INTRODUCTION TO THE GENERAL CHEMISTRY

CONTENTS                    PREFACE

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…..

LITERATURE..

(Хрусталёв В.В. и др. ВВЕДЕНИЕ В ОБЩУЮ ХИМИЮ …2014

PREFACE
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[16]). 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.

 

LESSON 1
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.
Questions:
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[3]. 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»[29] 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 (), and theSun 太陽 represents Yang (). 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 ) 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 (), Mars 火星 is Fire (), Saturn 土星 is Earth (),Venus 金星 is Metal (), and Mercury 水星 is Water ().  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)[7] 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.[2]  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 chloridelead 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,[3] and the Haber process in 1909–1910 for ammonia  synthesis.[4]

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.[5] 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.[6]

1.2 ATOMS AND MOLECULES
The  modern chemistry began with theory of the Conservation of mass and Antoine Lavoisier,  discovery of the chemical elements and atomic theory,

Atom is the smallest piece of an element that maintains the identity of that
element.
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
substance.
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
polyatomic molecules.
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).
Questions:
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.

Exercises:
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,
Mn(NO2)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,
HCl.

 

By the time of Antoine Lavoisier, a list of elements [30]  correspond more closely to four of the states of matter: solid, liquid, gas and  plasmaLaw 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.[13]

led to atomic theory in chemistry[12] — 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:

  1. Elements are made of extremely small particles called atoms.
  2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.
  3. Atoms cannot be subdivided, created, or destroyed.
  4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
  5. 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).

nuclear fusion and fission are nuclear reactions and not chemical reactions.  different isotopes of an element have slightly varying weights. Lavoisier‘s oxygen-based chemistry

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.

Various atoms and molecules as depicted in John Dalton’s A New System of Chemical Philosophy (1808 —  listed twenty elements and seventeen simple molecules. See also  Atomic mass unit (dalton)

(Dalton Minimum – a period of low solar activity  Daltonism  DemocritusPneumatic chemistry)

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.[44]

LESSON 2

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
a molecule.

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.
Questions:
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
the same?

Exercises:
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)…

 

3.1 VALENCE
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
valences.
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
obey.
Exercises:
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↑
10

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:

 

LESSON 4
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.[45]

Compound

Carbon dioxide (CO2), an example of a Chemical compound

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.[46]

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.[47] Inorganic compounds are named according to the inorganic nomenclature system.[48]

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.

LESSON 16

16.1 OXIDES

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

routes.

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

sodium hydroxide.

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

anymore.

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).

Questions:

  1. Give examples of basic, acidic and amphoteric oxides.
  2. How do the properties of oxides change if we go from left to right in  a period of the Periodic table?
  3. How do the properties of oxides change if we go from top to bottom in a group of the Periodic table?
  4. What is anhydride?
  5. What is volatility?

Exercises:

  1. 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).

  1. 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?
  2. 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.

  1. Write down reactions between i) calcium oxide and sulfuric acid; ii) iron oxide (II) and nitric acid; iii) chrom oxide (III) and sulfuric acid;
  2. iv) aluminum oxide and phosphoric acid.
  3. 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.

  1. Imagine how can the following oxides be produced: CaO; Al2O3; CO2; SiO2?
  2. Calculate the mass of iron in 500 g of i) FeO; ii) Fe2O3; iii) Fe3O4.

LESSON 17

17.1 BASES

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)

17.2 ALKALIS

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

Chemical properties.

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]

Questions:

  1. Give a definition of base.
  2. Give a definition of alkali.
  3. Give a definition of hydroxide.
  4. What metals are active?

Exercises:

  1. Imagine how can the following substances be produced: i) KOH;  ii) LiOH; iii) Fe(OH)3; iv) Mg(OH)2; v) Zn(OH)2.
  2. How can Zn(OH)2 be produced from the set of the following  substances: Na; H2SO4; ZnO; H2O?

HCl; Fe2O3; H2O?

  1. What substances react with KOH? Write down equations of possible reactions.
  2.  Al2O3; Ca(HCO3)2; HCl; CO2; H2SO4; KNO3; Zn(OH)2; NaHCO3;K2CO3; SO3; H3PO4; NH4Cl.
  3. 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.

  1. Finish chemical reactions:   NaOH + (i)  P2O5 →, ii. + Cr2(SO4)3→

iii. NaOH + CO2 →,   NaOH + NH4NO3 →,  NaOH + Cr(OH)3 →

  1. 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 →

LESSON 18

18.1 ACIDS

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

When hydrogen bromide (HBr), pictured, is dissolved in water, it forms the strong acid hydrobromic acid

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.[59] 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.[60]

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↑

Table 18.1

Formulas and names of some acids

Formula Name

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

Chemical properties.

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.

Questions:

  1. Give the formula for each acid.
  2. perchloric acid
  3. hydriodic acid
  4. hydrosulfuric acid
  5. phosphorous acid
  6. Name each acid.
  7. HF
  8. HNO3
  9. H2C2O4
  10. H2SO4
  11. H3PO4
  12. HCl
  13. Name some properties that acids have in common.

Exercises:

  1. 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?

  1. 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.

  1. 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:

  1. P2O5 + H2O →
  2. HCl + Mg(OH)2 →  iii. Al(OH)3 + HCl →
  3. Fe + HCl→
  4. SO2 + H2O →
  5. Fe2O3 + H2SO4 →

vii. Fe + HI →

viii. N2O5 + H2O →

  1. CaO + H3PO4 →
  2. LiCl + H3PO4 →
  3. What is the mass of ZnSO4 which was produced from 9.8 g of sulfuric

acid and 8.1 g of ZnO?

  1. What is the mass of K3PO4 which was produced from 49 g of

phosphoric acid and 80 g of KOH?

LESSON 19

19.1 SALTS

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).

Neutral salts

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

carbonate.

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

Table 19.1

Solubility chart

Table

 

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

series”:

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.

Questions:

  1. Give a definition of the term “salt”.
  2. What is the difference between normal, acidic and basic salts?
  3. How can the reactivity series of metals be used?
  4. How can the solubility chart be used?

Exercises:

  1. 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

phosphate.

  1. What substances react with AgNO3? Write down equations of possible

reactions. HCl; FeSO4; CaCl2; BaCl2; H2SO4; NaI; KBr; K3PO4; Na2CO3;

Ba(NO3)2.

  1. How can the following salts be produced: KCl; Fe2(SO4)3; Zn(NO3)2?
  2. Which salts react with each other: i) CaCO3 + KCl; ii) MgCl2 +Na2CO3; iii) NaCl + K2CO3; iv) BaCl2 + K2SO4.
  3. 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?

  1. What kind of salt is produced in the reaction between 10.5 g of calcium

hydroxide and 13.7 g of nitric acid?

  1. 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.

Exercise:

  1. CaCO3 → CaO → Ca(OH)2 → Ca(HCO3)2 → CaCO3
  2. Mg-  MgO ®Mg(OH)2 ® MgCl2 ®Mg(NO3)2
  3. Ba

®Ba(OH)2 ® BaOHCl ® BaCl2 ®BaCO3

  1. K® KOH ®KHSO4 ® K2SO4 ® KCl
  2. Al → Al2(SO4)3 → Na[Al(OH)4] → Al(NO3)3 → Al(OH)3
  3. NaHCO3 → CO2 → CaCO3 → Ca(HCO3)2 → CaO
  4. Cu → CuS → CuSO4→ CuCl2 → Cu(NO3)2
  5. CuSO 4 ® SO3 ® NaHSO4 ® Na2SO4 ® NaHSO4
  6. Fe ® FeCl2 ® Fe(NO3)2 ® FeSO4 ® Fe(OH)2
  7. 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”).

Exercise:

  1. Calculate the sum of molar masses for compounds A, B, C and D from

the chain of chemical reactions.

H2→A→  B → C →  D

  1. Calculate the sum of molar masses for compounds B and D from

the chain of chemical reactions.

Na → A → B → C → D

  1. 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

  1. Calculate the molar mass of compound D from the chain of chemical

reactions.

Al, → A →B, → C → D

  1. Calculate the sum of molar masses for compounds C and D from

the chain of chemical reactions.

Ca → A → B → C → D

LESSON 21

Sample ticket for control task #3 on main types of inorganic chemical  compounds

  1. 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®

  1. P2O5®H3PO4®Ca(H2PO4)2®Ca3(PO4)2®H3PO4
  2. Ca ®Ca(OH)2®CaOHCl®CaCl2®CaCO3
  3. Modern chains of chemical reactions

Calculate the sum of molecular masses for compounds A, B, C and D

from chains of chemical reactions:

  1. C → A → B→C→D
  2. Ca3(PO4)2 → A → B ( )→C ( )→D

LESSON 22

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.

Table 22.1

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.

Questions:

  1. Define  solute and solvent.
  2. Define  saturated, unsaturated, and supersaturated solutions.
  3. Differentiate between polar and nonpolar solvents.
  4. 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?

Exercises:

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.

LESSON 23

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)

23.2MASS PERCENTAGE

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.

Exercises:

  1. 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.

  1. What is the molality of a solution made by dissolving 332 g of C6H12O6  in 4.66 kg of water?
  2. How many moles of MgCl2 are present in 0.0331 L of a 2.55 M water solution? Density is 1.2 g/mL.
  3. What is the mass percentage of MgCl2?
  4. 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?

  1. 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.

  1. Calculate the mass percentage of nitrogen in NH4NO3.
  2. Calculate the mass percentage of hydrogen in H2C2O4·2H2O.
  3. Calculate the mass percentage of oxygen in CaSO4·.H2O.

LESSON 24

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

molecules.

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.

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

 

 

 

 

 

 

Chemical equilibrium

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

Law of conservation of energy leads to the important concepts of equilibriumthermodynamics, and kinetics.  Hess’s law

1s, 2s, 2px,2py, and 2pz.
The shapes of the first five atomic orbitals using color to depict the phase of the wave function

LITERATURE

  1. Барковский, Е. В. Неорганическая химия: пособие-репетитор : теоретические

основы. Примеры решения типовых задач. Тесты для самоконтроля / Е. В. Барковский.

Минск : Аверсэв, 2008. 416 с.

  1. 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.

Part 2

CONTENTS Preface.

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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ч.1. Общая химия


  • Общая химия (начала), Предмет химии:  

1.1 Химия: 1. Предмет 2. Основные понятия,  химические элементы и соединения.  Расчеты.   

1.2 Физика: 3. Строение атома.  4. Периодическая система элементов.   5. Химическая связь.   6. Химические реакции и тепловой эффект.   7. Скорость и равновесие.  

1.3 Вещества, классы и их связь: 8. Растворы.  9. Электролиты, диссоциация и реакции ионов.   10.Оксиды.  11. Основания.   12. Кислоты.  13.Соли.  14.Взаимосвязь неорганических соединений.

ч.2 Неорганика

  • Неорганика ( 13 )
      Учебник — минимум: краткий курс химии (часть 2. Химия элементов: Свойства элементов, металлов, неметаллов и их соединений).

2.1. Металлы: 15. Металлы. 16. Свойства металлов и их соединений (система). 17. Соединения, получение и роль металлов в природе.

2.2. Неметаллы:  18. Водород.   19. Галогены.   20. VIА группа. Кислород.   21.Вода.   22. Сера.   23. VА группа.  Фосфор.   24. Азот.   25. IVА группа. Кремний.  26. Углерод.

ч.3. Органическая химия

— химия соединений углерода, в отличии от простых неорганических образующих длинные цепи и “сложные радикалы”, сохраняющиеся в реакциях (они показываются в формулах, начиная с СН3ОН и НСООН вместо СН4О и Н2СО2, показывая и функциональные группы и отличия связей О-Н и С-Н)*. Устойчивость их определяет огромное разнообразие и значение их комбинаций (>20 млн.соединений при 2 млн. всех неорганических), уровень более высокоорганизованной материи, промежуточный между неорганическим и органическим (сохранения уже не соединений, а их реакций и жизни, систем организмов)

“Основное различие органических веществ зависит от разнообразия не составных элементов, как в минеральной химии, а пропорций и способов соединения их”, “В неорганической химии все радикалы — просты, в органической же химии все элементы — сложны” (Либих, Дюма, 1836).

Предмет органической химии.

Органические соединения составляют нашу пищу, одежду, топливо, дерево и синтетические материалы. В живой природе и организме их источником является фотосинтез растений из СО2 через глюкозу (в.42, Д4). Из природных источников получают г.о. сложные органические вещества (в.4247), а более простые, в т.ч. высокомолекулярные, полимеры синтезируют из простейших углеводородов природного газа и нефти (в.2834).

Лавуазье и Берцелиус ввели понятие органической химии и радикалов, Дюма и Жерар — гомологии и типов, Кекуле — четырехвалентности и цепей углерода, Купер и Бутлеров — химических связей и строения. Вант-Гофф установил пространственное представление их (стереохимию), а Н.Морозов и Г.Льюис — электронные (пар электронов), развитые в современную электронную теорию химического строения (Л.Полинг, К.Ингольд и др.).

  • Органика ( 22 )
    Учебник химии (минимум, для поступающих, с тестами)

3.1. Углеводороды: 27.Теория химического строения.  28. Алканы. 29. Алкены.  30. Полимеры.   31. Диены,каучук.   32. Алкины.  33. Бензол. 34. Источники и применение углеводородов.

3.2. Кислородсодержащие соединения:  35. Спирты.   36. Многоатомные спирты.   37. Фенолы.   38. Альдегиды.   39. Кислоты.   40. Сложные эфиры.

3.3. Природные соединения:  41. Жиры.   42. Углеводы.   43. Ди- и поли-сахариды.   44. Азот- и гетеро-содержащие соединения. 45. Амины.   46. Аминокислоты.   47. Белки.   48. Взаимосвязь органических соединений.