Chapter 8: Covalent Bonding Notes
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Chemical Bonds
o Atoms bond to become more stable
o Lower energy states are more stable than higher energy states
o Energy is released when a bond forms
Covalent Bonds
o Sharing of valence electrons
o Valence atomic orbitals overlap end-to-end
 s-s orbital overlap, s-p orbital overlap, or p-p orbital overlap
o Involves attractive and repulsive forces
o Form molecules – two or more atoms bonded covalently
o Primarily occur between nonmetals
o Occur when both atoms can have a full octet
o Each bonded atom equally attracts the pair of shared electrons
o Unbonded pairs of electrons are known as lone pairs
o Examples of covalently bonded molecules: Cl2 (chlorine), O2 (oxygen), N2 (nitrogen), H2O (water), CO2
(carbon dioxide), CH4 (methane)
Single Covalent Bonds
o One pair of electrons is shared
o Single covalent bonds are also called sigma bonds (σ)
o Electron-dot structures: show valence electrons of atoms
 Examples of electron-dot structures:
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Lewis structures: show arrangement of electrons in a molecule
 Steps to drawing a Lewis Structure:
 (1) Calculate total # of valence electrons
o Ex) CO2 = 4 + 2(6) = 16 valence electrons
 (2) Determine central atom
o Usually listed first or is the least electronegative (furthest to the left on the
periodic table)
 (3) Draw skeletal structure
 (4) Connect every bonded pair of atoms by a dash (represents 2 electrons)
 (5) Distribute remaining electrons to atoms surrounding the central atom to satisfy the
Octet Rule
 (6) Distribute remaining electrons to central atom (may require multiple bonds)
 Examples of Lewis structures:
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Double Covalent Bonds
o Two pairs of electrons are shared
o Consist of one sigma and one pi bond
 Pi bond (π) – forms when parallel orbitals overlap and share electrons below and above where
the two atoms are joined together
 Drawing of a sigma and pi bond:
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Example of double bonded molecule: O2
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Triple Covalent Bonds
o Three pairs of electrons are shared
o Consist of one sigma and two pi bonds
o Example of triple bonded molecule: N2
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Strength of Covalent Bonds
o Bond length
 The distance between the two bonded nuclei at the position of maximum attraction
 Trends:
 The shorter the bond length, the stronger the bond
 Triple bonds are shorter and stronger than double bonds and double bonds are shorter
and stronger than single bonds
o Bond-dissociation Energy (Bond Energy)
 Amount of energy required to break a covalent bond
 Positive value
 Trends:
 The smaller the bond length, the stronger the bond, and the greater the bonddissociation energy
 A triple bond is the shortest, is the most difficult to break, and has the highest bond
energy.
 A single bond is the longest, is the easiest to break, and has the lowest bond energy.
 The sum of the bond-dissociation energy values for all of the bonds in a molecule is the amount
of chemical potential energy in a molecule of that compounds
 Total energy change of a chemical reaction is determined from the energy of the bonds broken
and formed
 Endothermic reaction
o Greater amount of energy is required to break the existing bonds in the
reactants than is released when the new bonds form in the products
 Exothermic reaction
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Structural Formulas
o Molecular formula – chemical formula of a molecule
 Ex) H2O – water
o Structural formula – uses letter symbols and bonds to show relative positions of atoms
 Structural formula example:
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Easy to get structural formulas by drawing Lewis structures
Know that there are a few exceptions to the octet rule
 Suboctets
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More energy is released during product bond formation than is required to
break bonds in the reactants
Expanded octets
Lewis Structure Practice Problems:
Resonance Structures
o Possible to have more than one correct Lewis structure
o Show Lewis structure possibilities by drawing resonance structures
o Resonance – occurs when more than one valid Lewis structure can be written for a molecule or ion
o Examples of Resonance Structures: