Chemistry Paper 1 Flashcards
1
Q
What are all substances made up of?
A
Atoms.
2
Q
Define atom.
A
The smallest part of an element that can exist.
3
Q
What is the chemical symbol for sodium?
A
Na
4
Q
What is the chemical symbol for chlorine?
A
Cl
5
Q
What is the chemical symbol for hydrogen?
A
H
6
Q
What is the chemical symbol for iron?
A
Fe
7
Q
Roughly how many different elements are there?
A
100
8
Q
In which table can we find the elements?
A
The periodic table.
9
Q
What type of reactions form compounds from elements?
A
Chemical reactions.
10
Q
Define compound.
A
A substance made from two or more different elements that have been chemically joined together.
11
Q
We can write reactions as word equations and___________ equations.
A
Symbol.
12
Q
Write the chemical formula for a hydroxide ion.
A
OH-
13
Q
Write the chemical formula for a nitrate ion
A
NO3-
14
Q
Write the chemical formula for a carbonate ion
A
CO32-
15
Q
Write the chemical formula for a sulphate ion
A
SO42-
16
Q
Define element
A
A substance made of only one type of atom
17
Q
Define mixture.
A
A substance made up of two or more elements or compounds not chemically joined together.
18
Q
When a mixture is formed, do the chemical properties of the substances used to make the mixture change?
A
No.
19
Q
What type of processes separate the components of a mixture?
A
Physical processes.
20
Q
State the five main separating techniques.
A
1) Filtration
2) Crystallisation
3) Simple distillation
4) Fractional distillation
5) Evaporation
21
Q
What does filtration involve?
A
It involves separating an insoluble solid from a liquid.
22
Q
During evaporation, what is evaporated, the solvent or the solute?
A
The solvent.
23
Q
When salty water is heated and the liquid evaporated, what is left behind?
A
Crystals of salt.
24
Q
Three liquids with very different boiling points are present in a mixture. Which separation technique would you use to separate them?
A
Simple distillation.
25
When would fractional distillation be used instead of simple distillation?
When the liquids making up a mixture have similar boiling points.
26
What did John Dalton describe atoms as?
Round spheres that could not be divided.
27
Describe J J Thomson's plum pudding model of the atom.
The atom is a ball of positive charge with negative electrons embedded in it.
28
What was the name of the experiments carried out by Rutherford and Marsden?
Alpha particle scattering experiments
29
What did Rutherford and Marsden fire alpha particles at?
A thin sheet of gold foil.
30
What did Rutherford and Marsden observe from their experiments that surprised them?
Most alpha particles went straight through the foil. But a few were scattered in different directions. Some even deflected backwards
31
What did Rutherford and Marsden conclude about the structure of the atom from their experiments?
That the mass of an atom was concentrated at the centre (nucleus), and that the nucleus was positively charged. Most of the atom is empty space.
32
Who adapted the nuclear model by suggesting that electrons orbit the nucleus at specific distances?
Niels Bohr.
33
The experimental work of James Chadwick provided the evidence for which subatomic particle?
The neutron.
34
What is the relative charge of a proton?
+1
35
What is the relative charge of an electron?
-1
36
What is the relative charge of a neutron?
0
37
Why do atoms have no overall charge?
1) Protons are +1 and electrons have -1
| 2) Atoms have the same number of protons and electrons so the charges cancel out
38
Do atoms of different elements have the same or different number of protons?
Different
39
Do atoms of the same elements have the same or different number of protons?
Same
40
What does the atomic number of an element tell us?
The number of protons in an atom of that element.
41
How many protons does Mg have?
12
42
Have many protons does Cl have?
17
43
Roughly what is the radius of an atom?
1 x 10-10 m
44
Roughly what is the radius of the nucleus of an atom?
1 x 10-14 m
45
Almost all of the mass of an atom is situated where?
The nucleus.
46
What is the relative mass of an electron?
Very small.
47
What is the relative mass of a neutron?
1
48
What is the relative mass of a proton?
1
49
What does the mass number of atom tell us?
The total number of protons and neutrons in the atom.
50
Atoms of the same element can have different number of neutrons. What are they called?
Isotopes.
51
How many neutrons are there in an atom on Na?
12
52
How many protons are there in an atom of Na?
11
53
How many electrons are there in an atom of Na?
11
54
How many neutrons are there in an atom on F?
10
55
How many protons are there in an atom of F?
9
56
How many electrons are there in an atom of F?
9
57
Cu - 63 has an abundance of 69.2% and Cu-65 has an abundance of 30.8%. Calculate the relative atomic mass of copper to 1 decimal place.
63.6
58
Cl-35 has an abundance of 75% and Cl-37 has an abundance of 25%. Calculate the relative atomic mass of copper to 1 decimal place.
35.5
59
What is the maximum number of electrons that can go in the first electron shell?
2
60
What is the maximum number of electrons that can go in the second electron shell?
8
61
What is the maximum number of electrons that can go in the third electron shell?
8
62
Write the electronic structure for sodium.
2, 8, 1
63
Write the electronic structure for nitrogen.
2, 5
64
Write the electronic structure for silicon.
2, 8, 6
65
Write the electronic structure for neon.
2, 8
66
The elements in the periodic table are arranged in order of…
Atomic number
67
Elements in the same group have….
Similar properties
68
What are the vertical columns called?
Groups
69
What are the horizontal rows called?
Periods
70
Elements in the same group have a different number of electrons in their outer shell. True or False?
False.
71
Elements with the same number of outer electrons have similar chemical properties. True or False?
True.
72
Would you expect magnesium and nitrogen to have similar properties, why?
No, they are in different groups and have a different number of outer shell electrons.
73
Would you expect fluorine and chlorine to have similar properties, why?
Yes, because they are in the same group and have the same number of outer shell electrons.
74
Before the discovery of electrons, neutrons protons, scientists tried to classify elements by arranging them in order of what?
Their atomic weight.
75
State one problems of the early periodic tables.
Some elements were placed in inappropriate groups.
76
Which scientist overcame some of the problems of the early periodic tables?
Mendeleev.
77
State two changes Mendeleev made to overcome some of the problems of the early periodic tables.
1) He left gaps for elements that he thought had not been discovered
2) In some places he also changed the order based on atomic weights (swap places)
78
State two discoveries that supported Mendeleev's changes to the early periodic tables.
1) Elements with properties predicted by Mendeleev were discovered and filled the gaps.
2) The discovery of isotopes made it possible to explain why the order based on atomic weights was not always correct.
79
What name is given to elements that react to form positive ions?
Metals
80
What name is given to elements that do not form positive ions?
Non-metals
81
Are the majority of elements metal or non-metals?
Metals
82
Where are metals found on the periodic table?
To the left and towards the bottom of the periodic table
83
Where are non-metals found on the periodic table?
Towards the right and top of the periodic table
84
State three physical properties of metals
1) Strong/hard
2) Malleable
3) Good electrical and thermal conductors
5) Shiny
4) High melting and boiling points
85
State two physical properties of non-metals
1) Tend to be dull looking
2) normally brittle
3) Don't generally conduct electricity
86
What name is given to elements in group 0?
The noble gases
87
The elements in group 0 are inert, what does this mean?
They are unreactive.
88
Why are the elements in group 0 inert?
They have a stable arrangement of electrons.
89
How many outer shell electrons do most of the elements in group 0, what is the exception?
8, the exception is helium, it has 2.
90
How does the boiling point change as you go down group 0?
It increases.
91
Why does the boiling point change as you go down group 0?
Increase in the number of electrons so greater intermolecular forces
92
What are the elements in group 1 known as?
The alkali metals
93
Why do group 1 metals have similar properties?
The all have one electron in their outer shell
94
How are alkali metals different to other metals?
They are soft and have a low density
95
How does the reactivity change as you go down group 1?
The reactivity increases
96
How does the boiling melting point change as you go down group 1?
It decreases
97
Name the products when sodium reacts with water.
sodium hydroxide and hydrogen
98
Write the chemical formula of the products formed when potassium reacts with water.
KOH and H2
99
Write a word equation for the reaction between lithium and oxygen.
lithium + oxygen -> lithium oxide
100
Explain why potassium is more reactive than lithium (3 marks)
More shells/electrons, less shielding (or weaker attraction from nucleus), easier to lose electrons
101
What are the elements in group 7 known as?
The halogens
102
Why do elements in group 7 have similar reactions?
They all have 7 electrons in their outer shell.
103
Are group 7 elements metals or non-metals?
Non-metals
104
The halogens consist of molecules made up of pairs of atoms. True or False?
True.
105
What state is fluorine at room temperature?
Gas
106
What state is chlorine at room temperature?
Gas
107
What state is bromine at room temperature?
liquid
108
What state is iodine at room temperature?
solid
109
What is a displacement reaction?
A reaction in which a more reactive element takes the place of a less reactive element in a compound
110
Explain why the following reaction does not proceed: KBr + I₂
Iodine is less reactive than bromine so cannot displace it
111
Balance the equation and explain why it is a displacement reaction: KBr + Cl₂ → KCl + Br₂
2KBr + Cl₂ → 2KCl + Br₂ , chlorine has displaced bromine as it is more reactive
112
Explain why fluorine is more reactive than chlorine
Fewer shells/electrons, less shielding (or stronger attraction from nucleus), easier to gain electrons
113
How do the boiling and melting point change as you go down group 7?
The boiling and melting point increase
114
How does reactivity change as you go down group 7?
The reactivity decreases
115
State the three types of chemical bonds.
Ionic, covalent and metallic.
116
Fill in the blank. For _______ bonding the particles are oppositely charged ions.
Ionic
117
Fill in the blank. For ________ bonding the particles are atoms which share pairs of electrons.
Covalent
118
Fill in the blank. For _________ bonding the particles are atoms which share delocalised electrons.
Metallic
119
What type of elements does ionic bonding occur between?
Metals and non-metals
120
What type of elements does covalent bonding occur between?
Non-metals
121
Define ionic bonding.
The electrostatic force of attraction between positive metal ions and negative non-metal ions.
122
Do metal atoms for positive or negative ions?
positive
123
Do non-metal atoms form positive or negative ions?
negative
124
Na has an electronic structure of 2, 8, 1. State the electronic structure of a sodium ion
2, 8
125
State the charge for the ions of group 1
+1
126
State the charge for the ions of group 2
+2
127
State the charge for the ions of group 7
-1
128
State the charge for the ions of group 7
-2
129
What type of bonding would occur between group 1 and group 7 ions?
Ionic bonding
130
State the charge of the ions present in MgCl2
Mg2+ and Cl-
131
Cl has an electronic structure of 2, 8, 7. State the electronic structure of a sodium ion
2, 8
132
What type of forces hold ions in an ionic compounds together?
Electrostatic forces of attraction
133
An ionic compound is a giant structure of ions? True or false?
True.
134
What name is given to structures formed by ionic bonding?
Giant ionic lattice
135
Describe the direction of the forces in an ionic structure.
In all directions
136
How many shared pair of electrons are there in one covalent bond?
1
137
How many electrons are there in one covalent bond?
2
138
How many electrons are there in 2 covalent bonds?
4
139
When atoms share electrons what type of bonds do they form?
Covalent bonds
140
What type of bonding is present in a molecule of NH3
Covalent
141
How many covalent bonds can carbon form?
4
142
How many covalent bonds are there in CH4?
4
143
Covalently bonded substances never consist of small molecules. True or False?
False.
144
Polymers are examples of very large molecules that have covalent bonds. True or False?
True.
145
Describe the structure of metals.
Consist of metal ions arranged in a regular pattern in a sea of delocalised electrons.
146
What does delocalised mean?
Free to move.
147
Which electrons of a metal are delocalised? (hint - shells)
The electrons in the outer shell
148
What are the three states of matter?
Solid, liquid and gas
149
What two physical processes take place at the melting point?
Melting and freezing
150
What two physical processes take place at the boiling point?
Condensation and boiling
151
Describe the layout of particles in a solid.
They are tightly packed in neat rows.
152
What happens to the particles of a solid as it is heated?
The particles gain energy and start to vibrate more.
153
The amount of energy needed to change state depends on the strength of what?
Forces between particles
154
The stronger the forces between particles the _________ the melting and boiling point of a substance.
Higher
155
State a limitation of using spheres to represent particles in the particle model.
1) Particles aren't solid spheres
| 2) Particles aren't inelastic
156
Oxygen has a melting point of -219oC and a boiling point of -183oC. What state is oxygen at -200oC?
Liquid
157
Why is energy required for a liquid to boil into a gas?
To overcome the forces of attraction holding the particles in a liquid together.
158
Which state of matter does not have a fixed volume and shape?
Gas
159
Which state of matter has a fixed volume, but does not have a fixed shape?
Liquid
160
What state symbol is used for solids?
(s)
161
What state symbol is used for liquids?
(l)
162
What state symbol is used for gases?
(g)
163
What does the state symbol (aq) mean?
Aqueous solution (dissolved in water)
164
What type of structure do ionic compounds have?
Giant ionic lattice
165
Why do ionic compounds have a high melting and boiling point?
A large amount of energy is needed to break the many strong ionic bonds.
166
When melted or dissolved in water, ionic compounds conduct electricity, why?
The ions are free to move and they carry a charge.
167
What type of forces exist between the ions in an ionic lattice?
Electrostatic forces of attraction
168
The attractive forces between ions in an ionic lattice are in all directions. True or false?
True.
169
What forces acting between small molecules hold them together?
Weak intermolecular forces
170
Small molecules are usually liquids or gases at room temperature. True or False?
True.
171
When a small molecule melts or boils, which forces are overcome? Intermolecular forces or covalent bonds?
Intermolecular forces
172
What type of bonding occurs between the atoms of a small molecule?
Covalent
173
Why don't small molecules conduct electricity?
They have no overall charge.
174
Which is stronger, the forces between small molecules or the bonds between the atoms of small molecules?
The bonds between the atoms of small molecules.
175
Are polymers large or small molecules?
Large.
176
What type of bonding holds the atoms in a polymer together?
Covalent.
177
Why are polymers normally solid at room temperature, whereas small molecules are normally liquids or gases?
Polymers have stronger intermolecular forces between molecules. They require a larger amount of energy to overcome.
178
What is a polymer?
Many small molecules joined together in a chain to form a large molecule
179
What is the name given to the structure of diamond, graphite and silicon dioxide?
Giant covalent
180
What bonding exists between atoms in giant covalent structures?
Covalent
181
What state are giant covalent structures at room temperature?
Solid
182
Do giant covalent structures have a low or high melting point?
High
183
Give three examples of giant covalent structures
1) Diamond
2) Silicon dioxide (silica)
3) Graphite
184
What type of bonds must be overcome to melt or boil giant covalent structures?
Covalent
185
Explain why pure metals can be bent and shaped
Layers of metal ions are free to slide over each other
186
What is an alloy?
A mixture of two or more metals
187
Explain why alloys can be harder than pure metals
Different size of atoms distort the layers to stop them sliding over each other
188
What type of bonding exists in metals?
Metallic bonding
189
Why are metals good conductors of electricity?
They have delocalised electrons which carry a charge through the metal
190
Why are metals good thermal conductors?
Energy is transferred by the delocalised electrons
191
What type of structure is diamond?
Giant covalent
192
What type of bonding exists in diamond?
Covalent
193
How many bond does each carbon atom in diamond form?
4
194
Diamond is very hard. True or false?
True.
195
Diamond does not conduct electricity, why?
No free ions or electrons that carry a charge.
196
Why does diamond have a high melting point?
Giant structure, Strong covalent bonds between the atoms, requires a lot of energy to break
197
In graphite, how many covalent bonds does each carbon form?
3
198
There are covalent bonds between the layers of hexagonal rings in graphite. True or false?
False.
199
Explain why graphite conducts electricity.
There is one delocalised electron from each carbon atom, this carries a charge.
200
How is graphite similar to metals?
Has delocalised electrons
201
Why is graphite often used as a lubricant?
There are weak forces of attraction between the hexagonal layers, so they can slide over each other
202
What is graphene?
A single layer of graphite.
203
Why is graphene useful in electronics?
It has delocalised electrons so can conduct electricity
204
Why is graphene very strong?
Because of the network of strong covalent bonds between the carbon atoms.
205
What atoms is graphene made up of?
Carbon.
206
Why is graphene often added to composite materials?
Because it is strong and light, so can add strength without adding much weight
207
What are fullerenes?
They are molecules of carbon atoms with hollow shapes
208
The first fullerene to be discovered was Buckminsterfullerene (C60), what shape does it have?
Spherical
209
What are carbon nanotubes?
They are cylindrical fullerenes.
210
Carbon nanotubes have a very high length to diameter ratio. True or false?
True.
211
Why are nanotubes used in electronics?
They have delocalised electrons which carry a charge - conduct electricity
212
Suggest why carbon nanotubes are used to make tennis racket frames
Nanotubes are strong and very light, so can add strength to the tennis racket without adding much weight.
213
Which law states that no atoms are lost or made during a chemical reaction?
(The law of) conservation of mass
214
In a chemical reaction, the mass of the products sometimes equals the mass of the reactants. True or False?
False.
215
Why do we balance symbol equations?
So that the numbers of atoms of each element on both sides of the equation are the same.
216
What is the Mr of CO2?
12 + 16 + 16 = 44
217
What is the Mr of NaOH?
24 + 16 + 1 = 40
218
What is the Mr of Mg(OH)2?
58
219
What is the Mr of Na2CO3?
106
220
What % of H2O is O?
16/18 x 100 = 88.9%
221
What % of Na2CO3 is Na?
46/106 x 100 = 43%
222
Why do some reactions appear to involve an increase in mass?
At least one of the reactants is a gas and its mass has not been taken into account
223
Why do some reactions appear to involve a decrease in mass?
At least one of the products is a gas and its mass has not been taken into account
224
An unsealed container is placed on a weighing balance. The metal inside is reacted with oxygen. Would the scale show an increase or decrease in mass?
An increase in mass
225
An unsealed container is placed on a weighing balance. The metal carbonate inside is thermally decomposed to produce a metal oxide and carbon dioxide. Would the scale show an increase or decrease in mass?
A decrease in mass
226
Calculate the mean for the following numbers 24.0, 24.5, 23.5, 25.0, 23.0.
24
227
Calculate the range for the following numbers 24.0, 24.5, 23.5, 25.0, 23.0.
25 - 23 = 2
228
The following numbers are results from a scientific experiment, with a mean of 24cm, calculate the uncertainty of the mean: 24.0cm, 24.5cm, 23.5cm, 25.0cm, 23.0cm.
Range = 2cm therefore uncertainty = range /2…. 2/2 =1cm
| 24cm +- 1cm
229
What is the symbol for the unit of mole?
mol
230
The mass of one mole of a substance in grams is numerically equal to its relative formula mass. True or false?
True.
231
State the value of Avogadro's constant.
6.02 x 10^23
232
I have 2 moles of Na atoms, how many Na atoms do I have?
1.204 x 10^24
233
How many oxygen atoms are there in 1 mole of carbon dioxide?
1.204 x 10^24
234
How many moles does 24g of carbon have?
24 / 2 = 2
235
What is the mass of 3 moles of Lithium?
7 x 3 = 21g
236
What mass is 1 mole of magnesium?
24g
237
What mass is one mole of MgO?
24 + 16 = 40g
238
How many moles are there in 66g of CO2?
66 / 44 = 1.5mol
239
What mass of carbon is there in 4 moles of carbon dioxide?
4 x 12 = 48g
240
What is the mass of 1.7 moles of NaCl?
58.5 x 1.7 = 99.45g
241
Mg + 2HCI -> MgCI2 + H2
| How many moles of hydrochloric acid react would with one mole of magnesium?
2
242
Mg + 2HCI -> MgCI2 + H2
| How many moles of hydrochloric acid would react with 2 moles of magnesium?
4
243
Mg + 2HCI -> MgCI2 + H2
| How many moles of magnesium chloride would I produce if I had 6 moles of hydrochloric acid and 3 moles of magnesium?
3
244
Mg + 2HCI -> MgCI2 + H2
| How many grams of Mg would be needed to react with 18.25g of HCl?
18.25 / 36.5 = 0.5moles of HCl
molar ratio is 2:1...
0.5moles / 2 = 0.25 moles of Mg
0.25 x 24 = 6g of Mg
245
Mg + 2HCI -> MgCI2 + H2
| How many grams of Hydrogen would be produced if 48g of magnesium reacted completely?
48g / 24 = 2moles of magnesium
molar ratio is 1:1 therefore 2 moles of hydrogen produced
2 x 1 = 2g of hydrogen produced
246
8.1g of ZnO reacts completely with 0.6g of C to form 2.2g of CO2 and 6.5g of Zn. Write a balanced symbol equation for this reaction.
Mr for each substance: ZnO = 81, C = 12, CO2 = 44, Zn = 65
divide mass by Mr to get moles...
ZnO = 0.1, C = 0.05, CO2 = 0.05, Zn = 0.1
Divide by the smallest number of moles to get...
ZnO = 2, C = 1, CO2 = 1, Zn = 2
2Zn + C -> CO2 + 2Zn
247
In a chemical reaction involving two reactants, it is common to use an excess of one of the reactants. What is the name given to the other reactant?
Limiting reactant.
248
If the amount of limiting reactant is doubled, what happens to the mass of product formed?
The mass of product also doubles.
249
If the amount of limiting reactant is halved, what happens to the mass of product formed?
The mass of product also halves.
250
In a reactant that involves two reactants, why is it common to use an excess of one of the reactants?
To ensure the second reactant has been completely used up.
251
Many chemical reactions take place in solutions. True or False?
True.
252
Give the units for concentration of a solution. (It must contain g/grams)
g/dm3
253
What is the formula to calculate concentration. (It must contain g/grams)
concentration = mass of dissolved substance / volume of solvent
254
30g of sodium chloride is dissolved in 0.2dm3 of water. What's the concentration of this solution in g/dm3?
0.3/0.2 = 150g/dm3
255
Convert 400cm3 to dm3
400/1000 = 0.4dm3
256
15g of salt is dissolved in 500cm3 of water. What's the concentration of this solution in g/dm3?
```
500/1000 = 0.5dm3
15/0.5 = 30g/dm3
```
257
A solution of magnesium chloride has a concentration of 24g/dm3. What mass of magnesium chloride is there in 0.4dm3 of this solution?
```
mass = concentration x volume
mass = 24 x 0.4 = 9.6g
```
258
What type of compound is produced when metals react with oxygen?
Metal oxides.
259
Define oxidation in the context of loss and gain of oxygen
Oxidation is the gain of oxygen
260
Define reduction in the context of loss and gain of oxygen
Reduction is the loss of oxygen
261
Write the word equation for when magnesium reacts with oxygen.
Magnesium + oxygen -> magnesium oxide
262
Write the balanced word equation for when sodium reacts with oxygen.
4Na + O2 -> 2Na2O
263
When metals react with other substances, what type of ion do the metal atoms form?
Positive ions
264
The reactivity of a metal is related to its…
tendency to form positive ions
265
What is the reactivity series?
A list in which metals are arranged in order of their reactivity. Most reactive at the top and least reactive at the bottom.
266
Which two non-metal elements are normally included in the reactivity series?
Hydrogen and carbon
267
A less reactive metal can displace a more reactive metal from a compound. True or false?
False.
268
Why is potassium more reactive than magnesium?
Potassium has a greater tendency to form positive ions than magnesium
269
Name a metal that is found by itself and not in ores.
Gold, silver
270
What is the name for a reaction where oxygen is removed from a compound?
Reduction
271
Explain why zinc can be extracted from zinc oxide with carbon but magnesium cannot be extracted from magnesium oxide with carbon
Magnesium is more reactive than carbon (1), Zinc is less reactive than carbon (1), Carbon can therefore remove oxygen from zinc oxide but not magnesium oxide (1)
272
Which chemical reaction can be used to extract metals less reactive than carbon from their oxides?
Reduction (using carbon)
273
Explain why gold and silver can be found naturally in the Earth's crust
They are very unreactive
274
Complete and balance the equation: CuO + C ->
2CuO + C -> 2Cu + CO2
275
What type of reaction occurs when copper oxide reacts with carbon?
Reduction
276
Complete and balance the equation: Fe2O3 + C ->
Fe2O3 + 3C -> 4Fe + 3CO2
277
Define oxidation in the context of loss and gain of electrons
Oxidation is the loss of electrons
278
Define reduction in the context of loss and gain of electrons
Reduction is the gain of electrons
279
Balance this ionic equation: Ca + Na⁺ → Ca²⁺ + Na
Ca + 2Na⁺ → Ca²⁺ + 2Na
280
Balance this ionic equation: Mg²⁺ + Na → Na⁺ + Mg
Mg²⁺ + 2Na → 2Na⁺ + Mg
281
Balance this ionic equation: Al + Ca²⁺ → Al³⁺ + Ca
2Al + 3Ca²⁺ →2Al³⁺ + 3Ca
282
In this reaction, what has been oxidised and what has been reduced? Mg²⁺ + 2Na → 2Na⁺ + Mg
Na has been oxidised, Mg²⁺ has been reduced
283
In this reaction, what has been oxidised and what has been reduced? 2Al + 3Ca²⁺ → 2Al³⁺ + 3Ca
Al has been oxidised, Ca²⁺ has been reduced
284
In this reaction, what has been oxidised and what has been reduced? K + Na⁺ → K⁺ + Na
K has been oxidized, Na⁺ has been reduced
285
Write the ionic equation for: Fe(s) + CuSO4(aq) -> FeSO4(aq) + Cu(s)
Fe(s) + Cu2+(aq) -> Fe2+(aq) + Cu(s)
286
Some metals such as zinc can react with acids to produce which two substances?
Salt and hydrogen
287
Complete the word equation: Magnesium + hydrochloric acid ->
Magnesium + hydrochloric acid -> Magnesium chloride + Hydrogen
288
Write the balanced symbol equation for the following: zinc reacts with sulphuric acid to produce zinc sulphate and hydrogen
H2SO4 + Zn -> ZnSO4 + H2
289
Fe + 2HCl -> FeCl2 + H2
| What is being reduced?
Hydrogen ions are being reduced. 2H+ + 2e- -> H2
290
Fe + 2HCl -> FeCl2 + H2
| What is being oxidised?
Iron atoms are being oxidised. Fe -> Fe2+ + 2e-
291
How do you know a redox reaction is taking place?
The reaction will have substances that gain, and substances that lose electrons.
292
Complete the equation: Acid + Metal oxide ->
Acid + Metal oxide -> Salt + Water
293
Complete the equation: Acid + Metal hydroxide ->
Acid + Metal hydroxide -> Salt + Water
294
Complete the equation: Acid + Metal carbonate ->
Acid + Metal carbonate -> Salt + carbon dioxide + water
295
Hydrochloric acid + copper oxide ->
Hydrochloric acid + copper oxide -> copper chloride + water
296
Sulphuric acid + potassium hydroxide ->
Sulphuric acid + potassium hydroxide -> potassium sulphate + water
297
Nitric acid + sodium hydroxide ->
Nitric acid + sodium hydroxide -> sodium nitrate + water
298
Hydrochloric acid + sodium carbonate ->
Hydrochloric acid + sodium carbonate -> Sodium chloride + carbon dioxide + water
299
Sulphuric acid + calcium carbonate ->
Sulphuric acid + calcium carbonate -> Calcium sulphate + carbon dioxide + water
300
Write the balanced symbol equation for the following reaction: Hydrochloric acid + sodium oxide -> sodium chloride + water
HCl + NaOH -> H2O+NaCl
301
Write the balanced symbol equation for the following reaction: Sulphuric acid + lithium hydroxide -> lithium sulphate + water
H2SO4 + 2LiOH -> Li2SO4 + 2H2O
302
Write the balanced symbol equation for the following reaction: Sulphuric acid + magnesium carbonate -> Magnesium sulphate + carbon dioxide + water
MgCO3 + H2SO4 -> MgSO4 + CO2 + H2O
303
Name an acid and insoluble base that can be used to produce soluble copper chloride.
copper hydroxide OR Copper oxide and hydrochloric acid
304
When preparing a soluble salt from an acid and insoluble base, which reactant should be added in excess?
The insoluble base
305
Which two examples of insoluble base can react with acid to produce a soluble salt?
An insoluble metal oxide or an insoluble metal hydroxide
306
How is copper sulphate made from copper oxide and sulphuric acid?
Add excess copper oxide to warmed sulphuric acid.
Filter out the excess solid.
Heat the solution to evaporate some of the water and until crystals start to form.
Allow the solution to cool and more crystals to form.
Filter the crystals out of the solution and leave to dry.
307
What ions do acids produce in aqueous solutions?
H+
308
What ions do aqueous solutions of alkali contain?
OH-
309
What numbers are displayed on the pH scale?
0-14
310
A solution with a pH of 7 is…
Neutral
311
Aqueous solutions of acids have a pH of ….
Less than 7
312
Aqueous solutions of alkalis have a pH of …
Greater than 7
313
Write the ionic equation for neutralisation, include state symbols.
H+ + OH- à H2O
314
Define base.
A substance with a pH of greater than 7
315
Write the word equation for neutralisation
Acid + Base à Salt + Water
316
State the pH of the products of neutralisation.
7
317
What is an indicator?
It is a dye that changes colour depending on whether it is above or below a certain pH
318
What is universal indicator?
It is a mixture of dyes that gradually change colour over a broad range of pH
319
What is a pH probe?
It is a piece of equipment used to measure pH electronically.
320
What is the difference between an alkali and a base?
A base is a substance that has a pH greater than 7. An alkali is a base that can dissolve in water to produce OH- ions.
321
A solution causes universal indicator to turn red. Is it an acid or alkali?
Acid
322
A solution causes universal indicator to turn purple. Is it an acid or alkali?
Alkali
323
Define strong acid.
An acid that completely ionises in aqueous solution.
324
Define weak acid.
An acid that only partially ionises in aqueous solution.
325
Name three strong acids.
Hydrochloric acid, nitric acid and sulphuric acid.
326
Name three weak acids.
Carbonic acid, citric acid and ethanoic acid.
327
For a given concentration, the strong the acid, the lower the pH. True or false?
True.
328
For a given concentration, the weaker the acid, the lower the pH. True or false?
False.
329
The pH of a solution changes from 4 to 3. How did the hydrogen ion concentration change?
It increased by a factor of 10.
330
The pH of a solution changes from 8 to 5. How did the hydrogen ion concentration change?
It increased by a factor of 1000.
331
What is pH the measure of?
Hydrogen ion concentration.
332
How is acid concentration different to acid strength?
Concentration refers to the amount of substance in a given volume, whereas acid strength refers to the degree of ionisation.
333
What is the relationship between the strength of an acid and its pH?
As an acid increases in strength the pH decreases.
334
What is a concentrated acid?
An acid where there are lots of acid particles in a given volume.
335
What is a dilute acid?
An acid where there are fewer acid particles in a given volume.
336
What does electrolysis mean?
Splitting using electricity.
337
What types of compounds can be separated using electrolysis?
Ionic compounds.
338
What happens to the ions in an ionic substance when it is melted or dissolved in water?
The ions become free to move around
339
What is the name for the positive electrode?
The anode
340
What is the name for the negative anode?
The cathode
341
Do positive ions move to the anode or the cathode?
Cathode
342
Do negative ions move to the anode or the cathode?
Anode
343
Define electrolyte.
A liquid or solution able to conduct electricity.
344
What happens to the ions in an electrolyte when a current is passed through it?
The ions move to the electrodes.
345
At which electrode would Zn²+(aq) turn into Zn(s)?
Cathode
346
At which electrode would Cl⁻(aq) turn into Cl₂(g)?
Anode
347
When ions are discharged at the electrodes, what are produced?
Elements
348
What will be the products for the electrolysis of molten iron bromide?
Iron and bromine
349
What will be the products for the electrolysis of molten zinc oxide?
Zinc and oxygen
350
When a molten ionic compound is electrolysed, what type of element forms at the cathode?
Metal
351
When a molten ionic compound is electrolysed, what type of element forms at the anode?
Non-metal
352
When molten sodium chloride is electrolysed, what element is formed at the cathode?
Sodium
353
When molten sodium chloride is electrolysed, what element is formed at the anode?
Chlorine
354
When molten magnesium oxide is electrolysed, what element is formed at the cathode?
Magnesium
355
When molten magnesium oxide is electrolysed, what element is formed at the anode?
Oxygen
356
When is electrolysis the preferred method to extract a metal from a molten compound?
When the metal is more reactive than carbon.
357
Why is a large amount of energy needed for the electrolysis of molten ionic compounds?
1) To melt the ionic compound
| 2) To produce the electrical current
358
Why is aluminium oxide mixed with cryolite before melting?
Adding cryolite reduced the melting point of aluminium oxide, saving energy.
359
Why is the positive electrode continually replaced during the electrolysis of molten aluminium oxide?
The electrode is made of carbon. The oxygen produced at the positive electrode reacts with the carbon to produce carbon dioxide.
360
Write the equation that takes place at the negative electrode when molten aluminium oxide is electrolysed.
Al³⁺ + 3e⁻ → Al
361
Write the equation that takes place at the positive electrode when molten aluminium oxide is electrolysed.
2O²⁻ → O₂ + 4e⁻
362
What element is produced at the negative electrode when molten aluminium oxide is electrolysed?
Aluminium
363
What element is produced at the positive electrode when molten aluminium oxide is electrolysed?
Oxygen
364
How many different ions are found in an aqueous solution that is being electrolysed?
4
365
As well as the metal and non-metal ions of the ionic compound, what other ions are present in an aqueous solution being electrolysed?
H+ and OH-
366
When electrolysing an aqueous solution, how do you determine which element is formed at the negative electrode (cathode)?
The least reactive element between the metal and hydrogen.
367
When electrolysing an aqueous solution, how do you determine which element is formed at the positive electrode (anode)?
If halide ions are present then a halogen is produced, otherwise oxygen is.
368
A solution of copper sulphate is electrolysed. Name the element produced at the cathode.
Copper
369
A solution of copper sulphate is electrolysed. Name the element produced at the anode.
Oxygen
370
A solution of sodium chloride is electrolysed. Name the element produced at the cathode.
Hydrogen
371
A solution of sodium chloride is electrolysed. Name the element produced at the anode.
Chlorine
372
A solution of copper sulphate is electrolysed. Write the half-equation that takes place at the cathode.
Cu2+ + 2e- -> Cu
373
A solution of copper sulphate is electrolysed. Write the half-equation that takes place at the anode.
4OH- -> 2H2O + O2 + 4e-
374
During electrolysis describe what happens at the cathode in terms of electrons.
Positively charged ions gain electrons - become reduced
375
During electrolysis describe what happens at the anode in terms of electrons.
Negatively ions lose electrons - become oxidised
376
Molten sodium chloride is electrolysed. Write the half-equation that takes place at the cathode.
Na+ + e- -> Na
377
Molten sodium chloride is electrolysed. Write the half-equation that takes place at the anode.
2Cl- -> Cl2 + 2e-
378
A solution of sodium chloride is electrolysed. Write the half-equation that takes place at the cathode.
2H+ + 2e- -> H2
379
A solution of sodium chloride is electrolysed. Write the half-equation that takes place at the anode.
2Cl- -> Cl2 + 2e-
380
State the law of conservation of energy.
Energy cannot be created or destroyed, it can only transferred from one place to another.
381
What type of reaction transfers energy to the surroundings?
Exothermic.
382
What type of reaction take in energy from the surroundings?
Endothermic.
383
What type of reactions show an increase in temperature?
Exothermic.
384
What type of reactions show a decrease in temperature?
Endothermic.
385
Give three examples of exothermic reactions.
Combustion, many oxidation reactions and neutralisation.
386
Give two examples of endothermic reactions.
Thermal decomposition and photosynthesis.
387
Give two everyday uses of exothermic reactions.
Self heating cans and hand warmers.
388
Give an everyday use of endothermic reactions.
Sports injury cooling packs.
389
Two solutions are mixed together. A reaction occurs increasing the temperature of the solution. What type of reaction occurred?
Exothermic.
390
Two solutions are mixed together. A reaction occurs decreasing the temperature of the solution. What type of reaction occurred?
Endothermic.
391
Define activation energy.
The minimum amount of energy that particles must have to react.
392
In order to react, particles do not need to collide. True or false?
False.
393
What do reaction profiles show?
They show if reactants have more or less energy than the products.
394
What type of reaction would the reactants have more energy than the products?
Exothermic.
395
What type of reaction would the reactants have less energy than the products?
Endothermic.
396
Is bond breaking endothermic or exothermic?
Endothermic. Chemical bonds are strong so require energy to break.
397
Is bond making endothermic or exothermic?
Exothermic. Energy is released when chemical bonds are formed.
398
How do we work out the overall energy change of a reaction?
Energy needed to break all the bonds in the reactants - the energy released to form all the bonds in the products.
399
If your overall bond energy is negative, what does that mean and why?
Exothermic reaction, energy is lost to the surroundings.
400
For an exothermic reaction, describe the difference between the energy needed to break bonds and the energy release when making bonds.
More energy is released making bonds than is needed to break bonds.
401
For an endothermic reaction, describe the difference between the energy needed to break bonds and the energy release when making bonds.
More energy is needed to break bonds than is released when making bonds.
402
Using the bond energy given below, calculate the energy change for the reaction between H2 and Cl2 forming HCl: H-H + Cl-Cl -> H-Cl H-Cl
H-H: +436k kJ/mol, Cl-Cl: +242 kL/mol, H-Cl: +431 kJ/mol
Energy required to break bonds: 436 + 242 = 678
Energy released when making bonds: 2 x 431 = 862
Energy required to break bonds - energy released when making bonds…
678 - 862 = -184 kJ/mol
