Acidic substances. Acids: classification and chemical properties
| Oxygen-free: | Basicity | Name of salt |
| HCl - hydrochloric (hydrochloric) | monobasic | chloride |
| HBr - hydrobromic | monobasic | bromide |
| HI - hydroiodide | monobasic | iodide |
| HF - hydrofluoric (fluoric) | monobasic | fluoride |
| H 2 S - hydrogen sulfide | dibasic | sulfide |
| Oxygen-containing: | ||
| HNO 3 – nitrogen | monobasic | nitrate |
| H 2 SO 3 - sulfurous | dibasic | sulfite |
| H 2 SO 4 – sulfuric | dibasic | sulfate |
| H 2 CO 3 - coal | dibasic | carbonate |
| H 2 SiO 3 - silicon | dibasic | silicate |
| H 3 PO 4 - orthophosphoric | tribasic | orthophosphate |
Salts – complex substances that consist of metal atoms and acidic residues. This is the most numerous class of inorganic compounds.
Classification. By composition and properties: medium, acidic, basic, double, mixed, complex
Medium salts are products of complete replacement of the hydrogen atoms of a polybasic acid with metal atoms.
Upon dissociation, only metal cations (or NH 4 +) are produced. For example:
Na 2 SO 4 ® 2Na + +SO
CaCl 2 ® Ca 2+ + 2Cl -
Acid salts are products of incomplete replacement of hydrogen atoms of a polybasic acid with metal atoms.
Upon dissociation, they produce metal cations (NH 4 +), hydrogen ions and anions of the acid residue, for example:
NaHCO 3 ® Na + + HCO « H + +CO .
Basic salts are products of incomplete replacement of OH groups - the corresponding base with acidic residues.
Upon dissociation, they give metal cations, hydroxyl anions and an acid residue.
Zn(OH)Cl ® + + Cl - « Zn 2+ + OH - + Cl - .
Double salts contain two metal cations and upon dissociation give two cations and one anion.
KAl(SO 4) 2 ® K + + Al 3+ + 2SO
Complex salts contain complex cations or anions.
Br ® + + Br - « Ag + +2 NH 3 + Br -
Na ® Na + + - « Na + + Ag + + 2 CN -
Genetic relationship between different classes of compounds
EXPERIMENTAL PART
Equipment and utensils: rack with test tubes, washing machine, alcohol lamp.
Reagents and materials: red phosphorus, zinc oxide, Zn granules, slaked lime powder Ca(OH) 2, 1 mol/dm 3 solutions of NaOH, ZnSO 4, CuSO 4, AlCl 3, FeCl 3, HСl, H 2 SO 4, universal indicator paper, solution phenolphthalein, methyl orange, distilled water.
Work order
1. Pour zinc oxide into two test tubes; add an acid solution (HCl or H 2 SO 4) to one and an alkali solution (NaOH or KOH) to the other and heat slightly on an alcohol lamp.
Observations: Does zinc oxide dissolve in an acid and alkali solution?
Write equations
Conclusions: 1.What type of oxide does ZnO belong to?
2. What properties do amphoteric oxides have?
Preparation and properties of hydroxides
2.1. Dip the tip of the universal indicator strip into the alkali solution (NaOH or KOH). Compare the resulting color of the indicator strip with the standard color scale.
Observations: Record the pH value of the solution.
2.2. Take four test tubes, pour 1 ml of ZnSO 4 solution into the first, CuSO 4 into the second, AlCl 3 into the third, and FeCl 3 into the fourth. Add 1 ml of NaOH solution to each test tube. Write observations and equations for the reactions occurring.
Observations: Does precipitation occur when alkali is added to a salt solution? Indicate the color of the sediment.
Write equations occurring reactions (in molecular and ionic form).
Conclusions: How can metal hydroxides be prepared?
2.3. Transfer half of the sediments obtained in experiment 2.2 to other test tubes. Treat one part of the sediment with a solution of H 2 SO 4 and the other with a solution of NaOH.
Observations: Does precipitate dissolution occur when alkali and acid are added to precipitates?
Write equations occurring reactions (in molecular and ionic form).
Conclusions: 1. What type of hydroxides are Zn(OH) 2, Al(OH) 3, Cu(OH) 2, Fe(OH) 3?
2. What properties do amphoteric hydroxides have?
Obtaining salts.
3.1. Pour 2 ml of CuSO 4 solution into a test tube and dip a cleaned nail into this solution. (The reaction is slow, changes on the surface of the nail appear after 5-10 minutes).
Observations: Are there any changes to the surface of the nail? What is being deposited?
Write the equation for the redox reaction.
Conclusions: Taking into account the range of metal stresses, indicate the method of obtaining salts.
3.2. Place one zinc granule in a test tube and add HCl solution.
Observations: Is there any gas evolution?
Write the equation
Conclusions: Explain this method obtaining salts?
3.3. Pour some slaked lime powder Ca(OH) 2 into a test tube and add HCl solution.
Observations: Is there gas evolution?
Write the equation the reaction taking place (in molecular and ionic form).
Conclusion: 1. What type of reaction is the interaction between a hydroxide and an acid?
2.What substances are the products of this reaction?
3.5. Pour 1 ml of salt solutions into two test tubes: into the first - copper sulfate, into the second - cobalt chloride. Add to both test tubes drop by drop sodium hydroxide solution until precipitation forms. Then add excess alkali to both test tubes.
Observations: Indicate the changes in the color of precipitation in the reactions.
Write the equation the reaction taking place (in molecular and ionic form).
Conclusion: 1. As a result of what reactions are basic salts formed?
2. How can you convert basic salts to medium salts?
Test tasks:
1. From the listed substances, write down the formulas of salts, bases, acids: Ca(OH) 2, Ca(NO 3) 2, FeCl 3, HCl, H 2 O, ZnS, H 2 SO 4, CuSO 4, KOH
Zn(OH) 2, NH 3, Na 2 CO 3, K 3 PO 4.
2. Indicate the formulas of the oxides corresponding to the listed substances H 2 SO 4, H 3 AsO 3, Bi(OH) 3, H 2 MnO 4, Sn(OH) 2, KOH, H 3 PO 4, H 2 SiO 3, Ge( OH) 4 .
3. Which hydroxides are amphoteric? Write down reaction equations characterizing the amphotericity of aluminum hydroxide and zinc hydroxide.
4. Which of the following compounds will interact in pairs: P 2 O 5 , NaOH, ZnO, AgNO 3 , Na 2 CO 3 , Cr(OH) 3 , H 2 SO 4 . Write down equations for possible reactions.
Laboratory work No. 2 (4 hours)
Subject: Qualitative analysis of cations and anions
Target: master the technique of conducting qualitative and group reactions on cations and anions.
THEORETICAL PART
The main task of qualitative analysis is to establish chemical composition substances found in various objects (biological materials, medicines, food products, objects environment). This work examines the qualitative analysis of inorganic substances that are electrolytes, i.e., essentially the qualitative analysis of ions. From the entire set of occurring ions, the most important in medical and biological terms were selected: (Fe 3+, Fe 2+, Zn 2+, Ca 2+, Na +, K +, Mg 2+, Cl -, PO, CO, etc. ). Many of these ions are part of various medicines and food products.
Not all are used in qualitative analysis possible reactions, but only those that are accompanied by a clear analytical effect. The most common analytical effects: the appearance of a new color, the release of gas, the formation of a precipitate.
There are two fundamentally different approaches to qualitative analysis: fractional and systematic . In systematic analysis, group reagents are necessarily used to separate the ions present into separate groups, and in some cases into subgroups. To do this, some of the ions are transferred into insoluble compounds, and some of the ions are left in solution. After separating the precipitate from the solution, they are analyzed separately.
For example, the solution contains A1 3+, Fe 3+ and Ni 2+ ions. If this solution is exposed to excess alkali, a precipitate of Fe(OH) 3 and Ni(OH) 2 precipitates, and [A1(OH) 4 ] - ions remain in the solution. The precipitate containing iron and nickel hydroxides will partially dissolve when treated with ammonia due to the transition to 2+ solution. Thus, using two reagents - alkali and ammonia, two solutions were obtained: one contained [A1(OH) 4 ] - ions, the other contained 2+ ions and a Fe(OH) 3 precipitate. With the help of characteristic reactions, the presence of certain ions in solutions and in the precipitate, which must first be dissolved, is proven.
Systematic analysis is used mainly for the detection of ions in complex multicomponent mixtures. It is very labor-intensive, but its advantage lies in the easy formalization of all actions that fit into a clear scheme (methodology).
To carry out fractional analysis, only characteristic reactions are used. Obviously, the presence of other ions can significantly distort the results of the reaction (overlapping colors, unwanted precipitation, etc.). To avoid this, fractional analysis mainly uses highly specific reactions that give an analytical effect with a small number of ions. For successful implementation reactions, it is very important to maintain certain conditions, in particular pH. Very often in fractional analysis it is necessary to resort to masking, i.e., converting ions into compounds that are not capable of producing an analytical effect with the selected reagent. For example, dimethylglyoxime is used to detect nickel ion. The Fe 2+ ion gives a similar analytical effect to this reagent. To detect Ni 2+, the Fe 2+ ion is transferred to a stable fluoride complex 4- or oxidized to Fe 3+, for example, with hydrogen peroxide.
Fractional analysis is used to detect ions in simpler mixtures. Analysis time is significantly reduced, but at the same time the experimenter is required to have a deeper knowledge of the flow patterns chemical reactions, since to take into account in one specific technique all possible cases The mutual influence of ions on the nature of the observed analytical effects is quite difficult.
In analytical practice, the so-called fractional-systematic method. With this approach, a minimum number of group reagents is used, which makes it possible to outline analysis tactics in general outline, which is then carried out using the fractional method.
According to the technique of conducting analytical reactions, reactions are distinguished: sedimentary; microcrystalscopic; accompanied by the release of gaseous products; conducted on paper; extraction; colored in solutions; flame coloring.
When carrying out sedimentary reactions, be sure to note the color and nature of the sediment (crystalline, amorphous), and, if necessary, carry out additional tests: check the precipitate for solubility in strong and weak acids, alkalis and ammonia, and excess reagent. When carrying out reactions accompanied by the release of gas, its color and smell are noted. In some cases, additional tests are carried out.
For example, if the gas released is suspected to be carbon monoxide (IV), it is passed through an excess of lime water.
In fractional and systematic analyses, reactions during which a new color appears are widely used, most often these are complexation reactions or redox reactions.
In some cases, it is convenient to carry out such reactions on paper (droplet reactions). Reagents that do not decompose under normal conditions are applied to paper in advance. Thus, to detect hydrogen sulfide or sulfide ions, paper impregnated with lead nitrate is used [blackening occurs due to the formation of lead(II) sulfide]. Many oxidizing agents are detected using iodine starch paper, i.e. paper soaked in solutions of potassium iodide and starch. In most cases, the necessary reagents are applied to paper during the reaction, for example, alizarin for the A1 3+ ion, cupron for the Cu 2+ ion, etc. To enhance the color, extraction into an organic solvent is sometimes used. For preliminary tests, flame color reactions are used.
7. Acids. Salt. Relationship between classes of inorganic substances
7.1. Acids
Acids are electrolytes, upon the dissociation of which only hydrogen cations H + are formed as positively charged ions (more precisely, hydronium ions H 3 O +).
Another definition: acids are complex substances consisting of a hydrogen atom and acid residues (Table 7.1).
Table 7.1
Formulas and names of some acids, acid residues and salts
| Acid formula | Acid name | Acid residue (anion) | Name of salts (average) |
|---|---|---|---|
| HF | Hydrofluoric (fluoric) | F − | Fluorides |
| HCl | Hydrochloric (hydrochloric) | Cl − | Chlorides |
| HBr | Hydrobromic | Br− | Bromides |
| HI | Hydroiodide | I − | Iodides |
| H2S | Hydrogen sulfide | S 2− | Sulfides |
| H2SO3 | Sulphurous | SO 3 2 − | Sulfites |
| H2SO4 | Sulfuric | SO 4 2 − | Sulfates |
| HNO2 | Nitrogenous | NO2− | Nitrites |
| HNO3 | Nitrogen | NO 3 − | Nitrates |
| H2SiO3 | Silicon | SiO 3 2 − | Silicates |
| HPO 3 | Metaphosphoric | PO 3 − | Metaphosphates |
| H3PO4 | Orthophosphoric | PO 4 3 − | Orthophosphates (phosphates) |
| H4P2O7 | Pyrophosphoric (biphosphoric) | P 2 O 7 4 − | Pyrophosphates (diphosphates) |
| HMnO4 | Manganese | MnO 4 − | Permanganates |
| H2CrO4 | Chrome | CrO 4 2 − | Chromates |
| H2Cr2O7 | Dichrome | Cr 2 O 7 2 − | Dichromates (bichromates) |
| H2SeO4 | Selenium | SeO 4 2 − | Selenates |
| H3BO3 | Bornaya | BO 3 3 − | Orthoborates |
| HClO | Hypochlorous | ClO – | Hypochlorites |
| HClO2 | Chloride | ClO2− | Chlorites |
| HClO3 | Chlorous | ClO3− | Chlorates |
| HClO4 | Chlorine | ClO 4 − | Perchlorates |
| H2CO3 | Coal | CO 3 3 − | Carbonates |
| CH3COOH | Vinegar | CH 3 COO − | Acetates |
| HCOOH | Ant | HCOO − | Formiates |
Under normal conditions, acids can be solids (H 3 PO 4, H 3 BO 3, H 2 SiO 3) and liquids (HNO 3, H 2 SO 4, CH 3 COOH). These acids can exist both individually (100% form) and in the form of diluted and concentrated solutions. For example, H 2 SO 4 , HNO 3 , H 3 PO 4 , CH 3 COOH are known both individually and in solutions.
A number of acids are known only in solutions. These are all hydrogen halides (HCl, HBr, HI), hydrogen sulfide H 2 S, hydrogen cyanide (hydrocyanic HCN), carbonic H 2 CO 3, sulfurous H 2 SO 3 acid, which are solutions of gases in water. For example, hydrochloric acid is a mixture of HCl and H 2 O, carbonic acid is a mixture of CO 2 and H 2 O. It is clear that using the expression “solution of hydrochloric acid" wrong.
Most acids are soluble in water; silicic acid H 2 SiO 3 is insoluble. The overwhelming majority of acids have a molecular structure. Examples of structural formulas of acids:
In most oxygen-containing acid molecules, all hydrogen atoms are bonded to oxygen. But there are exceptions:
Acids are classified according to a number of characteristics (Table 7.2).
Table 7.2
Classification of acids
| Classification sign | Acid type | Examples |
|---|---|---|
| Number of hydrogen ions formed upon complete dissociation of an acid molecule | Monobase | HCl, HNO3, CH3COOH |
| Dibasic | H2SO4, H2S, H2CO3 | |
| Tribasic | H3PO4, H3AsO4 | |
| The presence or absence of an oxygen atom in a molecule | Oxygen-containing (acid hydroxides, oxoacids) | HNO2, H2SiO3, H2SO4 |
| Oxygen-free | HF, H2S, HCN | |
| Degree of dissociation (strength) | Strong (completely dissociate, strong electrolytes) | HCl, HBr, HI, H2SO4 (diluted), HNO3, HClO3, HClO4, HMnO4, H2Cr2O7 |
| Weak (partially dissociate, weak electrolytes) | HF, HNO 2, H 2 SO 3, HCOOH, CH 3 COOH, H 2 SiO 3, H 2 S, HCN, H 3 PO 4, H 3 PO 3, HClO, HClO 2, H 2 CO 3, H 3 BO 3, H 2 SO 4 (conc) | |
| Oxidative properties | Oxidizing agents due to H + ions (conditionally non-oxidizing acids) | HCl, HBr, HI, HF, H 2 SO 4 (dil), H 3 PO 4, CH 3 COOH |
| Oxidizing agents due to anion (oxidizing acids) | HNO 3, HMnO 4, H 2 SO 4 (conc), H 2 Cr 2 O 7 | |
| Reducing agents due to anion | HCl, HBr, HI, H 2 S (but not HF) | |
| Thermal stability | Exist only in solutions | H 2 CO 3, H 2 SO 3, HClO, HClO 2 |
| Easily decomposes when heated | H2SO3, HNO3, H2SiO3 | |
| Thermally stable | H 2 SO 4 (conc), H 3 PO 4 |
All general chemical properties of acids are due to the presence in their aqueous solutions of excess hydrogen cations H + (H 3 O +).
1. Due to the excess of H + ions, aqueous solutions of acids change the color of litmus violet and methyl orange to red (phenolphthalein does not change color and remains colorless). IN aqueous solution weak carbonic acid, litmus is not red, but pink; a solution over a precipitate of very weak silicic acid does not change the color of the indicators at all.
2. Acids interact with basic oxides, bases and amphoteric hydroxides, ammonia hydrate (see Chapter 6).
Example 7.1.
To carry out the transformation BaO → BaSO 4 you can use: a) SO 2; b) H 2 SO 4; c) Na 2 SO 4; d) SO 3.
Solution. The transformation can be carried out using H 2 SO 4:
BaO + H 2 SO 4 = BaSO 4 ↓ + H 2 O
BaO + SO 3 = BaSO 4
Na 2 SO 4 does not react with BaO, and in the reaction of BaO with SO 2 barium sulfite is formed:
BaO + SO 2 = BaSO 3
Answer: 3).
3. Acids react with ammonia and its aqueous solutions to form ammonium salts:
HCl + NH 3 = NH 4 Cl - ammonium chloride;
H 2 SO 4 + 2NH 3 = (NH 4) 2 SO 4 - ammonium sulfate.
4. Non-oxidizing acids react with metals located in the activity series up to hydrogen to form a salt and release hydrogen:
H 2 SO 4 (diluted) + Fe = FeSO 4 + H 2
2HCl + Zn = ZnCl 2 = H 2
The interaction of oxidizing acids (HNO 3, H 2 SO 4 (conc)) with metals is very specific and is considered when studying the chemistry of elements and their compounds.
5. Acids interact with salts. The reaction has a number of features:
a) in most cases, when a stronger acid reacts with a salt of a weaker acid, a salt of a weak acid and a weak acid are formed, or, as they say, a stronger acid displaces a weaker one. The series of decreasing strength of acids looks like this:
2HCl + Na 2 CO 3 = 2NaCl + H 2 O + CO 2
H 2 CO 3 + Na 2 SiO 3 = Na 2 CO 3 + H 2 SiO 3 ↓
2CH 3 COOH + K 2 CO 3 = 2CH 3 COOK + H 2 O + CO 2
3H 2 SO 4 + 2K 3 PO 4 = 3K 2 SO 4 + 2H 3 PO 4
Do not interact with each other, for example, KCl and H 2 SO 4 (diluted), NaNO 3 and H 2 SO 4 (diluted), K 2 SO 4 and HCl (HNO 3, HBr, HI), K 3 PO 4 and H 2 CO 3, CH 3 COOK and H 2 CO 3;
b) in some cases, a weaker acid displaces a stronger one from a salt:
CuSO 4 + H 2 S = CuS↓ + H 2 SO 4
3AgNO 3 (dil) + H 3 PO 4 = Ag 3 PO 4 ↓ + 3HNO 3.
Such reactions are possible when the precipitates of the resulting salts do not dissolve in the resulting dilute strong acids (H 2 SO 4 and HNO 3);
c) in the case of the formation of precipitates that are insoluble in strong acids, a reaction may occur between strong acid and a salt formed by another strong acid:
BaCl 2 + H 2 SO 4 = BaSO 4 ↓ + 2HCl
Ba(NO 3) 2 + H 2 SO 4 = BaSO 4 ↓ + 2HNO 3
AgNO 3 + HCl = AgCl↓ + HNO 3
Example 7.2.
Indicate the row containing the formulas of substances that react with H 2 SO 4 (diluted).
1) Zn, Al 2 O 3, KCl (p-p); 3) NaNO 3 (p-p), Na 2 S, NaF; 2) Cu(OH) 2, K 2 CO 3, Ag; 4) Na 2 SO 3, Mg, Zn(OH) 2.
Solution. All substances of row 4 interact with H 2 SO 4 (dil):
Na 2 SO 3 + H 2 SO 4 = Na 2 SO 4 + H 2 O + SO 2
Mg + H 2 SO 4 = MgSO 4 + H 2
Zn(OH) 2 + H 2 SO 4 = ZnSO 4 + 2H 2 O
In row 1) the reaction with KCl (p-p) is not feasible, in row 2) - with Ag, in row 3) - with NaNO 3 (p-p).
Answer: 4).
6. Concentrated sulfuric acid behaves very specifically in reactions with salts. This is a non-volatile and thermally stable acid, therefore it displaces all strong acids from solid (!) salts, since they are more volatile than H2SO4 (conc):
KCl (tv) + H 2 SO 4 (conc.) KHSO 4 + HCl
2KCl (s) + H 2 SO 4 (conc) K 2 SO 4 + 2HCl
Salts formed by strong acids (HBr, HI, HCl, HNO 3, HClO 4) react only with concentrated sulfuric acid and only when in a solid state
Example 7.3.
Concentrated sulfuric acid, unlike dilute one, reacts:
BaO + SO 2 = BaSO 3
3) KNO 3 (tv);
Solution. Both acids react with KF, Na 2 CO 3 and Na 3 PO 4, and only H 2 SO 4 (conc.) react with KNO 3 (solid). Methods for producing acids are very diverse.
- Anoxic acids
receive:
by dissolving the corresponding gases in water:
- HCl (g) + H 2 O (l) → HCl (p-p)
H 2 S (g) + H 2 O (l) → H 2 S (solution)
from salts by displacement with stronger or less volatile acids:
FeS + 2HCl = FeCl 2 + H 2 S
Oxygen-containing acids Methods for producing acids are very diverse.
- by dissolving the corresponding acidic oxides in water, while the degree of oxidation of the acid-forming element in the oxide and acid remains the same (with the exception of NO 2):
N2O5 + H2O = 2HNO3
SO 3 + H 2 O = H 2 SO 4
P 2 O 5 + 3H 2 O 2H 3 PO 4
- oxidation of non-metals with oxidizing acids:
S + 6HNO 3 (conc) = H 2 SO 4 + 6NO 2 + 2H 2 O
- by displacing a strong acid from a salt of another strong acid (if a precipitate insoluble in the resulting acids precipitates):
Ba(NO 3) 2 + H 2 SO 4 (diluted) = BaSO 4 ↓ + 2HNO 3
AgNO 3 + HCl = AgCl↓ + HNO 3
- by displacing a volatile acid from its salts with a less volatile acid.
For this purpose, non-volatile, thermally stable concentrated sulfuric acid is most often used:
NaNO 3 (tv) + H 2 SO 4 (conc.) NaHSO 4 + HNO 3
KClO 4 (tv) + H 2 SO 4 (conc.) KHSO 4 + HClO 4
- displacement of a weaker acid from its salts by a stronger acid:
Ca 3 (PO 4) 2 + 3H 2 SO 4 = 3CaSO 4 ↓ + 2H 3 PO 4
NaNO 2 + HCl = NaCl + HNO 2
K 2 SiO 3 + 2HBr = 2KBr + H 2 SiO 3 ↓
Acids are chemical compounds that are capable of donating an electrically charged hydrogen ion (cation) and also accepting two interacting electrons, resulting in the formation of a covalent bond.
In this article we will look at the main acids that are studied in the middle classes of secondary schools, and also learn many interesting facts about a variety of acids. Let's get started.
Acids: types
In chemistry, there are many different acids that have very different properties. Chemists distinguish acids by their oxygen content, volatility, solubility in water, strength, stability, and whether they belong to the organic or inorganic class. chemical compounds. In this article we will look at a table that presents the most famous acids. The table will help you remember the name of the acid and its chemical formula.
So, everything is clearly visible. This table presents the most famous chemical industry acids. The table will help you remember names and formulas much faster.
Hydrogen sulfide acid
H 2 S is hydrosulfide acid. Its peculiarity lies in the fact that it is also a gas. Hydrogen sulfide is very poorly soluble in water, and also interacts with many metals. Hydrogen sulfide acid belongs to the group of “weak acids”, examples of which we will consider in this article.
H 2 S has a slightly sweet taste and also a very pungent odor rotten eggs. In nature, it can be found in natural or volcanic gases, and it is also released during protein decay.
The properties of acids are very diverse; even if an acid is indispensable in industry, it can be very harmful to human health. This acid is very toxic to humans. When a small amount of hydrogen sulfide is inhaled, a person awakens headache, severe nausea and dizziness begin. If a person inhales a large number of H 2 S, it can lead to seizures, coma or even instant death.
Sulfuric acid
H 2 SO 4 is a strong sulfuric acid, which children are introduced to in chemistry lessons in the 8th grade. Chemical acids such as sulfuric acid are very strong oxidizing agents. H 2 SO 4 acts as an oxidizing agent on many metals, as well as basic oxides.
H 2 SO 4 in contact with skin or clothing causes chemical burns, however, it is not as toxic as hydrogen sulfide.
Nitric acid
Strong acids are very important in our world. Examples of such acids: HCl, H 2 SO 4, HBr, HNO 3. HNO 3 is a well-known nitric acid. She found wide application in industry, as well as in agriculture. It is used to make various fertilizers, in jewelry, in photograph printing, in the production of medicines and dyes, as well as in the military industry.
Such chemical acids, like nitrogen, are very harmful to the body. HNO 3 vapors leave ulcers and cause acute inflammation and respiratory irritation.
Nitrous acid
Nitrous acid is often confused with nitric acid, but there is a difference between them. The fact is that it is much weaker than nitrogen, it has completely different properties and effects on the human body.
HNO 2 is widely used in the chemical industry.
Hydrofluoric acid
Hydrofluoric acid (or hydrogen fluoride) is a solution of H 2 O with HF. The acid formula is HF. Hydrofluoric acid is very actively used in the aluminum industry. It is used to dissolve silicates, etch silicon and silicate glass.
Hydrogen fluoride is very harmful to the human body and, depending on its concentration, can be a mild narcotic. In case of contact with the skin, there are no changes at first, but after a few minutes it may appear. sharp pain and chemical burn. Hydrofluoric acid is very harmful to the environment.
Hydrochloric acid
HCl is hydrogen chloride and is a strong acid. Hydrogen chloride retains the properties of acids belonging to the group of strong acids. The acid is transparent and colorless in appearance, but smokes in air. Hydrogen chloride is widely used in the metallurgical and food industries.
This acid causes chemical burns, but getting into the eyes is especially dangerous.
Phosphoric acid
Phosphoric acid (H 3 PO 4) is a weak acid in its properties. But even weak acids can have the properties of strong ones. For example, H 3 PO 4 is used in industry to restore iron from rust. In addition, phosphoric (or orthophosphoric) acid is widely used in agriculture - many different fertilizers are made from it.
The properties of acids are very similar - almost each of them is very harmful to the human body, H 3 PO 4 is no exception. For example, this acid also causes severe chemical burns, nosebleeds, and chipping of teeth.
Carbonic acid
H 2 CO 3 is a weak acid. It is obtained by dissolving CO 2 (carbon dioxide) in H 2 O (water). Carbonic acid is used in biology and biochemistry.
Density of various acids
The density of acids is important place in theoretical and practical parts of chemistry. By knowing the density, you can determine the concentration of a particular acid, solve chemical calculation problems, and add the correct amount of acid to complete the reaction. The density of any acid changes depending on the concentration. For example, the higher the concentration percentage, the higher the density.
General properties of acids
Absolutely all acids are (that is, they consist of several elements of the periodic table), and they necessarily include H (hydrogen) in their composition. Next we will look at which are common:
- All oxygen-containing acids (in the formula of which O is present) form water upon decomposition, and also oxygen-free ones decompose into simple substances(for example, 2HF decomposes into F 2 and H 2).
- Oxidizing acids react with all metals in the metal activity series (only those located to the left of H).
- They interact with various salts, but only with those that were formed by an even weaker acid.
According to their own physical properties acids differ sharply from each other. After all, they can have a smell or not, and also be in a variety of physical states: liquid, gaseous and even solid. Solid acids are very interesting to study. Examples of such acids: C 2 H 2 0 4 and H 3 BO 3.
Concentration
Concentration is a value that determines the quantitative composition of any solution. For example, chemists often need to determine how much pure sulfuric acid is present in dilute acid H 2 SO 4. To do this they pour a small amount of dilute acid into a measuring cup, weigh it and determine the concentration using the density table. The concentration of acids is closely related to density; often, when determining the concentration, there are calculation problems where you need to determine the percentage of pure acid in a solution.
Classification of all acids according to the number of H atoms in their chemical formula
One of the most popular classifications is the division of all acids into monobasic, dibasic and, accordingly, tribasic acids. Examples of monobasic acids: HNO 3 (nitric), HCl (hydrochloric), HF (hydrofluoric) and others. These acids are called monobasic, since they contain only one H atom. There are many such acids, it is impossible to remember absolutely every one. You just need to remember that acids are also classified according to the number of H atoms in their composition. Dibasic acids are defined similarly. Examples: H 2 SO 4 (sulphuric), H 2 S (hydrogen sulfide), H 2 CO 3 (coal) and others. Tribasic: H 3 PO 4 (phosphoric).
Basic classification of acids
One of the most popular classifications of acids is their division into oxygen-containing and oxygen-free. How to remember, without knowing the chemical formula of a substance, that it is an oxygen-containing acid?
All oxygen-free acids do not contain important element O is oxygen, but it contains H. Therefore, the word “hydrogen” is always attached to their name. HCl is a H 2 S - hydrogen sulfide.
But you can also write a formula based on the names of acid-containing acids. For example, if the number of O atoms in a substance is 4 or 3, then the suffix -n-, as well as the ending -aya-, is always added to the name:
- H 2 SO 4 - sulfur (number of atoms - 4);
- H 2 SiO 3 - silicon (number of atoms - 3).
If the substance has less than three oxygen atoms or three, then the suffix -ist- is used in the name:
- HNO 2 - nitrogenous;
- H 2 SO 3 - sulfurous.
General properties
All acids taste sour and often slightly metallic. But there are other similar properties that we will now consider.
There are substances called indicators. The indicators change their color, or the color remains, but its shade changes. This occurs when the indicators are affected by other substances, such as acids.
An example of a color change is such a familiar product as tea, and lemon acid. When lemon is added to tea, the tea gradually begins to noticeably brighten. This is due to the fact that lemon contains citric acid.
There are other examples. Litmus, which in a neutral environment has purple colour turns red when hydrochloric acid is added.
When the tensions are in the tension series before hydrogen, gas bubbles are released - H. However, if a metal that is in the tension series after H is placed in a test tube with acid, then no reaction will occur, there will be no gas evolution. So, copper, silver, mercury, platinum and gold will not react with acids.
In this article we examined the most famous chemical acids, as well as their main properties and differences.
Acids- electrolytes, upon dissociation of which only H + ions are formed from positive ions:
HNO 3 ↔ H + + NO 3 - ;
CH 3 COOH↔ H + +CH 3 COO — .
All acids are classified into inorganic and organic (carboxylic), which also have their own (internal) classifications.
Under normal conditions, significant amounts of inorganic acids exist in liquid state, some are in the solid state (H 3 PO 4, H 3 BO 3).
Organic acids with up to 3 carbon atoms are highly mobile, colorless liquids with a characteristic pungent odor; acids with 4-9 carbon atoms - oily liquids with unpleasant smell, and acids with a large number of carbon atoms are solids that are insoluble in water.
Chemical formulas of acids
Let us consider the chemical formulas of acids using the example of several representatives (both inorganic and organic): hydrochloric acid - HCl, sulfuric acid - H 2 SO 4, phosphoric acid - H 3 PO 4, acetic acid - CH 3 COOH and benzoic acid - C 6 H5COOH. The chemical formula shows the qualitative and quantitative composition of the molecule (how many and which atoms are included in a particular compound). Using the chemical formula, you can calculate the molecular weight of acids (Ar(H) = 1 amu, Ar(Cl) = 35.5 amu. amu, Ar(P) = 31 amu, Ar(O) = 16 amu, Ar(S) = 32 amu, Ar(C) = 12 a.m.):
Mr(HCl) = Ar(H) + Ar(Cl);
Mr(HCl) = 1 + 35.5 = 36.5.
Mr(H 2 SO 4) = 2×Ar(H) + Ar(S) + 4×Ar(O);
Mr(H 2 SO 4) = 2×1 + 32 + 4×16 = 2 + 32 + 64 = 98.
Mr(H 3 PO 4) = 3×Ar(H) + Ar(P) + 4×Ar(O);
Mr(H 3 PO 4) = 3×1 + 31 + 4×16 = 3 + 31 + 64 = 98.
Mr(CH 3 COOH) = 3×Ar(C) + 4×Ar(H) + 2×Ar(O);
Mr(CH 3 COOH) = 3×12 + 4×1 + 2×16 = 36 + 4 + 32 = 72.
Mr(C 6 H 5 COOH) = 7×Ar(C) + 6×Ar(H) + 2×Ar(O);
Mr(C 6 H 5 COOH) = 7 × 12 + 6 × 1 + 2 × 16 = 84 + 6 + 32 = 122.
Structural (graphic) formulas of acids
The structural (graphic) formula of a substance is more visual. It shows how atoms are connected to each other within a molecule. Let us indicate the structural formulas of each of the above compounds:
Rice. 1. Structural formula of hydrochloric acid.
Rice. 2. Structural formula of sulfuric acid.
Rice. 3. Structural formula of phosphoric acid.
Rice. 4. Structural formula of acetic acid.
Rice. 5. Structural formula of benzoic acid.
Ionic formulas
All inorganic acids are electrolytes, i.e. capable of dissociating in an aqueous solution into ions:
HCl ↔ H + + Cl - ;
H 2 SO 4 ↔ 2H + + SO 4 2- ;
H 3 PO 4 ↔ 3H + + PO 4 3- .
Examples of problem solving
EXAMPLE 1
| Exercise | With complete combustion 6 g organic matter 8.8 g of carbon monoxide (IV) and 3.6 g of water were formed. Determine the molecular formula of the burned substance if it is known that its molar mass is 180 g/mol. |
| Solution | Let’s draw up a diagram of the combustion reaction of an organic compound, designating the number of carbon, hydrogen and oxygen atoms as “x”, “y” and “z”, respectively: C x H y O z + O z →CO 2 + H 2 O. Let us determine the masses of the elements that make up this substance. Values of relative atomic masses taken from the Periodic Table of D.I. Mendeleev, round to whole numbers: Ar(C) = 12 amu, Ar(H) = 1 amu, Ar(O) = 16 amu. m(C) = n(C)×M(C) = n(CO 2)×M(C) = ×M(C); m(H) = n(H)×M(H) = 2×n(H 2 O)×M(H) = ×M(H); Let's calculate the molar masses of carbon dioxide and water. As is known, the molar mass of a molecule is equal to the sum of the relative atomic masses of the atoms that make up the molecule (M = Mr): M(CO 2) = Ar(C) + 2×Ar(O) = 12+ 2×16 = 12 + 32 = 44 g/mol; M(H 2 O) = 2×Ar(H) + Ar(O) = 2×1+ 16 = 2 + 16 = 18 g/mol. m(C) = ×12 = 2.4 g; m(H) = 2 × 3.6 / 18 × 1 = 0.4 g. m(O) = m(C x H y O z) - m(C) - m(H) = 6 - 2.4 - 0.4 = 3.2 g. Let's determine the chemical formula of the compound: x:y:z = m(C)/Ar(C) : m(H)/Ar(H) : m(O)/Ar(O); x:y:z= 2.4/12:0.4/1:3.2/16; x:y:z= 0.2: 0.4: 0.2 = 1: 2: 1. This means the simplest formula of the compound is CH 2 O and the molar mass is 30 g/mol. To find the true formula of an organic compound, we find the ratio of the true and resulting molar masses: M substance / M(CH 2 O) = 180 / 30 = 6. This means that the indices of carbon, hydrogen and oxygen atoms should be 6 times higher, i.e. the formula of the substance will be C 6 H 12 O 6. This is glucose or fructose. |
| Answer | C6H12O6 |
EXAMPLE 2
| Exercise | Derive the simplest formula of a compound in which the mass fraction of phosphorus is 43.66%, and the mass fraction of oxygen is 56.34%. |
| Solution | Mass fraction element X in a molecule of composition HX is calculated using the following formula: ω (X) = n × Ar (X) / M (HX) × 100%. Let us denote the number of phosphorus atoms in the molecule by “x”, and the number of oxygen atoms by “y” Let's find the corresponding relative atomic masses elements of phosphorus and oxygen (relative atomic mass values taken from D.I. Mendeleev’s Periodic Table, rounded to whole numbers). Ar(P) = 31; Ar(O) = 16. We divide the percentage content of elements into the corresponding relative atomic masses. Thus we will find the relationship between the number of atoms in the molecule of the compound: x:y = ω(P)/Ar(P) : ω (O)/Ar(O); x:y = 43.66/31: 56.34/16; x:y: = 1.4: 3.5 = 1: 2.5 = 2: 5. This means that the simplest formula for combining phosphorus and oxygen is P 2 O 5 . It is phosphorus(V) oxide. |
| Answer | P2O5 |
Acids- complex substances consisting of one or more hydrogen atoms that can be replaced by metal atoms and acidic residues.
Classification of acids
1. By the number of hydrogen atoms: number of hydrogen atoms ( n ) determines the basicity of acids:
n= 1 monobase
n= 2 dibase
n= 3 tribase
2. By composition:
a) Table of oxygen-containing acids, acid residues and corresponding acid oxides:
|
Acid (H n A) |
Acid residue (A) |
Corresponding acid oxide |
|
H 2 SO 4 sulfuric |
SO 4 (II) sulfate |
SO3 sulfur oxide (VI) |
|
HNO 3 nitrogen |
NO3(I)nitrate |
N 2 O 5 nitric oxide (V) |
|
HMnO 4 manganese |
MnO 4 (I) permanganate |
Mn2O7 manganese oxide ( VII) |
|
H 2 SO 3 sulfurous |
SO 3 (II) sulfite |
SO2 sulfur oxide (IV) |
|
H 3 PO 4 orthophosphoric |
PO 4 (III) orthophosphate |
P 2 O 5 phosphorus oxide (V) |
|
HNO 2 nitrogenous |
NO 2 (I) nitrite |
N 2 O 3 nitric oxide (III) |
|
H 2 CO 3 coal |
CO 3 (II) carbonate |
CO2 carbon monoxide ( IV) |
|
H 2 SiO 3 silicon |
SiO 3 (II) silicate |
SiO 2 silicon(IV) oxide |
|
HClO hypochlorous |
ClO(I) hypochlorite |
C l 2 O chlorine oxide (I) |
|
HClO 2 chloride |
ClO 2 (I) chlorite |
C l 2 O 3 chlorine oxide (III) |
|
HClO 3 chloric |
ClO 3 (I) chlorate |
C l 2 O 5 chlorine oxide (V) |
|
HClO 4 chlorine |
ClO 4 (I) perchlorate |
C l 2 O 7 chlorine oxide (VII) |
b) Table of oxygen-free acids
|
Acid (H n A) |
Acid residue (A) |
|
HCl hydrochloric, hydrochloric |
Cl(I) chloride |
|
H 2 S hydrogen sulfide |
S(II) sulfide |
|
HBr hydrogen bromide |
Br(I) bromide |
|
HI hydrogen iodide |
I(I)iodide |
|
HF hydrogen fluoride, fluoride |
F(I) fluoride |
Physical properties of acids
Many acids, such as sulfuric, nitric, and hydrochloric, are colorless liquids. solid acids are also known: orthophosphoric, metaphosphoric HPO 3, boric H 3 BO 3 . Almost all acids are soluble in water. Example insoluble acid– silicon H2SiO3 . Acid solutions have a sour taste. For example, many fruits are given a sour taste by the acids they contain. Hence the names of acids: citric, malic, etc.
Methods for producing acids
|
oxygen-free |
oxygen-containing |
|
HCl, HBr, HI, HF, H2S |
HNO 3, H 2 SO 4 and others |
|
RECEIVING |
|
|
1. Direct interaction of nonmetals H 2 + Cl 2 = 2 HCl |
1. Acidic oxide + water = acid SO 3 + H 2 O = H 2 SO 4 |
|
2. Exchange reaction between salt and less volatile acid 2 NaCl (tv.) + H 2 SO 4 (conc.) = Na 2 SO 4 + 2HCl |
|
Chemical properties of acids
1. Change the color of the indicators
|
Indicator name |
Neutral environment |
Acidic environment |
|
Litmus |
Violet |
Red |
|
Phenolphthalein |
Colorless |
Colorless |
|
Methyl orange |
Orange |
Red |
|
Universal indicator paper |
Orange |
Red |
2. React with metals in the activity series up to H 2
(excl. HNO 3 -Nitric acid)
Video "Interaction of acids with metals"
Me + ACID = SALT + H 2 (r. substitution)
Zn + 2 HCl = ZnCl 2 + H 2
3. With basic (amphoteric) oxides – metal oxides
Video "Interaction of metal oxides with acids"
Fur x O y + ACID = SALT + H 2 O (exchange ruble)
4. React with bases – neutralization reaction
ACID + BASE= SALT+ H 2 O (exchange ruble)
H 3 PO 4 + 3 NaOH = Na 3 PO 4 + 3 H 2 O
5. React with salts of weak, volatile acids - if acid forms, precipitates or gas evolves:
2 NaCl (tv.) + H 2 SO 4 (conc.) = Na 2 SO 4 + 2HCl ( R . exchange )
Video "Interaction of acids with salts"
6. Decomposition of oxygen-containing acids when heated
(excl. H 2 SO 4 ; H 3 P.O. 4 )
ACID = ACID OXIDE + WATER (r. expansion)
Remember!Unstable acids (carbonic and sulfurous acids) - decompose into gas and water:
H 2 CO 3 ↔ H 2 O + CO 2
H 2 SO 3 ↔ H 2 O + SO 2
Hydrogen sulfide acid in products released as gas:
CaS + 2HCl = H 2 S+CaCl2
ASSIGNMENT TASKS
No. 1. Distribute chemical formulas acids in the table. Give them names:
LiOH, Mn 2 O 7, CaO, Na 3 PO 4, H 2 S, MnO, Fe (OH) 3, Cr 2 O 3, HI, HClO 4, HBr, CaCl 2, Na 2 O, HCl, H 2 SO 4, HNO 3, HMnO 4, Ca (OH) 2, SiO 2, Acids
Bes-sour-
native
Oxygen-containing
soluble
insoluble
one-
basic
two-basic
three-basic
No. 2. Write down the reaction equations:
Ca+HCl
Na+H2SO4
Al+H2S
Ca + H3PO4
Name the reaction products.
No. 3. Write down reaction equations and name the products:
Na 2 O + H 2 CO 3
ZnO + HCl
CaO + HNO3
Fe 2 O 3 + H 2 SO 4
No. 4. Write down reaction equations for the interaction of acids with bases and salts:
KOH + HNO3
NaOH + H2SO3
Ca(OH) 2 + H 2 S
Al(OH) 3 + HF
HCl + Na 2 SiO 3
H2SO4 + K2CO3
HNO3 + CaCO3
Name the reaction products.
EXERCISES
Trainer No. 1. "Formula and names of acids"
Trainer No. 2. "Establishing correspondence: acid formula - oxide formula"
Safety precautions - First aid if acids come into contact with skin
Safety precautions -