The following are high-level thinking skills (KBAT) questions designed based on the learning outcomes provided:
8.1 Half-cell and redox equations
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.
What is electrochemistry used for?
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?
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:
How do you determine which is oxidized and reduced?
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 – reducing agent (highest tendency to donate electron)- negative. An Ox Negative.
Reduction (gain electrons) – cathode- oxidising agent (highest tendency to accept electron) – 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.
SEK vs SKL
– susunan berdasarkan kereaktifan logam/ hidrogen/ karbon bertindak balas dengan oksigen.
What are the types of electrochemistry?
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 the difference between Daniell cell and galvanic cell?
What is present in a Daniell cell?
How does Daniell cell work?

Daniell cell. 19th-century illustration of an assembled Daniell cell (far left), with the disassembled components next to it extending to far right. The Daniell cell is a type of electrochemical cell (battery). It was invented in 1836 by British chemist John Frederic Daniell. It consisted of a copper pot, a copper sulphate solution, sulphuric acid, and a zinc electrode. The chemical reactions in the assembled battery produced electrical power. This illustration is from ‘Physique Populaire’ (Emile Desbeaux, 1891).
What is voltaic (galvanic) cell notation (cell diagram)?
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.
What is Volta’s battery, column type?
What is a lead acid battery and how does it work?
8.2 Standard electrode potential
What is the electric potential of a cell?
What is the standard cell potential E cell?
How to calculate standard reduction potential?
How is E° cell value calculated?
What is the overall cell potential’s formula?
What is the 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.
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?
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)
Draw and label the example of a standard hydrogen electrode, SHE diagram?
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.
What does a higher reduction potential mean?
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.
What is the driving force in an electrochemical reaction?
What increases cell voltage?
How does pH affect reduction potential?
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?
8.3 Non-standard cell potentials
What is the Nernst equation used for?
What does it mean when Ecell is positive?
What does the negative value of E cell indicate?
What does a positive voltage indicate about the oxidation reduction reaction?
Is E cell always positive or negative Ecell?
Which has the highest reduction potential?
Is reduction potential positive or negative?
8.4 Fuel cells
8.5 Electrolysis
What is an electrolysis? What kind of electricity is used in electrolysis?
Electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell.
Can electrolysis use alternating electric current (AC)?
Electrolysis depends on controlling the voltage and current. Even if it were current-limited, alternating current (AC) would not be appropriate for electrolysis. Because the “cathode” and “anode” are constantly switching places, AC produces explosive mixtures of hydrogen and oxygen.
Why direct current is used in electrolysis?
The direct current helps to deposit the anions in the anode and the cations in the cathode. If alternate current was used, then the direction of current will go on changing and hence this would lead to uneven deposition of ions in the electrodes.
Electrolysis of Water
Electrolytic cell vs electrochemical cell
Predict the products formed during electrolysis
Faraday’s first law of electrolysis
Faraday’s second laws of electrolysis
Relationship between the Faraday constant, the Avogadro constant and the electronic charge
Calculate the quantity of electricity used, the mass of material and/ or gas volume liberated during electrolysis.
8.6 Applications of electrochemistry
What is dry cell?
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.
How can iron be protected from corrosion?
One way to keep iron from corroding is to keep it painted. The layer of paint prevents the water and oxygen necessary for rust formation from coming into contact with the iron. As long as the paint remains intact, the iron is protected from corrosion. Other strategies include alloying the iron with other metals.
Why does Aluminium not undergo corrosion but iron does?
Aluminum unlike iron and steel,does not rust or corrode in moist conditions. Its surface is protected by a natural layer of aluminium oxide. This prevents the metal below from coming into contact with air and oxygen.
* note that metal at upper the SKL (Siri kereaktifan logam) is stronger reducing agent that donate electron to the lower metal. Upper metal will get oxidised not the lower metal when they get contact. When Al-Fe, Fe will not corrode (oxidised) but when Fe-Sn, Fe will corrode rapidly. Logam di atas akan berkorban untuk logam di bawah.
Kedudukan karbon dalam siri kereaktifan logam?
Which metals can prevent the corrosion of iron?
The corrosion-prone iron alloy steel is commonly coated with zinc, a more active metal, in a process known as galvanizing. Corrosion of the sacrificial zinc results in its oxidation; the iron is reduced, which renders it cathodic and inhibits its corrosion.
What is iron coated with zinc?
Galvanization or galvanizing (also spelled galvanisation or galvanising) is the process of applying a protective zinc coating to steel or iron, to prevent rusting. The most common method is hot-dip galvanizing, in which the parts are submerged in a bath of molten hot zinc.
What is aluminum zinc alloy coated steel?
U. S. Steel GALVALUME® Steel Sheet is carbon steel sheet coated with aluminum-zinc alloy by a continuous hot-dip process. The nominal coating composition is 55% aluminum and 45% zinc. A small but important addition of silicon is included in the coating alloy.
Why are iron washers coated with zinc?
Galvanization is the process of applying a protective zinc coating to steel or iron in order to prevent premature rust and corrosion. … The corrosion of zinc is very slow, which gives it an extended life while it protects the base metal. Due to the alloying of the Zinc to the iron, cathodic protection occurs.
Why does zinc not rust?
The zinc layer acts as a sacrificial metal for the steel. This means that the zinc layer will combine with the oxygen more readily than the iron in the steel will. This creates a zinc oxide layer that prevents the formation of iron oxide, thus eliminating the possibility of rust forming.
Does zinc alloy rust easily?
Like all ferrous metals, zinc corrodes when exposed to air and water. However, zinc corrodes at a rate of 1/30 of that for steel. Also like other ferrous metals, zinc corrodes or rusts at different rates depending on its environment. The patina layer is the products of zinc corrosion and rust.
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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), E°;
(c) calculate the standard cell potential using the E° 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 E° 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
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.
…