The document summarizes different types of chemical bonding: 1. Ionic bonding results from the attraction between oppositely charged ions 2. Covalent bonding results from the sharing of electron pairs between atoms 3. Metallic bonding allows for electron delocalization and mobility in metal solids due to overlapping vacant orbitals, contributing to metals' electrical and thermal conductivity properties.

1. Chapter 6 –
Modern
CHEMICAL BONDING Chemistry
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2. CHEMICAL BONDING
 In a chemical bond between atoms, their valence electrons
are redistributed in ways that make the atoms more stable.
 The way that atoms are redistributed determines the type of
bonding.
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3. T YPES OF CHEMICAL BONDING
1. Ionic Bonding – chemical bonding that results from the
electrical attraction between large numbers of cations
(+, metals) and anions ( -, nonmetals).
2. Covalent Bonding – results from the sharing of electron
pairs between two atoms.
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4. 11/19/2010 S.Martinez 4

5. IONIC OR COVALENT???
 Electronegativity is a measure of an atom’s ability to attract
electrons.
 By calculating the dif ference in the bonded elements’
electronegativity, the type of bond can be determined(ionic or
covalent)….generally bonds are not 100% ionic or covalent but
a mixture.
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6. COVALENT OR IONIC???
Type of Bond Difference in Percentage
electronegati of ionic
vities character
Nonpolar- 0 to 0.3 0 to 5%
covalent
bonds
Polar- 0.3 to 1.7 5 to 50%
covalent
bonds
Ionic bonds 1.7 to 3.3 50 to 100%
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7. 11/19/2010 S.Martinez 7

8. T YPES OF COVALENT BONDING
1. Nonpolar Covalent – a covalent bond in which the bonding
electrons are shared equally by the bonded atoms, resulting
in a balanced distribution of electrical charge.
2. Polar Covalent – is a covalent bond in which the bonded
atoms have an unequal attraction for the shared
electrons…..uneven distribution of charge.
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9.  Example: When Hydrogen and Chlorine combine the
dif ference in their electronegativites is 3.0 – 2.1 = 0.9,
indicating a polar covalent bond. The electrons in this bond
are closer to the more electronegative chlorine atom that to
the hydrogen atom. The chlorine end of the bond has a
partial negative charge, indicated by the symbol δ - . The
hydrogen end has a partial positive charge, δ + .
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10. PRACTICE PROBLEMS
 See periodic table on page 151 for values of electronegativity.
 Classify the following bonds as either ionic, nonpolar
covalent, or polar covalent.
Chlorine & Calcium
Chlorine & Oxygen
Chlorine & Bromine
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11. Bonding Between Electronegativity Bond type More-negative
Chlorine and…. difference atom
Calcium 3.0 – 1.0 = 2.0 Ionic Chlorine
Oxygen 3.5 -3.0 = 0.5 Polar-covalent Oxygen
Bromine 3.0 – 2.8 = 0.2 Nonpolar-covalent Chlorine
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12. CHAPTER 6 SECTION 1
 Page 163, 1 – 4
 Answer the questions in your composition book.
 Do not write the questions unless you feel that it is necessary.
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13. COVALENT BONDING
Chapter 6
AND MOLECULAR Section 2
COMPOUNDS
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14. COVALENT BONDING & MOLECULAR
COMPOUNDS
Many chemical compounds are
composed of molecules.
Molecule – is a neutral group of
atoms that are held together by
covalent bonds….are capable of
existing on its own.
Example: Oxygen (O 2), water (H 2O),
sugar (C 12H22O11)
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15. Molecular compound – is a chemical
compound whose simplest units are
molecules.
Chemical formulas give the
composition of a compound
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16. Chemical formulas also indicate the
relative numbers of atoms of each
kind in a chemical compound by
using atomic symbols and numerical
subscripts….the chem. formula of a
molecular compound is called a
molecular formula.
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17. A molecular formula show the types
and numbers of atoms combined in
a single molecule of a molecular
compound.
A diatomic molecules is a molecule
containing only two atoms..ex:
O 2, N 2, Cl 2, I 2, Br 2, F 2, H 2
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18. FORMATION OF COVALENT BONDS
Most atoms are at a lower potential
energy when they are bonded to
other atoms than they are at as
independent particles….this is what
makes them more stable.
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19. There are three different types of forces
at work when atoms come together.
(electron-electron, proton-proton, &
electron-proton).
A bond length is the distance between
two bonded atoms at their minimum
potential energy….this occurs at a
distance when the forces of attraction =
the forces of repulsion.
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20. http://www.youtube.com/user/TTUchem10
10#p/u/22/z3F7LjTvdX0
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21.  Bond energy is the energy required to break a chemical bond
and form neutral isolated atoms…this is the same amount of
energy that is released as atoms change from isolated
individual atoms to part of a molecule.
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22. Bond Bond length Bond Energy Bond Bond length Bond Energy
(pm) (kJ/mol) (pm) (kj/mol)
TABLE 6-1, PAGE 168
H-H 74 436 C-C 154 346
F-F 141 159 C-N 147 305
Cl-Cl 199 243 C-O 143 358
Br-Br 228 193 C-H 109 418
I-I 267 151 C-Cl 177 327
H-F 92 569 C-Br 194 285
H-Cl 127 432 N-N 145 163
H-Br 141 366 N-H 101 386
H-I 161 299 O-H 96 459
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23.  Bond lengths and energies vary with the type of atoms that
have combined…and can vary with the same type of atoms
bonding together. Bond length will decrease as bond energy
or strength increases.
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24. THE OCTET RULE
Chemical compounds tend to form
so that each atom, by
gaining, losing, or sharing
electrons, has an octet of electrons
in its highest occupied energy level.
(this rule applies to both ionic and
covalent compounds)
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25. Noble-gas atoms exist independently
in nature….because of their electron
configurations. This stability is a
result of their outermost s and p
orbitals being completely filled with
8 electrons….their octet is full.
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26. Other main-group elements (s and p
block elements) are able to obtain a
stable configuration through the use
of the octet rule.
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27. 11/19/2010 S.Martinez 27

28. EXCEPTIONS TO THE OCTET RULE
 Hydrogen forms bonds in which it is only surrounded by two
electrons.
 Boron, has just three valence electrons, so it tends to form
bonds in which it is only surrounded by 6 electrons. Example:
BF 3
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29. ELECTRON-DOT NOTATION
 Shows the electron-configuration notation in which only the
valence electrons of an atom of a particular element are
shown, indicated by dots placed around the element’s symbol.
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30. LEWIS STRUCTURES
 Uses dashes to show covalent bonds between atoms and uses
dots to show the remaining valence or the unshared pairs of
electrons.
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31. STRUCTURAL FORMULA
 Indicates the kind, number, arrangement, and bonds but not
the unshared pairs of the atoms in a molecule.
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32. Draw the Lewis structure
of iodomethane, CH 3I
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33. Draw the Lewis
structure of
ammonia, NH 3
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34. Draw the Lewis
structure for hydrogen
sulfide, H 2S
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35. Draw the Lewis
structure for methanal
(aka
formaldehyde), CH 2O
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36.  PCl5
C6H10Cl2OH
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37. BOND LENGTHS AND ENERGIES FOR SINGLE
AND MULTIPLE BONDS
Bond Bond Bond Bond Bond Bond
Length (pm) Energy Length (pm) Energy
(kJ/mol) (kJ/mol)
C-C 154 346 C-O 143 358
C=C 134 612 C=O 120 732
C C 120 835 C O 113 1072
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38. RESONANCE
 Refers to bonding in molecules that cannot be correctly
represented by a single Lewis structure…..the true structure
lies somewhere between the two resonance structures. A
double arrow is used indicate a molecule’s resonance
structure.
 O=O-O O-O=O
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39. IONIC BONDS AND Chapter 6
Section 3
COMPOUNDS
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40. IONIC BONDING & IONIC COMPOUNDS
Most of the rocks & minerals in the
Earth’s crust are composed of positive
& negative ions held together by bonds.
Ionic compound – composed of positive
and negative ions that are combined so
that the numbers of positive and
negative charges are equal…generally
crystals.
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41. Chem. Formula for an ion represents
the simplest ratio of the combined
ions that gives electrical neutrality.
Formula unit shows the simplest
collection of atoms from which an
ionic compound’s formula can be
established.
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42. Example: Rock Salt, sodium
chloride, NaCl is composed of Na + & Cl -
Example: calcium fluoride(the liquid
form is used to melt ice on highways or
in the use of oxygen sensitive
applications such as the making of
metal alloys), CaF 2 is composed of Ca 2+
& F-
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43. ELECTRON-DOT NOTATION….
Can be used to demonstrate the
changes that occur in ionic bonding.
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44. CHARACTERISTICS OF IONIC BONDING
Since nature
favors
arrangements
where PE is @ a
min. ionic crystals
are arranged in a
crystal lattice.
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45. FORCES AT ACT IN THE CRYSTAL LATTICE
Forces of attraction Forces of repulsion =
= oppositely charged between like
ions & those charged ions &
between the nuclei those between
& electrons of electrons of
adjacent ions. adjacent ions.
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46. IONIC VS. MOLECULAR
Ionic:  Molten state can
conduct since ions are
the forces that hold free to move.
it together are very  Some ionic compounds
strong. do not dissolve in water
because the attraction
Higher melting & of water cannot
boiling pts. Also, do overcome the attraction
not vaporize @ room between ions.
temp.  Solid state ions cannot
move so they do not
 Hard but brittle due to conduct.
a large build-up of
repulsion when a shift
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occurs.

47. Molecular – covalent
bonds are also
strong but not in
comparison to ionic.
Melt at lower
temperatures or are
gases @ room temp.
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48. POLYATOMIC IONS
Is a charged group Examples:
of covalently Ammonium
bonded NH 4+, Nitrate NO 3-
atoms…can be , Sulfate SO 42- &
represented by Phosphate PO 42-
Lewis structures
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49. Chapter 6
METALLIC BONDING Section 4
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50. METALLIC BONDING
Explains why they This is due to the
are such excellent highly mobile
conductors of heat & valence electrons of
electricity in the the atoms that make
solid state compared up a metal.
to molten ionic
compounds.
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51. This mobility is not
possible in
molecular or ionic
compounds since
they are localized or
bound to individual
ions that are held in
place.
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52. The vacancy that is These electrons are
observed in s-block considered to be
and d-block metals delocalized,
allows for an overlap meaning they do not
of the vacant belong to an
orbitals that in turn individual atom.
allows the outer These mobile
electrons to roam. electrons are
referred to as a sea
of electrons.
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54. This freedom Also, since metallic
accounts for  bonding is the same
electrical & thermal in all directions this
conductivity. accounts for why
Metals can absorb metals are both
wide range of light malleable
frequencies (Flame (hammered into thin
Test) responsible for sheets) or ductile
shiny appearance. (drawn into a wire).
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55. METALLIC BOND STRENGTH
Varies with nuclear The amt of  (heat)
charge and the # of required to vaporize
electrons in the the metal is a
metal’s electron sea. measure of the
This is reflected in strength of the
the heat of bonds that hold the
vaporization of a metal together.
metal.
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56. TABLE 6-4 HEATS OF VAPORIZATION OF
SOME METALS (KJ/MOL)
Period Element Element Element
Second Li – 147 Be – 297
Third Na – 97 Mg – 128 Al – 294
Fourth K – 77 Ca – 155 Sc – 333
Fifth Rb – 76 Sr – 137 Y – 365
Sixth Cs – 64 Ba – 140 La - 402
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57. MOLECULAR GEOMETRY
 VSEPR Theory (valence-shell, electron-pair repulsion) states
that the repulsion between the sets of valence -level electrons
surrounding an atom causes these sets to be oriented as far
apart as possible.
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58. VESPR MODEL
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59. VESPR MODEL
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60.  Hybridization is the mixing of two or more atomic orbitals of
similar energies on the same atom to produce new orbitals of
equal energies.
 Example: Methane CH 4 1s 2 2s 2 2p 2 (see board)
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61. GEOMETRY OF HYBRID ORBITALS
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62. VESPR AND MOLECULAR GEOMETRY
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63. VESPR AND MOLECULAR GEOMETRY
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64. PREDICT THE MOLECULAR GEOMETRY FOR
THE FOLLOWING:
 HI
 CBr 4
 AlBr 3
 CH 2 Cl 2
 NH 3
 H 2O
 PCl 5
 SF 6
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65.  Intermolecular Forces – are the forces of attraction between
molecules…
 The stronger the forces are the higher the boiling point will
be..
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67. PREDICTING POLARIT Y
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68. PREDICTING POLARIT Y
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69. MOLECULAR POLARIT Y
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