Chapter 11 Chemical Bonds: The Formation of Compounds from Atoms

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Chapter 11Chemical BondsThe Formation of
Compounds from Atoms

• Objectives
• Describe the trends in the periodic table
• Know how to draw Lewis Structures of atoms
• Understand and predict the formation of ionic
  bonds
• Understand and predict covalent bonds
• Describe electronegativity
• Know how to draw complex lewis structures of
  compounds
• Understand the formation of compounds containing
  polyatomic ions
• Describe molecular shape, including the VSEPR
  model

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Periodic Trends in Atomic Properties

• Periodic table designed to show trends
• Use trends to predict properties and reactions
  between elements
• Trends include
• Metals, nonmetals, metalloids
• Atomic radius
• Ionization energy
• Electronegativity

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Metals, Nonmetals and Metalloids

• Metals
• Lustrous, malleable, good conductors of heat and
  electricity
• Left-hand side of table
• Most elements are metals
• Tend to lose electrons and form positive ions

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Metals, Nonmetals and Metalloids

• Nonmetals
• Nonlustrous, brittle, poor conductors
• (Hydrogen displays nonmetallic properties under
  normal conditions but is UNIQUE element)

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Metals, Nonmetals and Metalloids

• Metalloids
• Found along border between metals and nonmetals
• Metal Nonmetal
• Usually electrons are transferred from metal to
  nonmetal

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Atomic Radius

• Increases down each group
• Decreases from left to right across a period
• Increase in positive charge stronger pull on
  electrons gradual decrease in atomic radius

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Atomic Radius
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Ionization Energy

• The energy required to remove an electron from
  the atom

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Ionization Energy

• Ionization energy in Group A elements
• Ionization energy
• Nonmetals tend to gain electrons (rather than
  give them up)

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Ionization Energy
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Lewis Structures

• Diagram that shows valence electrons
• Dots number of s and p electrons
• Paired dots
• Simple way of showing electrons

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Lewis Structures

• When drawing
• 3, 6, 9
• Just like orbital filling diagram
• Examples draw Lewis Structures of B, N, F, Ne

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Lewis Structures
F
B

• N

Ne
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The Ionic Bond

• Ionic bond
• Attraction between electrostatic charges is a

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The Ionic Bond
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The Ionic Bond

• NOT A MOLECULE
• Bond not just between

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The Ionic Bond

• Typically metal nonmetal

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Predicting Formulas of Ionic Compounds

• In almost all stable chemical compounds of
  representative elements, each atom attains a
  noble gas electron configuration. This concept
  forms the basis for our understanding of chemical
  bonding.

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Predicting Formulas of Ionic Compounds

• How many electrons must be gained or lost to
  achieve noble gas configuration?

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Predicting Formulas of Ionic Compounds

• Elements in a family usually form compounds with
  the same atomic ratios

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Predicting Formulas of Ionic Compounds

• The formula for sodium oxide is Predict the
  formula for
• Sodium sulfide
• Sodium Ne3s1 must
• Sulfur Ne3s23p4 must
• Soformula must

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Predicting Formulas of Ionic Compounds

• Rubidium Oxide
• Rubidium Kr5s1 must
• Oxygen He2s22p4
• Soformula must be
• This makes sense b/c rubidium is in same family
  as sodium

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The Covalent Bond

• A pair of electrons
• Most common type of bond
• Electron orbital expands to include both nuclei

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The Covalent Bond
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The Covalent Bond

• Atoms may share more than one pair of electrons
• Double bond
• Triple bond
• Multiple bonds are
• Covalent bonding between identical atoms means
  electrons are
• Covalent bonding between different atoms leads to
  

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Electronegativity

• The attractive force that an atom of an element
  has for shared electrons
• Atoms have different electronegativities
• Electrons will spend more time near atom with
  stronger (larger) electronegativity
• Soone atom assumes a
• The other assumes a

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Electronegativity

• Electronegativity trends and periodic table
• See table 11.5 page 237
• Generally increases from left to right
• Decreases down a group
• Highest is fluorine (4.0)
• Lowest is francium (0.7)

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Electronegativity
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Electronegativity

• Polarity is determined by difference in
  electronegativity
• Nonpolar covalent
• Polar covalent
• Ionic compound

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Electronegativity
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Electronegativity

• If the electronegativity difference is greater
  than 1.7-1.9 then the bond will be more ionic
  than covalent
• Above 2.0
• Below 1.5
• See Continuum on page 239

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Electronegativity

• Polar bonds form between two atoms
• Molecules can also be polar or nonpolar
• Dipole
• Polar
• Nonpolar

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Lewis Structures of Compounds

• Convenient way of showing ionic or covalent bonds
• Usually the single atom in a formula is the
  central atom

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The Ionic Bond

• LEWIS STRUCTURES of ionic bonds

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The Covalent Bond

• LEWIS STRUCTURES of covalent bonds
• Use dashes instead of dots

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The Covalent Bond






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Lewis Structures of Compounds

• Obtain the total number of valence electrons
• Add the valance electrons of all atoms
• Ionic add one electron for each negative charge
  and subtract one electron for each positive charge

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Lewis Structures of Compounds

• Write the skeletal arrangement of the atoms and
  connect with a single covalent bond
• Subtract two electrons for each single bond
• This gives you the net number of electrons
  available for completing the structure

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Lewis Structures of Compounds

• Distribute pairs of electrons around each atom to
  give each atom a noble gas structure
• If there are not enough electrons then try to
  form double and triple bonds

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Lewis Structures of Compounds

• Write the Lewis Structure for methane CH4
• Total number of valence electrons is eight
• Draw skeletal structure
• Dashes equal two electrons being shared
• Subtract the eight electrons shown as dashes
• Check that all atoms have a noble gas structure

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Lewis Structures of Compounds

• Methane, CH4

H
H
H
C
H
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Lewis Structures of Compounds

• Carbon Dioxide, CO2
• Total valence electrons 16

O C O
Not Enough! Must try double bonds
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Complex Lewis Structures

• Some molecules and polyatomic ions have strange
  behaviors
• No single Lewis structure is consistent
• If multiple structures are possible the molecule
  shows resonance
• Resonance structures show all possibilities

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Complex Lewis Structures

• Carbonate ion, CO32-

2-
2-
2-
Carbon only has 6 electrons try double bonds
more than one location..form resonant structures
O C O O
O C O O
O C O O
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Compounds ContainingPolyatomic Ions

• Polyatomic ion stable group of atoms that has a
  positive or negative charge
• Behaves as a single unit in many chemical
  reactions
• Sodium carbonate (Na2CO3)
• Carbonate ion (co3) has covalent bonds
• Sodium atoms are ionically bonded to carbonate ion

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Compounds ContainingPolyatomic Ions

• Easier to dissociate ionic bond than break
  covalent bond
• More in chapters 6 and 7

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Molecular Shape

• Three-dimensional shape of molecule important
• Explains
• Helpful to know how to predict the geometric
  shape of molecules
• Linear?
• V-shaped?
• Trigonal planar?
• Tetrahedral?

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The VSEPR Model

• Valence Shell Electron Pair Repulsion Model
• Make predictions about shape
• Electron pairs will

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The VSEPR Model

• Linear Structure
• 180o apart

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The VSEPR Model

• Trigonal Planar
• 120o apart

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The VSEPR Model

• Tetrahedral structure
• 109.50 apart
• When drawing
• Wedged line to show atom protruding from page
  dashed line to show atom receding from page

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The VSEPR Model

• Pyramidal shape
• Four pairs of electrons on central atom BUT only
  three shared
• Electrons are tetrahedral but actual shape is
  more of a pyramid

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The VSEPR Model

• Electron pairs determine shape BUT name for shape
  is determined by position of atoms

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The VSEPR Model

• V-shaped or bent
• Four electron pairs but only two shared
• Electron arrangement is
• But, molecule is
• Water
• Helps explain some properties

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The VSEPR Model

• Predict the shape for .
• Draw the Lewis Structure
• Count the electron pairs and determine the
  arrangement that will minimize repulsions
• Determine the positions of the atoms and name the
  structure