1. Lesson № 1 Basic theoretical concepts on the structure of atom.
The structure of compounds.
Theory of chemical bond.
Structure of molecule.
Basic theoretical concepts about the structure of atom. Contemporary interpretation of the periodic law of D.I. Mendeleev on the basis of electronic theory of atom.

2. Part 1-2: 1. CHEMICAL FORMULAS AND THE MOLE CONCEPT 2. THE ATOM

3. Chemistry (from Egyptian kēme (chem), meaning "earth") is the science concerned with the composition, structure, and properties of matter, as well as the changes it undergoes during chemical reactions.
Химия (от египетского Кеме (хим), что означает «земля») это наука о составе, структуре и свойствах вещества, а также изменений, которые вещество претерпевает в ходе химических реакций.

4. Chemistry – Х И М И Я
Science – Н А У К А
Matter – М А Т Е Р И Я, В Е Щ Е С Т В О
Chemical reactions – Х И М И Ч Е С К И Е Р Е А К Ц И И
Concerned – С О С Т А В
Structure – С Т Р О Е Н И Е
Properties – С В О Й С Т В А

5. A chemical element is a pure chemical substance consisting of a single type of atom distinguished by its atomic number, which is the number of protons in its atomic nucleus.
chemical element
metals
metalloids
non-metals
Металлы
Амфотерные элементы
Неметаллы

6. A chemical element is a pure chemical substance consisting of a single type of atom distinguished by its atomic number, which is the number of protons in its atomic nucleus.
Химический элемент — совокупность атомов с одинаковым зарядом ядра и числом протонов, совпадающим с атомным номером.

7. Chemical element - Химический элемент
Atom – Атом
Atomic number – Атомный номер
Number of protons – Число протонов
Relative atomic mass – Относительная атомная масса (АR)

8. Names of the first 20 elements
Atomic Name Symbol Relative
Number atomic
mass
Атомный Название Символ Относительная
Номер атомная
масса
hydrogen H
helium He
lithium Li
beryllium Be
boron B
carbon C
nitrogen N
oxygen O
fluorine F
neon Ne
водород H
гелий He
литий Li
бериллий Be
бор B
углерод C
азот N
кислород O
фтор F
неон Ne

9. Names of the first 20 elements
Atomic Name Symbol Relative
Number atomic
mass
Атомный Название Символ Относительная
Номер атомная
масса
sodium Na
magnesium Mg
aluminium Al
silicon Si
phosphorus P
sulfur S
chlorine Cl
argon Ar
potassium K
calcium Ca
натрий Na
магний Mg
алюминий Al
кремний Si
фосфор P
сера S
хлор Cl
аргон Ar
калий K
кальций Ca

10. Compounds
Some substances are made up of a single element although there may be more than one atom of the element in a particle of the substance. Oxygen is diatomic, that is, a molecule of oxygen contains two oxygen atoms.
О2 = 2О
A compound contains more than one element, for example a molecule of water contains two hydrogen atom and one oxygen atom.
Н2О
Water is a compound not an element because it can be broken down chemically into its constituent elements: hydrogen and oxygen.

11. (literally the formula obtained by experiment)
Formulas of compounds
Compounds can be described by different chemical formulas.
Empirical formula
Э м п и р и ч е с к а я ф о р м у л а
(literally the formula obtained by experiment)
This show the whole number ratio of atoms of each element in a particle of the substance.
It can be obtained be either knowing the mass of each element in the compound or from the percentage composition by mass of the compound.
The percentage composition can be converted directly into mass by assuming 100 g of the compound are taken.

12. Example: A compound contains 40. 00% carbon, 6. 73% hydrogen and 53
Example: A compound contains 40.00% carbon, 6.73% hydrogen and 53.27% oxygen by mass, determine the empirical formula.
АR (С) = 12,01; АR (Н) = 1,01; АR (О) = 16,00
Amount/mol Ratio
C /12.01= Empirical
H /1.01= formula = CH2O
O /16.00=

13. Formulas of compounds Molecular formula
М о л е к у л я р н а я ф о р м у л а
For molecules this is much more useful as it shows the actual number of atoms of each element in a molecule of the substance. It can be obtained from the empirical formula if the molar mass of the compound is also known.
Methanol CH2O (Mr=30), ethanoic acid C2H4O2 (Mr=60) and glucose C6H12O6 (Mr=180) are different substances with different molecular formulas but all with the same empirical formula CH2O. Note that subscripts are used to show the number of atoms of each element in the compound.

14. Formulas of compounds
Structural formula С т р у к т у р н а я ф о р м у л а This show the arrangement of atoms and bonds within a molecule and is particularly useful in organic chemistry. The three different formulas can be illustrated using ethane: CH2 C2H4 H2C=CH2 Empirical Molecular Structural formula formula formula

15. Empirical formula - Эмпирическая формула
Molecular formula – Молекулярная формула
Structural formula – Структурная формула

16. Mole concept and Avogadro′s constant
A single atom of an element has an extremely small mass.
For example an atom of c arbon-12 has a mass of * g. this is far too small to weigh.
A more convenient amount to weigh is g/ g of carbon-12 contains 6.02*1023 atoms of carbon-12. This number is known as Avogadro′s constant (NA or L).

17. In chemistry the Avogadro constant (symbols: L, NA) is defined as the number of constituent particles (usually atoms or molecules) per mole of a given substance, where the mole is one of the seven base units in the International System of Units (SI).
The Avogadro constant has dimensions of reciprocal mol and its value is equal to (27)×1023 mol−1.

18. Chemists measure of substances in moles
Chemists measure of substances in moles. A mole is the amount of substance that contains L particles of that substance.
The mass of one mole of any substance is known as the molar mass and has the symbol M.
For example, hydrogen atoms have 1/12 of the mass of carbon-12
so a mole of hydrogen atoms contains 6.02*1023 hydrogen atoms and
has a mass of g.

19. The relative atomic mass of an element Ar is the weighted mean of all the naturally occurring isotopes of the element relative to carbon-12. This explains why the relative atomic masses given for the elements are not whole numbers.
The units of molar mass are g/mol but relative molar masses Mr have no units. For molecules relative molecular mass is used.
For example, the Mr of glucose, C6H12O6 = (6*12.01)+(12*1.01)+(6*16.00)=
For ionic compounds the term relative molecular mass is used.

20. Avogadro′s constant – Постоянная Авогадро Постоянная Авогадро = NA
Relative molecular mass – Относительная молекулярная масса = Mr
Mr =ΣАR
Molar mass – Молярная масса = М
М = г/моль

21. . THE ATOM Composition of atoms
The smallest part of element is atom.
It used to be thought that atoms re indivisible but they can be broken down into many different sub-atomic particles.

22. Structure of an Atom
subatomic particles
electrons
neutrons
protons

23. All atoms, with exception of hydrogen, are made up of three fundamental sub-atomic particles – protons, neutrons, and electrons.

24. Size and structure of atoms
Atoms have a radius in the order of m. Almost all of the mass of an atoms is concentrated in the nucleus which has a very small radius in the order of m. All the protons and neutrons (collectively called nucleons) are located in the nucleus. The electrons are to be found in energy levels or shells surrounding the nucleus. Much of the atom is empty space.

25. SUMMARY OF RELATIVE MASS AND CHARGE
Particle Relative mass Relative charge
Proton
Neutron
Electron *

26. Equal to the number of protons and neutrons in the nucleus.
Mass number (A)
Equal to the number of protons and neutrons in the nucleus.

27. Atomic number (Z)
Equal to the number of protons in the nucleus and to the number of electrons in the atom.
The atomic number defines which element the atom belongs to and consequently its position in the Periodic Table.

28. classical physics predicts that electron falls into nucleus
mass number
17p
18n
17e 35 17 Cl
18 neutrons
atomic number
Why are atoms stable? e-
classical physics predicts that electron falls into nucleus

29. Charge Atoms have no charge so n=0 and this is left blank.
However by losing one or more electrons atoms become positive ions, or by gaining one or more electrons form negative ions.

30. Isotopes
All atoms of the same element must contain the same number of protons, however they may contain a different number of neutrons.
Such atoms are known as isotopes. Chemical properties are related to the number of electrons so isotopes of the same element have identical chemical properties.
Since their mass is different their physical properties such as density and boiling point are different.

31. Example of isotopes:
11 H, 12 H, 13 H
612 C, 614 C
1735 Cl, 1737 Cl

32. Relative atomic mass
The two isotopes of chlorine occur in the ratio of 3:1.
That is, naturally occurring сhlorine contains 75% 1735Cl and 25% 1737Cl.
The weighted mean molar mass is thus:
((75*25+25*37))/100=35.5 g/mol
And the relative atomic mass is 35.5.

33. The planetary model of atom and its contradiction
By 1911 the components of the atom had been discovered. The atom consisted of subatomic particles called protons and electrons. However, it was not clear how these protons and electrons were arranged within the atom. J.J. Thomson suggested the "plum pudding" model. In this model the electrons and protons are uniformly mixed throughout the atom:
Rutherford's Planetary Model
of the Atom

34. Rutherford tested Thomson's hypothesis by devising his "gold foil" experiment. Rutherford reasoned that if Thomson's model was correct then the mass of the atom was spread out throughout the atom. Then, if he shot high velocity alpha particles (helium nuclei) at an atom then there would be very little to deflect the alpha particles. He decided to test this with a thin film of gold atoms. As expected, most alpha particles went right through the gold foil but to his amazement a few alpha particles rebounded almost directly backwards.
Rutherford's Planetary Model
of the Atom

35. These deflections were not consistent with Thomson's model
These deflections were not consistent with Thomson's model. Rutherford was forced to discard the Plum Pudding model and reasoned that the only way the alpha particles could be deflected backwards was if most of the mass in an atom was concentrated in a nucleus. He thus developed the planetary model of the atom which put all the protons in the nucleus and the electrons orbited around the nucleus like planets around the sun.
Rutherford's Planetary Model
of the Atom
Prof. N. De Leon

36. Bohr's postulates 1. Electrons in an atom exist in STATIONARY STATES.
Bohr stated that electrons orbit the nucleus WITHOUT emitting EM radiation.
Any permanent change in their motion must be accompanied by a complete transition from one stationary state to another.
Note that Bohr could NOT explain this.
2. Transmission between stationary states produces/absorbs EM Radiation.
When an electron moves between stationary states, it is accompanied by the emission or absorption of a photon.
This photon's energy is given by ΔE=hf
Thus, EM radiation is produced by the movement of electrons between energy states which account for Planck's "atomic oscillators.“

37. Bohr's postulates
3. The ANGULAR MOMENTUM of a stationary electron is QUANTISED
To explain, recall from Rutherford's model of the atom, in which the electrons orbited in circular orbits at ANY RADII.
If this coincided with Bohr's 2nd postulate, it would mean that every element can emit a FULL SPECTRUM as any transition would be possible.
As this is not the case, it would imply that electrons orbited at FIXED RADII.
Thus, this 3rd postulate allowed Bohr to explain the distinct spectra lines.
This can be represented mathematically by:
Note that you will never be asked to do equations with this, just remember it.

38. WAVE-PARTICLE DUALITY
large pieces of matter are mainly particle-like
small pieces of matter are mainly wave-like
MASS
Baseball
Proton
Electron
Photon
Particle-like
Wave-like

39. 1. light behaves like wave and particle
2. electron behaves like wave and particle
3. electrons are constituents of atoms
4. light is emitted/absorbed from atoms in discrete quantities (quanta)

40. QUANTUM NUMBERS principle quantum number
2. angular momentum quantum number
3. magnetic quantum number
4. spin quantum number

41. The principal quantum number (n) describes the size of the orbital.
QUANTUM NUMBERS
principle quantum number: the principal quantum number (n) describes the size of the orbital.
The principal quantum number (n) describes the size of the orbital.
The principal quantum number therefore indirectly describes the energy of an orbital.

42. n 1. principle quantum number n = 1, 2, 3, 4, 5…
hydrogen atom: n determines the energy of an atomic orbital
measure of the average distance of an electron from nucleus
n increases → energy increases
n increases → average distance increases

43. maximum numbers of electrons in each shell
K L M N O P
‘shell’
maximum numbers of electrons in each shell
n = 4
n = 3
n = 2
n = 1
2 n2
Fit waves into orbits e-
EXP8

44. QUANTUM NUMBERS
2. angular momentum quantum number: (l) describes the shape of the orbital.
Orbitals have shapes that are best described as spherical (l = 0), polar (l = 1), or cloverleaf (l = 2).
They can even take on more complex shapes as the value of the angular quantum number becomes larger.

45. define the ‘shape’ of the orbital
2. angular momentum quantum number
l = 0, 1, … (n-1)
l =
s p d f g h
define the ‘shape’ of the orbital
EXPIX

46. 1s, 2s, 3s 1s 2s 3s

47. QUANTUM NUMBERS
3. magnetic quantum number: in atomic physics, the magnetic quantum number is the third of a set of quantum numbers (the principal quantum number, the azimuthal quantum number, the magnetic quantum number, and the spin quantum number) which describe the unique quantum state of an electron and is designated by the letter m. The magnetic quantum number denotes the energy levels available within a subshell.

48. 3. magnetic quantum number
ml = -l, (-l + 1), … 0…… (+l-1) +l
defines orientation of an orbital in space

49. 2px, 3px, 4px
2px
3px
4px

50. d orbitals

51. QUANTUM NUMBERS
4. spin quantum number: In atomic physics, the spin quantum number is a quantum number that parametrizes the intrinsic angular momentum (or spin angular momentum, or simply spin) of a given particle. The spin quantum number is the fourth of a set of quantum numbers (the principal quantum number, the azimuthal quantum number, the magnetic quantum number, and the spin quantum number), which describe the unique quantum state of an electron and is designated by the letter s. It describes the energy, shape and orientation of orbitals.

52. 4. spin quantum number
ms = -1/2; + 1/2

53. ORBITALS AND QUANTUM NUMBERS
principle quantum number
2. angular momentum quantum number
3. magnetic quantum number
4. spin quantum number
n = 1, 2, 3, 4, 5…
l = 0, 1, … (n-1)
ml = -l, (-l + 1), … 0…… (+l-1) +l
ms = -1/2; + 1/2

54. (n, l, ml, ms) ATOMIC ORBITALS n l ml orbitals designation 1 1s 2 2s1s 2 2s
-1,0,+1 3
2px,2py,2pz 3s
3px,3py,3pz
-2,-1,0,+1,+2 5
3dxy,3dyz,3dxz,
3dx2-y2,3dz2 4 …

55. Principles and rules of filling of atomic orbitals by electrons
The principle of lowest energy
Electrons fill orbitals in a way to minimize the energy of the atom.
The electrons in an atom therefore fill the principal energy levels in order of increasing energy (the electrons are getting farther from the nucleus).
The order of levels filled looks like this: 
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p

56. Principles and rules of filling of atomic orbitals by electrons
Pauli Exclusion Principle
The Pauli exclusion principle states that no two electrons can have the same four quantum numbers. The first three (n, l, and ml) may be the same, but the fourth quantum number must be different.
A single orbital can hold a maximum of two electrons, which must have opposing spins; otherwise they would have the same four quantum numbers, which is forbidden. One electron is spin up (ms = +1/2) and the other would spin down (ms = -1/2).
This tells us that each subshell has double the electrons per orbital. The s subshell has 1 orbital that can hold up to 2 electrons, the p subshell has 3 orbitals that can hold up to 6 electrons, the d subshell has 5 orbitals that hold up to 10 electrons, and the f subshell has 7 orbitals with 14 electrons.

57. Pauli Exclusion Principle

58. Principles and rules of filling of atomic orbitals by electrons
Hund's rule states that:
Every orbital in a sublevel is singly occupied before any orbital is doubly occupied.
All of the electrons in singly occupied orbitals have the same spin (to maximize total spin).

59. Part 3: THEORY OF CHEMICAL BOND

60. IONIC BONDING
Ionic compounds are formed when electrons are transferred from one atom to another to form ions with complete outer shells of electrons.

61. In an ionic compound the positive and negative ions are attracted to each other by strong electrostatic forces, and build up into a strong lattice.
Ionic compounds have high melting points as considerable energy is required to overcome these forces of attraction.

62. The classic example of an ionic compound is sodium chloride Na + Cl –, formed when sodium metal burns in chlorine.
The crystal structure of sodium chloride, NaCl, a typical ionic compound.
The purple spheres are  sodium cations, Na+, and the green spheres are chlorideanions, Cl−.

63. Chlorine is a covalent molecule, so each atom already has an inert gas configuration.
However, the energy given out when the ionic lattice is formed is sufficient to break the bond in the chloride molecule to give atoms of chlorine.
Each sodium atom then transfers one electron to a chlorine atom to form the ions.
The charge carried by an ion depends on the number of electrons the atom needed to lose or gain to achieve a full outer shell.

65. Thus magnesium chloride two chlorine atoms each gain one electron from a magnesium atom to form Mg2+ Cl-2.

66. Cations Anions Group 1 Group 2 Group 3 Group 4 Group 5 Group Li+Na+K+ Mg2+Ca2+ Al3+ N3- P3- O2- S2- F- Cl- Br-

67. Formulas of ionic compounds
Lithium fluoride Li+F-
Magnesium chloride Mg2+Cl-2
Sodium oxide Na+2O2-
Calcium sulfide Ca2+S2-
Potassium nitride K+3N3-
Calcium phosphide Ca2+3P3-2
Aluminium bromide Al3+Br-3
Iron (lll) oxide Fe3+2O2-3
Iron (ll) oxide Fe2+O2-

68. Формулы ионных соединений
Фторид лития Li+F-
Магния хлорид Mg2+, Cl-2
Оксид натрия Na+2O2-
Кальция сульфид Ca2+S2-
Калия нитрид K+3N3-
Кальций фосфористый Ca2+3P3-2
Алюминий бромид Al3+Br-3
Железо (III) оксид Fe3+2O2-3
Железо (ll) оксид Fe2+O2-

69. COVALENT BONDING Single covalent bonds
Covalent bonding involves the sharing of one or more pairs of electrons so that each atom in the molecule achieves an inert gas configuration.

70. The simples covalent molecule is hydrogen.
Each hydrogen atom has one electron in its outer shell.
The two electrons are shared and attracted electrostatically by both positive nuclei resulting in a directional bond between the two atoms to form a molecule.
When one pair of electrons is shared the resulting bond is known as a single covalent bond is chlorine, Cl2.

71. Covalent bonds are affected by the electronegativity of the connected atoms.
Two atoms with equal electronegativity will make nonpolar covalent bonds such as H–H and Сl–Сl.
An unequal relationship creates a polar covalent bond such as with H−Cl.

72. Co-ordinate (dative) bonds
The electrons in the shared pair may originate from the same atom. This is known as a coordinate covalent bond. Carbon monoxide (CO), ammonium ion (NH4+), hydroxonium ion (H3O+). Sulfur dioxide and sulfur trioxide are bond sometimes shown as having a coordinate bond between sulfur and oxygen or they are show as having double between the sulfur and the oxygen. Both are acceptable.

74. METALLIC BONDING
The valence electrons in metals become detached from the individual atoms so that metal consist of a close packed lattice of positive ions in a sea of delocalized electrons.

75. A metallic bonds is the attraction that two neighbouring positive ions have for the delocalized electron between them.
Metals are malleable, that is, they can be bent and reshaped under pressure. They are also ductile, which means they can be drawn into a wire.