381. The oxidation state of an element is called:
382. What is the valency of an atom with the sign of its electrovalency called:
383. What is the algebraic sum of the oxidation states of all atoms that make up the molecule:
384. Reactions as a result of which the oxidation states of elements change are called:
385. Oxidizing agent and reducing agent:
386. The amount of oxidizing agent that adds 1 mole of electrons in a given redox reaction is called:
387. Which reaction is redox:
388. What is the oxidation state of chlorine in potassium perchlorate (KClO 4):
389. What is the oxidation state of the chromium atom in the Cr 2 (SO 4) 3 molecule:
390. What is the oxidation state of Mn in the compound KMnO 4:
391. What is the oxidation state of the chromium atom in the K 2 Cr 2 O 7 molecule:
392. Determine the oxidation state of Mn in the compound K 2 MnO 4:
393. Which of the redox reactions is a disproportionation reaction:
394. Which of the redox reactions is intramolecular:
395. The process ClO 3 - ® Cl - is:
396. Specify the final product of the transformation of the MnO ion in an alkaline medium:
397. Specify the final product of the transformation of the MnO ion in an acidic medium:
398. Indicate the final product of the transformation of the MnO ion in a neutral medium:
399. What is the number of electrons involved in the half-reaction of oxidation of the sulfite ion SO to the sulfate ion SO:
400. What is the number of electrons involved in the half-reaction of oxidation of the sulfide ion S2- to the sulfate ion SO:
401. What is the number of electrons involved in the half-reaction of the reduction of the sulfite ion SO to the sulfide ion S 2-:
402. What is the number of electrons involved in the half-reaction of the reduction of the MnO ion to the Mn 2+ ion:
403. What is the number of electrons involved in the half-reaction of oxidation of the S 2- ion to the SO ion:
404. The coefficient in front of the oxidizing formula in the reaction equation between aluminum and bromine is equal to:
405. The coefficient in front of the reducing agent formula in the reaction equation between aluminum and bromine is equal to:
406. Coefficients before the formulas of the reducing agent and oxidizing agent in the reaction equation, the scheme of which is P + KClO 3 = KCl + P 2 O 5:
407. Coefficient in front of the reducing agent formula in the reaction equation, the scheme of which is Mg + HNO 3 = N 2 O + Mg(NO 3) 2 + H 2 O:
408. In the reaction equation, the scheme of which is P + HNO 3 + H 2 O = H 3 PO 4 + NO, the coefficient in front of the reducing agent formula is equal to:
409. What is the equivalent of a reducing agent in a redox reaction: 2H 2 S + H 2 SO 3 = 3S + 3H 2 O:
410. What is the equivalent mass of the reducing agent in the reaction HNO 3 + Ag = NO + AgNO 3 + H 2 O:
411. Specify the equivalent of the oxidizing agent for the reaction HNO 3 + Ag = NO 2 + AgNO 3 + H 2 O:
412. When concentrated nitric acid reacts with sodium metal, the following products are formed:
413. To what substance is concentrated nitric acid reduced when it reacts with silver:
414. With non-metals, dilute nitric acid is reduced to form:
415. Specify the products of the reaction of dilute nitric acid with phosphorus:
416. The products of the interaction of dilute sulfuric acid with copper are:
417. Which metals displace hydrogen in the reaction of their interaction with dilute sulfuric acid:
Electrochemistry
418. What does electrochemistry study:
419. What is the basis of electrochemical phenomena:
420. Components of the simplest electrochemical system:
421. Conductors of the 1st kind in an electrochemical system are:
422. Conductors of the 2nd kind in an electrochemical system can be:
423. The external circuit of the electrochemical system is:
424. Electricity meters (coulometers, current integrators) and other devices are created on the basis of the laws:
425. The wording: “The amount of substance formed on the electrode during electrolysis is directly proportional to the amount of current passing through the electrolyte” is a reflection of:
426. According to Faraday’s law, how much electricity must be consumed to release one gram equivalent of any substance during electrolysis:
427. Oxidation processes in electrochemistry are called:
428. Cathodic processes in electrochemistry are called:
429. Electrodes on which oxidation processes are carried out:
430. Electrodes on which reduction processes are carried out:
431. The total chemical reaction occurring in a galvanic cell is called:
432. How is the interface between a conductor of the first and second kind indicated when schematically writing a galvanic element:
433. How is the interface between conductors of the second kind indicated when schematically writing a galvanic cell:
434. The maximum potential difference of the electrodes that can be obtained during operation of a galvanic cell:
435. The maximum voltage value of a galvanic cell corresponding to a reversible reaction is called:
436. The standard electrode potential (φ°) is called:
437. If we select the processes Ме z + + Zе = Ме from a number of standard electrode potentials, then we obtain the values forming:
438. The Nernst formula, reflecting the dependence of the electrode potential of a metal on various factors, has the following mathematical reflection:
439. Change in electrode potential when current passes:
440. What does electrochemical kinetics study:
441. Single-use device that converts the energy of chemical reactions into electrical energy:
442. The components of the simplest galvanic cell are:
443. A current of 2.5 A passing through an electrolyte solution releases 2.77 g of metal from the solution in 30 minutes. What is the equivalent mass of the metal:
444. A current of 6 A was passed through an aqueous solution of sulfuric acid for 1.5 hours. What is the mass of decomposed water (g):
445. A current of 6 A was passed through an aqueous solution of sulfuric acid for 1.5 hours. What is the volume (l) of hydrogen released (normal conditions):
446. A current of 6 A was passed through an aqueous solution of sulfuric acid for 1.5 hours. What is the volume (l) of the released oxygen (normal conditions):
447. During the operation of which galvanic cell the processes take place Zn -2e = Zn 2+ ; Cu 2+ + 2e = Cu:
448. Specify the diagram of an iron-copper galvanic cell:
449. Diagram of a zinc-magnesium galvanic cell:
450. Specify the circuit of a nickel-copper galvanic cell:
451. The chemical reaction underlying the anodic process when charging an acid battery:
452. The chemical reaction underlying the cathodic process when charging an acid battery:
453. What process during the operation of a lead battery is reflected by the chemical reaction PbO 2 + 2H 2 SO 4 = PbSO 4 + SO 2 + 2H 2 O:
454. What process during the operation of an acid battery is reflected by the chemical reaction Pb + H 2 SO 4 = PbSO 4 + H 2:
455. The chemical reaction underlying the cathodic process when charging an acid battery:
456. The chemical reaction underlying the anodic process when charging an acid battery:
457. In alkaline batteries, a 20% solution serves as an ionic conductor:
458. The general name of a battery in which the current-generating reaction is 2NiOOH + Cd + 2H 2 O → 2Ni(OH) 2 + Cd(OH) 2:
459. The positive electrode in alkaline batteries contains:
460. Negative plates in an alkaline battery, where the current-forming reaction Ni OOH + Fe + 2H 2 O → 2Ni(OH) 2 + Fe(OH) 2 occurs
461. On both electrodes, when an acid battery is discharged, the following is formed:
462. What metal are the positive plates of cadmium-nickel alkaline batteries made of:
463. Negative platinum cadmium-nickel alkaline batteries consist of:
464. The positive plates of a silver-zinc alkaline battery are made from:
465. What metal are the negative platins of a silver-zinc alkaline battery made of:
466. In what cases is a porous partition - a diaphragm - introduced into the electrolyzer:
467. What is the material for making the diaphragm during the operation of the electrolyzer:
468. What process occurs at the cathode during the electrolysis of a solution of potassium sulfate K 2 SO 4:
469. What process occurs on an inert anode during the electrolysis of sodium sulfate Na 2 SO 4:
470. Specify the salt, during the electrolysis of which free oxygen is released at the anode:
471. Ionic equation of the cathodic process 2H 2 O + 2e = H 2 + 2OH - possible during the electrolysis of salt:
472. The ionic equation of the anodic process 2H 2 O - 4e = O 2 + 4H + is possible during the electrolysis of salt:
473. Nickel plates are immersed in aqueous solutions of the salts listed below. What salts will nickel react with?
474. Zinc plates are immersed in aqueous solutions of the salts listed below. Which salt will zinc react with:
475. Indicate the property of iron that negatively affects its use in technology:
476. A cleaned iron nail is dipped into a blue solution of copper (II) chloride, which quickly becomes covered with a coating of copper. The solution acquires a greenish color, which is explained by:
477. The light on the device for testing substances for electrical conductivity will light up when the electrodes are immersed in:
478. How will the glow of a light bulb change in a device for testing the electrical conductivity of solutions if its electrodes are immersed in lime water through which carbon monoxide (IV) is passed? Why?
479. Specify a metal characterized by complete thermodynamic stability to electrochemical corrosion:
480. Until recently, tin cans were made from so-called tinplate (an iron body coated with a protective layer of tin). It is not recommended to store food in open cans, since if the protective layer is scratched, the can will quickly rust. Indicate the reactions underlying this process.
481. Electronic equation of the anodic process of atmospheric corrosion of tinned iron:
482. Electronic equation of the cathodic process of atmospheric corrosion of tinned iron:
Polymers
483. The process of formation of polymers from low molecular weight substances, accompanied by the release of a by-product (water, ammonia, hydrogen chloride, etc.).
A ready-to-use lead acid battery consists of a lattice of lead plates, some of which are filled with lead dioxide and others with metal sponge lead. The plates are immersed in a solution at this concentration; the specific electrical conductivity of the sulfuric acid solution is maximum.
When a battery operates - when it is discharged - an oxidation-reduction reaction occurs in it, during which the metal lead is oxidized
and lead dioxide is reduced:
Electrons given up by metallic lead atoms during oxidation are accepted by lead atoms during reduction; electrons are transferred from one electrode to another through an external circuit.
Thus, metal lead serves as the anode in a lead battery and is negatively charged, and serves as the cathode and is positively charged.
In the internal circuit (in solution) during battery operation, ion transfer occurs. Ions move towards the anode and ions move towards the cathode. The direction of this movement is determined electric field, arising as a result of electrode processes: anions are consumed at the anode, and cations are consumed at the cathode. As a result, the solution remains electrically neutral.
If we add up the equations corresponding to lead oxidation and reduction, we get the total equation for the reaction occurring in a seinium battery during its operation (discharge):
E.m.f. of a charged lead-acid battery is approximately 2 V. As the battery charges, its cathode and anode (Pb) materials are consumed. Sulfuric acid is also consumed. At the same time, the voltage at the battery terminals drops. When it becomes less than value allowed by operating conditions, the battery is charged again.
To charge (or charge), the battery is connected to an external current source (plus to plus and minus to minus). In this case, current flows through the battery in the direction opposite to that in which it passed when the battery was discharged. As a result of this, the electrochemical processes on the electrodes are “reversed”.
The lead electrode now undergoes a reduction process
i.e. this electrode becomes the cathode.
The electrolyte of a lead battery is a solution of sulfuric acid containing a relatively small amount of ions. The concentration of hydrogen ions in this solution is much greater than the concentration of lead ions. In addition, lead comes before hydrogen in the voltage series. However, when charging a battery, it is lead, not hydrogen, that is reduced at the cathode. This occurs because the hydrogen evolution overvoltage on lead is particularly high (see Table 20 on page 295).
Electrochemistry
Zaylobov L. T., graduate student of Tashkent State pedagogical university them. Nizami (Uzbekistan)
DEMONSTRATION OF REDOX REACTION PROCESSES TAKE PLACE IN A LEAD BATTERY USING INNOVATIVE TECHNOLOGIES
An animation model is presented to demonstrate the processes of redox reactions taking place in a lead battery using innovative technologies. This article is recommended for students of academic lyceums and colleges with in-depth study of chemistry.
Key words: redox reactions, galvanic cell, battery, lead accumulator, H2S04 solution, electrode, animation model, lead metal, outcome electric current- discharge, recovery - charge, ions, electrical conductivity.
DEVELOPMENT OF EDUCATION ON OXIDATION-REDUCTION REACTIONS OCCURRING IN LEAD CELLS USING INNOVATIVE TECHNOLOGIES
Is presented animation model of the development of the cost of oxidizing-reconstruction reactions passing in plumbum battery with applying of innovation technologies. This article is recommended for taken into account academic lyceums and colleges with the in-depth studies of chemistry.
Keywords: oxidizing-reconstruction reactions, galvanic element, a batterie, leaden battery, solution H2S04, electrode, animation model, metallic lead, upshot of the electric current - a category, reconstruction - a charge, Ions, conduction.
Currently, widely used galvanic cells - batteries and accumulators are an integral part of our lives. Oxidation and reduction processes that take place in batteries are one of the most difficult topics in general chemistry. Explaining this topic without visual aids and chemical experiments is the main reason for this problem.
The periodic movements of electrons in oxidation and reduction reactions taking place in galvanic cells can only be demonstrated with the help of innovative technologies. A dynamic model of these processes is demonstrated using a computer. Ready-made electronic data and conducting computer lessons based on animation and demonstrating them to students improve the quality of the lesson.
Lead acid battery. The following reactions take place in the elements: At the enode: Pb+SO43^PbSO4+24
At the cathode: Pb O2+ SO42+24^ PbSO4+2H2O The battery has the property of reversibility (can be recharged), since the product of the reactions occurring with it - lead sulfate formed on both electrodes - settles on the plates, and does not diffuse or fall off from them. One cell of the lead-acid battery shown here produces a voltage of about 2 V; in 6 or 12 V batteries, three or six of the described cells are connected in series.
The first functional lead-acid battery was invented in 1859 by French scientist Gaston Plante. The battery design consisted of sheet lead electrodes separated by linen separators, which were rolled into a spiral and placed in a vessel with a 10% sulfuric acid solution. The disadvantage of the first lead-acid batteries was their low capacity.
As an example, consider a ready-to-use lead-acid battery. It consists of lattice lead plates, some of which are filled with lead dioxide and others with metal sponge lead. The plates are immersed in a 35-40% H2804 solution; at this concentration, the specific electrical conductivity of the sulfuric acid solution is maximum.
When the battery is operating - when it is discharged - an oxidation-reduction reaction occurs in it, during which the metal lead is oxidized:
Pb+804-2=Pb804+2е or Pb-2е=Pb+2
And lead dioxide is reduced:
Pb02+2H2804=Pb(804)2+2H20
Pb(804)2+2e= Pb804+ 80^2 or Pb+4+2e=Pb
The electrons given up by metallic lead atoms during oxidation are accepted by lead atoms PbO2 during reduction; electrons are transferred from one electrode to another through an external circuit.
Thus, the chemical processes taking place in batteries were created and tested in the form of an animation model. It shows the outcome of electric current - discharge and recovery - charge. The occurrence of each reaction is explained by the movement of ions in the solution.
р-1.23-1.27 g/ml
In the internal circuit (in the H2804 solution) when the battery is operating, a transfer occurs
ions. 804 ions move to the anode, and H+ ions move to the cathode. The direction of this movement is determined by the electric field resulting from the occurrence of electrode processes: anions are consumed at the anode, and cations are consumed at the cathode. As a result, the solution remains electrically neutral.
If we add up the equations corresponding to the oxidation of lead and the reduction of Pb02, we obtain the total equation for the reaction occurring in a lead battery during its operation (discharge):
Pb + Pb02 + 4H++ 2B04
2Pb04 + 2H2O
E.m.f. of a charged lead-acid battery is approximately 2V. As the battery discharges, the materials of its cathode (PbO2) and anode (Pb) are consumed. Sulfuric acid is also consumed. At the same time, the voltage at the battery terminals drops. When it becomes less than the value allowed by operating conditions, the battery is charged again.
To charge (or charge), the battery is connected to an external current source (plus to plus and minus to minus). In this case, current flows through the battery in the direction opposite to that in which it passed when the battery was discharged. As a result, the electrochemical processes on the electrodes are “reversed.” The reduction process now occurs on the lead electrode:
Pb804 + 2H++2e = H2B04 + Pb i.e. this electrode becomes the cathode. An oxidation process takes place on the Pb02 electrode:
Pb804+2H+-2e=Pb02+H2804+2H+
Therefore, this electrode is now the anode. The ions in the solution move in directions opposite to those in which they moved when the battery was operating.
Adding the last two equations, we obtain the equation for the reaction that occurs when charging the battery:
2Pb04 + 2Sh0^Pb + Pb02 + 2H2B04
It is easy to see that this process is the opposite of the one that occurs during battery operation: when the battery is charged, the substances necessary for its operation are again obtained in it.
Lead-acid batteries are the most common among all currently existing chemical power sources. Their large-scale production is determined both by the relatively low price, due to the comparative scarcity of starting materials, and by the development different options these batteries meeting the requirements wide range consumers.
The use of a visual demonstration of the processes taking place in a given lead-acid battery and the use of an animation model makes it easier for students to understand such a difficult-to-understand topic.
LITERATURE
1. R. Dickerson, G. Gray, J. Haight. Basic laws of chemistry. Publishing house "Mir" Moscow 1982. 653 p.
2. Deordiev S.S. Batteries and their care. K.: Technology, 1985. 136 p.
3. Electrical reference book. In 3 volumes. T.2. Electrical products and devices/under general. ed. professors of Moscow Power Engineering Institute (editor-in-chief I.N. Orlov) and others. 7th ed. 6 rev. and additional M.: Energoatomizdat, 1986. 712 p.
Redox reactions– reactions that occur with a change in the oxidation states of elements.
Oxidation– process of electron donation.
Recovery– the process of adding electrons.
Oxidant– an atom, molecule, or ion that accepts electrons.
Reducing agent– an atom, molecule, or ion that donates electrons.
Oxidizing agents, accepting electrons, go into a reduced form:
F2 [approx. ] + 2ē → 2F¯ [restored].
Reductants, giving up electrons, go into the oxidized form:
Na0 [reduced ] – 1ē → Na+ [ok.].
The equilibrium between the oxidized and reduced forms is characterized by Nernst equations for redox potential:
Where E0– standard value of redox potential; n– number of transferred electrons; [restored ] and [approx. ] are the molar concentrations of the compound in reduced and oxidized forms, respectively.
Values of standard electrode potentials E0 are given in tables and characterize the oxidative and reduction properties of compounds: the more positive the value E0, the stronger the oxidizing properties, and the more negative the value E0, the stronger the restorative properties.
For example, for F2 + 2ē ↔ 2F¯ E0 = 2.87 volts, and for Na+ + 1ē ↔ Na0 E0 =-2.71 volts (the process is always recorded for reduction reactions).
A redox reaction is a combination of two half-reactions, oxidation and reduction, and is characterized by an electromotive force (emf) Δ E0: Δ E0 = Δ E0ok – Δ E0resist, Where E0ok and Δ E0resist– standard potentials of the oxidizing agent and reducing agent for this reaction.
E.m.f. reactions Δ E0 associated with change free energy Gibbs ΔG and reaction equilibrium constant TO:
ΔG = – nF Δ E0 or Δ E = (RT/nF) ln K.
E.m.f. reactions at non-standard concentrations Δ E equal to: Δ E =Δ E0 – (RT/nF) × Ig K or Δ E =Δ E0 –(0,059/n)lg K .
In the case of equilibrium, ΔG = 0 and ΔE = 0, whence Δ E =(0.059/n)lg K And K = 10nΔE/0.059.
For the reaction to proceed spontaneously, the following relations must be satisfied: ΔG< 0 или K >> 1, which corresponds to the condition Δ E0> 0. Therefore, to determine the possibility of a given redox reaction, it is necessary to calculate the value of Δ E0. If Δ E0 > 0, the reaction is in progress. If Δ E0< 0, no response.
Chemical current sources
Galvanic cells– devices that convert energy chemical reaction into electrical energy.
Daniel's galvanic cell consists of zinc and copper electrodes immersed in ZnSO4 and CuSO4 solutions, respectively. Electrolyte solutions communicate through a porous partition. In this case, oxidation occurs on the zinc electrode: Zn → Zn2+ + 2ē, and reduction occurs on the copper electrode: Cu2+ + 2ē → Cu. In general, the reaction goes: Zn + CuSO4 = ZnSO4 + Cu.
Anode– electrode on which oxidation occurs. Cathode– the electrode on which the reduction takes place. In galvanic cells, the anode is negatively charged and the cathode is positively charged. On element diagrams, metal and mortar are separated by a vertical line, and two mortars are separated by a double vertical line.
So, for the reaction Zn + CuSO4 = ZnSO4 + Cu, the circuit diagram of the galvanic cell is written: (-)Zn | ZnSO4 || CuSO4 | Cu(+).
The electromotive force (emf) of the reaction is equal to Δ E0 = E0ok – E0restore = E0(Cu2+/Cu) – E0(Zn2+/Zn) = 0.34 – (-0.76) = 1.10 V. Due to losses, the voltage created by the element will be slightly less than Δ E0. If the concentrations of solutions differ from the standard ones, equal to 1 mol/l, then E0ok And E0resist are calculated using the Nernst equation, and then the emf is calculated. corresponding galvanic cell.
Dry element consists of a zinc body, NH4Cl paste with starch or flour, a mixture of MnO2 with graphite and a graphite electrode. During his work in progress reaction: Zn + 2NH4Cl + 2MnO2 = Cl + 2MnOOH.
Element diagram: (-)Zn | NH4Cl | MnO2, C(+). E.m.f. element - 1.5 V.
is a current source in which the chemical energy of the active substances of spatially separated electrodes as a result of redox reactions is converted into electrical energy. You can buy a high-quality lead-acid battery from SSC. You can be confident in the quality of batteries if you purchase them from a trusted company that has status and many positive reviews among knowledgeable people. In lead batteries, the positive electrodes consist of lead dioxide Pb0 2, the negative electrodes are made of sponge lead. The electrolyte is an aqueous solution of sulfuric acid H 2 SO 4.
The main current-generating process, in accordance with the generally accepted theory of double sulfation in a lead battery, is described by the following reaction:
Рb + Рb0 2 + 2H 2 S0 4 2PbS0 4 + 2H 2 0, (1.1)
Reaction (1.1) is total and is determined by the following processes occurring on the positive and negative electrodes. The process at the negative electrode is expressed as:
Pb + HS0 4 PbSQ 4 + it + 2e, (1.2)
On the positive:
Pb0 2 +HSO 4 - +3H 3 + 2e PbSO 4 + 2H 2 O, (1.3)
Thus, when lead batteries are discharged, practically insoluble lead sulfate is formed on both electrodes due to the reduction of lead dioxide on the positive electrode and oxidation of lead on the negative electrode. When charging, on the contrary, Pb02 is formed on the positive electrode and sponge lead on the negative. A schematic representation of the main processes occurring in a lead-acid battery is presented in Figure 1.1.
As can be seen, during discharge the electrolyte solution is diluted. During long-term discharge modes, the electrolyte density can decrease to a value of 1.02-1.03 g/cm3. This is typical for batteries from any manufacturer and distributor, but only if you can.
Figure 1.1 Schematic representation of the main redox processes occurring in a lead-acid battery
The electromotive force of this electrochemical system is described by the well-known Nernst equation:
where: E is the standard value of e. d.s, a and - activity of a solution of sulfuric acid and water, v = 2.3,
R, T, z, F are known thermodynamic quantities.
The E° value can be easily calculated from thermodynamic data.
E° = 2.041 V.
Thus, the equation of electromotive force in a lead battery is:
shows that e. d.s. depends on the concentration of the sulfuric acid solution.
When charging lead batteries, in addition to current-generating ones, side processes of gas formation occur, caused by the decomposition of water and reducing the utilization rate of the charging current. Hydrogen is released on the negative electrodes, and oxygen is released on the positive electrodes. If the release of hydrogen begins when the battery is almost fully charged, then the release of oxygen begins much earlier. In addition, when using positive current leads made of lead-antimony alloys on negative electrodes, due to the electrical transfer of antimony from positive to negative electrodes, the formation of toxic antimonous hydrogen SbH3 (stibine) occurs.
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