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
• Right side of table
• Tend to gain electrons and form negative ions
• (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
• Show characteristics of both
• Metal Nonmetal
• Usually electrons are transferred from metal to
  nonmetal

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Atomic Radius

• Increases down each group
• Each step down additional energy level
• More energy levels greater distance from
  nucleus large average size
  
• Decreases from left to right across a period
• Electrons added to the same energy level (also
  means protons added)
  
• 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
  
• More energy required to remove 2nd, 3rd, 4th,
  5th, etc. electron
  
• Noble gas structure is stable so takes large
  amount of energy to remove an electron
  

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Ionization Energy

• Ionization energy in Group A elements decreases
  as you move down a group
  
• Ionization energy increases from left to right
  across a period
  
• Metals some give up electrons more easily than
  others
  
• 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
• American chemist Gilbert N. Lewis
• Dots number of s and p electrons in outermost
  energy level
  
• Paired dots paired electrons
• Simple way of showing electrons
• Most reactions involve only outermost electrons

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Lewis Structures

• When drawing
• Use electron configuration
• Move in clockwise direction
• 12 s orbital
• 3, 6, 9 p orbitals fill each with ONE
  electron before filling with pairs
  
• 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 the attraction between oppositely
  charged ions
  
• Transfer of electrons from one atom to another
• Ions form with or charges
• Attraction between electrostatic charges is a
  strong force which holds
  atoms together

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The Ionic Bond
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The Ionic Bond

• NOT A MOLECULE
• Bond not just between (for example) one sodium
  and one chloride
  
• One sodium ion attracts 6 chlorine ions

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The Ionic Bond

• Typically metal nonmetal
• Metals usually lose electrons
• Nonmetals usually gain electrons

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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?
  
• Ba must lose 2 electrons Xe6s2
• Forms the ion Ba2
• S must gain 2 electrons Ne3s23p4
• Forms the ion S2-
• Somust be ratio of 1 to 1 when barium and sulfur
  combine
  
• BaS

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Predicting Formulas of Ionic Compounds

• Elements in a family usually form compounds with
  the same atomic ratios
  
• Because they have the same number of valence
  electrons
  
• Must gain or lose the same number of electrons
• See table 11.4 pg 233

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Predicting Formulas of Ionic Compounds

• The formula for sodium oxide is Na2O. Predict
  the formula for
  
• Sodium sulfide
• Sodium Ne3s1 must lose one electron
• Sulfur Ne3s23p4 must gain two electrons
• Soformula must be Na2S
• Sulfur is in same family as oxygenso same ratio

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Predicting Formulas of Ionic Compounds

• Rubidium Oxide
• Rubidium Kr5s1 must lose one electron
• Oxygen He2s22p4 must gain two electrons
• Soformula must be Rb2O
• This makes sense b/c rubidium is in same family
  as sodium

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The Covalent Bond

• A pair of electrons shared between two atoms
• Most common type of bond
• Stronger than ionic bond
• Electron orbital expands to include both nuclei
• most often found between two nuclei
• Negative charges allow positive
  nuclei to be drawn close to each other

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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 two pairs being shared
• Triple bond three pairs being shared
• Multiple bonds are stronger than single bonds
• Harder to break
• Covalent bonding between identical atoms means
  electrons are shared equally
  
• Covalent bonding between different atoms leads to
  unequal sharing (polar covalent bond)

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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 partial positive charge
• The other assumes a partial negative charge

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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
• Electronegativities are equal
• Electrons shared equally
• Polar covalent
• One electronegativity is larger.
• Ionic compound
• Electronegativity is so great that no
  sharing occurs
  
• Electrons are lost or gained instead

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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 ionic bond
• Below 1.5 nonpolar covalent
• See Continuum on page 239

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Electronegativity

• Polar bonds form between two atoms
• Molecules can also be polar or nonpolar
• Dipole
• Electrically asymmetrical
• Oppositely charged at two points
• Polar
• Usually molecules with only one polar bond
• Part of molecule has charge, part has charge
• Nonpolar
• Molecules with multiple bonds
• Dipoles cancel each other by acting in opposite
  directions
  
• Overall has no charge

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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 vond
  
• 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

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Lewis Structures of Compounds

• Carbon Dioxide, CO2
• Total valence electrons 16

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-

Carbon only has 6 electrons try double bonds
more than one location..form resonant structures
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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 molecular interactions
• 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 from Lewis
  structures
  
• Electron pairs will repel each other
  electrically
  
• Try to minimize this repulsion
• Arranged as far apart as possible around a
  central atom

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The VSEPR Model

• Linear Structure
• Two pairs of electrons surrounding the central
  atom
  
• 180o apart

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The VSEPR Model

• Trigonal Planar
• Three pairs of electrons around the central atom
• 120o apart

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The VSEPR Model

• Tetrahedral structure
• Four pairs of electrons on central atom
• 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 tetrahedral
• But, moledule is bent
• Water
• Helps explain some properties

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The VSEPR Model

• Predict the shape for CF4, NF3, and BeI2.
• 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

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The VSEPR Model

• CF4
• Lewis Structure (on board)
• Number of electron pairs 4
• Electron pair arrangement tetrahedral
• Molecular shape tetrahedral
• NF3
• Lewis Structure (on board)
• Number of electron pairs 4
• Electron pair arrangement tetrahedral
• Molecular shape pyramidal

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The VSEPR Model

• BeI2
• Lewis Structure (on board)
• Number of electron pairs 2
• Electron pair arrangement linear
• Molecular shape linear

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Homework

• Questions 1-6 14-16
• Paired Exercises 19-29 odd 33-57 odd
• Additional Exercises 59 65-66 69-72