The following are high-level thinking skills (KBAT) questions designed based on the learning outcomes provided:
8 Electrochemistry
8.1 Half-cell and redox equations
– Define the redox reaction?
A redox (or oxidation-reduction) reaction is a type of chemical reaction that involves a transfer of electrons between two species. An oxidation-reduction reaction is any chemical reaction in which the oxidation number of a molecule, atom, or ion changes by gaining or losing an electron. Redox reactions are common and vital to some of the basic functions of life, including photosynthesis, respiration, combustion, and corrosion or rusting.
– Draw the baseball analogy of redox reaction?
– Define electrochemistry?
Electrochemistry is the study of chemical processes that cause electrons to move. This movement of electrons is called electricity, which can be generated by movements of electrons from one element to another in a reaction known as an oxidation-reduction (“redox”) reaction.
– How do you determine which is oxidized and reduced?
Identifying the Elements Oxidized and Reduced
1. Assign oxidation numbers to all atoms in the equation.
2. Compare oxidation numbers from the reactant side to the product side of the equation.
3. The element oxidized is the one whose oxidation number increased.
4. The element reduced is the one whose oxidation number decreased.
– What is electrochemistry used for?
Electrochemistry has a number of different uses, particularly in industry. The principles of cells are used to make electrical batteries. In science and technology, a battery is a device that stores chemical energy and makes it available in an electrical form.
– What is the different between reactivity series and electrochemical series?
The reactivity series is a list of metals (often with hydrogen and carbon also included) based on their ability to reduce other chemicals. The electrochemical series is built up by arranging various redox equilibria in order of their standard electrode potentials (redox potentials).
– S E K vs S K L
Ramai yang confius bila nak explain guna SEK atau SKL kan. Sebagai contoh, terangkan tindak balas berikut :
1. Tindak balas larutan kuprum(II) sulfat dan serbuk zink.
❌ Zink lebih reaktif daripada kuprum.
✔ Zink lebih elektropositif daripada kuprum. Kedudukan zink lebih tinggi daripada kuprum dalam Siri Elektrokimia.
2. Tindak balas kuprum(II) oksida dan serbuk zink.
❌ Zink lebih elektropositif daripada kuprum.
✔ Zink lebih reaktif daripada kuprum. Kedudukan zink lebih tinggi daripada kuprum dalam Siri Kereaktifan Logam.
***
Yang ❌ tu pernyataan yang salah.
Yang ✔ tu pernyataan yang sepatutnya.
SIRI KEREAKTIFAN LOGAM
– susunan berdasarkan kereaktifan logam/ hidrogen/ karbon bertindak balas dengan oksigen.
SIRI ELEKTROKIMIA
– susunan berdasarkan keupayaan logam/ hidrogen untuk membebaskan elektron dan menjadi ion.
– What is dry cell?
A dry cell is a type of electric battery, commonly used for portable electrical devices. It was developed in 1886 by the German scientist Carl Gassner, after development of wet zinc-carbon batteries by Georges Leclanché in 1866. The modern version was developed by Japanese Yai Sakizo in 1887. The dry cell is one of many general types of electrochemical cells. A dry cell has the electrolyte immobilized as a paste, with only enough moisture in it to allow current to flow. Unlike a wet cell, a dry cell can operate in any orientation without spilling, as it contains no free liquid. Compared to other voltaic cells, the electrolyte in a dry cell is not liquid. Hence the name dry cell. In a lead acid battery, the electrolyte is dilute sulphuric acid. In case of dry cell, the electrolyte is a paste of ammonium chloride (in the old days). Dry Cell : It is a primary cell based on Leclanche cell invented by G. Leclanche in 1868. In a primary cell, the electrode reactions cannot be reversed by an external source of electrical energy. In this cell, the cell reaction takes place only once i.e., this cell is not rechargeable.
– What are the types of electrochemistry?
There are two types of electrochemical cells: galvanic, also called Voltaic, and electrolytic. Galvanic cells derives its energy from spontaneous redox reactions, while electrolytic cells involve non-spontaneous reactions and thus require an external electron source like a DC battery or an AC power source.
– What is a galvanic or voltaic cell?
A galvanic cell or voltaic cell is a device in which a redox reaction. A voltaic cell is an electrochemical cell that uses a chemical reaction to produce electrical energy. The important parts of a voltaic cell: The anode is an electrode where oxidation occurs.
– What is an electrolytic cell?
An electrolytic cell is a kind of electrochemical cell. It is often used to decompose chemical compounds, in a process called electrolysis—the Greek word lysis means to break up. Important examples of electrolysis are the decomposition of water into hydrogen and oxygen, and bauxite into aluminium and other chemicals. An electrolytic cell is a cell in which electrical energy from an external source drives a non-spontaneous reaction. This is different from the spontaneous redox reactions that occur in voltaic cells. An electrolytic cell has a positive ΔG and a negative Ecell.
– Is an electrolytic cell a battery?
An Electrolytic cell is one kind of battery that requires an outside electrical source to drive the non-spontaneous redox reaction. Rechargeable batteries act as Electrolytic cells when they are being recharged.
– What is the difference between Daniell cell and galvanic cell?
A galvanic cell is also referred to as a voltaic cell or Daniell cell. The common household battery is an example of a galvanic cell. The flow of electrons from one chemical reaction to another happens through an outside circuit resulting in current. In a Daniell cell electrons flow from zinc electrode to copper electrode to copper electrode through an external circuit, while metal ions form one half cell to the other through the salt bridge. Daniell cell is a reversible cell while a voltaic cell may be reversible or irreversible.
– What is present in a Daniell cell?
The Daniell cell is a type of electrochemical cell invented in 1836 by John Frederic Daniell, a British chemist and meteorologist, and consists of a copper pot filled with a copper (II) sulfate solution, in which is immersed an unglazed earthenware container filled with sulfuric acid and a zinc electrode.
– How does Daniell cell work?
The principle behind the Daniell cell is redox reaction. In the process of the reaction, electrons can be transferred from the corroding zinc to the copper through an electrically conducting path as a useful electric current. The principle used in daniell’s battery & Volta’s battery is the same.
– What is Volta’s battery, column type?
It is made of copper and zinc parts, separated by felt disks soaked in an acid solution. It is supplied with a bottle of acid solution.
– What is a lead acid battery and how does it work?
Each cell is made up of a set of positive and negative plates immersed in a dilute sulfuric acid solution known as electrolyte, and each cell has a voltage of around 2.1 volts when fully charged. The six cells are connected together to produce a fully charged battery of about 12.6 volts.
– What is the Nernst equation used for?
Nernst Equation. The Nernst Equation enables the determination of cell potential under non-standard conditions. It relates the measured cell potential to the reaction quotient and allows the accurate determination of equilibrium constants (including solubility constants).
– What is voltaic cell notation?
Cell notation or line notation in chemistry is a shorthand way of expressing a certain reaction in an electrochemical cell. The cell anode and cathode (half-cells) are separated by two bars or slashes representing a salt bridge, with the anode on the left and cathode on the right.
– Give a general of cell diagram (cell notation) of the Daniell cell?
The cell anode and cathode (half-cells) are separated by two bars or slashes representing a salt bridge, with the anode on the left and cathode on the right. This cell is very famous: the Daniell cell. If the electrodes are connected, a spontaneous reaction takes place.
– How to remember reduction cathode, oxidising anode?
First, we have to know the type of cell. In galvanic cell, we determine,
Oxidation (lose electrons) – anode – negative. An Ox Negative.
Reduction (gain electrons) – cathode – positive. Red Cat Positive.
Electron flow from anode to cathode.
An ox is negatively gives electrons to a positive red cat.
But, when we use electrolysis cell, we determine,
Oxidation (lose electrons) – anode – positive. An Ox Positive.
Reduction (gain electrons) – cathode – negative. Red Cat Negative.
Electron flow from anode to cathode.
An ox is positively gives electrons to a negative red cat.
At the anode, is oxidation, where the chemicals lose electrons and at the cathode is the reduction where the chemicals gain electrons. Remember LEO the lion goes GER. Every electrochemical cell has a cathode and an anode.
Reduction happens at the cathode, oxidation happens at the anode.
Positive is anode, negative is cathode.
Electron flow from anode to cathode.
– Why don’t directly attach the negative cable to the negative terminal while jumping a car battery?
Caution: Don’t attach the negative cable to the negative terminal of the weak battery when jumping a car battery! This common mistake could ignite hydrogen gas directly over the battery. Battery explosions can cause serious injury. … Finally, remove the positive cable from the car with the weak battery.
– With give several of specific examples, construct redox equations?
Oxidation-Reduction or “redox” reactions occur when elements in a chemical reaction gain or lose electrons, causing an increase or decrease in oxidation numbers. The Half Equation Method is used to balance these reactions. The equation is separated into two half-equations, one for oxidation, and one for reduction.
– How to balance simple redox equations?
Follow these rules to balance simple redox equations:
1. Write the oxidation and reduction half-reactions for the species that is reduced or oxidized.
2. Multiply the half-reactions by the appropriate number so that they have equal numbers of electrons.
3. Add the two equations to cancel out the electrons.
8.2 Standard electrode potential
– What is the standard cell potential E cell?
The potential of the cell under standard conditions (1 M for solutions, 1 atm for gases, pure solids or liquids for other substances) and at a fixed temperature (25°C) is called the standard cell potential (E°cell). Only the difference between the potentials of two electrodes can be measured.
– What is the electric potential of a cell?
The cell potential, Ecell, is the measure of the potential difference between two half cells in an electrochemical cell. The potential difference is caused by the ability of electrons to flow from one half cell to the other.
– How to determine stability ions from standard reduction potential half cell?
Oxidising agent with more positive E value, will be reduced. So, the oxidation number will decreased.
Redusing agent with more negative E value, will be oxidised. So, the oxidation number will increased.
– What is the driving force in an electrochemical reaction?
The transfer of electrons through the external wire create a current that can do work. The driving force pushing the electrons through the wire is the difference in the attraction for electrons in the two half-cells. This voltage difference is called the cell potential (Ecell) and is measured in volts.
– What does it mean when Ecell is positive?
A positive value indicates the oxidation-reduction reaction is a spontaneous reaction. That means without the help of an external agency. Here reduction at the cathode and oxidation at the anode takes place. If it is a negative value, it means only the reverse reaction is spontaneous.
– What does a positive voltage indicate about the oxidation reduction reaction?
A positive voltage that forms across the electrodes of a voltaic cell indicates that the oxidation-reduction reaction is a spontaneous reaction for reduction at the cathode and oxidation at the anode.
– Is E cell always positive or negative Ecell?
In order for delta G to be negative, which indicates that the reaction is a spontaneous one, E cell must be positive. For electrolytic cells, which are reactions that occur only with the input of an external energy source, E cell is negative because they are non-spontaneous.
– What increases cell voltage?
In an electrochemical cell, increasing the concentration of reactants will increase the voltage difference, as you have indicated. A higher concentration of reactant allows more reactions in the forward direction so it reacts faster, and the result is observed as a higher voltage.
– Draw and label the example of a standard hydrogen electrode, SHE diagram?
– Explain Standard hydrogen electrode?
Standard hydrogen electrode. The standard hydrogen electrode (abbreviated SHE), is a redox electrode which forms the basis of the thermodynamic scale of oxidation-reduction potentials. Hydrogen electrode is based on the redox half cell:
2 H+ (aq) + 2 e− –> H2(g)
– State the characteristic of Standard hydrogen electrode electrolyte pH, act as anode/ cathode, temperature, standard reduction potential?
Electrolyte pH = 0,
Can act as both anode cathode,
temperature is 25oC,
standard reduction potential is 0 V.
– How does pH affect reduction potential?
Reduction potentials are the energy of a half reaction per electron transferred. pH is related to the concentration of protons/ hydroxide ions in solution, so this is why it can modify the driving force of reactions involving protons/ hydroxide ions, and thus the reduction potential as well. But, in half cell other than H+ and OH- is independent of pH.
– How do you determine which is a stronger reducing agent?
Reducing agents can be ranked by increasing strength by ranking their reduction potentials. The reducing agent is stronger when it has a more negative reduction potential and weaker when it has a more positive reduction potential.
– Which has the highest reduction potential?
Fluorine gas is one of the best oxidizing agents there are and it is at the top of the table with the biggest most positive standard potential (+2.87 V). Reducing Agents: At the other end, are reactions with negative standard potentials. This means that the desired path of the reaction is actually the reverse reaction.
– Is reduction potential positive or negative?
Negative values of E∘ (Reduction potential) will lead to positive values of ΔG and vice versa. ΔG is positive, the reaction is non-spontaneous, and when ΔG is negative, the reaction is spontaneous.
– With give several of specific examples, show the use the standard hydrogen electrode to determine standard electrode potential (standard reduction potential), Eº?
By setup using galvanic cell.
– With give several of specific examples, show how to calculate the standard cell potential using the E° values, and write the redox equations? How do you calculate the potential of a cell?
At the standard state.
Write the half-reactions for each process.
Look up the standard potential for the reduction half-reaction.
Look up the standard reduction potential for the reverse of the oxidation reaction and change the sign.
Add the cell potentials to get the overall standard cell potential.
– How to Calculate Emf?
– What is the overall cell potential’s formula?
The half-cell with more positive Eo will form the cathode (positive terminal), while the half-cell with less positive Eo will form the anode (negative terminal).
E°cell = E°reduction −E° oxidation
E°cell = E°cathode −E° anode
The emf of a cell must be positive. Otherwise no current will flow
– How is E cell value calculated?
E cell = E reduced – E oxidised = + 0.34 V – (-0.76 V) = + 1.10 V. As the E cell value calculated is positive, the reaction is thermodynamicaly spontaneous.
– How to calculate standard reduction potential?
The standard reduction potential can be determined by subtracting the standard reduction potential for the reaction occurring at the anode from the standard reduction potential for the reaction occurring at the cathode. The minus sign is necessary because oxidation is the reverse of reduction.
– What does the negative value of E cell indicate?
A positive value indicates the oxidation-reduction reaction is a spontaneous reaction. If it is a negative value, it means only the reverse reaction is spontaneous. It means oxidation at the cathode and reduction at the anode.
– Is E cell always positive or negative Ecell?
In order for delta G to be negative, which indicates that the reaction is a spontaneous one, E cell must be positive. For electrolytic cells, which are reactions that occur only with the input of an external energy source, E cell is negative because they are nonspontaneous.
– What is the standard reduction potential?
The standard reduction potential is the reduction potential of a molecule under specific, standard conditions. Standard reduction potentials can be useful in determining the directionality of a reaction. The reduction potential of a given species can be considered to be the negative of the oxidation potential.
– How to measure standard reduction potential?
A galvanic cell can be used to determine the standard reduction potential. The standard reduction potential can be determined by subtracting the standard reduction potential for the reaction occurring at the anode from the standard reduction potential for the reaction occurring at the cathode.
– What does a higher reduction potential mean?
A solution with a higher (more positive) reduction potential than the new species will have a tendency to gain electrons from the new species (i.e. to be reduced by oxidizing the new species) and a solution with a lower (more negative) reduction potential will have a tendency to lose electrons to the new species (i.e. to be oxidized by reducing the new species).
– How we would predict the power of oxidising and reducing agents from Eº values?
The species at the top left have the greatest “potential” to be reduced, so they are the strongest oxidizing agents.
The strongest oxidizing agent in the list is F2, followed by H2O2, and so on down to the weakest oxidizing agent, Li+.
Standard Reduction Potential Table
– How the stability of aqueous ions from Eº values would be predicted? How we would predict the feasibility of a reaction from ‘E note cell’ value, E°cell and from the combination of various electrode potentials: spontaneous and non-spontaneous electrode reactions? What does the negative value of E cell indicate? Spontaneous change?
A positive value indicates the oxidation-reduction reaction is a spontaneous reaction. If it is a negative value, it means only the reverse reaction is spontaneous. It means oxidation at the cathode and reduction at the anode.
8 Electrochemistry
8.1 Half-cell and redox equations
Candidates should be able to:
(a) explain the redox processes and cell diagram (cell notation) of the Daniell cell;
(b) construct redox equations.
8.2 Standard electrode potential
Candidates should be able to:
(a) describe the standard hydrogen electrode;
(b) use the standard hydrogen electrode to determine standard electrode potential (standard reduction potential), Eº;
(c) calculate the standard cell potential using the Eo values, and write the redox equations;
(d) predict the stability of aqueous ions from Eº values;
(e) predict the power of oxidising and reducing agents from Eº values;
(f) predict the feasibility of a reaction from ‘E note cell’ value and from the combination of various electrode potentials: spontaneous and non-spontaneous electrode reactions.
8.3 Non-standard cell potentials
Candidates should be able to:
(a) calculate the non-standard cell potential, Ecell, of a cell using the Nernst equation.
8.4 Fuel cells
Candidates should be able to:
(a) describe the importance of the development of more efficient batteries for electric cars in terms of smaller size, lower mass and higher voltage, as exemplified by hydrogen-oxygen fuel cell.
8.5 Electrolysis
Candidates should be able to:
(a) compare the principles of electrolytic cell to electrochemical cell;
(b) predict the products formed during electrolysis;
(c) state the Faraday’s first and second laws of electrolysis;
(d) state the relationship between the Faraday constant, the Avogadro constant and the electronic charge;
(e) calculate the quantity of electricity used, the mass of material and/ or gas volume liberated during electrolysis.
8.6 Applications of electrochemistry
Candidates should be able to:
(a) explain the principles of electrochemistry in the process and prevention of corrosion (rusting of iron);
(b) describe the extraction of aluminium by electrolysis, and state the advantages of recycling aluminium;
(c) describe the process of anodisation of aluminium to resist corrosion;
(d) describe the diaphragm cell in the manufacture of chlorine from brine;
(e) describe the treatment of industrial effluent by electrolysis to remove Ni2+, Cr3+ and Cd2+;
(f) describe the electroplating of coated plastics.
8 Electrochemistry
analogy
8.1 Half-cell and redox equations
Candidates should be able to:
(a) explain the redox processes and cell diagram (cell notation) of the Daniell cell;
(b) construct redox equations.
8.2 Standard electrode potential
Candidates should be able to:
(a) describe the standard hydrogen electrode (SHE);
(b) use the standard hydrogen electrode (SHE) to determine standard electrode potential (standard reduction potential), Eº;
(c) calculate the standard cell potential using the Eo values, and write the redox equations;
(d) predict the stability of aqueous ions from Eº values;
(e) predict the power of oxidising and reducing agents from Eº values;
(f) predict the feasibility of a reaction from ‘E note cell’ value and from the combination of various electrode potentials: spontaneous and non-spontaneous electrode reactions.
8.3 Non-standard cell potentials
Candidates should be able to:
(a) calculate the non-standard cell potential, Ecell, of a cell using the Nernst equation.
8.4 Fuel cells
Candidates should be able to:
(a) describe the importance of the development of more efficient batteries for electric cars in terms of smaller size, lower mass and higher voltage, as exemplified by hydrogen-oxygen fuel cell.
8.5 Electrolysis
Candidates should be able to:
(a) compare the principles of electrolytic cell to electrochemical cell;
(b) predict the products formed during electrolysis;
(c) state the Faraday’s first and second laws of electrolysis;
(d) state the relationship between the Faraday constant, the Avogadro constant and the electronic charge;
(e) calculate the quantity of electricity used, the mass of material and/or gas volume liberated during electrolysis.
8.6 Applications of electrochemistry
Candidates should be able to:
(a) explain the principles of electrochemistry in the process and prevention of corrosion (rusting of iron);
(b) describe the extraction of aluminium by electrolysis, and state the advantages of recycling aluminium;
(c) describe the process of anodisation of aluminium to resist corrosion;
(d) describe the diaphragm cell in the manufacture of chlorine from brine;
(e) describe the treatment of industrial effluent by electrolysis to remove Ni2+, Cr3+ and Cd2+;
(f) describe the electroplating of coated plastics.
…
