Ticket 1.
1. Basic chemical concepts (using the example of any chemical formula).
1. Complex substance - consists of different chemical elements.
2. 5 (coefficient) molecules of a complex substance.
3. Qualitative composition of a complex substance - consists of hydrogen and oxygen.
4. Quantitative composition of 1 molecule: 2 H atoms and one O atom; 5 molecules: 10 H atoms and 5 O atoms.
5. Molar mass M (H 2 O) = 1 * 2 + 16 = 18 g/mol
6. Mass of 5 molecules m (H 2 O) = 5 * 18 = 90 g
7. Mass fraction of hydrogen in the molecule: w = = = 0.3333 (33.33%)
2.
Elements of the oxygen subgroup - oxygen O, sulfur S, selenium Se, tellurium Te, polonium Ro- have a common name “chalcogens”, which means “giving birth to ores”.
Structure and properties of atoms.
Sulfur atoms, like oxygen atoms and all other elements of the main subgroup of group VI of D.I. Mendeleev’s Periodic Table, contain 6 electrons in the outer energy level, of which 2 are unpaired electrons.
Simple substances. Allotropy of oxygen is the simple substances oxygen O 2 and ozone O 3.
Sulfur, like oxygen, is characterized by allotropy. This is rhombic and plastic sulfur.
Chemical properties. Sulfur can be both an oxidizing agent and a reducing agent.
1. In relation to reducing agents - hydrogen, metals, sulfur exhibits oxidizing properties and acquires an oxidation state of -2. Under normal conditions, sulfur reacts with all alkali and alkaline earth metals, copper, mercury, silver, for example:
H 2 + S = H 2 S.
2. However, compared to oxygen and fluorine, sulfur is a reducing agent, forming compounds with an oxidation state of +4, +6.
Sulfur burns with a bluish flame, forming sulfur oxide (IV):
S + O 2 = SO 2.
This compound is commonly known as sulfur dioxide
3.
Ca + N 2 ®Ca 3 N 2
Cu + H 2 SO 4 (conc) ® CuSO 4 + SO 2 + H 2 O
Ticket 2.
1. Discovery by D.I. Mendeleev's Periodic Law. Periodic table of chemical elements.
D. I. Mendeleev arranged all the chemical elements known at the time of the discovery of the Periodic Law in a row, according to increasing atomic masses, and marked segments in it - periods , in which the properties of the elements and the substances formed by them changed in a similar way, namely (in modern terms):
1) metallic properties weakened;
2) non-metallic properties were enhanced;
3) the oxidation state of the element in higher oxides increased from +1 to +7;
4) oxides from basic through amphoteric were replaced by acidic ones;
5) hydroxides from alkalis through amphoteric hydroxides were replaced by increasingly stronger acids.
Based on these observations, D.I. Mendeleev made a conclusion in 1869 and formulated the Periodic Law:
properties of chemical elements and those formed by them substances are in a periodic depending on their atomic weights. In modern formulation atomic masses of elements replaced by nuclear charge.
2. Carbon subgroup: structure and properties of carbon atoms, simple substances formed by carbon, chemical properties of carbon.
Carbon subgroup (group 4 A) – carbon, silicon, germanium, tin, lead.
Carbon C is the first element of the main subgroup of group IV of D. I. Mendeleev’s Periodic Table. Its atoms contain 4 electrons in the outer energy level, so they can accept four electrons, acquiring an oxidation state of -4, i.e., exhibit oxidizing properties and give up their electrons to more electronegative elements, i.e., exhibit reducing properties, acquiring at This oxidation state is +4.
Carbon is a simple substance. Carbon forms allotropic modifications - diamond And graphite. They have a structure similar to graphite soot And charcoal. Coal, due to its porous surface, has the ability to absorb gases and dissolved substances. This property of some substances is called adsorption.
Chemical properties of carbon.
Diamond and graphite combine with oxygen at very high temperatures. Soot and coal interact with oxygen much more easily, burning in it. But in any case, the result of such interaction is the same - carbon dioxide is formed:
C + O 2 = CO 2
When heated, carbon forms carbides with metals, for example:
4Al + 3C = Al 4 C 3
3. Prove the presence of carbonate ion in sodium carbonate using a characteristic reaction.
CO 3 2- + H + (any acid) ® CO 2 +H 2 O
A heavy, colorless gas is released, which extinguishes the burning match.
Ticket 3.
1. Theory of atomic structure: planetary model of atomic structure, distribution of electrons across energy levels using the example of an element of the main and secondary subgroups.
Planetary model of the atom (Rutherford model)
 |
Nucleus: protons (p +) and neutrons (n 0).
The concept of the electron shell of an atom (electronic layers, energy levels)
In the electron shell, there are layers on which electrons with different amounts of energy will be located, which is why they are also called energy levels.
The number of these levels in an atom of a chemical element = the corresponding period number in D.I. Mendeleev’s table:
The Al atom, an element of period 3, has three levels. Each level can accommodate a certain maximum number of electrons: 1st - 2e - , 2nd - 8e - , and although the maximum number of electrons that can fit on the 3rd level is 18, atoms of elements of this period can place on it, like atoms of elements of period 2, only 8e - .
Energy levels containing the maximum number of electrons are called completed. If they contain fewer electrons, then these levels are incomplete.
Elements of side subgroups always have 2 electrons on the outer level (with the exception of Cr and Cu, they have 1 electron). Lastly, the pre-external level is filled:
2. Subgroup of halogens: structure and properties of atoms.
Elements of the main subgroup of group VII of the Periodic Table of D. I. Mendeleev, united under the common name halogens, fluorine F, chlorine Cl, bromine Br, iodine I, astatine At (rarely found in nature) are typical non-metals. This is understandable, because their atoms contain seven electrons in the outer energy level, and they only need one electron to complete it. Halogen atoms, when interacting with metals, accept an electron from the metal atoms. In this case, salts are formed. This is where the general name of the subgroup “halogens” comes from, i.e. “giving birth to salts”.
Halogens are very strong oxidizing agents. Fluorine in chemical reactions exhibits only oxidizing properties, and is characterized only by the oxidation state of -1 in compounds. The remaining halogens can also exhibit reducing properties when interacting with more electronegative elements - fluorine, oxygen, nitrogen. In this case, their oxidation states can take the values +1, +3, +5,
7. The reducing properties of halogens increase from chlorine to iodine, which is associated with an increase in the radii of their atoms: chlorine atoms are approximately one and a half times smaller than those of iodine.
Halogens are simple substances. All halogens exist in a free state in the form of diatomic molecules F 2, Cl 2, Br 2, I 2. Fluorine and chlorine are gases, bromine is a liquid, iodine is a solid. From F 2 to I 2 the color intensity of the halogens increases. Iodine crystals have a metallic sheen.
3. Prove the presence of sulfate ion in sodium sulfate using a characteristic reaction.
SO 4 2- + Ba 2+ (soluble barium salt) ® BaSO 4 ¯
White fine crystalline precipitate
Ticket 4.
1. Rules for determining oxidation states.
Elements that have a constant oxidation state:
1. Group I A: Li +, Na +, K +, Rb +, Cs +.
2. II group A: Be +2, Mg +2, Ca +2, Zn +2, Sr +2, Cd +2, Ba +2.
3. III group A: Al +3
6. H +1 (MeH -1)
7. In simple substances, s.o. = 0.
For the remaining elements, s.o. consider
H 2 +1 S X O 4 - 2 : so sulfur does not have a constant s.o., so we take it as X.
+1 *2 + X + (-2 ) * 4 = 0
Higher s.o. = Group No. (except O, F)
Lowest s.o. = Group No. – 8 (Me does not have a lower s.o.)
2. Chemical properties of halogens - simple substances.
The chemical activity of halogens, like non-metals, weakens from fluorine to iodine.
Each halogen is the strongest oxidizing agent in its period. The oxidizing properties of halogens are distinct when they interact with metals. In this case, salts are formed. Thus, fluorine already reacts under normal conditions with most metals, and when heated, it also reacts with gold, silver, and platinum, which are known for their chemical passivity. Aluminum and zinc ignite in a fluorine atmosphere:
0 0 +2 -1
Zn + F 2 = ZnF 2.
The remaining halogens react with metals mainly when heated.
The decrease in the oxidative properties and the increase in the reducing properties of halogens from fluorine to iodine can also be judged by their ability to displace each other from salt solutions.
Thus, chlorine displaces bromine and iodine from solutions of their salts, for example:
Cl 2 + 2NaBr = 2NaCl + Br 2.
3. Make up molecular and ionic equations for reactions between substances: lead (II) nitrate and potassium sulfate, iron (III) chloride and silver nitrate.
Ticket 5.
1. Classification of chemical reactions according to the number of starting substances and reaction products.
2. Hydrogen halides and hydrohalic acids and their salts.
N 2 + G 2 = 2NG
(G is the conventional chemical designation for halogens).
All hydrogen halides (their general formula can be written as NG) are colorless gases with a pungent odor and are toxic. They dissolve very well in water and smoke in humid air, as they attract water vapor in the air, forming a foggy cloud.
Solutions of hydrogen halides in water are acids, these are HF - hydrofluoric, or hydrofluoric, acid, HC1 - hydrochloric, or hydrochloric acid, HBr - hydrobromic acid, HI - hydroiodic acid. The strongest of the hydrohalic acids is hydroiodic acid, and the weakest is hydrofluoric acid.
Salts of hydrohalic acids. Hydrohalic acids form salts: fluorides, chlorides, bromides and iodides. Chlorides, bromides and iodides of many metals are highly soluble in water.
To determine chloride, bromide and iodide ions in solution and distinguish them, a reaction with silver nitrate is used.
3. Calculate the mass fraction of oxygen in sodium sulfate.
| Given: Na 2 SO 4 | Solution: W O = =  = W O = 0.451 =45.1% |
| W O - ? % |
Answer: mass fraction of oxygen 45.1%.
Ticket 6.
1. Electrolytes and non-electrolytes.
According to the conductivity of electric current, all substances are divided into electrolytes and non-electrolytes.
Electrolytes are substances whose solutions conduct electric current. These include acids, bases, and salts. These substances conduct current because can dissociate into a cation and anion:
Acids: HAnH + + An -
Bases: MON M + + OH -
Salts: МAn→ М + + An -
The index after a simple ion or parenthesis becomes a coefficient
Ca 3 (PO 4) 2 → 3Ca 2+ + 2 (PO 4) 3-
Non-electrolytes include all others - simple substances, oxides, almost all organic substances.
2.
The physical properties of metals are determined by their structure: the presence of free electrons in the crystal lattice. Thanks to free electrons, all metals have electrical conductivity, thermal conductivity, and a metallic luster.
Electro- And thermal conductivity. Electrons moving chaotically in a metal under the influence of an applied electrical voltage acquire directional movement, resulting in the generation of an electric current. Silver, copper, as well as gold, aluminum, and iron have the highest electrical conductivity; the smallest - manganese, lead, mercury.
Most often, the thermal conductivity of metals changes in the same sequence as electrical conductivity. It is due to the high mobility of free electrons, which, colliding with vibrating ions and atoms, exchange energy with them. Therefore, the temperature quickly equalizes throughout the entire piece of metal.
Metallic shine. The electrons that fill the interatomic space reflect light rays rather than transmit them like glass, which is why all metals in the crystalline state have a metallic luster.
The remaining properties - hardness, density, fusibility, plasticity - are different.
3. Describe one of the elements - metals (sodium, calcium, aluminum or iron) (all optional).
CHARACTERISTICS OF A METAL ELEMENT USING THE EXAMPLE OF ALUMINA
1. Position in the Periodic Table.Aluminum(serial number 13 ) is an element 3 period, main subgroups 3
2. Number of protons in an atom aluminum equals 13 , number of electrons - 13 , number of neutrons in the isotope 27 13 Al - 27-13 =14, nuclear charge +13 , electron level distribution 2, 8, 3 .
3. Simple substance.Aluminum- This amphoteric metal. Atoms aluminum show restorative properties.
4. Higher oxide, its character. Aluminum forms a higher oxide, the formula of which is Al2O3. According to its properties it is amphoteric oxide.
4. Higher hydroxide, its character. Aluminum forms a higher hydroxide, the formula of which is Al(OH)3. By properties amphoteric base.
Ticket 7.
1. The concept of strong and weak electrolytes.
Electrolytes include salts, acids, and bases.
Salts are all strong electrolytes, i.e. conduct electricity well. Therefore, in the dissociation equation they put only one arrow in the direction of disintegration into ions
МAn→ М + + An -
Strong bases are alkalis, i.e. water-soluble bases.
Ca(OH) 2 → Ca 2+ +2(OH) -
Insoluble and slightly soluble are weak, therefore, when writing the dissociation equation, they put a reversibility sign (in addition to ions, there are molecules)
MON M + + OH -
Strong acids include HCl, HBr, HI, H2SO4, HNO3, HClO4, HClO3.
2. Alloys.
These are materials with characteristic properties, consisting of two or more components, at least one of which is metal.
In metallurgy, iron and all its alloys are divided into one group called black metals; other metals and their alloys have a technical name non-ferrous metals.
The vast majority of iron (or ferrous) alloys contain carbon. They are divided into cast iron and steel.
Cast iron- an iron-based alloy containing more than 2% carbon, as well as manganese, silicon, phosphorus and sulfur. Cast iron is much harder than iron, it is usually very brittle, cannot be forged, and breaks when hit. This alloy is used for the manufacture of various massive parts by casting, the so-called cast iron, and for processing into steel - pig iron.
Depending on the state of carbon in the alloy, gray and white cast iron are distinguished.
Steel is an iron-based alloy containing less than 2% carbon. Based on their chemical composition, steel is divided into two main types: carbon And alloyed.
Examples of non-ferrous alloys can be: nichrome, tertiary solder, pobedit, duralumin.
Duralumin- an alloy of aluminum (95%), magnesium, copper and manganese. Very light and durable alloy. It is equal in strength to steel, but three times lighter. Used in aircraft construction.
3. Describe one of the elements - non-metals (chlorine, sulfur, phosphorus, nitrogen, carbon, silicon) (all optional).
CHARACTERISTICS OF A NON-METAL ELEMENT USING THE EXAMPLE OF SULPHUR
1. Position in the Periodic TableSulfur(serial number 16 ) is an element 3 period, main subgroups 6 groups of the Periodic Table.
2.The structure of the atom, its properties. The number of protons in a sulfur atom is 16 , number of electrons - 16 , number of neutrons in the isotope 32 16 S - 32-16 =16, nuclear charge +16 , distribution of electrons across levels 2, 8, 6.
3. Simple substance. Sulfur is non-metal. Sulfur atoms exhibit oxidative properties.
3.Higher oxide, its character. Sulfur forms a higher oxide, the formula of which is SO 3. According to its properties it is acid oxide.
4.Higher hydroxide, its character. Sulfur forms a higher hydroxide, the formula of which is H2SO4. By properties acid.
Ticket 8.
1. Oxides: their composition, classification and names.
Oxides- these are binary compounds, in second place is oxygen with an oxidation state of -2.
Depending on which element comes first, oxides are divided into three groups:
1) Basic. These are oxides in which the metal comes first: CaO, Na 2 O.
2) Acidic. These are oxides in which a non-metal comes first: P 2 O 5.
3) Amphoteric. These are oxides in which the first element is an amphoteric element (transition metal): Al 2 O 3, Fe 2 O 3
Basic oxides correspond to bases. For example, Na 2 O - NaOH. Acid oxides correspond to acids: P 2 O 5 - H 3 PO 4.
The names are composed of the name of oxygen (in Latin) - oxide, and the name of the first element indicating the oxidation state (if variable)
P 2 +5 O 5 phosphorus (V) oxide, Fe 2 +3 O 3 iron (III) oxide
2. Oxygen subgroup: structure and properties of atoms, simple substances, chemical properties of sulfur.
For the answer, see ticket 1, question 2.
3. Prove the presence of chloride ion in potassium chloride using a characteristic reaction.
Cl - + Ag + (soluble silver salt) ® Ag Cl ¯
White curdled sediment
Ticket 9.
1. Acids. Names and formulas of acids.
Acids are complex inorganic substances consisting of hydrogen cation and an acid residue anion.
HCl – hydrochloric
HNO 3 – nitrogen
H 2 SO 4 – sulfuric
H 2 CO 3 – coal
H 3 PO 4 – phosphoric
2. Alloys.
For the answer, see ticket 7, question 2.
3. Describe one of the elements - metals (lithium, magnesium, potassium or aluminum) (all optional).
For a sample answer, see ticket 6, question 3.
Ticket 10.
1. The position of metals in the periodic table of chemical elements D.I. Mendeleev, the structure of their atoms and crystals.
Me are simple substances that easily give up electrons. For the main subgroups:
Me includes all elements of secondary subgroups. This position of Me in the periodic table is associated with their structure: a small number of electrons on the outer level (1-3), which for the main subgroups is determined by the group number, and for the side ones - always 2 electrons. The second characteristic for Me is a large radius (increases in the table from top to bottom).
In the crystal lattice, Me has free electrons, which are responsible for the main physical properties of Me:
2. Foundations in the light of TED; their classification and chemistry. properties.
Bases are electrolytes that, when dissociated, form a metal cation and an acidic anion.
Classification:
1. Water-insoluble bases.
2. Alkalis – soluble in water.
Typical base reactions
1 . Base + acid® salt + water.
(exchange reaction)
Hl + NaOH = NaCl + H 2 O
H + + OH - = H 2 O (neutralization reaction).
2. Base + acid oxide®salt + water.
(exchange reaction)
2NaOH + N2O5 = 2NaNO3 + H2O
2OH - + N 2 O 5 = 2NO 3 - + H 2 O;
3 . Lye + salt ® new base + new salt.
(exchange reaction)
2KOH + CuSO 4 = = Cu(OH) 2 ¯+ K 2 SO 4
Cu 2+ + 2OH - = = Cu(OH) 2 ¯
4. Bases insoluble in water decompose when heated into metal oxide and water, which is not typical for alkalis, for example:
Cu(OH) 2 ¯ = CuO + H 2 O
3. Arrange the coefficients in reaction schemes using the electronic balance method. Indicate the oxidizing agent and reducing agent, oxidation and reduction processes.
Al + O 2 ® Al 2 O 3
HNO 3 + P® H 3 PO 4 + NO 2 + H 2 O
When preparing for the exam, see the solution in the laboratory journal - practical work No. 2.
Ticket 11.
1. Electronic balance method.
Al 0+ O2 0 ® Al 2 +3 O 3 -2
We write down the elements that changed the s.o.
Al 0 – 3e - → Al +3 4 Al 0 – reducing agent, oxidation process
O 2 0 +2*2e - →2O -2 3 O 2 0 – oxidizing agent, reduction process
Note. If a simple substance has an index (2), then it is transferred to the electronic balance.
We equalize the reaction using the coefficients from the electronic balance (4, 3):
4Al +3O 2 ® 2 Al 2 O 3
2. General chemical properties of metals. Electrochemical voltage series of metals and the interaction of metals with solutions of acids and salts.
Metals are reducing agents. Reductive properties are exhibited in reactions with simple and complex substances.
I. With simple – non-metals
2Na + S = Na 2 S sodium sulfide
II. With complex: water, acids, salt solutions (substitution reactions). When writing all these reactions, it is necessary to take into account the activity series (electrochemical series) of metals.
K, Ca, Na, Mg, Al, Zn, Fe, Ni, Sn, Pb, (H 2), Cu, Hg, Ag, Au.
1. Metals standing in the voltage series to the left of hydrogen displace it from acid solutions, and those standing to the right, as a rule, do not displace hydrogen from acid solutions:
Zn + 2HCl = ZnCl 2 + H 2.
2. Each metal displaces from salt solutions other metals located to the right of it in the stress series, and can itself be displaced by metals located to the left, for example:
Fe + CuSO 4 = FeSO 4 + Cu,
Сu + HgCl 2 = Hg + CuCl 2.
3. Determine the mass of carbon monoxide (IV) by the amount of substance 2mmol.
Answer: 88 mg carbon monoxide (IV).
Ticket 12.
1. Hydrolysis of salts by cation.
МAn + HOH = MOH + HАn
Salt base acid
A salt undergoes hydrolysis if it is formed by at least one weak ion. If the cation is weak (from a weak base), then hydrolysis is called according to the cation.
Weak bases are insoluble in water.
For example, FeCl 3 is a salt formed by a strong acid (HCl) and a weak base (Fe(OH) 3)
FeCl3Û Fe 3+ +3Cl -
weak cation
Fe 3+ + H + OH - Û Fe OH 2+ + H+
4. Determine whether the solution is acidic
This is the case hydrolysis by cation.
2. General physical properties of metals.
See the ticket for the answer. 6 , question 2.
3. Carry out reactions confirming that sulfuric acid contains hydrogen cations and sulfate anions.
H 2 SO 4 Û 2H + + SO 4 2-
H+ - methyl orange (will turn red), or litmus (will turn red)
SO 4 2- + Ba 2+ ® Ba SO 4 ¯ (white fine-crystalline precipitate)
Ticket 13.
1. Hydrolysis of salts by anion.
Salt hydrolysis is the interaction of a soluble salt with water.
МAn + HOH = MOH + HАn
Salt base acid
A salt undergoes hydrolysis if it is formed by at least one weak ion. If the anion is weak (from a weak acid), then hydrolysis is called according to the anion.
Strong acids: H 2 SO 4, HNO 3, HClO 3, HClO 4, HCl, HBr, HI
The rest are weak.
For example, Na 2 CO 3 - a salt is formed by a weak acid and a strong base
1. Write down the salt dissociation equation. Na 2 CO 3Û 2Na + + CO 3 2-
weak anion
2. Select a weak ion: cation or anion.
3. Record its interaction with water. CO 3 2- + H + OH - Û HCO 3 - + HE -
4. Determine the solution environment: HE -- alkaline environment, H + - acidic environment, absence of H + and OH - neutral.
This is the case hydrolysis at the anion.
2. General chemical properties of metals.
For the answer, see ticket 11, question 2.
3. How many grams of iodine and alcohol do you need to take to prepare 30 g of a 5% solution of iodine tincture?
When preparing for the exam, see the solution in the laboratory journal - practical work No. 1.
Ticket 14.
1 . Drawing up formulas of chemical substances according to the degree of oxidation.
1. Enter the oxidation states:
For the first element, the constant is the highest (by group number), or variable (indicated in the name of the substance)
For the second - the lowest (-(8-No. gr.)), or according to the solubility table (for a group of elements);
2. Cross the oxidation states to get the indices (reduce if necessary).
For example.
1) make aluminum oxide: Al 2 +3 O 3 -2
2) compose lead(IV) sulfide: Pb 2 +4 S 4 -2 → PbS 2
3) make calcium sulfate: Ca +2 SO 4 -2
2. Subgroup of halogens.
When preparing for the exam, see the answer in ticket 3, question 2.
3. Carry out reactions to confirm the qualitative composition of barium chloride.
BaCl 2 Û Ba 2+ + 2Cl -
Ba 2+ + SO 4 2- ® Ba SO 4 ¯ (white fine-crystalline precipitate)
Сl - + Ag + ® Ag Сl ¯ (white cheesy sediment)
Ticket 15.
1. Ion exchange reactions.
In order to record an ion exchange reaction, you must adhere to the following algorithm.
1. Write a molecular equation for the reaction
Fe(NO 3) 3 + 3NaOH = Fe(OH) 3 + 3NaNO 3
2. Check the possibility of the reaction occurring (reaction products: sediment, gas or water)
Fe(NO 3) 3 + 3NaOH = Fe(OH) 3↓ + 3NaNO 3
3. Write down the ionic equation of the reaction, and do not forget:
· We leave it in the form of a molecule - a weak electrolyte (H 2 O) and a non-electrolyte, sediment or gas;
· The coefficient in front of the formula of a substance refers to both ions!!!
· The formulas of polyatomic (complex) ions do not break: OH -, CO3 2-, PO4 3-, etc.
· The index after a simple ion or bracket goes into the coefficient in front of it in the ionic equation
Fe 3+ + 3(NO 3) - + 3Na + + 3OH - = Fe(OH) 3↓ + 3Na + + NO 3 -
4. “Reduce” similar ones
Fe 3+ + 3NO 3 - + 3Na++ 3OH - = Fe(OH) 3↓ + 3Na+ + NO 3 -
5. Rewrite the abbreviated ionic equation
Fe 3+ + 3OH - = Fe(OH) 3↓
2. General characteristics of alkali metals: atomic structure and physical properties of simple substances.
Chemical properties of medium salts
Interaction of medium salts with metals
The reaction of a salt with a metal occurs if the initial free metal is more active than the one that is part of the original salt. You can find out which metal is more active by using the electrochemical series of metal voltages.
For example, iron interacts with copper sulfate in an aqueous solution, since it is more active than copper (to the left in the activity series):
At the same time, iron does not react with a solution of zinc chloride, since it is less active than zinc:
It should be noted that such active metals as alkali and alkaline earth metals, when added to aqueous solutions of salts, will primarily react not with the salt, but with the water included in the solutions.
Interaction of medium salts with metal hydroxides
Let us make a reservation that in this case metal hydroxides mean compounds of the type Me(OH) x.
In order for the middle salt to react with the metal hydroxide, it must simultaneously (!) two requirements must be met:
- sediment or gas must be detected in the intended products;
- the original salt and the original metal hydroxide must be soluble.
Let's look at a couple of cases in order to understand this rule.
Let's determine which of the reactions below occur and write the equations for the reactions that occur:
- 1) PbS + KOH
- 2) FeCl 3 + NaOH
Consider the first interaction of lead sulfide and potassium hydroxide. Let's write down the supposed ion exchange reaction and mark it on the left and right with “curtains”, indicating in such a way that it is not yet known whether the reaction actually occurs:
In the supposed products we see lead (II) hydroxide, which, judging by the solubility table, is insoluble and should precipitate. However, the conclusion that the reaction is proceeding cannot yet be made, since we have not checked the satisfaction of another mandatory requirement - the solubility of the original salt and hydroxide. Lead sulfide is an insoluble salt, which means the reaction does not proceed, since one of the mandatory requirements for the reaction between the salt and the metal hydroxide to occur is not met. Those.:
Let's consider the second proposed interaction between iron(III) chloride and potassium hydroxide. Let's write down the expected ion exchange reaction and mark it on the left and right with “curtains”, as in the first case:
In the supposed products we see iron (III) hydroxide, which is insoluble and must precipitate. However, it is not yet possible to draw a conclusion about the course of the reaction. To do this, you must also ensure the solubility of the original salt and hydroxide. Both starting materials are soluble, which means we can conclude that the reaction is proceeding. Let's write down its equation:
Reactions of medium salts with acids
A medium salt reacts with an acid when a precipitate or weak acid is formed.
It is almost always possible to recognize a precipitate among the expected products using the solubility table. For example, sulfuric acid reacts with barium nitrate, since insoluble barium sulfate precipitates:
It is impossible to recognize a weak acid from the solubility table, since many weak acids are soluble in water. Therefore, the list of weak acids should be memorized. Weak acids include H 2 S, H 2 CO 3, H 2 SO 3, HF, HNO 2, H 2 SiO 3 and all organic acids.
For example, hydrochloric acid reacts with sodium acetate to form a weak organic acid (acetic acid):
It should be noted that hydrogen sulfide H2S is not only a weak acid, but is also poorly soluble in water, and therefore is released from it in the form of a gas (with the smell of rotten eggs):
In addition, you should definitely remember that weak acids - carbonic and sulfurous - are unstable and almost immediately after formation they decompose into the corresponding acid oxide and water:
It was said above that the reaction of a salt with an acid occurs if a precipitate or weak acid is formed. Those. if there is no precipitate and a strong acid is present in the intended products, then the reaction will not proceed. However, there is a case that does not formally fall under this rule, when concentrated sulfuric acid displaces hydrogen chloride when acting on solid chlorides:
However, if you take not concentrated sulfuric acid and solid sodium chloride, but solutions of these substances, then the reaction really will not work:
Reactions of medium salts with other medium salts
The reaction between intermediate salts occurs if simultaneously (!) two requirements are met:
- the original salts are soluble;
- the expected products contain sediment or gas.
For example, barium sulfate does not react with potassium carbonate because, although the intended products contain a precipitate (barium carbonate), the solubility requirement for the original salts is not met.
At the same time, barium chloride reacts with potassium carbonate in solution, since both original salts are soluble, and there is a precipitate in the products:
A gas is formed during the interaction of salts in the only case - if a solution of any nitrite is mixed with a solution of any ammonium salt when heated:
The reason for the formation of gas (nitrogen) is that the solution simultaneously contains NH 4 + cations and NO 2 - anions, forming thermally unstable ammonium nitrite, which decomposes in accordance with the equation:
Reactions of thermal decomposition of salts
Carbonate decomposition
All insoluble carbonates, as well as lithium and ammonium carbonates, are thermally unstable and decompose when heated. Metal carbonates decompose to metal oxide and carbon dioxide:
and ammonium carbonate produces three products - ammonia, carbon dioxide and water:
Nitrate decomposition
Absolutely all nitrates decompose when heated, and the type of decomposition depends on the position of the metal in the activity series. The decomposition diagram of metal nitrates is presented in the following illustration:
So, for example, in accordance with this scheme, the decomposition equations for sodium nitrate, aluminum nitrate and mercury nitrate are written as follows:
It should also be noted the specificity of the decomposition of ammonium nitrate:
Decomposition of ammonium salts
The thermal decomposition of ammonium salts is most often accompanied by the formation of ammonia:
If the acid residue has oxidizing properties, instead of ammonia, some product of its oxidation is formed, for example, molecular nitrogen N2 or nitric oxide (I):
Chemical properties of acid salts
The ratio of acid salts to alkalis and acids
Acidic salts react with alkalis. Moreover, if the alkali contains the same metal as the acid salt, then medium salts are formed:
Also, if in the acidic residue of an acid salt there are two or more mobile hydrogen atoms left, as, for example, in sodium dihydrogen phosphate, then the formation of both an average is possible:
and another acidic salt with a smaller number of hydrogen atoms in the acid residue:
It is important to note that acid salts react with any alkalis, including those formed by another metal. For example:
Acid salts formed by weak acids react with strong acids in a similar way to the corresponding medium salts:
Thermal decomposition of acid salts
All acidic salts decompose when heated. As part of the Unified State Exam program in chemistry, you should learn from the decomposition reactions of acid salts how bicarbonates decompose. Metal bicarbonates decompose already at temperatures above 60 o C. In this case, metal carbonate, carbon dioxide and water are formed:
The last two reactions are the main cause of scale formation on the surface of water heating elements in electric kettles, washing machines, etc.
Ammonium bicarbonate decomposes without a solid residue to form two gases and water vapor:
Chemical properties of basic salts
Basic salts always react with all strong acids. In this case, intermediate salts can be formed if an acid with the same acidic residue as in the main salt was used, or mixed salts if the acidic residue in the basic salt differs from the acidic residue of the acid reacting with it:
Also, basic salts are characterized by decomposition reactions when heated, for example:
Chemical properties of complex salts (using the example of aluminum and zinc compounds)
As part of the Unified State Examination program in chemistry, one should learn the chemical properties of such complex compounds of aluminum and zinc as tetrahydroxoaluminates and tetrahydroxoaluminates.
Tetrahydroxoaluminates and tetrahydroxozincates are salts whose anions have the formulas - and 2-, respectively. Let's consider the chemical properties of such compounds using sodium salts as an example:
These compounds, like other soluble complex compounds, dissociate well, while almost all complex ions (in square brackets) remain intact and do not dissociate further:
The action of an excess of strong acid on these compounds leads to the formation of two salts:
When they are exposed to a lack of strong acids, only the active metal passes into the new salt. Aluminum and zinc in the hydroxides precipitate:
Precipitation of aluminum and zinc hydroxides with strong acids is not a good choice, since it is difficult to add the strictly required amount of strong acid without dissolving part of the precipitate. For this reason, carbon dioxide is used, which has very weak acidic properties and therefore is not able to dissolve the hydroxide precipitate:
In the case of tetrahydroxoaluminate, hydroxide precipitation can also be carried out using sulfur dioxide and hydrogen sulfide:
In the case of tetrahydroxozincate, precipitation with hydrogen sulfide is impossible, since zinc sulfide precipitates instead of zinc hydroxide:
When solutions of tetrahydroxozincate and tetrahydroxoaluminate are evaporated, followed by calcination, these compounds transform into zincate and aluminate, respectively.
SOME REFERENCES ON CHEMISTRY
Basic characteristics of elementary particles
|
Particle and its designation |
Weight |
Charge |
Note |
|
Proton - p+ |
The number of protons is equal to the atomic number of the element |
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|
Neutron - n 0 |
The number of neutrons is found by the formula: N=A-Z |
||
|
Electron - e |
1:1837 |
The number of electrons is equal to the atomic number of the element. |
The maximum (largest) number of electrons located at an energy level can be determined by the formula: 2n 2 , where n is the level number.
Simple substances
|
Metals |
Nonmetals |
|
1.Solids(except mercury - Hg) |
1. Solid(sulfur - S, red phosphorus and white phosphorus - P4, iodine - I2, diamond and graphite - C), gaseous substances(Oxygen - O2, ozone - O3, nitrogen - N2, hydrogen - H2, chlorine - Cl2, fluorine - F2, noble gases) and liquid (bromine - Br2) |
|
2. Have a metallic sheen. |
2. They do not have a metallic luster (exceptions are iodine-I2, graphite-C). |
|
3. Electrically and thermally conductive |
3. Most do not conduct electricity (conductors are, for example, silicon, graphite) |
|
4. Malleable, plastic, viscous |
4. In a solid state - brittle |
Changes in indicator color depending on the environment
|
Indicator name |
Indicator color |
||
|
in a neutral environment |
in an alkaline environment |
in an acidic environment |
|
|
Litmus |
Purple |
Blue |
Red |
|
Methyl orange |
Orange |
Yellow |
Red-pink |
|
Phenolphthalein |
Colorless |
Raspberry |
Colorless |
When dissolved sulfuric acid need to pour it in a thin stream into the water and stir.
Nomenclature of salts
|
Acid name (formula) |
Name of salts |
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Nitrogenous (HNO2) |
Nitrites |
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Nitrogen (HNO3) |
Nitrates |
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Hydrochloric (hydrochloric) HCl |
Chlorides |
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Sulphurous (H2SO3) |
Sulfites |
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Sulfuric (H2SO4) |
Sulfates |
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Hydrogen sulfide (H2S) |
Sulfides |
|
Phosphoric (H3PO4) |
Phosphates |
|
Coal (H2CO3) |
Carbonates |
|
Silicon (H2SiO3) |
Silicates |
Calcium carbonate CaCO3 is a water-insoluble salt from which marine animals (molluscs, crayfish, protozoa) build the coverings of their bodies - shells; calcium phosphate Ca3(PO4)2 is a water-insoluble salt, the basis of phosphorite and apatite minerals.
Substances with atomic crystal lattice: crystal thief, silicon and germanium, as well as complex substances, for example those containing silicon oxide (IV) - SiO2: silica, quartz, sand, rock crystal.
Molecular crystal lattice: HCl, H2O - polar bonds; N2, O3 - non-polar bonds; solid water-ice, solid carbon monoxide (IV) - “dry ice”, solid hydrogen chloride and hydrogen sulfide, solid simple substances formed by one- (noble gases), two- (H2, O2, Cl2, I2), three- (O3 ), four- (P4), eight-atomic (S8) molecules.
Chemical analysis - determination of the composition of mixtures.
Particularly pure substances- substances in which the content of impurities affecting their specific properties does not exceed one hundred thousandth or even one millionth of a percent.
The relationship between some physical and chemical quantities and their units
|
Unit |
Weight(m) |
Amount of substance (n) |
Molar mass (M) |
Volume (V) |
Molar volume (V) |
Number of particles (N) |
|
Most often used in the study of chemistry |
mole |
g/mol |
l/mol |
Avogadro's number N= 6x10 23 |
||
|
1000 times bigger |
kg |
kmol |
kg/kmol |
m 3 |
m 3 /kmol |
6x10 26 |
|
1000 times smaller |
mg |
mmol |
mg/mmol |
ml |
ml/mmol |
6x10 20 |
Classification of acids
|
Signs of classification |
Acid groups |
|
Presence of oxygen in the acid residue |
A) oxygen: phosphorus, nitrogen B) oxygen-free: hydrogen sulfide, chlorine, hydrogen bromide |
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Basicity |
A) monobasic: chlorine, nitrogen B) dibasic: sulfur, coal, hydrogen sulfide B) tribasic: phosphoric |
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Solubility in water |
A) soluble: sulfuric, nitrogenous, hydrogen sulphide B) insoluble: silicon |
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Volatility |
A) volatiles: chlorine, nitrogen, hydrogen sulfide B) non-volatile: sulfur, silicon, phosphorus |
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Degree of electrolytic dissociation |
A) strong: sulfuric, chloric, nitrogenous B) weak: hydrogen sulfide, sulfur, coal |
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Stability |
A) stable: sulfuric, phosphoric, chloric B) unstable: sulfur, coal, silicon |
Typical acid reactions
1. Acid + base = salt + water (exchange reaction)
2. Acid + metal oxide = salt + water (exchange reaction)
3. Acid + metal = salt + hydrogen (substitution reaction)
4. Acid + salt = new acid + new salt (exchange reaction)
Classification of bases
|
Signs of classification |
Base groups |
|
Solubility in water |
A) soluble (alkalis): sodium hydroxide, potassium hydroxide, calcium hydroxide, barium hydroxide B) insoluble bases: copper (II) hydroxide, iron (II) hydroxide, iron (III) hydroxide |
|
Acidity (number of hydroxo groups) |
A) monoacid: sodium hydroxide (caustic soda), potassium hydroxide (caustic potash) B) diacid: iron(II) hydroxide, copper(II) hydroxide |
Typical base reactions
1. Base + acid = salt + water (exchange reaction)
2. Base + non-metal oxide = salt + water (exchange reaction)
3. Alkali + salt = new base + new salt (exchange reaction)
Insoluble bases decompose when heated into metal oxide and water, which is not typical for alkalis, for example: Fe(OH)2 = FeO + water
Typical reactions of basic oxides
1. Basic oxide + acid = salt + water (exchange reaction)
2. Basic oxide + acidic oxide = salt (compound reaction)
3. Basic oxide + water = alkali (compound reaction). This reaction occurs if a soluble base, an alkali, is formed. For example, CuO + water - the reaction does not occur, because Copper(II) hydroxide is an insoluble base.
Typical acid oxide reactions
1. Acid oxide + base = salt + water (exchange reaction)
2. Acidic oxide + basic oxide = salt (compound reaction)
3. Acidic oxide + water = acid (compound reaction). This reaction is possible if the acid oxide is soluble in water. For example: silicon (IV) oxide practically does not interact with water.
Typical salt reactions
1. Salt + acid = another salt + another acid (exchange reaction)
2. Salt + alkali = another salt + another base (exchange reaction)
3. Salt1 + salt2 = salt3 + salt 4 (exchange reaction: two salts react, resulting in two other salts)
4. Salt + metal = another salt + another metal (substitution reaction), you need to see the position of the metal in the electrochemical voltage series of metals.
Rules for a range of metal stresses
1. Metals that are located to the left of hydrogen interact with acid solutions. This extends to the ability of metals to displace other metals from salt solutions. For example, copper can be replaced from solutions of its salts by metals such as magnesium, aluminum, zinc and other metals. But copper is not replaced by mercury, silver, and gold, because These metals are located to the right in the voltage series than copper. But copper displaces them from salt solutions.
The first rule of the series of stresses of metals about the interaction of metals with solutions of acids does not apply to concentrated sulfuric acid and nitric acid of any concentration: these acids interact with metals in the series of stresses both before and after hydrogen, being reduced to sulfur oxide (IV ), NO, etc. For example, when dilute nitric acid reacts with copper, it produces copper(II) nitrate, nitric oxide (II) and water.
2. Each metal displaces other metals located to the right of it in the stress series from salt solutions. This rule is observed if the following conditions are met:
Both salts (before and after the reaction - reacting and formed) must be soluble;
Metals should not interact with water, therefore the metals of the main subgroups of groups I and II (for the latter, starting with calcium) do not displace other metals from salt solutions.
Redox reactions
Reductant - atoms, ions, molecules, giving electrons.
The most important reducing agents: metals; hydrogen; coal; carbon monoxide (II) CO; hydrogen sulfide; ammonia; hydrochloric acid, etc.
The process of giving up electrons by atoms, ions and molecules is oxidation.
Oxidizing agent - atoms, ions, molecules, hosts electrons.
The most important oxidizing agents: halogens; nitric and sulfuric acids; potassium permanganate, etc.
The process of adding electrons by atoms, ions and molecules is reduction.
“Determination of ammonium salts” - Obtaining ammonium. Properties of ammonium salts. Ammonium salts. Relation to heating. Ability to decompose. Ammonium. Chemical properties of ammonium salts. Physical properties of ammonium salts. Application of ammonium salts. Application of ammonium salts in agriculture. Preparation of ammonium salts.
“Chemical properties of salts” - Salt of a weak base and a strong acid. Non-metal. Na2CO3 + 2HCl. Salt of a strong base and a strong acid. Electrolytes. Determination of salts. Row of metal. Genetic relationships between classes of inorganic compounds. Chemical properties of salts. Universal indicator. Complex substances. Classification of salts.
"Salt" - Salt. Calcite. Chemical formula - CaCO3. Calcium carbonate (calcium carbonate) is a salt of carbonic acid and calcium. formula of nitric acid HNO3 acidic residue NO3- - nitrate Let's create the formulas of salts: NaNO3 - Using the solubility table, we determine the charges of the ions. - NaCl. Calcium carbonate CaCO3. Marble. Therefore, Berthollet salt is used in pyrotechnics in the production of fireworks, sparklers, and matches.
“Salt of nitric acid” - Chemical properties of nitrates. What conclusions did the young chemist come to? Interesting story. Specify the oxidizing agent and reducing agent. Know and be able to. What substances are called salts? A solution of nitric acid reacts with each of the substances. Given pairs of substances, create possible reaction equations. Decomposition of copper(II) nitrate.
“Salts, acids and bases” - 8. Alkali + salt. Obtaining salts. CuO+H2SO4 ?CuSO4+H2O. Fe?FeO?FeSO4?Fe(OH)2? Fe(OH)Cl?FeCl2 2Fe+O2?2FeO; 2) FeO+H2SO4?FeSO4+H2O; Preparation of basic salts. 3) Insoluble. 2. Medium salt1+alkali?basic salt+medium salt2. CaCO3+CO2+H2O ?Ca(HCO3)2 Na2SO4+H2SO4 ?2NaHSO4. 2NaOH+Mg(NO3)2 ? 2NaNO3+Mg(OH)2?.
“Substance salt” - Plan for compiling hydrolysis: For example: NaHS - sodium hydrosulfide. The first type of hydrolysis. Simple. For salts of oxygen-free acids, the suffix - id is added to the name of the non-metal. D) ni. Hydrated. Let's consider an example of an ionic crystal lattice: Inorganic. Physical properties of salts. The fourth type of hydrolysis.
There are a total of 22 presentations in the topic
Typical reactions of acids, bases, oxides, salts (conditions for their implementation)
Typical acid reactions
1 . Acid + base → salt + water
2 . Acid + metal oxide → salt + water
3 . Acid + metal → salt + hydrogen (conditions: a) the metal must be in the electrochemical voltage series to the left of hydrogen; b) a soluble salt should be obtained; c) insoluble acid – silicic acid does not react with metals; d) concentrated sulfuric and nitric acids react differently with metals, hydrogen is not released)
4 . Acid + salt → new acid + new salt. (condition: the reaction occurs if a precipitate or gas is formed)
Typical base reactions
1 . Base + acid → salt + water
2 . Base + non-metal oxide → salt + water (condition: non-metal oxide – acid oxide)
3 . Alkali + salt → new base + new salt (condition: precipitate or gas is formed)
Typical reactions of basic oxides
1 . Basic oxide + acid → salt + water
2 . Basic oxide + acidic oxide → salt
3 . Basic oxide + water → alkali (condition: a soluble alkali base is formed)
Typical acid oxide reactions
1 . Acidic oxide + base → salt + water
2 . Acidic oxide + basic oxide → salt
3 . Acidic oxide + water → acid (condition: the acid must be soluble)
Typical salt reactions
1 . Salt + acid → another salt + another acid (condition: if a precipitate or gas is formed)
2 . Salt + alkali → another salt + another base (condition: if a precipitate or gas is formed)
3 . Salt 1 + salt 2 → salt 3 + salt 4 (condition: a precipitate is formed)
4 . Salt + metal → another salt + another metal (condition: each metal displaces from salt solutions all other metals located to the right of it in the voltage series; both salts must be soluble)