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Periodic table

- The elements are classified into four types :-

* 1. Noble gases

They are found in p- block - in group zero.
They are stable ( completed last level).
k L

Example:

[10Ne] 1S2 2S2 2p6

* 2. The representative elements

They are found in s- block and p- block except (np6 and He).
They undergo reactions to complete their outer level.
k L M

Example:

(11Na) 1S2 2S2 2P6 3S1 3P 3d

* 3. The Transition elements

They are found in d- block.

They are incomplete the two highest levels nd,
(n+1) s

Example:

(21Sc)1S2 2S2 2p6 3S2 3p6 3d1 4S2 4p 4d 4f

* 4. The inner transition elements

They are found in f- block.

They are incomplete the three highest levels nf, (n+1)d , (n+2)S.

Example:

57La

* Description of the periodic table

It is consisted of 7 horizontal periods and 18 vertical groups.

Each period begins with new energy level and ends with Nobel gas.

Elements of the same group has the same configuration except the principle number.

1. The First period

It contains two elements-hydrogen 1S1 and He 1s2

The energy level is K.

2. The Second period

It consists of 8 element, 2 in s-block and 6 in p-block.

It begins with - Li3 (1 S2 , 2S1 )………. and ends with Ne10 (1 S2 , 2S2 ,2P6).

3- The Third period

It consists of 8 elements 2 in S and 6 in P

It begins with Na11, (1 S2 , 2S2, 2P6 ,3S1 )and ends with 18Ar

The sublevel 3d transferred to period 4 ( because it has higher energy than 4S) .

4- The fourth period

Its elements fill the forth level N and sublevel 3d - while 4d is transferred to period 5

They are 18 elements

It begins with k19 = [Ar18], 4S1 and Ca20 [Ar18 ], 4S2 Then 10 elements in d-block ( Transition elements).

Transition elements begin with Sc21=[Ar18],4S2 ,3d1 and end with 30Zn = [Ar18] , 4S2 , 3d10

Then 6 elements in p- block end with Kr36 = [Ar18] , 4S2, 3d10 , 4p6.

5- The fifth period

It begins with Rb37 = [ Kr36], 5S1 ,and end with Xe54 [Kr36], 5S2 , 4d10 , 5P6

They are l8 elements 2 in 5S -10 in 4d and 6 in 5p

6- The Sixth period

It has 32 elements -2 in 6S, - 14 in 4f (they are found under the table in order to be less length) -10 in 5d and 6 in 6P.

It begins with Cs55 = [ Xe54] , 6S1 and ends with Rn86 = [Xe54] , 6S2 , 4f14 , 5d10 , 6P6

The 14 elements in 4F are called Lanthanides (from inner transition elements).

7- The Seventh period

It has 2 elements in 7S and 14 in 5f(actinides ) and 6 in 6 d

It begins with Fr87 = [ Rn86] , 7S1 ….

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Trends and periodicity of properties

* 1- The atomic radius

- It is “Half the distance between the centers of two similar atoms in a diatomic molecule.

- Bond length

The distance between the nuclei of two bonded atoms.

We can measure the bond length by:
a)X-ray
b) Electron diffraction.

- In ionic bond - The bond length is the distance between the centers of cation and anion and it is equal the sum of the two ionic radii.

Ex

The bond length in Cl - Cl =1.98 A0
The bond length in C - C1 = 1 .76 A0
Find the atomic radius of carbon

Solution

The atomic radius of chlorine = = 0.99 A0

Bond length=atomic radius of carbon+a.r.of chlorine
1.76 = a.r. of carbon + 0.99
The atomic radius of carbon = 1.76 - 0.99 = 0.77 Aº

* Gradiation of atomic radius

a) In periods

The atomic radius decreases when the atomic number increases ( as we go right) Because the number of positive charge increases and the attraction force of the nucleus to electrons increases so the
radius decreases.

b) In groups

The atomic radii increase with the increase of atomic numbers

Because of :
1) Addition of an extra shell.
2) Screening effect of the inner filled orbital.

3) Increasing repulsive force between electrons.

N.B.

1- The addition of a new level exceeds the addition of one electron
2- The cation’s radius is smaller than that of the atom because of increasing pull of the nucleus.
3- The anion’s radius is bigger than that of the atom because of increasing the number of electrons.

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* 2- Ionization potential ( energy)

“The amount of energy required to remove the most loosely bound electron completely from an isolated gaseous atom.

- It is determined from spectral measurements

a) In Periods

The ionization energy is inversely proportional to atomic radius, it increases as the atomic number increases, because the valence electrons become more strongly bound to the nucleus and need higher energy to be separated

b) In groups

It decreases as the atomic number increases because
1- Addition of new shell and increasing the atomic radius
2- This completed shells screen the nuclear charge
- These facilitate the removal of outer electrons

N.B.

1- The ionization energy of noble gase : are very high because it is difficult to remove electron from a completely filled shell

2- The second ionization energy is greater than the first ionization energy due to increasing nuclear charge.

3- The third ionization energy is much bigger because it may need to break up completely filled shell.

Ex

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* 3- Electron affinity :

“The amount of energy released when an extra electron is added to a neutral gaseous atom”.

a) In periods

- The electron affinity increases as the atomic number increases

exception

-

The filled or half- filled orbital give the atom some stability so it is difficult to attract new electron and the electron affinity decreases

(As in 4Be = 1S2,2S2 , 7N = IS2,2S2,2p3, 10Ne =lS2,2S2,2p6)

b) In groups:

-

The electron affinity decreases as the atomic number increases due to increasing atomic radius increasing the repulsive force between the electrons and decreasing attraction force .

exception

Electron affinity of fluorine is less than that of chlorine due to the smallest Of the radius of flourine which causes high repulsive force between 9 electron of flourine.

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* 4- Electro-negativity

-

The average of the ionization energy and electron affinity.

The Electronegativity is

the tendency of an atom to attract the electrons of the chemical bond to itself”

Electron affinity refers to the single atom while electronegativity refers to the combined atom in molecule

a) In periods

- It increases with the increase of atomic number.

b) In groups

It decreases with the increase of atomic number.

The difference in the electronegativity determines the nature of
bond between two atoms.
Fluorine is the most electronegative elements

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* 5- Metallic and non metallic property

* Metals

Their valence shell has less than half its capacity

They are electropositive elements because they lose electrons and become positive ions(cations)Their valence electrons can move from one position to another in the atom so they conduct electricityThey have large atomic radii - but small ionization energy and small electron affinity.

Example

Na 1S2 2S2 2p6 3S1

* Non - Metals

Their valence shell has more than half its capacity .They are electronegative elements because they gain electrons and they become negative ions ( anions ) .Their valence electrons can’t move so they don’t conduct electricity .They have small atomic radii - but big ionization energy and big electron affinity

Example

Cl 1S2 2S2 2P6 3S2 3P5

* Metalloids

Their valence shell half filled
Their properties intermediate between metals, and non metals
Metalloids can unite with metals or none metals.
They conduct electricity less than metals but more than non - metals.
They are used as semiconductors in electronic instruments

Example

14Si 1S2 2S2 2P2 3S2 3P2

* Trends of the propety

a) In periods

The period begins with the highest metallic character then the metallic properties decrease as the atomic number increase, and non metallic properties increase.The highest metallic character in the left side while the highest non - metallic properties in the right side

Metalloids are found in the middle.

b) In groups

The highest non-metallic properties are found at the top in the right side and decrease as we go down.
The group in the left begins with metallic properties and increase as we move down.

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* 6- Acidic and basic properties

Acid

Non-metal oxide dissolves in water forming acid

Alkali

Metal oxide soluble in water forming alkali

Notes

1) Acid reacts with alkali forming salt and Water

2) Some oxides can react with both acids and
alkalis so they are called” Amphoteric oxides ”

Ex:

Al2O3

A)In period

-As the atomic number increases the basic properity decreases while the acidic property increases.

In groups

- The basic and acidic properties are both increase when the atomic numbers increase ( moves down )

- The strongest bases at the end left of the table

- The strongest acids at the end right of the table..

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* (7) Oxidation numbers:
- Valency
The number of single (unpaired)electrons in the valency shell of the atom.

Ex:
Nitrogen (N7) trivalent
- Unpaired electrons are counted in Valency because they only electrons of bonds.

* Oxidation number

“ The number that refers to the electric charges the atom would have in compound”.

- Rules for assigning oxidation numbers
* 1- In ionic compounds

· In cations: It is equal to its Valency preceded by (+ sign)
· In Anion: It is equal to its Valency preceded by - sign.
· (+ ve) sign refers to loosing electrons .
(- ve) sign refers to gaining electrons .
Ex

K1+ , Mg2+ , Cl1- , O2-
* 2- In Covalent Compounds

a) Diatomic molecules of similar atoms in electronegativity the oxidation number of each atom = zero (why?).
Diatomic molecules of different atoms in electronegativity the shared electrons are assigned to the more negative atom

N.B.

1- Oxidation number of hydrogen + 1 (except metal hydrides)

Ex:
NaH - KH. it equals -l

2- Oxidation number of oxygen = -2
(except few compounds such as H2O2 and Na2O2 = -1
,KO2 it equals -1/2 , OF2 = +2


* Basic rules

1) In neutral molecule algebraic sum of oxidation numbers of all atoms = 0

2) The oxidation number of an element in elementary form = 0

3) The oxidation number of polyatomic ion = The charge on the ion.

Oxidation number is preferred to Valency as it helps to explain the change in the electron structure of the reacted atoms and helps to tell the type of chemical change during chemical reactions

N.B.
Oxidation loosing electrons(increase(+ve)charges).
Reduction Gaining electrons(increase(-ve)charges).

Ex.1
- Explain the type of change(oxidation and reduction).

fecl2-----------fecl3
(1) Iron

Fe Cl2 = 0 reactants.

= Fe + 2C1 = 0 as Cl = -1

= Fe + [2 (-l)] = 0



FeCl3= 0
Fe + (3x-1) = 0





So the iron has been oxidized

Ex.2

Find the type of change that occurred to Cr and Fe in :-
K2Cr2O7 + 6FeCl2 + 14HCl → 2KCl+2CrCl3 + 6FeCl3 +7H2O
1-Chromium

2-Iron

* Gradation of oxidation numbers

1) Groups IA, IIA, IIIA
Their electron configuration nS1 , nS2 , np1
They lose the valence electrons to form positive ions
Their oxidation numbers are + 1, +2, +3
Ex.: Na1+Cl - Mg2+ Cl2 - A13+Cl3

2) Groups IVA, VA, VIA, VII A
( gain electrons forming -ve ions.)

1- The highest oxidation number equals group number

2- The lowest oxidation number = group number -8

3) Group zero: Their oxidation number = Zero




* The discovery of freon

Ammonia, sulphur dioxide and propane were used as cooling agents.

But :- Ammonia is poisonous
Sulphur dioxide is poisonous and corrosive.
propane is vigorouslly inflammable.

- Thomas Medgely

discover the freon andobserved that(in studying periodic table)
Non-metals forms gaseous compounds at normal temp. and their ability to ignite decrease on moving to the right in the periodic table
In addition halogen compounds [ CCl4 ] were used to put out fire.
Medgely synthesized freons

- Freons
Compounds containing carbon - fluorine and chlorine such as CF4, CCl2 F2.
Freons are used in fridges and air conditioners.

N.B
Freons Cause the decay of the ozone layer which protects the earth from harmful rays (ultraviolet)
- so scientist look for other agents

[masa654]

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