Covalent Bonding

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Covalent Bonding

• Sch3u1
• General Panet HS

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• Sections 8.1, 8.2, 9.3, and 8.4
• Sections 6.1 16.1, 6.5 and 16.3

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Remember

• Ionic bonds form between
• An ionic bond happens when one atom and the
  other atom

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Remember

• The definition of ionic bond is
• The chemical formula of an ionic compound
  represents a

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Covalent Bonds

• The four sentences above represent four of the
  most essential differences between covalent and
  ionic compounds.
• These differences are so important that

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Essay Question Test 3-1

• Define ionic bond and covalent bond. Outline and
  define, in detail, four major differences between
  ionic compounds and molecular (covalent)
  compounds.

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Molecular Compounds

• Formed by covalent bonds
• Ionic compounds are generally crystalline solids
  at room temperature.
• Molecular compounds (CO2 and water, for example,
  have VERY different properties.)

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Molecular Compounds

• Molecular compounds are formed through covalent
  bonds.
• Covalent bonds are created when atoms SHARE
  electrons, instead of gaining and losing them.

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Vocabulary

• Molecule group of atoms joined by covalent bonds
• Diatomic molecules molecules consisting of two
  atoms
• Molecular formula shows how many atoms of each
  element a molecule contains

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Think About It

• Chlorine is a diatomic element, meaning that it
  exists in its atomic state as two bonded atoms.
• Draw two chlorine atoms.
• Is the bond between these two atoms ionic or
  covalent? How do you know?

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Properties of Molecular Compounds

• Covalent bonds usually occur between
• Often are gases or liquids at room temperature
• Images will show atoms stuck to one another

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Properties of Molecular Compounds

• In general, melting and boiling points of
  molecular compounds are lower than ionic compounds

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Molecular Formulas

• Molecular formula of a molecular compound shows
  how many atoms of each element are in ONE
  MOLECULE of the compound.
• (Contrast this with the chemical formula of ionic
  compounds, which show only the ratio of elements
  in the compound.)

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Molecular Formulas

• Example
• IONIC Calcium chloride CaCl2
• Means that in the compound there are two chloride
  ions for every one calcium ion
• COVALENT Carbon dioxide CO2
• Means that each carbon dioxide molecule consists
  of one carbon atom bonded to two oxygen atoms

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Molecular Compounds

• Molecular compounds can be significantly larger
  than ionic compounds.
• Benzoic Acid C7H6O2
• 2,4-Dichlorophenoxyacetic acid C8H6Cl2O3

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Molecular Compounds

• Formulas not always in lowest terms
• Example Ethane C2H6
• Formulas do not give molecules structure. (It
  must be inferred.)

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Structure Diagrams

• Molecular Formula
• Structural Formula
• Ball-and-stick model
• Space Filling Model
• Perspective drawing

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Forming Covalent Bonds
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Octet Rule

• In covalent bonds, atoms share electrons so that
  they fill their valence levels
• Usually 8 (but only 2 for hydrogen)

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Single Covalent Bonds

• Atoms held together by sharing one pair of
  electrons are said to form a SINGLE COVALENT BOND
• Each atom donates one electron to the bond

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Single Covalent Bonds
Cl
Cl
Cl
Cl
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Single Covalent Bonds
Cl
Cl
Cl
Cl
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Single Covalent Bonds
Single Bond
Cl
Cl
Cl
Cl
Lone Pairs
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Covalent Bonds

• Electrons that do not take part in the bond are
  called lone pairs or unshared pairs

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Covalent Bonds
There are exceptions!

• Different elements can form different numbers of
  bonds
• Group 7A elements need one more electron, and can
  form one bond
• Group 6A elements need two more electrons and can
  form two bonds
• Group 5A three bonds
• Group 4A four bonds

Hydrogen, too!
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Working With Covalent Bonds

1. Draw the electron dot structures.
2. Determine arrangement.
3. Replace shared pairs of electrons with a line.
   (Leave lone pairs.)

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Draw Structural Formulas

• NH3
• H2S
• PBr3

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Draw Structural Formulas

• H2O
• CH4
• OF2

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Draw Structural Formulas

• SCl2
• N2H4
• CCl4

• CHCl3
• C2H6
• HF

Usually, the atom that can form MORE bonds will
be in the center of the molecule!
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Draw Structural Formulas, Part 2

• OBr2
• P2H4
• CI4

• CHBr3
• C2Cl6
• HCl

Usually, the atom that can form MORE bonds will
be in the center of the molecule!
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Double and Triple Bonds
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Double Covalent Bonds

• Atoms attain noble gas configuration by sharing
  two pairs of electrons (four)
• Bond length is shorter

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Double Covalent Bonds

• Oxygen has 6 valence electrons
• O (Group 6A) can form two bonds

O
O
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Double Covalent Bonds

• Oxygen has 6 valence electrons
• O (Group 6A) can form two bonds

O
O
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Double Covalent Bonds

• OCTET RULE NOT FULFILLED!

O
O
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Double Covalent Bonds

• OCTET RULE NOT FULFILLED!

O
O
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Double Covalent Bonds

• OCTET RULE FULFILLED!

O
O
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Other molecules with double covalent bonds are

• CO2
• Ethene, C2H4
• Carbonyl, COH2

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Double Covalent Bonds

• When counting number of valence electrons, double
  bonds count as 4 shared electrons.
• Hydrogen will not form double covalent bonds why?

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Triple Covalent Bonds

• Atoms attain noble gas configuration by sharing
  three pairs of electrons (six)
• Bond length is even shorter

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Triple Covalent Bonds

• Nitrogen has 5 valence electrons
• N (Group 5A) can form three bonds

N
N
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Triple Covalent Bonds

• Nitrogen has 5 valence electrons
• N (Group 5A) can form three bonds

N
N
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Triple Covalent Bonds

• OCTET RULE NOT FULFILLED!

N
N
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Triple Covalent Bonds

• OCTET RULE NOT FULFILLED!

N
N
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Triple Covalent Bonds

• OCTET RULE NOT FULFILLED!

N
N
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Triple Covalent Bonds

• OCTET RULE NOT FULFILLED!

N
N
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Triple Covalent Bonds

• OCTET RULE FULFILLED!

N
N
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Other molecules with triple covalent bonds are

• Acetylene, C2H2
• Hydrogen Cyanide, HCN
• Propyne, C3H4

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Triple Covalent Bonds

• When counting number of valence electrons, triple
  bonds count as 6 shared electrons.

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Diatomic Elements
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Diatomic Elements

• Diatomic elements exist in their atomic forms as
  binary molecular compounds, since covalent bonds
  form between the atoms
• i.e. a molecule of oxygen gas is O2, not O

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Diatomic Elements

• F2
• Cl2
• Br2
• I2

• H2
• N2
• O2

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Exceptions To The Octet Rule
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Exceptions

• Compounds cannot satisfy the Octet Rule for all
  atoms if the total number of valence electrons is
  odd.
• NO2 total number of valence electrons is 17

O
O
N
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More Exceptions

• Nonmetals in the third period and beyond can form
  more than 4 bonds, since they have empty d
  orbitals where they can promote or store
  extra s or p electrons.
• Ex. Phosphorus can form 5 bonds.

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Naming Binary Molecular Compounds
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Naming Molecular Compounds

• CO and CO2 are very different compounds
• How can we distinguish them in their names?

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Naming Molecular Compounds

1. Confirm that the compound is molecular, not
   ionic.
2. Name the elements in the order listed in the
   formula.
3. Add prefixes to identify the numbers of each atom
   in the compound.

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Prefixes Used
Mono- 1
Di- 2
Tri- 3
Tetra- 4
Penta- 5
Hexa- 6
Hepta- 7
Octa- 8
Nona- 9
Deca- 10
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Naming Molecular Compounds

1. Omit the prefix mono- on the first element in
   the name.
2. Add -ide as a suffix at the end of the second
   elements name.

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Examples

• N2O
• Nitrogen oxygen
• Dinitrogen monoxygen
• DINITROGEN MONOXIDE

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Practice Write the Molecular Formula

• Nitrogen trichloride
• Carbon tetrabromide
• Diphosphorus trisulfide

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Practice Write the Name

• Cl2O8
• PH3
• N2O4
• SF6

• H2O
• S2F10
• PCl5
• N2F6

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Polar Bonds and Molecules
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Electronegativity

• A measure of how well an atom attracts electrons
• Measured in Paulings
• In a molecule, some atoms more forcefully attract
  electrons than others

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Electronegativity

• Decreases from top to bottom
• Increases from left to right

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Polar Bonds

• Polar bond covalent bond in which electrons are
  shared UNEQUALLY
• Difference in electronegativity values controls
  whether bond is nonpolar, polar, or ionic

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Polar Bonds

• Differences
• 0.0-0.4 ? nonpolar covalent
• 0.4-2.0 ? polar covalent
• 2.0 ? ionic

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Polar Bonds

• Greek letter Delta (d) represents the partial
  charge acquired by atoms in a polar bond
• H2O
• HF
• CO2

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Intermolecular Attractions

• Polar molecules attracted to one another (called
  dipole interactions)
• Hydrogen bonds are attractions that occur between
  hydrogen and unshared electrons on another
  molecule

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Test Review
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Test Review

• Covalent bonds definitions
• Molecular vs. ionic compounds
• Writing structural formulas
• Writing molecular formulas (from name or from
  structure)
• Writing compound names
• Information on Polar Bonds (pg. 237-240)