ATOMIC STRUCTURE AND PERIODIC SYSTEM ELEMENTS

A. ATOMIC STRUCTURE
1). Atomic Model Dalton
a) Atom described as a very small solid ball.
b) Atoms are the smallest particles that can not be broken again.
c) Atoms of an element have the same same, while the atoms of different elements, different in mass and nature.
d) Compounds are formed when atoms combine with each other.
e) A chemical reaction is the reorganization of atoms, so there is no atom has changed due to chemical reactions.
Dalton's atomic theory is supported by the second law of nature, namely:
1. Law of Conservation of Mass (Lavoisier law): the mass of substances before and after the reaction is the same.
2. Keep Comparative Law (law Proust): comparison of mass elements that make up a substance is fixed.
Weakness Atomic Model Dalton:
According to Dalton's atomic theory of number 5, no atom has changed due to chemical reactions. Now it turns out the reactions of nuclear chemistry, an atom can change into other atoms.
Example:
2). Thomson Atomic Model
a) After the discovery of the electron by JJ Thomson, Thomson atomic model formulated as an improvement Dalton atomic model.
b) Atoms consist of a positively charged material and in which the electrons are scattered like raisins in raisin bread.
3). Rutherford Atom Model
a) Rutherford found evidence that the atomic nuclei are positively charged, are smaller than the size of the atomic mass of an atom but almost entirely from the point masses.
b) Atoms consist of a positively charged nucleus and are at the center of the atom and the electrons move through the nucleus (like planets in the solar system).
Weakness Rutherford Atomic Model:
• Inability to explain why the electrons do not fall into the nucleus due to electrostatic attractive forces to the core electrons.
• According to Maxwell's theory, if the electrons as charged particles around the core that has the opposite charge then the trajectory will spiral and will lose power / energy in the form of radiation and eventually fell to the core.
4). Niels Bohr Atom Model
• atomic model based on quantum theory to explain the spectrum of hydrogen gas.
• According to Bohr, the line spectrum indicates that the electrons only occupy certain energy levels in the atoms.
He said:
a) Atoms consist of a positively charged nucleus and the surrounding circulating electrons are negatively charged.
b) Electrons orbit around the nucleus of an atom in a particular orbit known as the stationary state of motion (fixed), hereinafter referred to as the main energy level (electron shell), which is expressed by the principal quantum number (n).
c) During the electrons are in a stationary orbit, the energy will remain so no light is emitted.
d) The electrons can only move from the lower stationary trajectory to a higher stationary trajectory if absorbing energy. Conversely, if the electrons move from a higher stationary trajectory to lower the release of energy.
e) In the normal state (without outside influence), the electrons occupy the lowest energy level (called the basic level = ground state).
Niels Bohr Atom Model Weaknesses:
1. Just to explain the spectrum of the atom or ion containing one electron and is not in accordance with the spectrum of electron atoms or ions that much.
2. Not being able to explain that atoms can form molecules through chemical bonds.
5). Modern Atomic Model
Developed based on the theory of quantum mechanics is called wave mechanics; initiated by 3 experts:
a) Louis Victor de Broglie
Stating that the material has the properties of a material dualism, and as waves.
b) Werner Heisenberg
Put forward the principle of uncertainty for the material is of a particle and a wave. Distance or location of electrons that surround the nucleus can only be determined by the possibility - probability alone.
c) Erwin Schrodinger (refining models Bohr Atom)
Successfully prepared for electron wave equation using the principle of wave mechanics. Electrons surrounding the nucleus contained in an area that is 3-dimensional orbital around the nucleus where electrons with a specific energy can be found with the greatest possible.
• Orbital described as a cloud of electrons, namely: the forms of space where an electron is likely to be found.
• The meeting of the electron cloud, the more likely the electron was discovered and vice versa.
ATOM BASIC DEVELOPMENT OF PARTICLE
(Learn Chemistry 1A Package Book page 88 to 96!)
Notation Particle Mass Charge
Indeed Relative Relative thd thd fact proton proton
Proton 1.67 x 10-24 g 1 sma 1.6 x 10-19 C +1
Neutron 1.67 x 10-24 g 1 sma 0 0
Electron 9.11 x 10-28 g sma -1 -1.6 x 10-19 C
Note: the mass of elementary particles expressed in atomic mass units (sma).
1 sma = 1.66 x 10-24 g

ATOMIC NUMBER
• State the number of protons in an atom.
• For neutral atoms, the number of protons = number of electrons (atomic number also said the number of electrons).
• Given a symbol of the letter Z
o The atom loses electrons turn into a positive ion, which receives electrons instead turned into negative ions.
Example: 19K
Meaning ..............
MASS NUMBER
v Shows the number of protons and neutrons in the atomic nucleus.
v Protons and neutrons as particles making up the nucleus are called nucleons.
v The number of nucleons in an atom of an element expressed as mass number (given the symbol of the letter A), so that:
A = mass number
= Number of protons (p) + number of neutrons (n)
A = p + n = Z + n
v Writing a single atom is equipped with atomic numbers on the left bottom and mass number in the upper left of the atomic symbol. The notation is called the nuclides.
Description:
X = symbol of atom A = mass number
Z = atomic number Example:
ION STRUCTURE
v An atom can lose / release electrons or receive / accept an additional electron.
Atoms that lose v / release electrons, will become positive ions (cations).
v Atomic receiving / accepting electrons, would be a negative ion (anion).
v In the Ion, that has changed is the number of electrons, while the number of protons and neutrons remain.
Example:
Species Proton Neutron Electron
Atomic Na 11 11 12
Ion 11 10 12
Ion 11 12 12
The general formula for calculating the number of protons, neutrons and electrons:
1). For neutral atoms nuclides:
: P = Z
e = Z
n = (A-Z)
2). For nuclides cation:
: P = Z
e = Z - (+ y)
n = (A-Z)
3). For nuclides anion:
: P = Z
e = Z - (-y)
n = (A-Z)
Isotopes, Isobar AND Isoton
1). Isotope
Are atoms of the same element (having the same atomic number) but different mass numbers.
Example:;;
2). Isobar
Are atoms of different elements (having different atomic number) but have the same mass number.
Example: the
3). Isoton
Are atoms of different elements (having different atomic number) but have the same number of neutrons.
Example: the
ELECTRON CONFIGURATION
(Learn Chemistry 1A Package Book pages 83 to 88!)
ü Distribution of electrons in atomic shells are called configurations.
ü Skin atoms of the first (closest to the nucleus) was given the symbol K, skin-to-2 given the symbol L ff.
ü The maximum number of electrons in each shell meets 2N2 formula (n = number of skin).
Example:
K shell (n = 1) up to 2 x 12 = 2 electrons
Skin L (n = 2) a maximum of 2 x 22 = 8 electrons
Leather M (n = 3) Maximum 2 x 32 = 18 electrons
Leather N (n = 4) a maximum of 2 x 42 = 32 electrons
Skin O (n = 5) a maximum of 2 x 52 = 50 electrons
Note:
Although skin O, P and Q can accommodate more than 32 electrons, but in fact the skins have not been filled.
Step-by-Step Writing Electron Configuration:
1. Stuffed shells from the K shell, then L ff.
2. Especially for the main group (group A):
Total number of periods = skin
Number of valence electrons = group number
1. The maximum number of electrons in the outer shell (valence electrons) is 8.
• The valence electrons contribute to the formation of bonds between atoms in a compound form.
• The chemical properties of an element determined by the valence electrons. Therefore, the elements that have the same valence electrons, will have similar chemical properties.
Example:
Atomic Number Element K L M N O
He 2 2
Li 3 2 1
Ar 18 2 8 8
Ca 20 2 8 8 2
Sr 38 2 8 18 8 2
Consider Table 3.3 Chemical Package Book 1A page 85!
Note:
• The electron configuration for elements of class B (class transitions) is slightly different from group A (main group).
• No additional electrons to fill the outer shell, but it fills the two outermost skin; such that the two outer skin that contains 18 electrons.
Example:
Atomic Number Element K L M N
Sc 21 2 8 9 2
Ti 22 2 8 10 2
Mn 25 2 8 13 2
Zn 30 2 8 18 2
Solved Problem Solved:
Do Exercise 3.3 from Book Packages page 88 number 1 - number 5!
RELATIVE ATOMIC MASS (Ar)
(Learn Chemistry 1A Package Book pages 38 to 39 and pages 100 to 103!)
• Is the ratio between the mass of one atom to another atom.
• In general, consists of several isotopes of an element in the determination of the relative atomic mass (Ar) used the average mass of the isotopes-isotopes.
• According to the IUPAC, as the comparison used the C-12 atom of mass 1 atom C-12; thus formulated:
Ar element X = ........................ (1)
• Due: mass of 1 atom C-12 = 1 sma; then:
Ar element X = ........................ (2)
Relative molecular mass (Mr)
• Is the ratio between the mass of a molecule with a standard.
• The magnitude of relative molecular mass (Mr) of a substance = number of relative atomic mass (Ar) of the atoms making up the molecules of these substances.
• Especially for ionic compounds used term Relative formula mass (Mr) for ionic compounds are not made up of molecules.
• Mr = S Ar
Example:
Given: the relative atomic mass (Ar) H = 1 and C = 12, N = 14 and O = 16.
What is the relative molecular mass (Mr) of CO (NH2) 2
Answer:
Mr. CO (NH2) 2 = (1 x Ar C) + (1 x Ar O) + (2 x Ar N) + (4 x Ar H)
= (1 x 12) + (1 x 16) + (2 x 14) + (4 x 1)
= 60
1. B. DEVELOPMENT OF PERIODIC SYSTEM ELEMENTS
(Learn Chemistry 1A Package Book page 72 to 82!)
1). Law Triade Dobereiner
• Presented by Johan Wolfgang Dobereiner (German).
• The elements are grouped into groups of three elements called Triade.
• Essentially: the similarity of the physical and chemical properties of these elements.
Type Triade:
1. Triade Lithium (Li), Sodium (Na) and Potassium (K)
Atomic Mass Elements of Being
Li 6.94 Solid
Solid Na 22.99
K 39.10 Solid
Atomic Mass Na (Na Ar) == 23.02
1. Triade Calcium (Ca), strontium (Sr) and Barium (Ba)
2. Triade Chlorine (Cl), Bromine (Br) and iodine (I)
2). Newlands Law of Octaves
v Presented by John Newlands (UK).
v The elements grouped by rising relative atomic mass (Ar).
v-8 element to have chemical properties similar to the first element; Element 9 has properties similar to the second element and so on.
v The properties of elements found after 8 regular or periodic element called the Law of Octaves.
H Li Be B C N O
F Na Mg Al Si P S
Cl K Ca Ti Cr Mn Fe
Based on the list above Newlands Octave; elements of H, F and Cl have similar properties.
3). Mendeleev Periodic System (Short Periodic System)
ü Two chemist, Lothar Meyer (Germany) and Dmitri Ivanovich Mendeleev (Russia) based on the principle of Newlands, perform classification elements.
ü Lothar Meyer prefers the chemical properties of elements while Mendeleev prefers atomic mass increases.
ü According to Mendeleev: the properties of elements are a periodic function of the relative atomic mass. Meaning: if the elements are arranged according to their relative atomic mass increases, the specific properties will be repeated periodically.
ü The elements that have similar properties are placed on a single upright column, called Group.
ü The horizontal rows, for elements based on the relative atomic mass increases and called Period.
4). Modern Periodic System (System Periodic Length)
• Presented by Henry G Moseley, who argued that the properties of elements are periodic functions of their atomic numbers.
• Meaning: the basic properties of an element are determined by its atomic number rather than by the relative atomic mass (Ar).
1. C. PERIOD AND GROUPS IN MODERN SPU
1). Period
• A horizontal rows of the periodic table.
• Modern SPU consists of 7 periods. Each period stated amount / number of atomic shell elements that occupy those periods.
Period Number = Total Skin Atom

So:
• The elements that have 1 skin (skin K only) located in period 1 (row 1), the elements of which have 2 skins (skins K and L) lies in the second period and so on.
Example:
9F: 2, 7-second period
12mg: 2, 8, 2 the 3rd period
31Ga: 2, 8, 18, 3 period-4
Note:
a) Period 1, 2 and 3 is called the short period since it contains relatively few elements.
b) Period 4 and so called long period.
c) Period 7 called period is incomplete because not up to class VIII A.
d) To determine the period number of an element by atomic number, you only need to know the atomic number of the element that start each period.
2). Group
• periodic system consists of 18 vertical columns are divided into 8 main groups (group A) and 8 classes of transition (group B).
• The elements that have the same valence electrons are placed in the same group.
• For the elements of class A according to its place in the periodic system:

Group Number = Number of Valence Electrons

• The elements of class A has another name, namely:
1. Group IA Alkali class =
2. Group IIA = class of Alkali Soil
3. Group IIIA = class of Boron
4. Group IVA class = Carbon
5. Group VA = class of Nitrogen
6. Group VIA = class of Oxygen
7. Group VIIA = class Halide / Halogen
8. Group VIIIA class = Noble Gases
1. D. PERIODIC PROPERTIES ELEMENTS
Includes:
1). Jari-Jari Atom
• It is the distance from the nucleus to the electrons in the outer shell.
• The size of the atomic radius is influenced by the size of the atomic number of the element.
• The larger the atomic number of the elements of a class, the more the number of electron shell, so the larger the atomic radius.
So: in one group (from top to bottom), the atomic radius increases.
• In a period (from left to right), atomic number increases, which means the increasing nuclear charge, while the number of electrons fixed skin. As a result of the pull of the nucleus, the greater the outer electrons, resulting in increasingly smaller atomic radius.
So: in a period (from left to right), the atomic radius gets smaller.
2). Ionization Energy
ü It is the minimum energy required in the form of a gas of neutral atoms to release one electron to form charged ions +1.
ü If the atom releases two electrons that all it will take more energy (called the second ionization energy), and so on.
EI 1 <EI 2 <EI 3 ff
ü In one group (from top to bottom), EI is getting smaller as the atomic radius increases, so the core of the gravity of the smaller outer electrons. As a result, the outer electrons are easily released.
ü In a period (from left to right), EI greater because the atomic radius so small that gravity core to the outer electrons bigger / stronger. As a result, the outer electrons more difficult to remove.
3). Electron Affinity
• It is the energy released or absorbed by neutral atoms in gas form when receiving an electron to form negative ions.
• The price of a negative electron affinity, the easier atom receive / pull electrons and also the more reactive element.
• electron affinity is not the opposite of ionization energy.
• In one group (from top to bottom), the price of the smaller electron affinity.
• In a period (from left to right), the price of the greater electron affinity.
• main group elements have electron affinities are negative, except for groups IIA and VIIIA.
• has the largest electron affinity group VIIA.
4). Electronegativity
• It is the ability of an element to attract electrons in a molecule of a compound (in bond).
• Measured using Pauling scale of between 0.7 (electronegativity Cs) to 4 (electronegativity F).
• Elements that have a large electronegativity prices, will tend to accept electrons and form negative ions.
• Elements that have small electronegativity prices, will tend to lose electrons and form positive ions.
• In one group (from top to bottom), the price of the smaller electronegativity.
• In a period (from left to right), the price of the greater electronegativity.