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نتائج و أخبار متعلقة بالتعليم بكافة مراحلة : الابتدائي والاعدادي والثانوي واخبار النتائج
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كيمياء لغات - الباب الرابع| Elements of group (I A) [alkali metals] للإطلاع على الدرس أضغط هنا
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كيمياء لغات - الباب الثالث| CHEMICAL COMBINATION
02 Theories Of Covalent Bondig
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كيمياء لغات - الباب الثالث| CHEMICAL COMBINATION
CHEMICAL COMBINATION
Introduction: for an element to be stable, its outermost energy level should be completely filled, so that all elements (except noble gases) undergo chemical reactions to acquire an identical electronic configuration. as that of the nearest noble gas, that’s by losing , gaining or sharing, through forming chemical bonds.
*If a chemical bond is not formed or broken amongst the atoms , there is no
chemical reaction
For example :- *If Iron fillings mixed with sulphur powder , the product
is not a chemical compound but it is a mixture
*If this mixture is heated to high enough temp. it forms new
compound and this is called chemical reaction
Chemical reaction:-
Def :- Breaking bonds between atoms in reactants and forming new bonds
between atoms in products
These bonds are: - 1) Ionic bond 2) Covalent bond 3) Coordinate bond
4) Hydrogen bond 5) Metallic
1) IONIC BOND
Def : - * It’s a bond formed by transfer one or more electrons
from metal to non-metal
* Electrostatic attraction force takes place between cation and anion. So ionic bond has no materialistic existence.
Steps 1) Metal loses one or more electrons to form positive ion[ cation]
2) Non-metal gains that electron to form negative ion (anion)
3) Cations and anions are attracted to each other forming ionic
compound.
EX.
Na 1S2 , 2S2, 2P6 , 3S1 17Cl 1S2 , 2S2, 2P6 , 3S2,3P5
Na+1 1S2 , 2S2 ,2P6 Cl-1 1S2 , 2S2 , 2P6 , 3S2 , 3P6
Attraction force
Na+1 ــــــــــــ Cl-1
G.R.F
1- Ionic bond is formed between the two extremes of the periodic table.
2-Elements of group 1A form Ionic bond with elements of group 7A
*Because elements at the left hand side of the periodic table
characterized by a large volume and a small ionization energy so electrons
easily loosed from the outer level forming cation.
but elements at the right hand side of the periodic table characterized by
small volume and high electron affinity so electrons easily gained [anion].
* Effect of electronegatirity on the strength of the ionic bond
* There is a relation between the difference in electronegativity between atoms in ionic compound and the strength of ionic bond.
* When the difference in electronegativity increases the strength of ionic bond increases. as shown in the table
* Ionic compounds are formed if the difference in electronegativity between the two bonded atoms is more than (1.7).
Group I II III
Element Sodium Magnesium Aluminium
Electro negativity 0.9 1.2 1.5
Chloride compound NaCl MgCl2 AlCl3
Diff in E.n 3- 0.9 = 2.1 3 –1.2 = 1.8 3 – 1.5 = 1.5
M.p 810C0 714 C0 190 C0
B.p 1465 C0 1412 C0 Sublimes'
Conductivity Very good conductor Good conductor Does not conduct
2) COVALENT BOND
Def :- It’s a bond formed between two non-metals by sharing electrons in order to reach the stable electronic configuration as the nearest noble gas.
* Types of covalent bonds
Pure covalent bond Polar covalent bond
It is formed between two similar atoms have the same electronegativity
EX H2 , O2 , N2 It is formed between two different atoms have the difference in electronegativity
HCl , NH3 , H2O
Diff. in electronegativity between atoms equals zero Diff in electronegativity between atoms is less than 1.7
Every atom in the molecule has the same ability to attract electrons of the chemical bond The atom of more electronegativity has greater ability to attract electrons of the chemical bond
The electron pair spends the same time in the vicinity of each atom The electron pair spends more time in the vicinity of each atom
*When difference in electronegativity equals zero bond will be pure covalent
*When difference in electronegativity is less than 1.7 bond will be Polar covalent
*When difference in electronegativity more than 1.7 bond will be ionic
G.R.F
1- The bond formed between two Fluorine ,Chlorine or Hydrogen atoms in
a diatomic molecule is a pure covalent bond.
*Because each atom in the molecule has the same ability to attract the two shared electrons of bond so the electron pair spends the same time in the vicinity of each atom and the net charge = zero.
4- The bond formed between Hydrogen and Chlorine atoms in Hydrogen
Chloride molecule is a polar covalent bond.
*Because Chlorine has a higher electronegativity than Hydrogen so it carries a partial negative charge and Hydrogen carries a partial positive charge due to unequal sharing of the electrons pair .
* Comparison between ionic bond and covalent bond
Ionic-bond Covalent-bond
1 Formed between metal and non-metal Formed between non-metals only
2 one or more electrons transfer from metal to non-metal No transferring electrons but sharing takes place between two non-metals.
3 Ions are formed ions are not formed
4 Electrostatic attraction takes place between cations and anions. Sharing of electrons takes place between the bonded atoms.
5 Has no materialistic existence (just attraction) G.R.F Has a materialistic existence by the shared electrons.
6 Difference in electronegativity is more than 1.7 Difference in electronegativity is less than 1.7
7 Strong bond G.R.F Weak bond
THEORIES OF COVALENT BONDING
1) Octet-rule “Electronic Theory of Valence” E.T.V
“With the exception of Hydrogen , Lithium and Beryllium ,atoms of all elements tend to reach the octet structure ”
By scientists named (Kossel and Lewis 1916 )
EX
H2 1S1 H H EX H2O 8O 1S2 , 2S2 , 2P4
H O H
EX 17Cl 1S2,2S2,2P6,3S2,3P5
Cl Cl EX NH3 7N 1S2,2S2,2P3
H N H
H
* Inadequacies of octet rule :- [Exam 2007]
1) It can’t explain the bonding in some molecules such as
boron trifluoride BF3 and phosphorus pentachloride. PCl5
2) It is not sufficient to explain many of covalent properties
such as the stereo structure and angle between bonds
G.R.F Boron triflouride and phosphorus penta chloride disobey octet rule.
Which of the following compounds obey or disobey the octet rule
Sulphur trifluoride SF3
Sulphur hexafluoride SF6
Sulphur dioxide SO2
Phosphorus pentaoxide P2O5
Phosphorus trichloride PCl3
Carbon dioxide CO2
2) Valence Bond Theory V.B.T
*This theory was based on the conclusions of quantum mechanics that consider electron not only as negative particle that moves in a definite orbit , but as a material with wave property which can exist in any position in the space around the nucleus
Def.:- It explains the formation of the covalent bond as a result of overlap of an atomic orbital of one atom contains single electron with orbital of another atom contains single electron.
* HYBRIDIZATION
*It’s the mixing of orbitals of the same atom closed in energy forming number of equivalent hybridized orbitals which take part in a chemical bonding.
* Types of Hybridization
Point of comp. Methane CH4 Ethylene C2 H4 Acetylene C2 H2
Type of hyb. Sp3 Sp2 Sp
Number of hyb. orbitals S + 3(P) = 4SP3 S + 2(P) = 3SP2 S + P = 2SP
Shape tetrahedral pyramid
[tetrahedron] planner triangle linear planner
Angle between the hybridized orb. 109o.28 120o 180o
Structure
H
H C H
H H H
C C
H H H ــ C C ــ H
Bonds between 2C - have one pi and one sigma in c-c two pi bonds and one sigma bond
number of bonds four sigma bonds five sigma bonds and one pi bond three sigma bonds and two pi bonds
Note:-
*The hybridized molecular orbitals must protrude to the outside
to be more capable of overlapping than the pure atomic orbital.
HYBERIDIZATION OF METHANE CH4
Ground state 2P2
2S2
6C 1S2
Excited state 2P3
2S1
1S2
Hyb. state
1S2
HYBERIDIZATION OF ETHYLENE C2 H4
2P2 2P2
2S2 2S2
6C 1S2 1S2 6C
2P3 2P3
2S1 2S1
6C 1S2 1S2
pz Pz
1S2 1S2
H H
C C
H H
HYBERIDIZATION OF ACETYLLENE
2P2 2P2
2S2 2S2
6C 1S2 1S2 6C
2P3 2P3
2S1 2S1
6C 1S2 1S2 6C
1S2
Shape:-
Linear planner
Angle:-
180o
Type:- SP
EX. Acetylene or Ethyne
*Conditions of hyberidization
1-Hybridization takes place between orbitals of the same atom.
2-Hybridization takes place between orbitals which closed in energy.
3-Number of hybridized orbitals equals orbitals which share in
hybridization and takes its symbols.
G.R.F
1- Angle between hybridized orbitals in Methane equals to 109.28
2- Angle between hybridized orbitals in Ethylene equals to 1200
3- Angle between hybridized orbitals in Acetylene equals to 1800
* Because the orbitals must go a part as far as possible from the
other orbitals to decreases the repulsion force between orbitals and
at this angle the molecule will be more stable. [ 2006 ]
SP3 hyb. :- it is the common hybridization takes place on mixing one
S orbital and 3p orbitals
3) MOLECULAR ORBITAL THEORY
M.O.T
Def:- All the atomic orbitals of the combined atoms are mixed to form molecular orbitals. which have the symbols (sigma , pi , delta )
V.B.T M.O.T
The molecule is formed of two atoms or more. The molecule is a one unit as a big atom with multi nuclei.
Hybridization occurs between some atomic orbitals. Hybridization occurs between all atomic orbitals.
* Compare between ( sigma ) and ( pi ) bonds.
Sigma bond Pi bond
1 Usually formed due to overlap of hybridized orbitals (head to head) Usually formed due to overlap of hybridized orbitals (side by side)
2 Strong bond G.R.F Weak bond
3 Due to greater orbital overlap Due to less orbital overlap
4 The overlapped orbital are on
one line (collinear-overlap)
along one axis Overlap between two parallel orbitals (collateral overlap)
3) COORDINATION BOND
Def :- It’s a special type of covalent bond in which the two electrons required for the bond are donated by one of the two atoms. to the other atom.
Or It’s a special type of covalent bond formed between two atoms one of them is donor and other is acceptor.
Or It’s a special type of covalent bond formed between two atoms one of them has lone pair of electron and other has empty orbital.
The atom, which gives that pair of electrons is called donor atom.
The atom, which gains that pair of electrons is called acceptor atom.
EX. (1) – Hydroxonium[Hydronium] ion [ H3O]+1
H O +H+ [ H O H+]
H H 8O 1S2 , 2S2 , 2P4
h4
* Oxygen is a donor atom
* Hydrogen ion is an acceptor
EX. (2) Ammonium ion [ NH4]+
H H
H N +H+ [ H N H+
H H 7N 1S2 , 2S2 , 2P3
* Nitrogen is a donor atom
* Hydrogen ion is an acceptor atom
Compare between Covalent bond and Coordination bond
Covalent bond Coordination bond
DEFINITION
SOURCE OF ELECTRONS OF CHEMICAL BOND
Types of chemical bonds in ammonium ion
1- Polar covalent bond 2-Coordination bond
Types of chemical bonds in ammonia solution
1- Polar covalent bond between Nitrogen and Hydrogen atom
2- Coordination bond between Nitrogen atom and Hydrogen ion .
3- Ionic bond between positive ion and negative ion.
4) THE HYDROGEN BOND
Def :- It’s a bond formed between hydrogen in a polar molecule and atom in other molecule has a high electronegativity such as Nitrogen , Oxygen and Fluorine.
*So Hydrogen is a bridge between two bonds one of them is
a polar covalent and other is hydrogen bond
Shapes of hydrogen bond
( linear – closed – open net shaped )
EX : [ water , liquid hydrogen fluoride , ammonia. ]
1) H2O Hـــ O…… Hــــ O ……… Hـــــ O……… HـــO
H H H H
F
2) HF F F F F F H H
H H H H H F F
H Closed ring
H H H
3) NH3 H N H N H N
H H H
G.R.F H2O (18) M.w H2S (34) M.w
Boiling point of water is higher than boiling point of H2S (hydrogen sulphids) OR [Anomalous of boiling point of water]. G.R.F
*Oxygen atom has a small atomic radius and high electronegativity[3.5] than
hydrogen [2.1] So water is a polar molecule So hydrogen bonds between
water molecules are formed which need amount of energy for broken
*Compare between covalent bond and hydrogen bond .
Point of comparison Covalent bond Hydrogen bond
Length 1A0 [ Shorter ] 3A0 [ Longer ]
Energy 418 Kj [ stronger ] 21 Kj [ weaker ]
Def.
5) METALLIC BOND
*The positive ions of the metal are hold together by the moving electron cloud.
Def :- It’s produced from the electron clouds of valence electrons which decreases the repulsion forces between the positive metal ions in the crystal lattice.
11Na 1S2 , 2S2 , 2P6 , 3S1
* It contains (1) electron in delay level
12Mg 1S2 , 2S2 , 2P6 , 3S2
* It contains (2) electrons in delay level
13Al 1S2 , 2S2 , 2P6 , 3S2 , 3P1
* It contains (3) electrons in delay level
G.R.F * Aluminum is hard metal * Manganese is mild * Sodium is soft
If No. of valance electrons increases “strength of metallic bond increases” and melting point increases. [Sodium has one electron in the outer level but Magnesium has two electrons but Aluminum has three electrons]
G.R.F M.P of Al greater than Magnesium greater than Sodium
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كيمياء لغات - الباب الثانى | Strength of oxygenated acids
01 Strength of Oxygenated Acids
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كيمياء لغات - الباب الثانى | Strength of oxygenated acids
Strength of oxygenated acids
*It depends on the number of Oxygen atoms which does not bind with
Hydrogen atoms
*Oxygenated acid have the formula MOn (OH)m
M atom of an element
On atoms of Oxygen which do not bind with Hydrogen atom
m number of oxygen bind with hydrogen
Acid
MOn (OH)m Oxygen atoms which not bind with Hydrogen Strength of acid Chemical formula
Si (OH)4
ortho siliconic _____ weak acid H4SiO4
PO (OH)3
orthophosphoric 1 moderate H3PO4
SO2 (OH)2
Sulphuric 2 strong H2SO4
ClO3 (OH)
Perochloric 3 very strong HClO4
G.R.F 1-Perochloric acid HClO4 more stronger than orthophosphoric acid.
2- Perochloric acid more stronger than sulphuric acid. Exam 2012
3- HClO4 is stronger than H4SiO4 [ orthosiliconic acid]
………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………
4- Strength of oxygenated acids increases from left to right across the period.
………………………………………………………………………….....................
7) Graduation of oxidation number in the periodic table.
* The old definition of valence: -
Number of hydrogen atoms which combines with one atom of element.
Number of hydrogen atoms that can be replaced by one atom of element.
Ex. HF H2O NH3
(Flourine is monovalent) (Oxygen is divalent) (Nitrogen is trivalent)
* The modern definition of valence: -
Number of ( unpaired ) electrons in the valence shell of an atom.
7N 1S2 , 2S2 , 2P3 Nitrogen is trivalent because
It has three unpaired electrons G.R.F.
8O 1S2 ,2S2 , 2P4 Oxygen, is divalent because
it has two unpaired electrons.
•
* Oxidation number: -
Def * It’s a number refers to the electric charges
(positive or negative) of elements in compound.
* Difference between [oxidation number and valence]
* Whose meaning is similar to valence. but the term oxidation number is preferred than valence because it helps us to explain the change in the electronic structure of atoms in compounds. G.R.F.
Example Valence Oxidation no. Example Valence Oxidation no.
Na+1 1 +1 Cl-1 1 -1
Mg+2 2 +2 O-2 2 -2
Calculation’s rules of oxidation numbers
1) The oxidation number of an atom in its elemental state equals zero.
Ex. [ O2 , O3 , N2 , S , H2 , O3 , S8 , P4 , Cl2 ]
2) Sum. oxidation numbers of any compound is zero H2 S O4
3) The oxidation number of Oxygen equals –2 except that
Oxygen diflouride Hydrogen peroxide Sodium peroxide
O F2 H2O2 Na2O2
O F2 H2 O2 Na2 O2
O = +2 O = -2/2 = -1 O = -2/2 = -1
4) The oxidation number of Hydrogen (+1) except [metal hydrides]
Ca H2 NaH
Ca H2 Na H
H = -2/2 = -1 H = -1
*On electrolysis of Sodium hydrides , Hydrogen gas evolves at anode
because oxidation number of Hydrogen in hydrides = -1 G.R.F
5) In the formula for an atomic group (radical) the sum of oxidation numbers equals the charge of the group. For example, in the radical.
(Cr2 O7 ) , the sum of all oxidation number is = -2
Calculate oxidation number in each of the following
1- Oxygen in {OF2 , KO2 , Na2O2 , Li2O , O3 , O2 }
2- Chlorine in { NaCl , NaClO4 , NaClO3 , NaClO2 , NaClO }
3- Nitrogen in { NO2 , HNO2 , NO , N2O , N2 , NH3 , ( NO3)-1,(NO2)-1
4- Sulphur in {Na2S2O3 , K2S , SO2 , NaHSO3 , H2SO4 , (SO4)-2
5- Manganese in {KMnO4 , MnO2 , (MnO4)-1 , MnCl2 }
6- Iron in {FeSO4 , Fe2(SO4)3 , }
7- Sulphur in { ferric sulphate , ferrus sulphate }
Fe2(SO4)3 FeSO4
*Oxidation number of some elements:-
Group 1A 2A 3A
O.N +1 +2 +3
Ex Li ,Na , K , Rb , Cs Be , Mg ,Ca , Ba Al
Oxidation :-
*It is the process of losing electrons leading to increasing the positive
charges or decreasing the negative charges
Reduction :-
*It is the process of gaining electrons leading to increasing the negative
charges or decreasing the positive charges
Oxidizing agent :-
Substance which takes electron , takes hydrogen or gives oxygen
Reducing agent :-
Substance which gives electron , takes oxygen or gives hydrogen
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كيمياء لغات - الباب الثانى | Trends the metallic and non-metallic properties
01 Trends The Metallic and non metallic Properties
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كيمياء لغات - الباب الثانى | Trends the metallic and non-metallic properties
5) Trends the metallic and non-metallic properties
*Berzelius classified elements into metals and non-metals
Points of comparison Metals Non-metals
1) No. of electrons in the valence level. less than half the maximum (Less than 4 ) more than half the maximum (more than 4)
2) Position in the periodic table. at the beginning of periods. at the end of periods.
3) Atomic radius. relatively large. relatively small.
4) Ionization potential. relatively small. relatively large.
5) Electron affinity. relatively small. relatively large.
6) Electronegativity. relatively small. relatively large.
7) Losing or gaining electrons. atom loses electrons of the valence shell to reach the stable structure of nearest noble gas. atom accept electrons of the valence shell to reach the stable structure of the nearest noble gas.
8) Kind of ions. the ions are (+ve) (Electro positive elements) the ions are (-ve) (Electro negative elements)
9) Electric conductivity high electric conductivity bad electric conductivity
Metalloids:- Elements have the metallic appearance of metals but the most properties of non-metals at the same time.
Properties of metalloids: -
1) Their valency shell is nearly half filled.
2) Electronegativity is intermediate between metals and non-metals .
3) Their electrical conductivity is less than metals but more than
non-metals , so they are called semiconductors
4) They have important uses in electric instruments such as
semiconductors and are known as transistors .
Ex:- ( Boron – Silicon – Germanium )
G.R.F
1) Metals are good conductors of electricity
Because: - They have free valence electrons (easily transferred).
2) Non-metal are bad conductors of electricity or insulators .
Because: -Their valence electrons difficult transferred(strongly bounded).
3) Metals have small values for ionization energy and electron affinity.
Because: - Metals have relatively large atomic radius.
(found at the beginning of periods)
4) Non-metals have a high values of ionization energy and electron affinity.
Because: - Non-metals have relatively small atomic radius
( found at the end of periods )
5) Metals are called electropositive elements while non-metals are called electronegative elements.
Because: - Metals tend to lose electrons forming ( + ve ions ).
but non-metals tend to gain electrons forming ( -ve ions ).
Variation of metallic and non-metallic properties in the periods:-
* Begin with a strong metal finish with a strong non-metal
* Between them, metalloids are present so metallic property decreases from left to right across the periods .and non –metallic property increases.
In groups: * The metallic property increases and non-metallic property
decreases from up to down.
G.R.F 1- Caesium has the most metallic properties
but flourine has the most non-metallic properties
*Because Caesium lies at the bottom on the left hand side of the periodic table So its atom lose the valence electron easily .but fluorine lies at the top right so it gains electron to complete its outer level.
* The strongest metal lies at the bottom of the left hand side of the
periodic table but the strongest non-metal lies at the top of right
hand side of the periodic table. Group 7A [ Halogens ]
6- Trends the acidic and basic properties
• Compare between acidic oxides and basic oxides
Acidic oxides Basic oxides
They are non-metallic oxides such as
CO2 , SO2 , P2 O5 , SO3 G.R.F They are metallic oxides such as
Na2O ,CaO , K2O , CuO, , MgO G.R.F
They dissolve in water forming acids Some dissolves in water forming alkalis
CO2 + H2O H2CO3 carbonic acid
SO3 + H2O H2SO4 sulphuric acid Na2O + H2O 2NaOH sod. hydroxide
K2O + H2O 2KOH
They can’t react with acids They can’t react with alkalis
They react with alkalis forming salt and water .
CO2 + 2NaOH Na2 CO3 + H2O They react with acid to forming salt and water .
Na2O + 2HCl 2NaCl + H2O
MgO + H2SO4 MgSO4 + H2O
* Alkali :- Water soluble basic oxide
Amphoteric oxides
Def * They are oxides which reacts with acid as a basic and react with base as acidic forming salt and water.
For example: (Al2 O3 , SnO , ZnO , Sb2O3 )
Aluminium oxide , Tin oxide , Zinc oxide ,Antimony oxide
*In the periods :Acidic property in oxides increases and basic property
decreases when atomic number increases from left to right.
* In the group 1A :-
Basic property increases from up to down due to
increasing atomic size of element while the charge remains constant
Explaining :-
*Acids and alkalis are considered hydroxy compounds
*They have the fromula MOH
*They are ionized by two different methods
1-It may be ionized giving OH- and considered base
MOH ( base ) M+ + OH-
2-It may be ionized giving H+ and considered an acid.
HMO ( acid) MO¬¬- + H+
M+
Repulsion force attraction force
H+ O-
attraction force
* When attraction force between M+ and O- greater than attraction
force between O- and H+ it is ionized giving H+
* When attraction force between H+ and O- greater than attraction
force between O- and M+ it is ionized giving OH-
*When the two attraction forces are equals it is ionized as acid or as
base depending on medium of reaction.
[ reacts with acid as base and with base as acid ]
*The above attraction force depends on atom of element according to
its size and its electric charge
* in alkali metals as Na
It is characterized by a large size and has one positive charge
[ small attraction force] so attraction force with O- decreases but attraction force between O- and H+ increases so it is ionized into OH-
* in Non-metals as Chlorine [ from left to right ]
It is characterized by a small size and has more charges
so attraction force with O- increases but attraction force between O- and H+ decreases so it is ionized into H+
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كيمياء لغات - الباب الثانى | C- TYPES OF ELEMENTS IN THE PERIODIC TABLE
01 Trends And Periodicity Of Properties in the periodic table
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كيمياء لغات - الباب الثانى | C- TYPES OF ELEMENTS IN THE PERIODIC TABLE
C- TYPES OF ELEMENTS IN THE PERIODIC TABLE
1- Noble gases (inert gases)
* Elements in which all their levels are completely filled.
so they are more stable. and chemically inactive. G.R.F.
* They are called zeroth group
* They are elements which have the electronic configuration np6 except Helium which has electronic configuration 1S2 .
2- Representative elements [ main groups elements]
*Elements which all their levels are completely filled except
the highest level , their electronic configuration ends by nS1-2 , nP1-5
* They are elements of S-block and P-block except zeroth group.
*Their highest levels tend to reach the completely electronic configuration
[nS2, nP6 ] by losing , gaining or sharing electrons.
3- Main transition elements
* They are elements of the d-block. (B Groups)
* d-sublevel of these elements are successively filled.
Classification of main transition elements
1st series 2nd series 3rd series
Transition series (3d) Transition series (4d) Transition series (5d)
*It includes the elements in which the sublevel 3d is filled successively.
* Placed in the fourth period.
* Consists of 10 elements. * It includes the elements in which the sublevel 4d is filled successively.
* Placed in the fifth period.
Consists of 10 elements. *It includes the elements in which the sublevel 5d is filled successively.
* Placed in the sixth period.
* Consists of 10 elements.
4- Inner-transition elements nf, (n+1)d and ( n+2 ) S
*They are elements of f-block , f-sublevel filled successively.
* These elements are divided into lanthanides and actinides
Lanthanides
have incomplete orbitals 4f,5d and 6S Actinides
have incomplete orbitals 5f,6dand 7S
* Placed in the sixth period.(after 57La)
* Consists of 14 (rare earth’s elements)
* Includes elements in which sublevel 4f is filled successively. * Placed in the seventh period.(after 89Ac)
* Consists of 14( radioactive elements)
* Includes elements in which sublevel 5f is filled successively.
G.R.F
1) Lanthanides are called rare earth’s elements
Because: They are very difficult to be separated
2) lanthanides are similar in behavior [ chemical properties ].
Because: They have the same valence electrons (6S2 )
3) Actinides are called radioactive elements
Because: Their nuclei are unstable
Exercise :-
* Write electronic configuration then determine :-
[ block – period – type – sublevels - orbitals ]
or determine [the location of elements in periodic table]
26Fe , 36Kr , 13Al , 20Ca , 2He , 19K , 18Ar
Write the electronic configuration according to nearest inert gas:-
58Ce
64Gd
90th
93NP
Types and numbers of elements in the periodic table
NUMBER OF PERIOD SUBLEVELS WHICH FILLED SUCC. NUMBER OF ELEMENTS TYPES OF ELEMENTS IN
EACH PERIOD
First 1S2 2 Representative and noble [2]
Second 2S2 , 2P6 8 Representative and noble [2]
Third 3S2 and 3P6 filled succ. 8 Representative and noble [2]
Fourth 4S2 , 3d10 and 4P6 filled succ. 18 Representative , and main transition
[first series] and noble [3]
Fifth 5S2 , 4d10 and 5P6 filled succ 18 Representative , and main transition [second series]
and noble [3]
Sixth 6S2, 4f14 , 5d10 and 6P6 filled succ 32 Representative , and main transition [third series ]
and inner transition [lanthanides ] and noble [4]
D- TRENDS AND PERIODICITY OF PROPERTIES IN THE PERIODIC TABLE
1) The atomic radius
* It’s incorrect to define the atomic radius as the distance from the nucleus to the farthest electron. G.R.F
Because: wave mechanics proved that :-
* It is impossible to determine exactly the location of electrons around
the nucleus. So the atomic radius cannot be defined and cannot be
physically measured
Atomic radius: * 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.
• Covalent atomic radius = bond length / 2.
Molecule Bond length in Angstrom Covalent atomic radius
H - H 0.60 Ao 0.30A0
F - F 1.28A0 0.64A0
Cl - Cl 1.98A0 0,99A0
Br - Br 2.28A0 1.14A0
I - I 2.66A0 1.33A0
Ionic crystals* Ionic crystals consists of positive ions and negative ions such as sodium chloride crystals and the radius between them called ionic radius.
Ionic radius :- Sum of the two radius of cation and anion
and it depends on the number of electrons losed or gained.
problems
Bond length of = covalent atomic radius of A + covalent atomic radius of B
compound A-B
Ex.1: The bond length in the chlorine molecule (Cl – Cl) equals
1.98 Ao and the bond length between carbon and chlorine atom (C – Cl) = 1.76 Ao calculate the atomic radius of carbon.
Solution
Bond length in C - Cl = covalent radius of C + covalent radius of Cl
1.76 = ?? + 1.98/2
1.76 = ?? + 0.99
1.76 - 0.99 = ??
0.77Ao = covalent atomic radius (of C )
Ex.2: The bond length in chlorine ( Cl – Cl ) = 1.98 Ao and the bond length in hydrogen chloride molecule ( H – Cl ) = 1.29 Ao
Calculate the covalent atomic radius of hydrogen
Ex.3: The bond length in nitrogen molecule ( N – N ) = 1.96 Ao and the bond length in hydrogen molecule = 0.62 Ao calculate:
a) Atomic radius of nitrogen. b)Atomic radius of hydrogen
c) Bond length between ( N ) and ( H ) in ammonia molecule NH3
Graduation of atomic radius in the periodic table
* We can observe the following from the periodic table.
In the horizontal periods :-
Atomic radius decreases from left to right across a period. G.R.F
Because :- * number of positive charges inside the nucleus increase and
attraction force of nucleus increases.
So atoms of group 1A have the largest size and group 7A have
the smallest size which called halogens.
In the vertical groups :-
Atomic radius increases as the atomic number increases
from up to down G.R.F
Because:
* Number of energy levels increases .
* Screening effect of an extra shell takes place.
* Repulsion force between electrons in the levels increases .
* Attraction force of nucleus decreases.
* Atomic radius increases from up to down across a group.
Obs. Increasing atomic radius when transfer from period to another more
than decreasing atomic radius when transfer from group to another.
Because the effect of increasing energy level more than the effect of
increasing one positive charge across a period.
Changes of cations and anions from their atoms
Na Na+1
1.75 Ao 0.95 Ao
A) Cations are smaller than their atoms
Or cationic radius smaller than atoms
Answer: due to increasing the pull of the nuclear charge on the remaining electrons in the cation.
*When number of positive charges of an ion increases , attraction force of the nucleus increases and atomic radius decreases. [ as in Iron ]
Fe
Fe+2 Fe+3
1.17 Ao 0.75 Ao 0.60 Ao
G.R.F
Radius of Fe+3 smaller than that of Fe+2 smaller than that of Fe atom
*Because when number of positive charges increase, attraction force of nucleus increases so atomic radius decreases.
b)Radius of anions are larger than their atoms
The radius of chloride ion is larger than the radius of its atom. G.R.F
Answer: due to increasing the number of electrons without increasing
the nuclear charge. [ tug of war ]
*Metals :- Element in which its ionic radius smaller than its atom.
* Non-metal :- Element in which its ionic radius larger than its atom.
2) Ionization energy (Ionization pot.)
I.E I.P
Def.: It’s the amount of energy required to remove the most loosely bound electron completely from an isolated gaseous atom.
*and it is determined by a spectral measurements so :-
*It’s possible to remove one, two, three or more electrons from atoms
*So, there are first, second, third ………. etc ionization energy.
* First ionization energy:
- It’s the energy required to convert an atom into ion with one positive charge. M M+1 + e * H = + ve KJ/ mol
* Second ionization energy: M+1 M+2 + 2 e-
- It’s the energy required to convert M+1 into M+2
G.R.F Second ionization energy is higher than first ionization energy
Variation of ionization energy in periods
Ionization energy increases from left to right across period G.R.F
. Because from left to right
*Attraction force of nucleus increases and atomic radius decreases.
Variation of ionization energy in groups
Ionization energy decreases from up to down. G.R.F
Because * No. of levels increases (repulsion force of electrons increase)
* Energy levels will spread out.
* Screening effect takes place
* Attraction force of the nucleus decreases
* Atomic radius increases..
1) Ionization energy of 4Be, 12Mg is high. G.R.F
or elements of group II A have high ionization energy )
Answer: 4Be 1S2, 2S2 12Mg 1S2, 2S2, 2P6, 3S2
* Because the outer most sublevel S is filled with 2 electrons so atom will be more stable so it must need large amount of energy to remove electron.
2) Ionization energy of Nitrogen (7N) , ( 15P ) is high
or elements of group 5A have high ionization energy.
Answer: 7N 1S2, 2S2, 2P3 15P 1S2, 2S2, 2P6, 3S3, 3P3
*Because the outer most sublevel P is half filled, [p3] so atom will be more stable so it must need large amount of energy to remove electron.
3) The first ionization energy of noble gases of zeroth group is very high.
4). The second ionization energy of elements of 1A group is high.
5)- The third ionization energy of elements of 2A group is very high.
*Because it must need large amount of energy to break down the completely filled level due to the stability of atom.
Ionization energies of Magnesium.
(2,8,2)Mg Mg+ + e H +737 k.j / mol
(2,8,1)Mg+ Mg+2 + e H = +1450
(2,8)Mg+2 Mg+3 + e H = +7730
* Ionization energy is inversely proportional to atomic radius.
*Compare between the first ionization energy and the second ionization energy
3) Electron affinity الميل الإلكترونى
Def: It is the amount of energy released (produced) when an extra
electron is added to a neutral gaseous atom .
X + e- X-1 + energy H= - ….. k.j mol-1
Variation in periods: * Electron affinity increases from left to right
*Due to decreasing atomic radius leading to easy to attract new electron
Variation in-groups: - Electron affinity decreases from up to dow
. *Due to increasing atomic radius leading to increasing the distance
between nucleus and valence level
* The magnitude of the electron affinity is high when the added electrons make the orbitals half filled or completely filled due to increasing the stability of the atom.
G.R.F
1) Beryllium has a relatively small electron affinity.
Answer:- Due to the stability of its atom., it has filled orbitals. 4Be 1S2,2S2
2) Nitrogen has a relatively small electron affinity
Answer:- Due to the extra stability of its atom. The outer sublevel is half filled. 7N 1S2 ,2S2 , 2P3
3) Neon has very low E.Aff
*Due to great stability of atoms because all levels are completely filled
4) Electron affinity of fluorine -339.9 kj/mol is less than that of chlorine
-361.7 although electron affinity decrease in vertical groups from up to down.
Because:- Fluorine has small size [ top right of periodic table ]so repulsion force takes place between the new electron and original nine electrons
4) Electronegativity
Def: - It’s the tendency of an atom to attract the electrons of
chemical bond to itself.
*Fluorine is considered to be the most electronegative element
* Diff. in electronegativity between elements plays an important role in determining the nature of chemical bonds [ chapter 3]
Electronegativity increases from left to right Electronegativity decreases from up to down
* Increasing the atomic number.
* Increasing the +ve nuclear charge.
* Increasing the attraction force.
* Decreasing the atomic radius.
* Increasing electron affinity.
* Increasing electronegativity. * Increasing the electron levels.
* Increasing the repulsion among electrons.
* Decreasing attraction force.
* Increasing the atomic radius.
* Decreasing electron affinity.
* Decreasing electronegativity.
1- Electronegativity of Oxygen is higher than nitrogen ( 8O ., 7N )
2- Electronegativity of Fluorine is higher than Bromine ( 9F , 35Br )
• Compare between electronaffinity and electronegativity
Electronaffinity Electronegativity
* Refers to the atom in its single state *Refers to combined atoms in molecule
*Term of energy *Not term of energy
* It is the amount of energy released when ………………………………. * It is the tendency of an atom to attract …………………………..
* Its unit is Kj / mole * unitless
Compare between ionization energy and electron affinity
الصف الثالث الثانوى كيمياء لغات الباب الثانى| كيمياء لغات
كيمياء لغات - الباب الثانى | The periodic table
01 Classification Of Elements And Periodic Table
الدرس في شكل نص مقروء
كيمياء لغات - الباب الثانى | The periodic table
table
*Arrangement of elements in an ascending order according to their atomic numbers and it agrees with sequence of building up principle
2P
3P
4P
5P
6P
7P
1S
2S
3S
4S
5S
6S
7S
3d
4d
5d
6d
4f
5f
A- BLOCKS OF THE PERIODIC TABLE
S- Block P- Block d- Block f- Block
It includes the elements whose outer most electrons occupy the s-sublevel
nS1,nS2 It includes the elements whose outer most electrons occupy the p-sub level nP1 to np6 It includes the elements whose outer most electrons occupy the d-sub level It includes the elements whose outer most electrons occupy the f-sub level
It Consists of the groups (1A) and (2A) It Consists of the groups (3A) , (4A) (5A) , (6A) , (7A),
(zero group) It Consists of the groups B which called (main transition elements) It Consists of lanthanides and actinides
[ Inner T.E ]
It’s placed in the left hand side of the table It’s placed in the right hand side of the table It’s placed in the middle of the table It’s placed below the table G.R.F
To reduce the table
B- DESCRIPTION OF PERIODIC TABLE
* It consists of seven ( horizontal ) periods and 18 vertical groups
Period one: - * It consists of two elements, hydrogen and helium.
Period two: - * It consists of eight elements G.R.F
Because: the sublevels 2S2 and 2P6 will be filled succ.
Period three: - * It consists of eight elements G.R.F
Because: the sublevels 3S2 and 3P6 will be filled succ.
Period four: - * It consists of 18 elements G.R.F
Because: the sublevels 4S2, 3d10 and 4P6 will be filled succ.
Period five: - * It consists of 18 elements G.R.F
Because: The sublevels 5S2, 4d10 and 5P6 will be filled succ.
Period six: - * It consists of 32 elements G.R.F
Because: 6S2 ,4F14, 5d10 and 6P6 will be filled succ.
Period :-
*Several elements are arranged horizontally according to their atomic numbers
which increases by one from left to right.
* Each period begins by filling a new energy level and ends by an inert gas.
Elements of the same period similar in number of energy levels but different in atomic number , valency level and chemical properties
Group:-
*Several elements arranged vertically and they are similar in valency level and chemical properties but different in the principal quantum number.
الصف الثالث الثانوى كيمياء لغات الباب الأول - كيمياء لغات
كيمياء لغات - الباب الأول| The principles[modifications ] of modern atomic theory
01 The Principles Of Modern Atomic Theory
الدرس في شكل نص مقروء
كيمياء لغات - الباب الأول| The principles[modifications ] of modern atomic theory
The principles[modifications ] of modern atomic theory
* modern atomic theory based on some essential modifications on Bohr's model . the most important modifications are :- [exam 2007]
1) Dual nature of electron. (De Broglie)
2) Uncertainty principle. (Heisenberg)
3) Finding the mathematical expression which describes the wave motion of electron, its shape and energy. (schrodanger)
1- The dual nature of electron
Def *Electron is a particle has a wave properties.
* Every moving body such as electron, nucleus of an atom or all molecules or tennis ball is associated with a wave motion called “matter waves”.
* The matter waves differs from electromagnetic waves in two points.
1) They are not separating from the moving body.
2) Their speed doesn’t equal the speed of electromagnetic waves.
*Compare between matter waves and electromagnetic waves.
2) Heisenberg uncertainty principle [probability]
He proved that by using quantum mechanics
“It’s impossible to determine both of the location and the speed of electron
partically at the same time.”
* Because the apparatus used will change either speed or location of
the electron.
3) Schrodinger's wave equation
* He depended on planck , Einestein , De Broglie and Heisenberg to put the
wave equation
By solving this equation: -
* It can be used to determine the allowed energy levels of the electron.
*It can be used to define the region of space around the nucleus where it’s most probable to find the electron in each energy level.
As the result of this equation.
a) Our concept of the orbit around the nucleus changed to use the term orbital vail (electron cloud) to describe orbital.
b) By mathematical solution of schrodinger equation quantum numbers are introduced.
Orbit concept in Bohr’s theory Orbital concept in schrodinger’s wave mechanics concept.
• It’s a stable circular orbit of particular radii in which the electron orbits.
* It’s the region of space around the nucleus in which great prabability finding electron and the electron cloud is used to describe any orbital
*Compare between orbit and orbital [ def and draw ]
[ draw energy graph of orbit and orbital ] [ Exam 2006 ]
Orbital : - It’s the region of space around the nucleus with a great probability Electron cloud of finding the electron.
* Presence of electron in all directionsa and dimentions around
nucleus
2-THE FOUR QUANTUM NUMBERS”
Def: - These are numbers, which define the volume [ maximum prbability for finding electron ] beside shape, energy, and direction of orbitals in an atom.
1) Principal quantum number (n) its values 1,2,3,4,…not zero
K , L , M , N , O , P , Q
n = 1 2 3 4 5 6 7
Def a- It is a quantum number which determine the order of principal energy
levels of the electron [ seven in the known heaviest atom in ground state ]
b- determine the number of electrons, which fill each
main energy level according to the rule
2n2 [ twice square the principal quantum number ]
G.R.F * The rule 2n2 can’t be applied on the energy levels higher than fourth
Because: - The atom becomes unstable if the number of electrons
exceeds 32 in any energy level.
G.R.F * The third energy level ( M ) is saturated with 18 electrons
Due to the rule 2n2 which determine the maximum number of electrons
* Each principal energy level is subdivided into a number of sublevels which defined by value of another quantum number called subsidiary Q.number
*Bohr had used the P.Q.n to explain the spectral line of hydrogen atom.
2)Subsidiary(orbital or azimuthal)quantum number (l)
Def: - * It is a quantum number which define the number of energy sublevels in each principal energy level.it is equals P.Q.N
“Each principal energy level contains energy sublevels equal its number”
* These sublevels take the symbols as follow.
Main energy level (N) No. of sublevels Type of sublevels No. of orbitals (n2) No. of electrons (2n2)
K
1 1S 1 2
L
2 2S
2P 4 8
M
3 3S
3P
3d 9 18
N
4 4S
4P
4d
4F 16 32
* Observation: - This is shown by the scientist Sommerfield when he used a spectroscope, which has a high resolving power than Bohr .
* He found that: - [The single spectral line has divided into number of fine spectral lines]so “main energy levels contains energy sublevels equal its number
3) Magnetic quantum number (m) [2L+1]
Def :- It is a quantum number which defines the number of orbitals in
each sublevel.
*Each sublevel contains odd number of orbitals.
*Each orbital is saturated with 2 electrons.
*Orbitals of the same sublevel are similar in shape , energy , size but
different in direction
* S has (one) orbital (spherical symmetrical shape around the nucleus)
*P has (three) orbitals ( Px , Py , Pz )( filled with six electrons)
* d has (five) orbitals (So it is filled with ten electrons) G.R.F
*F has (seven) orbitals (So it is filled with 14 electrons) G.R.F
* It has spherical symmetrical shape around the nucleus. * Each P orbital is perpendicular to each other .
* The electron cloud of each orbital takes the forming of two pears meeting head to head
at point called node has no electron density.
[dumb-bell shape ]
Draw the shapes of
S and P orbitals P
Px Py Pz
4) The spin quantum no. (ms)
Def :- It is a quantum number which defines the spin motion of electron around its
axis during its rotation around the nucleus.
* The electron can rotate around its axis in either a clockwise or anti-clockwise
Although the two electrons of the same orbital carry the same negative charge,
they don’t repel. G.R.F
Because one of them rotates clockwise and other rotates anticlockwise So
The magnetic field of one electron is opposite to the magnetic field of other.
* it is said that the two electrons are in a spin paired state
G.R.F d- sublevel is filled with ten electrons but f-sublevel is filled with 14 electrons
……………………………………………………………………………………….
………………………………………………………………………………….
principles of distributing electrons
Aufbau (building-up) principle Hund’s rule
1) Aufbau (building-up) principle: -
“Electrons occupy the sublevels in the order of increasing their energy" [ The lowest energy sublevels are filled first” ]
* The sequence of energy sublevels according to their increasing energy follows this order: -
* Write the electronic configuration of the following atoms: -
11Na , 20Ca , 30Zn , 27Co , 35Br , 48Cd
11Na : 1S2 , 2S2 , 2P6 , 3S1
20Ca : 1S2 , 2S2 , 2P6 , 3S1 , 3P6 , 4S2
30Zn : 1S2 , 2S2 , 2P6 , 3S2 , 3P6 , 4S2 , 3d10
2) Hund’s rule: - “No electron pairing takes place in a given sublevel until
each orbital contains one electron” G.R.F.
*Because on pairing electrons the repulsion force produced decreases the stability of of the atom So electrons prefere to be a single before pairing.
7N 1S2 , 2S2 , 2P3 7N 1S2 , 2S2 , 2P3
(Wrong configuration)
* In the oxygen atom (its electron configuration is shown in this diagram) we find that sublevel 2P has (4) electrons.
* Three of them are distributed first in the three orbitals according to Hund’s rule.
G.R.F 1- 4S sublevel is filled before 3d
*Because 4S sublevel less energy than 3d and according to building up priciple ………………...
2- Electrons prefer pairing in the same sublevel than
transfer to higher energy level.
*Because energy of pairing less than energy of transfer to higher energy level
write electronic configuration according to
1- principal energy levels 2- Building up principal
3- Hund's rule 4- The nearest inert gas.
6C , 13Al , 16S , 7N , 26 Fe , 17Cl , 1H , 35Br
The relation between principal Q.N , sublevels and orbitals
1- Number of sublevels equals the principal quantum number [n]
2- Number of orbitals in the same level equals square the pricipal quantum number.[n2]
Ex . Second level consists of four orbitals [ 2S , 2Px , 2Py and 2Pz ]
3- Number of electrons which fill each main energy level equals twice squar the
principal quantum number [ 2n2 ]
Ex Second level filled with 8 electrons distributed as follow
[ 2S2 , 2Px2 , 2Py2 and 2Pz2 ]
Complete the following table
Main level Principal Q.N Sublevels Orbitals Electrons
First K
Second L
Third M
Fourth N
Write electronic configuration then determine sublevels and orbitals
12Mg , 21Sc , 16S , 8O , 28 Ni , 19K , 1H , 25Mn
…………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………
1-Sublevel (P) is saturated with (6)electrons.
2-Electrons never fall inside the nucleus.
3-Electrons occupy the sublevel (4S) before (3d).
4-The line spectra of any element is a property distinguish it.
5-Electrons prefer to be paired in the same Sublevel than transfer to higher energy level.
6-The electronic configuration of 8O is 1S2,2S2,2P4 but not 1S2,2S2,2P3,3S1.
7-Electrons of the same sublevel prefer to be single before pairing.
8-Electron has a Dual nature .
9-It’s impossible to determine both location and speed of electron at the same time.
10-Electron cloud is prefered to describe any orbital.
11-Atom has a vast space.
12-The electronic configuration of 2He is 1S2 but not 1S1,2S1.
13-Importance of solving Schrodinger’s equation.
14-Atom is electrically neutral.
15-Mass of the atom is concentrated inside the nucleus.
16-Quantum of energy required to transfer an electron from one energy level to another is not equal.
17-The rule 2n2 is not applied on the levels more than fourth.
18-Electrons doesn’t move from energy level to another if the energy absorbed
or emitted is less than quantum.
19-Sublevel (F)is saturated with (14)electrons.
20-Sublevel (d)is saturated with (10)electrons.
21-Sublevel (S)is saturated with (2)electrons.
1-Quantum:
2-Electron cloud:
3-Quantum numbers:
4-Principal Q.N:
5-Subsidary Q.N:
6-Magnetic Q.N:
7-Spin Q.N:
8-Exited atom:
9-Uncertainty principle
10-Auf-bau principle:
11-Hund’s rule:
12-De-Broglie princple:
13-Matter waves:
14-Node
1-Number of orbitals in sublevel (3d) is…………….. May95;96 Aug97
2-Energy sublevel that consistes of three orbitals is ……… Aug 95
3-The number of electrons required to fill the fourth energy level(N)….. May96
4- number of orbitals in the main energy level (n)is equal to …… May92
5-………..quantum number determines the number of the electrons in any energy level. May 91.
1-Pairing of electrons is not happened in an energy sublevel unless all the single electrons occupied the empty orbitals first.( ) May95;98
2-Numbers that define the shape and directions of the orbitals in the space. ( ) Aug 95
3-Quantum number that defines the orbitals of a given energy sublevel and their orientation in space ( ) May 98
1-Aufbau principle May 96
2-Hund’s rule May 97
3-Quantum May 94
1-(S) sublevel is saturated with (2)electrons while (P) sublevel is saturated with (6) electrons. Aug 98
2- Electrons prefere to occupy alone one orbital before pairing can takes place in a given sublevel. Aug 98
Notes :.
1- Difference in energy between main energy levels is large.
2- Difference in energy between sublevels in the same orbital is small.
3-There is no diff in energy between orbitals of the same energy levels.
1- principal quantum number and subsidary quantum number
2- Spin quantum number and magnetic quantum number
3- Matter waves and electromagnetic waves.
1- The electron has a dual nature in the sense that it is a ……… which also has ………..
2- The exciteed atom which gained ……… and then one ………. or more is transferred from its energy level to …………
3- The maximum number of energy levels in the known atoms in their ground stase is ……….
4- The sublevel (S) saturated with ………. electrons while the sublevel (d) is saturated with ……….. electrons
5- Sublevel (f) consists of ……… orbitals
6- The fourth energy level is saturated with ……… electrons while the second energy level is saturated with ………. electrons
7- Matter waves differ from electromagnetic waves in ……… and ……….
8- ……….. defines the subleveles and its number within a principal
energy level
9- Electrons rotate around the nucleus in a ………. movment. two forces affect the electron during its rotation which are ……… force and …………force
10-An element with atomic number 19, its electrons can be distributed in a
number of sublevels equals ……… and a number of orbitals
equals …………
11-In 3d7 , 3means ……… , d means ………. and 7 means …………
12-Any sublevel contains an odd number of …………
1-The maximum number of electrons occupying an energy level of a
principal quantum number (n) is ………..
a)2n b) n2 c)2n2 d) (2n)2
2-The number of orbitals in the 3d sublevel is ……….
a) Three b) Five c) six d) seven
3-The energy of the main energy level (N) is ………. The energy of the
main energy level (m).
a) greter than b) less than c) equal
a) 4-The sublevel (S) ………… a) has onlt one orbital c ) a and b
b) its shape is spherical and symmetrical.
5-The atom is ………. In the ground state .
a) Electrically neutral b) Negativity charged c) Positivly charged
6-The number of energy sublevels for an element of atomic number 18
equals…………. a) 3 b) 10 c) 6 d) 5
7-The number of sublevels equals ……….
a) n b) 2n c) n2 d) 2n2
8-The scintict who introduced the presence of sublevels is ………..
a) Hund b) Bohr c) Heisenberg d) schrodinger
9-If the atom was in its ground state, the electrons…………
a) lose energy b) Gain energy
c) Emit light continously d) don’t lose or gain energy.
10-Anyu orbital can be saturated with ………… electrons.
a) 2 b) 5 c) 18 d) 32
11-The number of orbitals in the (p) sublevel is …………..
a) Three b) four c) five d) severn
12-The sublevel energy that consistsa of three orbitals is ……….
( s , p , d , f )
13- The sublevel energy that consists of seven orbitals …………
( s , p , d , f )
14-The sublevel energy that consists of one orbital is ……….
( s , p , d , f )
15-The energy of orbitals are equal in one the following cases:-
a) Orbitals of 1s = 2s = 3s = 4s b) Orbitals of the same sublevel.
b) Orbitals of the same principal energy level.
c) Orbitals which contain the same number of electrons.
16-The third energy l;evel is saturated with ………….. electrons
( 8 – 10 – 18 – 32 )
17-The fourth energy level is saturated with ……….. electrons
( 8 – 18 – 32 )
18-The number of sublevel equals ………….
a) The principal quntum number .b) Squar the principal quantum number.
b) Three times the principal quanyum number.
19-the second energy level is saturated with ……….
(8 – 10 – 18 – 32 )
20-The number which indicates the number of sublvels in each principal
energy level is known as………..
a) The principal quantum numner. B) The subsiadary quantum number.
b) The spin quantum number.
21- The scientist who used the term orbital to indentify regions of space
around the nucleus in which the electron most probable excists is …….
( Hund – Heisenberg – Boher – Schrodinger )
Write the scietific term:
1-The region of space around the nucleus where there is a great probability
of finding the elctron. […………………..]
2-Every morning body is associated with a wave motion which has some
properties of light waves. [……………………..]
3-The number which describes the spin motion of an electron around its
own axis. […………………….]
4-The amount of energy absorbed or emitted when an electron is transferred
from one energy level to another. [………………….]
5-The atom which gained an amount of energy enough to transfer
electrons from their original level to higher ones. […………….]
6-Eletrons occupy orbitals in the order of increasing orbital energy, the
lwoest energy orbitals are filled first. [………………………..]
7-Number that defines the shapes and the direction of the orbitals in the
space. […………………………….]
8-Number that defines the principal energy levels in the atoms.[…………]
9-No electron pairing takes place in a given sublevel until each orbital contains one electron. […………………]
10-It is impossible to determine the location and the velocity of the
electron at the same time. [……………………]
11-The sublevel which consists of 5 orbitals.[……………….]
12-The wave that asocites any moving particle such as electron.
Secondly: Write the electronic configuration of each of the following
elements according to the Aufbau priciple:
10Ne , 15P , 19K , 26Fe , 30Zn
Choose form column (B) what suits to column (A)
(A) (B)
1-No electron pairing takes place in a given sublevel until each orbital contains one electron.
2-The atom is not uniformly dense and there is a
fast space in it.
3-It is impossible to determine the location of electron and its velocity practically at the same time.
4-The difference in energy between the principal energy levels is not equal, therefore each energy level has its own energy.
5-Electrons occuy the orbitals in the order of increasing orbital’s energy where the lowest levels are filled first.
6-Every moving body is associated with a wave motion. 1-Rutherford.
2-Bohr.
3-Heisenbergy.
4-Hund.
5-DeBroglie.
6-Aufbau.
الصف الثالث الثانوى كيمياء لغات الباب الأول - كيمياء لغات
كيمياء لغات - الباب الأول| Greek philosophers
02 Thomsom's Model Of Atom Observations
الدرس في شكل نص مقروء
كيمياء لغات - الباب الأول| Greek philosophers
Greek philosophers
*They imagined the possibility of dividing any piece of matter to smaller parts
which is divided into smaller one until reach to A tom
[A] means No but [tom ] means divided.
* Aristotle
* He rejected the concept of [the atom] and believed that:-
*All matters consists of four components [water , air , dust and fire]
* It was belived that cheap metal as Iron or Copper can be changed into precious
one as a Gold by changing the ratio of the constituents.
2-Boyle *He refused Aristole's idea about the nature of substances.
*He put the first definition about the element.
Element: The simplest pure substance can't be divided into simple one
by a traditional chemical method . [ exam 2006]
3- Dalton's atomic theory
* Dalton carried out many researches and experiments.
* Dalton stated the first theory about the atomic structure.
*He postulated that
1-Matter is composed of very minute particles named atoms.
2- An atom is a one solid part and undividable.
3- Atoms of the different elements are different.
Discovery of cathode rays
From some exp. about electric discharge in 1897 scientists reached to :-
1-All gases under a normal conditions from pressure and temp. are insulators
So cathode ray tube must be evacuated G.R.F
2-If a glass tube is discharged from a gas and the pressure inside it ranges from
[ 0.01 - 0.001 ] mmHg then exposed to suitable potential difference
it becomes a good conductor
3-When a potential difference between the two poles increased to 10000 volt
The scientists observed that:-
*Stream of invisible rays were emitted from the cathode and the glass tube glow
*These rays are called cathode rays
Question Define :- [ Cathode rays]
………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………
properties of cathode rays
1-They consist of a very small particles.
2-They have a thermal effect.
3-They transfer in a straight lines .
4- They are effected by electric and magnetic field
5- They are negatively charged particles.
6-Their nature do not depend on
-Substance of cathode. or type of a gas inside the discharge tube
* Means they do not differ either in behavior or in nature of the material of cathode or the gas which is used
* This is strong evidence that.
So [ They are fundamental constituent of all matters.] G.R.F
4-Thomson's model of the atom [Exam 2007]
* He suggested a new atomic model of the atom that took into consideration the
existence of electron
*Atom is considered a symmetrical sphere from positive electricity contains
negative electrons inside it
So atom is electrically neutral.
Rutherford's exp.
*By Giger and Marceden
Experiment 1- A deep lead box with a small hole .
2- A piece of radioactive element which produces an alpha particle was placed inside the box.
3- A metal sheet in the form of incomplete circle was plated by
a layer of zinc sulphide then placed in the pass of alpha particles .
4- A very thin gold foil was placed between the beam of alpha
particles and the metal sheet.
Observations and conclusions
Observations Conclusions
Atom has a vast space and it is not uniformly dense. A) Most of the alpha particles appears on the same place.
The atom must contain a tiny part of a very high density [ was named the nucleus ] B) Very small % of alpha particles
was reflected back [ some flashes appeared in front of the foil ]
The dense part of the atom where most of the mass is present , appears to have a similar positive charge to that of alpha particles. C) Some of (alpha) particles penetrated the foil but deflected
* What is the role of rutherford ?
Rutherford’s model
* Postulates of the theory
The atom: - 1) It’s similar to Solar system .
2) It’s extremely small in size.
3) It has a very complicated structure.
The nucleus: - 1) At the center.
2) Very small in size compared to an atom.
(There is a vast space between it and electron’s orbits)
3) An atom is not uniformly dense .
4) The mass of an atom is concentrated in the nucleus G.R.F. Because mass of protons and neutrons inside the nucleus greater than mass of electrons which can be neglected
The electrons: - 1) Thery orbit around the nucleus in a great speeds at
different distances from it.
2) They have a negligible mass compared to the mass of
nucleus .
3) They have –Ve charges and the sum of –Ve charges equal to the sum of +Ve charges so that the atom is electrically neutral. G.R.F.
4) They keep their orbits around the nucleus due to
the centrifugal force = attraction force
OBJECTIONS ON RUTHERFORD’S MODEL
1- Rutherford's concept was contradicated by Maxwell's theory
[An electromagnetic theory based on classic Newton’s laws for slightly big particles. This is called Maxwell’s theory. ]
Maxwell’s theory: -
“when an electrically charged particle rotates around the opposite charged particle, it will lose some of its energy gradually by emission of radiations, leading to a gradual decrease of its speed and its orbit’s radius until hit in the centre.
* By applying this theory on the electron movement.
* We would expect that electrons are in a state of continous emission of radiation with a gradual decrease in the orbit radius, thus leading the electrons to spiral inwords until they hit in the nucleus.
* That was the obvious contradication classical mechanical laws and Rutherford's concept.
2- Failed to explain the spectral line of elements although it depends on atomic structure.
Atomic spectra and its explanation
* The study of atomic spectra is considered the key which solved the puzzle of
the atomic structure .
* On heating a mass of very closed atoms [ solid – liquid – compressed gases ]
to high temperature it radiates light .
* On examining this light by spectroscope we observe rainbow of mixed colours
without any separation in between.
* On heating gases or vapours under reduced pressure to high temp. or by electric
spark , they radiate another type of radiations .
* By examining this light by spectroscope it was found to be composed of a limited
numbers of coloured lines named line spectra.
* Spectral lines are essential characteristic for each element as finger print
* Study of line spectra of the sunlight indicated Hydrogen and Helium, are the main
components of the sun .
Spectroscope :- It is an optical instrument for light analysis.
Secondly: Bohr’s atomic model
Postulates from Rutherford Added other postulates Advantages of Bohr’s model Defects of Bohr’s model Notes
* He adapted some of Rutherford’s postulates: -
1) The nucleus is small in size , heavy and positively charged.
2) No. of +Ve charges = No. of –Ve charges so atom is (electrically neutral).
3) The centrifugal force = attraction force.
* And he added other postulates: -
4) Electrons orbit around nucleus in a definite allowed energy levels (they can’t be found in between).
5) Each electron has a definite energy, depending on the distance from the
nucleus. * This energy increases and gaps decrease further from nucleus.
6) It’s possible to determine both speed and location of electrons practically at the same time.
7) Electrons orbit around nucleus without emission or absorption of energy under normal conditions (ground state).
8) Electron remains in the lowest allowed energy level till it absorbs an amount of energy (quantum), by heating or by electric discharge, it becomes excited and jumps to higher enrgy level, depending on the amount of energy which gained. And return to its original one by losing the same amount of energy
which appears as spectral line.
* Excited atom :- An atom which gained amount of energy and electron jumps
to higher enegy level .
Quantum: - It’s the amount of energy absorbed or emitted when an electron jumps from a level to another.
* The quantum required to transfer an electron from one level to another is not equal means the gaps between the levels are not equal.
It decreases further from the nucleus.
* The electron doesn’t transfer to higher energy level unless the absorbed or emitted energy is equal to the difference between the two levels.
Question Compare between quantum and exitation energy
*Energy increases
*Gaps decrease
* Distance from nucleus increases
Adequacy [Advantages ]of Bohr’s model [ Result]: -
1) It explained the Hydrogen atom spectrum .
2) It introduced a new quantaized energy called (quantum) in an atom.
3) It emphasized that [electrons when moving around the nucleus in the ground state doesn’t radiate energy].
[ Bohr succeded in a reconciliation between Rutherford and Maxwell ]
* Inadequacy [disadvantages ]of Bohr’s atomic model: -
*Despite the great effort of Bohr to construct his atomic model the quantitative calculations of this theory didn't agree with all exp. Data.
1) It failed to explain the spectral line of any atom except Hydrogen.
[ till Helium atom which contains only two electrons.]
2) It considered the electron is a particle only, and did not considered that it also has wave properties.
3) It postulated that it’s possible to determine the location and speed of electron at the same time but this pracically impossible. G.R.F
4) It described that Hydrogen atom is a planar, but it has three dimentional coordinates ( x , y , z ).
1) Electrons never fall inside the nucleus
* That’s due to the presence of two forces equal in magnitude but opposite in direction which are: -
a) Centripetal force .[The attraction force between the nucleus and the electrons.]
b) Centrifugal force [ arising while the electron orbit the nucleus]
2) The quantum of energy required to transfer an electron from one
energy level to another is not equal.
* Because the difference in energy between them is not equal since the
energy gaps decrease further from the nucleus.
3) The electron does not move from energy level to another if the energy absorbed or emitted is less than quantum.
* Because quantum is a limited amount of energy can’t be divided or multiplied.
