Transcript
Chapter 6: Chemical Bonding 1
Learning outcomes •
At the end of the lesson, students should be able to understand and explain: •
Ionic bonding.
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Covalent bonding.
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Octet rule.
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Lewis structure
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Bond polarity and bond strength.
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Octet rule •
Almost all elements in its natural state are not stable.
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Only elements of group 8 are stable.
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The stability of group 8 elements are due to the fact that its valence shell are full. This is known as the octet rule. Elements that do not have fully filled valence shells will try to achieve octet stability by either: •
Donating electron Accepting electron
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Sharing electron
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Ionic bond •
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Ionic bonds are usually formed between metals and non metals. Metals: •
Tend to donate its valence electrons
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Forms positive ion or cation
Non-metals: •
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Tend to accept electrons Forms negative ion or anion
Ionic bond is the strong electrostatic forces of attraction between two oppositely charged ions.
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Covalent bond •
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Covalent bonds are usually formed between non-metals In order to achieve octet configuration, non-metals tend to share its valence electrons. Covalent bond is formed when non-metals’ valence shell overlaps to allow sharing of electrons. Rules of covalent bond: •
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The sharing of electron between atoms are mutual (if atom A shares one electron, atom B will also share one electron) An atom will share exactly the same amount of electrons it needs to achieve octet configuration (chlorine has 7 valence electrons, it needs 1 e to achieve octet stability, hence it will share only 1 e)
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Lewis structure •
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Lewis structure only shows the valence electrons of atoms that are involved in covalent bonding. ‘Dots and crosses’ are still used to represent different electrons.
Lewis structure is simplified as it removes the need to draw the other inner shells.
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Lewis structure
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Electronegativity •
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Electronegativity is the ability of an atom to attract electrons in a covalent bond Different atoms have different value of electronegativity The stronger the electronegativity, the greater the ability to attract electrons General pattern: •
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Non-metals have greater electronegativity compared to metals Electronegativity increases across the period Electronegativity decreases down the group
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Bond polarity •
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When two of the same atoms forms a covalent bond, the electrons are shared equally resulting in a non-polar bond When two different elements are bonded, the electrons are not shared equally The more electronegative atom will pull the electrons closer towards itself The electrons are unsymmetrically distributed The difference in electronegativity of atoms in a covalent bond results in polar bonds 9
Bond Polarity
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Bond Polarity •
The unequal distribution of electrons results in polar bonds
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The bond has a dipole indicated by •
+
(less electronegative atoms)
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– (more
electronegative atoms)
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Bond strength •
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Bond strength refers to the strength of a particular covalent bond. Bond strength can be measured based on bond energy. Bond energy can be defined as the energy needed to break one mole of a particular covalent bond. Bond energies varies from compound to compound. One of the factors that affect bond energy is the length of the bond. •
The shorter the bond, the higher the bond energy.
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Bond strength Bond
Bond length (nm)
Bond energy (kJ mol -1)
H-H
0.074
435
Cl-Cl
0.198
243
O=O
0.121
495
N≡N
0.110
941
H-Cl
0.109
414
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THE END 14
