EL

Question 1

Mass number definition

Answer 1

No of protons and neutrons

Question 2

Atomic number definition

Answer 2

Number of protons = number of electrons
Electrons relative mas - 1/2000

Question 3

What is an isotope?

Answer 3

Same number of protons different number of neutrons

Question 4

What is the empirical formula?

Answer 4

The simplest whole number ratio of atoms in a compound

Question 5

Percentage yield =

Answer 5

actual yield / theoretical yield
x 100

Question 6

Exceptions of electron configuration

Answer 6

Copper - 3d10 4s1
Chromium - 3d5 4s1
This means they’re more stable

Question 7

Periods 2 and 3 metal periodic trends

Answer 7

Melting points increase across the period because bonds get stronger
This is because there’s an increasing number of delocalised electrons and a decreasing ionic radius
This results in a higher charge density which attracts the ions more strongly

Question 8

Periods 2 and 3 giant covalent structures periodic trends (C and Si)

Answer 8

Have the highest melting points in their periods

Question 9

Periods 2 and 3 simple molecular structures periodic trends

Answer 9

Their melting points depend upon the strength of the intermolecular forces between their molecules
Intermolecular forces are weak and easily broken so they have low melting points

The noble gases have the lowest melting points because they’re monoatomic with very weak intermolecular forces

More atoms in a molecule means stronger intermolecular forces

Question 10

What is the first ionisation enthalpy?

Answer 10

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

X(g) –> X+(g) + e-

The lower the ionisation enthalpy the easier it is to remove an outer electron and form an ion

Question 11

What 3 things affect the size of ionisation enthalpies?

Answer 11

1) Atomic radius
2) Nuclear charge
3) Electron shielding

Question 12

What is electron shielding?

Answer 12

The inner electrons shield the outer shell electrons from the attractive force of the nucleus. More inner shells results in more shielding
This means the ionisation enthalpy will be lower

Question 13

What is atomic radius?

Answer 13

The further the outer shell electrons from the positive radius the less nuclear attraction

Ionisation enthalpy will be lower

Question 14

What is nuclear charge?

Answer 14

The positive charge of the nucleus, caused by the number of protons
The more protons, the more the nucleus will attract outer electrons so ionisation enthalpy will be higher

Question 15

First ionisation enthalpy trends with groups

Answer 15

Decrease down a group
Because there’s less nuclear attraction to the outer electrons
This shows that electrons are arranged in energy levels
Shielding increases which also decreases nuclear attraction

Question 16

First ionisation enthalpy with periods

Answer 16

Increase across a period
Because of the increasing number of protons so stronger nuclear attraction
There’s little extra shielding as all outer-shell electrons are roughly at the same energy level

Question 17

Why do s-block metals have low ionisation enthalpies?

Answer 17

Relatively low nuclear charges
So less nuclear attraction

P-block metals have higher nuclear charges because of the increase in protons across each period. So higher first ionisation enthalpy

Question 18

Group 2 elements reaction with water

Answer 18

Form metal hydroxides and hydrogen

Ca(s) + 2H2O(l) –> Ca(OH)2(aq) + H2(g)

Question 19

Reactivity trends down group 2

Answer 19

Increasingly more reactive because there’s less nuclear attraction so outermost electrons lost more easily

Question 20

Group 2 elements burning in oxygen

Answer 20

Form solid white oxides

2Ca(s) + O2(g) –> 2CaO(s)

Question 21

How do group 2 oxides and hydroxides form alkaline solutions?

Answer 21

Group 2 oxides react readily with water to form metal hydroxides which dissolve
The OH- ions make these solutions strongly alkaline
Mg oxide is an exception as it reacts slowly and the hydroxide isn’t fully soluble
The hydroxides get more soluble down the group so the oxides form more strongly alkaline solutions

Question 22

How do group 2 metals neutralise acids?

Answer 22

Group 2 oxides and hydroxides are bases
Both neutralise dilute acids forming salts

MgO(s) + 2HCl(aq) –> H2O(l) + MgCl2(aq)

Mg(OH)2(s) + 2HCl(aq) –> 2H2O(l) + MgCl2(aq)

Question 23

Solubility trends in group 2

Answer 23

Depend on the compound anion

If the compound contains singly charged negative ions (e.g. OH-) there’s an increase in solubility down the group

If the compound contains doubly charged negative ions (e.g. CO32- or SO42-) there’s a decrease in solubility down the group

Question 24

What do group 2 carbonates thermally decompose to form?

Answer 24

Carbon dioxide and metal oxides
The volume of CO2 produced decreases down the group as the metal gets bigger so a smaller number of MCO3 moles are contained in the same mass

Question 25

Thermal stability down group 2 carbonates

Answer 25

Thermal stability increases down the group Carbonate ions are large anions which can be made unstable in the presence of a cation (such as group 2 metal ion) The cation polarises the carbonate ion - the greater the distortion, the less stable the carbonate ion The larger the cation the less distortion as lower charge density So, the further down the group, cations get larger Lower charge density so less distortion and the more stable the carbonate anion

Question 26

What is charge density?

Answer 26

The charge on the ion relative to its volume Group two cations all have a charge of 2+ but the charge density decreases down the group

Question 27

Acid + base -->

Answer 27

salt + water

Question 28

Name the soluble salts

Answer 28

Nitrates Sodium Lithium Potassium Ammonium Most chlorides, iodides and bromides (except for silver halides) Sulfates

Question 29

Name the insoluble salts

Answer 29

Most hydroxides except lithium, sodium, potassium, calcium, strontium, barium, and ammonium Most carbonates except lithium, potassium, sodium and ammonium Insoluble carbonates form coloured precipitates

Question 30

Insoluble carbonate precipitates

Answer 30

Silver carbonate - yellow Copper carbonate - blue-green Iron (II) - off white Barium, calcium, lead (II) and zinc - white

Question 31

Making an insoluble salt

Answer 31

Use a precipitation reaction Pick two solutions that contain the ion you need Once the salt has precipitated filter from solution, wash and dry

Question 32

Making soluble salts using an alkali

Answer 32

Use a titration to add exactly the right amount of alkali to just neutralise the acid Repeat the titration, this time without indicator otherwise the salt will be contaminated Evaporate off the water and crystallise the salt

Question 33

Making soluble salts using a metal or insoluble base

Answer 33

To make chlorides use HCl, to make sulfates use H2SO4, to make nitrates use HNO3 Add the solid metal, metal oxide or hydroxide to the acid to neutralise acid. When all the acid has been neutralised no more solid will dissolve Filter out excess metal to get salt solution Crystillise to get pure solid crystals of the salt

Question 34

Exam question on electron energy levels

Answer 34

E=hv

Question 35

Order of electromagnetic spectrum in increasing frequency and decreasing wavelength

Answer 35

Radio Micro Infrared Visible UV X-rays Gamma rays

Question 36

Absorption vs emission spectra

Answer 36

Absorption is dark lines on coloured background Emission is coloured lines on dark background Both line spectra with lines in the same position for a given element Lines get closer together as frequency increases