Week 20: Redox and Electrolysis Leaving Certificate Chemistry Learning Intentions • Revise oxidation and reduction • Explore change in oxidation number • Balance equations using oxidation numbers • Define electrolysis, electrolyte, electroplating and electrochemical series • Investigate a number of experiments that use electrolysis (2 OL and 4HL) • Describe an experiment using a galvanic/voltaic cell Oxidation Reduction Is loss of electrons Is gaining of electrons Oxidation number will become more positive (increase) as negative charges are removed Oxidation number will become less positive (decrease) as negative charges are added Oxidation number: is the charge that an atom has (or appears to have) when electrons are distributed according to certain rules. Rules 1. Oxidation number of any single element is 0 2. Oxidation number of an ion is the same as it’s charge 3. Sum of oxidation is all the elements must add up to zero 4. Oxidation of oxygen is -2 (unless in a peroxide (-1) or when bound to a more electronegative element (OF 2 = +2) 5. Hydrogen = +1 (unless bonded to more e-negative element, then = -1) 6. Halogens = -1 (unless bonded to more e-negative element, then = +1) 7. If the compound has a charge the oxidation numbers must equal the charge Practice assigning oxidation numbers Na 2 Cr 2 O 7 MnO 4 - S 2 O 3 2- MnO 2 Balancing equations (using oxidation numbers ) – HL only If one compound in a reaction oxidised then something else must be reduced. The number of electrons that are lost must equal the number gained by another substance.
Section B [See page 1 for instructions regarding the number of questions to be answered]
4. Answer eight of the following items (a), (b), (c), etc. (50)
(a) Define relative atomic mass.
(b) Account for the difference in the shapes of the ammonia (NH3) and boron trifluoride (BF3) molecules.
(c) The boiling points of hydrogen and oxygen are 20.0 K and 90.2 K respectively. Account for the higher boiling point of oxygen.
(d) State Charles’ law.
(e) Write (i) the conjugate acid and (ii) the conjugate base of .
(f) How are heavy metals, e.g. mercury, removed from industrial waste before it is discharged into rivers, lakes or the sea?
(g) What is the oxidation number (i) of oxygen in H2O2 and (ii) of bromine in KBrO3?
(h) What is the percentage by mass of iron in iron(III) oxide (Fe2O3)?
(i) State and explain the colour observed at the negative electrode in the electrolysis of aqueous potassiumiodide, containing a little phenolphthalein indicator, using inert electrodes.
(j) How could the presence of sulfite ions in aqueous solution be detected?
(k) Answer part A or B.
A How is oxygen gas produced industrially?
or
B How does the anodising of aluminium protect it from corrosion? ____________________________
5. (a) Write the electron configuration (s, p, etc.) of the nitrogen atom. (5)
Show, using dot and cross diagrams, the bond formation in a nitrogen molecule. Describe the bonding in the nitrogen molecule in terms of sigma (σ) and pi (π) bonding. (9)
What type of intermolecular forces would you expect to find in nitrogen gas? Explain your answer. (6)
(b) Define first ionisation energy. (9)
There is a general increase in first ionisation energy across a period of the periodic table. State the two principal reasons for this trend. (6)
The table shows the first and second ionisation energies of nitrogen, oxygen, neon and sodium.
Account for the decrease in first ionisation energybetween nitrogen and oxygen.
Explain why the second ionisation energy of sodium is significantly (about nine times) higher than the first while the increase in the second ionisation energy of neon compared to its first is relatively small (less than twice the first). (15)
9. The alkenes are a homologous series. Ethene (C2H4) is the first member of the series.
(a) What is meant by a homologous series? (5)
(b) Ethene may be made in a school laboratory using the arrangement of apparatus drawn on the right.
(i) Give the name and formula of the solid A which is heated using the Bunsen burner. (6)
(ii) Identify the solid B which is used to keep the ethanol at the end of the test tube. (3)
(iii) What precaution should be observed when heating is stopped? Why is this necessary? (6)
(iv) Give one major use of ethene gas. (3)
(c) Describe the mechanism for the bromination of ethene. (9) State and explain one piece of experimental evidence to support this mechanism. (6)
(d) Draw the structures and give the systematic (IUPAC) names for two alkene isomers of molecularformula C4H8. (12)
____________________________ 10. Answer any two of the parts (a), (b) and (c) (2 × 25)
(a) (i) What are isotopes? (4)
(ii) Define relative atomic mass, Ar. (6)
(iii) What is the principle on which the mass spectrometer is based? (9)
(iv) Calculate the relative atomic mass of a sample of lithium, given that a mass spectrometershows that it consists of 7.4 % of 6Li and 92.6 % of 7Li. (6)
(b) Define oxidation in terms of change in oxidation number. (4)
What is the oxidation number of (i) chlorine in NaClO and (ii) nitrogen in NO3¯? (6)
State and explain the oxidation number of oxygen in the compound OF2. (6)
Using oxidation numbers or otherwise, identify the reducing agent in the reaction between acidified potassium manganate(VII) and potassium iodide solutions represented by the balanced equation below. Use your knowledge of the colours of the reactants and products to predict the colour change you would expect to see if you carried out this reaction. (9)
2MnO4¯ + 10I¯ + 16H+ 2Mn2+ + 5I2 + 8H2O
(c) The chart compares the boiling points of alkanes and primary alcohols containingfrom one to four carbon atoms.
(i) Give two reasons why each of thesealcohols has a higher boiling pointthan the corresponding alkane. (7)
(ii) Explain why the difference in boiling points between methane and methanol is 226.5 K while the difference in boiling points between butane and butanol is only 119 K. (6)
(iii) Describe, in general terms, the solubilities of methane, methanol, butane and butanolin water. (12)
11. Answer any two of the parts (a), (b) and (c). (2 × 25)
(a) (i) Define oxidation in terms of change in oxidation number. (4)
(ii) What is observed when chlorine gas is bubbled into an aqueoussolution of sodium bromide?
Explain your answer in terms of oxidation and reduction. (9)
(iii) A solution of acidified water (dilute sulfuric acid) is electrolysed by passing an electric current through it using inert electrodes.At which electrode A or B does oxidation occur?Which species is oxidised?Write a balanced half equation for the oxidation reaction. (12)
(b) (i) Define a mole of a substance. (7)
(ii) State Avogadro’s law. (6)
(iii) A foil balloon has a capacity of 10 litres. How many atoms of helium occupy this balloon when it is filled with a 10% (v/v) mixture of helium in air at room temperature and pressure? (12)
(c) Answer either part A or part B.
A
(i) What is meant by the term addition polymerisation? (7)
(ii) Name the Du Pont chemist pictured on the right who discovered poly(tetrafluoroethene), PTFE. (3)
(iii) Describe using an equation how poly(tetrafluoroethene) is produced from its monomers. (9)
(iv) Give two common uses of PTFE. (6)
or
B
(i) Account for the inert nature of nitrogen gas. (7)
(ii) What is meant by nitrogen fixation?
State two ways by which nitrogen is fixed in nature. (9)
(iii) The concentration of NO2 in the atmosphere has increased in the past fifty years. Describe with the aid of chemical equations how an increase in the number of cars has contributed to thischange. (9)
9. The alkenes are a homologous series. Ethene (C2H4) is the first member of the series.
(a) What is meant by a homologous series? (5)
(b) Ethene may be made in a school laboratory using the arrangement of apparatus drawn on the right.
(i) Give the name and formula of the solid A which is heated using the Bunsen burner. (6)
(ii) Identify the solid B which is used to keep the ethanol at the end of the test tube. (3)
(iii) What precaution should be observed when heating is stopped? Why is this necessary? (6)
(iv) Give one major use of ethene gas. (3)
(c) Describe the mechanism for the bromination of ethene. (9) State and explain one piece of experimental evidence to support this mechanism. (6)
(d) Draw the structures and give the systematic (IUPAC) names for two alkene isomers of molecularformula C4H8. (12)
____________________________ 10. Answer any two of the parts (a), (b) and (c) (2 × 25)
(a) (i) What are isotopes? (4)
(ii) Define relative atomic mass, Ar. (6)
(iii) What is the principle on which the mass spectrometer is based? (9)
(iv) Calculate the relative atomic mass of a sample of lithium, given that a mass spectrometershows that it consists of 7.4 % of 6Li and 92.6 % of 7Li. (6)
(b) Define oxidation in terms of change in oxidation number. (4)
What is the oxidation number of (i) chlorine in NaClO and (ii) nitrogen in NO3¯? (6)
State and explain the oxidation number of oxygen in the compound OF2. (6)
Using oxidation numbers or otherwise, identify the reducing agent in the reaction between acidified potassium manganate(VII) and potassium iodide solutions represented by the balanced equation below. Use your knowledge of the colours of the reactants and products to predict the colour change you would expect to see if you carried out this reaction. (9)
2MnO4¯ + 10I¯ + 16H+ 2Mn2+ + 5I2 + 8H2O
(c) The chart compares the boiling points of alkanes and primary alcohols containingfrom one to four carbon atoms.
(i) Give two reasons why each of thesealcohols has a higher boiling pointthan the corresponding alkane. (7)
(ii) Explain why the difference in boiling points between methane and methanol is 226.5 K while the difference in boiling points between butane and butanol is only 119 K. (6)
(iii) Describe, in general terms, the solubilities of methane, methanol, butane and butanolin water. (12)
Boiling points
050100150200250300350400450
1 2 3 4
Number of carbons
b.p./K Alkanes
Alcohols
solid B soaked with ethanol
solid A C2H4
Page 6 of 8
9. (a) Define the rate of a chemical reaction. Why does the rate of chemical reactions generally decrease with time? (8)
(b) The rate of reaction between an excess of marble chips (CaCO3) (diameter 11 – 15 mm) and 50 cm3
of 2.0 M hydrochloric acid was monitored by measuring the mass of carbon dioxide produced.
The table shows the total mass of carbon dioxide gas produced at stated intervals over 9 minutes.
Plot a graph of the mass of carbon dioxide produced versus time. (12)
Use the graph to determine (i) the instantaneous rate of reaction in grams per minute at 4.0 minutes,
(ii) the instantaneous rate of reaction at this time in moles per minute. (9)
(c) Describe and explain the effect on the rate of reaction of repeating the experiment using 50 cm3 of 1.0M hydrochloric acid and the same mass of the same size marble chips. (6)
(d) Particle size has a critical effect on the rate of a chemical reaction.
(i) Mark clearly on your graph the approximate curve you would expect to plot if the experiment were repeated using 50 cm3 of 2.0 M HCl and using the same mass of marble chips but this timewith a diameter range of 1 – 5 mm. (6)
(ii) Dust explosions present a risk in industry. Give three conditions necessary for a dust explosionto occur. (9)
_____________________________________
10. Answer any two of the parts (a), (b) and (c). (2 × 25)
(a) (i) Write the equilibrium constant (Kc) expression for the reaction (7)
N2(g) + 3H2(g) 2NH3(g)
(ii) Three moles of nitrogen gas and nine moles of hydrogen gas were mixed in a 1 litre vessel at a temperature T. There were two moles of ammonia in the vessel at equilibrium. Calculate the value of Kc for this reaction at this temperature. (12)
(iii) Henri Le Chatelier, pictured on the right, studied equilibrium reactions in industry in the late 19th century. According to Le Chatelier’s principle, what effect would an increase in pressure have on the yield of ammonia at equilibrium? Explain. (6)
(b) (i) State Avogadro’s law. (7)
(ii) Carbon dioxide is stored under pressure in liquid form in a fire extinguisher. Two kilograms of carbon dioxide are released into the air as a gas on thedischarge of the fire extinguisher. What volume does this gas occupy at a pressure of 1.01 × 105 Pa and a temperature of 290 K? (9)
What mass of helium gas would occupy the same volume at the same temperatureand pressure? (6)
(iii) Give one reason why carbon dioxide is more easily liquefied than helium. (3)
(c) The halogens are good oxidising agents.
(i) How does the oxidation number of the oxidising agent change during a redox reaction? (4)
(ii) Assign oxidation numbers in each of the following equations to show clearly that the halogen is the oxidising agent in each case. (12)
Br2 + 2Fe2+ → 2Br– + 2Fe3+
Cl2 + SO32– + H2O → Cl– + SO4
2– + H+
Hence or otherwise balance the second equation. (6)
(iii) Why does the oxidising ability of the halogens decrease down the group? (3)
Page 6 of 8
9. The alkenes are a homologous series of unsaturated hydrocarbons. Ethene (C2H4) is the first member of the series. Alkenes undergo addition reactions and polymerisation reactions.
(a) Draw a labelled diagram of an apparatus used to prepare ethene gas in the school laboratory. (8)
(b) Draw the structure of any one of the isomers of the third member of the alkene series. Indicate clearly which carbon atoms have planar bonding and which are bonded tetrahedrally. (12)
(c) Explain the term unsaturated. (6)
(d) The ionic addition mechanism for the reaction of ethene with bromine water involves the formation of an intermediate ionic species. Draw the structure of this species.Give the names or structural formulas of the three products that would be formed if the bromine water used in the reaction contained sodium chloride.
How does the formation of these three products support the mechanism of ionic addition? (18)
(e) Name the polymer formed when ethene undergoes addition polymerisation. Draw two repeating units of this polymer. (6)
____________________________
10. Answer any two of the parts (a), (b) and (c). (2 × 25)
(a) A student is given a bucket of seawater. (i) Describe how the student could determine by filtration the total suspended solids (expressed as
ppm) in the water. (9)
(ii) How could the student determine the total dissolved solids (expressed as ppm) in a sample of the filtered seawater? (9)
(iii) Describe a test to confirm the presence of the chloride ion in aqueous solution. (7)
(b) Define oxidation in terms of (i) electron transfer, (ii) change in oxidation number. (7)
(iii) For the redox reactions shown below, use oxidation numbers to identify the species oxidised in the first reaction and the oxidising reagent in the second reaction. (6)
ClO¯ + I¯ + H+ Cl¯ + I2 + H2O
I2 + S2O32– I¯ + S4O6
2–
(iv) Using oxidation numbers or otherwise balance both equations. (12)
(c) (i) Define energy level. (4)
(ii) Distinguish between ground state and excited statefor the electron in a hydrogen atom. (6)
The diagram shows how Bohr related the lines in the hydrogen emission spectrum to the existence of atomic energy levels.
(iii) Name the series of lines in the visible part of the line emission spectrum of hydrogen. (3)
(iv) Explain how the expression E2 – E1 = hf links the occurrence of the visible lines in the hydrogen spectrum to energy levels in a hydrogen atom. (12)
n = 6 n = 5
n = 4
n = 3
n = 2
n = 1
nucleus
purple blue/green red
Page 6 of 8
9. (a) Explain (i) activation energy, (ii) effective collision. (8) The effect of temperature on the rate of a chemical reaction was investigated using dilute solutions of hydrochloric acid and sodium thiosulfate. Suitable volumes and concentrations of the solutions were used. The reaction is represented by the following balanced equation.
2HCl + Na2S2O3 o 2NaCl + H2O + S + SO2
Describe how the time for the reaction between the solutions of hydrochloric acid and sodiumthiosulfate was obtained at room temperature. (6)
In a reaction mixture what effect, if any, does an increase in temperature of 10 K have on each of thefollowing: (i) the number of collisions, (ii) the effectiveness of the collisions, (iii) the activation energy. (9)
(b) The catalytic oxidation of methanol using platinum wire is illustrated in the diagram. State one observation made during the experiment. Name any two products of the oxidation reaction. What type of catalysis is involved in this reaction? (12)Explain one way in which the presence of the platinum catalyst speeds up the oxidation of the hot methanol. Explain how a catalyst poison interferes with this type ofcatalysis. (9) Give another example of a reaction which involves the sametype of catalysis, indicating clearly the reactant(s) and the catalyst. (6)
_____________________________________
10. Answer any two of the parts (a), (b) and (c). (2 u 25)
(a) State Avogadro’s law. (7)
Give two assumptions of the kinetic theory of gases. (6)
Give two reasons why real gases deviate from ideal gas behaviour. (6)
How many moles of gas are present in a sample containing 1.8 u 1024 atoms of chlorine at s.t.p.? (6)
(b) Define oxidation in terms of electron transfer. (4)
The electrolysis, using inert electrodes, of aqueous potassium iodide, KI, to which a few drops of phenolphthalein indicator have been added, is shown in the diagram.
(i) Name a suitable material for the electrodes. (3)
(ii) Write balanced half equations for the reactions that takeplace at the electrodes. (12)
(iii) Explain the colour change observed at the positiveelectrode (anode). (6)
(c) In 1922, Francis Aston, pictured right, was awarded the Nobel Prize inchemistry for detecting the existence of isotopes using the first mass spectrometer.
(i) What are isotopes? (7)
(ii) What is the principle of the mass spectrometer? (9)
(iii) Calculate, to two decimal places, the relative atomic mass of a sample of neon shown by mass spectrometer to be composed of 90.50% of neon–20 and 9.50% of neon–22. (9)
9. (a) Define (i) oxidation, (ii) reduction, in terms of electron transfer. (8)
(b) Identify (i) the substance oxidised, (ii) the oxidising agent, in the reaction between sodium and chlorine to produce sodium chloride. The equation for the reaction is: Na + ½Cl2 o NaCl (6)
(c) Name the type of chemical bond that exists between the chlorine atoms in a chlorine molecule. Give one other example of a molecule having this type of bond. State two properties of substances having this type of bond. (12)
(d) What type of bonding occurs in sodium chloride? Give one other example of a compound having this type of bonding. State two properties of substances which contain this type of bond. (12)
(e) Describe what you would observe when a small piece of sodium is put in a dish of water. Name one product of the reaction that occurs. (12)
10. Answer any two of the parts (a), (b) and (c). (2 × 25)
(a) The volume of oxygen liberated from liquid X in the presence of manganese(IV) oxide (MnO2) as catalyst was measured at one-minute intervals using the apparatus shown in the diagram. The results obtained are shown in the following table.
Time (min) 0 1 2 3 4 5 6 7
Volume of oxygen (cm3) 0 36 54 64 69 71 72 72
(i) Give the name and formula of liquid X. (7) (ii) Plot, on graph paper, a graph of volume (y-axis) versus time (x-axis). (12) (iii) From your graph estimate the volume of gas liberated after 2.5 minutes. (6)
(b) What do you understand by the relative molecular mass (Mr) of a compound? (7) Calculate the relative molecular mass of the glucose molecule (C6H12O6). (6) Find the percentage by mass of hydrogen in glucose. (6) How many moles are there in 9 grams of glucose? (6)
(c) The following scientists all played a part in the development of our knowledge of elements. Boyle Mendeleev Moseley The Greeks Davy
Write in your answer book the name corresponding to each of the numbers from 1 to 5 in the statements below. (25)
It was thought by 1 that matter was made up of four elements: earth, air, fire and water. In the 17th century 2 defined an element as a simple substance that cannot be split into simpler substances. Early in the 19th century the elements, sodium and potassium, were isolated by 3 . Later in the 19th century the known elements were arranged in a table by 4 . In the 20th century the table of the elements was modernised as
a result of the discovery of atomic number by 5 .
oxygen
MnO2
X
water
Page 7 of 8
11. Answer any two of the parts (a), (b) and (c). (2 × 25)
(a) Consider the following six terms. refluxing fractionation recrystallisation distillation chromatography steam distillation In your answer book, match any five of the terms above with their corresponding descriptions (A to F) in the table below. (25)
A Separation of the coloured components of a mixture using a solvent as the mobile phase and a suitable solid as the stationary phase
B Separation of two liquids on the basis of their different boiling pointsC Dissolving impure crystals in the minimum amount of hot solvent and then filtering the
mixture after allowing it to coolD Separation of several substances in crude oil on the basis of difference in boiling pointsE Boiling a reaction mixture without loss of material, thus giving time for reaction to occur F Extraction of clove oil from cloves (or similar extraction of another oil)
(b) Define (i) oxidation, (ii) reduction, in terms of electron transfer. (6)
When a zinc rod is dipped into a blue solution of copper(II) sulfate, the following oxidation-reduction reaction takes place. Zn + Cu2+ → Zn2+ + Cu
Identify (iii) the species oxidised, (iv) the species reduced, (v) the oxidising agent. (12) Is zinc (Zn) placed above or below copper (Cu) in the electrochemical series? Give a reason for your answer. (7)
(c) Answer part A or part B.
AChemical processes are usually either batch processes or continuous processes.Distinguish between the two. (6)
Suggest one way of reducing energy costs in a chemical industry. (4)
Answer the following questions in relation to the chemical process taken as your case study based on the Irish chemical industry.
(i) State the location in Ireland of the process that was the subject of your case study.Suggest two reasons in support of the suitability of this location. (9)
(ii) Name the product of the chemical process you studied. Give a use for this product. (6)
B Carbon exists in different physical forms (allotropes), e.g. diamond. Part of the structure of diamond, a covalent macromolecular crystal, is shown on the right.
(i) What particles occupy the lattice points in diamond crystals? (ii) Name another physical form (allotrope) of carbon. (10)
A small part of the cubic crystal lattice of sodium chloride is also shown. (iii) What particles occupy the lattice points in sodium chloride crystals? (iv) What are the binding forces in a sodium chloride crystal? (v) What technique is used to study the internal structures of crystals? (15)
Page 7 of 8�
X
charge
blast furnace
hot airmolten iron
11. Answer any two of the parts (a), (b) and (c). (2 × 25)
(a) The following words are omitted from the passage below. zinc gained lost oxidised reduced redox Write in your answer book the omitted word corresponding to each number (1 to 6).
In a 1 reaction electrons are transferred from one substance to another. When electrons are 2 the substance is said to be oxidised and when electrons are 3 the substance is
reduced. In the reaction Zn + CuSO4 � ZnSO4 + Cu
the copper ions are 4 and 5 acts as the reducing agent. Therefore zinc is more easily 6 than copper.
State one observation made when zinc reacts with copper sulfate. (25)
(b) A chemical equilibrium between nitrogen gas, hydrogen gas and ammonia gas is set up at a certain temperature and pressure according to the following balanced equation.
N2 + 3H2 � 2NH3
Explain the underlined term. (7)Write the equilibrium constant (Kc) expression for this reaction. (6) What effect would an increase in pressure have on the yield of ammonia at equilibrium?Explain your answer.This equilibrium is used in the Haber process to make ammonia industrially. Give one reason why extremely high pressures are not used in the Haber process. (12)
(c) Answer part A or part B.A
(i) Oxygen gas (O2) is produced industrially by liquefaction and fractional distillation of air. What other gas is the major product of this process?What is the approximate percentage of this second gas in the atmosphere?Give a widespread use for this second gas. (10)
(ii) Name the other form of oxygen gas (O3) formed in the stratosphere.What environmentally beneficial effect does O3 have?��
� � Identify a group of chemicals that damages O3. (15)
or B The diagram shows a blast furnace used to extract iron metal from iron ore. Iron ore is�added in the charge at the top while hot air is pumped in at the bottom as shown. Molten iron trickles down and is removed at the bottom. Molten impurities that float on the iron are removed at outlet X.
(i) Give the name or formula of an ore of iron.Name another substance added in the charge
with the iron ore. What name is given to the impurity removed at X? (15)Most of the iron produced is converted to an alloy of iron.
(ii) Name this alloy.What other element is always present in this alloy? (10)
Higher Level Chemistry 2014 Page 8 of 24
QUSESTION 4Eight items to be answered. Six marks to be allocated to each item and one additional mark to be addedto each of the first two items for which the highest marks are awarded.
(a) WHAT: (i) green (yellow-green) // (ii) crimson (red) (2 × 3)
(b) DESCRIBE: mass of positively-charged material //with electrons (small negative charges)scattered (embedded) in it
[Correctly labelled diagram acceptable.]
(2 × 3)
(c) WRITE: Fr87223 → 𝐑𝐚𝟖𝟖
𝟐𝟐𝟑 / + 𝒆−𝟏 𝟎 (2 × 3)
[Allow if francium not written.] [Accept β instead of e.][Beta-particle as reactant notacceptable unless placed with a minus sign after Fr on the left.]
(d) STATE: position (location) and momentum (velocity, speed, energy) //of an electron cannot be found (measured or known) simultaneously (at the sametime) / of a particle cannot be found simultaneously (at the same time) with accuracy
(2 × 3)
(e) HOW: (i) 1 sigma // (ii) 2 pi (2 × 3)
(f) WHAT: amount containing as many particles* as //the number of atoms in 0.012 kg (12 g) of carbon-12 /
or
amount containing the Avogadro number (Avogadro constant, L, 6 × 1023
of particles /) //
or
amount equal to the relative formula (molecular) mass (Mrexpressed in grams (2 × 3)
) //
[*Allow ‘atoms’, ‘molecules’, ‘ions’, ‘units’ for ‘particles’.]
