Embed Size (px)
Citation preview
Ed. I
The Norwood School Science Faculty
Retrieval Practice | Chemistry Paper 2
Quiz Questions | Edexcel
Science involves the recall and application of a lot of information.
You will not get a good grade unless you can remember this information.
This booklet will help you form and keep this information in your brain.
Regularly test yourself!
What this book won’t do: It is also very important that you are able to apply knowledge to new situations. This book will not help you practice this. You should ask your teacher for practice exam-style questions once you are confident with the content in this book.
2
Edexcel 9-1 Quiz Questions
Chemistry
Paper 2
Topic 1 – Key concepts in chemistry Topic 6 – Groups in the periodic table Topic 7 – Rates of reaction and energy changes Topic 8 – Fuels and Earth science
-------------------------------------------------------
3
Topic 1 - Key Concepts in Chemistry
Atomic Structure
1. Describe how Dalton's ideas about atoms have changed over time. 2. Describe how the subatomic particles are arranged in an atom. 3. Explain how atoms of different elements are different.
4. Recall the charges and relative masses of the three subatomic particles by completing the table:
5. Explain why all atoms have no overall charge.
6. Describe how the size of an atom compares to the size of its nucleus. 7. State where most of the mass of an atom is found. 8. State the meaning of relative atomic mass number. 9. Which number tells you how many protons are in the nucleus of an atom? 10. Which number tells you how many electrons an atom has?
11. Which numbers tell you how many neutrons are in nucleus of an atom? 12. The mass number of an atom is equal to the sum of what?
13. State the number of electrons in an atom of chlorine.
14. Calculate the numbers of protons, neutrons and electrons in an atom of potassium.
15. State what is meant by the term isotope.
16. Why are some relative atomic masses (Ar) not whole numbers?
17. [H] Recall the formula for calculating the relative atomic mass (Ar) of an element (given data relating to
isotopes) 18. [H] Calculate the relative atomic mass (Ar) of bromine using the following data:
Bromine consists of two isotopes, 50% 79Br and 50% 81Br, calculate the relative atomic mass, Ar, of bromine from the mass numbers.
Sub-atomic particle Relative mass Relative charge
Neutron
Proton
Electron
4
The Periodic Table 1. Describe how Mendeleev arranged elements into a periodic table.
2. Describe how Mendeleev predicted the existence and properties of some elements yet to be discovered.
3. Describe some problems Mendeleev had when ordering the elements.
4. Explain the meaning of the term ‘atomic number’.
5. Describe how the elements are arranged in the modern periodic table.
6. Recall the positions of metals and non-metals in the periodic table.
7. State what the term ‘electronic configuration’ means.
8. Write electronic configuration of an atom of oxygen.
9. Draw a diagram to show the electron configuration of oxygen. 11. Explain the links between an element’s position in the periodic table and its electronic configuration.
12. Explain the rules of filling electron shells of atoms.
5
Ionic Bonding 1. Define the term ion 2. What is the term given to a positively charged ion?
3. What is the term given to a negatively charged ion?
4. State the charge of a group 1 ion
5. State the charge of a group 2 ion
6. State the charge of a group 6 ion
7. State the charge of a group 7 ion
8. Define the term ionic bond
9. Explain how ionic bonds are formed
10. Describe the formation of a sodium ion.
11. Describe the formation of a chloride ion 12. Explain how the ionic compound magnesium oxide (MgO) forms.
13. Group 1 ions often bond with ions from which group?
14. Group 2 ions often bond with ions from which group?
15. Explain the use of -ide and -ate at the end of compound names
16. Draw a diagram, showing electron transfer, to explain the formation of sodium oxide (Na2O) from sodium and
oxygen atoms.
17. Identify the formula of a sulphate ion
18. Identify the formula of an oxide ion
19. Identify the formula of a hydroxide ion
20. Identify the formula of a chloride ion
21. Identify the formula of a carbonate ion
22. Identify the formula of a nitrate ion
23. Explain the physical structure of ionic compounds.
24. Describe two physical properties of ionic compounds
6
Covalent Bonding 1. Define the term covalent bond.
2. Explain how a covalent bond is formed.
3. Define the term molecule.
4. Which of the following is the best estimation of the size of an atom or small molecule?
a) 1.2 x 10-10 metres b) 1.2 x 10-3 metres c) 1.2 x 105 metres d) 1.2 x 10-18 metres
Types of Substance
1. Draw dot and cross diagrams to represent the following simple covalent molecules:
a. Hydrogen
b. Hydrogen chloride
c. Water
d. Methane
e. Oxygen
f. Carbon dioxide
2. Complete the table to identify the features of substances with different types of bonding.
Relative
melting and boiling point
Relative solubility in water
Ability to conduct
electricity as a solid
Ability to conduct electricity in solution (or
molten)
Ionic bonding
Simple molecular covalent
Giant covalent
Metallic
7
4. Explain why ionic substance have relatively high melting and boiling points.
5. Explain why simple molecular covalent substances have relatively low melting and boiling points.
6. Explain why giant covalent molecules have relatively high melting and boiling points.
7. Explain why metallic substances have relatively high melting and boiling points.
8. Explain why ionic substances can only conduct electricity when aqueous (in solution) or molten.
9. Explain why covalent substances are poor conductors of electricity.
10. Explain why metallic substances can always conduct electricity.
11. Use the data below to identify the type of bonding in each substance. Explain your reasons for each.
12. Graphite and diamond are different forms (allotropes) of which element?
13. Identify the structure and bonding of graphite and diamond.
14. Describe the structure of diamond.
15. Describe the structure of graphite
16. Identify a use of diamond. 17. Identify two uses of graphite
18. Explain why graphite is used to make electrodes and as a lubricant.
19. Explain why diamond is used in cutting tools.
20. C60 is an example of which group of substances?
21. Describe and explain the properties of fullerenes.
22. Using poly(ethene) as an example, describe the structure of a polymer.
23. Describe and explain the physical properties of metals.
24. Describe the bonding in metals.
25. Identify four types of representations used to model simple molecules
8
Calculations Involving Mass
The calculations in this unit will require lots of practice. Ask your teacher for more examples. 1. Relative formula mass can be calculated using which values?
2. Calculate the relative formula mass of the following
a) Sodium bromide, NaBr b) Water, H2O c) Magnesium nitrate, Mg(NO3)2
3. Define the term empirical formula 4. State the four stages of calculating the empirical formula of a compound. 5. Calculating empirical formula – example calculation:
3.2g of sulfur reacts with oxygen to produce 6.4 g of sulfur oxide. What is the formula of the oxide? 6. State the empirical formula of these compounds from their molecular formula.
a) H2O2 b) C4H8 c) Na2O
7. Explain how you can deduce the molecular formula of a compound from its empirical formula and its relative molecular mass.
8. A compound has the empirical formula CH2, it has a relative molecular mass (Mr) of 56. What is it’s molecular formula?
9. Describe an experiment to determine the empirical formula of a simple compound such as magnesium oxide.
10. Use two examples to explain the law of conservation of mass.
11. Find the theoretical mass of magnesium oxide (MgO) that can be obtained by reacting 60g of magnesium with oxygen.
12. Recall the formula for calculating concentration
13. How do you convert a number from cm3 to dm3? 14. [H] Define the term mole
15. [H] State Avogadro’s constant
16. [H] State the two equations involved in mole and mass calculations
17. [H] Calculate the moles of carbon in 18g
18. [H] Calculate the number of particles in 2 mol of sodium chloride, NaCl.
19. [H] Calculate the number of particles in 3g of magnesium chloride, MgCl2 20. [H] What is meant by the term ‘stoichiometry’?
***** If you are aiming for an 8 or a 9 grade, there are a number of more technical calculations that are required in this part of the course. Ask your teacher for lots of practice questions! *****
9
Topic 6 – Groups in The Periodic Table
1. Group 1 elements are also known by what name? 2. Group 7 elements are also known by what name? 3. Group 0 elements are also known by what name?
4. Describe two unique features of alkali metals 5. Which two products are made when group 1 metals react with water?
6. Complete the following word equations:
Lithium + water -> Sodium + water -> Potassium + water ->
7. Write a balanced equation, including state symbols, for the reaction between sodium and water.
8. Describe the pattern in reactivity of the alkali metals with water. 9. Explain the trend reactivity of alkali metals with water. 10. Recall the colour and physical states of the halogens (group 7) at room temperature.
11. Halogens are diatomic, explain what this means.
12. Describe the trend in melting and boiling points as you go down group 7.
13. Describe the chemical test for chlorine.
14. Halogens react with metals to form which compounds? 15. Complete the following word equations as examples of halogen + metal reactions.
Chlorine + sodium -> Fluorine + potassium ->
16. Halogens react with hydrogen to produce which type of compound? 17. Hydrogen halides will dissolve in water to form what type of solution? 18. What type of chemical reaction can show the relative reactivities of the halogens?
19. Describe, using examples, how displacement reactions can be used to show the relative reactivities of
the halogens.
20. [H] Explain why halogen displacement reactions are redox reactions. Use equations in you answer.
10
21. Explain the trend in reactivity of the halogens. 22. Describe the reactivity of the noble gases 23. Explain why noble gases are inert.
24. Describe four uses of noble gases 25. Explain how the properties of the noble gases make them suitable for their uses. 26. Give two examples of properties that show a trend within the group of noble gases
11
Topic 7 – Rates of Reaction and Energy Changes 1. Identify four types of chemical reactions where you can measure the temperature change. 2. Describe a method for measuring the temperature (and thus energy change) accompanying a
chemical reaction. 3. Define exothermic reaction and give two examples. 4. Define endothermic reaction and give two examples. 5. Complete the following sentences using either exothermic or endothermic:
‘Bond breaking is ___________________’ ‘Bond making is __________________’
6. Explain, using ideas about bond breaking and making, why the overall heat energy change for a
reaction may be exothermic. 7. Explain, using ideas about bond breaking and making, why the overall heat energy change for a
reaction may be endothermic. 8. Draw an energy profile diagram to show an exothermic reaction
9. Draw an energy profile diagram to show an endothermic reaction
10. Some reactions are slower and some are faster, we say they have different what? 11. Give an example of a reaction with a slow rate and an example of a reaction with a fast rate.
12. Which four factors may affect the rate of a chemical reaction? 13. Describe and explain the effect of an increased temperature on the rate of a typical reaction. 14. Describe and explain the effect of increasing concentration on the rate of a chemical reaction. 15. Describe and explain the effect of increasing surface area of a solid reactant on the rate of a chemical
reaction. 16. Describe and explain the effect of a suitable catalyst on the rate of a chemical reaction.
17. Explain how a catalytic convertor in a car works
18. Describe an experiment that could be completed in order to investigate the effect of temperature of
the rate of a reaction. 19. Describe an experiment that could be completed in order to investigate the effect of surface area (of
a solid) on the rate of a reaction. 20. Describe an experiment that could be completed in order to investigate the effect of concentration
on the rate of a reaction.
12
Topic 8 – Fuels and Earth Science
1. Define the term hydrocarbon
2. Describe four features of crude oil.
3. Describe the process by which crude oil is separated.
4. Recall the names and uses of the fractions in crude oil
5. Identify and explain four ways in which the hydrocarbons in different fractions of crude oil differ from
one another.
6. Most of the hydrocarbons found in crude oil are members of which homologous series?
7. Explain four features of a homologous series.
8. Explain the term complete combustion
9. Describe the term incomplete combustion
10. Explain how carbon monoxide behaves as a toxic gas.
11. Describe other problems associated with incomplete combustion (not carbon monoxide as a toxic gas).
12. Explain how sulfur dioxide is produced
13. Explain what causes acid rain
14. Explain some problems caused by acid rain
15. Explain why nitrogen oxides are produced by cars
16. Describe some problems caused by nitrogen oxides in the air.
17. Evaluate the advantages and disadvantages of using hydrogen, rather than petrol, as a fuel for cars.
18. Recall three non-renewable fossil fuels obtained from crude oil
19. Recall one non-renewable fossil fuel obtained from natural gas
20. Explain the process of cracking
21. Explain why cracking is necessary
13
Topic 9 - Earth and Atmospheric Science
1. The gases that formed the Earth’s early atmosphere were produced by what?
2. Describe the composition of Earth’s early atmosphere
3. Explain how the oceans formed
4. Explain why the amount of carbon dioxide in the early atmosphere decreased
5. Explain why the amount of oxygen in the atmosphere increased over time
6. Describe the chemical test for oxygen
7. Name three greenhouse gases
8. Describe the greenhouse effect
9. Evaluate the evidence for human activity as the cause of climate change.
10. Describe the composition of today’s atmosphere
11. Describe the potential effects of increased levels of carbon dioxide and methane generated by
human activity
12. Explain how the effects of climate change may be mitigated (this means limited or prevented)
14
Topic 1 - Key Concepts in Chemistry
Atomic Structure
1. Describe how Dalton's ideas about atoms have changed over time. John Dalton described atoms as solid spheres, stating that each different element had a different
‘sphere’. This idea was developed by J. J. Thompson who proposed the ‘plum pudding model’ that described
atoms as positively charged spheres containing small, negatively charged particles. Ernest Rutherford later showed this to be incorrect, the gold foil experiment involved firing alpha
particles at a thin sheet of gold. Three conclusion were made from this: a) Most particles passed straight through the foil, showing the atom to be mostly empty space. b) Some positively charged alpha particles were deflected more than expected, showing there was a positively charged part of the atom that was repelling the alpha particles. c) Finally, some particles were deflected backwards which showed that there was a central part of the atom that contained most of it’s mass.
Neils Bohr proposed a model that developed Rutherford's ideas further. Bohr suggested that electrons orbited the nucleus arranged in shells with each shell having a fixed energy.
2. Describe how the subatomic particles are arranged in an atom.
There is a very small, central nucleus containing the protons and neutrons. This is surrounded by shells (energy levels) of orbiting electrons.
3. Explain how atoms of different elements are different.
They have different proton numbers
4. Recall the charges and relative masses of the three subatomic particles by completing the table:
5. Explain why all atoms have no overall charge.
Because they contain equal number of protons (which are positively charged) and electrons (which are negatively charged).
6. Describe how the size of an atom compares to the size of its nucleus. The atom is extremely large in comparison to its nucleus.
7. State where most of the mass of an atom is found.
In the nucleus. 8. State the meaning of relative atomic mass number.
The relative atomic mass is the mass of an atom of an element relative to 1/12th the mass of an atom of carbon-12.
9. Which number tells you how many protons are in the nucleus of an atom?
The atomic number = proton number
Sub-atomic particle Relative mass Relative charge
Neutron 1 0
Proton 1 +1
Electron 1/2000 -1
15
10. Which number tells you how many electrons an atom has? The atomic number = electron number
11. Which numbers tell you how many neutrons are in nucleus of an atom?
The mass number – the atomic number = neutron number 12. The mass number of an atom is equal to the sum of what?
The protons and neutrons in the atom’s nucleus.
13. State the number of electrons in an atom of chlorine. 17
14. Calculate the numbers of protons, neutrons and electrons in an atom of potassium. Protons = 19 Neutrons = 20 Electrons = 19
15. State what is meant by the term isotope.
Atoms of the same element (same number of protons) with different number of neutrons. 16. Why are some relative atomic masses (Ar) not whole numbers?
Relative atomic mass numbers are an average of all existing isotopes of an element
17. [H] Recall the formula for calculating the relative atomic mass (Ar) of an element (given data relating to isotopes)
(mass of isotope 1 x abundance) + (mass of isotope 2 x abundance) Relative atomic mass (Ar) = _________________________________________________________
total abundance (100 if abundance is %) 18. [H] Calculate the relative atomic mass (Ar) of bromine using the following data:
Bromine consists of two isotopes, 50% 79Br and 50% 81Br, calculate the relative atomic mass, Ar, of bromine from the mass numbers.
(79 x 50) + (81 x 50)
Relative atomic mass (Ar) = _________________ 100
Answer = 80
16
The Periodic Table 1. Describe how Mendeleev arranged elements into a periodic table.
Mendeleev ordered the elements in periods according to their relative atomic mass and grouped them according to their chemical and physical properties.
2. Describe how Mendeleev predicted the existence and properties of some elements yet to be discovered.
Mendeleev noticed trends within periods and groups and hypothesised that some elements must not have yet been discovered. Mendeleev left gaps for these elements (e.g. gallium) and predicted their properties. When they were found, the properties of the new elements closely matched those predicted by Mendeleev and further improved the reputation of his period table.
3. Describe some problems Mendeleev had when ordering the elements.
The relative atomic masses of some elements didn’t fit the trend of ascending relative atomic and fit into groups of similar elements. Mendeleev made the choice to swap these elements (e.g. tellurium and iodine).
4. Explain the meaning of the term ‘atomic number’.
The number of protons in an element.
5. Describe how the elements are arranged in the modern periodic table. Elements are arranged in order of increasing atomic number in rows called periods. Elements with similar properties are placed in the vertical columns called groups.
6. Recall the positions of metals and non-metals in the periodic table.
Metals are found to the left and in the centre of the table, non-metals are to the right.
7. State what the term ‘electronic configuration’ means. It is how the electrons are arranged in an atom.
8. Write electronic configuration of an atom of oxygen.
2.6 9. Draw a diagram to show the electron configuration of oxygen.
10. Explain the links between an element’s position in the periodic table and its electronic configuration.
The number of outer shell electrons = the group of the element in the periodic table. The number of outer electrons determines an element’s chemical properties; this explains why elements in the same group have similar chemical properties. E.g. oxygen is in group 6 and has 6 electrons in it’s outer shell.
11. Explain the rules of filling electron shells of atoms.
The first shell must be full before moving on to the second and so forth. The first shell can fit up to 2 electrons The second shell can fit up to 8 electrons The third shell can fit up to 8 electrons
17
Ionic Bonding
1. Define the term ion An ion is an atom, or group of atoms, that have either gained or lost electrons in order to become charged.
2. What is the term given to a positively charged ion?
A cation 3. What is the term given to a negatively charged ion?
An anion 4. State the charge of a group 1 ion
+
5. State the charge of a group 2 ion 2+
6. State the charge of a group 6 ion 2-
7. State the charge of a group 7 ion -
8. Define the term ionic bond The electrostatic attraction between oppositely charged ions.
9. Explain how ionic bonds are formed
Ionic bonds are formed when electrons and transferred from one atom (which then becomes a positive ions) to another atom (which then becomes a negative ion).
10. Describe the formation of a sodium ion. The sodium atom loses the electron on it’s outer shell in order to become positively charged: Na -> Na+ + e-
11. Describe the formation of a chloride ion
The chlorine atom gains one electron onto it’s outer shell in order to become negatively charged: Cl + e- Cl-
12. Explain how the ionic compound magnesium oxide (MgO) forms.
The magnesium atom loses its two outer electrons to become an Mg2+ ion. The oxygen atom gains these two electrons to become an O2- ion. It is the attraction between these ions that forms the ionic bond.
13. Group 1 ions often bond with ions from which group? Group 7
14. Group 2 ions often bond with ions from which group?
Group 6 15. Explain the use of -ide and -ate at the end of compound names
-ide is used for any compound that has two different elements bonding together. -ate is used for any compound that has three or more elements bonding together including oxygen.
18
16. Draw a diagram, showing electron transfer, to explain the formation of sodium oxide (Na2O) from sodium and oxygen atoms.
17. Identify the formula of a sulphate ion
SO42-
18. Identify the formula of an oxide ion
O2-
19. Identify the formula of a hydroxide ion OH-
20. Identify the formula of a chloride ion Cl-
21. Identify the formula of a carbonate ion CO3
2-
22. Identify the formula of a nitrate ion NO3
-
23. Explain the physical structure of ionic compounds. They have a regular arrangements of ions in a lattice structure that are held together by strong electrostatic forces (ionic bonds) between oppositely charges ions.
24. Describe two physical properties of ionic compounds
They have high melting and boiling points and the conduct electricity when in solution (aqueous) or molten. They do not conduct when solid.
19
Covalent Bonding 1. Define the term covalent bond.
The electrostatic attraction between a nucleus and a shared pair of electrons.
2. Explain how a covalent bond is formed. A covalent bond is formed when atoms share electrons.
3. Define the term molecule. (Usually) a small group of atoms bonded covalently.
4. Which of the following is the best estimation of the size of an atom or small molecule? a) 1.2 x 10-10 metres b) 1.2 x 10-3 metres c) 1.2 x 105 metres d) 1.2 x 10-18 metres
Types of Substance
1. Draw dot and cross diagrams to represent the following simple covalent molecules:
Hydrogen
Hydrogen chloride
Water
20
Methane
Oxygen
Carbon dioxide
21
2. Complete the table to identify the features of substances with different types of bonding.
3. Explain why ionic substance have relatively high melting and boiling points. The strong electrostatic forces of attraction between oppositely charged ions with the giant ionic lattice structure require large amounts of energy to be overcome and separate them.
4. Explain why simple molecular covalent substances have relatively low melting and boiling points.
The weak intermolecular forces between molecules require little energy to be overcome and thus separate the molecules.
5. Explain why giant covalent molecules have relatively high melting and boiling points.
Strong covalent bonds require large amounts of energy to be overcome and separate them.
6. Explain why metallic substances have relatively high melting and boiling points. The strong electrostatic forces of attraction between the positive metal ions and the sea of delocalised electrons require large amounts of energy to be overcome.
7. Explain why ionic substances can only conduct electricity when aqueous (in solution) or molten.
When solid, the oppositely charged ions are fixed in the regular lattice structure and so cannot move in order to carry a charge. When the compound is in solution or molten, the ions break free from the lattice structure and become mobile; this means they can now carry a charge.
8. Explain why covalent substances are poor conductors of electricity. They do not have mobile electrons or ions. No charge carriers mean that a current cannot pass through the substance.
9. Explain why metallic substances can always conduct electricity.
The sea of negatively charged, delocalised electrons that surround the positive metal ions are able to move to form an electric current.
Relative
melting and boiling point
Relative solubility in water
Ability to conduct
electricity as a solid
Ability to conduct electricity in solution (or
molten)
Ionic bonding High Soluble No Yes
Simple molecular covalent
Low Insoluble No No
Giant covalent
High Insoluble No No (except graphite)
Metallic High Insoluble Yes Yes
22
10. Use the data below to identify the type of bonding in each substance. Explain your reasons for each.
Sodium chloride – ionic lattice because it has a relatively high boiling and melting points, does not conduct electricity when solid but does conduct electricity when in solution. These are all properties of ionic bonding.
Magnesium sulfate – ionic lattice because it has a high melting point, does not conduct electricity when solid but does conduct electricity when in solution. These are all properties of ionic bonding.
Hexane – simple molecular covalent because it has relatively low melting and boiling points and does not conduct electricity. These are both properties of simple molecular covalent bonding.
Liquid paraffin – simple molecular covalent because it has relatively low melting and boiling points and does not conduct electricity. These are both properties of simple molecular covalent bonding.
Silicon (IV) oxide – giant molecular covalent because it does not conduct electricity at all but does have high melting and boiling points. These are both properties of giant molecular covalent bonding.
Copper sulfate – ionic lattice, it does not conduct electricity when solid but does when in solution. These is a property of ionic bonding.
Sucrose – simple molecular covalent because it has a low melting point and does not conduct electricity when solid or in solution. These are both properties of simple molecular covalent bonding.
11. Graphite and diamond are different forms (allotropes) of which element? Carbon
12. Identify the structure and bonding of graphite and diamond.
They are giant covalent structures 13. Describe the structure of diamond.
Diamond is a form of carbon in which each carbon atom is joined to four other carbon atoms, forming a giant, three-dimensional, covalent lattice structure.
14. Describe the structure of graphite
Graphite is a form of carbon in which the atoms are arranged in a giant, two-dimensional layered structure. Each carbon atom is bonded to three others. Each individual layer is called graphene.
15. Identify a use of diamond.
Diamond is used in cutting tools. 16. Identify two uses of graphite
Graphite is use to make electrodes and it is also used as a lubricant. 17. Explain why graphite is used to make electrodes and as a lubricant.
23
Graphite is able to conduct electricity because carbon atom is only bonded to three others (rather than four), this means that there are ‘spare’ electrons that become delocalised and can carry a current across the structure. This makes is suitable for use as an electrode. Graphite’s layered structure means that the layers can easily slide off of each other. The carbon atoms are only bonded in two dimensions. This makes graphite useful as a lubricant.
18. Explain why diamond is used in cutting tools.
Diamond is very hard due to strong covalent bonds operating in three dimensions. Each carbon atom within the structure is bonded to four other carbon atoms.
19. C60 is an example of which group of substances?
Fullerenes 20. Describe and explain the properties of fullerenes.
Fullerenes are molecules of carbon that are shaped as tubes or hollow balls. They often contain carbon atoms arranged in shapes such as hexagons or pentagons. Fullerenes can be used to keep other atoms inside it’s ‘cage’ which may have useful medical applications. Fullerenes also have a very large surface area which could make them effective catalysts. Nanotubes are an example of a fullerene; they are extremely small cylinders of graphene. They conduct electricity and have a very high tensile strength.
21. Using poly(ethene) as an example, describe the structure of a polymer. Polymers are made up of long chains of repeating units called monomers. Monomers are made of chains of carbon atoms covalently bonded to one another. For example, poly(ethene) is the polymer made when ethene, C2H4, is polymerised.
22. Describe and explain the physical properties of metals.
Metals are very strong. This is because of the strong electrostatic forces of attraction between the metal ions and the delocalised electrons within their structure. These forces require a lot of energy to be broken. Metals typically have very high melting and boiling points for the same reason as above. Metals are usually shiny because they reflect a lot of light. Metals are generally dense which means their ions are packed closely together in the metal lattice. Metal are malleable; this means they can be shaped without breaking. The reason metals can do this is because the layers of ions within their lattice structure can slide over one another. Metals are also good conductors of heat and electricity. This is due to the delocalised electrons in their structure.
23. Describe the bonding in metals.
Metals have a giant structure, within this structure are positive metal ions surrounded by a sea of delocalised electrons.
24. Identify four types of representations used to model simple molecules Dot and cross diagrams Ball and stick diagrams Two-dimensional diagrams Three-dimensional diagrams
24
Calculations Involving Mass
The calculations in this unit will require lots of practice. Ask your teacher for more examples.
1. Relative formula mass can be calculated using which values? The relative atomic mass numbers of atoms on the periodic table 2. Calculate the relative formula mass of the following a) Sodium bromide, NaBr b) Water, H2O c) Magnesium nitrate, Mg(NO3)2 NaBr = 103 g mol-1 H2O = 18 g mol-1 Mg(NO3)2 = 256 g mol-1 3. Define the term empirical formula The simplest whole-number ratio or atoms of each element in a compound. 4. State the four stages of calculating the empirical formula of a compound. 1. Find experimental masses M 2. Find Ar values A 3. Divide masses by Ar D 4. Find the ratio R 5. Calculating empirical formula – example calculation: 3.2g of sulfur reacts with oxygen to produce 6.4 g of sulfur oxide. What is the formula of the oxide?
6. State the empirical formula of these compounds from their molecular formula.
1. H2O2 2. C4H8 3. Na2O
HO, CH2, Na2O
25
7. Explain how you can deduce the molecular formula of a compound from its empirical formula and its relative molecular mass. 1. Calculate the mass of one empirical formula ‘unit’. 2. Divide the relative molecular mass by the ‘unit’. 3. Multiply the number of each atom by the factor given in part 2.
11. A compound has the empirical formula CH2, it has a relative molecular mass (Mr) of 56. What is it’s
molecular formula? 1. Calculate the mass of one empirical formula ‘unit’.
12 + (1 x 2) = 14
2. Divide the relative molecular mass by the ‘unit’. 56 / 14 = 4
3. C x 4 = C4
H2 x 4 = H8
8. Describe an experiment to determine the empirical formula of a simple compound such as magnesium oxide. 1. Add a strip of magnesium to a clean crucible. Weight it and record its mass. 2. Heat the crucible. Put the lid on the crucible but allow oxygen in by lifting the lid regularly. 3. Heat strongly until all of the magnesium has turned white. 4. Allow the crucible to cool and re-weigh it. 5. Continue to heat, cool and weigh until the crucible has reached a constant mass. This is called heating
to constant mass. 6. The mass of the crucible can now be used to calculate the mass of magnesium oxide inside it. 7. From this, it is possible to calculate how much oxygen reacted and thus work out the empirical
formula of the compound.
9. Use two examples to explain the law of conservation of mass. a. In a precipitation reaction with a closed system, new products can be seen to be formed but the overall
mass of the system remains constant. b. When dissolving a known mass into a solution it can be seen that the solution increases in mass by
exactly the mass of the added solute.
10. Find the theoretical mass of magnesium oxide (MgO) that can be obtained by reacting 60g of magnesium with oxygen.
11. Recall the formula for calculating concentration Concentration (g dm-3) = mass (g) / volume (dm3)
12. How do you convert a number from cm3 to dm3? Divide (÷) by 1000
26
13. [H] Define the term mole A number of particles (6.02 x 1023) equal to the number of particles in 12g of carbon-12.
14. [H] State Avogadro’s constant 6.02 1023
15. [H] State the two equations involved in mole and mass calculations
Moles = mass/Ar or Mr Number of particles = moles x Avogadro’s constant (6.02 x 1023)
16. [H] Calculate the moles of carbon in 18g Moles = mass/Ar Moles = 18g / 12 = 1.5 mol
17. [H] Calculate the number of particles in 2 mol of sodium chloride, NaCl. Number of particles = moles x Avogadro’s constant Number of particle = 2 x 6.02 x 1023
= 1.204 x 1024
18. [H] Calculate the number of particles in 3g of magnesium chloride, MgCl2
Find the moles: Moles = mass/Mr Moles = 3g / 95 Moles = 0.032 moles Number of particles = Moles x Avogadro’s constant Number of particles = 0.32 x 6.02 x 1023 Number of particles = 1.93 x 1023
19. [H] What is meant by the term ‘stoichiometry’?
It is the relationship between the molar quantities in a chemical reaction, as shown by the balanced symbol equation. ***** If you are aiming for an 8 or a 9 grade, there are a number of more technical calculations that are required in this part of the course. Ask your teacher for lots of practice questions! *****
27
Topic 6 – Groups in The Periodic Table
1. Group 1 elements are also known by what name?
The alkali metals 2. Group 7 elements are also known by what name?
The halogens 3. Group 0 elements are also known by what name?
The noble gases
4. Describe two unique features of alkali metals They are soft metals with comparatively low melting points
5. Which two products are made when group 1 metals react with water?
A metal hydroxide (an alkali) + hydrogen gas 6. Complete the following word equations:
Lithium + water -> Sodium + water -> Potassium + water -> Lithium + water -> lithium hydroxide + hydrogen Sodium + water -> sodium hydroxide + hydrogen Potassium + water -> potassium hydroxide + hydrogen
7. Write a balanced equation, including state symbols, for the reaction between sodium and water. 8. Describe the pattern in reactivity of the alkali metals with water.
As you go down group 1 the metals become more reactive. This means that they react more vigorously with water, for example lithium will fizz and move slowly on the surface of water whereas potassium burns, may explodes and moves quickly across the surface of water.
9. Explain the trend reactivity of alkali metals with water.
Group 1 metals reacts by losing the one electron in their outer shell. As you move down the group (increase atomic number) the electron is more easily lost which means the element is more reactive. The electron is more easily lost because down the group there are more electron shells in each atom which leads to a greater electron shielding effect. This means there is a weaker attraction between the outer electron and the positive nucleus.
10. Recall the colour and physical states of the halogens (group 7) at room temperature.
Chlorine is a green gas Bromine is a red/brown liquid Iodine is a grey solid
28
11. Halogens are diatomic, explain what this means.
It means that they are found as molecules with the formula X2 (e.g. F2, Cl2)
12. Describe the trend in melting and boiling points as you go down group 7. As you go down group 7, the melting and boiling points increase.
13. Describe the chemical test for chlorine. Chlorine gas will bleach damp, blue litmus paper.
14. Halogens react with metals to form which compounds? Metal halides
15. Complete the following word equations as examples of halogen + metal reactions.
Chlorine + sodium -> Fluorine + potassium ->
Chlorine + sodium -> sodium chloride Fluorine + potassium -> potassium fluoride
16. Halogens react with hydrogen to produce which type of compound?
Hydrogen halides 17. Hydrogen halides will dissolve in water to form what type of solution?
An acidic solution 18. What type of chemical reaction can show the relative reactivities of the halogens?
Displacement
19. Describe, using examples, how displacement reactions can be used to show the relative reactivities of the halogens. A displacement reaction is one where a more reactive element displaces (replaces) another, less reactive element. Halides dissolved in water (e.g. potassium iodide) can be tested with halogens (e.g. chlorine), if the halogen reacts with the halide then this demonstrates that it is more reactive then the halogen in the halide compound (e.g. chlorine will displace the iodide ion in potassium iodide). An equation to demonstrate this is:
Cl2(g) + 2KI(aq) -> I2(aq) + 2KCl(aq)
If there is no reaction then the halogen is not more reactive than the halogen in the halide compound (e.g. iodine will not react with potassium bromide because bromine is more reactive than iodine).
20. [H] Explain why halogen displacement reactions are redox reactions. Use equations in you answer.
Halogen displacement reactions are redox reactions because both oxidation and reduction occur at the same time. For example, in the reaction between bromine and potassium iodide, the bromine atoms gain electrons to become bromine ions (this is reduction). The iodide ions lose electrons to become iodine atoms (this is oxidation). Overall: Br2 + 2KI 2KBr + I2
Br2 + 2e- 2Br-
2I- I2 + 2e-
29
21. Explain the trend in reactivity of the halogens.
- As you go down group 7, the halogens get less reactive. - The reason for this is as the atoms get bigger (down the group) the atom has more electron
shells. - This means there is a greater shielding effect reducing the nuclear attraction between the
nucleus and the reacting electrons. - There is also a greater distance between the nucleus and the outer electrons, further reducing
the attractive ‘power’ of the nucleus. - The atoms lower down in the group are much less effective at attracting new electrons which
means that they are less reactive. 22. Describe the reactivity of the noble gases
Compared to other elements, the noble gases are very unreactive, they are said to be inert. 23. Explain why noble gases are inert.
Noble gases are inert (unreactive) because they have a full outer shell of electrons meaning that they will not easily gain or lose electrons (which is needed for a chemical reaction).
24. Describe four uses of noble gases
Argon is used in filament lamps and welding Helium is used in airships and part balloons
25. Explain how the properties of the noble gases make them suitable for their uses.
Argon is used in filament lamps and welding because it provides an inert atmosphere that stops reactions with oxygen. Helium is used in airships and part balloons because it has a lower density than air. This means that objects filled with helium will float in air.
26. Give two examples of properties that show a trend within the group of noble gases
Boiling point and density.
30
Topic 7 – Rates of Reaction and Energy Changes
1. Identify four types of chemical reactions where you can measure the temperature change. Salts dissolving in water Neutralisation reactions Displacement reactions Precipitation reactions
2. Describe a method for measuring the temperature (and thus energy change) accompanying a
chemical reaction. 1. Measure the temperature of the reactants 2. Mix the reactants inside a polystyrene cup (to reduce energy loss, the polystyrene is a heat
insulator) 3. Placing the polystyrene cup in a beaker of cotton wool further prevents heat loss 4. Putting a lid on the cup will also reduce energy loss by evaporation 5. Measure the final temperature of the solution at the end of the reaction
3. Define exothermic reaction and give two examples.
A reaction in which heat energy is given out, for example in combustion reactions or explosions 4. Define endothermic reaction and give two examples.
A reaction in which heat energy is taken in, for example photosynthesis or dissolving ammonium nitrate in water
5. Complete the following sentences using either exothermic or endothermic:
‘Bond breaking is ___________________’ ‘Bond making is __________________’ ‘Bond breaking is endothermic’ ‘Bond making is exothermic’
6. Explain, using ideas about bond breaking and making, why the overall heat energy change for a
reaction may be exothermic. A reaction will be exothermic if more heat energy is released making bonds in the products than is required to break bonds in the reactants.
7. Explain, using ideas about bond breaking and making, why the overall heat energy change for a
reaction may be endothermic. A reaction will be endothermic if less heat energy is released making bonds in the products than is required to break bonds in the reactants.
31
8. Draw an energy profile diagram to show an exothermic reaction
9. Draw an energy profile diagram to show an endothermic reaction
10. Some reactions are slower and some are faster, we say they have different what?
Rates of reaction
11. Give an example of a reaction with a slow rate and an example of a reaction with a fast rate. Rusting has a slow rate, explosion have a fast rate.
12. Which four factors may affect the rate of a chemical reaction? Temperature Surface area of a solid Concentration (pressure for gases) The presence of a catalyst
13. Describe and explain the effect of an increased temperature on the rate of a typical reaction.
Increasing temperature increases the rate of a reaction. This is because the particles have increased energy meaning they will move faster; this leads to a higher frequency of successful collisions between particles. In addition, more particles have energy greater than the activation energy which also increases the frequency of successful collisions.
32
14. Describe and explain the effect of increasing concentration on the rate of a chemical reaction. Increasing concentration increases the rate of a reaction. This is because there are more reactant particles within the same volume, leading to a higher frequency of successful collisions between particles.
15. Describe and explain the effect of increasing surface area of a solid reactant on the rate of a chemical
reaction. Increasing the surface area of a solid reactant will increase the rate of a chemical reaction. This is because there will be more reactant particles available to make contact with one another thus leasing to a higher frequency of successful collisions.
16. Describe and explain the effect of a suitable catalyst on the rate of a chemical reaction.
A catalyst will increase the rate of a chemical reaction by lowering the activation energy.
17. Explain how a catalytic convertor in a car works Catalytic convertors have a high surface area to increase the rate of reaction between carbon monoxide and unburnt fuel from exhaust gases with oxygen from the air to produce carbon dioxide and water. They work best at high temperatures.
18. Describe an experiment that could be completed in order to investigate the effect of temperature of
the rate of a reaction. The reaction between sodium thiosulfate and hydrochloric acid can be used. When reacted together a yellow precipitate of sulfur forms, the quicker this forms the faster the rate of the reaction. Using the same volumes of sodium thiosulfate and hydrochloric acid for each experiment (a control variable), the starting temperatures of the solutions can be varied (independent variable) and the rate of reaction can be measured by recording the time taken (dependent variable) for a black cross underneath the reaction flask to ‘disappear’.
33
19. Describe an experiment that could be completed in order to investigate the effect of surface area (of
a solid) on the rate of a reaction. The reaction between hydrochloric acid and marble chips (calcium carbonate) can be used. Using the same volume and concentration of acid for each experiment and using the same mass of calcium carbonate (both control variables), the two reactants can be placed in a conical and the amount of gas released from the reaction can be measured using a gas syringe. You can either measure the amount of has given off in a certain time, or how much gas is released every 10 seconds for a determined period of time (dependent variable). For each experiment, the surface area of the marble chips can be changed (e.g. powdered marble chips rather than large chunks).
20. Describe an experiment that could be completed in order to investigate the effect of concentration
on the rate of a reaction. The reaction between hydrochloric acid and marble chips (calcium carbonate) can be used. Using the same volume of acid for each experiment and using the same mass and surface area of calcium carbonate (both control variables), the two reactants can be placed in a conical and the amount of gas released from the reaction can be measured using a gas syringe. You can either measure the amount of has given off in a certain time, or how much gas is released every 10 seconds for a determined period of time (dependent variable). For each experiment, the concentration of the acid can be changed.
34
Topic 8 – Fuels and Earth Science
1. Define the term hydrocarbon
A hydrocarbon is a compound that contains carbon and hydrogen atoms only
2. Describe four features of crude oil. - It is a complex mixture of hydrocarbons - It contains molecules in which carbon atoms are in chains or rings - It is an important source of useful substances (e.g. fuels and feedstock for the petrochemical
industry) - It is a finite resource
3. Describe the process by which crude oil is separated.
Crude oil is separated into simpler, more useful mixtures called ‘fractions’, this is done using fractional distillation. Crude oil is put in to a fractionating column where it is heated. The molecules begin to evaporate and rise up the column. As they rise up the fractionating column the molecules cool until they reach their condensing point. When they condense they are drained off and kept as a separate fraction.
4. Recall the names and uses of the fractions in crude oil Gases – used for domestic heating and cooking Petrol – used as a fuel for cars Kerosene – used as fuel for aircraft Diesel oil – used as fuel for some cars and trains Fuel oil – used as fuel for large ships and in some power stations Bitumen – used to surface roads and roofs
5. Identify and explain four ways in which the hydrocarbons in different fractions of crude oil differ from one another. - The number of carbon and hydrogen atoms their molecules contain. As you go down the
fractionating column, the number of carbon and hydrogen atoms in the molecules increases. - The boiling points of the molecules. As you go down the fractionating column the boiling points
of the molecules increases. This is because larger molecules have greater intermolecular forces which require more energy to overcome.
- The ease of ignition. As you go down the fractionating column the ease of ignition of the fractions decreases.
- Viscosity. As you go down the fractionating column the viscosity of the molecules increases.
6. Most of the hydrocarbons found in crude oil are members of which homologous series? The alkanes
7. Explain four features of a homologous series. - They contain compounds with the same general formula (e.g. CnH2n+2 for alkanes) - They differ by CH2 in molecular formula from neighbouring compounds - They show a gradual variation in physical properties (e.g. boiling points) - They have similar chemical properties
35
8. Explain the term complete combustion Complete combustion is a chemical reaction in which a fuel reacts with oxygen in order to produce carbon dioxide and water. Energy is also given out during this reaction. It is called complete combustion when there is plenty of oxygen available to react with the fuel.
9. Describe the term incomplete combustion Incomplete combustion of hydrocarbons happens when there is little or no oxygen available to react with the fuel. The products of incomplete combustion can be either carbon monoxide or carbon as well as water.
10. Explain how carbon monoxide behaves as a toxic gas. When breathed in carbon monoxide binds to erythrocytes (red blood cells). When it does this, it stops oxygen from binding with the red blood cells and prevents the formation of oxyhaemoglobin which is a molecule used to transport oxygen to respiring cells around the body. Instead, carboxyhaemoglobin is formed. The lack of oxygen leads to cells being unable to respire and release energy for normal body processes to take place. This can eventually lead to death.
11. Describe other problems associated with incomplete combustion (not carbon monoxide as a toxic gas). Incomplete combustion produces soot, solid carbon that causes buildings and surfaces to look dirty and may cause breathing difficulties in inhaled, particularly for those with respiratory disease.
12. Explain how sulfur dioxide is produced
Sulfur dioxide is produced when fossil fuels that contain sulfur impurities are burned in power stations. The oxygen in the air reacts with the sulfur to produce sulfur dioxide. Sulfur + oxygen -> sulfur dioxide This sulfur dioxide then leaves the power station and enters the atmosphere.
13. Explain what causes acid rain Sulfur dioxide in the atmosphere dissolves in the water in clouds to form a mixture of acids, including sulfurous acid: H2O(l) + SO2(g) H2SO3 (aq) Sulfurous acid is then oxidised by oxygen in the air to form sulfuric acid: 2H2SO3(aq) + O2(g) 2H2SO4(aq) The mixture of sulphurous acid and sulfuric acid then falls in the rain water as acid rain.
14. Explain some problems caused by acid rain - It increases the rate of weathering of buildings and statues made with limestone - It increases the rate of corrosion of metals - It acidifies lakes and rivers which can kill fish and insects and stop their eggs from hatching - It can acidify soil which stop crops growing well
15. Explain why nitrogen oxides are produced by cars
Nitrogen oxides (NOx) are formed when oxygen and nitrogen from the air react together in the high temperatures of a car engine.
16. Describe some problems caused by nitrogen oxides in the air. Nitrogen oxides are pollutants that can: - Dissolve in water and contribute to acid rain - In high enough concentrations, nitrogen dioxide is a toxic gas that can cause respiratory disease
36
17. Evaluate the advantages and disadvantages of using hydrogen, rather than petrol, as a fuel for cars. Advantages - The combustion of hydrogen only produces water, no pollutant gases such as carbon dioxide
which are greenhouse gases linked to global warming and climate change. - Hydrogen can be obtained from water which is a renewable resource
Disadvantages
- The engines are very expensive - Producing hydrogen gas often requires using energy produced by burning fossil fuels that
release carbon dioxide - Hydrogen is difficult to store as it is explosive and comes in gaseous form - Hydrogen is not yet widely available, so the range of hydrogen vehicles may be low
18. Recall three non-renewable fossil fuels obtained from crude oil
Petrol, kerosene and diesel oil
19. Recall one non-renewable fossil fuel obtained from natural gas Methane
20. Explain the process of cracking Cracking is a process in which long-chained alkenes are broken down by heat into short-chained alkanes and alkenes. This involves passed vaporised hydrocarbons over a powdered catalyst at very high temperatures and pressures.
21. Explain why cracking is necessary Cracking is necessary in order to match the supply and demand of hydrocarbon molecules. Shorter-chained molecules such as petrol are in short supply but in high demand (because they are very useful). In order to balance the supply and demand, long-chained molecules (which are in lower supply due to being less useful) are used to make more of the short-chained molecules.
37
Topic 9 - Earth and Atmospheric Science
1. The gases that formed the Earth’s early atmosphere were produced by what?
Volcanic eruptions
2. Describe the composition of Earth’s early atmosphere - Large amounts of carbon dioxide, CO2 - Some water vapour - Very little or no oxygen - Small amounts of other gases
3. Explain how the oceans formed
As the Earth cooled, the water vapour condensed to form oceans.
4. Explain why the amount of carbon dioxide in the early atmosphere decreased The amount of carbon dioxide in the atmosphere decreased when carbon dioxide dissolved in the oceans as they formed.
5. Explain why the amount of oxygen in the atmosphere increased over time The growth of primitive plants that used carbon dioxide and released oxygen through photosynthesis increased the amount of oxygen in the atmosphere.
6. Describe the chemical test for oxygen Oxygen re-lights a glowing splint
7. Name three greenhouse gases Methane Carbon dioxide Water vapour
8. Describe the greenhouse effect A process by which heat radiation from the Sun in absorbed by the Earth and re-emitted. This re-emitted heat radiation is absorbed by the bonds in greenhouse gas molecules which absorb and re-emit the heat radiation back towards Earth. This process keeps the Earth warm but increases as there are more greenhouse gases in the atmosphere.
9. Evaluate the evidence for human activity as the cause of climate change. - There is a correlation between the change in atmospheric carbon dioxide concentration and the
consumption of fossil fuels. - There is a correlation between the change in atmospheric carbon dioxide concentration and the
Earth’s temperature - Many scientists believe that these are not just correlations but that there is casual link between
these factors. - There are uncertainties in the data and historical trends caused by the location where these
measurements were taken and their varying accuracy.
38
10. Describe the composition of today’s atmosphere 78% Nitrogen, N2 21% Oxygen, O2 1% Argon, Ar 0.04 % Carbon dioxde, CO2 And other gases
11. Describe the potential effects of increased levels of carbon dioxide and methane generated by human activity Carbon dioxide
- The human population is increasing which means there are more people respiring and releasing carbon dioxide in to the atmosphere.
- An increase in population also means that there is higher demand for electricity for lighting, heating and cooking. This electricity is often generated using fossil fuels which released carbon dioxide when they are burnt.
- More land is being used to build houses and grow food. This means that trees are being chopped in deforestation. This means that there are fewer trees on the planet to absorb carbon dioxide through photosynthesis.
- Carbon dioxide is produced in volcanic eruptions (not caused by human activity)
Methane - Livestock farming (particularly cattle) produces large amounts of methane. - Soil bacteria in landfill sites produce methane - Soil bacteria in rice ‘paddy’ fields also produce methane
12. Explain how the effects of climate change may be mitigated (this means limited or prevented)
- Using renewable resources reduces greenhouse gas emissions - Burning fewer fossil fuels - There are some new technologies that suggest we might be able to ‘capture’ the CO2 released in
reactions and bury it in rocks underground.
