Chapter 10 Chemical Bonding

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Chapter 10ChemicalBonding

• Tro, 2nd ed.

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LEWIS STRUCTURES OF ATOMS

• Metals form cations and nonmetals form anions to
  attain a stable valence electron structure.
• These rearrangements occur by losing, gaining or
  sharing electrons.
• The Lewis structure of an atom is a
  representation that shows the valence electrons
  for that atom.
• Valence electrons the electrons that occupy the
  outermost energy level of an atom.
• Valence electrons are responsible for the
  electron activity that occurs to form chemical
  bonds.

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The Lewis structure of an atom uses dots to show
the valence electrons of atoms.
B
2s22p1
The number of dots equals the number of s and p
electrons in the atoms outermost shell.
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Lewis Structures of the first 20 elements
Notice that Cs e- config is 2s22p2.
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CHEMICAL BONDING

• Atoms will do one of three things to get to a
  noble gas electron configuration
• 1. Take electrons from another atom
• 2. Give electrons to another atom
• 3. Share electrons with atom(s)
• In choices 1 2 cause ions to form, then ionic
  bonds
• In choice 3 sharing electrons results in
  covalent bonds
• With the exception of hydrogen helium, this
  structure consists of eight electrons in the
  outermost energy level (The Octet Rule)

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The Ionic Bond

• Cations () have given up e-s and anions (-) have
  gained e-s, and now have opposite electrical
  charges
• Results in strong electrostatic force of
  attraction
• All cations and anions exist in crystal lattices,
  defined geometric structures with repeating 3-D
  pattern

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The 3s electron of sodium transfers to the 3p
orbital of chlorine.
The force holding Na and Cl- together is an
ionic bond.
Lewis representation of sodium chloride formation.
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The forces holding Mg2 and two Cl- together are
ionic bonds.
Two 3s electrons of magnesium transfer to the 3p
orbitals of two chlorine atoms.
A magnesium ion (Mg2) and two chloride ions
(Cl-) are formed.
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NaCl is made up of cubic crystals.
In the crystal each sodium ion is surrounded by
six chloride ions.
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In the crystal each chloride ion is surrounded by
six sodium ions.
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The ratio of Na to Cl- is 11
There is no molecule of NaCl
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The Ionic Bond

• Using your Periodic Table, determine the cation
  and anion each atom is likely to form, then write
  the Lewis structures and made the compounds.
  Finish by writing the compounds chemical
  formula.
• Practice Al F, Mg O, Na O, Na
• N, Al O

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A sodium ion is smaller than a sodium atom
because
(1) the sodium atom has lost its outermost
electron.
(2) the 10 remaining electrons are now attracted
by 11 protons and are drawn closer to the nucleus.
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A chloride ion is larger than a chloride atom
because
(1) the chlorine atom has gained an electron
and now has 18 electrons and 17 protons.
(2) The nuclear attraction on each electron has
decreased, allowing the chlorine to expand.
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Transition Metals form cations a little
different

• Transition metals lose their s electrons first,
  because they are in the highest principle energy
  level, then they lose their d electrons.
• Zn ? Zn2 2 e- Cu ? Cu 1e-
• Ar4s23d10?Ar3d10 Ar4s13d10?Ar3d10

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COVALENT BONDING

• A covalent bond consists of a pair of electrons
  shared between two atoms.
• In the millions of chemical compounds that exist,
  the covalent bond is the predominant chemical
  bond.
• Substances which covalently bond exist as
  molecules.

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Carbon dioxide bonds covalently. It exists as
individually bonded covalent molecules containing
one carbon and two oxygen atoms.
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The term molecule is not used when referring to
ionic substances.
Instead they are called Formula Units.
Sodium chloride bonds ionically. It consists of
a large aggregate of positive and negative ions.
No molecules of NaCl exist.
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COVALENT BONDING

• Nonmetal Atoms have deficiency of electrons in
  outermost shell and want to gain electrons to get
  full shell
• Since two nonmetal atoms both want more
  electrons, they will share electrons to get full
  shell
• H has 1 e- and wants 2
• Cl has 7 e-s and wants 8
• Both satisfied if they share a pair of electrons
  between them
• Each contributes 1 e- to the pair and each gets
  to share the 2 e-s in the pair
• H has 2 e-s and Cl has 8 e-s and they are HAPPY

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COVALENT BONDING

• IMPORTANT in giving e-s to be shared, atom
  actually gains e-s
• The number of e-s an atom contributes to be
  shared is equal to the number of e-s it needs to
  have an octet! (or a full shell)
• A pair of shared e-s is called a covalent bond
• 1 pair of e-s between two atoms single bond
• 2 pairs of e-s betwn two atoms double bond
• 3 pairs
  triple bond

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LEWIS STRUCTURES OF COMPOUNDS

• In writing Lewis structures, the most important
  consideration for forming a stable compound is
  that the atoms attain a noble gas configuration.
• The most difficult part of writing Lewis
  structures is determining the arrangement of the
  atoms in a molecule or an ion.
• In simple molecules with more than two atoms, one
  atom will be the central atom surrounded by the
  other atoms.

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Cl2O has two possible arrangements.
The two chlorines can be bonded to each other.
Cl-Cl-O
The two chlorines can be bonded to oxygen.
Cl-O-Cl
Usually the single atom will be the central atom.
(also usually the leftist or lowest on the
Periodic Table)
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Valence Electrons of Group A Elements
Atom Group Valence Electrons
Cl VIIA 7
H IA 1
C IVA 4
N VA 5
S VIA 6
P VA 5
I VIIA 7
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Procedures for WritingLewis Structures

• Step 1. Obtain the total number of valence
  electrons to be used in the structure by adding
  the number of valence electrons in all the atoms
  in the molecule or ion.
• If you are writing the structure of an ion, add
  one electron for each negative charge or subtract
  one electron for each positive charge on the ion.
• Step 2. Write the skeletal arrangement of the
  atoms and connect them with a single covalent
  bond (two dots or one dash). Choose the leftist
  or lowest element as the central atom. Arrange
  terminal atoms symmetrically around the central
  atom.
• Hydrogen, which contains only one bonding
  electron, can form only one covalent bond.
• Oxygen atoms usually have a maximum of two
  covalent bonds (two single bonds, or one double
  bond).

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Procedures for WritingLewis Structures

• Step 3. Subtract two electrons for each single
  bond you used in Step 2 from the total number of
  electrons calculated in Step 1.
• This gives you the net number of electrons
  available for completing the structure by adding
  lone pairs of electrons to the terminal atoms
  until they have an octet. Any remaining
  electrons become lone pairs on the central atom.

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Procedures for WritingLewis Structures

• Step 4. Check that each atom is satisfied. If
  one atom doesnt have an octet, move lone pairs
  of electrons in as bond pairs to make multiple
  covalent bonds. Do this symmetrically.
• Step 5. Check the total number of electrons in
  the structure and make sure it matches the number
  of valence electrons in step 1.
• (Also learn the number of bonds an atom prefers
  to make
• H and F always 1 bond and terminal atom C mostly
  4 (and usually a central atom) halogens mostly
  1 O and S mostly 2 N and P mostly 3)

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Write the Lewis structure for H2O. The total
number of valence electrons is eight, two from
the two hydrogen atoms and six from the oxygen
atom. The two hydrogen atoms are connected to the
oxygen atom which is central. Write the skeletal
structure
Place two dots between the hydrogen and oxygen
atoms to form the covalent bonds. Subtract the
four electrons used from eight valence electrons
to obtain four electrons yet to be used around
the oxygen. (Why not the H?)
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Distribute the four remaining electrons in lone
pairs around the oxygen atom. (Hydrogen atoms
cannot accommodate any more electrons. NEVER
have more than 1 bond to H or have lone pairs
around H.)
The shape of the molecule is not shown by the
Lewis structure.
These arrangements are Lewis structures because
each atom has a noble gas electron structure.
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Covalent bonding with equal sharing of electrons
occurs in diatomic molecules formed from one
element.
chlorine
iodine
nitrogen
hydrogen
A dash may replace a pair of dots that represent
a bond
H-H
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Practice Lewis Structures

• Look in packet for practice sheet and work with
  one partner to draw the Lewis structures on
  separate paper.
• Bring them up to show on the document camera.

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Complex Lewis Structures

• Do the Lewis structures for the following with a
  partner
• HCN, CH4, SO3, CH3OH, SF6, PCl3, NO2
• Some will have EXCEPTIONS to the Octet Rule.

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Complex Lewis Structures

• Exceptions to Octet Rule
• Expanded valence shell any central atom with
  outermost e-s in period 3 or below has d orbitals
  available for bonding and can hold 10 or 12 e-s
• Electron deficient or free radical structures
  have less than 8 e-s and will be very reactive
  compounds

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Complex Lewis Structures

• There are some molecules and polyatomic ions for
  which no single Lewis structure consistent with
  all characteristics and bonding information can
  be written.
• When more than one structure satisfies the rules,
  we call them resonance structures.
• Real molecule is a hybrid of all possible Lewis
  structures.
• Resonance stabilizes the molecule.
• Try O3.

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DRAWING LEWIS STRUCTURES

• Multiple Bonds O2 and N2
• Multiple Central Atoms C2H6, N2H4, C3H8, C6H6,
  CH3NH2, CH3COOH

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Compounds ContainingPolyatomic Ions

• A polyatomic ion is a stable group of atoms that
  has either a positive or negative charge and
  behaves as a single unit in many chemical
  reactions.
• Practice NH4, SO32-, NO2-, NO3-, I3-,

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A scale of relative electronegativities was
developed by Linus Pauling. Electronegativity
decreases down a group for representative
elements.
Electronegativity generally increases left to
right across a period.
Metals are low in EN and nonmetals are high.
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ELECTRONEGATIVITY

• Electronegativity The relative attraction that
  an atom has for a pair of shared electrons in a
  covalent bond.
• If the two atoms that constitute a covalent bond
  are identical then there is equal sharing of
  electrons.
• This is called nonpolar covalent bonding.
• Ionic bonding and nonpolar covalent bonding
• represent two extremes.

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ELECTRONEGATIVITY

• If the two atoms that constitute a covalent bond
  are not identical then there is unequal sharing
  of electrons.
• This is called polar covalent bonding.
• One atom assumes a partial positive charge and
  the other atom assumes a partial negative
  charge.
• This charge difference is a result of the unequal
  attractions the atoms have for their shared
  electron pair.

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ELECTRONEGATIVITY

• H-H are the same atom, and have the same
  greediness, so the two atoms are forced to
  share equally.
• F-F same - forced to share equally.
• If two atoms have diff EN, the one with higher EN
  will take the e-s in the pair more often than
  the other atom.
• H-F are not the same atoms, and are not equal in
  greediness, F is far greedier, takes the e-s more
  than half the time.

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Polar Covalent Bonding in HF
Partial positive charge on hydrogen.
Partial negative charge on fluorine.


The shared electron pair is closer to fluorine
than to hydrogen.
Shared electron pair.
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Types of Covalent Bonding

• The polarity of a bond is determined by the
  difference in electronegativity values of the
  atoms forming the bond.
• If the electronegativity difference between two
  bonded atoms is greater than 1.9 to 2.0, the bond
  will be more ionic than covalent.
• If the electronegativity difference is greater
  than 2, the bond is strongly ionic.
• If the electronegativity difference is less than
  1.9 but greater than 0.5, the bond is polar
  covalent.
• If the electronegativity differences is 0.5 or
  less, the bond in nonpolar covalent.

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Types of Covalent Bonding

• Estimate whether a bond is polar cov, mostly pure
  or nonpolar cov or ionic by finding the absolute
  value of the difference in EN between the two
  atoms in the bond.
• __________________________________________
• 
  
• 0 .2 .4 .6 .8 1 1.2 1.4 1.6
  1.8 2.0 2.2 2.4 2.6
• nonpolar polar cov
  mostly ionic
• covalent
• Practice H-Cl, C-Cl, C-O and CO, N-Cl, Ca-N
• H-Cl C-Cl C-O CO N-Cl Ca-N
• 0.9 0.5 1.0 1.0 0.0 2.0
• pol cov nonpol ionic

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Molecular Geometry
(bent)
bent
One more geometry is trigonal pyramidal.
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Some Geometric Figures

• Linear
• 2 atoms on opposite sides of central atom
• 180 bond angles
• Trigonal Planar
• 3 atoms form a triangle around the central atom
• Planar
• 120 bond angles
• Tetrahedral
• 4 surrounding atoms form a tetrahedron
  around the central atom
• 109.5 bond angles

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Some Geometric Figures

• Trigonal Pyramidal
• 3 atoms form a triangular pyramid beneath the
  central atom
• Not planar
• 109 bond angles
• Derivative of tetrahedral geometry

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The Valence ShellElectron Pair Repulsion (VSEPR)
Model

• The VSEPR model is based on the idea that
  electron pairs will repel each other electrically
  and will seek to minimize this repulsion.
• To accomplish this minimization, the electron
  pairs will be arranged as far apart as possible
  around a central atom.
• The 3-dimensional arrangement of the atoms within
  a molecule determines molecular interactions
  (physical properties and chemical reactions).

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BeCl2 is a molecule with only two pairs of
electrons around beryllium, its central atom.
Its electrons are arranged 180o apart for maximum
separation.
LINEAR EP ARRANGEMENT MOLECULAR GEOMETRY
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• BF3 is a molecule with three pairs of electrons
  around boron, its central atom.

Its electrons are arranged 120o apart for maximum
separation.
This arrangement of atoms is called trigonal
planar.
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• CH4 is a molecule with four pairs of electrons
  around carbon, its central atom.

An obvious choice for its atomic arrangement is a
90o angle between its atoms with all of its atoms
in a single plane.
However, since the molecule is 3-dimensional the
molecular structure is tetrahedral with a bond
angle of 109.5o.
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• Ball and stick models of methane, CH4, and carbon
  tetrachloride, CCl4.

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Tetrahedral Shapes

• Tetrahedral
• 4 areas of electrons around the central atom
• 109.5 bond angles
• All Bonding tetrahedral
• 3 Bonding 1 Lone Pair trigonal pyramid
• 2 Bonding 2 Lone Pair bent

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Tetrahedral Derivatives
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• Ammonia, NH3, has four electron pairs around
  nitrogen.

The arrangement of electron pairs around nitrogen
is tetrahedral.
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NH3 has one lone pair of electrons.
The NH3 molecule is trigonal pyramidal.
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• Water has four electron pairs around oxygen.

The arrangement of electron pairs around oxygen
is tetrahedral.
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H2O has two lone pairs of electrons.
The H2O molecule is bent.
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The VSEPR Model

• Summary electron pair arrangement depends upon
  of bonded atoms and lone pairs around the central
  atom. Lone pairs exert more repulsion that bond
  pairs. Count of bonded atoms (B) and of LPs
  (E).
• EP Arrangements 2 Linear 3 Trigonal
  Planar 4 Tetrahedral
• Within these electron pair arrangements, the
  molecular geometry is based only on seeing the
  atoms. I call the molecular geometry the family
  members and the electron pair arrangement is the
  electron pair family the members are in.
• Electron Pair Families and their Molecular
  Geometry members
• Linear AB2, linear only
• Trigonal Planar AB3 trigonal planar AB2E bent
• Tetrahedral AB4 tetrahedral AB3E trigonal
  pyramidal, and AB2E2 bent

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VSEPR Practice

• Use practice sheets in packet to fill in Electron
  Pair Arrangement and Molecular Geometry and Bond
  Angles.

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Dipole Moments

• A dipole is a molecule with positively and
  negatively charged ends
• Polar covalent bonds or molecules have one end
  slightly positive, d and the other slightly
  negative, d-
• (not full charges, come from nonsymmetrical
  electron distribution)
• Dipole Moment is a measure of the size of the
  polarity. (We are NOT going to worry about the
  Debye unit or actual numbers for dipole moment,
  just whether a molecule has a dipole or not.)

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Polarity of Molecules

• For a molecule to be polar it must
• - have polar bonds
• electronegativity difference - theory
• bond dipole moments measured values
• - have an unsymmetrical shape
• vector addition
• Polarity affects the intermolecular forces of
  attraction

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DIPOLE
NO DIPOLE
Polar covalent bonds, but nonpolar
molecule, because vectors cancel
Polar covalent bonds and unsymmetrical shape
cause molecule to be polar
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Adding Dipole Moments
Table 10.3