Feroxide catalysts for raspberry powder, igniter composition, caramel fuel.
Method 1. Obtaining iron oxide Fe 2 O 3 from iron sulfate
Iron oxides are very often used as catalysts in pyrotechnic compounds. Previously, they could be purchased in stores. For example, iron oxide monohydrate FeOOH has been encountered as a pigment "iron oxide yellow pigment". Iron oxide Fe 2 O 3 was sold in the form of minium iron. At present, it is not easy to buy all this, as it turned out. I had to take care of getting it at home. I am no chemist, but life forced me. Check out recommendations on the net. Alas, normal, i.e. simple and safe, a recipe for home conditions was not easy to find. Only one recipe seemed to fit, but I couldn't find it again. The list of admissible components in a head was postponed. I decided to go my own way. Oddly enough, the result was very acceptable. The compound turned out with clear signs of iron oxide is very homogeneous and finely dispersed. Its use in raspberry powder and secondary igniter fully confirmed that what was needed was obtained.
So, we buy in a gardening store ferrous sulfate FeSO 4, in the pharmacy we buy pills hydroperita, three packs, and stock up in the kitchen drinking soda NaHCO 3. We have all the ingredients, let's start cooking. Instead of hydroperite tablets, you can use a solution hydrogen peroxide H 2 0 2, also happens in pharmacies.
In a glass dish with a volume of 0.5 liters, we dissolve about 80 g (one third of a pack) of ferrous sulfate in hot water. Add baking soda in small portions while stirring. Some kind of rubbish of a very nasty color is formed, which foams a lot.
FeSO 4 + 2NaHCO 3 \u003d FeCO 3 + Na 2 SO 4 + H 2 O + CO 2
Therefore, everything must be done in the sink. Add baking soda until foaming almost stops. Having slightly settled the mixture, we begin to slowly pour in the crushed tablets of hydroperite. The reaction again proceeds quite vigorously with the formation of foam. The mixture takes on a characteristic color and a familiar rusty smell.
2FeCO 3 + H 2 O 2 \u003d 2FeOOH + 2CO 2
We continue backfilling hydroperite again until the foaming, that is, the reaction, almost completely stops.
We leave our chemical vessel alone and see how a red precipitate falls out - this is our oxide, more precisely FeOOH oxide monohydrate, or hydroxide. It remains to neutralize the connection. We defend the sediment and drain the excess liquid. Then add clean water, defend and drain again. So we repeat 3-4 times. In the end, we dump the sediment on a paper towel and dry it. The resulting powder is an excellent catalyst and can already be used in the manufacture of stopins and secondary igniter composition, "raspberry" gunpowder and for catalyzing caramel rocket fuels. /25.01.2008, kia-soft/
However, the original recipe for "crimson" gunpowder prescribed the use of pure red oxide Fe 2 O 3. As experiments with caramel catalysis have shown, Fe 2 O 3 is indeed a somewhat more active catalyst than FeOOH. To obtain ferric oxide, it is enough to ignite the resulting hydroxide on a hot iron sheet, or simply in a tin can. The result is a red powder Fe 2 O 3 .
After making the muffle furnace, I carry out calcination in it for 1-1.5 hours at a temperature of 300-350°C. Very comfortably. /kia-soft 06.12.2007/
P.S.
Independent studies by the vega rocket scientist have shown that the catalyst obtained by this method has an increased activity compared to industrial feroxides, which is especially noticeable in the sugar caramel fuel obtained by evaporation.
Method 2. Obtaining iron oxide Fe 2 O 3 from ferric chloride
There is information about this possibility on the net, for example, oxide was obtained using bicarbonate on the forum of Bulgarian rocket scientists, this method was mentioned on the forum of chemists, but I did not pay much attention, since I did not have ferric chloride. Recently, a guest of my RubberBigPepper website reminded me of this option. Very timely, as I was actively engaged in electronics and stocked up on chloride. I decided to test this option for obtaining iron hydroxide. The method is financially somewhat more expensive, and the main component of ferric chloride is more difficult to obtain, but in terms of preparation it is easier. So we need ferric chloride FeCl 3 and drinking soda NaHCO 3. Ferric chloride is commonly used for etching printed circuit boards and is sold in radio shops.
Pour two teaspoons of FeCl3 powder into a glass of hot water and stir until dissolved. Now slowly add soda with constant stirring. The reaction proceeds vividly with bubbling and foaming, so there is no need to rush.
FeCl 3 + 3NaHCO 3 \u003d FeOOH + 3NaCl + 3CO 2 + H 2 O
Rash until the bubbling stops. We defend and get the same FeOOH hydroxide in the sediment. Next, we neutralize the compound, as in the first method, by several drains of the solution, topping up with water and settling. Finally, the precipitate is dried and used as a catalyst or to obtain iron oxide Fe 2 O 3 by calcination (see method 1).
Here's an easy way. The yield is very good, from two teaspoons (~15 g) of chloride, 10 g of hydroxide is obtained. Catalysts obtained by this method have been tested and are in good agreement. /kia-soft 11.03.2010/
P.S.
I cannot guarantee the 100% accuracy of the equations of chemical reactions, but in fact they correspond to ongoing chemical processes. Especially dark is the case with Fe(III) hydroxide. According to all the canons, Fe (OH) 3 should precipitate. But in the presence of peroxide (method 1) and at elevated temperature (method 2), in theory, the trihydroxide is dehydrated to FeOOH monohydrate. On the surface, this is exactly what is happening. The resulting hydroxide powder looks like concrete rust, and the main component of rust is FeOOH. ***
Many substances have special properties, which in chemistry are called oxidizing or reducing.
Some chemicals exhibit the properties of oxidizing agents, others - reducing agents, while some compounds can exhibit both properties at the same time (for example, hydrogen peroxide H 2 O 2).
What is an oxidizing agent and a reducing agent, oxidation and reduction?
The redox properties of a substance are associated with the process of giving and receiving electrons by atoms, ions or molecules.
An oxidizing agent is a substance that accepts electrons during a reaction, i.e., is reduced; reducing agent - gives up electrons, i.e., is oxidized. The processes of transferring electrons from one substance to another are usually called redox reactions.
Compounds containing atoms of elements with a maximum degree of oxidation can only be oxidizing agents due to these atoms, because they have already given up all their valence electrons and are only able to accept electrons. The maximum oxidation state of an atom of an element is equal to the number of the group in the periodic table to which the element belongs. Compounds containing atoms of elements with a minimum oxidation state can only serve as reducing agents, since they are only capable of donating electrons, because the external energy level of such atoms is completed by eight electrons
In the process of a redox reaction, the reducing agent gives up electrons, that is, it is oxidized; The oxidizing agent gains electrons, that is, it is reduced.
Redox reactions, or OVR for short, are one of the foundations of the subject of chemistry, as they describe the interaction of individual chemical elements with each other. As the name implies, these reactions involve at least two different chemicals, one of which acts as an oxidizing agent and the other as a reducing agent.
To learn how to correctly determine the role of a particular chemical element in a reaction, you need to clearly understand the following basic concepts. Oxidation is the process of giving off electrons from the outer electron layer of a chemical element.
Typical reducing agents are metals and hydrogen: Fe, K, Ca, Cu, Mg, Na, Zn, H). The less they are ionized, the greater their reducing properties. For example, partially oxidized iron that has donated one electron and has a charge of +1 will be able to donate one less electron compared to "pure" iron. Let's define an oxidizing agent and a reducing agent using the example of a simple reaction of the interaction of the interaction of sodium with oxygen.
Therefore, sodium is a reducing agent and oxygen is an oxidizing agent. To do this, you need to know what the degree of oxidation is. Learn to determine the oxidation state of any atom in a chemical compound.
The former are reducing agents, the latter are oxidizing agents. In addition, you can see what degree of oxidation the elements are in (suddenly somewhere it is minimal or vice versa maximum). Chemical reactions can be divided into two types. The first type includes ion-exchange reactions. In them, the oxidation state of the elements that make up the interacting substances remains unchanged.
REDOX REACTIONSTerms, definitions, concepts
This group of reactions is called redox. In cases of interaction of typical oxidizing agents and reducing agents, you can immediately determine that we are talking about a redox reaction. For example, this is the interaction of alkali metals with acids or halogens, combustion processes in oxygen. Similarly, you determine that the oxidation state of sulfur in potassium sulfide is (+4). Three oxygen atoms take 6 electrons, and two potassium atoms donate two electrons.
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And you can conclude that this reaction is redox. Reactions that occur with a change in the oxidation states of the atoms that make up the reactants are called redox reactions. The change in oxidation states occurs due to the transfer of electrons from the reducing agent to the oxidizing agent. The oxidation state is the formal charge of the atom, assuming that all bonds in the compound are ionic.
When compiling an equation for a redox reaction, it is necessary to determine the reducing agent, oxidizing agent and the number of given and received electrons
If an element is an oxidizing agent, its oxidation state decreases. The process of receiving electrons by substances is called reduction. The oxidizing agent is reduced during the process. The reducing agent has an increased oxidation state.
The reducing agent is oxidized during the process. Using this reaction as an example, let's consider how to draw up an electronic balance. However, no coefficient was placed before the hydrochloric acid formula, since not all chloride ions participated in the redox process. The electron balance method allows you to equalize only the ions involved in the redox process.
Namely, potassium cations, hydrogen and chloride anions. A “copper” coin was placed in a glass with 10 ml of acid. The entire space above the liquid turned brown, brown vapors poured out of the glass. The solution turned green. The reaction was constantly accelerating. After about half a minute, the solution turned blue, and after two minutes the reaction began to slow down.
The green color of the solution in the initial stage of the reaction is due to the products of the reduction of nitric acid. 4. Equalize the number of given and received electrons. When redox reactions occur, the final products depend on many factors.
In a neutral medium, MnO2 is formed and the color changes from red-violet to brown. This includes the production of metals, combustion, the synthesis of sulfur and nitrogen oxides in the production of acids, and the production of ammonia. Hello! I'm wondering if you have any problems doing your homework. We have a lot of people to help you here. Also, my last question was solved in less than 10 minutes :D Anyway, you can just log in and try adding your question.
In turn, the oxidizing agent will be an atom, molecule or ion that accepts electrons and thereby lowers the degree of its oxidation, which is restored. During the lesson, the topic "Oxidation-reduction reactions" was studied.
Chapter 10
Redox reactions.
Redox reactions – these are reactions that occur with a change in the oxidation states of the atoms of the elements that make up the molecules of the reacting substances:
2Mg + O 2 2MgO,
2KClO 3 2KCl + 3O 2 .
Recall that oxidation state – this is the conditional charge of an atom in a molecule, arising from the assumption that the electrons are not displaced, but are completely given to an atom of a more electronegative element.
The most electronegative elements in the compound have negative oxidation states, and the atoms of elements with less electronegativity are positive.
The oxidation state is a formal concept; in some cases, the value of the oxidation state of an element does not coincide with its valence.
To find the oxidation state of the atoms of the elements that make up the reactants, the following rules should be borne in mind:
1. The oxidation state of the atoms of elements in the molecules of simple substances is zero.
For example:
Mg0, Cu0.
2. The oxidation state of hydrogen atoms in compounds is usually +1.
For example: +1 +1
Exceptions: in hydrides (compounds of hydrogen with metals), the degree of oxidation of hydrogen atoms is –1.
For example:
NaH -1 .
3. The oxidation state of oxygen atoms in compounds is usually -2.
For example:
H 2 O -2, CaO -2.
Exceptions:
The oxidation state of oxygen in oxygen fluoride (OF 2) is +2.
the degree of oxidation of oxygen in peroxides (H 2 O 2 , Na 2 O 2) containing the –O–O– group is –1.
4. The oxidation state of metals in compounds is usually a positive value.
For example: +2
5. The oxidation state of non-metals can be both negative and positive.
For example: –1 +1
6. Amount c the oxidation states of all atoms in the molecule is zero.
Redox reactions are two interrelated processes - the oxidation process and the reduction process.
Oxidation process – it is the process of donating electrons by an atom, molecule or ion; in this case, the oxidation state increases, and the substance is a reducing agent:
– 2ē 2H + oxidation process,
Fe +2 – ē Fe +3 oxidation process,
2J – – 2ē oxidation process.
The reduction process is the process of adding electrons, while the oxidation state decreases, and the substance is an oxidizing agent:
+ 4ē 2O –2 reduction process,
Mn +7 + 5ē Mn +2 reduction process,
Cu +2 +2ē Cu 0 reduction process.
Oxidizer – a substance that accepts electrons and is reduced in the process (the element's oxidation state is reduced).
Reducing agent – a substance that donates electrons and is oxidized at the same time (the oxidation state of an element decreases).
It is possible to make a reasonable conclusion about the nature of the behavior of a substance in specific redox reactions based on the value of the redox potential, which is calculated from the value of the standard redox potential. However, in some cases, it is possible, without resorting to calculations, but knowing the general laws, to determine which substance will be an oxidizing agent and which one will be a reducing agent, and make a conclusion about the nature of the redox reaction.
Typical reducing agents are:
some simple substances:
metals: e.g. Na, Mg, Zn, Al, Fe,
non-metals: for example, H 2 , C, S;
some complex substances: for example, hydrogen sulfide (H 2 S) and sulfides (Na 2 S), sulfites (Na 2 SO 3), carbon monoxide (II) (CO), hydrogen halides (HJ, HBr, HCI) and salts of hydrohalic acids (KI, NaBr), ammonia (NH 3);
metal cations in lower oxidation states: for example, SnCl 2 , FeCl 2 , MnSO 4 , Cr 2 (SO 4) 3 ;
cathode during electrolysis.
Typical oxidizing agents are:
some simple substances - non-metals: for example, halogens (F 2, CI 2, Br 2, I 2), chalcogens (O 2, O 3, S);
some complex substances: for example, nitric acid (HNO 3), sulfuric acid (H 2 SO 4 conc.), potassium premanganate (K 2 MnO 4), potassium dichromate (K 2 Cr 2 O 7), potassium chromate (K 2 CrO 4), manganese (IV) oxide (MnO 2), lead (IV) oxide (PbO 2), potassium chlorate (KCIO 3), hydrogen peroxide (H 2 O 2);
Anode during electrolysis.
When compiling the equations of redox reactions, it should be borne in mind that the number of electrons donated by the reducing agent is equal to the number of electrons accepted by the oxidizing agent.
There are two methods for compiling equations of redox reactions - electron balance method and electron-ion method (half-reaction method) .
When compiling the equations of redox reactions by the electron balance method, a certain procedure should be followed. Consider the procedure for compiling equations by this method using the example of the reaction between potassium permanganate and sodium sulfite in an acidic medium.
We write down the reaction scheme (indicate the reagents and reaction products):
We determine the oxidation state of the atoms of elements that change its value:
7 + 4 + 2 + 6
KMnO 4 + Na 2 SO 3 + H 2 SO 4 → MnSO 4 + Na 2 SO 4 + K 2 SO 4 + H 2 O.
3) We draw up an electronic balance diagram. To do this, we write down the chemical signs of the elements whose atoms change their oxidation state, and determine how many electrons give or add the corresponding atoms or ions.
We indicate the processes of oxidation and reduction, the oxidizing agent and the reducing agent.
We equalize the number of given and received electrons and, thus, determine the coefficients for the reducing agent and oxidizing agent (in this case, they are respectively equal to 5 and 2):
5 S +4 - 2 e- → S +6 oxidation process, reducing agent
2 Mn +7 + 5 e- → Mn +2 reduction process, oxidizing agent.
2KMnO 4 + 5Na 2 SO 3 + 8H 2 SO 4 \u003d 2MnSO 4 + 5Na 2 SO 4 + K 2 SO 4 + 8H 2 O.
5) If hydrogen and oxygen do not change their oxidation states, then their number is counted last and the required number of water molecules is added to the left or right side of the equation.
Redox reactions are divided into three types: intermolecular, intramolecular and self-oxidation reactions - self-healing (disproportionation).
Reactions of intermolecular oxidation - reduction redox reactions are called, in which the oxidizing agent and reducing agent are represented by molecules of different substances.
For example:
2Al + Fe 2 O 3 \u003d 2Fe + Al 2 O 3,
Al 0 - 3e - → Al +3 oxidation, reducing agent,
Fe +3 +3e – → Fe 0 reduction, oxidizing agent.
In this reaction, the reducing agent (Al) and the oxidizing agent (Fe +3) are part of different molecules.
Reactions of intramolecular oxidation – recovery reactions are called in which the oxidizing agent and the reducing agent are part of the same molecule (and are represented either by different elements or by one element, but with different oxidation states):
2 KClO 3 \u003d KCl + 3O 2
2 CI +5 + 6e – → CI –1 reduction, oxidizer
3 2O –2 – 4e – → oxidation, reducing agent
In this reaction, the reducing agent (O -2) and the oxidizing agent (CI +5) are part of the same molecule and are represented by different elements.
In the reaction of thermal decomposition of ammonium nitrite, the atoms of the same chemical element, nitrogen, which are part of one molecule, change their oxidation states:
NH 4 NO 2 \u003d N 2 + 2H 2 O
N –3 – 3e – → N 0 reduction, oxidizer
N +3 + 3e - → N 0 oxidation, reducing agent.
Reactions of this type are often called reactions. counterproportionation .
Auto-oxidation reactions– self-healing(disproportionation) - These are reactions in the course of which the same element with the same oxidation state itself both increases and lowers its oxidation state.
For example: 0 -1 +1
Cl 2 + H 2 O \u003d HCI + HCIO
CI 0 + 1e – → CI –1 reduction, oxidizer
CI 0 - 1e - → CI +1 oxidation, reducing agent.
Disproportionation reactions are possible when the element in the original substance has an intermediate oxidation state.
The properties of simple substances can be predicted by the position of the atoms of their elements in the periodic system of elements D.I. Mendeleev. So, all metals in redox reactions will be reducing agents. Metal cations can also be oxidizing agents. Non-metals in the form of simple substances can be both oxidizing and reducing agents (excluding fluorine and inert gases).
The oxidizing power of non-metals increases in a period from left to right, and in a group from bottom to top.
Restorative abilities, on the contrary, decrease from left to right and from bottom to top for both metals and non-metals.
If the redox reaction of metals occurs in solution, then to determine the reducing ability, use a range of standard electrode potentials (activity series of metals). In this series, metals are arranged as the reducing ability of their atoms decreases and the oxidizing ability of their cations increases ( see table. 9 apps ).
The most active metals, standing in a series of standard electrode potentials up to magnesium, can react with water, displacing hydrogen from it.
For example:
Ca + 2H 2 O \u003d Ca (OH) 2 + H 2
When interacting metals with salt solutions, it should be borne in mind that each more active metal (not interacting with water) is able to displace (restore) the metal behind it from a solution of its salt.
So, iron atoms can restore copper cations from a solution of copper sulfate (CuSO 4):
Fe + CuSO 4 \u003d Cu + FeSO 4
Fe 0 - 2e - \u003d Fe +2 oxidation, reducing agent
Cu +2 + 2e - = Cu 0 reduction, oxidizing agent.
In this reaction, iron (Fe) is located in the activity series before copper (Cu) and is a more active reducing agent.
The reaction, for example, of silver with a solution of zinc chloride will be impossible, since silver is located in the series of standard electrode potentials to the right of zinc and is a less active reducing agent.
All metals that are in the activity series up to hydrogen can displace hydrogen from solutions of ordinary acids, that is, restore it:
Zn + 2HCl \u003d ZnCI 2 + H 2
Zn 0 - 2e - \u003d Zn +2 oxidation, reducing agent
2H + + 2e – → reduction, oxidizing agent.
Metals that are in the activity series after hydrogen will not reduce hydrogen from solutions of ordinary acids.
To determine if there might be oxidizing agent or reducing agent complex substance, it is necessary to find the degree of oxidation of the elements that make it up. The elements that are in highest oxidation state , can only lower it by accepting electrons. Consequently, substances whose molecules contain atoms of elements in the highest oxidation state will only be oxidizing agents .
For example, HNO 3 , KMnO 4 , H 2 SO 4 in redox reactions will only function as an oxidizing agent. The oxidation states of nitrogen (N +5), manganese (Mn +7) and sulfur (S +6) in these compounds have maximum values (coincide with the group number of this element).
If the elements in the compounds have the lowest oxidation state, then they can only increase it by donating electrons. At the same time, such substances containing elements in the lowest oxidation state will only function as a reducing agent .
For example, ammonia, hydrogen sulfide and hydrogen chloride (NH 3, H 2 S, HCI) will only be reducing agents, since the oxidation states of nitrogen (N -3), sulfur (S -2) and chlorine (Cl -1) are the lowest for these elements .
Substances that contain elements with intermediate oxidation states can be both oxidizing and reducing agents., depending on the specific reaction. Thus, they may exhibit redox duality.
Such substances include, for example, hydrogen peroxide (H 2 O 2), an aqueous solution of sulfur oxide (IV) (sulphurous acid), sulfites, etc. Similar substances, depending on environmental conditions and the presence of stronger oxidizing agents (reducing agents), may exhibit in some cases, oxidizing properties, and in others - reducing.
As you know, many elements have a variable degree of oxidation, being part of various compounds. For example, sulfur in the compounds H 2 S, H 2 SO 3, H 2 SO 4 and sulfur S in the free state has oxidation states -2, +4, +6 and 0, respectively. Sulfur refers to the elements R-electron family, its valence electrons are located on the last s- and R-sublevels (...3 s 3R). The sulfur atom with the oxidation state - 2 valence sublevels is fully equipped. Therefore, a sulfur atom with a minimum oxidation state (–2) can only donate electrons (oxidize) and be only a reducing agent. A sulfur atom with an oxidation state of +6 has lost all of its valence electrons and in this state can only accept electrons (recover). Therefore, the sulfur atom with the maximum oxidation state (+6) can only be an oxidizing agent.
Sulfur atoms with intermediate oxidation states (0, +4) can both lose and gain electrons, that is, they can be both reducing agents and oxidizing agents.
Similar reasoning is valid when considering the redox properties of atoms of other elements.
The nature of the course of the redox reaction is affected by the concentration of substances, the environment of the solution and the strength of the oxidizing agent and reducing agent. Thus, concentrated and dilute nitric acid react differently with active and inactive metals. The depth of nitrogen reduction (N+5) of nitric acid (oxidant) will be determined by the activity of the metal (reductant) and the concentration (dilution) of the acid.
4HNO 3 (conc.) + Cu \u003d Cu (NO 3) 2 + 2NO 2 + 2H 2 O,
8HNO 3 (razb.) + 3Cu \u003d 3Cu (NO 3) 2 + 2NO + 4H 2 O,
10HNO 3 (conc.) + 4Mg \u003d 4Mg (NO 3) 2 + N 2 O + 5H 2 O,
10HNO 3 (c. razb.) + 4Mg \u003d 4Mg (NO 3) 2 + NH 4 NO 3 + 3H 2 O.
The reaction of the medium has a significant influence on the course of redox processes.
If potassium permanganate (KMnO 4) is used as an oxidizing agent, then, depending on the reaction of the solution medium, Mn +7 will be reduced in different ways:
in an acidic environment (up to Mn +2) the reduction product will be a salt, for example, MnSO 4,
in a neutral environment (up to Mn +4) the reduction product will be MnO 2 or MnO (OH) 2,
in an alkaline environment (up to Mn +6) the reduction product will be a manganate, for example, K 2 MnO 4 .
For example, when reducing a solution of potassium permanganate with sodium sulfite, depending on the reaction of the medium, the corresponding products will be obtained:
sourWednesday –
2KMnO 4 + 5Na 2 SO 3 + 3H 2 SO 4 = 5Na 2 SO 4 + 2MnSO 4 + K 2 SO 4 + H 2 O
neutralWednesday –
2KMnO 4 + 3Na 2 SO 3 + H 2 O \u003d 3Na 2 SO 4 + 2MnO 2 + 2KOH
alkalineWednesday –
2KMnO 4 + Na 2 SO 3 + 2NaOH \u003d Na 2 SO 4 + Na 2 MnO 4 + K 2 MnO 4 + H 2 O.
The temperature of the system also affects the course of the redox reaction. So, the products of the interaction of chlorine with an alkali solution will be different depending on the temperature conditions.
When chlorine reacts with cold alkali solution The reaction proceeds with the formation of chloride and hypochlorite:
Cl 2 + KOH → KCI + KCIO + H 2 O
CI 0 + 1e – → CI –1 reduction, oxidizer
CI 0 - 1e - → CI +1 oxidation, reducing agent.
If you take hot concentrated KOH solution, then as a result of interaction with chlorine we get chloride and chlorate:
0 t° -1 +5
3CI 2 + 6KOH → 5KCI + KCIO 3 + 3H 2 O
5 │ CI 0 + 1e – → CI –1 reduction, oxidizer
1 │ CI 0 - 5e - → CI +5 oxidation, reducing agent.
10.1. Questions for self-control on the topic
1. What reactions are called redox reactions?
2. What is the oxidation state of an atom? How is it defined?
3. What is the degree of oxidation of atoms in simple substances?
4. What is the sum of the oxidation states of all atoms in a molecule?
5. What process is called the oxidation process?
6. What substances are called oxidizing agents?
7. How does the oxidation state of an oxidizing agent change in redox reactions?
8. Give examples of substances that are only oxidizing agents in redox reactions.
9. What process is called the recovery process?
10. Define the term "reductant".
11. How does the oxidation state of the reducing agent change in redox reactions?
12. What substances can only be reducing agents?
13. What element is an oxidizing agent in the reaction of dilute sulfuric acid with metals?
14. What element is an oxidizing agent in the interaction of concentrated sulfuric acid with metals?
15. What is the function of nitric acid in redox reactions?
16. What compounds can be formed as a result of the reduction of nitric acid in reactions with metals?
17. What element is an oxidizing agent in concentrated, dilute and very dilute nitric acid?
18. What role can hydrogen peroxide play in redox reactions?
19. How are all redox reactions classified?
10.2. Tests for self-control of knowledge of the theory on the topic "Oxidation-reduction reactions"
Option number 1
1) CuSO 4 + Zn = ZnSO 4 + Cu,
2) CaCO 3 + CO 2 + H 2 O \u003d Ca (HCO 3) 2,
3) SO 3 + H 2 O \u003d H 2 SO 4,
4) FeCl 3 + 3NaOH \u003d Fe (OH) 3 + 3NaCl,
5) NaHCO 3 + NaOH = Na 2 CO 3 + H 2 O.
2. Based on the structure of atoms, determine under what number the formula of the ion is indicated, which can only be an oxidizing agent:
1) Mn 
, 2) NO 3– , 3) Br – , 4) S 2– , 5) NO 2– ?
3. Under what number is the formula of the substance that is the most powerful reducing agent, from among the following:
1) NO 3–, 2) Сu, 3) Fe, 4) Ca, 5) S?
4. What number indicates the amount of substance KMnO 4, in moles, which interacts with 10 mol of Na 2 SO 3 in the reaction represented by the following scheme:
KMnO 4 + Na 2 SO 3 + H 2 SO 4 → MnSO 4 + Na 2 SO 4 + K 2 SO 4 + H 2 O?
1) 4, 2) 2, 3) 5, 4) 3, 5) 1.
5. What number is the disproportionation reaction (self-oxidation - self-recovery)?
1) 2H 2 S + H 2 SO 3 \u003d 3S + 3H 2 O,
2) 4KClO 3 \u003d KCl + 3KClO 4,
3) 2F 2 + 2H 2 O \u003d 4HF + O 2.
4) 2Au 2 O 3 \u003d 4Au + 3O 2,
5) 2KClO 3 \u003d 2KCl + 3O 2.
Option number 2
1. Under what number is the equation of the redox reaction given?
1) 4KClO 3 \u003d KCl + 3KClO 4,
2) CaCO 3 \u003d CaO + CO 2,
3) CO 2 + Na 2 O \u003d Na 2 CO 3,
4) CuOHCl + HCl \u003d CuCl 2 + H 2 O,
5) Pb (NO 3) 2 + Na 2 SO 4 = PbSO 4 + 2NaNO 3.
2. Under what number is the formula of a substance that can only be a reducing agent:
1) SO 2, 2) NaClO, 3) KI, 4) NaNO 2, 5) Na 2 SO 3?
3. Under what number is the formula of the substance, which is the most powerful oxidizing agent, among the given:
1) I 2 , 2) S, 3) F 2 , 4) O 2 , 5) Br 2 ?
4. Under what number is the volume of hydrogen in liters under normal conditions, which can be obtained from 9 g of Al as a result of the following redox reaction:
2Al + 6H 2 O \u003d 2Al (OH) 3 + 3H 2
1) 67,2, 2) 44,8, 3) 33,6, 4) 22,4, 5) 11,2?
5. What number is the scheme of the redox reaction that takes place at pH > 7?
1) I 2 + H 2 O → HI + HIO,
2) FeSO 4 + HIO 3 + ... → I 2 + Fe(SO 4) 3 + ...,
3) KMnO 4 + NaNO 2 + ... → MnSO 4 + ...,
4) KMnO 4 + NaNO 2 + ... → K 2 MnO 4 + ...,
5) CrCl 3 + KMnO 4 + ... → K 2 Cr 2 O 7 + MnO (OH) 2 + ....
Option number 3
1. Under what number is the equation of the redox reaction given?
1) H 2 SO 4 + Mg → MgSO 4 + H 2,
2) CuSO 4 + 2NaOH →Cu(OH) 2 + Na 2 SO 4,
3) SO 3 + K 2 O → K 2 SO 4,
4) CO 2 + H 2 O → H 2 CO 3,
5) H 2 SO 4 + 2KOH → K 2 SO 4 + 2H 2 O.
2. Based on the structure of the atom, determine the number under which the formula of the ion is given, which can be a reducing agent:
1) Ag + , 2) A l3+ , 3) C l7+ , 4) Sn 2+ , 5) Zn 2+ ?
3. What is the recovery process number?
1) NO 2– → NO 3–, 2) S 2– → S 0, 3) Mn 2+ → MnO 2,
4) 2I – → I 2 , 5) 
→ 2Cl - .
4. Under what number is the mass of the reacted iron given, if as a result of the reaction represented by the following scheme:
Fe + HNO 3 → Fe(NO 3) 3 + NO + H 2 O
formed 11.2 L NO(n.o.)?
1) 2,8, 2) 7, 3) 14, 4) 56, 5) 28.
5. Under what number is the scheme of the reaction of self-oxidation-self-recovery (dismutation)?
1) HI + H 2 SO 4 → I 2 + H 2 S + H 2 O,
2) FeCl 2 + SnCl 4 → FeCl 3 + SnCl 2,
3) HNO 2 → NO + NO 2 + H 2 O,
4) KClO 3 → KCl + O 2,
5) Hg(NO 3) 2 → HgO + NO 2 + O 2.
See the answers to the test questions on p.
10.3. Questions and exercises for self-study
research work on the topic.
1. Indicate the number or sum of the conditional numbers under which the schemes of redox reactions are located:
1) MgCO 3 + HCl MgCl 2 + CO 2 + H 2 O,
2) FeO + P Fe + P 2 O 5,
4) H 2 O 2 H3O + O 2, 8) KOH + CO 2 KHCO 3.
2. Indicate the number or sum of conditional numbers under which redox processes are located:
1) electrolysis of sodium chloride solution,
2) pyrite firing,
3) hydrolysis of sodium carbonate solution,
4) lime slaking.
3. Indicate the number or sum of conditional numbers under which the names of groups of substances are located, characterized by an increase in oxidizing properties:
1) chlorine, bromine, fluorine,
2) carbon, nitrogen, oxygen,
3) hydrogen, sulfur, oxygen,
4) bromine, fluorine, chlorine.
4. Which of the substances - chlorine, sulfur, aluminum, oxygen– is a stronger reducing agent? In your answer, indicate the value of the molar mass of the selected compound.
5. Indicate the number or sum of conditional numbers under which only oxidizing agents are located:
1) K 2 MnO 4, 2) KMnO 4, 4) MnO 3, 8) MnO 2,
16) K 2 Cr 2 O 7, 32) K 2 SO 3.
6. Indicate the number or sum of conditional numbers under which the formulas of substances with redox duality are located:
1) KI, 2) H 2 O 2, 4) Al, 8) SO 2, 16) K 2 Cr 2 O 7, 32) H 2.
7. Which of the compounds - iron oxide(III) chromium oxide(III) sulfur oxide(IV) nitrogen oxide(II) nitrogen oxide(V) - can only be an oxidizing agent? In your answer, indicate the value of the molar mass of the selected compound.
8. Indicate the number or sum of conditional numbers, under which are the formulas of substances that have an oxygen oxidation state - 2:
1) H 2 O, Na 2 O, Cl 2 O, 2) HPO 3, Fe 2 O 3, SO 3,
4) OF 2 , Ba(OH) 2 , Al 2 O 3 , 8) BaO 2 , Fe 3 O 4 , SiO 2 .
9. Which of the following compounds can only be an oxidizing agent: sodium nitrite, sulfurous acid, hydrogen sulfide, nitric acid? In your answer, indicate the value of the molar mass of the selected compound.
10. Which of the following nitrogen compounds is NH 3; HNO3; HNO2; NO 2 - can only be an oxidizing agent? In your answer, write down the value of the relative molecular weight of the selected compound.
11. Under what number, among the names of substances listed below, is the strongest oxidizing agent indicated?
1) concentrated nitric acid,
2) oxygen,
3) electric current at the anode during electrolysis,
12. Which of the following nitrogen compounds is HNO 3; NH3; HNO2; NO - can only be a reducing agent? In your answer, write down the molar mass of the selected compound.
13. Which of the compounds is Na 2 S; K2Cr2O7; KMnO 4 ; NaNO 2 ; KClO 4 - can it be both an oxidizing agent and a reducing agent, depending on the reaction conditions? In your answer, write down the molar mass of the selected compound.
14. Indicate the number or sum of conditional numbers, where ions are indicated that can be reducing agents:
1) (MnO 4) 2–, 2) (CrO 4) –2, 4) Fe +2, 8) Sn +4, 16) (ClO 4) –.
15. Indicate the number or sum of conditional numbers, under which only oxidizing agents are located:
1) K 2 MnO 4, 2) HNO 3, 4) MnO 3, 8) MnO 2, 16) K 2 CrO 4, 32) H 2 O 2.
16. Indicate the number or sum of conditional numbers, under which are located only the names of substances between which redox reactions are not possible:
1) carbon and sulfuric acid,
2) sulfuric acid and sodium sulfate,
4) hydrogen sulfide and hydrogen iodide,
8) sulfur oxide (IV) and hydrogen sulfide.
17. Indicate the number or sum of conditional numbers under which the oxidation processes are located:
1) S +6 S -2, 2) Mn +2 Mn +7, 4) S -2 S +4,
8) Mn +6 Mn +4, 16) O 2 2O -2, 32) S +4 S +6.
18. Indicate the number or sum of conditional numbers under which the recovery processes are located:
1) 2I -1 I 2, 2) 2N +3 N 2, 4) S -2 S +4,
8) Mn +6 Mn +2, 16) Fe +3 Fe 0, 32) S 0 S +6.
19. Specify the number or sum of conditional numbers under which the recovery processes are located:
1) C 0 CO 2, 2) Fe +2 Fe +3,
4) (SO 3) 2– (SO 4) 2–, 8) MnO 2 Mn +2.
20. Indicate the number or sum of conditional numbers under which the recovery processes are located:
1) Mn +2 MnO 2, 2) (IO 3) - (IO 4) -,
4) (NO 2) - (NO 3) -, 8) MnO 2 Mn +2.
21. Indicate the number or the sum of conditional numbers under which the ions that are reducing agents are located.
1) Ca +2, 2) Al +3, 4) K +, 8) S –2, 16) Zn +2, 32) (SO 3) 2–.
22. Under what number is the formula of a substance, in the interaction with which hydrogen acts as an oxidizing agent?
1) O 2, 2) Na, 3) S, 4) FeO.
23. Under what number is the reaction equation in which the reducing properties of the chloride ion appear?
1) MnO 2 + 4HCl = MnCl 2 + Cl 2 + 2H 2 O,
2) CuO + 2HCl = CuCl 2 + H 2 O,
3) Zn + 2HCl \u003d ZnCl 2 + H 2,
4) AgNO 3 + HCl \u003d AgCl + HNO 3.
24. When interacting with which of the following substances - O 2, NaOH, H 2 S - sulfur oxide (IV) exhibits the properties of an oxidizing agent? Write the equation of the corresponding reaction and in the answer indicate the sum of the coefficients of the starting substances.
25. Indicate the number or sum of conditional numbers under which the disproportionation reaction schemes are located:
1) NH 4 NO 3 N 2 O + H 2 O, 2) NH 4 NO 2 N 2 + H 2 O,
4) KClO 3 KClO 4 + KCl, 8) KClO 3 KCl + O 2.
26. Draw an electronic balance diagram and indicate how much potassium permanganate is involved in the reaction with ten moles of sulfur oxide (IV). The reaction proceeds according to the scheme:
KMnO 4 + SO 2 MnSO 4 + K 2 SO 4 + SO 3.
27. Draw an electronic balance diagram and indicate how much potassium sulfide substance interacts with six moles of potassium permanganate in the reaction:
K 2 S + KMnO 4 + H 2 O MnO 2 + S + KOH.
28. Draw an electronic balance diagram and indicate how much potassium permanganate substance interacts with ten moles of iron (II) sulfate in the reaction:
KMnO 4 + FeSO 4 + H 2 SO 4 MnSO 4 + Fe 2 (SO 4) 3 + K 2 SO 4 + H 2 O.
29. Draw an electronic balance diagram and indicate how much potassium chromite (KCrO 2) reacts with six moles of bromine in the reaction:
KCrO 2 + Br 2 + KOH K 2 CrO 4 + KBr + H 2 O.
30. Draw an electronic balance diagram and indicate how much of the manganese (IV) oxide substance interacts with six moles of lead (IV) oxide in the reaction:
MnO 2 + PbO 2 + HNO 3 HMnO 4 + Pb (NO 3) 2 + H 2 O.
31. Write the reaction equation:
KMnO 4 + NaI + H 2 SO4 I 2 + K 2 SO 4 + MnSO 4 + Na 2 SO 4 + H 2 O.
32. Write the reaction equation:
KMnO 4 + NaNO 2 + H 2 O MnO 2 + NaNO 3 + KOH.
In your answer, indicate the sum of the stoichiometric coefficients in the reaction equation.
33. Write the reaction equation:
K 2 Cr 2 O 7 + HCl conc. KCl + CrCl 3 + Cl 2 + H 2 O.
In your answer, indicate the sum of the stoichiometric coefficients in the reaction equation.
34. Draw an electronic balance diagram and indicate how much sodium nitrite (NaNO 2) substance interacts with four moles of potassium permanganate in the reaction:
KMnO 4 + NaNO 2 + H 2 SO 4 MnSO 4 + NaNO 3 + K 2 SO 4 + H 2 O.
35. Draw an electronic balance diagram and indicate how much hydrogen sulfide substance interacts with six moles of potassium permanganate in the reaction:
KMnO 4 + H 2 S + H 2 SO 4 S + MnSO 4 + K 2 SO 4 + H 2 O.
36. What amount of iron substance in moles will be oxidized by oxygen with a volume of 33.6 liters (n.o.) in the reaction proceeding according to the scheme below?
Fe + H 2 O + O 2 Fe (OH) 3.
37. Which of the following metals - Zn, Rb, Ag, Fe, Mg - does not dissolve in dilute sulfuric acid? In your answer, indicate the value of the relative atomic mass of this metal.
38. Which of the following metals - Zn, Rb, Ag, Fe, Mg - does not dissolve in concentrated sulfuric acid? In your answer, indicate the ordinal number of the element in the periodic system of D.I. Mendeleev.
39. Indicate the number or sum of conditional numbers under which the metals are passivated in concentrated solutions of oxidizing acids.
1) Zn, 2) Cu, 4) Au, 8) Fe, 16) Mg, 32) Cr.
40. Indicate the number or sum of conditional numbers under which are the chemical signs of metals that do not displace hydrogen from a dilute solution of sulfuric acid, but displace mercury from solutions of Hg 2+ salts:
1) Fe, 2) Zn, 4) Au, 8) Ag, 16) Cu.
41. Under what number are the chemical signs of metals, each of which does not react with nitric acid?
1) Zn, Ag; 2) Pt, Au; 3) Cu, Zn; 4) Ag, Hg.
42. Under what number is the method of obtaining chlorine in industry indicated?
1) electrolysis of sodium chloride solution;
2) the action of manganese oxide (1V) on hydrochloric acid;
3) thermal decomposition of natural chlorine compounds;
4) the action of fluorine on chlorides.
43. Under what number is the chemical formula of the gas that is predominantly released during the action of a concentrated solution of nitric acid on copper?
1) N 2, 2) NO 2, 3) NO, 4) H 2.
44. Under what number are the formulas of the reaction products of the combustion of hydrogen sulfide in air with a lack of oxygen?
1) SO 2 + H 2 O, 2) S + H 2 O,
3) SO 3 + H 2 O, 4) SO 2 + H 2.
Give the number of the correct answer.
45. Write an equation for the reaction of the interaction of concentrated sulfuric acid with copper. In your answer, indicate the sum of the coefficients in the reaction equation.
10.4. Answers to tasks of tests for self-control
knowledge of the theory on the topic.
"Redox Reactions"
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Option number 1 |
Option number 2 |
Option number 3 |
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5oxidation Document Increases 4) the degree of oxidation of iron decreases Oxidatively-restorative reaction connection proceeds between: 1) hydrogen chloride and ... potassium bichromate K2Cr2O7 can perform in oxidizing-restorative reactions Function: 1) Both oxidizer and... “Compilation of reaction equations in molecular and ionic forms. Calculation tasks for calculating the mass fraction of a substance in a solution. TargetDocument... oxidizing-restorative reactions, working out a practical skill in drawing up equations oxidizing-restorative reactions electronic balance method. Theory. Oxidatively-restorative called reactions ... | |||||
