Atomic structure and the periodic table

Question 1

What are the masses of the subtatomic particles?

Answer 1

p = 1
n = 1
e- = 0.0005

Question 2

What are isotopes? atoms of an element with the same number of protons but different number of neutrons

Answer 2

atoms of an element with the same number of protons but different number of neutrons

Question 3

What do the physical and chemical properties of an atom depend on?

Answer 3

• The number and arrangement of electrons decides the chemical properties of an element.
• Isotopes have the same configuration of electrons, so they’ve got the same chemical properties.
• Isotopes of an element do slightly different physical properties though, (e.g. densities, rates of diffusion).
• Physical properties depend on the mass of an atom.

Question 4

Relative atomic mass definition

Answer 4

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

Question 5

Relative isotopic mass definition

Answer 5

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

Question 6

Relative molecular mass definition

Answer 6

The average mass of a molecule compared to 1/12th of the mass of carnon-12.

Question 7

What’s the difference between relative molecular mass and relative formula mass?

Answer 7

Relative molecular mass: Used for simple molecules. To find the relative molecular mass, add up the relative atomic masses of all the atoms in the molecule.
Relative formula mass: Used for ionic or giant covalent molecules. To find the relative formula mass, add up the relative atomic masses of al the ions/atoms in the formula unit.

Question 8

How can you calculate the relative atomic mass of an element from its isotopic abundance?

Answer 8

When given the isotopic abundances in %, multiply each relative isotopic mass by its % relative isotopic abundance, and add up the results. Divide by 100.

Question 9

What is mass spectrometry?

Answer 9

• Mass spectra are produced by mass spectrometers, they tell us about relative isotopic masses and abundances of different elements.

Question 10

What are the y and x-axis for in a mass spectra?

Answer 10

y-axis: Gives the abundance of ions, often as a %. The height of each peak gives the relative isotopic abundance.
x-axis: The units are given as an M/Z value, which is a mass/charge ratio. Since the charge on the ions is mostly 1+, you can often assume the x-axis is simply the relative isotopic mass.

Question 11

How can you work out the relative isotopic mass from mass spectra graph?

Answer 11

1. Multiply each relative isotopic mass (m/z) by its relative isotopic abundance (height of peak) and add up the results.
2. Divide by the sum of isotopic abundances.

Question 12

How can you predict the mass spectra for diatomic molecules? E.g. 2 chlorine isotopes

Answer 12

1. When given 2 isotopes and their masses, express each % as a decimal.
2. Make a table showing all the different Cl2 molecules. For each molecule, mutiply the decimal abundances, to get the relative abundance of each one. You should get 4 values.
3. Look for any molecules in the table that are the same and add up their abundances. In this case 35Cl-37Cl and 37Cl-35Cl are the same, so add up their abundances.
4. Divide all the relative abundances by the smallest relative abundance to get the smallest whole number ratio. And work out the relative molecular mass by adding both atomic masses.

Question 13

How can you find the molecular mass from a mass spectra graph?

Answer 13

• Find the M peak (peak with the highest M/Z value, ignoring any small M+1 that occur due to the presence of any atoms of carbon-13)
• The M/Z value of this peak is the molecular mass.

Question 14

What are orbitals?

Answer 14

The region of an atom that can hold up to 2 electrons with opposite spins.

Question 15

How many orbitals and therefore maximum electrons in the s subshell?

Answer 15

1 orbital
2 electrons

Question 16

How many orbitals and therefore maximum electrons in the p subshell?

Answer 16

3 orbitals
6 electrons

Question 17

How many orbitals and therefore maximum electrons in the d subshell?

Answer 17

5 orbitals
10 electrons

Question 18

How many orbitals and therefore maximum electrons in the f subshell?

Answer 18

7 orbital
14 electrons

Question 19

How many subshells, and therefore total number of electrons in the 1st shell?

Answer 19

1 subshell
2 electrons

Question 20

How many subshells, and therefore total number of electrons in the 1st shell?

Answer 20

2 subshells
8 electrons

Question 21

How many subshells, and therefore total number of electrons in the 1st shell?

Answer 21

3 subshells
18 electrons

Question 22

How many subshells, and therefore total number of electrons in the 1st shell?

Answer 22

4 subshells
32 electrons

Question 23

What’s the shape of an s-orbital?

Answer 23

Spherical

Question 24

What’s the shape of a p-orbital?

Answer 24

Dumbbell
3 types: Px, Py and Pz orbital, which are: horizontal, verticle and diagonal

Question 25

What's the order that subshells fill in? Is there an exception?

Answer 25

Fill up the lowest energy subshells 1st. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f... The 4s has a lower energy than the 3d, even though it's principle quantum number is bigger, so fills up 1st.

Question 26

How do electrons fill orbitals?

Answer 26

They fill them up singly before they pair up.

Question 27

How can you use the periodic table to work out the electronic configurations?

Answer 27

- The periodic can be split into: s-block, p-block and d-block. - s-block elements end their electronic configuration with s1/s2 - p-block elements end their electronic configuration with s2p1-s2p6

Question 28

How are chromium and copper electron configurations different to other atoms' electron configurations?

Answer 28

They donate 1 of their 4s electrons to the 3d subshell. Because they are more stable with a full/half full d-subshell.

Question 29

How does the frequency and wavelength change along the EM spectrum?

Answer 29

Radiation increases in frequency and decreases in wavelength

Question 30

What are electron shells sometimes called?

Answer 30

Quantum shells or energy levels.

Question 31

What is an emission spectrum and how is it formed?

Answer 31

- An emision spectrum shows the frequencies of light emitted when electrons drop down to a lower energy level. They appear as coloured lines on a dark background. - In their ground state, atoms have their electrons in their lowest possible energy levels - If an electron takes in energy from their surroundings, they can move to higher energy levels, further from the nucleus. - At higher energy levels, electrons are said to be excited. - Electrons release energy by dropping from a higher energy level to a lower one. The energy levels all have certain fixed values, they're discrete. - Each element has a different electron arrangement, so the frequencies of radiation absorbed/emitted are different. - Spectrum for each element is unique.

Question 32

Which part of the spectrum does the series of lines appear at: n = 1 (ground state) n = 2 (2nd energy level) n = 3 (3rd energy level)

Answer 32

n = 1 UV n = 2 visible light (what we see in the emission spectrum) n = 3 Infrared

Question 33

Each set of lines on emission spectra get closer as the ____ increases.

Answer 33

Frequency

Question 34

What are the 4 basic principles when it comes to electron shells?

Answer 34

- Electrons can only exist in fixed orbits/shells, and not anywhere in between. - Each shell has a fixed energy - When an electron moves between shells, EM radiation is emitted/absorbed - Due to the energy of the shells being fixed, the radiation will have a fixed frequency.

Question 35

How does emission spectra support the idea of quantum shells?

Answer 35

The emission spectrum of an atom has clear lines for different energy levels. This supports the idea that energy levels are discrete, i.e. not continuous. It means that an electron doesn't 'move' from 1 energy level to the next, it jumps, with no in-between stage at all.

Question 36

1st ionisation energy definition

Answer 36

The energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.

Question 37

Is the 1st ionisation energy exothermic or endothermic?

Answer 37

Endothermic - energy is put in to ionise an atom/molecule.

Question 38

How does nuclear charge affect ionisation energy?

Answer 38

The more protons in the nucleus, the more positively charged the nucleus is, and the stronger the attraction for the electrons. So 1st IE increases.

Question 38

What 3 factors affect ionisation energy?

Answer 38

Nuclear charge Electron shell Shielding

Question 39

How does the number of electron shells affect ionisation energy?

Answer 39

Attraction decreases rapidly with distance. So the further away the electron from the nucleus, the weaker its attraction, so, the lower the 1st IE.

Question 40

How does shielding affect ionisation energy?

Answer 40

As the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nuclear charge. As shielding increases, 1st IE decreases.

Question 41

What does a high ionisation energy mean?

Answer 41

There's a strong attraction between the electron and the nucleus, so more energy is needed to overcome the attraction and remove the electron.

Question 42

How does 1st ionisation change down a group and across a period?

Answer 42

Down a group: 1st IE decreases Across a period: 1st IE increases

Question 43

Why does 1st IE decrease down a group?

Answer 43

- Number of electron shells increases, so outer electrons are further away from the nucleus, reducing their attraction to the nucleus. - The extra inner shells shield the outer electrons from the attraction of the nucleus. - The +ve charge of the nucleus does increase as you go down the group (due to the extra protons), but this effect is overridden by the effect of the extra shells.

Question 44

2nd ionisation energy definition

Answer 44

The energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions.

Question 45

Why does successive ionisation energy increase as the number of electrons removed increases?

Answer 45

- Within each shell, successive ionisation energies increase. This is because electrons are being removed from an increasingly +ve ion, there's less repulsion amongst the remaining electrons, so they're held more strongly by the nucleus. - The big jumps in ionisation energy happen when a new shell is broken into - an electron is being removed from a shell closer to the nucleus.

Question 46

What can a graph of successive ionisation energies tell you?

Answer 46

- The number of electrons removed before the 1st big jump tells you the group number. - By counting the number of points before each big jump, (right to left) you can find the number of electrons in each shell.

Question 47

How is the periodic table arranged?

Answer 47

- The periodic table is arranged into groups/periods. - All elements in a period have the same number of electron shells. - All elements within a group have the same number of electrons in their outer shell. Which means they have similar chemical properties.

Question 48

Periodicity definition

Answer 48

The repeating trends in physical and chemical properties of the elements across each period

Question 49

Why are group 0 inert?

Answer 49

They have completely filled s and p subshells and don't need to bother gaining/losing/sharing electrons.

Question 50

Which electrons do d-block elements tend to lose to form +ve ions.

Answer 50

s and d electrons

Question 51

Why does ionisation energy increase across a period?

Answer 51

- Number of protons increases, the nuclear attraction increases. This means the electrons are pulled in slightly. Making the atomic radius smaller. - No new shells are added, so shielding doesn't change.

Question 52

Why is there a drop in 1st ionisation energy between groups 2 and 3?

Answer 52

- Aluminium's outer electron is in a 3p orbital rather than a 3s. The 3p orbital has a slightly higher energy than the 3s, so the electron is, on average, found further from the nucleus. - The 3p orbital has additional shielding provided by the 3s2 electrons. - Both facters are strong enough to overide the effect of the increased nuclear charge, resulting in the IE dropping slightly. - This provides evidence for the theory of electron subshells.

Question 53

Why is there a drop in 1st ionisation energy between groups 5 and 6?

Answer 53

- Generally, elements with singly filled/full subshells are more stable than those with partially filled subshells, so have higher 1st IE's - Shielding is identical in the phosphorous and sulfur atoms, and the electron is removed from the same orbital. - Electron is removed from a singly-occupied orbital in phosphorous, and a full orbital in sulfur. - The repulsion between 3 electrons in an orbital means that electrons are easier to remove from shared orbitals. - So S has a lower IE than Ph

Question 54

Explain the trend for BPT across period 2 and 3 for the metals: Li, Be, Na, Mg and Al

Answer 54

- MPT and BPT increase across the period due to the metallic bonds getting stronger. - This is because the metal ions have an increasing number of delocalised electrons and a decreasing radius. - This means there's a stronger attraction between the metal ions and delocalised electrons, so stronger metallic bonding.

Question 55

Why do the elements C and Si have the highest BPT in their periods?

Answer 55

- They have giant covalent lattice structures with strong covalent bonds linking all their atoms together. - A lot of energy is needed to break all of these bonds.

Question 56

Explain the trend for BPT for the simple molecular structures: N2, O2, F2, and P4, S8, Cl2

Answer 56

- Their MPT's depend upon the strength of the London forces between their molecules. - London forces are weak and easily overcome, so these elements have low MPT's and BPT's.

Question 57

Why do the noble gases have the lowest MPT and BPT in their period?

Answer 57

They exist as individual atoms (monatomic) resulting in very weak London forces.