لخّصلي

Introduction

o A chemical bond is the force that holds a group of atoms together.15/29

Evaluation of bond polarity

I The electronegativity scale, developed by L.Pauling in the 1930s, was based on measurements of the strengths of covalent bonds between di

erent elements

I Fluorine was arbitrarily set an electronegativity of 4.0 (now refined to 3.98), thereby creating a scale in which all elements have values between 0 and 4.0 o Bond polarity is evaluated by determining the electronegativity di erence between two adjacent atoms that compose the bond

o Regarding the electronegativity di

erence (

EN) between them, a

bond can be classified as: I ionic if

EN O 2.0 I polar (covalent) if 0.4<

EN < 2.0

I covalent if

EN AE 0.4

16/29

Example 5

Given the relative electronegativity values (EN) below: Atoms H C K F Cl EN 2.1 2.5 0.8 4.0 3.0

classify the following bonds as ionic, polar or covalent : (a) HCl , (b) KF, and (c) C!=C and H!=C bonds in C2H6

17/29

Dipole moment

I The asymmetrical charge distribution in a polar substance such as HCl produces a dipole moment (u).For example: I C(C)C bond has a length of 120.3 pm in C2H2 molecule I C=C bond has a length of 133.9 pm in C2H4 molecule I C-C bond has a length of 153.5 pm in C2H6 molecule

19/29

Lewis structure

o A Lewis structure represents the covalent bonding in a molecule in which : I only valence electrons are shown I shared electrons between two atoms are shown either as lines or as pairs of dots I lone pairs are shown as pairs of dots

20/29

Drawing Lewis structure

o Lewis structures of compounds such as HF or Cl2 are simple to draw; however, many complexes molecules and ions (CH2O, CO2!= 3

... ) are not o Therefore, basic steps should be followed to draw these complexes compounds: 1.o The ions are oppositely charged; therefore they attract each other to form ionic networks, lattices o Generally, ionic compound are formed from the reaction of metals with non-metals: I electrons are transferred from the metal to the non metal I the metal becomes + ion and the non-metal becomes != ion o For example, Li (Z=3) reacts with F (Z=9) as follow:

o the VE 2s1 of Li is transferred to the F atom to form the LiF ionic compound.o For example the elements lithium (Z = 3), beryllium (Z=4) and boron (Z=5) are written using the Lewis dot symbols o The Lewis dot symbols explains why : I lithium, with one unpaired valence electron, tends to form one bond I beryllium, with two unpaired valence electrons, tends to form two bonds I boron, with three unpaired valence electrons, tends to form three bonds

5/29

Example 1

1- Draw the Lewis dot symbols for these elements C, N, O, F, Ne. 2- Predict the number of bonds that can be formed in a compound from these elements ? I and metallic bonds: Cu, Fe ... (solid pure metals)

o In the 20th century, a system has been devised to predict the number of bonds formed by most elements in their compounds, the Lewis dot symbols

4/29

The Lewis dot symbols

o The Lewis symbols consists of the chemical symbol for the element with a dot for each valence electron.o The covalent bond is: I non-polar in H2 molecule because the electrons are equally shared I polar in HF molecule because the electrons are not equally shared o Bond polarity is evaluated by determining the electronegativity di erence between two adjacent atoms that compose the bond.o With regard to the electronegativity di erence (

EN) between two atoms in a compound, the formed bond can be: I ionic (Na-Cl), I polar (H-Cl) I or non-polar (Cl-Cl, H-H, C-C,...)

14/29

Electronegativity variation

o In a periodic table, electronegativity increases from lower left to upper right.o For example, the formation reaction of the molecule Cl2 is as follow :

I each Cl atom has 7 VEs; but only the unpaired electron is involved in the covalent bond I the non-bonding electrons are called lone pairs, which are not involved in covalent bond I Each Cl in Cl2 molecule has three lone pairs of electrons

11/29

Example 4

1.I atoms such as F, Cl and Br (p block) have higher electronegativities I atoms such as Li, Na and K (s block) have lower electronegativities I electronegativity of the d block elements is intermediate between the alkali and the halogens elements.o A covalent bond between many-electron atoms involves only the VEs.Explain 2.

Introduction

• A chemical bond is the force that holds a
group of atoms together.
• The chemical bonds are classified into
three principal types: .
I ionic bonds: NaCl, KCl ...
I covalent bonds: CH4 , C2H6, C5H12,
C6H6 ....
I and metallic bonds: Cu, Fe ... (solid
pure metals)

• In the 20th century, a system has been
devised to predict the number of bonds
formed by most elements in their
compounds, the Lewis dot symbols

4/29

The Lewis dot symbols

• The Lewis symbols consists of the
chemical symbol for the element with a
dot for each valence electron.
• For example the elements lithium (Z = 3),
beryllium (Z=4) and boron (Z=5) are
written using the Lewis dot symbols
• The Lewis dot symbols explains why :
I lithium, with one unpaired valence
electron, tends to form one bond
I beryllium, with two unpaired valence
electrons, tends to form two bonds
I boron, with three unpaired valence
electrons, tends to form three bonds

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Example 1

1- Draw the Lewis dot symbols for these elements C, N, O, F, Ne.
2- Predict the number of bonds that can be formed in a compound from
these elements ?

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Example 1 - answer

1- Draw the Lewis dot symbols for these elements C, N, O, F, Ne.
dots representing VE are placed first, one at a time. Up to 4 dots
are placed above, below, to the left, and to the right of the
symbol, in any order. The next dots, for elements with > 4 VEs,
are again distributed one at a time, each paired with one of the
first four.
2- Predict the number of bonds that can be formed in a compound from
these elements ?
The unpaired dots are used to predict the number of bonds that
an element will form in a compound. N with 3 unpaired VEs,
tends to form three bonds. C, with four unpaired VEs, tends to
form four bonds, and so on. However, Ne with no unpaired VEs
will not form bonds, as observed with all noble gases

7/29

The Octet Rule

• The octet rule states that an atom other than H tends to form
bonds until it is surrounded by eight valence electrons.
• An atom that have less than eight valence electrons tends to react
to form a more stable compound.
I because all orbitals will be full and have noble electronic structure :

s
2
p
6

• d or f electrons are not considered in this rule.
I hence, only the s and p electrons are involved in the octet rule

8/29

The Ionic Bond

• An ionic bond is the electrostatic force that holds ions together in
an ionic compound.
• The ions are oppositely charged; therefore they attract each other to
form ionic networks, lattices
• Generally, ionic compound are formed from the reaction of metals
with non-metals:
I electrons are transferred from the metal to the non metal
I the metal becomes + ion and the non-metal becomes ≠ ion
• For example, Li (Z=3) reacts with F (Z=9) as follow:

• the VE 2s1 of Li is transferred to the F atom to form the LiF ionic
compound.
• In the ionic compound, the electron configuration of Li and F atoms
become respectively 1s2 (of He atom) and 1s22s22p6 (of Ne atom).

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Example 2

• Using the Lewis symbols, explain the formation of :

1. Sodium chloride NaCl


2. Aluminum chloride AlCl3


3. Aluminum oxide Al2O3
   Na (Z=11); Al (Z=13) ; O (Z=8)

10/29

The Covalent Bond

• A covalent bond is a chemical bond in which two electrons are
shared by two atoms.

I Put simply, the covalent bond is better written H≠H.
• In covalent bonding, each electron in a shared pair is attracted to
the nuclei of both atoms.
• A covalent bond between many-electron atoms involves only the
VEs.
• For example, the formation reaction of the molecule Cl2 is as follow :

I each Cl atom has 7 VEs; but only the unpaired electron is involved in
the covalent bond
I the non-bonding electrons are called lone pairs, which are not
involved in covalent bond
I Each Cl in Cl2 molecule has three lone pairs of electrons

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Example 4

1. Does the the formation of Cl2 verify the Octet rule ? Explain


2. Can we say that the unpaired electron involved in chemical bonding
   is transferred to the other chlorine ? Explain

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Bond polarity

• Bond polarity tells us if the electron pairs in a covalent bond are
equally shared or not.
• For example, the sharing of electron pairs in H2 and HF molecules is
not the same.

• The covalent bond is:
I non-polar in H2 molecule because the electrons are equally shared
I polar in HF molecule because the electrons are not equally shared
• Bond polarity is evaluated by determining the electronegativity
di
erence between two adjacent atoms that compose the bond.

13/29

Electronegativity • Electronegativity (EN) represents the
ability of an atom to attract the electrons
in a chemical bond. I EN of an element depends on its

I and

A

• For example : I Chlorine (Cl), with a high electron a

nity and a high ionization energy, has
a high electronegativity.
I Sodium (Na), with a low electron a
nity
and a low ionization energy, has a low
electronegativity.

• With regard to the electronegativity
di
erence (

EN) between two atoms in a
compound, the formed bond can be: I ionic (Na-Cl), I polar (H-Cl) I or non-polar (Cl-Cl, H-H, C-C,...)

14/29

Electronegativity variation

• In a periodic table, electronegativity increases from lower left to
upper right.

I atoms such as F, Cl and Br (p block) have higher electronegativities
I atoms such as Li, Na and K (s block) have lower electronegativities
I electronegativity of the d block elements is intermediate between the
alkali and the halogens elements.

15/29

Evaluation of bond polarity

I The electronegativity scale, developed by L.Pauling in the 1930s,
was based on measurements of the strengths of covalent bonds
between di

erent elements

I Fluorine was arbitrarily set an electronegativity of 4.0 (now refined
to 3.98), thereby creating a scale in which all elements have values
between 0 and 4.0
• Bond polarity is evaluated by determining the electronegativity
di
erence between two adjacent atoms that compose the bond

• Regarding the electronegativity di

erence (

EN) between them, a

bond can be classified as:
I ionic if

EN Ø 2.0
I polar (covalent) if 0.4<

EN < 2.0

I covalent if

EN Æ 0.4

16/29

Example 5

Given the relative electronegativity values (EN) below:
Atoms H C K F Cl
EN 2.1 2.5 0.8 4.0 3.0

classify the following bonds as ionic, polar or covalent : (a) HCl , (b) KF,
and (c) C≠C and H≠C bonds in C2H6

17/29

Dipole moment

I The asymmetrical charge distribution in a polar substance such as
HCl produces a dipole moment (μ).
I Dipole moments are vectors, they possess both a magnitude and a
direction. In HCl, for example, the dipole moment is indicated as an
arrow that shows the direction of electron flow :
I The dipole moment is defined as the product of the partial charge Q
on the bonded atoms and the distance r between the partial charges:

μ = Q ◊ r

I where Q is measured in coulombs (C) and r in meters
I The unit for dipole moments is the debye (D):
1D = 3.336 ◊ 10≠30 C.m

18/29

Multiple bonds

• A multiple bond is a covalent bond which is shared between two
atoms by more than one electron pair
I single bond: two atoms are held together by one electron pair
I double bond: two atoms are held together by two electron pairs

I triple bond: two atoms are held together by three electron pairs

• For the same pair of atoms, triple bonds are shorter than double
bonds, and double bonds are shorter than single bonds. For example:
I C©C bond has a length of 120.3 pm in C2H2 molecule
I C=C bond has a length of 133.9 pm in C2H4 molecule
I C-C bond has a length of 153.5 pm in C2H6 molecule

19/29

Lewis structure

• A Lewis structure represents the covalent bonding in a molecule in
which :
I only valence electrons are shown
I shared electrons between two atoms are shown either as lines or as
pairs of dots
I lone pairs are shown as pairs of dots

20/29

Drawing Lewis structure

• Lewis structures of compounds such as HF or Cl2 are simple to
draw; however, many complexes molecules and ions (CH2O, CO2≠
3

... ) are not
• Therefore, basic steps should be followed to draw these complexes
compounds:

1. write the skeletal structure of the compound by placing the least
   electronegative atom in the central position


2. count the total number of VEs in the molecule or ion
   I if anion, electrons must be added according to its charge
   I for a cation, electrons must be subtracted according to its charge


3. draw electron pairs between the central atom and each of the
   surrounding atoms


4. add enough electrons to each surrounding atom to give give them an
   octet; (exception with H)


5. the remaining VEs are placed on the central atom


6. add multiple bonds to the central atom, using lone pairs of the
   terminal atoms, to complete the octet