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Calculation of equivalent weight.






Equivalent weight of different types of substances can be calculate using formula: Мeq = feq • М

where

Мeq – equivalent molar mass

feq – equivalence factor

М – molar mass

 

Equivalence factors:

for an acids = , for a bases: , for a salts:

where n – number of metal atoms, Z – valency.

 

MOLE. MOLAR MASS.

 

In chemical processes involve the smallest particles - molecules, atoms, ions, electrons. The number of these particles even in small quantity of substances is very large. Therefore, to avoid mathematical operations with large numbers, to characterize the quantity of a substance involved in a chemical reaction, a special unit is used - the mole.

Mole - is the amount of the substance, which contains a number of particles (molecules, atoms, ions), equal to the Avogadro's constant:

NA = 6.0221367•1023 mol-1

This number is about the number of grains of sand that would fi t into a sphere the size of the Earth! Remember the value of this number to at least three signifi cant digits.

Avogadro’s constant (or Avogadro’s number) NA is defined as the number of atoms contained in 12 g of carbon isotope 12C:

 

Thus, one mole of substance contains 6.02 • 1023 particles of this substance.

Any amount of a substance can be expressed as a certain number of moles of ν (nu):

 

υ =

where N – number of particles of the substance;

NA – Avogadro’s constant.

The molar mass of the substance (M) – is the mass of 1 mole of the substace.

The molar mass of any substance can be calculated, when mass (m) and amount of substance (ν) are known:

M=

Accordingly, knowing the mass and molar mass of the substance, it is possible to calculate the number of moles:

υ =

or find the weight of substance from the number of moles and molar mass:

m = ν • M

Thus, mole is a quantity of a substance containing the same number of particles, but having a different weight for the different substances, because substance particles (atoms and molecules) has different weight.

For example: calculation the molar mass of nitric acid HNO3.

Mr (HNO3) = Ar (H) + Ar (N) + 3•Ar (O) = 1 + 14 + 3•16 = 63 g/mol

15 сураккк Atomos, the Greek root of the word atom, means “indivisible.” It was originally believed that the atom was the ultimate indivisible particle of which all matter was composed. Lord Rutherford showed in 1911 that the atom is not homogeneous, but rather has a dense, positively charged center surrounded by electrons. Subsequently, scientists have learned that the nucleus of the atom can be subdivided into particles called neutrons and protons. In fact, in the past two decades it has become apparent that even the protons and neutrons are composed of smaller particles called quarks.For most purposes, the nucleus can be regarded as a collection of nucleons (neutrons and protons), and the internal structures of these particles can be ignored. The number of protons in a particular nucleus is called the atomic number (Z), and the sum of the neutrons and protons is the mass number (A). Atoms that have identical atomic numbers but different mass number values are called isotopes. However, we usually do not use the singular form isotope to refer to a particular member of a group of isotopes. Rather, we use the term nuclide. In 1803, John Dalton (1766–1844) proposed his atomic theory, including the following postulates: 1. Matter is made up of very tiny, indivisible particles called atoms.

2. The atoms of each element has mass, but the mass of the atoms of one element is different from the mass of the atoms of every other element.

3. Atoms combine to form molecules. When they do so, they combine in small, whole-number ratios.

4. Atoms of some pairs of elements can combine with each other in different small, whole-number ratios to form different compounds.

5. If atoms of two elements can combine to form more than one compound, the most stable compound has the atoms in a 1: 1 ratio. (This postulate was quickly shown to be incorrect.)

The fi rst three postulates have had to be amended, and the fi fth was quickly abandoned altogether. But the postulates explained the laws of chemical combination known at the time, and they caused great activity among chemists, which led to more generalizations and further advances in chemistry. The postulates of Dalton’s atomic theory explained the laws of chemical combination very readily.

1. The law of conservation of mass is explained as follows: Because atoms only exchange “partners” during a chemical reaction and are not created or destroyed, their mass is also neither created nor destroyed. This, mass is conserved during a chemical reaction.

2. The law of defi nite proportions is explained as follows: Because atoms react in definite integral ratios (postulate 3), and atoms of each element have a definite mass (postulate 2), the mass ratio of one element to the other(s) must also be definite.

3. The law of multiple proportions is explained as follows: Because atoms combine in different ratios of small whole numbers (postulate 4), for a given number of atoms of one element, the number of atoms of the other element is in a small, whole-number ratio. A given number of atoms of the fi rst element implies a given mass of that element, and a small, whole-number ratio for the atoms of the second element (each of the same mass) implies a small, whole-number ratio of masses of the second element. Structure of an atom. Atoms are composed of many types of subatomic particles, but only three types will be important to chemists. Protons and neutrons exist in the atom’s nucleus, and electrons reside in orbitals around the nucleus. Protons have a positive (+) charge, neutrons have no charge - they are neutral. Electrons have a negative charge (-). The nucleus (plural, nuclei) is incredibly small, with a radius about one ten-thousandth of the radius of the atom itself. (If the atom were the size of a car, the nucleus would be about the size of the period at the end of this sentence.) The nucleus does not change during any ordinary chemical reaction. The protons, neutrons, and electrons have the properties listed in table below. The atom is the smallest unit that has the characteristic composition of an element, and in that sense, it is the smallest particle of an element. An uncombined atom is neither positive nor negative but electrically neutral, and thus the number of protons (p) must equal the number of electrons (e-): Number of protons = number of electrons (For a neutral atom)Because neutrons are neutral (has no charge), the number of neutrons (n) does not affect the charge on the atom. The number of protons in an atom determines the element’s identity. All atoms having the same number of protons are atoms of the same element. Atoms with different numbers of protons are atoms of different elements.The number of neutrons in the nuclei of atoms of the same element can differ. If two atoms have the same number of protons and different numbers of neutrons, they are atoms of the same element (they have the same atomic number). However, they have different masses because of the different numbers of neutrons. Such atoms are said to be isotopes of each other. Each isotope of an element is usually identifi ed by its mass number (A), which is defi ned as the sum of the number of protons and the number of neutrons in the atom: A = p + n = Z + n

16.Quantum numbers. Each electron in an atom is associated with a set of four quantum numbers: The principal quantum number (главное квантовое число).The angular momentum quantum number (орбитальное квантовое число).The magnetic quantum number (магнитное квантовое число).The spin projection quantum number (спиновое квантовое число).

Example: The quantum numbers used to refer to the outermost valence electrons of the Carbon (C) atom, which are located in the 2p atomic orbital, are; n = 2 (2nd electron shell), = 1 (p orbital subshell), m = 1, 0 or − 1, ms = ½ (parallel spins). The principal quantum number (n) can have any positive integral value, but the electrons in atoms in their most stable states have principal quantum numbers with values from 1 through 7 only. The most stable electronic state of an atom is called its ground state. Any higher energy state is called an excited state. (Unless “excited state” is specifi ed in later discussions, ground state is implied.). For each value of n, the angular momentum quantum number (l) for an electron can have integral values from zero to (n-1) it cannot be as large as n. The angular momentum quantum number has a small role in determining the energy of the electron, and it determines the shape of the volume of space that the electron can occupy. For each value of the angular momentum quantum number (l) the magnetic quantum number (ml) has values ranging from -l through zero to +l in integral steps. The spin quantum number (ms) may have values of -1/2 or +1/2 only. The value of does not depend on the value of any other quantum number. The spin value gives the orientation of the magnetic fi eld associated with the electron. Another important limitation on the quantum numbers of electrons in atoms, is the Pauli exclusion principle. This principle states that no two electrons in an atom can have the same set of four quantum numbers. This is like the business law that states that no two tickets to a rock concert can have the same set of date and section, row, and seat numbers. The row number may depend on the section number, and the seat number may depend on the row number, but the date does not depend on any of the other three. Similarly, the spin quantum number is independent of the other three quantum numbers. Hund’s rule states that the electrons within a given subshell remain as unpaired as possible. Or, the lowest energy confi guration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals. Ionisation Energy (I) is the energy required to remove an electron from a gaseous atom or ion: X(g) → X+(g) + e-where the atom or ion is assumed to be in its ground state. To introduce some of the characteristics of ionization energy, we will consider the energy required to remove several electrons in succession from aluminum in the gaseous state. The ionization energies are

In a stepwise ionization process, it is always the highest-energy electron (the one bound least tightly) that is removed fi rst. The fi rst ionization energy I1 is the energy required to remove the highest-energy electron of an atom. The fi rst electron removed from the aluminum atom comes from the 3p orbital. The second electron comes from the 3s orbital. Note that the value of I1 is considerably smaller than the value of I2, the second ionization energy.

Electron affinity (Ea) is defined as the amount of energy released when an electron is added to a neutral atom or molecule in the gaseous state to form a negative ion: X(g) + e- → X-(g)

This property is measured for atoms and molecules in the gaseous state only, since in a solid or liquid states their energy levels would be changed by contact with other atoms or molecules. Other theoretical concepts that use electron affinity include electronic chemical potential and chemical hardness. Another example, a molecule or atom that has a more positive value of electron affinity than another is often called an electron acceptor and the less positive an electron donor. Together they may undergo charge-transfer reactions. Atomic Radius

The atomic radius of a chemical element is a measure of the size of its atoms, usually the mean or typical distance from the centre of the nucleus to the boundary of the surrounding cloud of electrons. Since the boundary is not a well-defined physical entity, there are various non-equivalent definitions of atomic radius. Three widely used definitions of atomic radius are Van der Waals radius, ionic radius, and covalent radius.Atomic radius is generally stated as being the total distance from an atom’s nucleus to the outermost orbital of electron. In simpler terms, it can be defined as something similar to the radius of a circle, where the center of the circle is the nucleus and the outer edge of the circle is the outermost orbital of electron. Electronegativity. The different affi nities of atoms for the electrons in a bond are described by a property called electronegativity: the ability of an atom in a molecule to attract shared electrons to itself. Electronegativity is one of the most fundamental atomic parameters which expresses numerically the tendency to attract electrons to atoms in a molecule. It is very useful in explaining differences in bonding, structure, and reaction from the standpoint of atomic properties. Various schemes have been proposed to explain the theoretical basis of the power of electron attraction, and studies are still actively seeking new numerical values of electronegativity.One of them, R. Mulliken defined electronegativity χ M as the average of the ionization energy I and electron affinity Ae as follows

17.Development of the Periodic Table In a room where chemistry is taught or practiced, a chart called the periodic table is almost certain to be found hanging on the wall. This chart shows all the known elements and gives a good deal of information about each. Many atomic masses were determined as a direct result of Dalton’s postulates and the work that they stimulated, and scientists attempted to relate the atomic masses of the elements to the elements’ properties. This work culminated in the development of the periodic table by Dmitri Mendeleyev (1834-1907) and independently by Lothar Meyer (1830-1895). Because Mendeleyev did more with his periodic table, he is often given sole credit for its development. Mendeleyev put the elements known in the 1860s in ascending order according to their atomic masses (atomic numbers had not yet been defined) and noticed that the properties of every seventh known element were similar. He arranged the elements in a table, with elements having similar properties in the same group.

The periodic law was developed independently by Dmitri Mendeleev and Lothar Meyer in 1869. Mendeleev created the first periodic table and was shortly followed by Meyer. They both arranged the elements by their mass and proposed that certain properties periodically reoccur. Meyer formed his periodic law based on the atomic volume or molar volume, which is the atomic mass divided by the density in solid form. Mendeleev's table is noteworthy because it exhibits mostly accurate values for atomic mass and it also contains blank spaces for unknown elements

The classic Periodic Law formulation of D.I.Mendeleev is as states: the properties of elements and therefore the properties of the simple and complex substances they form are in the periodic dependence on their atomic weight.

Nowadays the formulation is: the properties of elements and therefore the properties of the simple and complex substances they form are in the periodic dependence on the charge of the atomic nuclei.

18) Periodic trends are specific patterns that are present in the periodic table, which illustrate different aspects of a certain element, including its size and its properties with electrons. The main periodic trends include: electronegativity, ionization energy, electron affinity, atomic radius, melting point, and metallic character. The periodic trends that arise from the arrangement of the periodic table provide chemists with an invaluable tool to quickly predict an element's properties. These trends exist because of the similar atomic structure of the elements within their respective group families or period and the periodic nature of the elements. For example: Ionization Energy Trends. Ionization Energy is the amount of energy required to remove an electron from a neutral atom in its gaseous phase. Conceptually, ionization energy is considered the opposite of electronegativity. The lower this energy is, the more readily the atom becomes a cation. Therefore, the higher this energy is, the more unlikely the atom becomes a cation. Generally, elements on the right side of the periodic table have a higher ionization energy because their valence shell is nearly filled. Elements on the left side of the periodic table have low ionization energies because of their willingness to lose electrons and become cations. Thus, ionization energy increases from left to right on the periodic table.

  • The ionization energy of the elements within a period generally increases from left to right. This is due to valence shell stability.
  • The ionization energy of the elements within a group generally decreases from top to bottom. This is due to electron shielding.
  • The noble gases possess very high ionization energies because of their full valence shell. Note that Helium has the highest ionization energy of all the elements.

19. A chemical bond is an attraction between atoms that allows the formation of chemical substances that contain two or more atoms. The bond is caused by the electrostatic force of attraction between opposite charges, either between electrons and nuclei, or as the result of a dipole attraction. Basic characteristics of the chemical bond are: bond energy, bond length and bond angle.The bond energy is a measure of the amount of energy needed to break apart one mole of covalently bonded gases. The SI units used to describe bond energy is kilojoules per mole of bonds (kJ/mol). Bond length is the distance between the centers of two covalently bonded atoms. The length of the bond is determined by the number of bonded electrons (the bond order). Generally, the length of the bond between two atoms is approximately the sum of the covalent radii of the two atoms, X + Y. Bond length is given in picometers. Therefore, the bond length of a triple bond < double bond < single bond.Bond angle is the angle that is formed between two adjacent bonds on the same atom. Covalent BondingMetal atoms can donate electrons to nonmetal atoms, but nonmetal atoms do not form monatomic positive ions because they would have to donate too many valence electrons to form octets. (Single nonmetal atoms do not donate electrons at all, but some groups of nonmetal atoms can. This will be discussed later in this section.) Nonmetal atoms can accept electrons from metal atoms if such atoms are present; otherwise, they can attain an octet by electron sharing. A covalent bond consists of shared electrons. One pair of electrons shared between two atoms constitutes a single covalent bond, generally referred to as a single bond. An unshared pair of valence electrons is called a lone pair. Elements or compounds bonded only by covalent bonds form molecules. Donor-Acceptor Bond (DAB) – special type of covalent bondDAB (also coordination bond), a term denoting one of the ways in which a chemical covalent bond is formed. The ordinary covalent bond between two atoms is due to the interaction of two electrons, one from each atom. The donor-acceptor bond is formed by a pair of electrons from one atom (the donor) and a free (unfilled) orbital from another (the acceptor). The difference can be expressed schematically as

Covalent bond Donor-acceptor bond

A + B → A: B A: + B → A: B

Hydrogen bonding is the intermolecular force of attraction between a hydrogen atom in one molecule and a small, highly electronegative atom with an unshared pair of electrons in another molecule. The most electronegative atoms have a great affi nity for electrons.Ionic BondingElectrons can be transferred from metal atoms to nonmetal atoms to achieve a more stable, lower energy state. The noble gases are composed of stable atoms; no reactions of the fi rst three (He, Ne, Ar) have been discovered, and the others (Kr, Xe, Rn) are almost completely unreactive. The stability of the noble gases is due to the 8 electrons in the outermost shell of each atom (2 electrons in the case of helium). In fact, 8 electrons in the outermost shell is a stable confi guration for most main group atoms. Atoms other than those of the noble gases tend to form ionic or covalent bonds (or both) with other atoms to achieve this electronic confi guration. The 8 electrons in the outermost shell are called an octet. The tendency of atoms to be stable with 8 electrons in the outermost shell is called the octet rule. In some compounds, one (or more) of the atoms does not obey the octet rule. Metallic bonding is the electromagnetic interaction between delocalized electrons, called conduction electrons, gathered in an " electron sea", and the metallic nuclei within metals. Understood as the sharing of " free" electrons among a lattice of positively charged ions (cations), metallic bonding is sometimes compared with that of molten salts; however, this simplistic view holds true for very few metals.


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