Download What is Organic Chemistry?

What is Organic Chemistry?
Organic chemistry (a science of less than 200 years old) is
the study of the compounds of carbon, which are the central
substances of which all living things on this planet are made.
- Proteins (blood, muscle, and skin.).
- DNA (contain all the genetic information).
- Enzymes (catalyze the reactions that occur in our bodies).
(clothing, household items, medicines, pesticides, insecticides,....etc)
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Atomic Structure
Shell
proton
+
electron
N
N
+
-
neutron
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Atomic Number and Atomic Mass
The atomic number (Z) is the number of protons in the atom's nucleus.
or the number of electrons around the nucleus.
The mass number (A) is the number of protons plus neutrons
(All the atoms of a given element have the same atomic number)
Isotopes are atoms of the same element that have different numbers of
neutrons and therefore different mass numbers
The atomic mass (atomic weight) of an element is the weighted average
mass in atomic mass units (amu) of an element’s naturally occurring
isotopes atomic weight
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Properties of subatomic particles
Symbol
Relative
Charge
electron
e-
-1
negative
0
9.11x10-28
proton
p+
+1
positive
1
1.67x10-24
neutron
n0
0
neutral
1
1.67x10-24
Particle
Relative
Mass
Actual
Mass (g)
Is very small
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Atomic Structure
Quantum mechanics: describes
electron energies and locations by a
wave equation in atoms in terms of:
 Main or principal energy levels (shells) (n)
 Energy sublevels (subshells) (orbitals)
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Principal Energy Levels (shells) (n)
Contain electrons that are close in energy.
 Have values of n = 1, 2, 3, 4, 5, 6…..
 The 1st shell (n = 1) is the lowest in energy, 2nd
level is higher, and so on 1<2<3<4
 Maximum number of electrons = 2n2
n =1
2(1)2 = 2
n =2
2(2)2 =8
n=3
?????
Each shell contains subshells known as atomic
orbitals
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 What is orbital? is a region of space where the
probability of finding an electron is large.
 Why it is important to understand location of
electrons?
(Arrangment of e- creates bonds between atoms,
these bonds will be involved in chemical reactions)
 Electron cloud has no specific boundary so we
show the most probable area
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Shapes of Atomic Orbitals
 Four different kinds of orbitals for electrons,
denoted s, p, d, and f
 s orbitals: spherical
 p orbitals: pairs of dumbbells aligned along x, y
and z axis at 90° to each other
 d orbitals: complex shapes
 f orbitals: very complex
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p-Orbitals
 In each shell there are three perpendicular p
orbitals, px, py, and pz, of equal energy
 Lobes of a p orbital are separated by region
of zero electron density, a node
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Orbitals and Shells
 Orbitals are grouped in shells of increasing size and
energy
 (Sublevel energy: s<p<d<f)
 Different shells contain different numbers and kinds of
orbitals
 Each orbital can be occupied by 2 electrons
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s
Number of
orbitals
1
Electrons
allowed
2
p
3
6
d
5
10
f
7
14
Sublevel
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Electron Configurations
Ground state electron
configuration: Z electrons
(Z = atomic number of the
atom) are placed serial
into the orbitals according
to the following rules:
Aufbau principle“build-up”:
electrons go into lowest
energy orbitals first.
1s→2s→2p→3s→3p→4s→3d
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Pauli exclusion principle:
each orbital can hold a
max. of two electrons
having opposite spins.
(Electron spin can have only
two orientations, up ↑ and
down ↓)
Hund’s rule:
When filling sublevels other than s, electrons are
placed in individual orbitals of equal energy before
they are paired up.
O8
1S2 2S22P4
↑↓
↑↓
↑↓
↑
↑
1s
2s
2px
2py
2pz
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Electron configuration and orbital diagram
Electron
Configuration
Li3
1S2 2S1
C6 1S2 2S22P2
N7 1S2 2S22P3
O8 1S2 2S22P4
Orbital Diagram
↑↓
↑
1s
2s
2px
2py
2pz
↑↓
↑↓
↑
↑
1s
2s
2px
2py
2pz
↑↓
↑↓
↑
↑
↑
1s
2s
2px
2py
2pz
↑↓
↑↓
↑↓
↑
↑
1s
2s
2px
2py
2pz
Exercise:
Give the electronic configuration for the following:
Na+
Cl-
Mg+2
(knowing that Na11, Mg12, Cl17)
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Write the electronic configuration for the following elements:
a)
Ca
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b)
Na
40
1s2,2s2,2p6,3s2,3p6, 4s2
d)
Cl
17
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1s2,2s2,2p6,3s2,3p5
11
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c)
e)
Si
28
1s2,2s2,2p6,3s2,3p2
16
1s2,2s2,2p4
1s2,2s2,2p6,3s1
14
O
8
f)
B
5
11
1s2,2s2,2p1
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Variation of atomic properties with
electronic structure
1.
2.
3.
Atomic and ionic size
Ionization Energy (IE)
Electron Affinity (EA)
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1. Atomic and ionic size
A. Atomic size, Atomic Radii
Higher effective nuclear charge
Electron held more tightly
Atomic size decrease
Electrons held less tightly
Orbital size increase
Atomic size increase
Representation of Periodic Table
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Size increase form top to bottom within a group?
because of larger orbitals and electrons are added
farther from the nucleus, there is less attraction.
Size decrease form left to right within a period?
Because of increasing positive nuclear charge increase the attraction of
the valence electron inward
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B. Ionic size
 When electrons are added
to an atom (creating a
negative ion) the repulsion
between them increases,
and causes them to push
away from each other,
therefore making the size
of the ion larger than the
neutral atom.
 When electrons are removed
from an atom (creating a
positive ion) the repulsion
decrease, and the remaining
electrons are drawn closer
to the nucleus, therefore
making the size of the ion
smaller than the neutral
atom.
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2. Ionization Energy
Ionization energy (IE)
is the energy required to remove an electron
from an atom (in the gas phase) or ion in its
ground state.
Mg (g) + 738 kJ ---> Mg+ (g) + e1st ionization energy
X → X+ + e2nd ionization energy
X+ → X2+ + e3rd ionization energy
X2+ → X3+ + e-
(Generally the (n+1)th ionization energy is larger than the n th
ionization energy. Always, the next ionization energy involves
removing an electron from an orbital closer to the nucleus. Electrons
in the closer orbitals experience greater forces of electrostatic
attraction; thus, their removal requires increasingly more energy.)
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3. Electron Affinity
Electron affinity (EA)
is the potential energy change associated with the
addition of an electron to an atom (in gas phase) or
ion in its ground state.
A(g) + e- ---> A-(g)

 
O atom

+ electron
EA = - 141 kJ
E.A. = ∆E

 

O ˗ ion
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Summary
Size, IE, and EA in periodic table
Larger
Atomic Size
Larger
Larger
Ionization
energy
Larger
Larger
Electron
affinity
Larger
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Chemical Bonding
The octet rule
Why do atoms bond together?
(has less energy)
more stable
 Atoms tend to attain electron configuration of
noble gases (why?) by:
1) Losing electrons
2) Gaining electrons
3) Sharing electrons
Electron configuration of noble gases is very
stable, because e- have filled up their orbitals,
having very high IE, thus it’s extremely difficult to
break an e- away.
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1- Ionic Bonds
 formed by the transfer of one or more electrons
from one atom to another to create ions.
 It is an attractive force between oppositely
charged ions and the products are called salts.
 Atoms involved widely differ in electronegativity.
(Electronegativity measures the ability of an
atom to attract electrons)
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2- Covalent Bonds
 When 2 or more atoms of the same
electronegativities react, a complete transfer of
electrons does not occur. In these instances the
atom achieve noble gas configuration by sharing
electrons . Covalent bonds form between the
atoms, and the products are called molecules.
 Covalent bond forms when two atoms approach
each other closely so that a singly occupied
orbital on one atom overlaps a singly occupied
orbital on the other atom
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Q. In a tabular form,
compare between ionic
bonds and covalent
bonds.
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