the negative outer electron that is being removed, the weaker the attraction and so less energy is needed.

Question 24

How does nuclear charge affect ionisation energies?

Answer 24

The more protons in the nucleus, the bigger the positive charge and so the bigger the attraction between nucleus and electrons therefore more energy is needed to remove an electron.

Question 25

How does atomic size affect ionisation energies?

Answer 25

The bigger the atom, the further away the outer electrons are from the nucleus and so the attractive force between them is less so its easier to remove electrons.

Question 26

Why do ionisation energies decrease as you go down the group?

Answer 26

* Down a group the no of shells increases so atomic radius increases, the outer electrons are further from the nucleus (attractive force is weaker so energy required decreases) * Extra inner shells shield outer electrons from nucleus (more shielding) * Positive charge of nucleus increasing is overidden by shielding effect

Question 27

Why does atomic radius decrease across a period?

Answer 27

* Increased nuclear charge (more protons) so more attractive force, pullling electrons closer to the nucleus * Electrons are added to the same shell so shielding is similar.

Question 28

Why does atomic radius increase down a group?

Answer 28

* More quantum shells, so more shielding. * This offsets the fact that there's an increasing nuclear charge.

Question 29

Why does ionisation energy increase as we go across a period?

Answer 29

* Increasing number of protons increases attractive forces towards the nucleus. * Shielding and size of the atom are similar.

Question 30

Why is there a dip in Al?

Answer 30

Aluminium has its outer shell electron in a higher energy subshell, further away from the nucleus than magnesium. Al - ...3s2, **3p1** Mg - ...3s2 Therefore the force of attraction of the electron with the nucleus is weaker and so less energy is needed.

Question 31

Why is there a dip in S?

Answer 31

Both P and S have lectrons in 3p sub shell so shielding is the same. However, sulfur has an orbital with two electrons while phosphorus doesn't. There's repulsion between electrons in the same orbitals so less energy is needed to remove the electron

Question 32

Why is there an increase from Na to Al?

Answer 32

* Increasing positive charge of ions * Increaseing number of delocalised electrons. * Therefore bonds are stronger and so more energy is needed to overcome them.

Question 33

Why is there a peak in Si?

Answer 33

Giant covalent (macromolecular) structure. Many strong covalent bonds hold silicon together so a large amount of energy is required.

Question 34

Why is there a decrease from Si to P?

Answer 34

P₄ is a simple molecular formula so less energy is needed to overcome the intermolecular forces than in the giant covalent in Si.

Question 35

Why is there an increase from P to S?

Answer 35

S₈ larger simple molecular structure so london forces are stronger than in P₄ so more energy is needed to overcome intermolecular forces.

Question 36

Why is there a decrease from S to Cl?

Answer 36

Cl₂ smaller molecular structure so decrease london forces so weaker intermolecular forces and so less energy is needed.

Question 37

Why is there a decrease from Cl to Ar?

Answer 37

Ar exists as individual atoms so weaker forces of attraction.

Question 38

difference between relative molecular mass and relative formula mass

Answer 38

molecular mass used when talking about simple molecules formula mass used for ionic or giant covalent molecules to work out add up atomic numbers

Question 39

how to work out relative atomic mass from mass spectrum

Answer 39

isotopic mass x relative isotopic abundance then divide by sum of isotopic abundances

Question 40

how to predict mass spectra for diatomic molecules

Answer 40

1. turn percentage abundances to decimal 2. make a table showing possible combinations e:g Cl-35xCl35 2x2 table with abdundance x abundance 3. add up any that are the same 4. divide all by the smallest to get whole number ratio, work out rmm of each and plot mass spectrum

Question 41

orbital

Answer 41

region within an atom that holds upto 2 electrons with opposite spins

Question 42

what does group number tell you

Answer 42

the number of electrons in the outer shell is the same for each group so have similar chemical properties

Question 43

what does period number tell you?

Answer 43

the number of electron shells

Question 44

number of electrons in each shell

Answer 44

4 shells 2, 8, 18, 32

Question 45

periodicity

Answer 45

repeating trends in physical and chemical properties