Bell work
• Compare and contrast ionic bonds to metallic bonds. What
makes them different and what makes them similar?
• Write down your answer in your notebook.
COVALENT BONDING
Chemistry Ch. 8
Today’s Objective
I can describe and model covalently bonded molecules.
I can explain what diatomic elements and the law of
multiple proportions are.
Bond Review
• Ionic bonds: one ion (the anion) has high electronegativity while
the other (the cation) has low electronegativity
• One wants the electrons and the other doesn’t
• Metallic bonds: both ions have low electronegativity
• nobody wants the electrons
• What about if both ions want the electrons?
Covalent Bonds
• When two nonmetals bond, they both pull
somewhat
equally on the electrons. Thus it is easier to share the
electrons than to take or lose them.
• This is due to their relatively high electronegativity.
• Covalent bonds: formed by
the sharing of electrons.
Molecular Compounds
• Molecular Compound: multiple atoms bonded
together by sharing electrons.
• Only covalent bonds
• The electrons are shared to give an atom a full octet
• Eight valence electrons
• Molecular compounds are formed by nonmetals
only!!!!
Molecular Formulas
• Molecular formula: chemical formula of a molecular
compound.
• Shows how many atoms of each element are in a
molecule.
• Ex. H2O and CO2
Representing Compounds
Orbital Representation
Representing Molecular Compounds
• Electron Dot Structure (dots for bonds)
• Structural Formula (dashes for bonds)
Types of Covalent Bonds
Single Covalent Bond: Formed when one pair of
electrons is shared between two atoms.
Ex. F2
Types of Covalent Bonds
Double Covalent Bond: Formed when two pairs of
electrons are shared between two atoms.
Ex. O2
Types of Covalent Bonds
Triple Covalent Bond: Formed when three pairs of
electrons are shared between two atoms.
Ex. N2
Diatomic Elements
• Diatomic element: elements that, in their pure form,
exist in nature covalently bonded in pairs.
• 7 diatomic elements
• H2 N2 O2 F2 Cl2 Br2 I2
Law of Multiple Proportions
• Law of Multiple Proportions: When covalently bonded,
two elements can sometimes come together in multiple
whole number ratios.
• Ex. CO2 and
CO
Bell-work
•
Please draw the Lewis dot structure for the following molecular compounds in your notes.
• NH3
• C2H4Cl2
• HCN
• H3C2HO
Today’s Objective
• I can draw a Lewis Dot Structure for an atom that is an exception
to the octet rule.
• I can name covalently bonded molecules.
Exceptions to the octet rule
• Most elements want 8 valence electrons, but some elements will
settle for other numbers of electrons
• usually occur when an atom bonds to an atom with high
electronegativity (F, Cl, etc)
• Ex:
Naming Molecular Compounds
1.
Name the first atom with the appropriate prefix that
tells how many of that atom are present
1-mono 2-di 3-tri 4-tetra 5-penta 6-hexa
7-hepta 8-octa 9-nona 10-deca
2.
Name the second atom with the appropriate prefix and
an –ide ending
3.
If there is only one of the first atom, mono is not needed
Naming Molecular Compounds
•CO2
• Carbon dioxide
•H2O
• Dihydrogen monoxide
•C3H8
• Tricarbon octahydride
Bond Length
• The more the bonds between two atoms, the
closer the two atoms are to each other.
• Length: Single > Double > Triple
Bond Dissociation Energy
• The energy needed to break a
covalent bond.
• Measured in kJ/mol
• Shorter bonds require more
energy to break.
Challenge… if you are finished with your quiz
• Create a valid Lewis dot structure for the following molecules.
C6H12O6
C 6H 6
Today’s Objective
• I can identify the molecular geometries of molecules
according to VSEPR.
VSEPR Theory (Valence Shell Electron Pair Repulsion)
• All electrons have a negative charge, so they repel each
other.
• Electron pairs in molecules will repel themselves so that
the valence electron pairs are as far apart as possible.
Today’s Objective
I can identify the polarity of a molecule.
Polar Bonds
• In covalent bonding, electrons are shared between two
atoms.
• Depending on the electronegativity of each atom in the
bond, the distribution of electrons can cause a partial
charge at each end of the bond.
Polar Covalent Bonds
• Polar Covalent Bonds: occur when bonded electrons
are not shared equally
• Br ( 3.0 ) takes on a slightly negative charge
• H ( 2.1 ) takes on a slightly positive charge
Nonpolar Covalent Bonds
• Nonpolar Covalent Bonds: occur when bonded
electrons are shared equally
• Happens in bonds between like atoms
Two Ways of Representing Polar Bonds
1.
An arrow pointing towards the more electronegative
atom with a “+” sign at the tail of the arrow
1.
δ+ or δ- signs at appropriate ends of the bond
Determining Bond Type by
Electronegativity
• The difference in electronegativities can tell you what
type of bond is formed
Table of Electronegativity Differences and Bond Type
Difference
Bond Type
Example
0.0-0.4
Nonpolar Covalent
H-H (0)
0.5-2.0
Polar Covalent
H-F (1.9)
> 2.0
Ionic
Na+Cl- (2.1)
Polarity of Molecules
• Dipole: a molecule that has a more positive end and a
more negative end
• If bond polarities cancel each other out, the molecule is
nonpolar.
Polarity of Molecules
Today’s objective
• I can explain the impact various intermolecular forces (IMF) have
on the attraction of molecules.
London Dispersion Forces
Dipole-Dipole Interactions
Hydrogen Bonding
Properties of Compounds
•Strength of Bonds
•Strongest
•
covalent ionic
Weakest
intermolecular forces
• Network Solids: a substance where all atoms
within the solid are covalently bonded
• Ex: Diamonds
Properties of Compounds
•Melting Point
•Highest
• Network solids
Lowest
ionic solids
substances held by
intermolecular forces