This file for education only.

3. Atom is a small part of element that takes part in chemical
reactions. It is made up of three subatomic structures
called Protons, Neutrons, and Electrons.
Subatomic structures:
Atomic model

4. Ernest rutherford
30 AUG 1871

6. 2.1 Atomic structure
The nucleus, in the center of an atom, consist of protons and neutrons.
Orbiting around the nucleus are the electrons. Each unique element has an Atomic
Number equal to the number of protons it contains. There are 94 naturally occurring
elements (1-94) and others which have been artificially created (95+..) Each
element has an Atomic Weight for the most commonly found isotope. Atomic
Weight = number of protons + number of neutrons. See: The Periodic table In a
stable uncharged atom the number of electrons will equal the number of protons. If
the number of electrons is changed the atom will become ionized and gain either a
positive (fewer electrons) or negative (greater electrons) charge.

7. Elementary particles of atoms
Particle: Electricity (c) Mass (g)
𝑒−
-1,602× 10−19
9,109× 10−28
𝑝+
1,602× 10−19
1,67× 10−24
N 0 1,675 × 10−24
The Nucleus, in the center of the atom, consists of protons
and neutrons. Orbiting around the nucleus are the electrons.

8. Atomic structure model

9. 2.2 Atomic number, atomic mass and isotope mass
2.2.1 Atomic Number
The atomic number or proton number (symbol Z ) of
a chemical element is the number of protons found in the nucleus of
every atom of that element. The atomic number uniquely identifies a
chemical element. It is identical to the charge number of the nucleus.
In an uncharged atom, the atomic number is also equal to the number
of electrons.
Atom mass
Atom number

10. The atomic mass (ma) is the mass of an atom. Its unit is the
unified atomic mass unit (symbol: u) where 1 unified atomic mass
unit is defined as ​1⁄12 of the mass of a single carbon-12 atom, at
rest.[1] The protons and neutrons of the nucleus account for nearly
all of the total mass of atoms, with the electrons and nuclear
binding energy making minor contributions. Thus, the atomic mass
measured in u has nearly the same value as the mass number.

11. 𝑴 𝒂 =
Weight of 1 atom
𝟏
𝟐
𝒘𝒆𝒊𝒈𝒉𝒕 𝑪 𝟏𝟐 𝟏𝒂𝒕𝒐𝒎
element mass Atomic mass
𝑯 𝟏
𝟏
1.06 × 10−24
1.0
𝑯𝒆 𝟐
𝟒
6.64 × 10−24
4.0
𝑪 𝟔
𝟏𝟐
1.99 × 10−23
12.0 (the standard)
𝑶 𝟖
𝟏𝟔
2.66 ×−23
16.0
𝑵𝒂 𝟏𝟏
𝟐𝟑
3.82 × 10−23
23.0

12. ❖ Atomic structure:
Electrons are the subatomic particles that orbit the nucleus of
an atom. They are generally negative in charge and are much smaller
than the nucleus of the atom and atom radius have a 10 × 10−15
𝑚 .

13. ❖ Radiation

14. ❖ molecule
A molecule is the smallest particle in a chemical element or
compound that has the chemical properties of that element or
compound. Molecules are made up of atoms that are held together
by chemical bonds. These bonds form as a result of the sharing or
exchange of electrons among atoms

15. Molecule structure

16. ❖ Neutron
The neutron is a subatomic particle, symbol n or n0
, with no net electric charge and a mass slightly greater than that of a
proton. Protons and neutrons constitute the nuclei of atoms.

17. ❖ What’s a proton
A proton is a subatomic particle, symbol p or p+, with a
positive electric charge of +1e elementary charge and a mass slightly
less than that of a neutron. Protons and neutrons, each with masses of
approximately one atomic mass unit, are collectively referred to as
"nucleons".

18. Isotope
Periodic table of the elements

19. 2.2.3 Isotope
Isotopes are variants of a particular chemical element which
differ in neutron number, and consequently in nucleon number.
All isotopes of a given element have the same number of protons but
different numbers of neutrons in each atom this isotope.

20. Electronic Structure of Atoms the arrangement of electron
around the nucleus of the atom.
electron
nuclear
Proton + neutron

21. 2.3.1 hydrogen atom
Niels Bohr introduced the atomic Hydrogen model in 1913. He
described it as a positively charged nucleus, comprised of protons and
neutrons, surrounded by a negatively charged electron cloud. In the
model, electrons orbit the nucleus in atomic shells. The atom is held
together by electrostatic forces between the positive nucleus and
negative surroundings.

22. But if hydrogen is stimulated with a higher energy, there
will be a different energy absorption between the two energy
levels, but on the other hand, if the energy level is reduced it will
cause an energy release which is in the form of a magnetic
radiation called a photon.

23. The energy used to change the electron energy level ∆𝐸
will be directly related to the frequency of the photons, which
follows the law of the Planck's equation.
∆𝐸 = ℎ𝑣 (2.1)
ເມື່ອ ℎ = 𝑃𝑙𝑎𝑛𝑘′
𝑠 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 = 6. 63 × 10−34
𝐽. 𝑆
∆𝐸 = Energy Conversion(eV)
𝑣 = Frequency
For magnetic radiation, write the equation:
𝑐 = 𝜆𝑣 (2.2)
𝑐 = Speed of light 3 × 108
𝑚/𝑠
𝜆 = Wavelength of light (𝑛𝑚)
This the transformative power can be written:
∆𝐸 =
ℎ𝑐
𝜆
(2.3)

24. 2.3.2 Quantum number
Modern atomic theory explains A total of four quantum
numbers are used to describe completely the movement and
trajectories of each electron within an atom.
1. The principal quantum number (n) describes the size of the orbital.
Orbitals for which n = 2 are larger than those for which n = 1, for
example. Because they have opposite electrical charges, electrons are
attracted to the nucleus of the atom. Energy must therefore be
absorbed to excite an electron from an orbital in which the electron is
close to the nucleus (n = 1) into an orbital in which it is further from
the nucleus (n = 2).

27. 2. Angular Momentum (Secondary,
Azimunthal) Quantum Number
(l): l = 0, ..., n-1.
Specifies the shape of an orbital with
a particular principal quantum
number. The secondary quantum
number divides the shells into smaller
groups of orbitals
called subshells (sublevels). Usually,
a letter code is used to identify l to
avoid confusion with n:

28. 3. Magnetic Quantum Number (ml): ml = -l, ..., 0, ..., +l.
Specifies the orientation in space of an orbital of a given energy (n) and shape (l).
This number divides the subshell into individual orbitals which hold the electrons;
there are 2l+1 orbitals in each subshell. Thus the s subshell has only one orbital,
the p subshell has three orbitals, and so on.

29. 4. Spin Quantum Number (ms): ms = +½ or -½.
Specifies the orientation of the spin axis of an electron. An
electron can spin in only one of two directions (sometimes
called up and down).
The Pauli exclusion principle (Wolfgang Pauli, Nobel Prize 1945)
states that no two electrons in the same atom can have identical
values for all four of their quantum numbers. What this means is
that no more than two electrons can occupy the same orbital, and
that two electrons in the same orbital must have opposite spins.

30. 4. Spin Quantum Number (ms): ms = +½ or -½.

31. ❖ Electron arrangement
Refers to the number of layers of an electron with 7 levels and 7
cycles, the symbol representing the number of levels is: “The
eigenvector is an e: K = 1; L = 2; M = 3; N = 4; O = 5; P = 6; & Q = 7
and has 4 power levels symbolically: s = 2; p = 8; d = 18; & f = 32. By
finding the energy level from the following formula: 𝑛 = 2 (𝑛) ^ 2. The
maximum electron energy will only be 32 and it will come down to the
lower energy levels again. It is also possible to know the final electron
number as well as the number of cycles of a slave by the electron
arrangement of the atoms as follows picture:

32. ❖ Electron arrangement