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  1. 1. CHEMISTRY ELECTROLYSIS Issue 1 11 Chemistry In this issue  Electrochemical cells  Parts of an electrochemical cell; electrolytes, salt bridge, electrodes.  Redox Reactions  Electrolytic cells  Applications of electrolysis From the lectures of Ms. Sadaf GulzarElectrochemistryElectrochemistryBy: Fatima Laraib By Fatima LaraibElectrochemistry is the study of the relations cathode of a galvanic cell is itsbetween electrical and chemical phenomena positive terminal. In both galvanicin terms of chemical changes produced by and electrolytic cells, oxidationelectrical current and the production of takes place at the anode andelectricity by chemical reactions electrons flow from the anode to the cathode.An electrochemical reaction is any chemicalreaction which is either caused or  The anode and cathode of an Parts of a cellaccompanied by the passage of an electric electrolytic cell are connected to a This section will deal with the components of ancurrent and involving the transfer of battery or other direct current electrochemical cell in particular. For e.g, the saltelectrons between two substances. source, whereas a simple cell serves bridge. as a source of electrical energy.Electrochemical Cells Page 2  In electrolytic cells, electricalElectrochemistry involves electrochemical energy from an external sourcecells which is a device capable of either causes non-spontaneous redoxderiving electrical energy from chemical reactions to occur, whereas inreactions or facilitating chemical reactions simple cells, spontaneous redoxthrough the introduction of electrical energy. reactions produce electricity.It is also a device where electron transfer isforced to take an external pathway instead  In electrolytic cells, the electrodesof going directly between the reactants. can be of the same or different metals but in simple cells, theThe differences between a simple cell and an electrodes have to be of differentelectrolytic cell are: elements only.  The anode of an electrolytic cell is  In electrolytic cell Ions are Redox Reactions positive (cathode is negative), since discharged on both the electrodes. This section will talk about the preliminary ideas behind the anode attracts anions from the While in a galvanic cell ions are only one of the most important electrochemical reactions. solution. However, the anode of a discharged at the cathode i.e. redox reactions. galvanic cell is negatively charged, since the spontaneous oxidation at  The electrons flow from the cathode Page 3 the anode is the source of the cells to anode in an electrolytic cell. This electrons or negative charge. The is not true in case of galvanic cells.
  2. 2. CHEMISTRY ELECTROLYSIS | Issue 1 2 Parts of an electrochemical cell Electrolytes The electrolyte is a molten ionic compound or an aqueous solution that conducts electricity. It provides the medium for transfer of ions inside the cell between the anode and cathode. It only conducts electricity because it contains charged particles (ions) that are mobile. It should be a non-conductor of electrons to avoid self-discharge of the cell. e.g:- acids, alkalis, aqueous solutions of salts, molten salts. Strong electrolytes KEY POINT!! ELECTRICAL & ELECTROLYTICAL The substances, which ionize almost CONDUCTION Weak electrolytes Metals and ionic compounds are conductors of completely into ions in aqueous phase In aqueous phase environments, the environments, are known as strong electricity. Ionic compounds in their molten or substances which ionize to a small extent aqueous states are ELECTROLYTICAL ELECTRICAL & called electrolytic electrolytes. Strong electrolytes fall into into ions are known as weak electrolytes. The weak electrolytes include weak conductors because they conduct electricity by CONDUCTION three categories: strong acids, strong acids and weak bases. In this case, the movement of ions across the electrolyte. bases, and salts. (Salts are sometimes Thus the electrolytes are decomposed to form also called ionic compounds) the molecules are in equilibrium with new substances. In case of electrical conduction, The equation which involves their ions. The equation which involves electricity is conducted by the flow of electrons the ionization of weak electrolytes is the ionization of strong electrolytes is from one end of the conductor to the other end. represented with double headed represented with only single headed Thus metals and graphite remain chemically arrows. For example, arrow directed to the right. For unchanged when an electric current flows + - example, CH3COOH + H2O ⇌ H3O + CH3COO through them. + HCl + H2O ———-> H3O + Cl releasing or sending oppositely The anode or negative electrode is the charged ions into the solution. reducing electrode. It gives up electrons to the external circuit and is oxidised Electrodes during the electrochemical (discharge) reaction. It is generally a metal or an alloy Electrodes are metallic or graphite but hydrogen is also used. The anodic terminals which are responsible for process is the oxidation of the metal bringing in or carrying out electric reducing agent to form metal ions. current within or outside the electrolytic cell. There are two types The cathode or positive electrode is the of electrodes, active electrodes and oxidising electrode. It accepts electrons inert electrodes. from the external circuit and is reduced during the electrochemical (discharge)Salt bridge Active electrodes reaction. It is usually a metallic oxide or a Electrodes like copper and silver sulfide but oxygen is also used. The which participate in the cathodic process is the reduction of theIt is the separator which electrically isolates the electrochemical processes are known oxidising agent (oxide) to leave the metal.positive and negative electrodes. It is a device as active electrodes.used to connect the oxidation and reduction half-cells of a galvanic cell. Salt bridges usually come Inert Electrodes Inert electrodes such as graphite orin two types: glass tube and filter paper. The salt platinum are non-reactive and arebridge: not affected by the ions surrounding a) Acts as a link between the two aqueous them in an electrochemical reaction. solutions. b) Overcomes liquid junction potential. c) Maintains the electrical neutrality of the aqueous solution of the electrodes by
  3. 3. CHEMISTRY ELECTROLYSIS | Issue 1 3 REDOX REAGENTS The picture shows the rusting of iron which is a redox reactionPotassium Iodide testPotassium Iodide changing its colour from colourlessto brown in the presence of an oxidising agent Redox Reactions Redox reactions Oxidation State Redox reactions describe all chemical The charge of an element in a compound reactions in which atoms have their with respect to the number of electrons oxidation number (oxidation state) it has lost or gained. changed or a reaction in which electrons are transferred, thereby oxidizing some Principles governing oxidation atoms, and reducing others. statesPotassium Manganate Test Oxidation 1. All the elements/molecules inPotassium manganate changes its colour from purple to A substance is said to be oxidized when their atomic or diatomic formscolourless in the presence of a reducing agent. carry a charge of 0. it enters into a combination with oxygen, 2. The oxidation state of a simple loses hydrogen, loses electrons or ion is the same as the charge on FAST FACTS increases its oxidation state after a the ion. reaction. 3. The oxidation states of the atoms present in the formula of Decomposition of compounds to form a compound must add up to 0. elements are redox reactions. Reduction A substance is said to be reduced when it 4. The total of the oxidation states A single replacement reaction is of a polyatomic ion is equal to always a redox reaction because it loses oxygen, gains hydrogen, gains its charge involves an element that becomes electrons or decreases its oxidation state 5. Oxygen in all its oxides will incorporated into a compound and an after a reaction. carry a charge of -2 where as in element in the compound being peroxides will carry a charge of released as a free element. -1. A double displacement reaction is Oxidising agent (oxidant) 6. Hydrogen when bonded with a usually not a redox reaction. Is a substance which oxidises another non-metal will carry a charge of Neutralization reactions, acid- +1 whereas with metals it will substance and itself is reduced by carbonate reactions and precipitation exhibit the -1 charge. reactions are usually non-redox. accepting electrons from the other 7. Group I, II and III elements will reactant. E.g:- carry a charge of +1,+2 and +3  Potassium dichromate (VI) consecutively.  Potassium Manganate (VII) A half-cell reaction is a reduction or an  Concentrated Sulfuric acid oxidation reaction component of a redox reaction. Reducing agent (reductant) FAST FACT Is a substance which reduces another A half reaction is obtained by considering In voltaic cells the oxidation and reduction the CHANGE in oxidation states of the substance and itself is oxidised. Good reactions are separated via half cells. INDIVISUAL substances involved in a redox donors of electrons E.g:- reaction.  Metals  Hydrogen The atoms and charges must be balanced on  Sulfer dioxide both the sides of the equation.  Potassium Iodide Equations can be combined to give an  Ammonia overall picture of the reaction or left as individual equations to illustrate the oxidation or the reduction steps separately.
  4. 4. CHEMISTRY ELECTROLYSIS | Issue 1 4 Electrolytic Cell The electrolyte which is the substance Voltaic cells use a spontaneous Electrolysis depends over: being electrolyzed must contain free chemical reaction to drive an electric current through an external moving ions. 1. The electrochemical series circuit. These cells are important 2. Molarity/concentration of Connection of the electrodes, to a source solution. because they are the basis for the of direct electric current renders one of 3. Types of electrodes. batteries that fuel modern society. them negatively charged and the other 4. Type of electrolyte; molten or But they arent the only kind of positively charged. Positive ions in the aqueous. electrochemical cell. It is also electrolyte migrate to the negative possible to construct a cell that electrode (cathode) and they combine does work on a chemical system by with one or more electrons in order to driving an electric current through gain their stable atomic state back. Thus the system. These cells are reduction occurs at the cathode. called electrolytic cells. Simultaneously, negative ions migrate to The redox reaction in an electrolytic the positive electrode (anode) and cell is nonspontaneous. Electrical transfer one or more electrons to it, losing energy is required to induce the their charge and becoming neutral electrolysis reaction. particles. Thus, oxidation occurs at the anode. Electrolysis refers to the decomposition of a substance by an The process of gaining or losing electrons electric current. It involves the at the electrodes is called discharge. reactions within the electrolyte and When ions are discharged at the the reactions at the surface of the electrodes they form atoms or ions. electrodes. Factors affecting discharge of ions in the Electrolyte The factors that influence the discharge of ions are 1. Relative positions of the ions in the reactivity series. 2. Concentration of the ions in the electrolytes. 3. Nature of the electrodes RULES FOR PREDICTING THE RULES FOR PREDICTING THE SELECTIVE SELECTIVE DISCHARGE FOR DISCHARGE FOR CATIONS ANIONS Positive ions from the metals lowest in the OH- ions from water are reactivity series are discharged at the preferentially discharged when the cathode in preference to any other ions solutions are dilute to form O2 present in the solution. Negative ions such as chloride, Ions of less reactive metals are bromide and iodide can be preferentially discharged. (e.g copper and preferentially discharged when silver) their concentrations are high enough when compared to OH – Otherwise, H+ ions from water will be discharged/reduced to form hydrogen gas. When sulfate and nitrate ions are present in water, it is the 2H+(aq) + 2eH2(g) hydroxide ions that are preferentially discharged. Ions of very reactive metals cannot be discharged in the presence of water.
  5. 5. CHEMISTRY ELECTROLYSIS | Issue 1 5 METAL EXTRACTION Applications of electrolysis Obtaining metals from their ores generally involve refining crude metal to obtain the pure metal. This is done via The numerous applications of electrolysis electrolysis. include: ELECTROPLATING In general the more reactive the metal The process of depositing a layer of is, the harder it is to extract the metal 1. Electrolytic purification metal on another substance using from its ore. Reactive metals such as 2. Electroplating electrolysis is called electroplating. sodium, potassium, calcium, magnesium 3. Metal Extraction and aluminum cannot be extracted by Both the anode and the cathode are reduction with Carbon. The compounds immersed in a solution which contains a of these metals are very difficult to split ELECTROLYTIC PURIFICATION dissolved metal salt (e.g., an ion of the up. Hence electricity is used to extract metal being plated) and other ions which these metals. Electrolysis is used to purify metals as in act to permit the flow of electricity the electrolytic purification of Copper. through the circuit. Direct current is CHLOR-ALKALI EXTRACTION supplied to the anode, oxidizing its Electrolysis of aqueous copper (II) metal atoms and dissolving them in the sulfate using copper electrodes. electrolyte solution. The dissolved metal ions are reduced at the cathode, plating At the Cathode the metal onto the item. The current through the circuit is such that Both Cu2+ and H+ ions are attracted to the rate at which the anode is dissolved the Cathode but Copper ions are is equal to the rate at which the cathode preferentially discharged and deposited is plated. on the Cathode as a brown layer of solid Hence: copper. The object to be plated is made Cu2+(aq)+2eCu(s) the cathode. The anode is the source of the At the Anode plating material. The electrolyte is an aqueous The chlorine-alkali (chloralkali) industry is an Both hydroxide and Sulfate ions are solution of a salt of the plating important part of the chemical industry, and attracted but NEITHER is DISCHARGED. material. produces chlorine and sodium hydroxide through the electrolysis of table salt The net result is the transfer of In fact, the copper anode dissolves to (NaCl). The main raw material is brine which is the plating material from the form Cu2+ ions in aqueous solution. a saturated solution of sodium chloride (NaCl) anode to the cathode. that is obtained from natural salt deposits (read Cu(s)  Cu2+(aq)+2e more at page 8) Electroplating of metals helps to prevent Overall reaction their corrosion and rusting and also improves their appearance. A common ALUMINIUM EXTRACTION 1. Cathode gains copper ions and example includes galvanizing; i.e. becomes larger. coating iron/steel with a layer of 2. No gas is evolved at the anode. another metal like zinc in order to In fact, the anode loses copper prevent rusting. ions and becomes smaller. Another example can be silver or gold 3. Concentration and colour of the plating which is used to coat a relatively copper(II) sulfate solution cheap metal to make it look more remains unchanged. expensive. 4. Amount of copper ions discharged to form copper Conditions favoring a good quality metal deposits at the cathode (from plating the solution) = Amount of The metal object to be plated copper atoms (from the anode) must be clean and free of Aluminum is extracted from the ore bauxite which ionizes and enters the (Al2O3) by electrolysis. The Al2O3 is insoluble, grease. solution as copper ions. so it is melted to allow the ions to move when The concentration of metal ions an electric current is passed through it. The in the electrolyte must be low. Hence to refine copper, impure copper is anodes are made from carbon and cathode is The electric current must be used as the anode. During electrolysis the carbon-lined steel case. not too large; otherwise the At the Cathode: Al3+(l)+3e-Al(l) the impure copper anode dissolves and a coating layer will form too At the anode: 2O2-(l) O2(g)+4e- layer of pure copper is deposited at the rapidly and peel off easily. The oxygen reacts with the carbon anodes to cathode. form carbon dioxide. The constantly need to be replaced because of this.
  6. 6. CHEMISTRY ELECTROLYSIS | Issue 1 6 ELECTROLYSIS OF Electrolysis of molten sodium As the hydrogen ions are lower in the chloride reactivity series than sodium, they accept BINARY COMPUNDS electrons more easily. The hydrogen ions Sodium chloride contains Sodium ions are discharged. Many ionic compounds are binary (Na+) and Chloride ions (Cl-). 2H+(aq)+2e-H2 compounds At the cathode: A binary compound is a compound At the anode: containing only two elements. It  Each Na+ ion gains one electrons Both the chloride ions and hydrogen ions contains a metal cation and a to form a sodium atom. It is migrate to the anode but the chloride ions non-metal anion. reduced are preferentially discharged because of their higher concentration. The electrolysis of a molten  Na+(l) + e- Na(l) 2Cl-(aq)Cl2(g)+2e binary compound will yield a Hence the colourless chlorine gas is given At the anode: metal and a non-metal as off. products.  Each Cl- ion gives up one electron As the hydrogen and chloride ions are to form a chlorine atom. It is discharged, sodium and hydroxide ions Electrolysis of molten lead bromide oxidised. Two chlorine atoms remain in the solution. The solution Lead(II) Bromide, PbBr2 is an ionic binary then combine to form a chlorine becomes sodium hydroxide. salt made of the ions Pb2+ and 2xBr- molecule. Electrolysis of aqueous Sodium  Cl-(l) Cl(l)+e- Chloride 2Cl(l) Cl2(g) An aqueous solution of sulfuric acid ELECTROLYSIS OF contains four different ions. Ions from NaCl: Na+(aq) and Cl-(aq) AQUEOUS SOLUTIONS Ions from water: H+(aq) and OH-(aq) An aqueous solution of a compound is a At the cathode: mixture of two electrolytes. The hydrogen and sodium ions are Ions discharged depend on their position in attracted to the cathode. H+ ions gain the electrochemical series. electrons from the cathode to form hydrogen gas. Electrolysis of Concentrated 2H+(aq)+2e-H2(aq) Sodium Chloride (Brine) Sodium ions remain in solution. The electrolytic cell used for electrolysis At the cathode: of concentrated sodium chloride solution At the anode: 2+ is designed to collect gaseous products The hydroxide and chloride ions are  Pb ions gain electrons from the attracted to the anode. Hydroxide ions at both electrodes. electrode to become lead atoms. give up electrons to the anode to form water and oxygen gas.  The Pb2+ ions are reduced 4OH-(aq)2H2O(l)+O2(g)+4e-  The Pb2+ have been discharged and molten greyish globules of Chloride ions remain in the solution. lead metals are formed below the electrolyte Hence the overall reaction is 2H20(l) 2H2(g)+O(g)  Electrode reaction at cathode: Pb2+(l)+2e Pb(l) Since water is being removed by decomposition into hydrogen and oxygen, At the anode: the concentration of sodium chloride solution increases gradually.  Br- ions lode electrons to become With electrolysis of aqueous solutions of bromine molecules. dilute acids or alkalis, the volume of The cathode can be platinum or hydrogen given off at the cathode is twice  The Br- ions are oxidised. carbon but the anode must be that of the oxygen gas at the anode. carbon to resist attack by chlorine.  Bromide ions are discharged NOTE: The test tube containing hydrogen forming an effervescence of Four ions are present in the mostly (more than half) appears to be pungent, red-brown bromine gas. solution: empty as compared to the one containing Cations: H+(aq) and Na+(aq) oxygen because hydrogen is completely  Electrode reaction at the anode: Anions: OH-(aq) and H+(aq insoluble in water while oxygen is slightly 2Br-(l) Br2 (g) + 2e soluble. At the cathode: The sodium and hydrogen ions move to the cathode.