During the reaction, compounds are formed. Chemical reactions
Many processes without which it is impossible to imagine our life (such as respiration, digestion, photosynthesis and the like) are associated with various chemical reactions of organic compounds (and inorganic). Let's look at their main types and dwell in more detail on a process called connection.
What is called a chemical reaction
First of all, it is worth giving a general definition of this phenomenon. The phrase under consideration means various reactions of substances of varying complexity, resulting in the formation of products different from the original ones. The substances involved in this process are referred to as “reagents”.
In a letter, the chemical reaction of organic compounds (and inorganic) is recorded using specialized equations. Outwardly, they slightly resemble mathematical examples of addition. However, instead of the equal sign ("\u003d"), arrows are used ("→" or "⇆"). In addition, sometimes there may be more substances on the right side of the equation than on the left. All that is before the arrow is substances before the start of the reaction (the left side of the formula). All that follows (the right side) is compounds formed as a result of a chemical process.
As an example of a chemical equation, we can consider water for hydrogen and oxygen under the influence of an electric current: 2Н 2 О → 2Н 2 + О 2. Water is the starting reagent, and oxygen and hydrogen are products.
As yet another, but already more complex example of the chemical reaction of compounds, we can consider a phenomenon familiar to every housewife who at least once baked sweets. It's about quenching baking soda with table vinegar. The action that is taking place is illustrated by the following equation: NaHCO 3 + 2 CH 3 COOH → 2CH 3 COONa + CO 2 + H 2 O. It is clear from it that during the interaction of sodium bicarbonate and vinegar, the sodium salt of acetic acid, water and carbon dioxide are formed.
By its nature, it occupies an intermediate place between the physical and nuclear.
Unlike the former, the compounds involved in chemical reactions are able to change their composition. That is, from the atoms of one substance, several others can be formed, as in the above equation for the decomposition of water.
Unlike nuclear reactions, chemical does not affect the nuclei of atoms of interacting substances.
What are the types of chemical processes
The distribution of reactions of compounds by species occurs according to different criteria:
- Reversibility / Irreversibility.
- The presence / absence of catalytic substances and processes.
- By absorption / evolution of heat (endothermic / exothermic reaction).
- By the number of phases: homogeneous / heterogeneous and their two hybrid varieties.
- By changing the degrees of oxidation of interacting substances.
Types of chemical processes in inorganic chemistry by the method of interaction
This criterion is special. With its help, four types of reactions are distinguished: compound, substitution, decomposition (decomposition) and exchange.
The name of each of them corresponds to the process that it describes. That is, they unite, in substitution - change into other groups, in decomposition several are formed from one reagent, and in exchange the reaction participants are changed by atoms.
Types of processes by the method of interaction in organic chemistry
Despite the great complexity, the reactions of organic compounds occur according to the same principle as inorganic ones. However, they have slightly different names.
So, the reactions of compound and decomposition are called “addition”, as well as “cleavage” (elimination) and directly organic decomposition (in this section of chemistry there are two types of splitting processes).
Other reactions of organic compounds are substitution (the name does not change), rearrangement (exchange), and redox processes. Despite the similarity of the mechanisms of their course, in organics they are more multifaceted.
Chemical reaction compounds
Having considered the various types of processes in which substances enter in organic and inorganic chemistry, it is worthwhile to dwell in more detail on the compound.
This reaction differs from all others in that, regardless of the number of reagents at its beginning, in the final they all combine into one.
As an example, we can recall the process of slaking lime: CaO + H 2 O → Ca (OH) 2. In this case, the compound of calcium oxide (quicklime) reacts with hydrogen oxide (water). As a result, calcium hydroxide (hydrated lime) is formed and warm steam is released. By the way, this means that this process is truly exothermic.
The reaction equation of the compound
Schematically considered process can be represented as follows: A + BV → ABV. In this formula, ABV is a newly formed A - a simple reagent, and BV is a variant of a complex compound.
It is worth noting that this formula is also characteristic of the process of joining and joining.
Examples of the reaction under consideration are the interaction of sodium oxide and carbon dioxide (NaO 2 + CO 2 (t 450-550 ° C) → Na 2 CO 3), as well as sulfur oxide with oxygen (2SO 2 + O 2 → 2SO 3).
Several complex compounds are also capable of reacting with each other: AB + HB → ABVG. For example, the same sodium oxide and hydrogen oxide: NaO 2 + H 2 O → 2NaOH.
Reaction conditions in inorganic compounds
As was shown in the previous equation, substances of various degrees of complexity are capable of entering into the interaction under consideration.
Moreover, for simple reagents of inorganic origin, redox reactions of the compound (A + B → AB) are possible.
As an example, we can consider the process of obtaining trivalent. For this, the reaction of the compound between chlorine and ferum (iron) is carried out: 3Cl 2 + 2Fe → 2FeCl 3.
If we are talking about the interaction of complex inorganic substances (AB + HB → ABVG), the processes in them can occur, both affecting and not affecting their valency.
As an illustration of this, it is worth considering an example of the formation of calcium hydrogen carbonate from carbon dioxide, hydrogen oxide (water) and white food coloring E170 (calcium carbonate): СО 2 + Н 2 О + СаСО 3 → Са (СО 3) 2. In this case, it has place a classic compound reaction. During its implementation, the valency of the reagents does not change.
A slightly more advanced (than the first) chemical equation 2FeCl 2 + Cl 2 → 2FeCl 3 is an example of a redox process in the interaction of simple and complex inorganic reagents: gas (chlorine) and salt (iron chloride).
Types of addition reactions in organic chemistry
As already indicated in the fourth paragraph, in substances of organic origin the reaction in question is called “addition”. As a rule, complex substances with a double (or triple) bond take part in it.
For example, the reaction between dibrom and ethylene leading to the formation of 1,2-dibromoethane: (C 2 H 4) CH 2 \u003d CH 2 + Br 2 → (C₂H₄Br₂) BrCH 2 - CH 2 Br. By the way, signs similar to equal and minus ("\u003d" and "-"), in this equation show the bonds between the atoms of a complex substance. This is a feature of writing formulas of organic substances.
Depending on which of the compounds act as reagents, several varieties of the joining process under consideration are distinguished:
- Hydrogenation (hydrogen molecules H are added via multiple bonds).
- Hydrohalogenation (hydrogen halide attached).
- Halogenation (addition of halogens Br 2, Cl 2 and the like).
- Polymerization (the formation of substances with a high molecular weight from several low molecular weight compounds).
Examples of the addition reaction (compound)
After listing the varieties of the process in question, it is worthwhile to learn in practice some examples of the reaction of the compound.
As an illustration of hydrogenation, one can pay attention to the equation for the interaction of propene with hydrogen, resulting in propane: (C 3 H 6) CH 3 —CH \u003d CH 2 + H 2 → (C 3 H 8) CH 3 —CH 2 —CH 3.
In organic chemistry, the reaction of the compound (addition) can occur between hydrochloric acid and ethylene with the formation of chloroethane: (C 2 H 4) CH 2 \u003d CH 2 + HCl → CH 3 - CH 2 —Cl (C 2 H 5 Cl). The presented equation is an example of hydrohalogenation.
As for halogenation, it can be illustrated by the reaction between dichloro and ethylene, leading to the formation of 1,2-dichloroethane: (C 2 H 4) CH 2 \u003d CH 2 + Cl 2 → (C₂H₄Cl₂) ClCH 2 -CH 2 Cl.
Many useful substances are formed due to organic chemistry. The reaction of the compound (addition) of ethylene molecules with a radical initiator of polymerization under the influence of ultraviolet radiation confirms this: n CH 2 \u003d CH 2 (R and UV light) → (-CH 2 -CH 2 -) n. The substance formed in this way is well known to everyone under the name of polyethylene.
Various types of packages, bags, dishes, pipes, insulation materials and much more are made from this material. A feature of this substance is the possibility of its secondary processing. Polyethylene owes its popularity to what does not decompose, which is why environmentalists have a negative attitude towards it. However, in recent years, a method has been found for the safe disposal of polyethylene products. For this, the material is treated with nitric acid (HNO 3). After that, certain types of bacteria are able to decompose this substance into safe components.
The reaction of the compound (accession) plays an important role in nature and human life. In addition, it is often used by scientists in laboratories to synthesize new substances for various important studies.
Compound reactions (formation of one complex substance from several simple or complex substances) A \u200b\u200b+ B \u003d AB
Decomposition reactions (decomposition of one complex substance into several simple or complex substances) AB \u003d A + B
Substitution reactions (between simple and complex substances in which atoms of a simple substance replace the atoms of one of the elements in a complex substance): AB + C \u003d AC + B
Exchange reactions (between two complex substances in which substances exchange their constituent parts) AB + SD \u003d AD + CB
1. Indicate the correct definition of the reaction of the compound:
A. The reaction of the formation of several substances from one simple substance;
B. A reaction in which one complex substance is formed from several simple or complex substances.
B. A reaction in which substances exchange their constituent parts.
2. Indicate the correct definition of the substitution reaction:
A. The reaction between base and acid;
B. A reaction between two simple substances;
B. A reaction between substances in which atoms of a simple substance replace the atoms of one of the elements in a complex substance.
3. Indicate the correct definition of the decomposition reaction:
A. A reaction in which several simple or complex substances are formed from one complex substance;
B. A reaction in which substances exchange their constituent parts;
B. Reaction to form oxygen and hydrogen molecules.
4. Indicate the signs of a metabolic reaction:
A. Water formation;
B. Only the formation of gas;
B. Only precipitation;
D. Precipitation, gas formation, or weak electrolyte formation.
5. What type of reaction is the interaction of acid oxides with basic oxides:
A. Exchange reaction;
B. The reaction of the compound;
B. Decomposition reaction;
D. Substitution reaction.
6. What type of reaction is the interaction of salts with acids or with bases:
A. Substitution reactions;
B. Decomposition reactions;
B. Exchange reactions;
D. Reaction compounds.
7. Substances whose formulas KNO3 FeCl2, Na2SO4 are called:
A) salts; B) the grounds; B) acids; D) oxides.
8 . Substances whose formulas are HNO3, HCl, H2SO4 are called:
9 . Substances whose formulas KOH, Fe (OH) 2, NaOH, are called:
A) salts; B) acids; C) the grounds; D) oxides. 10 . Substances whose formulas are NO2, Fe2O3, Na2O are called:
A) salts; B) acids; C) the grounds; D) oxides.
11 . Specify the alkali forming metals:
Cu, Fe, Na, K, Zn, Li.
Answers:
Na, K, Li.
7.1. The main types of chemical reactions
Transformations of substances, accompanied by changes in their composition and properties, are called chemical reactions or chemical interactions. In chemical reactions, there is no change in the composition of atomic nuclei.
Phenomena in which the form or physical state of substances changes or the composition of the nuclei of atoms changes are called physical. An example of physical phenomena is the heat treatment of metals, in which their shape changes (forging), metal melts, sublimates iodine, water turns into ice or steam, etc., as well as nuclear reactions, resulting in the formation of atoms from atoms of some elements. other elements.
Chemical phenomena can be accompanied by physical transformations. For example, as a result of chemical reactions, an electric current arises in a galvanic cell.
Chemical reactions are classified according to various criteria.
1. According to the sign of the thermal effect, all reactions are divided into endothermic (occurring with absorption of heat) and exothermic (proceeding with the release of heat) (see § 6.1).
2. According to the state of aggregation of the starting materials and reaction products are distinguished:
homogeneous reactionsin which all substances are in one phase:
2 KOH (p-p) + H 2 SO 4 (p-p) \u003d K 2 SO (p-p) + 2 H 2 O (g),
CO (g) + Cl 2 (g) \u003d COCl 2 (g),
SiO 2 (k) + 2 MgO (k) \u003d Si (k) + 2 MgO (k).
heterogeneous reactionssubstances in which they are in different phases:
CaO (k) + CO 2 (g) \u003d CaCO 3 (k),
CuSO 4 (solution) + 2 NaOH (solution) \u003d Cu (OH) 2 (q) + Na 2 SO 4 (solution),
Na 2 SO 3 (r-p) + 2HCl (r-p) \u003d 2 NaCl (r-p) + SO 2 (g) + H 2 O (g).
3. By the ability to flow only in the forward direction, as well as in the forward and reverse direction, they are distinguished irreversible and reversiblechemical reactions (see § 6.5).
4. The presence or absence of catalysts distinguish catalytic and non-catalytic reactions (see § 6.5).
5. According to the mechanism of occurrence, chemical reactions are divided into ionic, radical and others (the mechanism of chemical reactions involving organic compounds is considered in the course of organic chemistry).
6. According to the state of oxidation states of the atoms that make up the reacting substances, reactions occurring no change in oxidation state atoms, and with a change in the degree of oxidation of atoms ( redox reactions) (see § 7.2).
7. By changing the composition of the starting materials and reaction products, reactions are distinguished compounds, decomposition, substitution and exchange. These reactions can occur either with or without changes in the oxidation states of the elements, table . 7.1.
Table 7.1
Types of Chemical Reactions
General scheme |
Examples of reactions proceeding without changing the degree of oxidation of elements |
Examples of redox reactions |
|
Connections (from two or more substances one new substance is formed) |
HCl + NH 3 \u003d NH 4 Cl; SO 3 + H 2 O \u003d H 2 SO 4 |
H 2 + Cl 2 \u003d 2HCl; 2Fe + 3Cl 2 \u003d 2FeCl 3 |
|
Decomposition (several new substances are formed from one substance) |
A \u003d B + C + D |
MgCO 3 MgO + CO 2; H 2 SiO 3 SiO 2 + H 2 O |
2AgNO 3 2Ag + 2NO 2 + O 2 |
Substitutions (during the interaction of substances, atoms of one substance replace atoms of another substance in a molecule) |
A + BC \u003d AB + C |
CaCO 3 + SiO 2 CaSiO 3 + CO 2 |
Pb (NO 3) 2 + Zn \u003d Mg + 2HCl \u003d MgCl 2 + H 2 |
|
(two substances exchange their constituent parts, forming two new substances) |
AB + CD \u003d AD + CV |
AlCl 3 + 3NaOH \u003d Ca (OH) 2 + 2HCl \u003d CaCl 2 + 2H 2 O |
7.2. Redox reactions
As mentioned above, all chemical reactions are divided into two groups:
Chemical reactions that occur with a change in the oxidation state of the atoms that make up the reacting substances are called redox.
Oxidation Is the process of electron donation by an atom, molecule or ion:
Na o - 1e \u003d Na +;
Fe 2+ - e \u003d Fe 3+;
H 2 o - 2e \u003d 2H +;
2 Br - - 2e \u003d Br 2 o.
Recovery Is the process of electron attachment by an atom, molecule or ion:
S o + 2e \u003d S 2–;
Cr 3+ + e \u003d Cr 2+;
Cl 2 o + 2e \u003d 2Cl -;
Mn 7+ + 5e \u003d Mn 2+.
Atoms, molecules, or ions that accept electrons are called oxidizing agents. Reductants are atoms, molecules, or ions that donate electrons.
Accepting electrons, the oxidizing agent is reduced during the course of the reaction, and the reducing agent is oxidized. Oxidation is always accompanied by reduction and vice versa. In this way, the number of electrons given away by the reducing agent is always equal to the number of electrons taken by the oxidizing agent.
7.2.1. Oxidation state
The oxidation state is the conditional (formal) charge of an atom in a compound, calculated on the assumption that it consists only of ions. The oxidation state is usually denoted by an Arabic numeral on top of the symbol of the element with the sign “+” or “-”. For example, Al 3+, S 2–.
To find the degree of oxidation are guided by the following rules:
the degree of oxidation of atoms in simple substances is zero;
the algebraic sum of the degrees of oxidation of atoms in a molecule is zero, in a complex ion - the charge of the ion;
the degree of oxidation of alkali metal atoms is always +1;
the hydrogen atom in compounds with non-metals (CH 4, NH 3, etc.) exhibits an oxidation state of +1, and with active metals its oxidation state is –1 (NaH, CaH 2, etc.);
the fluorine atom in the compounds always exhibits an oxidation state of –1;
the oxidation state of the oxygen atom in the compounds is usually –2, except for peroxides (H 2 O 2, Na 2 O 2), in which the oxidation state of oxygen is –1, and some other substances (superperoxides, ozonides, oxygen fluorides).
The maximum positive oxidation state of elements in a group is usually equal to the group number. The exception is fluorine, oxygen, since their highest oxidation state is lower than the number of the group in which they are located. Elements of the copper subgroup form compounds in which their oxidation state exceeds the group number (CuO, AgF 5, AuCl 3).
The maximum negative oxidation state of elements in the main subgroups of the periodic system can be determined by subtracting the group number from eight. For carbon it is 8 - 4 \u003d 4, for phosphorus - 8 - 5 \u003d 3.
In the main subgroups, when passing from elements from top to bottom, the stability of the highest positive oxidation state decreases, in side subgroups, on the contrary, the stability of higher oxidation states increases from top to bottom.
The conventionality of the concept of oxidation state can be demonstrated by the example of some inorganic and organic compounds. In particular, in phosphinic (hypophosphorous) H 3 PO 2, phosphonic (phosphorous) H 3 PO 3 and phosphoric H 3 PO 4 acids, the oxidation states of phosphorus are +1, +3 and +5, respectively, while in all these compounds phosphorus is pentavalent. For carbon in methane CH 4, methanol CH 3 OH, formaldehyde CH 2 O, formic acid HCOOH and carbon monoxide (IV) CO 2, the oxidation states of carbon are –4, –2, 0, +2, and +4, respectively, while as the valency of the carbon atom in all of these compounds is four.
Despite the fact that the degree of oxidation is a conventional concept, it is widely used in the preparation of redox reactions.
7.2.2. The most important oxidizing agents and reducing agents
Typical oxidizing agents are:
1. Simple substances whose atoms are highly electronegative. These are, first of all, the elements of the main subgroups of the VI and VII groups of the periodic system: oxygen, halogens. Of the simplest substances, the most powerful oxidizing agent is fluorine.
2. Compounds containing some metal cations in high oxidation states: Pb 4+, Fe 3+, Au 3+, etc.
3. Compounds containing some complex anions, the elements of which are in high positive degrees of oxidation: 2–, -, and others.
Reducers include:
1. Simple substances, whose atoms have low electronegativity - active metals. Non-metals, for example, hydrogen and carbon, can also exhibit reducing properties.
2. Some metal compounds containing cations (Sn 2+, Fe 2+, Cr 2+), which, giving off electrons, can increase their oxidation state.
3. Some compounds containing such simple ions as, for example, I -, S 2–.
4. Compounds containing complex ions (S 4+ O 3) 2–, (НР 3+ O 3) 2–, in which elements, giving off electrons, can increase their positive oxidation state.
In laboratory practice, the following oxidizing agents are most often used:
potassium permanganate (KMnO 4);
potassium dichromate (K 2 Cr 2 O 7);
nitric acid (HNO 3);
concentrated sulfuric acid (H 2 SO 4);
hydrogen peroxide (H 2 O 2);
manganese (IV) and lead (IV) oxides (MnO 2, PbO 2);
molten potassium nitrate (KNO 3) and melts of some other nitrates.
The reducing agents that are used in laboratory practice include:
- magnesium (Mg), aluminum (Al) and other active metals;
- hydrogen (H 2) and carbon (C);
- potassium iodide (KI);
- sodium sulfide (Na 2 S) and hydrogen sulfide (H 2 S);
- sodium sulfite (Na 2 SO 3);
- tin chloride (SnCl 2).
7.2.3. Classification of redox reactions
Redox reactions are usually divided into three types: intermolecular, intramolecular and disproportionation reactions (self-oxidation-self-reduction).
Intermolecular reactions proceed with a change in the degree of oxidation of atoms that are in various molecules. For example:
2 Al + Fe 2 O 3 Al 2 O 3 + 2 Fe,
C + 4 HNO 3 (conc) \u003d CO 2 + 4 NO 2 + 2 H 2 O.
TO intramolecular reactions reactions include those in which the oxidizing agent and reducing agent are part of the same molecule, for example:
(NH 4) 2 Cr 2 O 7 N 2 + Cr 2 O 3 + 4 H 2 O,
2 KNO 3 2 KNO 2 + O 2.
AT disproportionation reactions (self-oxidation-self-reduction) atom (ion) of the same element is both an oxidizing agent and a reducing agent:
Cl 2 + 2 KOH KCl + KClO + H 2 O,
2 NO 2 + 2 NaOH \u003d NaNO 2 + NaNO 3 + H 2 O.
7.2.4. The basic rules for the preparation of redox reactions
The preparation of redox reactions is carried out according to the steps presented in table. 7.2.
Table 7.2
Stages of the preparation of equations of redox reactions
Act |
|
Identify the oxidizing agent and reducing agent. |
|
To establish the products of the redox reaction. |
|
To compose the balance of electrons and use it to arrange the coefficients of substances that change their oxidation state. |
|
Set the coefficients of other substances involved and formed in the redox reaction. |
|
Check the correct placement of the coefficients by counting the amount of matter of atoms (usually hydrogen and oxygen) located on the left and right sides of the reaction equation. |
We will consider the rules for the preparation of redox reactions using an example of the interaction of potassium sulfite with potassium permanganate in an acidic medium:
1. The definition of the oxidizing agent and reducing agent
The highly oxidized manganese cannot give off electrons. Mn 7+ will accept electrons, i.e. is an oxidizing agent.
The S 4+ ion can donate two electrons and go into S 6+, i.e. is a reducing agent. Thus, in the reaction under consideration, K 2 SO 3 is a reducing agent, and KMnO 4 is an oxidizing agent.
2. Establishment of reaction products
K 2 SO 3 + KMnO 4 + H 2 SO 4?
Giving two electrons an electron, S 4+ goes into S 6+. Potassium sulfite (K 2 SO 3), thus, passes into sulfate (K 2 SO 4). In an acidic medium, Mn 7+ accepts 5 electrons and forms a manganese sulfate (MnSO 4) in a solution of sulfuric acid (medium). As a result of this reaction, additional potassium sulfate molecules are also formed (due to the potassium ions that are part of the permanganate), as well as water molecules. Thus, the reaction in question is written in the form:
K 2 SO 3 + KMnO 4 + H 2 SO 4 \u003d K 2 SO 4 + MnSO 4 + H 2 O.
3. Balancing electrons
To balance the electrons, it is necessary to indicate the oxidation states that change in the reaction under consideration:
K 2 S 4+ O 3 + KMn 7+ O 4 + H 2 SO 4 \u003d K 2 S 6+ O 4 + Mn 2+ SO 4 + H 2 O.
Mn 7+ + 5 e \u003d Mn 2+;
S 4+ - 2 e \u003d S 6+.
The number of electrons given off by the reducing agent should be equal to the number of electrons taken by the oxidizing agent. Therefore, two Mn 7+ and five S 4+ should be involved in the reaction:
Mn 7+ + 5 e \u003d Mn 2+ 2,
S 4+ - 2 e \u003d S 6+ 5.
Thus, the number of electrons given away by the reducing agent (10) will be equal to the number of electrons taken by the oxidizing agent (10).
4. The arrangement of the coefficients in the reaction equation
In accordance with the electron balance, K 5 SO 3 must be set in front of K 2 SO 3 and 2 in front of KMnO 4. We put a coefficient of 6 in front of the potassium sulfate, since one molecule is added to the five K 2 SO 4 molecules formed during the oxidation of potassium sulfite K 2 SO 4 as a result of binding of potassium ions that are part of the permanganate. Since, as an oxidizing agent, two permanganate molecules, also formed on the right side two Manganese sulfate molecules. To bind the reaction products (potassium and manganese ions that are part of permanganate) it is necessary three sulfuric acid molecules, therefore, as a result of the reaction, three water molecules. Finally we get:
5 K 2 SO 3 + 2 KMnO 4 + 3 H 2 SO 4 \u003d 6 K 2 SO 4 + 2 MnSO 4 + 3 H 2 O.
5. Checking the correct placement of the coefficients in the reaction equation
The number of oxygen atoms on the left side of the reaction equation is:
5 · 3 + 2 · 4 + 3 · 4 \u003d 35.
On the right side, this number will be:
6 · 4 + 2 · 4 + 3 · 1 \u003d 35.
The number of hydrogen atoms on the left side of the reaction equation is six and corresponds to the number of these atoms on the right side of the reaction equation.
7.2.5. Examples of redox reactions involving typical oxidizing agents and reducing agents
7.2.5.1. Intermolecular redox reactions
Redox reactions proceeding with the participation of potassium permanganate, potassium dichromate, hydrogen peroxide, potassium nitrite, potassium iodide and potassium sulfide are considered below as examples. Redox reactions involving other typical oxidizing agents and reducing agents are discussed in the second part of the manual (“Inorganic Chemistry”).
Redox reactions involving potassium permanganate
Depending on the medium (acidic, neutral, alkaline), potassium permanganate, acting as an oxidizing agent, gives various reduction products, Fig. 7.1.
Fig. 7.1. The formation of products of the restoration of potassium permanganate in various environments
Below are the reactions of KMnO 4 with potassium sulfide as a reducing agent in various media, illustrating the scheme, Fig. 7.1. In these reactions, the product of the oxidation of the sulfide ion is free sulfur. In an alkaline medium, KOH molecules do not participate in the reaction, but only determine the product of the reduction of potassium permanganate.
5 K 2 S + 2 KMnO 4 + 8 H 2 SO 4 \u003d 5 S + 2 MnSO 4 + 6 K 2 SO 4 + 8 H 2 O,
3 K 2 S + 2 KMnO 4 + 4 H 2 O 2 MnO 2 + 3 S + 8 KOH,
K 2 S + 2 KMnO 4 – (KOH) 2 K 2 MnO 4 + S.
Redox reactions involving potassium dichromate
In an acidic environment, potassium dichromate is a strong oxidizing agent. A mixture of K 2 Cr 2 O 7 and concentrated H 2 SO 4 (chromic peak) is widely used in laboratory practice as an oxidizing agent. Interacting with a reducing agent, one molecule of potassium dichromate receives six electrons, forming compounds of trivalent chromium:
6 FeSO 4 + K 2 Cr 2 O 7 +7 H 2 SO 4 \u003d 3 Fe 2 (SO 4) 3 + Cr 2 (SO 4) 3 + K 2 SO 4 +7 H 2 O;
6 KI + K 2 Cr 2 O 7 + 7 H 2 SO 4 \u003d 3 I 2 + Cr 2 (SO 4) 3 + 4 K 2 SO 4 + 7 H 2 O.
Redox reactions involving hydrogen peroxide and potassium nitrite
Hydrogen peroxide and potassium nitrite exhibit predominantly oxidizing properties:
H 2 S + H 2 O 2 \u003d S + 2 H 2 O,
2 KI + 2 KNO 2 + 2 H 2 SO 4 \u003d I 2 + 2 K 2 SO 4 + H 2 O,
However, when interacting with strong oxidizing agents (such as, for example, KMnO 4), hydrogen peroxide and potassium nitrite act as reducing agents:
5 H 2 O 2 + 2 KMnO 4 + 3 H 2 SO 4 \u003d 5 O 2 + 2 MnSO 4 + K 2 SO 4 + 8 H 2 O,
5 KNO 2 + 2 KMnO 4 + 3 H 2 SO 4 \u003d 5 KNO 3 + 2 MnSO 4 + K 2 SO 4 + 3 H 2 O.
It should be noted that hydrogen peroxide, depending on the medium, is reduced according to the scheme, Fig. 7.2.
Fig. 7.2. Possible hydrogen peroxide reduction products
In this case, the reaction produces water or hydroxide ions:
2 FeSO 4 + H 2 O 2 + H 2 SO 4 \u003d Fe 2 (SO 4) 3 + 2 H 2 O,
2 KI + H 2 O 2 \u003d I 2 + 2 KOH.
7.2.5.2. Intramolecular oxidation-reduction reactions
Intramolecular redox reactions occur, as a rule, upon heating of substances in the molecules of which a reducing agent and an oxidizing agent are present. Examples of intramolecular reduction-oxidation reactions are the processes of thermal decomposition of nitrates and potassium permanganate:
2 NaNO 3 2 NaNO 2 + O 2,
2 Cu (NO 3) 2 2 CuO + 4 NO 2 + O 2,
Hg (NO 3) 2 Hg + NO 2 + O 2,
2 KMnO 4 K 2 MnO 4 + MnO 2 + O 2.
7.2.5.3. Disproportionation Reactions
As noted above, in disproportionation reactions, the same atom (ion) is both an oxidizing agent and a reducing agent. Consider the process of compiling this type of reaction on the example of the interaction of sulfur with alkali.
Characteristic degrees of sulfur oxidation: – 2, 0, +4 and +6. Acting as a reducing agent, elemental sulfur gives away 4 electrons:
So o – 4e \u003d S 4+.
Sulfur – the oxidizer takes two electrons:
S o + 2e \u003d S 2–.
Thus, as a result of the reaction of disproportionation of sulfur, compounds are formed, the oxidation state of the element in which – 2 and to the right +4:
3 S + 6 KOH \u003d 2 K 2 S + K 2 SO 3 + 3 H 2 O.
Disproportionation of nitric oxide (IV) in alkali yields nitrite and nitrate - compounds in which the oxidation states of nitrogen are +3 and +5, respectively:
2 N 4+ O 2 + 2 KOH \u003d KN 3+ O 2 + KN 5+ O 3 + H 2 O,
Disproportionation of chlorine in a cold solution of alkali leads to the formation of hypochlorite, and in hot - chlorate:
Cl 0 2 + 2 KOH \u003d KCl - + KCl + O + H 2 O,
Cl 0 2 + 6 KOH 5 KCl - + KCl 5+ O 3 + 3H 2 O.
7.3. Electrolysis
The redox process that occurs in solutions or melts when a constant electric current is passed through them is called electrolysis. In this case, anion oxidation occurs on the positive electrode (anode). At the negative electrode (cathode) cations are reduced.
2 Na 2 CO 3 4 Na + O 2 + 2CO 2.
During the electrolysis of aqueous solutions of electrolytes, along with transformations of the solute, electrochemical processes can take place involving hydrogen ions and hydroxide ions of water:
cathode (-): 2 Н + + 2е \u003d Н 2,
anode (+): 4 OH - - 4е \u003d О 2 + 2 Н 2 О.
In this case, the recovery process at the cathode occurs as follows:
1. Cations of active metals (up to and including Al 3+) are not reduced at the cathode; instead, hydrogen is reduced.
2. Metal cations located in a row of standard electrode potentials (in a row of voltages) to the right of hydrogen, during electrolysis are restored at the cathode to free metals.
3. Metal cations located between Al 3+ and Н + are reduced at the cathode simultaneously with the hydrogen cation.
The processes taking place in aqueous solutions at the anode depend on the substance of which the anode is made. There are insoluble anodes ( inert) and soluble ( active) As a material of inert anodes, graphite or platinum is used. Soluble anodes are made from copper, zinc and other metals.
During electrolysis of solutions with an inert anode, the following products may form:
1. During the oxidation of halide ions, free halogens are released.
2. During the electrolysis of solutions containing anions SO 2 2–, NO 3 -, PO 4 3– oxygen is released, ie not these ions are oxidized on the anode, but water molecules.
Given the above rules, let us consider as an example the electrolysis of aqueous solutions of NaCl, CuSO 4 and KOH with inert electrodes.
one). In solution, sodium chloride dissociates into ions.
9.1. What are the chemical reactions
Recall that we call chemical reactions any chemical phenomena of nature. In a chemical reaction, some are broken and other chemical bonds form. As a result of the reaction, other substances are obtained from certain chemicals (see Ch. 1).
Performing the homework to § 2.5, you became acquainted with the traditional isolation of the reactions of the four main types from the whole set of chemical transformations, then you proposed their names: reactions of compound, decomposition, substitution and exchange.
Examples of compounds reactions:
C + O 2 \u003d CO 2; (one)
Na 2 O + CO 2 \u003d Na 2 CO 3; (2)
NH 3 + CO 2 + H 2 O \u003d NH 4 HCO 3. (3)
Examples of decomposition reactions:
2Ag 2 O 4Ag + O 2; (four)
CaCO 3 CaO + CO 2; (5)
(NH 4) 2 Cr 2 O 7 N 2 + Cr 2 O 3 + 4H 2 O. (6)
Examples of substitution reactions:
CuSO 4 + Fe \u003d FeSO 4 + Cu; (7)
2NaI + Cl 2 \u003d 2NaCl + I 2; (8)
CaCO 3 + SiO 2 \u003d CaSiO 3 + CO 2. (9)
| Exchange reactions - chemical reactions in which the starting materials seem to exchange their constituent parts. |
Examples of metabolic reactions:
Ba (OH) 2 + H 2 SO 4 \u003d BaSO 4 + 2H 2 O; (10)
HCl + KNO 2 \u003d KCl + HNO 2; (eleven)
AgNO 3 + NaCl \u003d AgCl + NaNO 3. (12)
The traditional classification of chemical reactions does not cover all their diversity - in addition to the reactions of the four main types, there are many more complex reactions.
The isolation of two other types of chemical reactions is based on the participation of two most important non-chemical particles: an electron and a proton.
In the course of certain reactions, complete or partial transfer of electrons from one atom to another occurs. In this case, the oxidation states of the atoms of the elements that make up the starting materials change; from the examples given, these are reactions 1, 4, 6, 7, and 8. These reactions are called redox.
In another group of reactions, a hydrogen ion (H +), that is, a proton, passes from one reacting particle to another. Such reactions are called acid-base reactions or proton transfer reactions.
Among the examples given, such reactions are reactions 3, 10, and 11. By analogy with these reactions, redox reactions are sometimes referred to as electron transfer reactions. You will get acquainted with OVR in § 2, and with KOR in the following chapters.
COMPOUND REACTIONS, DECOMPOSITION REACTIONS, SUBSTITUTION REACTIONS, EXCHANGE REACTIONS, REDOX AND REDUCTION REACTIONS, ACID-BASIC REACTIONS.
Create the reaction equations corresponding to the following schemes:
a) HgO Hg + O 2 ( t); b) Li 2 O + SO 2 Li 2 SO 3; c) Cu (OH) 2 CuO + H 2 O ( t);
d) Al + I 2 AlI 3; d) CuCl 2 + Fe FeCl 2 + Cu; e) Mg + H 3 PO 4 Mg 3 (PO 4) 2 + H 2;
g) Al + O 2 Al 2 O 3 ( t); i) KClO 3 + P P 2 O 5 + KCl ( t); j) CuSO 4 + Al Al 2 (SO 4) 3 + Cu;
l) Fe + Cl 2 FeCl 3 ( t); m) NH 3 + O 2 N 2 + H 2 O ( t); m) H 2 SO 4 + CuO CuSO 4 + H 2 O.
Indicate the traditional type of reaction. Note the redox and acid-base reactions. In redox reactions, indicate which atoms of which elements change their oxidation state.
9.2. Redox reactions
Consider the redox reaction that occurs in blast furnaces in the industrial production of iron (more precisely, cast iron) from iron ore:
Fe 2 O 3 + 3CO \u003d 2Fe + 3CO 2.
Let us determine the oxidation states of the atoms that make up both the starting materials and the reaction products
| Fe 2 O 3 | + | = | 2fe | + |
As you can see, the oxidation state of carbon atoms as a result of the reaction increased, the oxidation state of iron atoms decreased, and the oxidation state of oxygen atoms remained unchanged. Therefore, the carbon atoms in this reaction underwent oxidation, i.e., they lost electrons ( oxidized), and iron atoms - to reduction, that is, electrons ( recovered) (see § 7.16). To characterize the OVR use the concepts oxidizer and reducing agent.
Thus, in our reaction, the oxidizing atoms are iron atoms, and the reducing atoms are carbon atoms.
In our reaction, the oxidizing agent is iron (III) oxide, and the reducing agent is carbon (II) oxide.
In cases where oxidizing and reducing atoms are part of the same substance (example: reaction 6 from the previous paragraph), the concepts of “oxidizing agent” and “reducing agent” are not used.
Thus, typical oxidizing agents are substances that include atoms that tend to attach electrons (in whole or in part), lowering their oxidation state. Of the simple substances, these are primarily halogens and oxygen, to a lesser extent sulfur and nitrogen. Of the complex substances - substances that include atoms in higher oxidation states that are not inclined to form simple ions in these oxidation states: HNO 3 (N + V), KMnO 4 (Mn + VII), CrO 3 (Cr + VI), KClO 3 (Cl + V), KClO 4 (Cl + VII), etc.
Typical reducing agents are substances that include atoms that tend to completely or partially donate electrons, increasing their oxidation state. Of the simple substances, these are hydrogen, alkali and alkaline earth metals, and also aluminum. Of the complex substances - H 2 S and sulfides (S – II), SO 2 and sulfites (S + IV), iodides (I – I), CO (C + II), NH 3 (N – III), etc.
In the general case, almost all complex and many simple substances can exhibit both oxidizing and reducing properties. For example:
SO 2 + Cl 2 \u003d S + Cl 2 O 2 (SO 2 is a strong reducing agent);
SO 2 + C \u003d S + CO 2 (t) (SO 2 - weak oxidizing agent);
C + O 2 \u003d CO 2 (t) (C is a reducing agent);
C + 2Ca \u003d Ca 2 C (t) (C is an oxidizing agent).
Let us return to the reaction we analyzed at the beginning of this section.
| Fe 2 O 3 | + | = | 2fe | + |
Please note that as a result of the reaction, oxidizing atoms (Fe + III) turned into reducing atoms (Fe 0), and reducing atoms (C + II) turned into oxidizing atoms (C + IV). But CO 2 under any conditions is a very weak oxidizing agent, and iron, although it is a reducing agent, is much weaker under these conditions than CO. Therefore, the reaction products do not react with each other, and the reverse reaction does not proceed. The given example is an illustration of the general principle that determines the direction of the flow of OVR:
Redox reactions proceed in the direction of the formation of a weaker oxidizing agent and a weaker reducing agent.
The redox properties of substances can only be compared under the same conditions. In some cases, this comparison can be done quantitatively.
Doing homework for the first paragraph of this chapter, you are convinced that it is quite difficult to select the coefficients in some reaction equations (especially OVR). To simplify this task in the case of redox reactions, the following two methods are used:
but) electronic balance methodand
b) electron-ion balance method.
You will study the electronic balance method now, and the electron-ion balance method is usually studied in higher education institutions.
Both of these methods are based on the fact that electrons in chemical reactions do not disappear anywhere and do not appear from anywhere, that is, the number of electrons received by atoms is equal to the number of electrons given by other atoms.
The number of electrons sent and received in the electronic balance method is determined by the change in the oxidation state of the atoms. When using this method, it is necessary to know the composition of both starting materials and reaction products.
Consider the application of the electronic balance method with examples.
Example 1We compose the equation for the reaction of iron with chlorine. It is known that the product of this reaction is iron (III) chloride. We write the reaction scheme:
Fe + Cl 2 FeCl 3.
Determine the oxidation state of the atoms of all the elements that make up the substances involved in the reaction:
Iron atoms give up electrons, and chlorine molecules take them. Express these processes electronic equations:
Fe - 3 e - \u003d Fe + III,
Cl 2 + 2 e - \u003d 2Cl –I.
In order for the number of electrons given to be equal to the number of received, you must multiply the first electronic equation by two, and the second by three:
| Fe - 3 e - \u003d Fe + III, Cl 2 + 2 e - \u003d 2Cl –I |
2Fe - 6 e - \u003d 2Fe + III, 3Cl 2 + 6 e - \u003d 6Cl –I. |
Introducing the coefficients 2 and 3 into the reaction scheme, we obtain the reaction equation:
2Fe + 3Cl 2 \u003d 2FeCl 3.
Example 2Let us draw up the equation for the combustion reaction of white phosphorus in excess of chlorine. It is known that under these conditions phosphorus (V) chloride is formed:
| + V –I | ||||
| P 4 | + | Cl 2 | PCl 5. |
White phosphorus molecules give up electrons (oxidize), and chlorine molecules take them (are reduced):
| P 4 - 20 e - \u003d 4P + V Cl 2 + 2 e - \u003d 2Cl –I |
1 10 |
2 20 |
P 4 - 20 e - \u003d 4P + V Cl 2 + 2 e - \u003d 2Cl –I |
P 4 - 20 e - \u003d 4P + V 10Cl 2 + 20 e - \u003d 20Cl –I |
The initially obtained factors (2 and 20) had a common divisor by which (as future coefficients in the reaction equation) they were divided. Reaction equation:
P 4 + 10Cl 2 \u003d 4PCl 5.
Example 3 Let us draw up the equation of the reaction proceeding during the burning of iron (II) sulfide in oxygen.
Reaction Scheme:
| + III –II | + IV –II | |||||
| + | O 2 | + |
In this case, both iron (II) atoms and sulfur atoms (- II) are oxidized. The composition of iron (II) sulfide atoms of these elements are in the ratio 1: 1 (see indices in the simplest formula).
Electronic balance:
| 4 | Fe + II - e - \u003d Fe + III S –II - 6 e - \u003d S + IV |
Total give 7 e – |
| 7 | O 2 + 4e - \u003d 2O –II |
Reaction equation: 4FeS + 7O 2 \u003d 2Fe 2 O 3 + 4SO 2.
Example 4. Let us draw up the equation of the reaction occurring during the burning of iron (II) disulfide (pyrite) in oxygen.
Reaction Scheme:
| + III –II | + IV –II | |||||
| + | O 2 | + |
As in the previous example, both iron (II) atoms and sulfur atoms are oxidized here, but with an oxidation state of I. The composition of pyrite includes the atoms of these elements in a ratio of 1: 2 (see indices in the simplest formula). It is in this respect that the iron and sulfur atoms react, which is taken into account when preparing the electronic balance:
| Fe + III - e - \u003d Fe + III 2S –I - 10 e - \u003d 2S + IV |
Total give 11 e – | |
| O 2 + 4 e - \u003d 2O –II |
Reaction equation: 4FeS 2 + 11O 2 \u003d 2Fe 2 O 3 + 8SO 2.
There are more complex cases of OVR, some of which you will meet while doing your homework.
ATOM-OXIDIZER, ATOM-REDUCER, SUBSTANCE-OXIDANT, SUBSTANCE-RESTORER, ELECTRON BALANCE METHOD, ELECTRON EQUATIONS.
1. Make an electronic balance sheet for each OVR equation given in the text of § 1 of this chapter.
2. Make up the equations of OVR that you discovered when completing the assignment to § 1 of this chapter. This time, use the electronic balance method to set the coefficients. 3. Using the electronic balance method, make up the reaction equations corresponding to the following schemes: a) Na + I 2 NaI;
b) Na + O 2 Na 2 O 2;
c) Na 2 O 2 + Na Na 2 O;
d) Al + Br 2 AlBr 3;
d) Fe + O 2 Fe 3 O 4 ( t);
f) Fe 3 O 4 + H 2 FeO + H 2 O ( t);
g) FeO + O 2 Fe 2 O 3 ( t);
i) Fe 2 O 3 + CO Fe + CO 2 ( t);
j) Cr + O 2 Cr 2 O 3 ( t);
l) CrO 3 + NH 3 Cr 2 O 3 + H 2 O + N 2 ( t);
m) Mn 2 O 7 + NH 3 MnO 2 + N 2 + H 2 O;
m) MnO 2 + H 2 Mn + H 2 O ( t);
o) MnS + O 2 MnO 2 + SO 2 ( t)
p) PbO 2 + CO Pb + CO 2 ( t);
c) Cu 2 O + Cu 2 S Cu + SO 2 ( t);
r) CuS + O 2 Cu 2 O + SO 2 ( t);
s) Pb 3 O 4 + H 2 Pb + H 2 O ( t).
9.3. Exothermic reactions. Enthalpy
Why do chemical reactions occur?
To answer this question, let us recall why individual atoms combine into molecules, why an ionic crystal is formed from isolated ions, why the principle of least energy acts when the atomic shell is formed. The answer to all these questions is the same: because it is energetically beneficial. This means that during such processes energy is released. It would seem that chemical reactions should proceed for the same reason. Indeed, many reactions can be carried out, during the course of which energy is released. Energy is released, as a rule, in the form of heat.
If the heat does not have time to be removed during the exothermic reaction, the reaction system heats up.
For example, in a methane combustion reaction
CH 4 (g) + 2O 2 (g) \u003d CO 2 (g) + 2H 2 O (g)
so much heat is released that methane is used as fuel.
The fact that heat is generated in this reaction can be reflected in the reaction equation:
CH 4 (g) + 2O 2 (g) \u003d CO 2 (g) + 2H 2 O (g) + Q.
This is the so-called thermochemical equation. Here is the symbol "+ Q"means that heat is generated during methane combustion. This heat is called thermal effect of the reaction.
Where does the released heat come from?
You know that chemical reactions break and form chemical bonds. In this case, bonds are broken between carbon and hydrogen atoms in CH 4 molecules, as well as between oxygen atoms in O 2 molecules. In this case, new bonds are formed: between the carbon and oxygen atoms in the CO 2 molecules and between the oxygen and hydrogen atoms in the H 2 O molecules. To break the bonds, you need to expend energy (see "binding energy", "atomization energy"), and when forming bonds energy is released. Obviously, if the "new" bonds are stronger than the "old" ones, then more energy will be released than absorbed. The difference between the released and absorbed energy is the thermal effect of the reaction.
The thermal effect (amount of heat) is measured in kilojoules, for example:
2H 2 (g) + O 2 (g) \u003d 2H 2 O (g) + 484 kJ.
Such a record means that 484 kilojoules of heat will be released if two moles of hydrogen react with one mole of oxygen and two moles of gaseous water (water vapor) are formed.
In this way, in thermochemical equations, the coefficients are numerically equal to the amounts of the substance of the reactants and reaction products.
What determines the thermal effect of each specific reaction?
The thermal effect of the reaction depends
a) from the state of aggregation of the starting materials and reaction products,
b) temperature and
c) on whether a chemical transformation occurs at a constant volume or at a constant pressure.
The dependence of the thermal effect of the reaction on the state of aggregation of substances is due to the fact that the processes of transition from one state of aggregation to another (like some other physical processes) are accompanied by the release or absorption of heat. It can also be expressed by a thermochemical equation. An example is the thermochemical equation for condensation of water vapor:
H 2 O (g) \u003d H 2 O (g) + Q.
In thermochemical equations, and, if necessary, in ordinary chemical equations, the state of aggregation of substances is indicated using letter indices:
(g) - gas,
(g) - liquid,
(t) or (cr) is a solid or crystalline substance.
The dependence of the thermal effect on temperature is associated with differences in specific heat
starting materials and reaction products.
Since the volume of the system always increases as a result of an exothermic reaction at constant pressure, part of the energy is spent on work to increase the volume, and the heat generated will be less than in the case of the same reaction at a constant volume.
The thermal effects of reactions are usually calculated for reactions occurring at a constant volume at 25 ° C and are indicated by Q o.
If energy is released only in the form of heat, and a chemical reaction proceeds at a constant volume, then the thermal effect of the reaction ( Q V) is equal to the change internal energy (D U) substances participating in the reaction, but with the opposite sign:
Q V \u003d - U.
Under the internal energy of the body is understood the total energy of intermolecular interactions, chemical bonds, the ionization energy of all electrons, the energy of nucleon bonds in nuclei and all other known and unknown forms of energy "stored" by this body. The sign “-” is due to the fact that when heat is released, the internal energy decreases. I.e
U= – Q V .
If the reaction proceeds at constant pressure, then the volume of the system may vary. Part of the internal energy is also spent on the work to increase the volume. In this case
U \u003d -(Q P + A) = –(Q P + P V),
where Q p - the thermal effect of the reaction proceeding at constant pressure. From here
Q P \u003d - U - P V .
Value equal to U + P Vgot the name enthalpy change and is denoted by D H.
H \u003d U + P V.
Hence
Q P \u003d - H.
Thus, with the release of heat, the enthalpy of the system decreases. Hence the old name for this quantity: "heat content".
In contrast to the thermal effect, a change in enthalpy characterizes the reaction regardless of whether it occurs at a constant volume or constant pressure. Thermochemical equations written using enthalpy changes are called thermochemical equations in thermodynamic form. The value of the change in enthalpy under standard conditions (25 ° C, 101.3 kPa), indicated by H o. For example:
2H 2 (g) + O 2 (g) \u003d 2H 2 O (g) H o \u003d - 484 kJ;
CaO (cr) + H 2 O (w) \u003d Ca (OH) 2 (cr) H o \u003d - 65 kJ.
The dependence of the amount of heat released in the reaction ( Q) from the thermal effect of the reaction ( Q o) and the amount of substance ( n B) one of the participants in the reaction (substance B - the starting material or reaction product) is expressed by the equation:
Here B is the amount of substance B, given by the coefficient in front of the formula of substance B in the thermochemical equation.
A task
Determine the amount of hydrogen substance burned in oxygen if 1694 kJ of heat is released.
Decision
2H 2 (g) + O 2 (g) \u003d 2H 2 O (g) + 484 kJ. |
|
|
Q \u003d 1694 kJ, 6. The thermal effect of the reaction of the interaction of crystalline aluminum with gaseous chlorine is 1408 kJ. Write down the thermochemical equation of this reaction and determine the mass of aluminum required to obtain 2816 kJ of heat using this reaction. 9.4. Endothermic reactions. Entropy In addition to exothermic reactions, reactions are possible, during which heat is absorbed, and if it is not supplied, the reaction system is cooled. Such reactions are called endothermic. The thermal effect of such reactions is negative. For example: Thus, the energy released during the formation of bonds in the products of these and similar reactions is less than the energy needed to break the bonds in the starting materials.
Take two flasks and fill one of them with nitrogen (colorless gas), and the other with nitrogen dioxide (brown gas) so that both the pressure and the temperature in the flasks are the same. It is known that these substances do not enter into a chemical reaction between themselves. We tightly connect the flasks with the necks and set them vertically, so that the flask with heavier nitrogen dioxide is at the bottom (Fig. 9.1). After some time, we will see that brown nitrogen dioxide gradually spreads to the upper flask, and colorless nitrogen penetrates the lower. As a result, the gases are mixed, and the color of the contents of the flasks becomes the same. In this way,
Equations of the relation between entropy ( S) and other quantities are studied in physics and physical chemistry courses. Unit of measurement of entropy [ S] \u003d 1 J / K. G \u003d H - T S The condition for spontaneous reaction: G< 0. At low temperatures, the factor determining the possibility of the reaction to a greater extent is the energy factor, and at high - the entropy. From the above equation, in particular, one can see why the decomposition reactions that do not occur at room temperature (entropy increases) begin to occur at elevated temperatures. ENDOTHERMAL REACTION, ENTROPY, ENERGY FACTOR, ENTROPY FACTOR, GIBBS ENERGY. 2CuO (cr) + C (graphite) \u003d 2Cu (cr) + CO 2 (g) is –46 kJ. Write down the thermochemical equation and calculate what energy needs to be spent for receiving 1 kg of copper on such reaction. CaCO 3 (cr) \u003d CaO (cr) + CO 2 (g) - 179 kJ formed 24.6 liters of carbon dioxide. Determine how much heat was used up uselessly. How many grams of calcium oxide formed? |
Definition
Chemical reaction called the transformation of substances in which there is a change in their composition and (or) structure.
Most often, chemical reactions are understood as the process of transformation of starting materials (reagents) into final substances (products).
Chemical reactions are written using chemical equations containing the formulas of the starting materials and reaction products. According to the law of conservation of mass, the number of atoms of each element in the left and right sides of the chemical equation is the same. Typically, the formulas of the starting materials are written on the left side of the equation, and the product formulas on the right. Equality of the number of atoms of each element in the left and right sides of the equation is achieved by arranging integer stoichiometric coefficients in front of the substance formulas.
Chemical equations may contain additional information about the characteristics of the reaction: temperature, pressure, radiation, etc., which is indicated by the corresponding symbol above (or “below”) the equal sign.
All chemical reactions can be grouped into several classes that have certain characteristics.
Classification of chemical reactions according to the number and composition of the starting and forming substances
According to this classification, chemical reactions are divided into reactions of compound, decomposition, substitution, exchange.
As a result compound reactions from two or more (complex or simple) substances one new substance is formed. In general terms, the equation of such a chemical reaction will look like this:
For example:
CaCO 3 + CO 2 + H 2 O \u003d Ca (NSO 3) 2
SO 3 + H 2 O \u003d H 2 SO 4
2Mg + O 2 \u003d 2MgO.
2 FeCl 2 + Cl 2 \u003d 2 FeCl 3
The reactions of the compound are in most cases exothermic, i.e. occur with the release of heat. If simple substances are involved in the reaction, then such reactions are most often redox (redox), i.e. occur with a change in the oxidation state of the elements. It is impossible to say unequivocally whether the reaction of the compound between complex substances cannot be related to OVR.
Reactions as a result of which several other new substances (complex or simple) are formed from one complex substance belong to decomposition reactions. In general terms, the equation for the chemical decomposition reaction will look like this:
For example:
CaCO 3 CaO + CO 2 (1)
2H 2 O \u003d 2H 2 + O 2 (2)
CuSO 4 × 5H 2 O \u003d CuSO 4 + 5H 2 O (3)
Cu (OH) 2 \u003d CuO + H 2 O (4)
H 2 SiO 3 \u003d SiO 2 + H 2 O (5)
2SO 3 \u003d 2SO 2 + O 2 (6)
(NH 4) 2 Cr 2 O 7 \u003d Cr 2 O 3 + N 2 + 4H 2 O (7)
Most decomposition reactions occur upon heating (1,4,5). Possible decomposition by electric current (2). The decomposition of crystalline hydrates, acids, bases and salts of oxygen-containing acids (1, 3, 4, 5, 7) proceeds without changing the oxidation states of the elements, i.e. these reactions are not related to OVR. OVR decomposition reactions include the decomposition of oxides, acids, and salts formed by elements in higher oxidation states (6).
Decomposition reactions are also found in organic chemistry, but under other names - cracking (8), dehydrogenation (9):
C 18 H 38 \u003d C 9 H 18 + C 9 H 20 (8)
C 4 H 10 \u003d C 4 H 6 + 2H 2 (9)
At substitution reactions a simple substance interacts with a complex, forming a new simple and new complex substance. In general terms, the equation of a chemical substitution reaction will look like this:
For example:
2Al + Fe 2 O 3 \u003d 2Fе + Al 2 O 3 (1)
Zn + 2CHl \u003d ZnCl 2 + H 2 (2)
2KBr + Cl 2 \u003d 2KSl + Br 2 (3)
2KSlO 3 + l 2 \u003d 2KlO 3 + Cl 2 (4)
CaCO 3 + SiO 2 \u003d CaSiO 3 + CO 2 (5)
Ca 3 (PO 4) 2 + ЗSiO 2 \u003d ЗСаSiO 3 + Р 2 О 5 (6)
CH 4 + Cl 2 \u003d CH 3 Cl + Hcl (7)
Most of the substitution reactions are redox (1 - 4, 7). Examples of decomposition reactions in which there is no change in the oxidation state are few (5, 6).
Exchange reactions called reactions occurring between complex substances in which they exchange their constituent parts. Typically, this term is used for reactions involving ions in an aqueous solution. In general terms, the equation for the chemical exchange reaction will look like this:
AB + CD \u003d AD + CB
For example:
CuO + 2HCl \u003d CuCl 2 + H 2 O (1)
NaOH + HCl \u003d NaCl + H 2 O (2)
NaHCO 3 + Hcl \u003d NaCl + H 2 O + CO 2 (3)
AgNО 3 + КВr \u003d АgВr ↓ + КNО 3 (4)
СrСl 3 + ЗNаОН \u003d Сr (ОН) 3 ↓ + ЗNаСl (5)
Metabolic reactions are not redox. A special case of these metabolic reactions is the neutralization reaction (the reaction of the interaction of acids with alkalis) (2). Exchange reactions proceed in the direction where at least one of the substances is removed from the reaction sphere in the form of a gaseous substance (3), a precipitate (4, 5) or a slightly dissociating compound, most often water (1, 2).
Classification of chemical reactions by changes in oxidation states
Depending on the change in the oxidation state of the elements that make up the reagents and reaction products, all chemical reactions are divided into redox (1, 2) and occurring without changing the oxidation state (3, 4).
2Mg + CO 2 \u003d 2MgO + C (1)
Mg 0 - 2e \u003d Mg 2+ (reducing agent)
C 4+ + 4e \u003d C 0 (oxidizing agent)
FeS 2 + 8HNO 3 (conc) \u003d Fe (NO 3) 3 + 5NO + 2H 2 SO 4 + 2H 2 O (2)
Fe 2+ -e \u003d Fe 3+ (reducing agent)
N 5+ + 3e \u003d N 2+ (oxidizing agent)
AgNO 3 + HCl \u003d AgCl ↓ + HNO 3 (3)
Ca (OH) 2 + H 2 SO 4 \u003d CaSO 4 ↓ + H 2 O (4)
Classification of chemical reactions by thermal effect
Depending on whether heat (energy) is released or absorbed during the reaction, all chemical reactions are conventionally divided into exo - (1, 2) and endothermic (3), respectively. The amount of heat (energy) released or absorbed during the reaction is called the thermal effect of the reaction. If the equation indicates the amount of released or absorbed heat, then such equations are called thermochemical.
N 2 + 3H 2 \u003d 2NH 3 +46.2 kJ (1)
2Mg + O 2 \u003d 2MgO + 602, 5 kJ (2)
N 2 + O 2 \u003d 2NO - 90.4 kJ (3)
Classification of chemical reactions in the direction of the reaction
In the direction of the reaction, there are reversible (chemical processes, the products of which are able to react with each other under the same conditions in which they are obtained, with the formation of starting materials) and irreversible (chemical processes, whose products are not able to react with each other with the formation of starting materials )
For reversible reactions, the equation in general form is usually written as follows:
A + B ↔ AB
For example:
СН 3 СОО + С 2 Н 5 ОН↔ Н 3 СООС 2 Н 5 + Н 2 О
The following reactions can serve as examples of irreversible reactions:
2KSlO 3 → 2KSl + ZO 2
С 6 Н 12 О 6 + 6О 2 → 6СО 2 + 6Н 2 О
Evidence of the irreversibility of the reaction can be the allocation as a reaction product of a gaseous substance, a precipitate or a slightly dissociating compound, most often water.
Classification of chemical reactions by the presence of a catalyst
From this point of view, catalytic and non-catalytic reactions are distinguished.
A catalyst is a substance that accelerates the course of a chemical reaction. Reactions proceeding with the participation of catalysts are called catalytic. Some reactions are generally impossible without the presence of a catalyst:
2H 2 O 2 \u003d 2H 2 O + O 2 (MnO 2 catalyst)
Often, one of the reaction products serves as a catalyst accelerating this reaction (autocatalytic reactions):
MeO + 2HF \u003d MeF 2 + H 2 O, where Me is a metal.
Examples of solving problems
EXAMPLE 1
