C13 Electrochemistry [PDF] - Past Paper Practice

Nan Hua High School Chemistry Department SECONDARY 4 Chapter 13 Electrochemistry Chapter Objectives Describe electrolysis as the conduction of electricity through an ionic compound (an electrolyte), when molten or dissolved in water, leading to chemical changes (...

Nan Hua High School Chemistry Department SECONDARY 4 Chapter 13 Electrochemistry Chapter Objectives Describe electrolysis as the conduction of electricity through an ionic compound (an electrolyte), when molten or dissolved in water, leading to chemical changes (including decomposition) at the electrodes Describe electrolysis as evidence for the existence of ions which are held in a lattice when solid but which are free to move when molten or in solution Describe, in terms of the mobility of ions present and the electrode products, the electrolysis of molten sodium chloride, using inert electrodes Predict the likely products of the electrolysis of a molten binary ionic compound using inert electrodes Chapter Objectives Apply the idea of selective discharge based on i. cations: linked to the reactivity series ii. anions: halides, hydroxides and sulfates (e.g. aqueous copper(II) sulfate and dilute sodium chloride solution (as essentially the electrolysis of water)) iii. concentration effects (as in the electrolysis of concentrated and diluted aqueous sodium chloride) (in all cases above, inert electrodes are used) Predict the likely products of the electrolysis of an aqueous electrolyte, given relevant information Construct ionic equations for the reactions occurring at the electrodes during the electrolysis, given relevant information Chapter Objectives Describe the electrolysis of aqueous copper(II) sulfate with copper electrodes as a means of purifying copper (no technical details are required) Describe the electroplating of metals, e.g. copper plating, and state one use of electroplating Describe the production of electrical energy from simple cells (i.e. two electrodes in an electrolyte) linked to the reactivity series and redox reactions (in terms of electron transfer) Describe hydrogen, derived from water or hydrocarbons, as a potential fuel, reacting with oxygen to generate electricity directly in a hydrogen fuel cell (details of the construction and operation of a fuel cell are not required)  13.1 What is electrolysis? What is electrolysis? What is electrolysis? Electrolysis involves passing an electric current through a compound to chemically separate its components. For example, the flow of electricity through water will cause it to break down into its elements, hydrogen and oxygen. Electrolysis is the process of using electricity to break down or decompose a compound (usually an ionic compound in the molten or aqueous state). Electrolysis takes place in an electrolytic cell which converts electrical energy into chemical energy. Parts of an electrolytic cell For electrolysis to occur, it requires the following: ❑ a power supply or a power source (e.g. battery) that drives the movement of charges around a circuit ❑ electrodes which are connected to opposite ends of the power supply; and ❑ an electrolyte in which the electrodes are immersed.  Battery is the power supply and it acts as an ‘electron pump’ as it causes electrons to move from the anode to the cathode. Electrons enter the battery from the anode and are “pumped out” to the cathode. Electrolyte is an electrically conductive substance in the molten or aqueous state. Conducts electricity due to mobile ions which act as mobile charge carriers to conduct electricity Examples of electrolytes: dilute sulfuric acid, molten sodium chloride, copper(II) sulfate solution + – Electrodes contains delocalised mobile electrons to conduct electricity. The anode is the positive electrode connected to the positive terminal of the power source. The cathode is the negative electrode connected to the negative terminal of the power source. Examples of electrodes: metal plates and carbon (graphite) rods Types of electrodes Electrodes must be electrical conductors that conduct electricity through the movement of delocalised mobile electrons found in their structures. There are two types of electrodes: Inert electrodes Reactive electrodes Electrodes that do not undergo Metal anodes undergo oxidation chemical changes and do not take during the electrolysis reaction. part in the electrolysis reaction. Example: other metals such as Example: graphite, platinum copper, silver The process of electrolysis In electrolysis, charges are carried through the external circuit by electrons and through the electrolyte by ions. Hence, the electrons and ions are known as charge carriers. In the external circuit: Electrons move from the negative terminal to the positive terminal of the power supply + – Within the electrolyte: The movement of ions inside the electrolyte “completes” the circuit.  At the cathode: At the anode: Cations are attracted to Anions are attracted to the the negatively-charged positively-charged anode. cathode. Anions can lose electrons Cations can gain electrons to the anode and become + – from the cathode and oxidised. become reduced. When anions and cations are oxidised and reduced respectively, they form atoms or molecules. We can say that the anions and cations are discharged at the electrodes. Electrolytes An ionic compound must be dissolved in water (aqueous state) or melted (molten state) before it can be an electrolyte and conduct electricity. Ionic compounds are electrolytic conductors Electrolytes undergo redox reactions at the electrodes to form new substances. Ions are held in fixed In the molten and aqueous state, the mobile ions positions and are immobile. enable the electrolyte to conduct electricity. Difference between electrical conductors & electrolytic conductors electrical conductors electrolytic conductors (e.g. metals / graphite) (electrolytes) Conduct electricity by flow of Conduct electricity by flow of mobile delocalised mobile electrons from one ions end of the conductor to the other end. Substance remains chemically Substance are broken down to form unchanged when an electric current new substances when an electric flow through current flow through  13.2 How do we predict the products of electrolysis? Molten binary ionic compounds A molten binary ionic compound is typically a salt containing only one cation and one anion in the liquid state. Examples of molten binary ionic compounds and their constituent ions: molten binary ionic compound ions present sodium chloride, NaCl(l) Na+(l) and Cl-(l) magnesium bromide, MgBr2(l) Mg2+(l) and Br-(l) aluminium oxide, Al2O3(l) Al3+(l) and O2-(l) iron(III) nitride, FeN Fe3+(l) and N3-(l) Example 1: electrolysis of molten sodium chloride Cathode Anode Ions present: Na+(l) Cl-(l) Reaction, Na+(l) + e- → Na(l) 2Cl-(l) → Cl2(g) + 2e- equation and observations: Sodium ions gain Chloride ions loses electrons and is electrons and is reduced to form oxidised to form sodium atom. Grey chlorine molecules. globules of sodium Yellow green chlorine obtained at gas obtained at cathode. anode. Overall 2NaCl(l) → 2Na(l) + Cl2(g) equation:  The number of electrons leaving the electrolyte via the anode is always equal to the number of electrons entering the electrolyte from the cathode. This allows us to determine the overall equation for electrolysis and the mole ratio of products formed. Step 1: Write the half-equations At the cathode: Na+(l) + e- → Na(l) involved. At the anode: 2Cl-(l) → Cl2(g) + 2e- Step 2: Balance the number of electrons At the cathode: 2Na+(l) + 2e- → 2Na(l) gained or lost. At the anode: 2Cl-(l) → Cl2(g) + 2e- Step 3: Combine the half equations from 2Na+(l) + 2e- + 2Cl-(l) → 2Na(l) + Cl2(g) + 2e- step 2. Step 4: Cancel out the electrons on both 2Na+(l) + 2e- + 2Cl-(l) → 2Na(l) + Cl (g) + 2e- 2 sides of the equation. Step 5: Write the overall equation 2NaCl(l) → 2Na(l) + Cl2(g) From the overall equation, the ratio of Na to Cl2 produced is 2 : 1. Inert electrodes for molten electrolytes Contamination of the products can occur if the electrodes take part in the process. Hence, inert electrodes are used as they are made up of materials that are usually unreactive. electrode advantages disadvantages High melting point Graphite will react with oxygen gas under high Will not melt when used in temperatures to produce carbon dioxide. graphite the electrolysis of molten Graphite anodes might have to be periodically binary ionic compounds replaced. Lower melting point than graphite Might melt when used in the electrolysis of molten Does not take part in the platinum binary ionic compounds electrolysis reaction Mainly used in the electrolysis of aqueous electrolytes Try it: Question 1 Complete the table for the Cathode Anode electrolysis of molten lead(II) Ions present: Pb2+(l) Br-(l) bromide using inert electrodes. Reaction, Pb2+(l) + 2e- → Pb(l) 2Br-(l) → Br2(g) + 2e- equation and Lead observations: ________ ions gain Bromide ions loses ________ electrons and is electrons and is reduced __________ to form oxidised to form ________ lead atom. chlorine molecules. Grey globules _________________ Red-brown bromine __________________ lead of ______ obtained gas _________ obtained at cathode. at anode. Overall PbBr2(l) → Pb(l) + Br2(g) equation: Try it: Question 2 The diagram below shows an electrolytic cell set-up. Yellow-green gas is liberated at carbon electrode A. (a) Identify the gas at electrode A. chlorine (a) Which electrode is the positive electrode? A (a) What is the other product from the electrolysis? Molten calcium (a) Write an overall equation for this electrolytic reaction. CaCl2(l) → Ca(l) + Cl2(g) Aqueous solutions Aqueous solutions are formed when a solute dissolves in water. Water can be viewed as having its own reservoir of hydrogen ions (H+) and hydroxide ions (OH-) that comes from the reversible dissociation of water molecules. H2O(l) ⇌ H+(aq) + OH-(aq) Hence, when an aqueous ionic substance is used as the electrolyte, it will contain two types of cations and two types of anions. For example: ions aqueous solution from dissolved substance from water sodium chloride, NaCl(aq) Na+(aq), Cl-(aq) H+(aq),OH-(aq) Selective discharge of cations Metals can be arranged according to the reactivity series. The more reactive the metal, the more stable its ion and the harder it is to convert the ion back to the metal. For example, if a solution contains Cu2+ and Ag+ K+ Na+ ions, then Ag+ ions will be selectively discharged at Ca2+ the cathode. Cu2+ions begin to be discharged only Mg2+ after all the Ag+ ions have discharged. Zn2+ Increasing ease of Fe2+ discharge Pb2+ It is practically impossible to discharge metal ions H+ that are above hydrogen in the series from an Cu2+ aqueous solution since all aqueous solutions Ag+ contain H+ ions. Selective discharge of anions The electrochemical series below ranks anions in terms of their electrochemical reactivity. In dilute solutions, the lower the position of the anion, the more likely the anion will be selectively discharged. For example, if a solution contains Cl- and OH- ions, SO42- then OH- ions will have a higher tendency to be NO3- Cl- Increasing ease of selectively discharged. This means that, in dilute Br- discharge aqueous solutions, it is difficult to discharge any I- anion other than OH-. OH- The ionic half equation for the discharge of OH- ions is: 4OH-(aq) → O2(g) + 2H2O(l) + 4e-  The concentration of a solution affects the selective discharge of anions. A high concentration increases the ease of discharge of Cl-, Br - or I- anions. This can sometimes override the electrochemical series. For example, if a concentrated solution of sodium chloride (NaCl) is electrolysed, then the Cl- ions present can be selectively discharged over the OH- ions. Cl2 will be evolved at the anode instead of O2. The ionic half-equation for the discharge of Cl- ions is: 2Cl-(aq) → Cl2(g) + 2e- Determining the products of electrolysis of aqueous solutions Apart from the products obtained at the anode and cathode, it is also important to look at the anions and cations left behind in the solution after electrolysis. These ions make up the product remaining in the electrolyte. The following steps can be used to guide you: Step 1: Identify the ions present in the electrolyte. Step 2: Determine the anion discharged at the anode. Write the half-equation involved. Step 3: Determine the cation discharged at the cathode. Write the half-equation involved. Step 4: Write the overall equation for the reaction. Step 5: Identify the anion and cation left behind to determine the products in the electrolyte. Example 2: electrolysis of dilute aqueous sodium chloride Cathode Anode Ions present: Na+(aq),H+(aq) Cl-(aq), OH-(aq) Ion discharged: H+(aq) OH-(aq) Reaction, equation and 2H+(aq) + 2e- → H2(g) 4OH-(aq) → O2(g) + 2H2O(l) + 4e- observations: Since hydrogen is Since OH- ions are lower than below sodium in the Cl- in the anion electrochemical reactivity series, H+ series, OH- ions are selectively ions are selectively discharged to form colourless discharged to form oxygen gas. colourless hydrogen gas. Determine the overall equation for the electrolysis of dilute aqueous sodium chloride. Step 1: Write the half-equations At the cathode: 2H+(aq) + 2e- → H2(g) involved. At the anode: 4OH-(aq) → O2(g) + 2H2O(l) + 4e- Step 2: Balance the number of At the cathode: 4H+(aq) + 4e- → 2H2(g) electrons gained or lost. At the anode: 4OH-(aq) → O2(g) + 2H2O(l) + 4e- Step 3: Combine the half equations 4H+(aq) + 4e- + 4OH-(aq) → 2H (g) + O (g) + 2H O(l) + 4e- 2 2 2 from step 2. Step 4: Cancel out the electrons on 4H+(aq) + 4e- + 4OH-(aq) → 2H (g) + O (g) + 2H O(l) + 4e- 2 2 2 both sides of the equation. Step 5: Write the overall equation 4H+(aq) + 4OH-(aq) → 2H2(g) + O2(g) + 2H2O(l) 4H2O(l) → 2H2(g) + O2(g) + 2H2O(l) 2H2O(l) → 2H2(g) + O2(g)  2H2O(l) → 2H2(g) + O2(g) From the overall equation and the diagram, the electrolysis of dilute sodium chloride is equivalent to the electrolysis of water. The ratio of H2 to O2 produced is 2 : 1. The solution remains neutral as H+ and OH- ions are being discharged. The water level drops as the electrolysis continues. Thus, the concentration of the sodium chloride solution increases. If enough water is lost, the solution might become concentrated enough for Cl- ions to be selectively discharged. Let’s take a look at the electrolysis of water! Example 3: electrolysis of concentrated sodium chloride Cathode Anode Ions present: Na+(aq),H+(aq) Cl-(aq), OH-(aq) Ion H+(aq) Cl-(aq) discharged: Reaction, 2H+(aq) + 2e- → H2(g) 2Cl- (aq) → Cl2(g) + 2e- equation and observations: Since hydrogen is Concentration effects below sodium in the apply since the reactivity series, H+ concentration of Cl- ions are selectively ions is greater than OH- discharged to form ions. Cl- ions are colourless hydrogen selectively discharged gas. to form yellow green chlorine gas.  Combining the two half equations: 2H+(aq) + 2Cl-(aq) → H2(g) + Cl2(g) ❑ The ratio of H2 to Cl2 produced is 1 : 1. ❑ The solution becomes alkaline as there is a net discharge of H+ ions. The remaining Na+ and OH- ions form sodium hydroxide, an alkaline solution. ❑ When a few drops of Universal Indicator are added to the electrolyte, the Universal Indicator changes from green to violet. Try it: Question 3 For the electrolytes given, state the ions that will be discharged at the cathode and anode. electrolyte ion discharged at cathode ion discharged at anode dilute HCl(aq) H+(aq) OH-(aq) dilute H2SO4(aq) H+(aq) OH-(aq) concentrated HCl(aq) H+(aq) Cl-(aq) dilute KCl(aq) H+(aq) OH-(aq) Try it: Question 4 State the ions present, ions discharged at the respective electrodes and the half equations at each electrode when aqueous copper(II) sulfate is electrolysed. Cathode Anode Ions present: Cu2+(aq), H+(aq) OH-(aq), SO42-(aq) Ion discharged: Cu2+(aq) OH-(aq) Half equation: Cu2+(aq) + 2e- → Cu(s) 4OH-(aq) → O2(g) + 2H2O(l) + 4e- Red-brown solid formed on the cathode. Try it: Question 4 When copper(II) sulfate solution undergoes electrolysis: After Cu2+ and OH- ions are discharged, H+(aq) and SO42-(aq) ions remains in the solution. Hence, the resulting electrolyte becomes increasingly acidic as the concentration of H+ is greater than the concentration of OH-. The concentration of Cu2+ ions decreases thus, the blue colour of the electrolyte gradually fades and eventually turns colourless (decolourises).  13.3 How is electrolysis used in industries?  Do you know? The electricity used in our daily life is carried by wires made of electrical conductors. Copper is most commonly used because of its good conductivity and relatively low price. However, copper wires must be pure as the presence of impurities will reduce the efficiency of electrical transfer. Electrolysis is used to purify copper! Electrolysis using reactive electrodes Reactive electrodes, unlike inert electrodes, can take part in the electrolysis process. In industries, reactive electrodes are often used to introduce a metal cation into the electrolyte. ❑ When a reactive metal is used as the anode, it is oxidised to form its cation which may then be reduced back to the original metal at the cathode. Only reactive anode will react during electrolysis, but not the reactive cathode. The cathode “grows” at the expense of the anode. Example 4: electrolysis of copper(II) sulfate using copper electrodes Cathode Anode Ions present: Cu2+(aq), H+(aq) SO42-(aq), OH-(aq) Ion Cu2+(aq) – discharged: Reaction, Cu2+(aq) + 2e- → Cu(s) Cu(s) → Cu2+(aq) + 2e- equation and observations: Since copper is below Copper is not an inert hydrogen in the electrode. Hence, reactivity series, Cu2+ copper anode is ions are selectively oxidised instead. discharged to form red- brown copper. Copper The copper anode cathode increases in dissolves to form Cu2+ mass. ions and decreases in mass. ❑ The copper deposited onto the cathode comes mainly from the electrolyte. ❑ The copper anode is constantly oxidising and dissolving into the electrolyte, replenishing the Cu2+ ions which are reduced at the cathode. ❑ Thus, the concentration of Cu2+ ions remains unchanged, and the colour intensity of the blue electrolyte remains constant. Application of electrolysis 1: metal purification The raw, impure copper is the anode while pure copper is the cathode. The electrodes are placed in an electrolyte that contains Cu2+ ions, like aqueous copper(II) sulfate. Can inert anode be used for metal purification? If the anode is replaced with an inert electrodes, some pure copper can still be deposited at the cathode. This copper is produced from the Cu2+ ions in the electrolyte. The colour intensity of the electrolyte gradually fades from blue to colourless. The deposition of copper at the cathode ceases once the Cu2+ ions are used up. It is hard to purify metals that are higher than hydrogen in the reactivity series using aqueous electrolytes as H+ ions will be selectively discharged. Application of electrolysis 2: electroplating Electroplating allows us to coat a thin layer of metal onto an object. Cathode: object to be coated with the metal Anode: plating metal Electrolyte: aqueous solution of a salt of the plating metal Cathode Anode Reaction and Cu2+(aq) + 2e- → Cu(s) Cu(s) → Cu2+(aq) + 2e- observations: Cathode is coated with a Anode dissolves to form layer of copper metal Cu2+ ions and decreases and increases in mass. in mass. Uses of electroplating Non-conductive objects can also be electroplated but they will first need to be coated with a layer of graphite before being immersed into the electrolyte. There are many uses of electroplating. These includes: ❑ Electroplating enhances the attractiveness of metals ❑ Electroplating reactive metals can protect them from corroding easily Try it: Question 5 e- e- (a) Draw the battery in the set-up provided and draw an arrow on the connecting wire to indicate the direction of electron flow. (a) Can the spoon be made of materials such as plastic or wood? Explain your answer. No, the spoon should be made of a material with electrical conductivity. Alternatively, it should be coated with a layer of graphite. Try it: Question 5 c) Write the half equations at the cathode and the e- e- anode. At anode: Ag(s) → Ag+(aq) + e- At cathode: Ag+(aq) + e- → Ag(s) c) State and explain if the concentration of the electrolyte changes. Concentration of the electrolyte remains unchanged. As 1 mole of Ag(s) oxidised to 1 mole of Ag+(aq) at the anode, 1 mole of Ag+(aq) is reduced to form 1 mole of Ag(s) at the cathode. Thus, there is no net change in concentration of Ag+ ions.  13.4 What are simple cells and hydrogen fuel cells? Simple cell Similar to electrolytic cells, some batteries contain aqueous solutions that act as electrolytes. Batteries power a lot of our everyday items through electrochemical reactions. Batteries are also known as simple cells. A simple cell is a device that converts chemical energy into electrical energy. Structure of a simple cell A simple cell requires the following: ❑ two metals of different reactivities connected to an external circuit; and ❑ an electrolyte Like in electrolytic cells, oxidation occurs at the anode and reduction occurs at the cathode. What happens in a simple cell? more reactive metal less reactive metal acts as the anode acts as the cathode is oxidised forming causes cations from cations that enter the the electrolyte to gain electrolyte electrons and be releases electrons that reduced flow through the external circuit This is the result of the movement of electrons. In simple cells, the electrons move spontaneously from where they are produced (anode) to where they are gained (cathode). This movement of electrons produces electrical energy. Difference between simple and electrolytic cells simple cell electrolytic cell electrical energy is supplied source of electrical electrical energy is produced by an external source (e.g. energy through chemical reactions battery) electrons move from the electrons move from the battery to the cathode, electron movement anode to the cathode through the electrolyte, and into the anode polarity of anode: negative (–) anode: positive (+) electrodes cathode: positive (+) cathode: negative (–) Measuring the potential difference To measure the electrical energy produced, a voltmeter is used. A voltmeter will tell us the potential difference of the cell. Unit for voltage: volts (V) Electrodes in simple cells Metals are used as electrodes in simple cell. The further apart the two metals are in the reactivity series, the greater the voltage produced. When copper is used as the cathode, the voltage generated is larger when magnesium is the anode compared to when iron is the anode. No current will flow if both electrodes are made of the same metal. Try it: Question 6 Given the following voltages, match the voltage produced to the simple cells. 2.7 V 1.1 V 0.8 V 0.5 V 0.0 V metal electrodes voltage / V zinc & copper 1.1 copper & copper 0.0 iron & copper 0.8 lead & copper 0.5 magnesium & copper 2.7 Try it: Question 7 Which of the following pairs of metals would produce the largest voltage when used as electrodes in a simple cell? A. calcium and zinc B. copper and zinc C. magnesium and silver D. silver and iron Hydrogen fuel cells Hydrogen fuel cells are used to power electric vehicles. A fuel cell is similar to a simple cell. It derives its power from a fuel that is continuously added at the anode and an oxidiser (or oxidising agent) that is continuously added at the cathode. The hydrogen fuel cell utilises: ❑ hydrogen as the fuel; and ❑ oxygen from air as the oxidiser. Hydrogen fuel cells do not directly contribute to climate change because their only product is water, besides electricity. What happens in a hydrogen fuel cell? The ratio of H2 to O2 consumed is 2 : 1. Excess hydrogen can be pumped back into the anode chamber to reduce wastage. The electrolyte “completes” the circuit between the anode and cathode. Cathode Anode O2(g) + 2H2O(l) + 4e- → 4OH-(aq) H2(g) + 2OH-(aq) → 2H2O(l) + 2e- Overall equation: 2H2(g) + O2(g) → 2H2O(l) Advantages of using hydrogen as a fuel Hydrogen is a renewable fuel and can be obtained via the electrolysis of water or from the cracking of hydrocarbon. Hydrogen fuel cells produce only water as a by-product. It is more efficient than fuel-burning electricity sources. A larger percentage of the chemical energy stored in the fuel ends up as useful electricity in a fuel cell. Disadvantages of using hydrogen as a fuel It is difficult to store and transport hydrogen safely as hydrogen is a highly flammable gas at room temperature and pressure. Hydrogen is often transported in high- pressure cylinders which can be dangerous to handle. Large amount of energy is needed to produce hydrogen from electrolysis. However, this can be mitigated by using renewable sources like solar and wind energy to power electrolysis.