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Chemistry > Elements of life > Flashcards

Elements of life Flashcards

1
Q

What are the 3 type of particles ?

A

Neutrons
Protons
Electrons

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2
Q

Describe the electron ?

A
  • It has a -1 charge
  • They wizz around the nucleus in
    shells. These shells take up the most of the volume of the atom.
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3
Q

Describe the nucleus ?

A
  • Most of the mass of the atom is concentrated in the nucleus
  • The diameter of the nucleus is rather tiny compared to the whole of the atom.
  • The nucleus is where you will find the protons and neutrons.
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4
Q

What is the relative mass and charge of the neutron ?

A

Relative mass - 1
Relative charge - 0

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5
Q

What is the relative mass and charge of the proton ?

A

Relative mass - 1
Relative charge - +1

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6
Q

What is the relative mass and charge of the electron ?

A

Relative mass - 1/1836
Relative charge - -1

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7
Q

What is the mass number ?

A

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

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8
Q

What is the atomic number ?

A

The number of protons in the nucleus.

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9
Q

What are isotopes ?

A

Isotopes of an element are atoms with the same atomic number but different mass numbers.

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10
Q

What determines the chemical properties of an element ?

A

The number and arrangement of electrons. Isotopes have the same configuration of electrons, so they have the same chemical properties.

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11
Q

Do isotopes of an element have slightly different physical properties ?

A

Yes - it often depends on the mass of the atom.

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12
Q

Where is the mass concentrated in an atom, and what makes up most of the volume of an atom ?

A

The mass is mostly concentrated in the nucleus of an atom. The electron shells make up the majority of the volume.

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13
Q

Define the term isotopes ?

A

Isotopes of an element are atoms with the same atomic number but different mass numbers.

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14
Q

Hydrogen, deuterium and tritium are all isotopes of each other.

Identify one similarity and one difference between these isotopes ?

A

Similarity - They have the same atomic number. (Same number of protons and electrons)
Difference - They have different mass numbers. (Different number of neutrons)

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15
Q

Hydrogen, deuterium and tritium are all isotopes of each other.

Deuterium can be written as 2H. Determine the number of protons, neutrons and electrons in a neutral deuterium atom ?

A

Protons - 1
Neutrons - 1
Electrons - 1

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16
Q

What did John Dalton describe atoms as ?

A

Solid spheres

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17
Q

What did J.J Thompson discover ?

A

Electrons

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18
Q

What was J.J Thomson’s model called, and why ?

A

The ‘Plum Pudding model’ - a positively charged sphere with negative electrons embedded in it.

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19
Q

What experiment did Ernest Rutherford with his students Hans Geiger and Ernest Marsden conduct ?

A

The Geiger-Marsden experiment.

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20
Q

What was the Geiger-Marsden experiment ?

A

They fired alpha particles (positively charged) at a thin sheet of gold.

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21
Q

What was the conclusions from the Geiger - Marsden experiment ?

A

They expected most of the alpha particles to be deflected very slightly by the positive charge that made up most of the atom.
But most of the alpha particles passed straight through, and a very small number were deflected backwards. This showed that the plum pudding model could not be right.

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22
Q

What model did Rutherford come up with ?

A

Th nuclear model

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23
Q

What is the nuclear model ?

A

A tiny, positively charged nucleus at the centre of the atom, where the majority of the mass is concentrated.
The nucleus is surrounded by a ‘cloud’ of freely orbiting negative electrons.
Most of the atom is empty space.

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24
Q

What particle did Rutherford discover ?

A

The proton (positively charged)

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25
Q

What particle did James Chadwick discover ?

A

The neutron

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26
Q

What are the 4 basic principles that Niels Bohr discovered ?

A
  • Electrons exist in fixed orbits
  • Each shell has a fixed energy
  • When electron moved between shells, electromagnetic radiation is emitted or absorbed - because the energy of shells is fixed, the radiation will have a fixed frequency.
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27
Q

Why did Bohr say that the noble gases are inert ?

A

Shells can only hold a fixed number of electrons, an element’s reactivity is due to its electrons. When an atom has full shells of electrons it is stable and does not react.

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28
Q

What particle did J J Thomas discover ?

A

Negatively charged particles - electrons

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29
Q

Describe the model of the atom that was adopted because of Thomson’s work ?

A

The ‘Plum Pudding model’ - a positively charged sphere with negative electrons embedded in it.

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30
Q

Who developed the ‘nuclear’ model of the atom ? AND What evidence did they have for it ?

A

Rutherford
They used the Geiger-Marsden experiment. (They fired alpha particles at a thin sheet of gold. Most of the particles passed straight through the gold atoms, and a very small number were deflected backwards > so Rutherford came up with a model that could explain this new evidence).

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31
Q

Scientific theories are constantly being revised in the light of new evidence. New theories are accepted because they have been successfully tested by experiments or because they help to explain certain observations.

Niels Bohr thought that the model of the atom proposed by Ernest Rutherford did not describe the electrons in an atom correctly. Why did he think this and how was his model of the atom different from Rutherford’s ?

A

Bohr knew that if an electron was freely orbiting the nucleus it would spiral into it, causing the atom to collapse. HIs model only allowed electrons to be in fixed shells and nor in between them.

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32
Q

Scientific theories are constantly being revised in the light of new evidence. New theories are accepted because they have been successfully tested by experiments or because they help to explain certain observations.

According to the Bohr model of the atom, what happens when electrons in an atom move from one shell to another

A

When a electron moves from one shell to another electromagnetic radiation is emitted or absorbed.

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33
Q

Scientific theories are constantly being revised in the light of new evidence. New theories are accepted because they have been successfully tested by experiments or because they help to explain certain observations.

How did the Bohr model explain the lack of reactivity of the noble gases ?

A

The shells of an atom can only hold fixed numbers of electrons. Noble gases have a full shell electrons and so do not react > the full shell of electrons makes an atom stable.

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34
Q

Scientific theories are constantly being revised in the light of new evidence. New theories are accepted because they have been successfully tested by experiments or because they help to explain certain observations.

How did the Bohr model explain the lack of reactivity of the noble gases ?

A

The shells of an atom can only hold fixed numbers of electrons. Noble gases have a full shell electrons and so do not react > the full shell of electrons makes an atom stable.

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35
Q

Is Relative atomic mass normally a whole number ?

A

No - because it is an average

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36
Q

Is relative isotopic mass normally a whole number ?

A

Yes.

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37
Q

What is the relative atomic mass ?

A

The average mass of an atom of an element on a scale where an atom of carbon-12 is 12.

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38
Q

What is the relative isotopic mass ?

A

The mass of an atom of an isotope of an element on a scale where an atom of carbon-12 is 12.

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39
Q

What is the relative molecular mass ?

A

The average mass of a molecule or formula unit on a scale where an atom of carbon-12 is 12.

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40
Q

How can relative masses be measured by ?

A

Using a mass spectrometer

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41
Q

What 4 things happen when a sample is squirted into to a mass spectrometer ?

A

1) Vapourisation - the sample is turned into gas using an electric heater.
2) Ionisation - the gas particles are bombarded with high-energy electrons to ionise them. Electrons are knocked off the particles, leaving positive ions.
3) Acceleration - the positive ions are accelerated by an electric field.
4) Detection - the time taken for the positive ions to reach the detector is measured. This depends on an ion’s mass and charge - light, highly charged ions will reach the detector first, while heavier ions with a smaller charge will take longer.

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42
Q

How can relative atomic mass be calculated from a mass spectrum ?

A

1) For each peak, read the % relative isotopic abundance from the y-axis and the relative isotopic mass from the x-axis. Multiply them together to get the total mass of each isotope.
2) Add up the totals
3) Divide by 100 (because % were involved)

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43
Q

What is the equation for % relative isotopic abundance ?

A

(relative abundance / total relative abundance) x 100

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44
Q

Explain what relative atomic mass and relative isotopic mass mean ?

A

Relative atomic mass - The average mass of an atom of an element on a scale where an atom of carbon-12 is 12.

Relative isotopic mass - The mass of an atom of an isotope of an element on a scale where an atom of carbon-12 is 12.

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45
Q

Explain the different between relative molecular mass and relative formula mass ?

A

Relative formula mass - It is used for compounds that are ionic, or giant covalent.

Relative molecular mass - to find the relative molecular mass you add up the relative atomic masses.

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46
Q

Explain the different between relative molecular mass and relative formula mass ?

A

Relative formula mass - It is used for compounds that are ionic, or giant covalent.

Relative molecular mass - to find the relative molecular mass you add up the relative atomic masses.

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47
Q

Describe how a mass spectrometer works ?

A

1) Vaporisation - the sample turns into a gas using an electric heater
2) Ionisation - the gas particles are bombarded with high-energy electrons to ionise them. Electrons are knocked off leaving positive ions.
3) Acceleration - the positive ions are being accelerated by an electric field.
4) Detection - the time taken for the positive ions to reach the detector is measured. This depends on the ions mass and charge - light, highly charged ions will reach the detector first, while the heavier ions with a smaller charge will take longer.

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48
Q

What is avogadro’s constant ?

A

6.02 x 10^23 particles

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49
Q

What is the equation for the number of moles ?

A

No. of particles you have / No. of particles in a mole

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50
Q

What is the equation for moles using mass and Mr?

A

Mass / Mr

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51
Q

How do you convert cm^3 to dm^3 ?

A

/ 1000

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52
Q

How do you find the no. of moles using concentration and volume ?

A

Concentration x volume.

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53
Q

Is a solution that has more moles per dm^3 concentrated or dilute ?

A

Concentrated

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54
Q

Is a solution that has less moles per dm^3 concentrated or dilute ?

A

Dilute

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55
Q

Balance this equation :
C2H6 + O2 –> CO2 + H2O

A

C2H6 + 3.5O2 –> 2CO2 + 3H2O

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56
Q

Can you use 0.5 to balance equations ?

A

Yes

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57
Q

Write an ionic for the equation :
HNO3 + NaOH –> NaNO3 + H2O

A

H+ + NO3- + Na+ + OH- –> Na+ +NO3- + H2O

(Cross out any ions that appear on both sides)

H+ + OH- –> H2O

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58
Q

What is one thing you have to check after you have written an ionic equation ?

A

To check that the charges balance

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59
Q

what do the state symbols mean;
1) s
2) l
3) g
4) aq

A

1) solid
2) liquid
3) gas
4) aqueous (solution in water)

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60
Q

Calculate the mass of iron oxide produced if 28.0 g of iron is burnt in air;
2Fe + 1.5O2 –> Fe2O3

A

The molar mass, Mr of Fe = 55.8 g Mol^-1
No. of moles in 28 g of Fe = Mass / Mr
= 28 / 55.8
= 0.502 moles

Mole ratio Fe : Fe2O3
2 : 1
0.502 : 0.251

Mr of Fe2O3 = (2 x 55.8) + (3 x 16)
= 159.6 g mol^-1

Mass = Moles x Mr
= 0.251 x 159.6
Mass = 40.1 g

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61
Q

Aluminium reacts with oxygen to form aluminium oxide, Al2O3. Calculate the number of grams of Al2O3 that could be produced if 2.50g of aluminium and 2.50g of oxygen were allowed to react ?

A
  • First you need to write the equation for the reaction:
    Al(s) + O2(g) –> Al2O3(s)
  • Balance the equation
    4Al(s) + 3O2(g) –> 2Al2O3(s)
  • Work out the moles of the aluminium and oxygen
    Mr of Al = 27 g mol^-1
    Mr of O2 = 32 g mol^-1

Moles = Mass / Mr
Moles of Al = 2.50 / 27 = 0.0926
Moles of O2 = 2.50 / 32 = 0.0781

  • Mole ratio Al : Al2O3 O2 : Al2O3
    4 : 2 3 : 2
    0.0926 : 0.0463 0.0781 : 0.0521

This mean that Al is the limiting reactant, as it produces fewer moles of Al2O3, so 2.50g of Al and 2.50g of O2 can only produce 0.0463 moles of Al2O3.

Mr of Al2O3 = 102 g mol^-1
Mass of Al2O3 = 0.0463 x 102 = 4.72 g

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62
Q

Write down the formula used to calculate the concentration of a solution ?

A

moles / volume

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63
Q

What is the state symbol for a solution of hydrochloric acid in water ?

A

Aqueous

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64
Q

Calculate the mass of propene required to produce 49.2g of bromopropane
C3H6 + HBr –> C3H7Br

A

Moles = Mass / Mr
49.2 / 122.9 = 0.400 moles of bromopropane
Mole ratio
C3H6 : C3H7Br
1 : 1
0.400 : 0.400
Mass = moles x Mr
= 0.400 x 42
Mass = 16.8 g

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65
Q

Balance this equation:
KI + Pb(NO3)2 –> PbI2 + KNO3

A

2KI + Pb(NO3)2 –> PbI2 + 2KNO3

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66
Q

Calcium carbonate (CaCO3) reacts with nitric acid (HNO3) in a neutralisation reaction. The products are carbon dioxide, water and a solution of calcium nitrate (Ca(NO3)2).

Write a balanced equation for this reaction, including state symbols ?

A

CaCO3(s) + 2HNO3(aq)
–> Ca(NO3)2(aq) + CO2(g) + H2O(l)

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67
Q

Calcium carbonate (CaCO3) reacts with nitric acid (HNO3) in a neutralisation reaction. The products are carbon dioxide, water and a solution of calcium nitrate (Ca(NO3)2).

Calculate the volume of 1.50 mol/dm^3 HNO3 need to neutralise 1.84g of CaCO3 in cm^3 ?

A

CaCO3(s) + 2HNO3(aq)
–> Ca(NO3)2(aq) + CO2(g) + H2O(l)

Volume = moles / concentration
Moles = mass / Mr
= 1.84 / 100.1
= 0.018 moles of CaCO3
Mole ratio
HNO3 : CaCO3
2 : 1
0.0368 : 0.018
Volume = 0.0368 / 1.50
Volume = 0.0245 dm^3
0.0245 x 1000 = 24.5 cm^3

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68
Q

What is the empirical formula ?

A

The smallest whole number ratio of atoms in a compound

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69
Q

What is a molecular formula ?

A

The actual number of atoms in a molecule

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70
Q

A molecule has an empirical formula of C4H3O2, and a molecular mass of 166. Work out its molecular mass ?

A
  • First find the relative mass of the atoms in the empirical formula:
    (4 x 12) + (3 x 1) + (2 x 16) = 48 + 3 + 32 = 83
  • But the molecular mass is 166, so there are 166/82 = 2 empirical units in a molecule.
  • The molecular formula must be the empirical formula x 2, so the molecular formula = C8H6O4.
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71
Q

When a hydrocarbon is burnt in excess oxygen, 4.4g of carbon dioxide and 1.8g of water are made. What is the empirical formula of the hydrocarbon ?

A

No. of moles of CO2 = Mass / Mr
= 4.4 / 44
= 0.1 mole

1 mole of CO2, contains 1 mole of carbon atoms, so you must have started with 0.1 moles of carbon atoms.

No. of moles of H2O = Mass / Mr
= 1.8 / 18
= 0.1 mole

1 mole of H2O contains 2 moles of hydrogen atoms, so you must have started with 0.2 moles of hydrogen atoms.

Ratio of C : H
0.1 : 0.2

Now divide both numbers by the smallest - here it is 0.1.

So, the ratio C : H
1 : 2

So the empirical formula must be CH2

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72
Q

A compound is found to have percentage composition 56.5% potassium, 8.70% carbon and 34.8% oxygen by mass. Calculate its empirical formula ?

A

In 100g of compound there are;
56.5 / 39.1 = 1.45 moles of K
8.70 / 12 = 0.725 moles of C
34.8 / 16 = 2.18 moles of O

  • Divide each number of moles by the smallest of these numbers - 0.725
    K : 1.45 / 0.725 = 2.00
    C : 0.725 / 0.725 = 1.00
    O: 2.18 / 0.725 = 3.01

The ratio of K : C : O
2 : 1 : 3

So the empirical formula is = K2CO3

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73
Q

A compound is found to have percentage composition 56.5% potassium, 8.70% carbon and 34.8% oxygen by mass. Calculate its empirical formula ?

A

In 100g of compound there are;
56.5 / 39.1 = 1.45 moles of K
8.70 / 12 = 0.725 moles of C
34.8 / 16 = 2.18 moles of O

  • Divide each number of moles by the smallest of these numbers - 0.725
    K : 1.45 / 0.725 = 2.00
    C : 0.725 / 0.725 = 1.00
    O: 2.18 / 0.725 = 3.01

The ratio of K : C : O
2 : 1 : 3

So the empirical formula is = K2CO3

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74
Q

What is the equation for percentage composition of an element ?

A

(total mass of element in compound) / (total mass of compound) x 100

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75
Q

What is the percentage composition of H in CH4 ?

A

(4 x 1) / (12 + (4 x 1) x 100
= 25 %

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76
Q

When 4.6g of an alcohol, with molar mass 46g, is burnt in excess oxygen, it produces 8.8g carbon dioxide and 5.4g of water. Calculate the empirical formula for the alcohol and then its molecular formula.

A

No. of moles in CO2 = Mass / Mr
= 8.8 / 44 = 0.20 moles
1 mole of CO2 contains 1 mole of C. So, 0.20 moles of CO2 contains 0.20 moles of C.

No. of moles H2O = Mass / Mr
= 5.4 / 18
= 0.30 moles
1 mole of H2O contains 2 moles of H. So, 0.30 moles of H2O contains 0.60 moles of H.

Mass of C = No.of moles x Mr
= 0.20 x 12 = 2.4 g
Mass of H = No. of moles x Mr
= 0.60 x 1 = 0.60 g
Mass of O = 4.6 - (2.4 + 0.60) = 1.6g
Number of moles O = Mass / Mr
= 1.6 / 16
= 0.10 moles

Molar ration: C : H : O
0.20 : 0.60 : 0.10
2 : 6 : 1

Empirical formula = C2H6O

Mass of empirical formula
(12 x 2) + (1 x 6) + 16 = 46g

Molecular mass = C2H6O

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77
Q

What is water of crystallisation ?

A

Water molecules that have been incorporated into compounds.

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78
Q

Do hydrated or anhydrous compounds contain water of crystallisation ?

A

Hydrated

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79
Q

When heated hydrated compounds lose their water of crystallization, to become what ?

A

anhydrous compounds

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80
Q

Heating 3.210g of hydrated magnesium sulfate, MgSO4.xH2O, forms 1.567g of anhydrous magnesium sulfate. Find the value of x and write the formula of the anhydrous salt ?

A

(FIND THE NO. OF MOLES LOST)
Mass of water lost:
3.210 - 1.567 = 1.643g
No. of moles of water lost:
Mass / Mr = 1.643 / 18 = 0.09127 moles

(FIND THE NO. OF MOLES OF ANHYDROUS SALT)
Mr mass of MgSO4:
24.3 + 32.1 + (4 x 16) = 120.4g/mol^1
No. of moles in (1.567g):
Mass / Mr = 1.567 / 120.4 = 0.01301

(NOW YOU WORK OUT THE RATIO OF MOLES OF ANHYDROUS SALT TO MOLES OF WATER IN THE FORM 1: n)
From the experiment,
0.01301 : 0.09127
1 : 7.015

x must be a whole number, and some errors are to be expected in any experiment, so you can safely round off your result - so the formula of the hydrated salt is MgSO4.7H2O

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81
Q

What is the theoretical yield ?

A

The mass of product that should be made in a reaction if no chemicals are ‘lost’ in the process.

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82
Q

Is the actual yield always less than the theoretical yield ?

A

Yes
Not all the starting reactants react fully.
Some chemicals are always lost.

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83
Q

What is the percentage yield ?

A

The actual amount of product you collect, written as a percentage of the theoretical yield.

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84
Q

What is the equation for percentage yield ?

A

(actual yield) / (theoretical yield) x 100

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85
Q

Ethanol can be oxidised to form ethanal:
C2H5OH + [O] –> CH3CHO + H2O. 9.2g of ethanol was reacted with an excess of oxidising agent and 2.1g of ethanal was produced. Calculate the theoretical yield and the percentage yield ?

A

Mr of C2H5OH:
(2 x 12) + (5 x 1) + 16 + 1 = 46 g/mol^1

Moles = Mass / Mr
Moles of C2H5OH
9.2 / 46 = 0.2 moles

1 mole of C2H5OH produces 1 mole of CH3CHO, so 0.2 moles of C2H5OH will produce 0.2 moles of CH3CHO.

Mr of CH3CHO
(2 x 12) + (4 x 1) + 16 = 44 g/mol^1

Theoretical yield (mass of CH3CHO):
Mass = Moles x Mr
0.4 x 44 = 8.8g

Actual yield = 2.1 g
Theoretical yield = 8.8g

Percentage yield = (2.1) / (8.8) x 100 = 24%

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86
Q

What is the difference between the molecular formula and the empirical formula ?

A

Empirical formulas give the smallest whole number ratio of atoms in a compound
Molecular formula gives the actual numbers of atoms in a molecule.

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87
Q

What is the difference between a hydrated and an anhydrous compound ?

A

Hydrated compounds contain water of crystallisation, and anhydrous don’t.

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88
Q

Explain what it is meant by percentage
yield ?

A

The actual amount of product you collect, written as a percentage of the theoretical yield.
Percentage yield:
(actual yield) / (theoretical yield) x 100

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89
Q

Hydrocarbon X has a molecular mass of 78. It is found to have 92.3% carbon and 7.70% hydrogen by mass. Calculate the empirical and molecular formula of x ?

A

No. of moles of C:
92.3 / 12 = 7.69 moles

No. of moles of H:
7.70 / 1 = 7.7 moles

Divide both by the smallest number

So the ratio of C : H
1 : 1
Empirical formula = CH

The mass of the atoms in the empirical formula = 12 + 1 = 13

No. of empirical units in molecule:
78 / 13 = 6

Molecular formula = C6H6

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90
Q

When 1.2g of magnesium ribbon is heated in air, it burns to form a white powder, which has a mass of 2g. What is the empirical formula of the powder ?

A

No. of moles Mg:
1.2 / 24 = 0.05 moles

No. of moles of O:
0.6 / 16 = 0.05 moles (MASS OF O IS EVERYTHING THAT Mg ISN’T)

Ratio of Mg : O
0.05 : 0.05
1 : 1

Empirical formula is MgO

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91
Q

When 19.8g of an organic acid, A, is burnt in excess oxygen, 33.0g of carbon dioxide and 10.8g of water are produced. Calculate the empirical formula for A and hence its molecular formula, if Mr(A) = 132 ?

A

No. of moles of CO2:
33 / 44 = 0.75 moles

Mass of C = 0.75 x 12 = 9g

No. of moles of H2O:
10.8 / 18 = 0.6 moles

0.6 moles of H2O = 1.2 moles of H

Mass of H 1.2 x 1 = 1.2g

Mass of O = 19.8 - (9 + 1.2) = 9.6 g

No. of moles of O:
9.6 / 16 = 0.6 moles

Mole ratio: C : H : O
0.75 : 1.2 : 0.6
1.25 : 2 : 1

(THE CARBON PART ISN’T A WHOLE NUMBER, SO YOU HAVE TO MULTIPLY THEM ALL UP UNTIL IT IS)

Mole ratio, 5 : 8 : 4
Empirical formula: C5H8O4
Empirical mass:
(12 x 5) + (1 x 8) + (16 x 4) = 132

Molecular formula: C5H8O4

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92
Q

42.9g of CuSO4.xH2O is heated until all the water has been driven off, leaving 27.3g of anhydrous copper sulfate. What is the value of x ?

A

Mr of CuSO4 = 63.5 + 32.1 + (4 x 16)
= 159.6g/mol^1
No. of moles of CuSO4 = 27.3 / 159.6
= 0.171 moles
Mass of water = 42.9 - 27.3 = 15.6g
Moles of water = 15.6 / 18 = 0.867 moles
(DIVIDE BY SMALLEST)
Cu = 0.171 / 0.171 = 1
H2O = 0.867 / 0.171 = 5.07
So, the value of x = 5

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93
Q

The equation below shows the reaction for the complete combustion of ethene.
C3H6 + O2 –> CO2 + H2O
If the reaction of 91.3g of C3H6 produces an 81.3% yield, how much grams of CO2 would be produced ?

A

C3H6 + 4.5O2 –> 3CO2 + 3H2O

Mr of C3H6 = (12 x 3) + (1 x 6) = 42
Moles of C3H6 = 91.3 / 42 = 2.17 moles

Mole ratio C3H6 : CO2
1 : 3
2.17 : 6.51
Mr of CO2 = 12 + (2 x 16) = 44

Theoretical yield = 44 x 6.51 = 287 g
Actual yield =
((Percentage yield) x (theoretical yield)) / 100
(81.3 x 287) / 100 = 233g

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94
Q

What does a standard solution have ?

A

A known concentration

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95
Q

How do you make a standard solution, using 250cm3 of a 6.00 mol/dm3 solution of sodium chloride ?

A
  • First work out how many moles of solute you need by using the formula:
    Moles = (conc.) x (vol.(cm3)) / 1000
    (6.00 x 250) / 1000 = 1.50 moles
  • Now work out how many grams of solute is needed using the formula:
    Mass = moles x Mr
    1.50 x 58.5 = 87.75g
  • Carefully weigh out this mass of solute - first weigh the beaker, note the weight, then add the correct mass.
  • Add a small amount of distilled water to the beaker and stir until all the solute has dissolved.
  • Tip the solution into a volumetric flask - make sure its the right size for the volume you’re making. Use a funnel to make sure it all goes in.
  • Rinse the beaker and stirring rod with distilled water and add that to the flask to. This makes sure there is no solute clinging to the beaker or rod.
  • Now top the flask up to the correct volume (250cm3) with the distilled water. Make sure the bottom of the meniscus reaches the line - when you get close to the line use a pipette to add water drop by drop. If you go over the line you’ll have to start all over again.
  • Stopper the bottle and invert a few time to make sure its mixed.
  • Check the meniscus again and add a drop or two of water if you need to.
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96
Q

What happens if the solution is to concentrated ?

A

Adding tiny amounts of water will cause large pH changes and results may become inaccurate.

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97
Q

How do you make standard solutions with a different concentration from a known concentration ?

A
  • First, divide the concentration you want by the concentration you have, and multiply by the volume you want. This gives you the volume of the concentrated solution to use.
  • Then accurately measure out this volume of solution and transfer it to a volumetric flask. You can’t use a balance as you would for measuring out solids. Instead, use a pipette or a burette, then top up the solution to the required volume with distilled water in the volumetric flask.
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98
Q

How do you find the volume to use ?

A

Vol. to use =
(final conc.) / (initial conc.) x vol. required.

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99
Q

What do titrations allow ?

A

Allow you to find out exactly how much acid is needed to neutralise a quantity of alkali.

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100
Q

How do you do a titration ?

A
  • Measure out some alkali using a pipette and put it in a flask, along with some indicator.
  • Do a rough titration to get an idea where the end point is.
  • To do this, take an initial reading to see how much acid is in the burette to start of with. Then, add the acid to the alkali - giving the flask a regular swirl.
  • Stop, when your indicator shows a permanent colour change. Record the final reading from your burette.
  • Now, do an accurate titration. Run the acid in to within 2cm3 from your end point, then add the acid dropwise. If you don’t notice exactly when the solution changed colour you’ve overshot and your result won’t be accurate.
  • Work out the amount of acid used to neutralise the alkali. This is just the final reading minus the initial reading. This volume is known as the titre.
  • It is best to repeat the titration a few times, making sure you get similar answers each time - your readings should be written 0.1cm3 of each other. The calculate a mean, ignoring any anomalous results. Remember to wash out the conical flask between each titration to remove any acid or alkali left in it.
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101
Q

What colour does the solution turn when you add methyl orange ?

A

Turns yellow to red when adding acid to alkali.

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102
Q

What colour does the solution turn when you add phenolphthalein ?

A

Turns red to colourless when adding acid to alkali

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103
Q

Calculate the volume of 1mol/dm3 hydrochloric acid solution needed to make 500cm3 of 0.5mol/dm3 hydrochloric acid solution ?

A

Volume
(0.5 / 1) x 500 = 250cm3

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104
Q

Describe how indicators are used and explain the importance of selecting an appropriate indicator when carrying out a titration. Include examples of indicators that would and would not be suitable for use in titrations.

A

Indicators change colour when the solution reaches a particular pH to mark an end point. They are used in acid/alkali titrations to mark the end point of the reaction. Indicators used in titrations need to change colour quickly over a very small pH range. A few drops of indicator solution are added to the analyte. The analyte can be place on a white surface to make a colour change easy to see. Methyl orange and phenolphthalein are both good indicators for titrations as they quickly change colour when the solution turns from alkali to acid. Universal indicator is a poor indicator to use for titrations as its colour changes gradually over a wide pH range.

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105
Q

25.0cm3 of 0.500mol/dm3 HCl was used to neutralise 35.0cm3 of NaOH solution. Calculate the concentration of the sodium hydroxide solution ?

A

HCl + NaOH –> NaCl + H2O
Volume HCl = 25.0cm3
Conc. of Hcl = 0.500mol/dm3

Volume of NaOH = 35.0cm3

Moles = conc. x vol.
Moles of HCl = (25.0 x 0.500) / 1000
Moles of HCl = 0.0125 moles

Mole ratio HCl : NaOH
1 : 1
0.0125 : 0.0125

Conc. = moles / volume
Conc. (0.0125 / 35.0) x 1000
Conc. of NaOH = 0.360mol/dm3

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106
Q

What is the equation for moles, using volume and concentration ?

A

Moles = Concentration x volume

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107
Q

20.4cm3 of a 0.500mol/dm3 solution of sodium carbonate reacts with 1.5mol/dm3 nitric acid. Calculate the volume of nitric acid required to neutralise the sodium
carbonate ?

A

Na2CO3 + 2HNO3 –> 2NaNO3 + H2O + CO2

Vol. of Na2CO3 = 20.4cm3
Conc. of Na2CO3 = 0.500mol/dm3

Conc of HNO3 = 1.5mol/dm3

Moles = Conc. x vol.
Moles of Na2CO3 = (20.4 x 0.500) / 1000
Moles of Na2CO3 = 0.0102 moles

Mole ratio Na2CO3 : HNO3
1 : 2
0.0102 : 0.0204

Vol. of HNO3 = (moles / conc.) x 1000
Vol. of HNO3 = (0.0204 / 1.5) x 1000
Vol. of HNO3 = 13.6cm3

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108
Q

What is the equation for moles, using the mass and Mr ?

A

Moles = Mass / Mr

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109
Q

Calculate the volume of 2.00mol/dm3 hydrochloric acid required to neutralise 3.86g of calcium hydroxide, Ca(OH)2 ?

A

Ca(OH)2 + 2HCl –> CaCl2 + 2H2O

Mr of Ca(OH)2
40.1 + (2 x 16) + (2 x 1) = 74.1g/mol

Moles = Mass / Mr
Moles = 3.86 / 74.1
Moles = 0.0521 moles

Mole ratio Ca(OH)2 : HCl
1 : 2
0.0521 : 0.104

Vol. of HCl = (moles / conc.) x 1000
= (0.104 / 2.00) x 1000
Vol. of HCl = 52.0cm3

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110
Q

What equation links the number of moles, concentration and volume ?

A

Moles = concentration x volume

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111
Q

Calculate the concentration (in mol/dm3) of a solution of ethanoic acid, CH3COOH, if 25.4cm3 of it is neutralised by 14.6cm3 of 0.500mol/dm3 sodium hydroxide solution.

CH3COOH + NaOH –> CH3COONa + H2O

A

Vol. of CH3COOH = 25.4cm3

Vol. of NaOH = 14.6cm3
Conc. of NaOH = 0.500mol/dm3

Moles of NaOH = (14.6 x 0.500) / 1000
Moles of NaOH = 0.0073 moles

Mole ratio of CH3COOH : NaOH
1 : 1
0.0073 : 0.0073

Conc. = (moles / vol.) x 1000
( 0.0073 / 25.4) x 1000
Conc. = 0.29mol/dm3

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112
Q

You are supplied with 0.750g of calcium carbonate and a solution of 0.250mol/dm3 sulfuric acid. What volume of acid will be needed to neutralise the calcium
carbonate ?

CaCO3 + H2SO4 –> CaSO4 + H2O + CO2

A

Mass of CaCO3 = 0.750g
Mr of CaCO3 = 100.1
Moles = Mass / Mr
Moles = 0.750 / 100.1
Moles = 7.49 x 10^-3

Mole ratio CaCO3 : H2SO4
7.49 x 10^-3 : 7.49 x 10^-3

Vol. of H2SO4 = (7.49 x 10^-3 / 0.250) x 1000
Vol. of H2SO4 = 29.96 cm3

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113
Q

In a titration, 17.1cm3 of 0.250mol/dm3 hydrochloric acid neutralises 25.0cm3 calcium hydroxide solution.

a) Write out a balanced equation for this reaction

b) Work out the concentration of the calcium hydroxide solution.

A

a) Ca(OH)2 + 2HCl –> CaCl2 + 2H2O

b) Moles = (Vol. x conc.) / 1000
= (17.1 x 0.250) / 1000
= 0.004275 moles of HCl

Mole ratio Ca(OH)2 : 2HCl
1 : 2
0.0021375 : 0.004275

Conc. of Ca(OH)2 = (0.0021375 / 25.0) x 1000
Conc. of Ca(OH)2 = 0.0855mol/dm3

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114
Q

What are electron shells made from ?

A

Sub-shells and orbitals

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115
Q

Which shells have a greater energy levels ?

A

Shells further to the nucleus have a greater energy level then shells closer to the nucleus.

116
Q

What is the maximum number of electrons a s-shell can hold ?

A

2

117
Q

What is the maximum number of electrons a p-shell can hold ?

A

6

118
Q

What is the maximum number of electrons a d-shell can hold ?

A

10

119
Q

What shape is an s orbital ?

A

A spherical shape

120
Q

What shape is a p orbital ?

A

dumbbell shape.

121
Q

How do you get a compound ?

A

When different elements join or bond together

122
Q

How are ions formed ?

A

Electrons are transferred from one atom to another

123
Q

What is ionic bonding ?

A

Electrostatic attraction between two oppositely charged ions.

124
Q

What structure does sodium chloride have ?

A

Giant ionic lattive

125
Q

When do ionic compounds conduct electricity ?

A

When they are molten or dissolved - not when they are solid. The ions in a liquid or a solution are free to move, in a solid they’re fixed in a position by the strong ionic bonds.

126
Q

Do ionic compounds have a high or low melting point ?

A

High melting point. The giant ionic lattices are held together by strong electrostatic forces, it takes loads of energy to overcome these forces, so melting points are very high.

127
Q

Are ionic compound soluble or insoluble ?

A

Soluble - Water molecules are polar, part of the molecule has a small negative charge, and the other bits have small positive charges. The water molecules pull the ions away from the lattice and cause it to dissolve.

128
Q

Whats a compound ?

A

When different elements join or bond together.

129
Q

What type of forces hold ionic substances together ?

A

Electrostatic attraction

130
Q

What happens when a current is passed through a dissolved ionic compound ?

A

They conduct electricity

131
Q

a) ions can be formed by electron transfer.
Explain this and give an example for a positive and negative ion.

b) Solid lead bromide does not conduct electricity, but molten lead bromide does.
Explain this with reference to ionic
bonding ?

A

a) Electrons move from one atom to another
Na+ and Cl-

b) In a solid, ions are held in place by strong ionic bonds. When the solid is heated to melting point, the ions gain enough energy to overcome the forces of attraction and become mobile, so can carry charge through substance.

132
Q

Which is the correct formula for aluminium chloride ?

a) AlCl
b) Al2Cl3
c) Al3Cl
d) AlCl3

A

D

133
Q

How are molecules formed ?

A

When two or more atoms bond together, and held together by covalent bonds

134
Q

What is covalent bonding ?

A

When two atoms share electrons, so they have both got full outer shells of electrons. Both the positive nuclei are attracted electrostatically to the shared electrons.

135
Q

What is covalent bonding ?

A

When two atoms share electrons, so they have both got full outer shells of electrons. Both the positive nuclei are attracted electrostatically to the shared electrons.

136
Q

What needs to happen to the forces for a covalent bond to be
maintained ?

A

The electrostatic forces of attraction between the nuclei and the electrons, there is also a repulsion between the positive nuclei.

137
Q

What are the typical properties for small molecules ?

A
  • Fairly low boiling point and melting point
  • Do not conduct electricity, because they have no charge carriers that are free to move.
  • Insoluble in water
138
Q

What is the other covalent bond ?

A

Dative.

139
Q

What is a dative covalent bond ?

A

One atom donates both electrons to a bond.

140
Q

Give an example of a dative covalent bond ?

A

Ammonium ion, NH4+.

141
Q

What types of atoms form covalent bonds ? Non-metals, metals or both.

A

Non-metals

142
Q

Describe 3 typical physical properties of a molecular substance ?

A
  • Fairly low melting and boiling point
  • Do not conduct electricity
  • Insoluble in water
143
Q

Describe two physical properties that you would expect silicon hydride ?

A
  • Non-conductor of electricity
  • Low melting point
144
Q

Methane melts at -183oC. Magnesium oxide melts at 2800oC. Explain this difference in terms of bonding ?

A

Methane is a simple molecular substance. When it melts only weak intermolecular forces are overcome. Magnesium oxide is ionic, so to melt it the strong electrostatic attractions must be broken, which requires more energy/higher temperatures.

145
Q

Predict whether ammonium chloride is lily to be soluble or insoluble in water. Explain your answer ?

A

Ammonium chloride is likely to be soluble in water. The polar water molecules can break apart an ionic lattice.

146
Q

When do covalent bonds form ?

A

When atoms share electrons with other atoms.

147
Q

Why is covalent bonded substances more stronger than the simple covalent molecules.

A

Covalently bonded substances have giant structures, which have many covalently bonded atoms, The electrostatic attraction holding the atoms together in these structures are much stronger than the electrostatic attraction between simple covalent molecules.

148
Q

What are some examples of giant covalent networks. Why ?

A

Carbon and silicon
This is because they can each form four strong, covalent bonds.

149
Q

What are some properties of giant structures for covalent bonding ?

A
  • Have a high melting point
  • Hard
  • Good thermal conductors
  • Insoluble in polar substances
  • Cannot conduct electricity
150
Q

What giant covalent structure can conduct electricity ?

A

Graphite

151
Q

What do most metal elements exist as ?

A

Giant metallic lattice structures

152
Q

What is metallic bonding ?

A

In metallic lattices, the electrons in the outermost shell of the metal atoms are delocalised - they are free to move. This leaves a positive metal ion. The positive metal ions are attracted to the delocalised negative electrons. They form a lattice of closely packed positive ions in a sea of delocalised electrons.

153
Q

What are some properties of metallic bonding ?

A
  • They have a high melting point
  • They can be shaped and ductile
  • Good thermal conductors
  • Good electrical conductors
  • Insoluble
154
Q

What are some properties of metallic bonding ?

A
  • They have a high melting point
  • They can be shaped and ductile
  • Good thermal conductors
  • Good electrical conductors
  • Insoluble
155
Q

Are the melting points of giant covalent lattices high or low? Explain why ?

A

HIgh melting and boiling points.
Need to break a lot of very strong bonds before the substances melts, which takes a lot of energy.

156
Q

Why won’t giant covalent structures dissolve ?

A

The covalent bonds mean atoms are more attracted to their neighbours in the lattice than to solvent molecules.

157
Q

Explain how the model of metallic bonding accounts for:
i) a metals ability to conduct electricity
ii) the relative high melting points of metals

A

i) The delocalised electrons are free to move and can carry a current.

ii) The strong metallic bonding, with the number of delocalised electrons per atoms affecting the melting point. The more electrons there are, the stronger the bonding will be and the higher the melting point.

158
Q

a) Explain what is meant by metallic bonding ?

b) Explain why calcium has a higher melting point than potassium ?

A

a) metallic bonding results from the attraction between positive metal ions and a sea of delocalised electrons between them.

b) Calcium has two delocalised electrons per atom, while potassium has only one delocalised electrons per atom. So calcium has more delocalised electrons and therefore a stronger metallic bonding.

159
Q

Carbon dioxide is a covalent compound that sublimes at -78oC. Silicon dioxide is a covalent compound that melts at 1610oC. Explain the difference in terms of bonding ?

A

Carbon dioxide is a simple molecular substance. Covalent bonds between carbon and oxygen are strong but when carbon dioxide sublimes it is the weak intermolecular forces that are overcome. Silicon dioxide has a giant covalent lattice structure, so to melt it the strong covalent bonds must be broken, which require more energy/higher tempuratures.

160
Q

Graphite is a giant covalent structure. However, unlike most giant covalent structures, it is able to conduct electricity. Explain why graphite is able to conduct electricity ?

A

Graphite consists of sheets of carbon atoms, where each carbon atom is bonded to three others. This means that each atom has one free electron not involved in bonds, and it is these free electrons that allow graphite to conduct electricity.

161
Q

Electrical grade copper must be 99.99% pure. If sulphur and oxygen impurities react with copper ions, its electrical conductivity is dramatically reduced. USe your knowledge of metallic and ionic bonding to explain this.

A

Copper is metallically bonded and so delocalised electrons are free to move. Oxygen and sulfur form copper oxide, fixing some electrons. This prevents them from moving and carrying charge.

162
Q

Carborundum (silicon carbide) has the formula SiC and is almost hard as diamond.
a) What sort of structure would you expect carborundum to have as a solid ?

b) Apart from hardness, give two other physical properties you would expect carborundum to have ?

A

a) Giant covalent

b) High melting point
Does not conduct electricity

163
Q

What is the electron pair repulsion theory ?

A
  • Electrons are all negatively charges, so electron pairs will repel each other as much as they can.
  • This sounds straightforward, but the type of the electron pair affects how much it repels other electron pairs. Lone pairs repel more then bonding pairs.
  • This means the greatest angles are between lone pairs of electrons, and bond angles between bonding pairs are often reduced because they are pushed together by lone-pair repulsion.
  • So the shape of the molecule depends on the type of electron pairs surrounding the central atom as well as the number.
164
Q

How is the periodic table arranged ?

A

In atomic number

165
Q

What do you call rows in a periodic table ?

A

Periods

166
Q

What do you call columns in a periodic table ?

A

Groups

167
Q

Do all the elements within the period have the same number of electron shells. Example ?

A

Yes
The elements of period one both have 1 electron shell
The elements in period two have two electron shells

168
Q

Do all the elements within a group have the same number of electrons in their outer shell. Example ?

A

Yes
Group one elements have 1 electron in their outer shell
Group four elements have 4 electrons in their outer shell

169
Q

Do all elements in the same group have similar physical and chemical properties ?

A

Yes

170
Q

Why do the melting points of metals increase across the period ?

A

The metal-metal bonds get stronger. The bonds get stronger because the metal ions have an increasing number of delocalised electrons and a decreasing ionic radius. This leads to a higher charge density, which attracts the ions together more strongly.

171
Q

Why do elements with giant covalent structure have a strong covalent bond ?

A

A lot of energy is needed to break these bonds. Carbon and silicon have the highest melting points in their period

172
Q

What do the melting points of the simple molecular substances depend on ?

A

They depend on the strength of the intermolecular forces between the molecules.

173
Q

Do simple molecules have a high or low melting point. Why ?

A

A low melting point because intermolecular molecules are weak and easily broken.

174
Q

What does more atoms in a molecule mean ?

A

The stronger intermolecular forces

175
Q

Why do noble gases have the lowest melting point ?

A

They exist as individual atoms, resulting in vry weak intermolecular forces.

176
Q

What is the first ionisation enthalpy ?

A

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

177
Q

What are some important points about ionisation
enthalpies ?

A
  • Must use the gas symbol
  • Always refer to 1 mole of atoms
  • The lower the ionization enthalpy, the easier it is to remove an outer electron and form an ion.
178
Q

What are the 3 main things that can affect the size ionisation enthalpies ?

A
  • The atomic radius
  • The nuclear charge
  • Electron shielding
179
Q

What is the atomic radius ?

A

The further the outer shell electrons are from the positive nucleus, the less they’ll be attracted towards the nucleus. So, the ionisation enthalpy will be lower.

180
Q

What is the nuclear charge ?

A

This is the positive charge on the nucleus caused by the presence of protons. The more protons there are in the nucleus, the more it’ll attract the outer electrons - it’ll be harder to remove the electrons, so the ionisation enthalpy will be higher.

181
Q

What is electron shielding ?

A

The inner electron shells shield the outer electrons from the attractive force of the nucleus. Because more inner shells mean more shielding, the ionisation enthalpy will be lower.

182
Q

Why does ionisation enthalpies decrease down the
group ?

A

As you go down the group, the outer electrons are in shells further from the nucleus, so they’re attracted to the nucleus less.
The amount of shielding increases because there are more filled inner shells. This means less nuclear attraction between the outer shell electrons.
Although the number of protons increases down the group, this doesn’t lead to an increase in ionisation enthalpy because it’s a less important factor then either shielding or the distance of the outer electrons from the nucleus.

183
Q

Why does ionisation enthalpies increase across a period ?

A

As you move across a period, the trend is for the ionisation enthalpies to increase, it gets harder to remove the outer electron.
This is because the number of protons is increasing, so the outer electrons are attracted more strongly to the nucleus.
And since all the outer-shell electrons are at roughly the same energy level, there’s generally little extra shielding effect or extra distance too lessen the attraction from the nucleus.
There are small drops between Group 2 and 3, and 5 and 6 but you don’t need to know about why these happen.

184
Q

Why does the s-block metals have the lowest ionisation enthalpy ?

A

The s-block metals have relatively low nuclear charges, so they have low first ionisation enthalpies compared to other elements in the same period. This means they lose their outer electron easily, as there is less attraction between the nucleus and the outer electrons. This makes the s-block metals reactive.
The p-block metals have higher nuclear charges than s-block metals in the same period. This is down to the increase in the number of protons across each period. This means hey have higher first ionisation enthalpies, so it is more difficult to lose their outer electron and they are less likely to lose an electron.

185
Q

In what order are the elements set out in the periodic table ?

A

By atomic number

186
Q

What is the name given to the columns in the periodic table ?

A

Groups

187
Q

Define the first ionisation enthalpy and give an equation as an example.

A

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-

188
Q

Describe the three main factor that affect ionisation enthalpies.

A
  • Atomic radius
  • Nuclear charge
  • Electron shielding
189
Q

Explain why first ionisation enthalpies show an overall tendency to increase across a period

A

As you move across a period, the number of protons increases, meaning a stronger pull from the positively charged nucleus making it harder to remove an electron from the outer shell. There are no extra inner electrons to add to the shielding affect.

190
Q

Explain why the melting point of magnesium is higher than that of sodium

A

Magnesium has more delocalised electrons per atom and the ion has a greater charge density. This gives magnesium a stronger metal-metal bond, resulting in a higher melting point.

191
Q

State and explain the trend in the first ionisation enthalpy of group 2 elements.

A

first ionisation enthalpies decreases down group 2 because the outer electrons are in shells further from the nucleus. The amount of shielding also increases because there are more filled inner shells. Both of these decrease the nuclear attraction for the outer shell electrons.

192
Q

What go group 2 metals create when they react with water ?

A

Metal hydroxides and hydrogen

193
Q

What do group 2 metals create when they react with oxygen ?

A

Oxides

194
Q

Does the group 2 elements increase or decrease reactivity when you go down the group ?

A

Increase reactivity

195
Q

Do group 2 oxides and hydroxides create an alkaline or acidic solution in water ?

A

Alkaline solution

196
Q

What makes solutions alkaline ?

A

-OH ions

197
Q

Why does the solutions get more strongly alkaline as you go down group 2 ?

A

The hydroxides get more soluble

198
Q

Are group 2 hydroxides and oxides bases ?

A

Bases

199
Q

Why do group 2 hydroxides and oxides neutralise acids. Give an example.

A

They are bases, both the oxides and hydroxides neutralise dilute acids, forming solutions of corresponding salts.
MgO + 2HCl –> H2O + MgCl2

200
Q

Do the Group 2 elements increase or decrease solubility down the group ?

A

Increase

201
Q

Do the group 2 hydroxides increase or decrease solubility down the group ?

A

Increase

202
Q

Do the group 2 carbonates increase or decrease solubility own the group ?

A

Decrease

203
Q

What do group 2 carbonates decompose to ?

A

CO2 and metal oxides

204
Q

What is thermal decomposition ?

A

When a substance breaks down when heated

205
Q

Does thermal stability increases or decreases as you go down group 2 ?

A

Increases down the group

206
Q

Why does thermal stability increase down group 2 ?

A

The carbonate ions are large anions and can be made unstable by the presence of a cation. The cation draws the electrons on the carbonate ion towards itself > polarises it. This distorts the carbonate ion. The greater the distortion, the less stable the carbonate ion.
Large cations cause less distortion than small cations, as they have a lower charge density. So the further down the group, the larger the cations, the less distortion caused and the more stable the carbonate anion.

207
Q

Which is the least reactive metal in group 2 ?

A

Beryllium

208
Q

Why does reactivity with water increase down group 2 ?

A

The outermost electrons are further from the nucleus, and so more easily lost

209
Q

Describe two reactions of a group 2 oxide that shows it to be a base.

A
  • They form alkaline solutions in water
    (The oxides of the group 2 metals reacts readily with water to form metal hydroxides, which dissolve)
  • They neutralise acids
    (The oxides neutralise dilute acids)
210
Q

Which is less soluble, barium hydroxide or magnesium hydroxide ?

A

Magnesium hydroxide

211
Q

Write a general equation for the thermal decomposition of a group 2 carbonate. Use M to represent the group 2 metal.

A

MCO3 –> MO + CO2

212
Q

Barium and calcium are both Group 2 elements.

a) Which of the two ions has the highest charge density ?

b) Write a balanced equation for the thermal decomposition of calcium carbonate including state symbols.

c) Barium and calcium both form carbonates. State whether barium carbonate or calcium carbonate is more thermally stable. Explain your answer.

A

a) Calcium

b) CaCO3(s) –> CaO(s) + CO2(g)

c) Barium carbonate is more thermally stable. This is because barium has a lower charge density han calcium, so it has weaker polarising powers. The weaker the polarising power of the barium ion causes less distortion of the carbonate ion.

213
Q

a) Write a balanced equation for the reaction of magnesium hydroxide with dilute hydrochloric acid.

b) Write balanced equation for the reaction of calcium oxide with water

c) Explain why calcium hydroxide is a stronger alkali then magnesium hydroxide.

A

a) Mg(OH)2 + 2HCl –> MgCl2 + 2H2O

b) CaO + H2O –> Ca(OH)2

c) Calcium hydroxide is more soluble in water than magnesium hydroxide, and so dissociation produces higher concentration of hydroxide ions in water.

214
Q

a) Write a balanced equation for the reaction of magnesium hydroxide with dilute hydrochloric acid.

b) Write balanced equation for the reaction of calcium oxide with water

c) Explain why calcium hydroxide is a stronger alkali then magnesium hydroxide.

A

a) Mg(OH)2 + 2HCl –> MgCl2 + 2H2O

b) CaO + H2O –> Ca(OH)2

c) Calcium hydroxide is more soluble in water than magnesium hydroxide, and so dissociation produces higher concentration of hydroxide ions in water.

215
Q

Are salts ionic compounds ?

A

Yes

216
Q

What are acids ?

A

Substances with a pH of less than 7

217
Q

What are bases ?

A

Substances that have a pH of more than 7

218
Q

What are neutralisation reactions ?

A

When acids react with bases to form a salt and water

219
Q

How are salts formed ?

A

Positively charged cations and negatively charged anions, so the product is neutral

220
Q

What are the soluble salts ?

A
  • Lithium, Sodium, Potassium and Ammonium
  • Nitrates
  • Most chlorides, bromides and iodides
  • Most sulfates
    -Lithium, Sodium, Potassium, Strontium, Calcium, Barium and Ammonium hydroxide
  • Lithium, Sodium, Potassium and Ammonium carbonate
221
Q

What are the insoluble salts ?

A
  • Silver halides (Copper iodide, Lead chloride and Lead bromide, Lead iodide)
  • Barium sulfate, Calcium sulfate and Lead sulfate
  • Most hydroxides
  • Most carbonates
222
Q

What is formed when you use a precipitation reaction ?

A

An insoluble salt

223
Q

How is an insoluble salt is formed using a precipitation reaction ?

A
  • Pick two solutions that contain the ions you need.
  • Once the salt has precipitated out, all you have to do is filter it from the solution, wash it and then dry it on filter paper
224
Q

How do you make a soluble salt, using a metal or insoluble base ?

A
  • Pick the right acid, plus a suitable metal or an insoluble
    base (a metal oxide or metal hydroxide).
  • Picking the right acid is easy - to make chlorides use hydrochloric acid to make sulfates use sulfuric acid and to make nitrates use nitric acid.
  • Add the solid metal or metal oxide or hydroxide to the acid - it will dissolve in the acid as it reacts. You will know when all the acid has been neutralised because no more solid will dissolve, so it will just sink to the bottom of the flask.
  • Then filter out the excess metal, metal oxide or metal hydroxide to get the salt solution.
  • To get pure, solid crystals of the salt, evaporate some of the water and then leave the rest to evaporate very slowly. This is called crystallisation.
225
Q

How do you make a soluble salt using an alkali ?

A
  • Add exactly the right amount of alkali to just neutralise the acid - the most accurate way to do this is to use a titration. This method uses an indicator to show you exactly how much alkali neutralises a known volume of acid.
  • Once you found out how much alkali you need to neutralise the acid, repeat the titration by combining these volumes again - just don’t add the indicator this time, otherwise it wil contaminate the salt.
  • Then just evaporate off the water to crystallise the salt as normal.
226
Q

What is the name of the salt formed from ammonium and sulfate ions ?

A

Ammonium sulfate

227
Q

Name one soluble salt and one insoluble salt ?

A

Soluble salt : Lithium hydroxide

Insoluble salt : Lead chloride

228
Q

What is a precipitation reaction and how is it important in making insoluble salts ?

A

A precipitation reaction is when a cation and anions combine to form an insoluble ionic salt, and the ionic salt creates a precipitate.

229
Q

Which aid would you use to make silver chloride ?

A

Hydrochloric acid

230
Q

What can you use to show when a neutralisation reaction is complete, if both the base and the salt are soluble in water ?

A

Indicator

231
Q

What is the formula of the soluble salt produced in the reaction between nitric acid and copper(II) hydroxide ?

A No3CU2
B CUOH2
C H2NO3
D Cu(NO3)2

A

D

232
Q

Which of the following compounds would react with iron(III) chloride to produce an insoluble salt ?

A Sodium Chloride
B Silver Nitrate
C Copper(II) Sulfate
D Calcium Sulfate

A

B

233
Q

Potassium sulfate can made by mixing potassium hydroxide with sulfuric acid. Describe the method you would use to produce pure, solid potassium sulfate in the lab ?

A

Titrate some potassium hydroxide with a known volume of sulfuric acid and add some drops of indicator to show when the reactions finished. Then repeat the titration using the same amount of acid and the exact amount of potassium hydroxide needed to neutralise it, so the salt isn’t contaminated with indicator. To get pure potassium sulfate, evaporate some of the water and then leave the rest to evaporate slowly.

234
Q

Describe the test for flame tests ?

A
  • Dip a wire loop in concentrated hydrochloric acid
  • Then dip the wire loop into the sample of the compound
  • Hold the loop in the clear blue part of the flame
  • Observe the colour change in the flame
235
Q

What are the flame tests for:
- Lithium
- Sodium
- Potassium
- Calcium
- Barium
- Copper

A

-Crimson
- Yellow
- Lilac
- Brick red
- Green
- Blue-green

236
Q

What is the test for metal ions ?

A

Add sodium hydroxide

237
Q

What precipitate is formed when you add sodium hydroxide to a silver ion. Give the ionic equation.

A

Brown
Ag+ + 2OH- –> Ag2O + H2O

238
Q

What precipitate is formed when you add sodium hydroxide to a calcium ion. Give the ionic equation.

A

White
Ca2+ + 2OH- –> Ca(OH)2

239
Q

What precipitate is formed when you add sodium hydroxide to a copper(II) ion. Give the ionic equation.

A

Blue
Cu2+ + 2OH- –> Cu(OH)2

240
Q

What precipitate is formed when you add sodium hydroxide to a lead(II) ion. Give the ionic equation.

A

White
Pb2+ + 2OH- –> Pb(OH)2

241
Q

What precipitate is formed when you add sodium hydroxide to a iron(II) ion. Give the ionic equation.

A

Green
Fe2+ + 2OH- –> Fe(OH)2

242
Q

What precipitate is formed when you add sodium hydroxide to an iron(III) ion. Give the ionic equation.

A

Reddish brown
Fe3+ + 3OH- –> Fe(OH)3

243
Q

What precipitate is formed when you add sodium hydroxide to a zinc ion. Give the ionic equation.

A

White.
Zn2+ + 2OH- –> Zn(OH)2
Then add excess NaOH to form a colourless solution.
Zn(OH)2 + OH- –> Zn(OH)3-

244
Q

What precipitate is formed when you add sodium hydroxide to an aluminium ion. Give the ionic equation.

A

White
Al3+ + 3OH- –> Al(OH)3
Then add excess NaOH to form a colourless solution.
Al(OH)3 + OH- –> Al(OH)4-

245
Q

What is the test for carbonates ?

A

Add hydrochloric acid

246
Q

Describe the tests for carbonates ?

A
  • Add dilute hydrochloric acid, carbonates will fizz because they produce carbon dioxide
  • Then add limewater to test for carbon dioxide
  • Carbon dioxide turns limewater cloudy > bubble the gas through a test tube of lime water. If the water goes cloudy you’ve identified a carbonate ion.
247
Q

What is the test for sulfates ?

A

Add HCl and Barium Chloride

248
Q

Describe the test for sulfates ?

A
  • Add dilute hydrochloric acid
  • Then add barium chloride
  • If a white precipitate of barium sulfate forms, it mean the original compound contained a sulfate.
249
Q

What is the test for ammonium compounds ?

A

Use litmus paper and NaOH

250
Q

Describe the test for ammonium compounds ?

A
  • Use a damp piece of red litmus paper, if ammonia is present them it will turn blue
  • Add sodium hydroxide to the substance and heat the mixture, this is the test for check for ammonium ions.
251
Q

What is the test for hydroxides ?

A

Use litmas paper

252
Q

Describe the test for hydroxides ?

A
  • Add a piece of red litmus paper into the solution
  • If a hydroxide if present, the paper will turn blue
253
Q

What is the test for halides ?

A

Add silver nitrate solution

254
Q

Give the ionic equation and the colour of precipitates when chlorides, bromides and iodides react with silver nitrate ?

A

Chlorides
White
Ag+ + CL- –> AgCl

Bromides
Cream
Ag+ + Br- –> AgBr

Iodides
Pale yellow
Ag+ + I- –> AgI

255
Q

What is the test for nitrates ?

A

Use NaOH and Aluminium

256
Q

Describe the test for nitrates ?

A
  • Warm the solution with sodium hydroxide and aluminium foil
  • The aluminium reduces the nitrate ions to ammonium ions
  • The ammonium ions then react with the hydroxide ions to produce ammonia gas and water
    NH4+ + OH- –> NH3 + H2O
  • So if ammonia is given off, your know your substance must contain nitrate ions.
  • You can test for ammonia gas, it should turn a piece of red litmus paper blue.
257
Q

Describe how to carry out a flame test ?

A
  • Dip a wire loop in concentrated hydrochloric acid.
  • Dip the wire loop into the sample of the compound.
  • Hold the loop in the clear blue part of the busen brner
  • Observe the colour change in the flame.
258
Q

Name two substances you need to add to a sample to test for sulfates.

A
  • Dilute hydrochloric acid
  • Barium Chloride
259
Q

How could you identify the presence of a halide ion.

A

Add dilver nitrate and dilute nitric acid

260
Q

In which two tests, would you use damp red litmus paper to test for ammonia gas.

A
  • To test for ammonia compounds
  • To test for nitrates
261
Q

Describe a test that can be used to test for carbonates in a solution ?

A

Add dilute hydrochloric acid to the solution and then test to see whether the gas given off its carbon dioxide by bubbling it through limewater. If the lime water goes cloudy, the solution contains carbonates.

262
Q

What color flame would be produced from a solution of potassium iodide.

A Yellow
B Brick red
C Lilac
D Crimson

A

C

263
Q

What colour precipitate would be produced from the reaction of calcium carbonates and sodium hydroxide ?

A Brick red
B White
C Blue
D Cream

A

B

264
Q

A student is given a solution of iron(II) sulfate. Describe how the student could prove that the solution actually is iron(II) sulfate.

A

Add a few drops of sodium hydroxide to a sample of the solution. If a green precipitate forms, then iron(II) ions are present. Then dilute hydrochloric acid followed by barium chloride solution to a sample of the solution. A white precipitate will form if sulfate ions are present.

265
Q

What is electromagnetic radiation ?

A

Energy that’s transmitted as waves

266
Q

Give the order of the electromagnetic radiation in order of increasing freqiency and decreasing wavelength ?

A

Radio Waves
Microwaves
Infrared
Visible light
Ultra-violet
X-rays
Gamma rays

267
Q

What happens if electrons take in energy from their surroundings ?

A

They move to higher energy levels, further from the nucleus. At a higher energy level, electrons are said to be excited.

268
Q

How canelectrons release energy ?

A

By dropping from a higher energy level to a lower energy level

269
Q

Do al; energy levels have a certain fixed value ?

A

Yes

270
Q

What is an absorption spectra ?

A

Dark lines on a coloured background

271
Q

What is an emission spectra ?

A

Bright lines non a black background

272
Q

Are the emission spectra different for different elements ?

A

Yes

273
Q

What does it mean if you have lots of lines in these spectra ?

A

Each set represents electrons moving to or form a different energy level.

274
Q

Is energy related to frequency ?

A

Yes

275
Q

What is the equation for energy ?

A

Difference in Energy = frequency x Planck’s constant

276
Q

When an electron falls from the third to the second energy level of a hydrogen atom it emits visible light with a frequency of 4.57 x 10^14 Hz.
Planck’s constant = 6.63 x 10^-34 Hz^-1. What is the difference in energy between the third and second energy level of a hydrogen atom ?

A

E = hv

E = (6.63 x 10^-34) x (4.57 x 10^14) = 3.03 x 10^-20 J

277
Q

What is the speed of light equation ?

A

Speed of light = frequency x wavelength

278
Q

What is nuclear fusion ?

A

When two smaller nuclei combine under high temperatures and pressure to make one larger nucleus.

279
Q

In stars, when hydrogen nuclei combine, what is formed ? and What is released ?

A

Helium nuclei
Los of energy uis released

280
Q

What happens when the hydrogen in a star core runs out ?

A

The temperature and pressure starts to rise

281
Q

WHat happens in a big star when it gets really hot ?

A

It will be able to fuse together heavier elements, starting with helium.

282
Q

Describe what the atomic emission spectrum of hydrogen shows ?

A

When electrons drops to lower energy levels, they give out certain amounts of energy.

283
Q

What equation links speed, wavelength and frequency of the waves of the electromagnetic spectrum ?

A

Speed of electromagnetic radiation = frequency x wavelength

284
Q

What is nuclear fusion ? Where does it hppen in nature ?

A

Nuclear fusion is when two smaller nuclei fuse together in high temperatures and high pressures to form a larger nuclei. It happens naturally in stars.

285
Q

The emission spectrum of the element sodium shows a set of lines in the visible part of the spectrum. There is a strong line at a frequency of 5.10 x 10^14 Hz, which corresponds to the colour yellow.

a) What is the energy of the electron transition responsible for this line ?
(Planck’s constant = 6.63 x 10^-34 J Hz^-1)

b) What is the wavelength of this radiation emitted ?
(Speed of electromagnetic radiation = 3.00 x 10^8 m/s)

A

a) E = hv
E = (6.63 x 10^-34) x (5.10 x 10^14)
E = 3.38 x 10^-19 J

b) c = Frequency x wavelength
Wavelength = speed of light / frequency
Wavelength = (3.00 x 10^8) / (5.10 x 10^14)
Wavelength = 5.88 x 10^-7 m/s

Chemistry (8 decks)

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