Covalent Bonding

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Covalent Bonding

• Chemistry
• Chapter 8

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Covalent Bonding

• Not all compounds result from the transferring of
  electrons (Ionic).
• Covalent bonding results when atoms are held
  together (bonded) by the sharing of electrons.

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Molecules Molecular Compounds

• Very rarely do elements exist by their self in
  nature most are found in the form of molecules.
• A molecule is a neutral group of atoms joined
  together by covalent bonds.
• A diatomic molecule - only two atoms
• there are 7 elements that occur naturally as
  diatomic molecules (HI BrONClF).
• Different elements also join to form molecules
• A compound formed of molecules is called a
  molecular compound.

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Molecules Molecular Compounds (cont)

• The molecules of a given molecular compound are
  all the same.
• Molecular compounds tend to have relatively lower
  melting and boiling points than ionic compounds.
• Most are gases or liquids at room temp.
• Covalent (molecular) compounds are all always
  composed of
• non-metals.

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Ionic Vs Molecular Compounds
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Molecular Formulas

• A molecular formula is the chemical formula of a
  molecular compound.
• Shows of atoms of each element a molecule
  contains (NH3)
• A molecular formula does not show either
• the arrangement of atoms
• which atoms are covalently bonded to one another
• A variety of diagrams and other molecular models
  are used for these purposes

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Examples of Different Models Formulas

• Different ways to represent a molecules stucture,
  i.e. ammonia (NH3)

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The Octet Rule in Covalent Bonding

• In forming covalent bonds, electrons sharing
  usually occurs so that atoms attain the electron
  configurations of noble gases.
• Combinations of non-metallic elements in groups
  4-7A are likely to form covalent bonds.

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Single Covalent Bonds

• Two atoms held together by sharing a pair of
  electrons are joined by a single covalent bond.
• An electron dot structure such a HH represents
  the shared pair of electrons of the covalent bond
  by two dots.
• The pair of dots is often represented by a dash
  - H-H. This type of representation is know as
  a structural formula and shows the arrangement of
  covalently bonded atoms (Dot diagram showing
  bonds).
• The molecular formula H2 shows only the number of
  atoms in a molecule.
• Valence electrons not shared are know as unshared
  pairs.

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Single covalent bond example

• Fluorine (F2)

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Single covalent bond example

• Water (H2O)

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Double Triple Covalent Bonds

• Sometimes atoms bond by sharing more than one
  pair of electrons!
• Atoms form double or triple covalent bonds if
  they can attain a noble gas structure by sharing
  two or three pairs of electrons.

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Double Covalent Bond, example

• Oxygen (O2)

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Double Covalent Bond, example

• Carbon dioxide (CO2)

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Triple Covalent Bond, example

• Nitrogen (N2)

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Coordinate Covalent Bonds

• A coordinate covalent bond is a covalent bond in
  which one atom contributes both bonding
  electrons, i.e. Carbon monoxide.

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Polyatomic Ions

• A polyatomic ion is a tightly bound (bonded)
  group of atoms that has a positive or negative
  charge and behaves as a unit, i.e. ammonium ion
  (NH4).

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Bond Dissociation Energies

• The energy required to break the bond between two
  covalently bonded atoms is known as the bond
  dissociation energy.
• Usually expressed as the energy needed to break
  one mole of bonds, or 6.02 x 1023 bonds.
• A large bond dissociation energy corresponds to a
  strong covalent bond.

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VSEPR Theory

• The valence-shell electron-pair repulsion theory,
  or VSEPR theory, explains the three-dimensional
  shapes of molecules.
• According to VSEPR theory, the repulsion between
  electrons pairs causes molecular shapes to adjust
  so that the valence-electron pairs stay as far
  apart as possible.

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Linear geometry
Bent triatomic (triangular)
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Common Molecular shapes
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Bond Polarity

• Nuclei of bonded atoms both pull on the shared
  pair.
• When they pull with the same force (Identical
  atoms), the bonding electrons are shared equally
  and the bond is a non-polar covalent bond.
• A polar covalent bond occurs when the electrons
  are shared unequally.
• The more electronegative atom attracts electrons
  more strongly and gains a slightly negative
  charge. The less electronegative atoms has a
  slightly positive charge.

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Bond Polarity (cont)

• Consider Hydrogen Chloride (HCl). H has an
  electroneg. Of 2.1 and chlorine has an EN of 3.0.
  Who attracts the electrons more? What does that
  do to charge of the molecule?

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Electronegativity Bond Types

• The electronegativity difference between two
  atoms can be used to determine what type of bond
  will be formed.
• EN difference between
• 0 0.5 non-polar covalent
• 0.5 2.1 polar covalent
• 2.1 - 3.0 ionic

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Polar Molecules

• In a polar molecule, one end of the molecule is
  slightly negative and the other end is slightly
  positive.
• A molecule that has two poles is called dipole.

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Attraction Between Molecules

• Molecules can be attracted by a variety of
  forces.
• Intermolecular attractions are weaker than either
  ionic or covalent bonds.
• Theses attractions are responsible for
  determining whether a covalent compound is a
  solid, liquid, or gas.

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Van der Waals Forces

• The 2 weakest forces or attractions between
  molecules are collectively called van der Waals
  forces
• Dipole interactions occur when polar molecules
  are attracted to one another.
• Similar but much weaker than ionic bonds.
• Dispersion forces are attractions between
  molecules caused by the electron motion of one
  molecule affecting the electron motion of another
  through electrical forces weakest of all
  interactions between molecules.

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Hydrogen Bonds

• Hydrogen bonds are attractive forces in which a
  hydrogen covalently bonded to a very
  electronegative atoms is also weakly bonded to an
  unshared electron pair of another electronegative
  atom.

The strong hydrogen bonding between water
molecules accounts for many properties of water,
such as the fact that water is a liquid rather
than a gas at ordinary temperatures.
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Intermolecular Attractions andMolecular
Properties

• The physical properties of a compound depend on
  the type of bonding (ionic or covalent) it
  displays.
• Most molecular solid compounds have weak
  attraction between molecules, however some do not
  melt until temps exceed 1000 F. These very
  stable substance are known as network solids (or
  crystals), solids in which all of the atoms are
  covalently bonded I.e. diamond.
• Most stable of all solids highest melting
  point!!!
• Melting a network solid would require breaking
  covalent bonds throughout the solid.

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Characteristics of Ionic Covalent Compounds
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Naming Writing Formulas forBinary Molecular
Compounds

• Binary Molecular Compounds (BMC) are also
  composed of two elements, but both are nonmetals
  and they are not ions!
• These differences effect the naming formulas

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Naming Binary Molecular Compounds

• Utilize prefixes in the name of BMCs to
  distinguish compounds containing different
  amounts of the same two elements.

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Naming Binary Molecular Compounds (cont)

• All names of BMCs end in -ide
• Vowel at the end of a prefix is often dropped
  when the element name begins with a vowel.
• Omit the prefix mono when the formula contains
  only one atom of the first element
• Examples
• CO
• Carbon monoxide (not monocarbon monoxide)
• CO2
• Carbon dioxide (not monocarbon dioxide)

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Practice Naming Binary Molecular Compounds

• N2O
• dinitrogen monoxide
• SF6
• Sulfur hexafluoride
• SO3
• Sulfur Trioxide
• NCl3
• Nitrogen Trichloride

• N2H4
• Dinitrogen tetrahydride
• N2O3
• Dinitrogen trioxide
• P4O6
• Tetraphosphorus hexoxide
• P4O7
• Tetraphosphorus Heptaoxide
• MgCl2
• Magnesium chloride