LESSON 1 …1.1 Physical and chemical properties…(and reactions) 1.2 Symbols of chemical elements 1.3 Atoms and molecules.
LESSON 2 ..2.1 Atomic mass and molecular mass… 2.2 The mole and molar mass….
LESSON 3 .. 3.1 Valence ….3.2 Chemical equations and their balancing
LESSON 4 … 4.1 Calculations using chemical equations… 4.2 Cases with limiting reactant..
LESSON 5 5.1 Molar volume of gases… 5.2 Relative densities of gases..
LESSON 6 .. Sample ticket for control task #1..
LESSON 7 . 7.1 The periodic table of elements…. 7.2 How to use the periodic table?…
LESSON 8 .. 8.1 Quantum numbers… (Nucleas and particles — add.
8.2 Electron configurations of atoms..
LESSON 9 ..9.1 Types of chemical bonds … 9.2 Electronegativity..
LESSON 10 ….10.1 Oxidation state .. 10.2 Oxidation state, oxidation number and valence.
LESSON 11 . 11.1 Several ways to classify chemical reactions.. 11.2 Redox reactions..
LESSON 12 ..12.1 Definitions of reduction and oxidation . 12.2 Balancing reduction-oxidation reactions. LESSON 13 ..13.1 Chemical equilibrium … 13.2 The law of mass action 13.3 Le Chatelier’s Principle LESSON 14 …14.1 The rate of chemical reaction…14.2 Factors influencing rate of reaction .. 14.3 Temperature coefficient of chemical reaction …
LESSON 15 .. Sample ticket for control task #2….
LESSON 16 … 16.1 Oxides…. 16.2 Basic and acidic anhydrides..
LESSON 17 .. 17.1 Bases… 17.2 Alkalis…
LESSON 18 …18.1 Acids…..18.2 Neutralization reaction…
LESSON 19 .. 19.1 Salts…19.2 Solubility chart of salts..
LESSON 20 .. 20.1 Classic chains of chemical reactions …. 20.2 Modern chains of chemical reactions LESSON 21 .. Sample ticket for control task #3…..
LESSON 22 ..22.1 Qualitative description of solutions…. 22.2 Solubilities of ionic compounds
LESSON 23 … 23.1 Molarity and Molality… 23.2 Mass percentage.
LESSON 24 .. 24.1 Theory of electrolytic dissociation … 24.2 Equations of stepwise dissociation
LESSON 25 .. 25.1 Ionic equations… 25.2 Examples of ionic equations…
LESSON 26 .. 26.1 When hydrolysis is impossible … 26.2 When hydrolysis is possible
LESSON 27 . Sample ticket for control task #4…..
(Хрусталёв В.В. и др. ВВЕДЕНИЕ В ОБЩУЮ ХИМИЮ …2014
The book provides an introduction into the General Chemistry. It is necessary for foreign students who are going to pass the Chemistry exam into the Belorussian State Medical Universityin English.
Actually, this book is a kind of compromise between translation of
chemistry text-books from Russian to English and popularizing material from original American sources. Authors hope that they combined the best Belorussian traditions with the best international points of view on Chemistry teaching (f.e. as the science of substances: their structure, their properties, and the reactions that change them — by Linus Pauling.[wiki15] or R.Chang (1998 — the study of matter and the changes it undergoes). Even though entrance exam into University usually requires some knowledge on very specific set of rules mostly based on simplifications and overestimations, this book not just deals with those “rules of thumb” but also provides modern explanations and interpretations.
Authors are looking forward to receive any feedback from readers and colleagues regarding style and content of the book.
1.1 PHYSICAL AND CHEMICAL PROPERTIES OF SUBSTANCES : WHAT IS THE DIFFERENCE? = What is chemistry? What is the subject you are starting (or, hopefully, continuing) to study? What is the difference between chemistry and physics?
Ancient have not differ physical and chemical phenomenon, but science in New Time have did it, include its bonds as special fields, as physical chemistry (from Boil to VantHof)*
Physical properties of a substance (for body — size, shape, mass) are: state of matter (solid, liquid, gas and plasma), density (the ratio between the mass and the volume), color (pink, blue, green, etc.), taste (sour, sweet, bitter, salty — more bio…)*, melting/freezing and boiling point (the temperature of crystallization and vaporization), solubility (the mass of substance that can be solved in water or other liquids at a given temperature). — See textbook/Shulyak…p.6-7, 13- f.e. for water, sugar
Chemical properties of a substance are described as its abilities to form other substances in different conditions.
In physical processes a substance changes at least one of its conditions: its volume, its shape, its position in the space, etc., while new substances are not formed. Phase transitions are also physical processes. There are several traditional examples of such physical processes: melting of the ice and crystallization of the water, boiling of the water and condensation of the vapor.
Chemical processes are described by chemical reactions. It is a process that leads to the transformation of one set of chemical substances to another. Substances from the first set are called “reactants”. Substances from the second set are called “products”.
For example, metalls (f.e. iron or copper) and oxygen or sulfur (the reactants) may react with each other and form compounds (as oxides or sulfides — the product).
Chemical equation for this reaction is as follows: Iron + Sulfur = Iron Sulfid, or Fe+S=FeS (others combinations: Copper Oxide, or CuO, non-metalls S + O2= SO2 (See 1.3. Symbols of chem.el. and the atomic theory of Dalton (1808), the law of definite proportions and certain relationship).
The rearrangement of atoms happens in chemical reactions, while atoms themselves stay the same.
Nuclear reactions are not chemical reactions, even though new substances are formed in them. In those nuclear reactions atoms of one chemical element turn to atoms of another chemical element. So, nuclear
reactions are studied in the course of physics and not chemistry.
There are several signs of chemical reactions: the thermal change (in some
cases the heat is produced in chemical reaction, in other cases the heat is
adsorbed from surroundings during the chemical reaction); the smell (for
example, hydrogen sulfide has a smell of rotten eggs); formation of a gas
without any characteristic smell; the change in color; precipitation (formation
of insoluble substance).
Finally, chemistry is the science on the interactions of matter with other
matter and with energy.
a. List several physical properties of water, sugar and salt.
b. A piece of (ice, salt) chalk has been dissolved in hydrochloric acid. Was it
a physical or chemical process?
c. What is the difference between reactants and products?
d. What are the signs of chemical reactions?How can you know that some
chemical reaction happened?
e. What is the subject ofchemistry?
*Odor is more biological property of substances.
+ History. Chemical reactions such as combustion in fire, fermentation and the reduction of ores to metals were known since antiquity. In East, Chinese philosophy the universe consists of heaven and earth (See Bible). Babylonian mythology, the cosmogony called Enûma Eliš (18-16th centuries BC), involves four gods that we might see as personified cosmic elements: sea, earth, sky, wind (m.b. independent of deities. Chinese had elements too, namely Fire, Water, Earth, Metal and Wood (word xing literally means something like «changing states of being», «permutations» or «metamorphoses of being» Sinologists wu xing is simply «the five changes») as part of 8-9=3×3 combination (like groups of Periodic system). Also, the Moon represents Yin (陰), and theSun 太陽 represents Yang (陽). Yin, Yang, and the five elements are associated with themes in the I Ching. In the bagua, metal is associated with the divination figure 兌 Duì (☱, the lake or marsh: 澤/泽 zé) and with 乾 Qián (☰, the sky or heavens… The five major planets are associated with and even named after the elements: Jupiter 木星 is Wood (木), Mars 火星 is Fire (火), Saturn 土星 is Earth (土),Venus 金星 is Metal (金), and Mercury 水星 is Water (水). Magic 3×3=2x2x2+1 In Taoism, qi functions similarly to pneuma in a prime matter (a basic principle of energetic transformation like en-trophy S=E/T=klnW) that accounts for both biological and inanimate phenomena.
West — Greek philosophers, Four-Element Theory of Empedocles or Five Elements (Plato and Aristotle, 360-350 BC) stating that any substance is composed of the ‘elements’ (stoicheia) in dialogue Timaeus, with the composition of inorganic and organic bodies and chemistry, particle of geometric shape: tetrahedron (fire), octahedron (air=T2), icosahedron (water=O2T=T5), cube (earth) and aether (these perfect polyhedra or Platonic solid composed of triangular faces the 30-60-90 and the 45-45-90 triangles, broken down into its component right-angled triangles, either isosceles or scalene, which could then be put together to form all of physical matter. The fifth element (i.e. Platonic solid) was the dodecahedron, whose faces are not triangular, most approximates a sphere, the shape into which God had formed the Universe. The creation of humans, soul, anatomy, perception, and transmigration of the soul follover Pythagoras. In the Middle Ages Alchemists attempted, in particular, to convert lead into gold, used reactions of metalls and alloys with sulfur. The production of new chemical substances, such as the synthesis of sulfuric and nitric acids attributed to Jābir ibn Hayyān, involved heating of minerals such as copper sulfate, alum and saltpeter. In the 17th century, Glauber produced hydrochloric acid and sodium sulfate by reacting sulfuric acid and sodium chloride, lead chamber in 1746 and the Leblanc process allow large-scale production of acid and sodium carbonate, chemical reactions became implemented into the industry, as the contact process in the 1880s, and the Haber process in 1909–1910 for ammonia synthesis.
From the 16th century, van Helmont, Boyle and Newton tried to establish theories of the experimentally observed chemical transformations. The phlogiston (Becher, 1667), «Oxy-gen» and «Hydro-gen» Lavoisier (1885) from the air and water replaced alchemic sulfur and as a fire-gen element released during combustion like electron in the modern explanation of Red-Ox reactions. The atomic theory and certain relationship of John Dalton, Joseph Proust and Gay-Lussac (1808) had developed the law of definite proportions, the concepts of stoichiometry and chemical equations.
Atom is the smallest piece of an element that maintains the identity of that
There are many substances that exist as two or more atoms connected
together. These combinations are called molecules. A molecule is the smallest
part of a substance that has the physical and chemical properties of that
Some elements exist in form of molecules. For example, hydrogen
and oxygen exist as two-atom molecules. Sulfur may exist as an eight-atom
molecule,S8, while phosphorus may exist as a four-atom molecule,P4. Other
elements, such as carbon (C), exist as individual atoms, ratherthan molecules.
In general, when nonmetal connects with other nonmetal, the compound
typicallyexists as molecule.
A chemical compoundis a chemical substance consisting of two or more
different chemical elements. Chemical compounds can be molecular
compounds held together by covalent bonds, salts held together by ionic bonds,
intermetallic compounds held together by metallic bonds, or complexes held
together by coordinate covalent bonds. Pure metals consist of atoms, positively
charged ions and freeelectrons (electron “gas”).
Pure chemical elements are not considered chemical compounds, even
if they consist of molecules that contain only multiple atoms of a single
element (such asH2,S8 etc.), which are called diatomic molecules or
Substances composed from atoms of the same element are historically
called “simple substances”. So, the term “pure chemical element” is a synonym
of the term “simple substance”.
Allotropy is the property of some chemical elements to exist in two or
more different forms, known as allotropes of these elements.
Coming back to carbon, the allotropes of that element include diamond
(where the carbon atoms are bonded together in a tetrahedral lattice
arrangement) and graphite (where the carbon atoms are bonded together in
sheets of a hexagonal lattice). The term allotropy is used for pure chemical elements (6)
only, and not for chemical compounds. Allotropy refers only to different
substances which exist as pure chemical elements within the same phase
(i. e. different solid, liquid or gas substances).
a. What is atom?
b. What is molecule?
c. Listsome substances which consist of molecules.
d. Listsome substances of nonmolecular structure.
e. Give a definition ofpure chemical element.
f. Give a definition of chemical compound.
g. Give a definition of simple substance.
h. What is allotropy?
i. Are oxygen (O2) and ozone (O3) allotropes?
The coefficient always goes before the compound or molecule, not after.
Coefficient shows the number of molecules, compounds or moles.
The subscript is written in small numbers by the bottom right corner of
the symbol. Subscript shows the number of certain atoms or groups of atoms in a given compound or molecule.
For example, 4H2O means four (coefficient) molecules of water. Water
consists of two (subscript) hydrogen atoms and a single oxygen atom.
Chemical formula 2Al2(SO4)3 means two (coefficient) compounds of
aluminum sulfate. Aluminum sulfate consists of two (subscript) aluminum ions and three
(subscript behind the brackets) sulfate anions. Each sulfate anion consists of a single sulfur
atom and four (subscript inside the brackets) oxygen atoms.
a. What does the coefficient show?
b. What does the subscript show?
c. Write chemical formulas of compounds made up from i) single iron
atom and three chlorine atoms; ii) two aluminum atoms and three oxygen
atoms; iii) single calcium atom, single carbon atom and three oxygen atoms.
d. Read the names of the following salts: CuSO4, CuSO3, CuS, Mg(NO3)2,
e. Read the names of the following hydroxides: Fe(OH)3, Ca(OH)2,
Ba(OH)2, KOH, NaOH.
f. Read the names of the following acids: H2SO4, H3PO4, HNO3, HNO2,
By the time of Antoine Lavoisier, a list of elements  correspond more closely to four of the states of matter: solid, liquid, gas and plasma. Law of conservation of mass continues to be conserved in isolated systems, even in modern physics. However, special relativity shows that due to mass–energy equivalence, whenever non-material «energy» (heat, light, kinetic energy) is removed from a non-isolated system, some mass will be lost with it. High energy losses result in loss of weighable amounts of mass, an important topic in nuclear chemistry.
Chemical reactions are governed by certain laws, which have become fundamental concepts in chemistry. Some Chemical law are: physical, for gas — Boyle’s law (1662, relating pressure and volume, PV=const), Charles’s law (1787, relating volume and temperature), Gay-Lussac’s law (1809, relating pressure and temperature), -> PV/T= const=nR=Rm/Mr, Avogadro’s law
(1803- law of partial pressures now known as Dalton’s law and researches on ethylene (olefiant gas) and methane (carburetted hydrogen) or by analysis of nitrous oxide (protoxide of azote) and nitrogen dioxide (deutoxide of azote), of Thomas Thomson.
led to atomic theory in chemistry — explanation of the law of multiple proportions to the idea that chemical combination consists in the interaction of atoms of definite and characteristic weight, the idea of atoms arose in his mind as a purely physical concept, forced upon him by study of the physical properties of the atmosphere and other gases (21/10/1803- 1805): Why does not water admit its bulk of every kind of gas alike? This question I have duly considered, and though I am not able to satisfy myself completely I am nearly persuaded that the circumstance depends on the weight and number of the ultimate particles of the several gases. The main points of Dalton’s atomic theory were:
- Elements are made of extremely small particles called atoms.
- Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.
- Atoms cannot be subdivided, created, or destroyed.
- Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
- In chemical reactions, atoms are combined, separated, or rearranged.
Dalton proposed an additional «rule of greatest simplicity» — When atoms combine in only one ratio, «..it must be presumed to be a binary one, unless some cause appear to the contrary» — caused him to assume that the formula for water was OH and ammonia was NH, quite different from our modern understanding (H2O, NH3).
Dalton proceeded to print his first published table of relative atomic weights. Six elements appear in this table, namely hydrogen, oxygen, nitrogen, carbon, sulfur, and phosphorus, with the atom of hydrogen conventionally assumed to weigh 1 (derived from analysis of water, ammonia, carbon dioxide, etc. by chemists of the time. chemical analysis particles of different weights, differentiated from Greeks, such asDemocritus and Lucretius.
The extension of this idea to substances in general necessarily led him to the law of multiple proportions: «The elements of oxygen may combine with a certain portion of nitrous gas or with twice that portion, but with no intermediate quantity» (11.1802-05)
He hypothesized the structure of compounds can be represented in whole number ratios. Many of the first compounds listed in the New System of Chemical Philosophy correspond to modern views, although many others do not.
- Law of definite composition, although in many systems (notably biomacromolecules and minerals) the ratios tend to require large numbers, and are frequently represented as a fraction.
- Law of multiple proportions
Modern science recognizes classes of elementary particles which have no substructure (or rather, particles that are not made of other particles, as e-) and composite particles having substructure (particles made of other particles, as nucleus). So:
Chemical element is a pure substance which is composed of a single type of atom, characterized by its particular number of protons in the nuclei of its atoms, known as the atomic number and represented by the symbol Z. The mass number is the sum of the number of protons and neutrons in a nucleus. Atoms of an element which have different mass numbers are known as isotopes. For example, all atoms with 6 protons in their nuclei are atoms of the chemical element carbon, but atoms of carbon may have mass numbers of 12, 13 or 14.
2.1 ATOMIC MASS AND MOLECULAR MASS
The atomic mass unit (u) is defined as one-twelfth of the mass of
a carbon-12 atom, an isotope of carbon that has six protons and six neutrons
in its nucleus. By this scale, the mass of a proton is 1.00728 u, the mass of
a neutron is 1.00866 u, and the mass of an electron is 0.000549 u (1835 times
lower than the mass of a proton, 1837 times lower than the mass of a neutron,
and 1821 times lower than the atomic mass unit). Approximate mass of
an atom may be estimated simply — by counting the total number of protons
and neutrons in the nucleus (that number is known as mass number).
The molecular mass is the sum of mass numbers of all the atoms in
2.2 THE MOLE AND MOLAR MASS
Chemists usually deal with millions, billions, and trillions of atoms
and molecules at a time. Mole is a unit of amount that relates quantities of
substances on a scale that is easy to interact with.
Chemistry uses a unit of amount called mole. A mole is a number of
things equal to the number of atoms in exactly 12 g of carbon-12. Experimental
measurements have determined that this number is very large:
1 mol = 6.02214179 × 1023
things (NA = Avogadro’s number)
In chemical calculations we usually use two digits after the point:
NA = 6.02·1023.
Once again, a mole means a number of things (6.02·1023
— a very big number), just like a dozen means a certain number of things (twelve).
Molar mass is the mass of 1 mole of a given substance. Molar mass is
measured in grams per mole.
The number of moles (n) is equal to the ratio between mass (m, measured
in grams) and molar mass (M, measured in gram/mole).
n = m / M
The numbers representing molar mass and molecular mass of the given
substance are exactly the same, even though molar mass is measured in grams
per mole, while molecular mass is measured in relative units.
a. Give the definition of the atomic mass unit.
b. What is the molecular mass?
c. How to determine a mass number of an atom?
d. How to determine molecular mass of a molecule?
e. What is Avogadro’s number?
f. What is the meaning of mole?
g. Give a definition of molar mass.
h. Why numbers representing molecular and molar masses are always
a. Calculate molecular masses of the following substances: HNO3 (Nitric
acid); KOH (Potassium hydroxide); ZnSO4 (Zinc sulfate); Fe2O3 (Iron (III)
oxide); MgCO3 (Magnesium carbonate); Mg3(PO4)…
Valence (it also may be written as valency or valence number) is
the number of chemical bonds a given atom has formed in a given molecule.
The number of bonds formed by a given element was originally thought
to be a fixed chemical property. In fact, in most cases this is not true.
For example, phosphorus often has a valence of three, but can also have other
Nowadays the definition of valence has become quite different from
the classic one. The current International Union of Pure and Applied Chemistry
(IUPAC) version of that term, adopted in 1994: “The maximum number of
univalent atoms (originally hydrogen or chlorine atoms) that may combine with
an atom of the element under consideration,or with a fragment, or for which
an atom of this element can be substituted”.
Valence is written in Roman numbers which has no sign (no plus or
minus). For example, the valence of hydrogen is always equal to “I”.
In the most of the substances the valence of oxygen is equal to “II”.
Each chemical bond is represented by a line in diagrams. Total number
of lines near the given atom is equal to its valence. For example, the valence
of sulfur in sulfur trioxide is equal to six…
3.2 CHEMICAL EQUATIONS AN
The law of conservation of matter says that matter cannot be created
or destroyed. In chemical equations, the number of atoms of each element
in the reactants must be the same as the number of atoms of eac
in the products. The mass of all the products should be the same as the mass
of all reactants (if we are not taking into account the famous Einstein’s
equation E = mc2 — the change in energy characteristic to every chemical
reaction should lead to some little changes in the mass of substances).
Coefficients are used to balance a chemical equation.
Coefficient is a number in a chemical equation indicating the number of
molecules (or moles) of the substance.
The sense of the balance is to make the
1. Balance the following chemical equations:
a) Na + H2O → NaOH + H
b) Al2O3 + HCl → AlCl3
c) P + O2 → P2O5
d) Mg + O2 → MgO
e) Cr2O3 + H2SO4 → Cr2
f) Fe(OH)3 + H3PO4 → FePO
g) Zn(OH)2 + HNO3 → Zn(NO
h) Fe2O3 + Al → Fe + Al
2. Finish chemical equations and then balance them:
a) Zn + Cl2 → ?
b) Fe + ? → FeCl3
c) Ca + HCl → ? + H2↑
d) Mg + HCl → MgCl2 + ?
e) FeO + HCl → ? + H2O
f) CuO + HNO3 → ? + H
g) Al(OH)3 → ? + H2O
h) CaCO3 → ? + CO2↑
Figure 3.1. Structures of SO3 and H3PO4
CHEMICAL EQUATIONS AND THEIR BALANCING
The law of conservation of matter says that matter cannot be created or destroyed.
In chemical equations, the number of atoms of each element in the reactants must be the same as the
number of atoms of each element in the products. The mass of all the products should be the same as the mass
of all reactants (if we are not taking into account the famous Einstein’s
the change in energy characteristic to every chemical
to some little changes in the mass of substances).
Coefficients are used to balance a chemical equation.
Coefficient is a number in a chemical equation indicating the number of
of the substance.
The sense of the balance is to make the law of conservation of matter
Balance the following chemical equations:
→ NaOH + H2↑
3 + H2O
2(SO4)3 + H2O
→ FePO4 + H2O
→ Zn(NO3)2 + H2O
→ Fe + Al2O3
Finish chemical equations and then balance them:
4.1 CALCULATIONS USING CHEMICAL EQUATIONS
Here is the example of simplest chemical calculation. What is the mass of
phosphoric acid (H3PO4) required for the complete neutralization of 100 g of calcium hydroxide (Ca(OH)2)?
There are at least two ways to make a calculation using chemical equation.
The first way allows calculations without direct referring to the quantity
of matter (number of moles). The given mass of one of the reactants or
products may be written just upon that chemical substance. The molar mass may be written under that substance. One should refer to the Periodic table to find atomic masses in case if he or she does not remember them by heart.
The molar mass of the substance with unknown mass may also be written under the formula. Remember that molar masses of substances should be multiplied by coefficients from the balanced equation. To find out the unknown mass of the substance A (H3PO4 in our case) one has to multiply the known mass of another substance B (Ca(OH)2) by the molar mass of the unknown substance A and by the coefficient before that substance. Then the result of the multiplication has to be divided by the molar mass of the substance B (previously multiplied by the coefficient before the substance B).
100 g X g?
3Ca(OH)2 + 2H3PO4 → Ca3(PO4)2 + 6H2O
3 ∙ (40 + 2 · (16 + 1)) = = 2 · 98 = 196 g
2 · ((3 · 1) + 31 + (4 · 16)) =
= 3 · 74 = 222 g
X = m(H3PO4) = (100 · 196) / 222 = 88.3 g
The second way to find the same answer is based on the quantity of
matter calculation. At first one has to divide the mass of substance A by its
molar mass to find out the number of moles. Then one has to multiply
the number of moles of substance A by the coefficient before the substance B
and divide the result by the coefficient before the substance A. The result
of this calculation is the number of moles of substance B. The last step is to
multiply the chemical quantity of substance B by its molar mass.
1) n(Ca(OH)2) = m(Ca(OH)2) / M (Ca(OH)2) = 100 / 74 = 1.35 mol
2) n(H3PO4) = (1.35 · 2) / 3 = 0.9 mol
3) m(H3PO4) = n(H3PO4) · M(H3PO4) = 0.9 · 98 = 88.2 g
In some cases it is easier to use the first method (when you jus
The standard presentation of the chemical elements is in the periodic table, which orders elements by atomic number. The periodic table is arranged in groups, or columns, and periods, or rows. for identifying periodic trends.
A compound is a pure chemical substance composed of more than one element. The properties of a compound bear little similarity to those of its elements.
The standard nomenclature of compounds is set by the International Union of Pure and Applied Chemistry(IUPAC). Organic compounds are named according to the organic nomenclature system. Inorganic compounds are named according to the inorganic nomenclature system.
Molecule is the smallest indivisible portion of a pure chemical substance that has its unique set of chemical properties, that is, its potential to undergo a certain set of chemical reactions with other substances. However, this definition only works well for substances that are composed of molecules, which is not true of many substances (see below).
Molecules are typically a set of atoms bound together by covalent bonds, such that the structure is electrically neutral and all valence electrons are paired with other electrons either in bonds or in lone pairs.
Thus, molecules exist as electrically neutral units, unlike ions. When this rule is broken, giving the «molecule» a charge, the result is sometimes named a molecular ion or a polyatomic ion. However, the discrete and separate nature of the molecular concept usually requires that molecular ions be present only in well-separated form, such as a directed beam in a vacuum in a mass spectrometer. Charged polyatomic collections residing in solids (for example, common sulfate or nitrate ions) are generally not considered «molecules» in chemistry.
When an atom gains an electron and thus has more electrons than protons, the atom is a negatively charged ion or anion. Cations and anions can form a crystalline lattice of neutral salts, such as the Na+ and Cl− ions forming sodium chloride, or NaCl. Examples of polyatomic ions that do not split up duringacid-base reactions are hydroxide (OH−) and phosphate (PO43−).
Plasma is composed of gaseous matter that has been completely ionized, usually through high temperature.
Redox (reduction-oxidation) reactions include all chemical reactions in which atoms have their oxidation state changed by either gaining electrons (reduction) or losing electrons (oxidation). Substances that have the ability to oxidize other substances are said to be oxidative and are known as oxidizing agents, oxidants or oxidizers. An oxidant removes electrons from another substance. Similarly, substances that have the ability to reduce other substances are said to be reductive and are known as reducing agents, reductants, or reducers.
A reductant transfers electrons to another substance, and is thus oxidized itself. And because it «donates» electrons it is also called an electron donor. Oxidation and reduction properly refer to a change in oxidation number—the actual transfer of electrons may never occur. Thus, oxidation is better defined as an increase in oxidation number, and reduction as a decrease in oxidation number.
An oxide is a binary chemical compound that contains oxygen in the oxidation state of –2 and other chemical element.
Different oxides of the same element are distinguished by Roman numerals denoting their oxidation number:
iron (II) oxide (FeO) versus iron (III) oxide (Fe2O3).
Oxides may be produced either by the way of combustion (of both pure
chemical elements and compounds), or by decomposition of certain acids,
bases and salts.
C + O2 → CO2, 4Al + 3O2 → 2Al2O3,
2H2S + 3O2 → 2H2O + 2SO2, 2CuS + 3O2→ 2CuO + 2SO2
H2SiO3 → H2O + SiO2, Ca(OH)2 → CaO + H2O, ZnCO3 → ZnO + CO2
Due to its electronegativity, oxygen forms stable chemical bonds with
almost all elements to give the corresponding oxides. Noble metals (such as
gold or platinum) are prized because they resist direct chemical combination
with oxygen, and substances like gold (III) oxide must be generated by indirect
Oxides of most metals adopt polymeric structures with M-O-M crosslinks.
Because these crosslinks are strong, the solids tend to be insoluble in solvents,
though they are attacked by acids and bases. The formulas are often deceptively
simple, while many of them are nonstoichiometric compounds.
Fe3O4 = FeO · Fe2O3, 3CrO2 = Cr2O3 · CrO3, 4VO2 = V2O3 · V2O5
Metal oxides are substances of ionic crystal structure, while the most of
nonmetal oxides are molecules. For example, carbon dioxide (CO2) and carbon
monoxide (CO) are molecular oxides. However, silicon oxide (SiO2) is
a substance with atomic crystal structure. Phosphorus pentoxide is a complex
molecular oxide with a deceptive name, the formula being P4O10 (P2O5·P2O5
Some polymeric oxides when heated depolymerize to give molecules, examples
being selenium dioxide (SeO2) and sulfur trioxide (SO3).
Oxides can be attacked by acids and bases. Those attacked only by acids
are basic oxides.
BaO + 2HCl → BaCl2 + H2O
Those attacked only by bases are acidic oxides.
P2O5 + 6KOH → 2K3PO4 + 3H2O
Oxides that react with both acids and bases are amphoteric.
Al2O3 + 6HCl →2AlCl3 + 3H2O
Al2O3 + 2NaOH → 2NaAlO2 + H2O
In water solution complex salt is formed.
Al2O3 + 2NaOH + 3H2O → 2Na[Al(OH)4]
Metals tend to form basic oxides, non-metals tend to form acidic oxides,
and amphoteric oxides are formed by elements near the boundary between
metals and non-metals (metalloids) and by elements from d-block.
Behavior of oxides for d-elements depends on their oxidation state.
Basic oxides: MnO and CrO
Amphoteric oxides: MnO2 and Cr2O3
Acidic oxides: Mn2O7 and CrO3
16.2 BASIC AND ACIDIC ANHYDRIDES
Oxides of more electropositive elements are called “basic anhydrides”.
Exposed to water, oxides of alkali and earth-alkali metals form basic
hydroxides. For example, sodium oxide is basic — when hydrated, it forms
Na2O + H2O → 2NaOH
Oxides of other metals cannot react with water, though their insoluble
hydroxides can be produced in reactions between their soluble salts and alkali.
CuCl2 + 2NaOH → Cu(OH)2↓ + 2NaCl
Oxides of more electronegative elements are called “acid anhydrides”;
adding water, they form oxoacids (oxygen containing acids).
N2O5 + H2O → 2HNO3
Basic oxides react with acidic oxides and form salts.
3CaO + P2O5 → Ca3(PO4)2
Less volatile acidic oxides react with salts of more volatile oxides.
Volatility is the tendency of a substance to vaporize. In other words, the more
volatile oxide simply flies away and cannot participate in the reverse reaction
CaCO3 + SiO2 → CaSiO3 + CO2↑
Some oxides do not show behavior as either acid or basic anhydrides
(for example: CO, N2O, NO). They are known under the name “neutral
oxides” (do not form salts).
- Give examples of basic, acidic and amphoteric oxides.
- How do the properties of oxides change if we go from left to right in a period of the Periodic table?
- How do the properties of oxides change if we go from top to bottom in a group of the Periodic table?
- What is anhydride?
- What is volatility?
- Write down formulas of the following compounds: potassium oxide;
sodium oxide; phosphorus oxide (III); phosphorus oxide (V); nitrogen oxide (III); nitrogen oxide (V);
iron oxide (II); iron oxide (III); manganese oxide (II); manganese oxide (IV); manganese oxide (VII).
- What are the formulas of oxides corresponding to the following hydroxides: Mg(OH)2; LiOH; Fe(OH)2; Fe(OH)3; Cu(OH)2; Cr(OH)3?
- Write down reactions between i) calcium oxide and sulfur oxide (VI); ii) iron oxide (III) and phosphorus oxide (V); iii) zinc oxide and nitrogen oxide
(V); iv) sodium oxide and carbon dioxide.
- Write down reactions between i) calcium oxide and sulfuric acid; ii) iron oxide (II) and nitric acid; iii) chrom oxide (III) and sulfuric acid;
- iv) aluminum oxide and phosphoric acid.
- Write down reactions between i) sodium hydroxide and carbon dioxide; ii) potassium hydroxide and nitrogen oxide (III); iii) barium hydroxide and
hydrochloric acid; iv) calcium oxide and hydrogen sulfide.
- Imagine how can the following oxides be produced: CaO; Al2O3; CO2; SiO2?
- Calculate the mass of iron in 500 g of i) FeO; ii) Fe2O3; iii) Fe3O4.
The Bronsted-Lowry theory defines bases as proton (hydrogen ion)
acceptors, while the more general Lewis theory defines bases as electron pair
donors, allowing other Lewis acids than protons to be included. The oldest
Arrhenius theory defines bases as substances which produce hydroxide anions
(OH−) in water solutions. By altering the autoionization equilibrium (H2O = H++ OH–),
bases give solutions with a hydrogen ion activity lower than that of pure water (pH > 7.0).
Bases can be thought of as the chemical opposite of acids. A reaction
between an acid and base is called neutralization — aqueous solutions of bases
react with aqueous solutions of acids to produce water and salts.
very weak acids in an acid-base reaction. Common examples of strong bases
are the hydroxides of alkali metals and alkaline earth metals like NaOH and
Ca(OH)2. Here is a list of strong bases:
– Lithium hydroxide (LiOH) – Sodium hydroxide (NaOH) – Potassium hydroxide (KOH)
– Rubidium hydroxide (RbOH) – Cesium hydroxide (CsOH)
– Calcium hydroxide (Ca(OH)2) – Strontium hydroxide (Sr(OH)2) – Barium hydroxide (Ba(OH)2)
Alkali is a base that dissolves in water.
The word “alkali” is derived from Arabic al qalīy (or alkali), meaning the calcined ashes,
referring to the original source of alkaline substances. A water-extract of burned plant ashes, called
potash and composed mostly of potassium carbonate (K2CO3), was mildly basic (K2CO3 + H2O = KHCO3 +
+ KOH). After heating this substance with calcium hydroxide (slaked lime),
a far more strongly basic substance known as caustic potash (potassium
hydroxide) was produced (K2CO3 + Ca(OH)2 → CaCO3↓ + 2KOH). Caustic
potash was traditionally used in conjunction with animal fats to produce soft
soaps, one of the caustic processes that rendered soaps from fats in the process
of saponification, known since antiquity. Plant potash lent the name to the
element potassium, which was first derived from caustic potash, and also gave
potassium its chemical symbol K (Kalium), which ultimately derives from al kali.
There are various definitions for alkali. Alkali is often defined as a subset
of base. Two examples of alkali definitions are given below.
– A hydroxide of alkali metal or alkaline earth metal (this includes
Mg(OH)2 but excludes NH4OH).
– Any base that is water-soluble and forms hydroxide ions or the solution
of a base in water (this excludes Mg(OH)2 but includes NH4OH).
Magnesium hydroxide is an example of an atypical alkali since it has low
solubility in water, while the dissolved portion is considered a strong base due
to complete dissociation of its ions.
Soluble hydroxides of alkali metals and alkaline earth metals (i. e. alkalis)
are often called “alkali salts”. There are commonly used names of some alkali
salts. “Caustic soda” is sodium hydroxide (NaOH). “Caustic potash” is
potassium hydroxide (KOH). “Limewater” is calcium hydroxide (Ca(OH)2).
Lime is a general term for calcium-containing inorganic materials, in which carbonates,
oxides and hydroxides predominate. The word “lime” originates with its
earliest use as building mortar and has the sense of “sticking or adhering”.
Alkali can be produced in the reaction between active metals (or their
oxides) and water. Active metals are: Cs, Rb, K, Na, Li, Ba, Sr, Ca.
Magnesium, aluminium and zinc can react with water, but the reaction is
usually very slow unless the metal samples are specially prepared to remove the
surface layer of oxide which protects the rest of the metal.
2Na + 2H2O →2NaOH + H2↑, Na2O + H2O → 2NaOH
Insoluble bases can be produced in reactions (metathesis) between soluble salts and alkalis.
ZnSO4 + 2KOH → Zn(OH)2↓ + K2SO4
In case of the excess of alkali complex soluble salts can be formed.
ZnSO4 + 4KOH → K2[Zn(OH)4] + K2SO4
Alkalis react with both acids and acidic oxides.
The products of those reactions are salts and water.
2NaOH + H2SO4 →Na2SO4 + 2H2O
2NaOH + SO3 → Na2SO4 +H2O
Alkalis react with certain salts (in case if one of the products of the
reaction is insoluble or volatile).
2KOH + FeCl2 →2KCl + Fe(OH)2↓
KOH + NH4Cl → KCl + NH3↑ + H2O
All bases, except NaOH and KOH, are decomposed at high temperatures.
The products of such decomposition are oxide and water.
Ca(OH)2 → CaO + H2O
Amphoteric hydroxides react with both acids and alkalis.
Al(OH)3 + 3HCl →AlCl3 + 3H2O
Al(OH)3 + NaOH → NaAlO2 + 2H2O
In water solution complex salt is formed.
Al(OH)3 + NaOH → Na[Al(OH)4]
- Give a definition of base.
- Give a definition of alkali.
- Give a definition of hydroxide.
- What metals are active?
- Imagine how can the following substances be produced: i) KOH; ii) LiOH; iii) Fe(OH)3; iv) Mg(OH)2; v) Zn(OH)2.
- How can Zn(OH)2 be produced from the set of the following substances: Na; H2SO4; ZnO; H2O?
HCl; Fe2O3; H2O?
- What substances react with KOH? Write down equations of possible reactions.
- Al2O3; Ca(HCO3)2; HCl; CO2; H2SO4; KNO3; Zn(OH)2; NaHCO3;K2CO3; SO3; H3PO4; NH4Cl.
- What substances react with Zn(OH)2? Write down equations of possible
reactions. KOH; H2O; H2SO4; CaCl2; NaCl; HNO3; CO2; SO2; SO3; HCl;NaOH; H3PO4.
- Finish chemical reactions: NaOH + (i) P2O5 →, ii. + Cr2(SO4)3→
iii. NaOH + CO2 →, NaOH + NH4NO3 →, NaOH + Cr(OH)3 →
- Fe(OH)3 + H2SO4 →, vii. Fe(OH)2 + HNO3 →
viii. Fe(OH)3 →
Ca(OH)2 + H3PO4 →, KOH + MgCl2 → , KOH + N2O5 →
xii. Al2(SO4)3 + KOH → xiii. Sr(OH)2 + H2SO4 →
xiv. Cr(OH)3 + HCl →, Al(OH)3 + NaOH →
An Arrhenius acid is a compound that increases the H+ ion concentration
in aqueous solution. The H+ ion is just a bare proton, and it is rather clear that
bare protons are not floating around in an aqueous solution. Instead, chemistry
has defined the hydronium ion (H3O+) as the actual chemical species that
represents an H+ ion. Classic Arrhenius acids can be considered ionic compounds
in which H+ is the cation.
If an acid is composed of only hydrogen and one other element, the name
is hydro- + the stem of the other element + -ic acid. For example,
the compound HCl(aq) is hydrochloric acid, while H2S(aq) is hydrosulfuric
acid. If these acids were not dissolved in water, the compounds would be called
hydrogen chloride and hydrogen sulfide, respectively. Both of these substances
are well known as molecular compounds; when dissolved in water, however,
they are treated as acids.
Acidity and basicity
A substance can often be classified as an acid or a base. There are several different theories which explain acid-base behavior. The simplest is Arrhenius theory, which states than an acid is a substance that produces hydronium ions when it is dissolved in water, and a base is one that produces hydroxide ions when dissolved in water. According to Brønsted–Lowry acid–base theory, acids are substances that donate a positive hydrogen ion to another substance in a chemical reaction; by extension, a base is the substance which receives that hydrogen ion (as NH3).
A third common theory is Lewis acid-base theory, which is based on the formation of new chemical bonds. Lewis theory explains that an acid is a substance which is capable of accepting a pair of electrons from another substance during the process of bond formation, while a base is a substance which can provide a pair of electrons to form a new bond. According to this theory, the crucial things being exchanged are charges. There are several other ways in which a substance may be classified as an acid or a base, as is evident in the history of this concept.
Acid strength is commonly measured by two methods. One measurement, based on the Arrhenius definition of acidity, is pH, which is a measurement of the hydronium ion concentration in a solution, as expressed on a negative logarithmic scale. Thus, solutions that have a low pH have a high hydronium ion concentration, and can be said to be more acidic. The other measurement, based on the Brønsted–Lowry definition, is the acid dissociation constant (Ka), which measures the relative ability of a substance to act as an acid under the Brønsted–Lowry definition of an acid. That is, substances with a higher Ka are more likely to donate hydrogen ions in chemical reactions than those with lower Ka values.
If a compound is composed of hydrogen ions and a polyatomic anion, then
the name of the acid is derived from the stem of the polyatomic ion’s name.
Typically, if the anion name ends in -ate, the name of the acid is the stem of
the anion name plus -ic acid; if the related anion’s name ends in -ite, the name
of the corresponding acid is the stem of the anion name plus -ous acid.
Acids may be produced in reactions between acidic oxides and water.
SO3 + H2O → H2SO4
Less volatile acids react with salts of more volatile acids.
FeS + 2HCl→ FeCl2 + H2S↑
Formulas and names of some acids
CH3COOH acetic acid
HCl hydrochloric acid
HClO3 chloric acid
HClO4 perchloric acid
HBr hydrobromic acid
HI hydroiodic acid
HF hydrofluoric acid
HNO2 nitrous acid
HNO3 nitric acid
H2C2O4 oxalic acid
H3PO4 phosphoric acid
H2SO4 sulfuric acid
H2SO3 sulfurous acid
Acids have some properties in common. They react with metals situated before hydrogen
in the electrochemical series of metals togive off H2 gas.
Zn + H2SO4 → ZnSO4 + H2↑
Electrochemical series of metals:
Li > K > Sr > Ca > Na > Mg > Al > Zn > Cr > Fe > Cd > Co > Ni > Sn >
> Pb >H > Cu > Ag > Hg > Pt > Au
Acids react with basic and amphoteric oxides.
H2SO4 + CuO → CuSO4 + H2O, 6HNO3 + Al2O3 → 2Al(NO3)3 + 3H2O
Acids react with salts in case if insoluble substances or gases are formed.
For example, they react with carbonate and hydrogen carbonate salts to give off
CO2 gas. CaCO3 + 2HCl → CaCl2 + CO2↑ + H2O
Ca(HCO3)2 + 2HCl → CaCl2 + 2CO2↑ + 2H2O
Acids that are ingested typically have a sour, sharp taste. The name
acid comes from the Latin word acidus, meaning “sour”.
18.2 NEUTRALIZATION REACTION
Acids and bases react with each other to make water and an ionic
compound called a salt. A salt, in chemistry, is any ionic compound made by
combining an acid with a base. A reaction between an acid and a base is called
a neutralization reaction and can be represented as follows:
acid + base → H2O + salt
The stoichiometry of the balanced chemical equation depends on
the number of H+ ions in the acid and the number of OH− ions in the base.
- Give the formula for each acid.
- perchloric acid
- hydriodic acid
- hydrosulfuric acid
- phosphorous acid
- Name each acid.
- Name some properties that acids have in common.
- Imagine how the following acids can be produced: H3PO4; HNO3;
HNO2; HCl; H2SO4; H2SO3; H2S.
2. What metals among the given list react with hydrochloric acid: Li; Ba;
Cu; Mg; Al; Au; Ag?
- What substances react with sulfuric acid? Write down equations of
possible reactions. CuCl2; Fe(OH)3; ZnO; HCl; Al(OH)3; SiO2; Pb(NO3)2;
KOH; BaCl2; CuO; Mg(OH)2; Zn.
- What substances react with nitric acid? Write down equations of
possible reactions. AgCl; Fe(OH)2; CuO; HBr; Zn(OH)2; CO2; NaNO3; KCl;
Ba(OH)2; H3PO4; Sr(OH)2; Cr.
Finish chemical reactions:
- P2O5 + H2O →
- HCl + Mg(OH)2 → iii. Al(OH)3 + HCl →
- Fe + HCl→
- SO2 + H2O →
- Fe2O3 + H2SO4 →
vii. Fe + HI →
viii. N2O5 + H2O →
- CaO + H3PO4 →
- LiCl + H3PO4 →
- What is the mass of ZnSO4 which was produced from 9.8 g of sulfuric
acid and 8.1 g of ZnO?
- What is the mass of K3PO4 which was produced from 49 g of
phosphoric acid and 80 g of KOH?
Salts are ionic compounds that result from the neutralization reaction of
an acid and a base. They are composed of such numbers of cations (positively
charged ions) and anions (negative ions) that the product is electrically neutral
(without a net charge). These component ions can be inorganic such as chloride
(Cl−), as well as organic such as acetate (CH3COO−) and monoatomic ions such
as fluoride (F−), as well as polyatomic ions such as sulfate (SO42-).
There are several varieties of salts. Salts that dissociate to produce
hydroxide ions when dissolved in water are basic salts (Fe(OH)2Cl) and salts
that dissociate to produce hydronium ions in water are acid salts (NaHSO4).
are those that are neither acid nor basic salts. Zwitterions contain
an anionic center and a cationic center in the same molecule but are not
considered to be salts. Examples include amino acids, many metabolites,
peptides, and proteins.
Acidic salts can be produced from neutral salts after the addition of acid.
CaSO4 + H2SO4 → Ca(HSO4)2
Basic salts can be produced from neutral salts after the addition of base.
CaSO4 + Ca(OH)2 → (CaOH)2SO4
19.2 SOLUBILITY CHART OF SALTS
Many ionic compounds can be dissolved in water. The exact combination
of ions involved makes each compound have a unique solubility in any solvent.
The solubility is dependent upon how well each ion interacts with the solvent,
so there are certain patterns.
For example, all salts of sodium, potassium and ammonium are soluble in
water, as are all nitrates and many sulfate salts except barium sulfate, calcium
sulfate (sparingly soluble) and lead (II) sulfate.
However, ions that bind tightly to each other and form highly stable
lattices are less soluble, because it is harder for these structures to break apart
for the compounds to dissolve. For example, most carbonate salts are not
soluble in water, such as lead carbonate and barium carbonate. Soluble
carbonate salts are: sodium carbonate, potassium carbonate and ammonium
Salts are formed by a chemical reaction between:
– A base and an acid, e.g., NH4OH + HCl → NH4Cl + H2O
– A metal and an acid, e.g., Mg + H2SO4 →MgSO4 + H2↑
– A metal and a non-metal, e.g., Ca + Cl2 → CaCl2
– A base and an acid anhydride, e.g., 2NaOH + CO2 → Na2CO3 + H2O
– An acid and a basic anhydride, e.g., 2HNO3 + Na2O → 2NaNO3 + H2O
– Salts can also be formed if solutions of different salts are mixed with
each other or with acid or alkali solutions. Their ions recombine, and in some
cases new salt (base or even acid) precipitates (see: the solubility chart below):
Pb(NO3)2 + Na2SO4 → PbSO4↓ + 2NaNO3. The same thing can be said about
reactions in which gases are formed: NH4Cl + KOH → KCl + NH3↑ + H2O
– Metal is able to substitute another metal in a salt in case if it is situated
before the second one in the reactivity series: CuSO4 + Zn → ZnSO4 + Cu
Reactivity series of metals
Active metals – those which react with water and acids
Cs Rb K Na Li Ba Sr Ca
Metals which react with acids and produce salts and H2
Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb
Metals which react with strong oxidizing acids only and don’t produce H2
Sb Bi Cu Hg Ag Au Pt
The reactivity series is sometimes quoted in the strict reverse order of
standard electrode potentials, when it is also known as the “electrochemical
Li > K > Sr > Ca > Na > Mg > Al > Zn > Cr > Fe > Cd > Co > Ni > Sn > Pb >
>H > Cu > Ag > Hg > Pt > Au
The positions of lithium and sodium are changed on such a series; gold
and platinum are also inverted, although this has little practical significance as
both metals are highly unreactive. Hydrogen is included in standard electrode
potentials order because the power of a reducing agent is measured relatively to
the standard hydrogen electrode.
- Give a definition of the term “salt”.
- What is the difference between normal, acidic and basic salts?
- How can the reactivity series of metals be used?
- How can the solubility chart be used?
- Write down the formulas of the following compounds: iron (III) sulfate;
sodium sulfate; barium dihydrogen phosphate; magnesium hydrogen carbonate;
magnesium hydroxy chloride; potassium sulfide; potassium hydroxy sulfide;
potassium hydroxy sulfate; iron (III) dihydroxy chloride; calcium hydrogen
- What substances react with AgNO3? Write down equations of possible
reactions. HCl; FeSO4; CaCl2; BaCl2; H2SO4; NaI; KBr; K3PO4; Na2CO3;
- How can the following salts be produced: KCl; Fe2(SO4)3; Zn(NO3)2?
- Which salts react with each other: i) CaCO3 + KCl; ii) MgCl2 +Na2CO3; iii) NaCl + K2CO3; iv) BaCl2 + K2SO4.
- Finish equations of chemical reactions: BaCl2 + Na3PO4 →, +Fe2(SO4)3 →
iii. H3PO4 + Mg(OH)2 →,HgSO4 + Zn →, BaCO3 + HCl →, AlOHSO4 + H2SO4 →
vii. Al + H2SO4 → viii. AlCl3 + NaOH → Na[Al(OH)4] + HCl →
Ca(OH)2 + H3PO4 →
What is the mass of silver chloride produced in the reaction between
5.85 g of sodium chloride and 33.8 g of silver nitrate?
What is the mass of barium sulfate produced in the reaction between
9.8 g of sulfuric acid and 41.6 g of barium chloride?
- What kind of salt is produced in the reaction between 10.5 g of calcium
hydroxide and 13.7 g of nitric acid?
- What kind of salt is produced in the reaction between 12.3 g of sodium
hydroxide and 26 g of copper chloride?
LESSON 20 20.1 CLASSIC CHAINS OF CHEMICAL REACTIONS
Chain of chemical reactions is one of the most commonly used types of
tasks in chemistry. In the classic type of that kind of task student has to write
down all the reactions from each chain. Each next substance must be somehow
produced from the previous substance. Student has a kind of freedom to choose
additional reactants and conditions to make reactions possible. In case if
the one-step reaction is impossible (for example, it is impossible to produce
insoluble base directly from the metal), two or even more reactions should be
written to complete one step in the chain.
- CaCO3 → CaO → Ca(OH)2 → Ca(HCO3)2 → CaCO3
- Mg- MgO ®Mg(OH)2 ® MgCl2 ®Mg(NO3)2
®Ba(OH)2 ® BaOHCl ® BaCl2 ®BaCO3
- K® KOH ®KHSO4 ® K2SO4 ® KCl
- Al → Al2(SO4)3 → Na[Al(OH)4] → Al(NO3)3 → Al(OH)3
- NaHCO3 → CO2 → CaCO3 → Ca(HCO3)2 → CaO
- Cu → CuS → CuSO4→ CuCl2 → Cu(NO3)2
- CuSO 4 ® SO3 ® NaHSO4 ® Na2SO4 ® NaHSO4
- Fe ® FeCl2 ® Fe(NO3)2 ® FeSO4 ® Fe(OH)2
- K → K2O → KOH → KCl → KNO3
20.2MODERN CHAINS OF CHEMICAL REACTIONS
Modern type of the chain of chemical reactions includes two tasks in one.
At first, student must write down all the reactions. All the reactants, conditions
and by-products are given (if they are not clearly obvious), while main products
are hidden behind the letters. At second, student must find out molar mass
example, for substance hidden behind the letter “B” and substance hidden
behind the letter “D”).
- Calculate the sum of molar masses for compounds A, B, C and D from
the chain of chemical reactions.
H2→A→ B → C → D
- Calculate the sum of molar masses for compounds B and D from
the chain of chemical reactions.
Na → A → B → C → D
- Calculate the sum of molar masses for compounds B, C and D from
the chain of chemical reactions.
Fe(OH)3 → A → B → C → D
- Calculate the molar mass of compound D from the chain of chemical
Al, → A →B, → C → D
- Calculate the sum of molar masses for compounds C and D from
the chain of chemical reactions.
Ca → A → B → C → D
Sample ticket for control task #3 on main types of inorganic chemical compounds
- Write down equations of chemical reactions between the following substances
in case if they are possible CaO + Na2O®
Na2O + NO®
CaO + N2O5®
K2O + Al2O3®
BaO + H2O®
N2O5 + H2O®
SiO2 + H2O®
Al2O3 + H2O®
KAlO2 + HCl®
Cu(OH)2 + Na2SO4®
BaCO3 + KOH®
- Ca ®Ca(OH)2®CaOHCl®CaCl2®CaCO3
- Modern chains of chemical reactions
Calculate the sum of molecular masses for compounds A, B, C and D
from chains of chemical reactions:
- C → A → B→C→D
- Ca3(PO4)2 → A → B ( )→C ( )→D
22.1 QUALITATIVE DESCRIPTION OF SOLUTIONS
The major component of a solution is called the solvent. The minor component
of a solution is called the solute. By major and minor we mean
whichever component has the greater presence by mass or by moles.
Sometimes this becomes confusing, especially with substances with very
different molar masses. For example, if the mass percentage of ethanol solution
in water is equal to 70%, the mole percentage of ethanol is still equal to 47.6%
(because molar mass of water is lower than that for ethanol). Obviously, such
expression as “98% ethanol” widely used for a solution in which there are 98 %
of ethanol and just 2% of water is not correct, because ethanol is still
mistakenly considered to be solute and not solvent.
Salt water is a solution of solid NaCl in liquid water; soda water is
a solution of gaseous CO2 in liquid water, while air is a solution of a gaseous
solute (O2) in a gaseous solvent (N2). In all cases, however, the overall phase of
the solution is the same phase as the solvent.
One important concept of solutions is in defining how much solute is
dissolved in a given amount of solvent.
Dilute describes a solution that has very little solute, while
concentrated describes a solution that has a lot of
solute. One problem is that these terms are qualitative; they describe more or
less but not exactly how much.
22.2 SOLUBILITIES OF IONIC COMPOUNDS
In most cases, only a certain maximum amount of solute can be dissolved
in a given amount of solvent. This maximum amount is called the
solubility of the solute. It is usually expressed in terms of the amount of solute that can
dissolve in 100 g of the solvent at a given temperature.
When the maximum amount of solute has been dissolved in a given amount
of solvent, we say that the solution is saturated with solute. When less than
the maximum amount of solute is dissolved in a given amount of solute, the solution
is unsaturated. A solution of 0.00019 g of AgCl per 100 g of H2O may be
saturated, but with so little solute dissolved, it is also rather dilute. A solution of
36.1 g of NaCl in 100 g of H2O is also saturated but rather concentrated.
Solubilities of Some Ionic Compounds
Solute Solubility (g per 100 g of H2O at 25 °C)
AgCl 0.00019 CaCO3 0.0006 NaCl 36.1 KBr 70.7 NaNO3 94.6
In some circumstances, it is possible to dissolve more than the maximum
amount of a solute in a solution. Usually, this happens by heating the solvent,
dissolving more solute than would normally dissolve at regular temperatures,
and letting the solution cool down slowly and carefully. Such solutions are
called supersaturated solutions and they are not stable; given an opportunity
(such as dropping a crystal of solute in the solution), the excess solute will
precipitate from the solution.
It should be obvious that some solutes dissolve in certain solvents but not
in others. NaCl, for example, dissolves in water but not in vegetable oil.
Beeswax dissolves in liquid hexane but not in water.
From experimental studies, it has been determined that if molecules of
a solute experience the same intermolecular forces that the solvent does,
the solute will likely dissolve in that solvent. So, NaCl — a very polar
substance because it is composed of ions — dissolves in water, which is very
polar, but not in oil, which is generally nonpolar. Nonpolar wax dissolves in
nonpolar hexane but not in polar water. This concept leads to the general
ancient alchemic rule that “like dissolves like” for predicting whether a solute
is soluble in a given solvent. However, this is a general rule, not an absolute
statement, so it must be applied with care.
- Define solute and solvent.
- Define saturated, unsaturated, and supersaturated solutions.
- Differentiate between polar and nonpolar solvents.
- Which solvent is Br2 more likely soluble in – CH3OH or C6H6?
e. Which solvent is NaOH more likely soluble in – CH3OH or C6H6?
f. Compounds with the formula CnH2n + 1OH are soluble in H2O when n is
small but not when n is large. Suggest an explanation for this phenomenon.
g. Glucose has the following structure: H(CHOH)5CHO
What parts of the molecule indicate that this substance is soluble in water?
a. A solution is prepared by combining 2.09 g of CO2 and 35.5 g of H2O.
Identify the solute and solvent.
b. A solution is prepared by combining 10.3 g of Hg(liquid) and 45.0 g of
Ag(solid). Identify the solute and solvent.
c. Decide if a solution containing 45.0 g of NaCl per 100 g of H2O is
unsaturated, saturated, or supersaturated.
d. Decide if a solution containing 0.000092 g of AgCl per 100 g of H2O is
unsaturated, saturated, or supersaturated.
e. Would the solution in Exercise “c” be described as dilute or concentrated?
Explain your answer.
f. Would the solution in Exercise “d” be described as dilute or concentrated?
Explain your answer.
23.1 MOLARITY AND MOLALITY
Molarity (molar concentration) is defined as the number of moles of solute
divided by the number of liters of solution:
molarity =moles of solute / liters of solution
C = n(solute) / V(solution)
Molarity is expressed in mol/L which can be simplified as just big letter “M”.
A similar in spelling but different in meaning unit of concentration is
molality, which is defined as the number of moles of solute per kilogram of
solvent, not per liter of solution:
molality = moles of solute / kilograms of solvent
Cm = n(solute) / m(solvent)
Another way to specify an amount of solute is percentage composition by
mass (or mass percentage, % m/m). It is defined as follows:
% m/m = (mass of solute / mass of entire sample)x 100 %
ω= m(solute) / m(solution)
Mass percentage has a wider sense than just a fraction of solute in
a solvent multiplied by 100. The same index is used to describe the mass
content of a compound. For example, the mass percentage of potassium (K) in
potassium oxide (K2O) is equal to the ratio between molar mass of potassium
multiplied by 2 (imagine that potassium is a solute) and the molar mass of
the whole compound (imagine that oxygen is a solvent). In more complicated
compounds “solvent” is everything else except atoms for which the mass
percentage has to be calculated.
- What is the molarity of a solution made by dissolving 13.4 g of NaNO3
in water? The final volume of that solution is equal to 345 mL.
- What is the molality of a solution made by dissolving 332 g of C6H12O6 in 4.66 kg of water?
- How many moles of MgCl2 are present in 0.0331 L of a 2.55 M water solution? Density is 1.2 g/mL.
- What is the mass percentage of MgCl2?
- How many moles of NH4Br are present in 88.9 mL of a 0.228 M water
solution? Density is 1.1 g/mL. What is the mass percentage of NH4Br?
- What volume of 5.56 M NaCl is needed to obtain 2L of 0.85 % NaCl
solution? Density is equal to 1 g/mL.
6. What mass of 96 % C2H5OH is needed to obtain 3L of 40 % C2H5OH
solution? Density is equal to 0.94 g/ml.
- Calculate the mass percentage of nitrogen in NH4NO3.
- Calculate the mass percentage of hydrogen in H2C2O4·2H2O.
- Calculate the mass percentage of oxygen in CaSO4·.H2O.
24.1 THEORY OF ELECTROLYTIC DISSOCIATION
Dissociation in chemistry is a general process in which ionic compounds
separate or split into smaller charged particles — ions. For example, when
a hydrogen chloride is put in water, a covalent bond between an electronegative
chlorine atom and a hydrogen atom is broken by heterolytic fission, which
gives a proton and a negatively charged ion. Dissociation process is frequently
confused with ionization. Although it may seem as a case of ionization, in
reality the ions of any salt already exist within the crystal lattice. When salt
is dissociated, its constituent ions are simply surrounded by water molecules
(i. e. solvation of ions happens) and their effects become visible (e.g.
solution becomes electrolytic). However, no transfer or displacement of
electrons occurs. Actually, the chemical synthesis of salt from metals and
nonmetals involves ionization.
Solvation, also sometimes called dissolution, is the process of attraction
and association of molecules of a solvent with molecules or ions of a solute. As
ions dissolve in a solvent they spread out and become surrounded by solvent
Polar solvents are those with a molecular structure that contains dipoles.
The polar molecules of these solvents can solvate ions because they can orient
the appropriate partially charged portion of the molecule towards the ion in
response to electrostatic attraction. This stabilizes the system and creates
a solvation shell (or hydration shell in figure 24.1.
Any acid that dissociates 100% into ions is called a strong acid.
If it does not dissociate 100 %, it is a weak acid:
CH3COOH -> CH3COO- + H+
The ratio between the number of dissolved compounds and the number of
dissociated compounds is called dissociation degree (α). It can be expressed in
percent as well. Namely, for 1 M
Although the concept of equilibrium is widely used across sciences, in the context of chemistry, it arises whenever a number of different states of the chemical composition are possible, as for example, in a mixture of several chemical compounds that can react with one another, or when a substance can be present in more than one kind of phase.
A system of chemical substances at equilibrium, even though having an unchanging composition, is most often not static; molecules of the substances continue to react with one another thus giving rise to a dynamic equilibrium. Thus the concept describes the state in which the parameters such as chemical composition remain unchanged over time
- Барковский, Е. В. Неорганическая химия: пособие-репетитор : теоретические
основы. Примеры решения типовых задач. Тесты для самоконтроля / Е. В. Барковский.
Минск : Аверсэв, 2008. 416 с.
- Ткачёв, С. В. Основы общей и неорганической химии : учеб.-метод. пособие /
С. В. Ткачёв. 12-е изд. Минск : БГМУ, 2014. 136 с.
3. Ball, D. W. Introductory Chemistry, v. 1.0. / D. W. Ball. Washington : Flat World
Education, Inc., 2014. 352 p.
4. Wilson, D. Kaplan AP Chemistry 2014–2015 / D. Wilson. New York : Kaplan
Publishing, 2014. 396 p.
Lesson 1 1.1 Hydrogen.. 1.2 Water
Lesson 2 .. 2.1 Halogens…. 2.2 Hydrochloric acid….
Lesson 3 .. 3.1 Oxygen and ozone …. 3.2 Oxygen compounds….
Lesson 4 . 4.1 Sulfur…. 4.2 Compounds of sulfur……………………………………………………………………….. 15 Lesson 5 ……………………………………………………………………………………………………… 18 5.1 Nitrogen and its compounds……………………………………………………………… 18 5.2 Properties of nitric acid ……………………………………………………………………. 19 Lesson 6 ……………………………………………………………………………………………………… 21 6.1 Phosphorus …………………………………………………………………………………….. 21 6.2 Phosphorus compounds……………………………………………………………………. 22 Lesson 7 ……………………………………………………………………………………………………… 24 7.1 Carbon …………………………………………………………………………………………… 24 7.2 Compounds of Carbon …………………………………………………………………….. 26 Lesson 8 ……………………………………………………………………………………………………… 28 8.1 Silicon……………………………………………………………………………………………. 28 8.2 Compounds of silicon………………………………………………………………………. 29 Lesson 9 ……………………………………………………………………………………………………… 30 9.1 Alkali metals…………………………………………………………………………………… 30 9.2 Compounds of alkali metals……………………………………………………………… 31 Lesson 10 ……………………………………………………………………………………………………. 33 10.1 Alkaline-earth metals …………………………………………………………………….. 33 10.2 Compounds of alkaline-earth metals………………………………………………… 34 Lesson 11 ……………………………………………………………………………………………………. 35 11.1 Aluminum and its compounds…………………………………………………………. 35 11.2 Iron and its compounds………………………………………………………………….. 37 Lesson 12 ……………………………………………………………………………………………………. 39 12.1 Sample ticket #1 for control task on the chemistry of the elements……… 39 12.2 Sample ticket #2 for control task on the chemistry of the elements ……… 39 Main sources of literature ……………………………………………………………………………… 40 The periodic table of the elements………………………………………………………………….. 41 The solubility chart ………………………………………………………………………………………. 42 The reactivity series of metals
Учебное издание Хрусталёв Владислав Викторович
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