EntranceChemistry tutor manual. Theory of electrical dissociation Electrolytic dissociation of substances and ion exchange reactions
During the lesson we will study the topic “ Electrolytic dissociation. Ion exchange reactions". Let's consider the theory of electrolytic dissociation and get acquainted with the definition of electrolytes. Let's get acquainted with the physical and chemical theory of solutions. Let's consider, in the light of the theory of electrolytic dissociation, the definition of bases, acids and salts, and also learn how to compose equations for ion exchange reactions and learn about the conditions for their irreversibility.
Topic: Solutions and their concentration, dispersed systems, electrolytic dissociation
Lesson: Electrolytic dissociation. Ion exchange reactions
Even at the dawn of the study of electrical phenomena, scientists noticed that not only metals, but also solutions can conduct current. But not all of them. Thus, aqueous solutions of table salt and other salts, solutions of strong acids and alkalis conduct current well. Solutions of acetic acid, carbon dioxide and sulfur dioxide conduct it much worse. But solutions of alcohol, sugar and most others organic compounds are not carried out at all electric current.
Electric current is the directed movement of free charged particles . In metals, such movement is carried out due to relatively free electrons, electron gas. But not only metals are capable of conducting electric current.
Electrolytes - These are substances whose solutions or melts conduct electric current.
Non-electrolytes - These are substances whose solutions or melts do not conduct electric current.
To describe the electrical conductivity of some solutions, it is necessary to understand what a solution is. TO end of the 19th century centuries, there were 2 main theories of solutions:
· Physical. According to this theory, the solution - it is a purely mechanical mixture of components, and there is no interaction between particles in it. She described the properties of electrolytes well, but had certain difficulties in describing electrolyte solutions.
· Chemical. According to this theory, during dissolution, a chemical reaction occurs between the solute and the solvent. This is confirmed by the presence of a thermal effect upon dissolution, as well as a change in color. For example, when white anhydrous copper sulfate is dissolved, a saturated blue solution is formed.
The truth is between these two extreme points. Namely , both chemical and physical processes occur in solutions.
Rice. 1. Svante Arrhenius
In 1887, the Swedish physical chemist Svante Arrhenius (Fig. 1), studying the electrical conductivity of aqueous solutions, suggested that in such solutions substances disintegrate into charged particles - ions, which can move to electrodes - a negatively charged cathode and a positively charged anode.
This is the reason for the electric current in solutions. This process is called (literal translation - splitting, decomposition under the influence of electricity). This name also suggests that dissociation occurs under the influence of an electric current. Further research showed that this is not the case: ions are onlycharge carriers in solution and exist in it regardless of whether it passes throughcurrent solution or not. With the active participation of Svante Arrhenius, the theory of electrolytic dissociation was formulated, which is often named after this scientist. The main idea of this theory is that electrolytes spontaneously disintegrate into ions under the influence of a solvent. And it is these ions that are charge carriers and are responsible for the electrical conductivity of the solution.
1. Electrolytes in solutions under the influence of a solvent spontaneously disintegrate into ions. This process is called electrolytic dissociation. Dissociation can also occur when solid electrolytes melt.
2. Ions differ from atoms in composition and properties. In aqueous solutions, ions are in a hydrated state. Ions in the hydrated state differ in properties from ions in gaseous state substances. This is explained as follows: ionic compounds already initially contain cations and anions. When dissolved, a water molecule begins to approach charged ions: positive pole - to negative ion, negative pole - to the positive. The ions are called hydrated (Fig. 2).
Rice. 2
3. In solutions or melts of electrolytes, ions move chaotically, but when an electric current is passed, the ions move directionally: cations - towards the cathode, anions - to the anode.
In the light of the theory of electrolytic dissociation, bases, acids and salts can be defined as electrolytes.
Grounds- these are electrolytes, as a result of the dissociation of which in aqueous solutions only one type of anion is formed: hydroxide anion: OH -.
NaOH ↔ Na + + OH −
The dissociation of bases containing several hydroxyl groups occurs in stages:
Ba(OH) 2 ↔ Ba(OH) + + OH − First stage
Ba(OH) + ↔ Ba 2+ + OH − Second stage
Ba(OH) 2 ↔ Ba 2+ + 2 OH − Summary equation
Acids - These are electrolytes, as a result of the dissociation of which in aqueous solutions only one type of cations is formed: H +. A hydrogen ion is precisely a hydrated proton and is designated H 3 O + , but for simplicity we write H + .
HNO 3 ↔ H + + NO 3 −
Polybasic acids dissociate stepwise:
H 3 PO 4 ↔ H + + H 2 PO 4 - First stage
H 2 PO 4 - ↔ H + + HPO 4 2- Second stage
HPO 4 2- ↔ H + + PO 4 3- Third stage
H 3 PO 4 ↔ 3H + + PO 4 3- Summary equation
Salts -
These are electrolytes that dissociate in aqueous solutions into metal cations and anions of the acid residue.
Na 2 SO 4 ↔ 2Na + + SO 4 2−
Medium salts - These are electrolytes that dissociate in aqueous solutions into metal cations or ammonium cations and acid residue anions.
Basic salts - These are electrolytes that dissociate in aqueous solutions into metal cations, hydroxide anions and acid residue anions.
Acid salts - These are electrolytes that dissociate in aqueous solutions into metal cations, hydrogen cations and acid residue anions.
Double salts - These are electrolytes that dissociate in aqueous solutions into cations of several metals and anions of an acidic residue.
KAl(SO 4) 2 ↔ K + + Al 3+ + 2SO 4 2
Mixed salts - these are electrolytes that dissociate in aqueous solutions into metal cations and anions of several acidic residues
Electrolytic dissociation to varying degrees - the process is reversible. But when some compounds are dissolved, the dissociation equilibrium is largely shifted towards the dissociated form. In solutions of such electrolytes, dissociation occurs almost irreversibly. Therefore, when writing dissociation equations for such substances, either an equal sign or a straight arrow is written, indicating that the reaction occurs almost irreversibly. Such substances are called strong electrolytes.
Weak are called electrolytes in which dissociation occurs insignificantly. When writing, use the reversibility sign. Table 1.
To quantify the strength of the electrolyte, the concept electrolytic degree dissociation .
The strength of an electrolyte can also be characterized using constants chemical equilibrium dissociation. It's called the dissociation constant.
Factors influencing the degree of electrolytic dissociation:
Nature of the electrolyte
Concentration of electrolyte in solution
· Temperature
As the temperature increases and the solution is diluted, the degree of electrolytic dissociation increases. Therefore, it is possible to evaluate the strength of an electrolyte only by comparing them under the same conditions. The standard is t = 18 0 C and c = 0.1 mol/l.
Homework
1. No. 6-8 (p. 48) Rudzitis G.E. Chemistry. Basics general chemistry. 11th grade: textbook for educational institutions: basic level/ G.E. Rudzitis, F.G. Feldman. - 14th ed. - M.: Education, 2012.
2. How can you prove that ions have a charge if they are colored?
3. What causes the crimson color of potassium permanganate?
Purpose of the work. Acquire skills in drawing up molecular and ionic equations for reactions occurring in electrolyte solutions. Learn to determine the direction of ionic reactions.
When some substances are dissolved in water (or other polar solvents), under the influence of solvent molecules, the molecules of the substance disintegrate into ions. As a result of this process, the solution contains not only solvent and solute molecules, but also the resulting ions. Solutions of substances that, when dissolved in water or other polar solvents, disintegrate into ions are called electrolytes.
The process of dissolution of molecules of a dissolved substance (electrolyte) into ions under the influence of polar solvent molecules is called electrolytic dissociation.
Electrolyte solutions have ionic electrical conductivity (in transfer electric charges ions are involved) and are conductors of the second kind.
A quantitative characteristic of the process of decomposition of a dissolved substance into ions is the degree of electrolytic dissociation – α. The degree of dissociation is the ratio of the number of molecules of a dissolved substance that have broken up into ions in a solution (n) to the total number of dissolved molecules (N):
The degree of electrolytic dissociation is determined experimentally and is expressed either in fractions of a unit or as a percentage. The degree of electrolyte dissociation depends on the nature of the electrolyte, concentration and temperature.
According to the degree of dissociation of the electrolyte in a solution with molar con-
equivalent concentration equal to 0.1 mol/l (0.1 N), solutions are conditionally
They are divided into three groups: strong, weak and medium electrolytes. If in
0.1 n. in an electrolyte solution α > 0.3 (30 \%) the electrolyte is considered a strong electrolyte, α ≤ 0.03 (3 \%) is considered a weak electrolyte. Electrolytes with intermediate values degrees of dissociation are considered average.
Strong electrolytes, if the solvent is water, are
– acids: HNO3, H2SO4, HCNS, HCl, HClO3, HClO4, HBr, HBrO3, HBrO4, HI, HIO3 HMnO4, H2SeO4, HReO4, HTcO4; as well as acids Н2СrO4, H4P2O7, H2S2O6, which are strong in the first stage of dissociation, i.e., when the first H+ ion is removed;
– bases: hydroxides of alkali (Li, Na, K, Rb, Cs, Fr) and alkaline earth metals (Ca, Sr, Ba, Ra): LiOH, NaOH, KOH, RbOH, CsOH, FrOH, Ca(OH)2 , Ba(OH)2, Sr(OH)2; Ra(OH)2; as well as TlOH;
– most salts. Exception: Fe(SCN)3, Mg(CN)2, HgCl2, Hg(CN)2.
Weak electrolytes include:
– acids: H2CO3, HClO, H2S, H3BO3, HCN, H2SO3, H2SiO3, CH3COOH, HCOOH, H2C2O4, etc. (Appendix, Table 2);
– bases (p- and d-elements): Be(OH)2, Mg(OH)2, Fe(OH)2, Zn(OH)2; ammonium hydroxide NH4OH, as well as organic bases – amines (CH3NH2) and ampholytes (H3N+CH2COOˉ).
Water is a very weak electrolyte (H2O) α = 2·10-9, i.e.
Water molecules can also break down into ions due to the interaction of molecules with each other.
Strong electrolytes are substances that, when dissolved in water, completely disintegrate into ions, i.e., dissociate almost completely. After the bond between the ions in the electrolyte molecule is broken under the action of water molecules, the resulting ions surround themselves with water molecules and therefore are in a hydrated state in solution. Taking into account the hydration of ions, the electrolytic dissociation equation could be written as follows:
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The dissociation equation for a strong electrolyte is written in a simplified manner,
For example:
NaCl → Na+ + Clˉ;
HNO3 → H+ + NO3ˉ;
Ba(OH)2 → Ba2+ + 2OHˉ
Weak electrolytes include substances that, when dissolved in water, partially dissociate into ions. An equilibrium is established between ions, whose concentration in the solution is low, and actually existing undissociated molecules:
CH3COOH ⇄ CH3COOˉ + H+; H2O ⇄ H+ + OHˉ.
This notation means that two things happen simultaneously in the solution.
process: the breakdown of molecules into ions and the formation of molecules from ions. Equilibrium in solutions of weak electrolytes is shifted towards the initial products, therefore weak electrolytes in solution exist predominantly in the form of molecules.
Chemical properties electrolyte solutions depend on the properties of the ions and molecules in the solution. The direction of reactions between ions and molecules in electrolyte solutions is determined by the possibility of the formation of poorly soluble substances or weak electrolytes. If the reaction does not result in the formation of a poorly soluble substance or a weak electrolyte, then such a reaction cannot proceed. For example, when merging solutions of sodium nitrate and potassium chloride, the reaction does not proceed, since exchange reaction Any poorly soluble substance or weak electrolyte cannot be formed from the ions in solution. These salts are strong electrolytes and are highly soluble in water, so the solution will contain
this mixture of ions:
Na+ + NO3ˉ + K+ + Clˉ,
of which the original substances consisted. Therefore, in this case it is impossible to write the molecular equation of the exchange reaction
NaNO3 + KCl ≠ KNO3 + NaCl.
The reaction occurring in solution can be represented as:
Molecular reaction equation;
Ion-molecular equation (full or abbreviated).
A reaction equation containing only the formulas of non-dissociated substances is called a molecular equation. The molecular form of the equation shows which substances and in what quantities are involved in the reaction. It allows you to make the necessary calculations related to this reaction. An equation containing the formulas of undissociated weak electrolytes and ions of strong electrolytes is called a complete ionic or ion-molecular reaction equation.
By reducing the same products on the left and right sides of the ionic-molecular reaction equation, we obtain an abbreviated or brief ionic reaction equation. An ionic equation that does not contain identical substances (ions or molecules) on the left and right sides of the reaction is called an abbreviated or short ionic equation of the reaction. This equation reflects the essence of the reaction taking place.
When writing ionic reaction equations, you must remember:
1) strong electrolytes should be written in the form of separate components
their constituent ions;
2) weak electrolytes and poorly soluble substances should be written down
pour in the form of molecules.
As an example, consider the interaction of soda with acid. In the molecular equation of a reaction, the starting materials and reaction products are written in the form of molecules:
Na2CO3 + H2SO4 = Na2SO4 + CO2 + H2O.
Taking into account that in an aqueous solution, electrolyte molecules
com decompose into ions, the complete ionic equation of this reaction has the form
CO2–In the ionic equation, weak electrolytes, gases, and poorly soluble substances are written as molecules. The ↓ sign at the formula of a substance means that this substance is removed from the sphere of reaction in the form
precipitate, and the sign indicates that the substance is removed from the reaction sphere in the form of a gas.
Substances whose molecules completely dissociate into ions (strong electrolytes) are written as ions. The sum of the electric charges on the left side of the equation must be equal to the sum of the electric charges on the right side.
When writing ionic equations, one should be guided by the table of solubility of acids, bases and salts in water, i.e., be sure to check the solubility of reagents and products, noting this in the equations, as well as the table of dissociation constants of weak electrolytes (Appendix, table. 1 and 2). Let's look at examples of recording some ion-molecular equations.
Example 1. Formation of poorly and poorly soluble compounds (precipitate).
a) Formation of barium sulfate
Molecular equation of the reaction:
BaCl2 + Na2SO4 = BaSO4↓ + 2NaCl.
Complete ionic (ionic-molecular) reaction equation:
Ba2+ + 2Clˉ + 2Na+ + SO4 ˉ = BaSO4↓ + 2Na
CO2–CO2 + H2O (abbreviated ionic equation).
Example 3. Formation of a weak electrolyte.
2Na+ + 2OH– +2H+ + SO 2–(full ionic equation)
2OH– + 2H+ = 2H2O (abbreviated ionic equation).
The reaction of neutralization of a strong acid with a strong base is reduced to the interaction of hydrogen ions with hydroxide ions;
b) weak acid:
2NaNO2 + H2SO4 = 2HNO2 + Na2SO4 (molecular equation)
NH+(full ionic equation)
NH4OH (short ionic equation).
Strong bases displace weak bases from their salts.
Example 4. When among the starting compounds and reaction products there is a weak electrolyte or a poorly soluble substance, then the equation uses -
There is a balance sign “⇄”. The equilibrium in the reaction shifts towards a weaker electrolyte or slightly soluble substance, which is indicated
icon (↷)..
a) CH3COOH + NaOH ⇄ CH3COONa + H2O
CH3COOH + OHˉ ⇄ CH3COOˉ + H2O (↷).
As a result of the reaction, a weaker electrolyte is formed - water. Equal-
This shifts towards a direct reaction.
b) CaSO4↓ + Na2CO3 ⇄ CaCO3↓ + Na2SO4;
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⇄ CaCO3↓ + 2 Na+
As a result of the reaction, a less soluble salt is formed - calcium carbonate.
tion. The equilibrium shifts towards the forward reaction.
Example 5. Write three possible molecular equations for the reaction,
corresponding to the abbreviated ionic equation: CH3COO– + H+ = CH3COOH.
Solution. The left side of the ionic equation shows the free ions CH3COO– and H+. These ions are formed during the dissociation of any soluble strong electrolytes. CH3COO– ions can be formed during the dissociation of, for example, salts KCH3COO, NaCH3COO, Mg (CH3COO)2; donors
new H+ can be any strong acid. Molecular reaction equations,
which this molecular-ionic equation corresponds to can be:
1. KCH3COO + HCl = CH3COOH + KCl;
2. NaCH3COO + HNO3 = CH3COOH + NaNO3;
3. Mg(CH3COO)2 + H2SO4 = 2 CH3COOH + MgSO4.
Safety precautions
1. Take special care when working with solutions of acids and alkalis, do not allow them to come into contact with skin and clothing.
2. If a toxic gaseous product is released during the experiment, be sure to conduct the experiment in a fume hood with ventilation running.
3. Be careful when working with toxic salts and their solutions (barium salts, chromium, copper, etc.).
Electrolytes and non-electrolytes
It is known from physics lessons that solutions of some substances are capable of conducting electric current, while others are not.
Substances whose solutions conduct electric current are called electrolytes.
Substances whose solutions do not conduct electric current are called non-electrolytes. For example, solutions of sugar, alcohol, glucose and some other substances do not conduct electricity.
Electrolytic dissociation and association
Why do electrolyte solutions conduct electric current?
The Swedish scientist S. Arrhenius, studying the electrical conductivity of various substances, came to the conclusion in 1877 that the cause of electrical conductivity is the presence in solution ions, which are formed when an electrolyte is dissolved in water.
The process of electrolyte breaking down into ions is called electrolytic dissociation.
S. Arrhenius, who adhered to the physical theory of solutions, did not take into account the interaction of the electrolyte with water and believed that there were free ions in solutions. In contrast, Russian chemists I.A. Kablukov and V.A. Kistyakovsky applied to the explanation of electrolytic dissociation chemical theory D.I. Mendeleev and proved that when an electrolyte is dissolved, a chemical interaction of the dissolved substance with water occurs, which leads to the formation of hydrates, and then they dissociate into ions. They believed that solutions contained not free, not “naked” ions, but hydrated ones, that is, “dressed in a coat” of water molecules.
Water molecules are dipoles(two poles), since the hydrogen atoms are located at an angle of 104.5°, due to which the molecule has an angular shape. The water molecule is shown schematically below.
As a rule, substances dissociate most easily with ionic bond and, accordingly, with an ionic crystal lattice, since they already consist of ready-made ions. When they dissolve, the water dipoles are oriented with oppositely charged ends around the positive and negative ions electrolyte.
Mutual attractive forces arise between electrolyte ions and water dipoles. As a result, the bond between the ions weakens, and the ions move from the crystal to the solution. It is obvious that the sequence of processes occurring during the dissociation of substances with ionic bonds (salts and alkalis) will be as follows:
1) orientation of water molecules (dipoles) near the ions of the crystal;
2) hydration (interaction) of water molecules with ions of the surface layer of the crystal;
3) dissociation (decay) of the electrolyte crystal into hydrated ions.
Simplified processes can be reflected using the following equation:
Electrolytes whose molecules have a covalent bond (for example, molecules of hydrogen chloride HCl, see below) dissociate similarly; only in this case, under the influence of water dipoles, the transformation of covalent polar connection to ionic; The sequence of processes occurring in this case will be as follows:
1) orientation of water molecules around the poles of electrolyte molecules;
2) hydration (interaction) of water molecules with electrolyte molecules;
3) ionization of electrolyte molecules (conversion of a covalent polar bond into an ionic one);
4) dissociation (decay) of electrolyte molecules into hydrated ions.
In a simplified way, the process of dissociation of hydrochloric acid can be reflected using the following equation:
It should be taken into account that in electrolyte solutions, chaotically moving hydrated ions can collide and recombine with each other. This reverse process called an association. Association in solutions occurs in parallel with dissociation, therefore the reversibility sign is put in the reaction equations.
The properties of hydrated ions differ from those of non-hydrated ions. For example, the unhydrated copper ion Cu 2+ is white in anhydrous crystals of copper (II) sulfate and has a blue color when hydrated, i.e., associated with water molecules Cu 2+ nH 2 O. Hydrated ions have both constant and variable number of water molecules.
Degree of electrolytic dissociation
In electrolyte solutions, along with ions, there are also molecules. Therefore, electrolyte solutions are characterized degree of dissociation, which is denoted by the Greek letter a (“alpha”).
This is the ratio of the number of particles decaying into ions (N g) to total number dissolved particles (N p).
The degree of electrolyte dissociation is determined experimentally and is expressed in fractions or percentages. If a = 0, then there is no dissociation, and if a = 1, or 100%, then the electrolyte completely disintegrates into ions. Different electrolytes have different degrees of dissociation, i.e. the degree of dissociation depends on the nature of the electrolyte. It also depends on the concentration: as the solution is diluted, the degree of dissociation increases.
According to the degree of electrolytic dissociation, electrolytes are divided into strong and weak.
Strong electrolytes- these are electrolytes that, when dissolved in water, almost completely dissociate into ions. For such electrolytes, the degree of dissociation tends to unity.
Strong electrolytes include:
1) all soluble salts;
2) strong acids, for example: H 2 SO 4, HCl, HNO 3;
3) all alkalis, for example: NaOH, KOH.
Weak electrolytes- these are electrolytes that, when dissolved in water, almost do not dissociate into ions. For such electrolytes, the degree of dissociation tends to zero.
Weak electrolytes include:
1) weak acids - H 2 S, H 2 CO 3, HNO 2;
2) aqueous solution of ammonia NH 3 H 2 O;
4) some salts.
Dissociation constant
In solutions of weak electrolytes, due to their incomplete dissociation, dynamic equilibrium between undissociated molecules and ions. For example, for acetic acid:
You can apply the law of mass action to this equilibrium and write down the expression for the equilibrium constant:
The equilibrium constant characterizing the process of dissociation of a weak electrolyte is called dissociation constant.
The dissociation constant characterizes the ability of an electrolyte (acid, base, water) dissociate into ions. The larger the constant, the easier the electrolyte breaks down into ions, therefore, the stronger it is. The values of dissociation constants for weak electrolytes are given in reference books.
Basic principles of the theory of electrolytic dissociation
1. When dissolved in water, electrolytes dissociate (break up) into positive and negative ions.
Ions is one of the forms of existence of a chemical element. For example, sodium metal atoms Na 0 vigorously interact with water, forming alkali (NaOH) and hydrogen H 2, while sodium ions Na + do not form such products. Chlorine Cl 2 has a yellow-green color and a pungent odor, and is poisonous, while chlorine ions Cl are colorless, non-toxic, and odorless.
Ions- these are positively or negatively charged particles into which atoms or groups of atoms of one or more are transformed chemical elements as a result of the donation or gain of electrons.
In solutions, ions move randomly in different directions.
According to their composition, ions are divided into simple- Cl - , Na + and complex- NH 4 + , SO 2 - .
2. The reason for the dissociation of an electrolyte in aqueous solutions is its hydration, i.e., the interaction of the electrolyte with water molecules and rupture chemical bond in it.
As a result of this interaction, hydrated ions are formed, i.e. associated with water molecules. Consequently, according to the presence of a water shell, ions are divided into hydrated(in solutions and crystalline hydrates) and unhydrated(in anhydrous salts).
3. Under the influence of an electric current, positively charged ions move to the negative pole of the current source - the cathode and are therefore called cations, and negatively charged ions move to the positive pole of the current source - the anode and are therefore called anions.
Consequently, there is another classification of ions - according to the sign of their charge.
The sum of the charges of cations (H +, Na +, NH 4 +, Cu 2+) is equal to the sum of the charges of anions (Cl -, OH -, SO 4 2-), as a result of which electrolyte solutions (HCl, (NH 4) 2 SO 4, NaOH, CuSO 4) remain electrically neutral.
4. Electrolytic dissociation is a reversible process for weak electrolytes.
Along with the dissociation process (decomposition of the electrolyte into ions), the reverse process also occurs - association(combination of ions). Therefore, in the equations of electrolytic dissociation, instead of the equal sign, the reversibility sign is used, for example:
5. Not all electrolytes dissociate into ions to the same extent.
Depends on the nature of the electrolyte and its concentration. The chemical properties of electrolyte solutions are determined by the properties of the ions that they form during dissociation.
The properties of weak electrolyte solutions are determined by the molecules and ions formed during the dissociation process, which are in dynamic equilibrium with each other.
The smell of acetic acid is due to the presence of CH 3 COOH molecules, the sour taste and color change of indicators are associated with the presence of H + ions in the solution.
The properties of solutions of strong electrolytes are determined by the properties of the ions that are formed during their dissociation.
For example, the general properties of acids, such as sour taste, changes in the color of indicators, etc., are due to the presence of hydrogen cations (more precisely, oxonium ions H 3 O +) in their solutions. General properties alkalis, such as soapiness to the touch, changes in the color of indicators, etc., are associated with the presence of hydroxide ions OH - in their solutions, and the properties of salts are associated with their decomposition in solution into metal (or ammonium) cations and anions of acid residues.
According to the theory of electrolytic dissociation all reactions in aqueous solutions of electrolytes are reactions between ions. This accounts for the high speed of many chemical reactions in electrolyte solutions.
Reactions occurring between ions are called ionic reactions, and the equations of these reactions are ionic equations.
Ion exchange reactions in aqueous solutions can occur:
1. Irreversible, to the end.
2. Reversible, that is, to flow simultaneously in two opposite directions. Exchange reactions between strong electrolytes in solutions proceed to completion or are practically irreversible when the ions combine with each other to form substances:
a) insoluble;
b) low dissociating (weak electrolytes);
c) gaseous.
Here are some examples of molecular and abbreviated ionic equations:
The reaction is irreversible, because one of its products is an insoluble substance.
The neutralization reaction is irreversible, because a low-dissociating substance is formed - water.
The reaction is irreversible, because CO 2 gas and a low-dissociating substance - water - are formed.
If among the starting substances and among the reaction products there are weak electrolytes or poorly soluble substances, then such reactions are reversible, that is, they do not proceed to completion.
In reversible reactions, the equilibrium shifts towards the formation of the least soluble or least dissociated substances.
For example:
The equilibrium shifts towards the formation of a weaker electrolyte - H 2 O. However, such a reaction will not proceed to completion: undissociated molecules of acetic acid and hydroxide ions remain in the solution.
If the starting substances are strong electrolytes, which upon interaction do not form insoluble or slightly dissociating substances or gases, then such reactions do not occur: when mixing solutions, a mixture of ions is formed.
Reference material for taking the test:
Periodic table
Solubility table
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Electrolytes are substances that dissociate (disintegrate) into ions in solution. Electrolyte solutions are capable of conducting electric current. For quantitative characteristics electrolytic dissociation, the concept of degree of dissociation was introduced.
The degree of dissociation is the ratio of the number of molecules dissociated into ions to the total number of molecules of the solute.
According to the degree of dissociation, all electrolytes are divided into strong and weak electrolytes. Strong electrolytes include compounds whose degree of dissociation is equal to unity, these are soluble salts, alkalis ( NaOH, KOH, LiOH, Ca(OH) 2 ), some acids ( HI, H 2 SO 4 , HCl, HBr, HNO 3 ). Weak electrolytes include substances whose degree of dissociation is significantly less than unity, such as water, insoluble and slightly soluble salts, insoluble bases, N.H. 4 OH, a series of acids ( CH 3 COOH, H 2 SO 3 , HNO 2 , H 2 S, HCN, H 3 P.O. 4 , H 2 CO 3 , H 2 SiO 3 , HF).
In solutions of weak electrolytes, the dissociation process is reversible, so the law of mass action can be applied to it. So in a solution of acetic acid the dissociation process is reversible:
CH 3 COOH CH 3 COO + H
the equilibrium constant will be equal to:
The equilibrium constant for the dissociation process is called dissociation constant, acidity constant, basicity constant, acid ionization constant, etc.
Polybasic acids undergo dissociation in stages, and each stage is characterized by its own dissociation constant:
H 2 CO 3 H + + HCO 3 −
HCO 3 − H + + CO 3 −
Water is also a weak electrolyte:
H 2 O H + + OH −
Table 1 shows the dissociation constants of a number of acids; for polybasic acids, the dissociation constants are presented in stages.
Table 1.
|
Name |
formula |
K d |
pK = -logK d |
|
Nitrogenous |
HNO2 |
6,9∙10 −4 |
3,16 |
|
Bornaya |
H3BO3 |
7.1∙10 −10 (K 1) |
9,15 |
|
Flint |
H2SiO3 |
1,3∙10 −10 |
9,9 |
|
Sulphurous |
H2SO3 |
1,4∙10 −2 |
1,85 |
|
Hydrogen sulfide |
H2S |
1,0∙10 −7 |
6,99 |
|
Coal |
H2CO3 |
4,5∙10 −7 |
6,35 |
|
Orthophosphoric |
H3PO4 |
7,1∙10 −3 5,0∙10 −13
|
2,15 12,0 |
|
Hydrogen cyanide |
HCN |
5,0∙10 −10 |
9,3 |
Drawing up equations for ion exchange reactions.
In electrolyte solutions, the direction of reactions is determined by the following rule: ionic reactions proceed towards the formation of poorly soluble substances, gases, weak electrolytes and complex ions; the reactions are practically irreversible. This rule is easily explained, because As a result of these reactions, one or more ions are removed from the reaction sphere, which, in accordance with La Chatelier's principle, leads to a more complete chemical reaction.
In such cases, it is recommended to write reaction equations in molecular-ionic form (molecular equation, abbreviated ionic equation), which allows you to better understand the essence of the process. In molecular ionic equations, soluble, strong electrolytes are written in the form of ions, and weak electrolytes and slightly soluble substances are written in the form of molecules.
Interaction of strong electrolytes with the formation of sediment:
Ba 2+ + 2Cl +2H + + SO 4 2 = BaSO 4 + 2H + + 2Cl (complete ionic equation)
Ba 2+ + SO 4 2- = BaSO 4 (shortened ionic equation)
2) The interaction of two strong electrolytes to form a weak electrolyte.
KCN + HCl = KCl + HCN
K + + CN + H + + Cl = K + + Cl + HCN
CN + H + = HCN
3) Interaction of a weak electrolyte with a strong one:
H 2 S + Pb(NO 3) 2 = PbS + 2HNO 3
H 2 S + Pb 2+ + 2NO 3 = PbS + 2H + + 2NO 3
H 2 S + Pb 2+ = PbS + 2H +
Interaction of sediment with acid:
CaCO 3 + 2H + + 2Cl = Ca 2+ + 2Cl + H 2 O + CO 2
CaCO 3 + 2H + = Ca 2+ + H 2 O + CO 2
Experimental part.
Experiment 1. Ion exchange reactions with the formation of precipitation.
Pour 5-6 drops of sodium phosphate into three test tubes and add 5-6 drops of cobalt nitrate to the first test tube, 5-6 drops of nickel sulfate to the second test tube, 5-6 drops of copper sulfate to the third test tube. Write reaction equations in molecular and ionic form.
Pour 5-6 drops of potassium dichromate into two test tubes, add 5-6 drops of barium chloride into the first, and 5-6 drops of bismuth nitrate into the second. Write the reaction equations in molecular and ionic form.
Based on the ionic equation, create a molecular equation and carry out the experiment:
Pb 2+ + 2I ─ = PbI 2
3Ca 2+ + 2PO 4 3– = Ca 3 (PO 4) 2
Based on the available reagents, obtain precipitates of copper, cobalt and nickel hydroxides. Write the reaction equations in molecular and ionic form.
Experiment 2. Ionic exchange reactions with the formation of a weak electrolyte.
Place some sodium acetate crystals in a test tube and add diluted sulfuric acid. Write the reaction equations in molecular and ion
Pour a few drops of ammonium chloride into the test tube and add potassium hydroxide, identify the gas released by smell (if there is no smell, the test tube can be slightly warmed). Write the reaction equation in molecular and ionic form.
Experiment 3. Shift of ionic equilibrium.
Pour 6-8 drops of ammonium hydroxide into two test tubes, add 2 drops of phenolphthalein. Then add 1 spatula of ammonium chloride to one of the test tubes and note the change in color intensity. Explain how the equilibrium in a solution shifts when ammonium chloride is added.
Pour 6-8 drops of acetic acid into two test tubes, add 2 drops of methyl orange, add 1 spatula of sodium acetate into one of the test tubes. Compare the color intensity in the test tubes. Note how the equilibrium in the solution shifts when salt is added.
Experiment 4. Dependence of the sequential precipitation of poorly soluble substances depending on their solubility product.
In one test tube you will get a precipitate of lead sulfate, in the second a precipitate of lead dichromate. Note the color of the precipitates formed. Add a few drops of potassium dichromate and sodium sulfate to the third test tube, mix the solution and add 2 drops of lead nitrate. Determine which substance precipitates first. Based on the product of the solubility of these salts, explain the sequence of their precipitation.