Reactivity 3.2 Electron Transfer Reactions PDF

Reactivity 3.2 Electron transfer reactions Date Created @October 30, 2024 5:38 PM 3. 2. 1 Oxidation and reduction Oxidation and reduction can be defined in several ways 1. Oxidation and reduction in te...

Reactivity 3.2 Electron transfer reactions Date Created @October 30, 2024 5:38 PM 3. 2. 1 Oxidation and reduction Oxidation and reduction can be defined in several ways 1. Oxidation and reduction in terms of oxygen gain/loss Oxidation is the reaction where a substance combines with oxygen. Reduction is the reaction where oxygen is removed from a substance 2. Oxidation and reduction in terms of hydrogen loss/gain Oxidation is the reaction where hydrogen is removed from a substance. reduction is the reaction where a substance combines with hydrogen. 3. Oxidation and reduction in terms of electron transfer Oxidation is losing an electron. Reduction is gaining an electron. One of the examples is the redox reaction in optometry (p. 583) 4. Oxidation and reduction in terms of oxidation state change The oxidation state represents the charge that an atom would have in a compound if the coumpound were composed of ions. Compound is oxidized if its oxidation state increases. It causes another specie to be reduces, so it is also a reducing agent. Reactivity 3.2 Electron transfer reactions 1  Compound is reduced if its oxidation state decreases. It causes another specie to be oxidized, so it is also a exidizing agent. 3. 2. 2 Half-equations Any redox reaction can be separated into two equations: reduction and oxidation. For example, reaction 2Na + Cl2 − > 2NaCl ​ can be separate into Na− > Na+ + e− Cl2 + 2e− − > 2Cl − ​ This may help to balance the equation of the redox reaction. Steps to balance the redox reaction are as follows 1. Identify the species being oxidized and reduced. 2. Separate the reaction into oxidation and reduction half-equations. 3. Balnce the equations. 4. Balance the half-equations so that the numbers of electrons from them are equal. 5. Add two half-equations together and cancel electrons. if the reaction is accuring in acidic solution, add H_2O to balance any oxygen and H^(+) to balance any hydrogen atom. 6. For neutral or basic solutions add OH^(-) to balance oxygen atoms and H_2O to balance hydrogen atoms. 3. 2. 3 Reactivity 3.2 Electron transfer reactions 2  Oxidation and reduction of metals and halogenes Relative ease of reduction of halogens Halogens can act as the oxidising agents. The reactivity of the halogens increases going up the group. Therefore, the weaker specie won’t oxidize the stronger specie. Relative ease of oxidation of metals Group 1 metal can act as reducing agent. The reactivity of the group 1 metals increases down the group. 3. 2. 4 Redox reactions of acids and metals Reactive metals (group 1 nd 2) are readily oxidized by strong acids producing salt and hydrogen gas where the metal is the reducing agent and the acid is the oxidising agent. 3. 2. 5 and 3. 2. 6 Electrochemical cells The electrochemical cell interconverts electrical and chemical energy. There are two type of electrochemical cells: Primary (voltaic) cells, secondary (rechargeable) cells and fuel cells - the chemical energy is converted into the electrical energy. Reactivity 3.2 Electron transfer reactions 3  Electrolytic cells - the electrical energy is converted into the chemical energy to drive the nonspontanous chemical reactions. How does it work? In redox reaction electrons flow from the oxidized sybstance to the reduced one. Most of redox reactions are exotherim, and the excess energy can be used to produce electrical energy. Redox reactions used in the primary cells are irreversible, but reactions in secondary cells can be reversed. Any two metals can produce the redox reaction. The specie higher in the activity series (in the booklet) is more likely to lose electrons and therefore to be oxidized. The species localised lower is rather reduced, as it gains electrons. These two processes can occur in two separated beakers called, half-cells. They cna be connected with a wire to form an electrochemical cell. Oxidation always occurs at the anode (on the left). Reduction always occurs at the cathode (on the right) RED CAT. This electrochemical cell can be represent by the cell diagram Reactivity 3.2 Electron transfer reactions 4  When the half-cells are connceted electrons from the Zn oxidation flow through the wire and reduce Cu. As a result Zn half-cell is slightly positive and Cu half-cell is slightly negative. Any further oxidation is prevented by these charges (positive charge doesnt want to be more positive and negative doesn’t want to be more negative). Copper repels electrons which remain in the Cu half-cell. Cell becomes polarized and the redox reaction stops. To get around this two cells are connected by the salt bridge. Salt bridge consist of the ionic salt solution. It allows the positive ions of Zn to flow towards the negatively charged Cu half-cell, and negative ions of Cu to flow towards the positively charged Zn half-cell. 3. 2. 7 Secondary cells A battery is a series of two or more electrochemical cells. In the primary cell at one point all materials will be used and the reaction is irreversable. It is because anothere or cathode are fully oxidized or reduced. Additionally the transfer of ions throught the salt bridge polarizes the cell and causes the chemical reaction to stop. Furthermore, polarization can cause a build-up of hydrogen bubbles on the surface of the anode. This can increase the internal resistance of the cell and reduce the electrical output. In the secondary (rechargeable) cell the chemical reaction that generates the electrisity can be reversed by applying an electric current to the cell. Thy can satisfy higher current demands, but they also have a higher rate of self- Reactivity 3.2 Electron transfer reactions 5  discharge. When the current is not applied the direction of electron flow is the same as in the primary cell. When current is applied this direction is reversed. Case study: Lead-acid battery Car batteries are made of secondary bateries. The electrical energy is used to power the engine and other systems. This is known as discharge. Some of the chemical energy from the combusion in the engene is used to reverse the chemical reaction, recharging the battery. Typical car uses a lead-acid battery, composed of lead anode Pb(s) and lead(IV) oxide cathode PbO2(s) and sulfuric acid H2SO4(aq). Sulfuric acid will ionize into hydrogen ions H^(+) and HSO4^(-). HSO4^(-) will oxdize Pb at the anode and H^(+) will reduce PbO2 at the cathode. Case study: Lithium-ion batteries Lithium-ion batteries use lithium atoms embedded in a lattice of graphite electrodes at the anode. The cathode is the lithium-cobalt oxide complex LiCoO2. The overal reaction during the discharge is Reactivity 3.2 Electron transfer reactions 6  The battery must be completely non-aqueous as Li reacts vigorously with water. When the battery is discharging electrons flow from anode to cathode through the external circut while lithium ions flow from the anode to the cathode through the polymer gel inside the cell. Fuel cells Fuel cells are electrochemical cells that convert hydrogen, methanol or ethanol and oxygen into water, carbon dioxide and heat. They cause little polution and are very efficient. They are irrechargable but they need constant supply of fuel and oxygen. In hydrogen fuel cell, hydrogen gas is supplied to the anode while oxygen gas is supplied to the cathode Hydrogen fuel cells use hydrogen gas as the fuel and don’t produce greenhouse gases. However, the hydrogen gas used must be very pure to prevent the poisoning of the catalyst. There are two main sources of the hydrogen: Product of the electrolysis of water. Reaction of hydrocarbons with steam. In the direct methanol fuel cell (DMFC) methanol is supplied to the anode Reactivity 3.2 Electron transfer reactions 7  A typical fuel cell has following key components Electrolyte or separator: prevents components from mixing. Electrodes: they are made of of catalyst that allow the reaction to occur. Anode is the oxidizing electrode and cathode is the reducing electrode. These catalysts are often made of expensive metals such as platinum Bipolar plate: Cunducts the electrical current from cell to cell and ensures uniform distribution of the fuel gas. 3. 2. 8 Electrolytic cells Electrolytic cells convert electrical energy into the chemical energy. The oxidation and reduction in these cells are not spontanious so they require the outer electrical energy. This proces is known as electrolysis. It consists of single contained filled with en electrolyte (solution of the ionic salt, mother salt or free moving cations and anions). Cathode and anode are dipped in the electrolyte and the direct current (DC) is connected to the electrodes. Reactivity 3.2 Electron transfer reactions 8  Electrons flow from the negative terminal of the DC connected to the cathode to the positive terminal connected to the anode. Therefore electrons from from the anode to the cathode (as on the picture) and reduce the cations in the electrolyte. The anions flow to the anode and undergo oxidation. 3. 2. 9 Oxidation of organic compounds 3. 2. 10 Reduction of organic compounds 3. 2. 11 Reduction of alkenes and alkynes 3. 2. 12 Standard electrode potentials Reactivity 3.2 Electron transfer reactions 9  Standard electron (reduction) potential E describes the ease of oxidation and reduction of a species in an electrochemical cell. It is defined relative to the hydrogen-based half-cell This cell is known as the standard hydrogen electrode (SHE). It has assigned potential of 0 under curtain conditions: T=298 K. p=100kPa. electrode uses inert platinum (or carbon, something inert) electrode. Solution contains 1M H^(+). Species with a more negative E (higher in the reactivity series) will have greater ease in oxidation. 3. 1. 12 Standard cell potentials Standard cell potwntial E is the voltage of the electrochemical cell For the reaction to be spontaneous E must be positive so Reactivity 3.2 Electron transfer reactions 10  3. 2. 14 Gibbs energy and standard cell potential The standard change in Gibbs energy of the electrochemical cell can be determined from the equation Where n is number of electrons transfered and F is the Faraday constant. Positive E and negative G contribute to the spontanous reaction. 3. 2. 15 Electrolysis of aqueous solutions The electrolysis of aqueous solutions of ionic salt introduce redox reactions involving water. If E of the salt cation is more negative then E of water, the water will be reduced instead of the salt cation and hydrogen gas will be formed at the cathode. The oxidation of water is as follows Reactivity 3.2 Electron transfer reactions 11  If the oxidation potential of salt anion is more negative then E of water oxidation the water will be oxidised instead of the salt anion and oxygen gas will be formed at the anode. 3. 2. 16 Electroplating If the electrodes are made from a reactive metal, the electrolysis of an aqueous solution of an ionic salt will add material to the cathode and take material away from the anode. There processes are known as plating and eroding, respectively. Reactivity 3.2 Electron transfer reactions 12