Pre-University Chemistry

1. 11
Chapter 6:Chapter 6:
Chemical BondingChemical Bonding
1.1. Use the periodic table to infer theUse the periodic table to infer the
number of valence electrons in annumber of valence electrons in an
atom and draw its electron dotatom and draw its electron dot
structure.structure.
2.2. Be able to explain the types of bondsBe able to explain the types of bonds
that atoms can form.that atoms can form.
3.3. List the characteristics of the differentList the characteristics of the different
types of chemical bonds.types of chemical bonds.
4.4. Define the vocabulary words.Define the vocabulary words.
5.5. Use electronegativity values toUse electronegativity values to
classify a bondclassify a bond

2. 2
Valence ElectronsValence Electrons
 Electrons in the highest occupied energy levelElectrons in the highest occupied energy level
of an element’s atomsof an element’s atoms
 For representative elements, the number ofFor representative elements, the number of
valence electrons is the same as the groupvalence electrons is the same as the group
number of that element (Page 414)number of that element (Page 414)
 Shown in electron dot structuresShown in electron dot structures
3 53 5
2 12 1
6 86 8
4 74 7
Right, left, top, bottom (1,2,3,4)Right, left, top, bottom (1,2,3,4)
Then 12 o’clock and counterclockwise (5,6,7,8)Then 12 o’clock and counterclockwise (5,6,7,8)
Symbol
of the
element

3. 3
Valence Electrons (cont’d)Valence Electrons (cont’d)
 Electrons in the highest occupied energy level of anElectrons in the highest occupied energy level of an
element’s atomselement’s atoms
 Can be figured out using the group numbers in the periodicCan be figured out using the group numbers in the periodic
table.table.
 Ex: The elements of Group 1A (hydrogen, lithium,Ex: The elements of Group 1A (hydrogen, lithium,
sodium, etc.) all have a valence number ofsodium, etc.) all have a valence number of 11, which, which
means there is 1 electron in the highest occupied energymeans there is 1 electron in the highest occupied energy
level. The elements of group 7A (fluorine, chlorine,level. The elements of group 7A (fluorine, chlorine,
bromine, etc.) havebromine, etc.) have 77 electrons in the outer energy level.electrons in the outer energy level.
 The valence numbers also tell us the likelyThe valence numbers also tell us the likely oxidationoxidation
statestate of that element. More on this later.of that element. More on this later.

4. 4
Oxidation States
The oxidation state of an atom is the charge it has when it gains or loses
electrons to form it’s most stable electron configuration.
Valence Number Oxidation State
(charge on the ion)
1 +1
2 +2
3 +3
5 -3
6 -2
7 -1
}
}
CATIONS
ANIONS

5. 5
The Octet RuleThe Octet Rule
 Gilbert Lewis used this to explain why atoms formGilbert Lewis used this to explain why atoms form
certain kinds of ions and molecule.certain kinds of ions and molecule.
 In forming compounds, atoms tend to achieve theIn forming compounds, atoms tend to achieve the
electron configuration of a noble gas (8 valence eelectron configuration of a noble gas (8 valence e--
))
 Recall that each noble gas (except He) has 8 electronsRecall that each noble gas (except He) has 8 electrons
in its highest energy level and a general electronin its highest energy level and a general electron
configuration of nconfiguration of nss22
nnpp66
 Exceptions: Molecules with an odd number ofExceptions: Molecules with an odd number of
electrons, more than an octet (PClelectrons, more than an octet (PCl55), and less than), and less than
an octet (very rare)an octet (very rare)
Example: NOExample: NO22 has seventeen valence electronshas seventeen valence electrons
[Nitrogen contributes five and each oxygen[Nitrogen contributes five and each oxygen
contributes 6 (2 x 6 =12)]contributes 6 (2 x 6 =12)]

6. 6
The Octet RuleThe Octet Rule
 An atom’s loss of an electron produces aAn atom’s loss of an electron produces a cationcation, or, or
positively charged ion. The most common cations are thosepositively charged ion. The most common cations are those
produced by the loss of valence electrons from the metals,produced by the loss of valence electrons from the metals,
since most of these atoms have 1-3 valence electrons.since most of these atoms have 1-3 valence electrons.
Let’s look at sodium (a 1A metal) as an example:Let’s look at sodium (a 1A metal) as an example:
Na 1Na 1ss22
22ss22
22pp66
33ss11
NaNa++
11ss22
22ss22
22pp66
-e-

7. 7
Practice ProblemsPractice Problems
 Write the electron dot structure for each of theWrite the electron dot structure for each of the
following:following:
1. Na1. Na
2. Al2. Al
3. N3. N
4. S4. S
5. Kr5. Kr
6. Chloride ion6. Chloride ion
7. Oxide ion7. Oxide ion
 Refer to pages 414, 417, and 418 for answers.Refer to pages 414, 417, and 418 for answers.

8. 8
Practice ProblemsPractice Problems
Please write thePlease write the oxidation numbersoxidation numbers of the following:of the following:
1)1) NaNa 11) Po11) Po
2)2) AlAl 12) Ga12) Ga
3)3) FF 13) Cr13) Cr
4)4) ClCl 14) N14) N
5)5) MgMg
6)6) PP
7)7) CaCa
8)8) SbSb
9)9) II
10)10) ScSc

9. 9
Common Polyatomic IonsCommon Polyatomic Ions
Hydroxide:Hydroxide: OHOH--
Permanganate: MnOPermanganate: MnO44
--
Bicarbonate:Bicarbonate: HCOHCO33
--
Ammonium: NHAmmonium: NH44
++
Carbonate:Carbonate: COCO33
2-2-
Acetate:Acetate: CC22HH33OO22
--
Sulfate:Sulfate: SOSO44
2-2-
Hydrogen-Hydrogen-
Sulfite:Sulfite: SOSO33
2-2-
Phosphate: HPOPhosphate: HPO44
2-2-
Phosphate:Phosphate: POPO44
3-3-
Dichromate:Dichromate: CrCr22OO77
2-2-
Perchlorate:Perchlorate: ClOClO44
--
Nitrate:Nitrate: NONO33
--
Nitrite:Nitrite: NONO22
--
Chlorate:Chlorate: ClOClO33
--
Cyanide:Cyanide: CNCN--

10. 10
Chemical BondingChemical Bonding
 Chemical energy & potential energy storedChemical energy & potential energy stored
in chemical bondsin chemical bonds
 Atoms prefer a low energy conditionAtoms prefer a low energy condition
 Atoms that are bonded have less energyAtoms that are bonded have less energy
than free atoms- more stable.than free atoms- more stable.
 To combine atoms: energy is absorbedTo combine atoms: energy is absorbed
 To break a bond: energy is released (AB)To break a bond: energy is released (AB)

11. 11
Chemical BondsChemical Bonds
 Created when two nuclei simultaneouslyCreated when two nuclei simultaneously
attract electronsattract electrons
 When electrons are donated or received,When electrons are donated or received,
creating an ion (anion, cation)creating an ion (anion, cation)
 In most elements, only valence electronsIn most elements, only valence electrons
enter chemical reactionsenter chemical reactions
 Atoms of everyday substances are heldAtoms of everyday substances are held
together by chemical bonds (water, salttogether by chemical bonds (water, salt
anti-freeze)anti-freeze)

12. 12
Types of Chemical BondsTypes of Chemical Bonds
1)1) Ionic BondIonic Bond:: chemical bonding that results fromchemical bonding that results from
the electrical attraction between cations and anionsthe electrical attraction between cations and anions
where atoms completely give their electron(s) awaywhere atoms completely give their electron(s) away

13. 13
Types of Chemical BondsTypes of Chemical Bonds
2)2) Covalent BondCovalent Bond:: chemical bonding that resultschemical bonding that results
from the sharing of electron pairs between two atoms.from the sharing of electron pairs between two atoms.
The electrons are “owned” equally by the two atoms.The electrons are “owned” equally by the two atoms.

14. 14
Relative Forces ofRelative Forces of
AttractionAttraction
 Ability of a nucleus to hold its valenceAbility of a nucleus to hold its valence
electrons (Group 7A has a greater abilityelectrons (Group 7A has a greater ability
to hold on to its valence electrons thanto hold on to its valence electrons than
Group 1A)Group 1A)
 Ionization energy: energy required toIonization energy: energy required to
lose an electron (As atomic numberlose an electron (As atomic number
increases down a group, the mostincreases down a group, the most
loosely bound electrons are more easilyloosely bound electrons are more easily
removed, so ionization energyremoved, so ionization energy
decreases. For the most part, itdecreases. For the most part, it
increases along each period.)increases along each period.)

15. 15
 Electron affinity – tendency to gain anElectron affinity – tendency to gain an
electron (Energy is released)electron (Energy is released)
 Electronegativity – measure of theElectronegativity – measure of the
electron attracting power of an atomelectron attracting power of an atom
when it bonds with another atomwhen it bonds with another atom
* Fluorine (4.0) is the highest* Fluorine (4.0) is the highest
* Cesium (0.7) is the lowest – least* Cesium (0.7) is the lowest – least
ability to attract bonding electrons andability to attract bonding electrons and
thus the greatest tendency to lose anthus the greatest tendency to lose an
electronelectron
* Noble gases are not assigned* Noble gases are not assigned
electronegativities because theseelectronegativities because these
elements do not generally form bondselements do not generally form bonds
(inert)(inert)

16. 16
 The periodic trend of theThe periodic trend of the
electronegativities is the same as that ofelectronegativities is the same as that of
the ionization energies. Thus, as thethe ionization energies. Thus, as the
atomic number increases along a period,atomic number increases along a period,
the electronegativity increases. As thethe electronegativity increases. As the
atomic number increases down a group,atomic number increases down a group,
the electronegativity decreases.the electronegativity decreases.
 In general, metals have a lowIn general, metals have a low
electronegativity and nonmetals have aelectronegativity and nonmetals have a
high electronegativityhigh electronegativity

17. 17
Electronegativity and Bond TypesElectronegativity and Bond Types
Covalent BondsCovalent Bonds: bonding between elements with an electro-: bonding between elements with an electro-
negativity difference ofnegativity difference of 1.7 or less1.7 or less..
Nonpolar-Covalent BondsNonpolar-Covalent Bonds: covalent bond in which electrons are: covalent bond in which electrons are
shared evenly by the bonded atoms with an electronegativityshared evenly by the bonded atoms with an electronegativity
difference ofdifference of 0 – 0.30 – 0.3..
Polar-Covalent BondsPolar-Covalent Bonds: covalent bond in which the bonded atoms: covalent bond in which the bonded atoms
have unequal attraction of the shared electrons, and have anhave unequal attraction of the shared electrons, and have an
electronegativity difference ofelectronegativity difference of 0.4 – 1.70.4 – 1.7
Ionic BondsIonic Bonds: bonding due to difference in electric charge of two: bonding due to difference in electric charge of two
elements due to loss/gain of electrons. Must have an electro-elements due to loss/gain of electrons. Must have an electro-
negativity ofnegativity of 1.8 – 4.01.8 – 4.0..

18. 18
Electronegativity and Bond TypesElectronegativity and Bond Types
Water is aWater is a polarpolar molecule, because the electronsmolecule, because the electrons
are not shared evenly by the hydrogen and oxygen.are not shared evenly by the hydrogen and oxygen.
Ionic BondsIonic Bonds: bonds in which electrons are donated from one: bonds in which electrons are donated from one
atom to another and have an electronegativity differenceatom to another and have an electronegativity difference
ofof 1.8 or higher1.8 or higher..

19. 19
Electronegativity and Bond TypesElectronegativity and Bond Types
Using the electronegativity values found on page 161 ofUsing the electronegativity values found on page 161 of
your book, predict the types of bonds the following will form.your book, predict the types of bonds the following will form.
1) O1) O22
2) NaCl2) NaCl
3) N3) N22
4) Knowing that the electronegativity of sulfur is 2.5, what type4) Knowing that the electronegativity of sulfur is 2.5, what type
of bond will sulfur form with:of bond will sulfur form with:
a) hydrogena) hydrogen
b) cesiumb) cesium
c) chlorinec) chlorine

20. 20
ElectronegativityElectronegativity
Electronegativity is a measure of how strongly an elementElectronegativity is a measure of how strongly an element
can remove an electron from another element.can remove an electron from another element.

21. 21
Ionic (Electrovalent)Ionic (Electrovalent)
BondsBonds The strongest chemical bondThe strongest chemical bond
Complete transfer of electron(s) from one element to anotherComplete transfer of electron(s) from one element to another
 Generally formed when metals combine withGenerally formed when metals combine with
nonmetals (Groups 1-2a w/ 5-7a)nonmetals (Groups 1-2a w/ 5-7a)
 Coulombic forces – electrostatic force in which twoCoulombic forces – electrostatic force in which two
oppositely charged ions are mutually attractedoppositely charged ions are mutually attracted
 Usually occurs when the difference inUsually occurs when the difference in
electronegativities is 1.8 or greaterelectronegativities is 1.8 or greater

22. 22
NaCl – Ionic BondNaCl – Ionic Bond
 DrawDraw

23. 23
Writing IonicWriting Ionic
CompoundsCompounds
 Beryllium fluorideBeryllium fluoride
 Calcium oxideCalcium oxide
 Scandium sulfideScandium sulfide
 Aluminum chlorideAluminum chloride

24. 24
Ionic SolidsIonic Solids
 Form crystal lattice (orderly, repeating,Form crystal lattice (orderly, repeating,
three-dimensional pattern)three-dimensional pattern)
 The charges and relative sizes of theThe charges and relative sizes of the
ions determines the crystal structureions determines the crystal structure
 The number of ions of opposite chargeThe number of ions of opposite charge
that surround the ion in a crystal is calledthat surround the ion in a crystal is called
thethe coordination numbercoordination number of the ion.of the ion.

25. 25
 Poor conductors of electricity (no freePoor conductors of electricity (no free
electrons)electrons)
 High melting pointHigh melting point
 High boiling pointHigh boiling point
 Brittle and break easily under stressBrittle and break easily under stress
 Liquid or aqueous: good conductors ofLiquid or aqueous: good conductors of
electricity but ionic bond is dissolvedelectricity but ionic bond is dissolved

26. 26
The NormalThe Normal
Arrangement of anArrangement of an
Ionic CrystalIonic Crystal
Opposite charges attractOpposite charges attract
-
-
-
-
-
-
-
- -
-
+
+
+
+
+
+
+
+
+
+

27. 27
Arrangement when StressArrangement when Stress
is Appliedis Applied
Adjacent to ions with same charge (repulsion)Adjacent to ions with same charge (repulsion)
+
+
+
-
-
-
-
+
-
+
-
-
+
+
+
+
-
-
+
-

28. 28
Crystal Lattice isCrystal Lattice is
DestroyedDestroyed
Crystal melts, vaporizes, or dissolves in waterCrystal melts, vaporizes, or dissolves in water
(ions free to move about)(ions free to move about)
Cleavage – splitting along a definite lineCleavage – splitting along a definite line
+
+
+
-
-
-
-
+
-
+
-
-
+
+
+
+
-
-
+
-

29. 29
 Electrons are sharedElectrons are shared
 One atom does not have enough pull on theOne atom does not have enough pull on the
electron to take it completely from the otherelectron to take it completely from the other
atomatom
 Occurs when electronegativity difference isOccurs when electronegativity difference is
less than 1.8less than 1.8
 Covalently Bonded Solids:Covalently Bonded Solids:
1. Softness1. Softness
2. Poor conductor of electricity and heat2. Poor conductor of electricity and heat
3. Low melting point3. Low melting point
Covalent BondingCovalent Bonding

30. 30
Lewis StructuresLewis Structures
 Single covalent bondSingle covalent bond – one shared pair– one shared pair
of electrons:of electrons:
 HH· + ·H· + ·H H H or H HH H or H H
 Double covalent bondDouble covalent bond – two shared pairs– two shared pairs
of electronsof electrons
O + OO + O O O or O OO O or O O
 Triple covalent bondTriple covalent bond – three shared pairs– three shared pairs
of electronsof electrons
N + NN + N N N or N NN N or N N
NoteNote: all of these obey the: all of these obey the octetoctet rulerule
:
: :
:
:.
.
. .
:
:
: : : : : :
:
:
:
:
.. . .
..
: : : :: : :

31. 31
 Coordinate covalent bondCoordinate covalent bond – one atom– one atom
contributes both bonding electronscontributes both bonding electrons
 NHNH33 + H+ H++
[NH[NH44]]++
ammonia hydrogen ionammonia hydrogen ion ammonium ionammonium ion
HH ++
H N H + HH N H + H++
H N HH N H
HH HH
The structural formula shows anThe structural formula shows an arrowarrow
that points from the atom donating thethat points from the atom donating the
electrons to the atom receiving them.electrons to the atom receiving them.
Refer to page 444Refer to page 444
: :
::
:
: :

32. 32
How to Construct Lewis StructuresHow to Construct Lewis Structures
Step 1Step 1:: Determine the type and # of atoms in moleculeDetermine the type and # of atoms in molecule
CHCH33II
has 1 Carbon, 3 Hydrogens and 1 Iodinehas 1 Carbon, 3 Hydrogens and 1 Iodine
Step 2Step 2:: Write electron dot notation for each type of atomWrite electron dot notation for each type of atom
CC HH· I· I
Step 3Step 3: Determine the total # of electrons available in: Determine the total # of electrons available in
the atoms to be combined.the atoms to be combined.
CC 1 x 4e1 x 4e--
= 4e= 4e--
II 1 x 7e1 x 7e--
= 7e= 7e--
HH 3 x 1e3 x 1e--
= 3e= 3e--
14 e14 e--
.
.
..
..
.....

33. 33
How to Construct Lewis StructuresHow to Construct Lewis Structures
Step 4Step 4:: Arrange the atoms to form a skeleton structureArrange the atoms to form a skeleton structure
for the molecule. Then connect the atoms byfor the molecule. Then connect the atoms by
electron-pair bonds.electron-pair bonds.
H C IH C I
Step 5Step 5: Add unshared pairs of electrons to each non-: Add unshared pairs of electrons to each non-
metal atom so that each is surrounded by 8.metal atom so that each is surrounded by 8.
HH C IC I
......
..
H
H......
.
H
H
....
..
.

34. 34
Electron Dot Practice: CompoundsElectron Dot Practice: Compounds
1)1) HH22OO 5) CCl5) CCl22HH22
2)2) HH22OO22 6) NH6) NH33
3)3) HCNHCN 7) N7) N22
4)4) AlFAlF33 8) CO8) CO22

35. 35
 A single water molecule is a good example of covalent bondingA single water molecule is a good example of covalent bonding
between atoms. The hydrogen atoms “share” their electrons withbetween atoms. The hydrogen atoms “share” their electrons with
the larger oxygen atom so that oxygen now has a full outer levelthe larger oxygen atom so that oxygen now has a full outer level
with 8 electrons and each hydrogen has a full outer level with 2with 8 electrons and each hydrogen has a full outer level with 2
electrons. Oxygen has a higher electronegativity than hydrogen,electrons. Oxygen has a higher electronegativity than hydrogen,
so there is actually an uneven sharing of electrons, resulting in aso there is actually an uneven sharing of electrons, resulting in a
polar molecule. More on this later.polar molecule. More on this later.
8p+
8n0
e-
e-
e-
e-
e-
e-
e-
e-
e-
e-
1p+
1p+
shared electrons shared electrons

36. 36
 Bond dissociation energyBond dissociation energy: total energy required to: total energy required to
break the bond between two covalently bonded atomsbreak the bond between two covalently bonded atoms
(remember that energy is measured in(remember that energy is measured in joulesjoules oror
kilojouleskilojoules))
HH–H + 435 kJ–H + 435 kJ H + HH + H
 Resonance StructuresResonance Structures: refers to bonding in molecules: refers to bonding in molecules
or ions that cannot be correctly represented by aor ions that cannot be correctly represented by a
single Lewis structure.single Lewis structure.
. .

37. 37
Bond Length vs. BondBond Length vs. Bond
EnergyEnergyThere is a correlation between bondThere is a correlation between bond lengthlength and the amountand the amount
of potential energy stored in that bond. For example:of potential energy stored in that bond. For example:
Bond Bond Length (pm) Bond energy (Kj/mol)Bond Bond Length (pm) Bond energy (Kj/mol)
C CC C 154154 346346
CC CC 134134 612612
CC CC 120120 835835
CC NN 147147 305305
CC NN 132132 615615
CC NN 116116 887887
NN NN 145145 163163
NN NN 125125 418418
NN NN 110110 945945

38. 38
 Molecular orbitals – when two atoms combine and theirMolecular orbitals – when two atoms combine and their
atomic orbitals overlapatomic orbitals overlap
 Sigma bond - molecular orbital that is symmetricalSigma bond - molecular orbital that is symmetrical
along the axis connecting two atomic nucleialong the axis connecting two atomic nuclei
In both of these examples, theIn both of these examples, the pp orbitals are overlappingorbitals are overlapping
and sharing electrons.and sharing electrons.

39. 39
 pi bond – weaker than sigma bond; usuallypi bond – weaker than sigma bond; usually
sausage-shaped regions above and below thesausage-shaped regions above and below the
bond axis (Page 445)bond axis (Page 445)

40. 40
Examples of Sigma and Pi bondsExamples of Sigma and Pi bonds
H3C – CH3
H2C = CH2
H3C – CH3
HC CH––
–

41. 41
VSEPR Theory (page 200)VSEPR Theory (page 200)
VSEPR TheoryVSEPR Theory ((VValencealence SShellhell EElectron-lectron-PPairair RRepulsion theory):epulsion theory):
states that repulsion between the sets of valence-level electronsstates that repulsion between the sets of valence-level electrons
surrounding an atom causes these sets to be oriented as farsurrounding an atom causes these sets to be oriented as far
away from each other as possible, thus determining the shapeaway from each other as possible, thus determining the shape
of molecules.of molecules.

42. 42
VSEPR TheoryVSEPR Theory
So then why is HSo then why is H22O bent, but BeFO bent, but BeF22 is linear?is linear?
The answer is theThe answer is the free electron pairsfree electron pairs. Oxygen has 2 pairs,. Oxygen has 2 pairs,
beryllium has none.beryllium has none.
HH22OO BeFBeF22
Be. .O:
:
.
.
1
2

43. 43
VSEPR ShapesVSEPR Shapes
LinearLinear
Trigonal-PlanerTrigonal-Planer
Bent/AngularBent/Angular
TetrahedralTetrahedral
Trigonal-PyramidalTrigonal-Pyramidal
Trigonal-BipyramidalTrigonal-Bipyramidal
OctahedralOctahedral
(#s 3, 5 and 7 are(#s 3, 5 and 7 are coordinate covalent bonds!)coordinate covalent bonds!)

44. 44
VSEPR TheoryVSEPR Theory
These free electron pairsThese free electron pairs repelrepel each other because they have aeach other because they have a
negative charge, and so they force those atoms that arenegative charge, and so they force those atoms that are
covalently bonded to be pushed as far away as possible.covalently bonded to be pushed as far away as possible.
.
. .
.

45. 45
 Hybridization – several atomic orbitalsHybridization – several atomic orbitals
mix to form the same total number ofmix to form the same total number of
equivalent hybrid orbitals (CHequivalent hybrid orbitals (CH44 – Page– Page
457457))
*Note: An sp*Note: An sp33
orbital is an example of aorbital is an example of a
hybrid.hybrid.

46. 46
Types of CovalentTypes of Covalent
BondingBonding
 1. Nonpolar – when atoms have the same or1. Nonpolar – when atoms have the same or
similar electronegativity; when the atoms insimilar electronegativity; when the atoms in
the bond pull equally and the bondingthe bond pull equally and the bonding
electrons are shared equally.electrons are shared equally.
(Generally a difference of 0.0- 0.4)(Generally a difference of 0.0- 0.4)
* Examples: Diatomic elements* Examples: Diatomic elements
(H(H22 , N, N22 , O, O22 , F, F22 , Cl, Cl22 , I, I22 , Br, Br22))
* Nonpolar Covalent:* Nonpolar Covalent:
Bonded Hydrogen AtomsBonded Hydrogen Atoms

47. 47
 2. Polar – unequal sharing of electrons2. Polar – unequal sharing of electrons
* Pairing of atoms when one has a stronger* Pairing of atoms when one has a stronger
attraction for the electronsattraction for the electrons
* Most compounds are polar covalent* Most compounds are polar covalent
Examples: HExamples: H22O , NHO , NH33 , HF , HCl, HF , HCl
* Polar covalent also called dipoles* Polar covalent also called dipoles
* Creates partial charges* Creates partial charges
Partially +Partially +
Partially -Partially -
Example: HCl (0.9 difference of the electro-Example: HCl (0.9 difference of the electro-
negativities) H (2.1) Cl (3.0)negativities) H (2.1) Cl (3.0)
Electronegativity difference is less than 1.7Electronegativity difference is less than 1.7

48. 48
Example:Example: waterwater..
An uneven distribution of the electrons results becauseAn uneven distribution of the electrons results because
the oxygen has a higher electron affinity than thethe oxygen has a higher electron affinity than the
hydrogens. Thus, you have a negative and positive endhydrogens. Thus, you have a negative and positive end
of the molecule: polarity. Because this molecule hasof the molecule: polarity. Because this molecule has 22
poles, it is called apoles, it is called a dipoledipole molecule.molecule.
8p+
8n0
e-
e-
e-
e-
e-
e-
e-
e-
e-
e-
1p+
1p+
Oxygen
Hydrogens
δδ --
δδ ++
δδ = delta or= delta or
overalloverall

49. 49
Attractions Between MoleculesAttractions Between Molecules
Molecules are often attracted to each other by a variety ofMolecules are often attracted to each other by a variety of
forces. Theforces. The intermolecularintermolecular attractions are weaker than eitherattractions are weaker than either
an ionic or covalent bond. These attractions are responsiblean ionic or covalent bond. These attractions are responsible
for determining whether a molecular compound is a gas, liquid,for determining whether a molecular compound is a gas, liquid,
or solid at a given temperature. Here is a list of these variousor solid at a given temperature. Here is a list of these various
attractions:attractions:
1.1. van der Waals forcesvan der Waals forces: weakest type of intermolecular: weakest type of intermolecular
attractions.attractions.
Dispersion (London) forcesDispersion (London) forces: weakest of all molecular: weakest of all molecular
interactions, caused by the motion of electrons.interactions, caused by the motion of electrons.
Increases as the # of electrons increases.Increases as the # of electrons increases.
Ex: Cl & F are gases at STP; Br is liquid at STP;Ex: Cl & F are gases at STP; Br is liquid at STP;
I is solid at STP.I is solid at STP.

50. 50
2.2. Dipole InteractionsDipole Interactions: attraction of polar molecules to one: attraction of polar molecules to one
another. Remember that polar molecules are like magnets;another. Remember that polar molecules are like magnets;
they have a positive and negative end.they have a positive and negative end.
A glucose moleculeA glucose molecule
in water has many dipolein water has many dipole
interactions since bothinteractions since both
water and glucose arewater and glucose are
polar. The positive polespolar. The positive poles
of the water molecule areof the water molecule are
attracted to the negativeattracted to the negative
poles on the glucose andpoles on the glucose and
vice versa.vice versa.

51. 51
3.3. Hydrogen BondsHydrogen Bonds: attractive forces in which a hydrogen: attractive forces in which a hydrogen
covalently bonded to a very electronegative atom is alsocovalently bonded to a very electronegative atom is also
weakly bonded to an unshaired pair of electrons on anotherweakly bonded to an unshaired pair of electrons on another
electronegative. Hydrogen bonding always involveselectronegative. Hydrogen bonding always involves
hydrogen. Hence the name. Duh.hydrogen. Hence the name. Duh.
The hydrogen bonding betweenThe hydrogen bonding between
water molecules dictates many ofwater molecules dictates many of
the properties of water. It alsothe properties of water. It also
explains why water is a liquidexplains why water is a liquid
rather than a gas at roomrather than a gas at room
temperature.temperature.

52. 52
Network SolidsNetwork Solids
 MacromoleculesMacromolecules
 Covalent network of atoms bondedCovalent network of atoms bonded
 Absence of molecules throughout theAbsence of molecules throughout the
solidsolid
 PropertiesProperties
1. Hardness1. Hardness
2. Poor conductor of electricity (electrical2. Poor conductor of electricity (electrical
insulation)insulation)
3. Poor conductor of heat3. Poor conductor of heat

53. 53
 Examples: diamond graphite (carbon)Examples: diamond graphite (carbon)
Does not melt - vaporizes to a gas at 3500Does not melt - vaporizes to a gas at 3500 °C°C
Carbon atomCarbon atom
Covalent bondCovalent bond
 Boron nitride (BN), asbestos, silicone carbideBoron nitride (BN), asbestos, silicone carbide
(SiC, grindstones), silicone dioxide (SiO(SiC, grindstones), silicone dioxide (SiO22 ,,
quartz)quartz)

54. 54
Metallic BondsMetallic Bonds
 Most metallic elements, except liquidMost metallic elements, except liquid
mercury, are solids at room temperaturemercury, are solids at room temperature
and exhibit a crystal structure (zinc)and exhibit a crystal structure (zinc)
 Arrangement of stationary positive metalArrangement of stationary positive metal
ions surrounded by a “sea of mobileions surrounded by a “sea of mobile
electrons””electrons””
- - - -- - - -
- - - -- - - -
- - - -- - - -
- - - -- - - -
+
+
+ +
+
+
+
+
+

55. 55
 Properties:Properties:
1. Malleability – ability to be hammered1. Malleability – ability to be hammered
into different shapesinto different shapes
2. Ductility – ability to be drawn into wire2. Ductility – ability to be drawn into wire
3. Conductor of heat3. Conductor of heat
4. Conductor of electricity4. Conductor of electricity
5. Luster – shine5. Luster – shine
6. Tenacity – structural strength6. Tenacity – structural strength
(resistance to being pulled apart)(resistance to being pulled apart)

56. 56
AlloysAlloys
 Mixtures composed of two or more elements, at leastMixtures composed of two or more elements, at least
one of which is a metalone of which is a metal
 Properties usually superior to those of the componentProperties usually superior to those of the component
elementselements
 Sterling silver – silver and copperSterling silver – silver and copper
 Bronze – copper and tinBronze – copper and tin
 Steel – iron, carbon, boron, chromium, manganese,Steel – iron, carbon, boron, chromium, manganese,
molybdenum, nickel, tungsten, vanadium (Interstitialmolybdenum, nickel, tungsten, vanadium (Interstitial
alloy)alloy)
 Interstitial alloy – smaller atoms fit into spaces betweenInterstitial alloy – smaller atoms fit into spaces between
larger atomslarger atoms
 Substitutional alloy – atoms of the components areSubstitutional alloy – atoms of the components are
about the same size (They can replace each other inabout the same size (They can replace each other in
the structure.)the structure.)

57. 57
Summary : Types ofSummary : Types of
BondsBonds
1. Ionic – complete transfer of electrons1. Ionic – complete transfer of electrons
2.2. Covalent – share electronsCovalent – share electrons
A. Nonpolar : same or similar electronegativityA. Nonpolar : same or similar electronegativity
B. Polar – unequal sharingB. Polar – unequal sharing
Electronegativity Difference:Electronegativity Difference:
CC << 2.02.0 ≤≤ II (Know exceptions)(Know exceptions)
*Know table on page 465*Know table on page 465
3.3. Network solids – covalent network of atomsNetwork solids – covalent network of atoms
(absence of molecules)(absence of molecules)
4.4. Metallic – positive ions around a “sea ofMetallic – positive ions around a “sea of
mobile electrons”mobile electrons”

58. 58
ElectronegativityElectronegativity
Differences and BondDifferences and Bond
TypesTypes
 0.0-0.3 Nonpolar covalent0.0-0.3 Nonpolar covalent
 0.4-1.0 Moderately polar covalent0.4-1.0 Moderately polar covalent
 1.0-1.8 Very polar covalent1.0-1.8 Very polar covalent
 1.8 or greater Ionic1.8 or greater Ionic

59. 59
General Trends of theGeneral Trends of the
Representative ElementsRepresentative Elements
 Group 1A - lose one electronGroup 1A - lose one electron
 Group 2A - lose two electronsGroup 2A - lose two electrons
 Group 3A - lose three electronsGroup 3A - lose three electrons
 Group 4A - share, lose or gain 4 e-Group 4A - share, lose or gain 4 e-
 Group 5A - share, gain three electronsGroup 5A - share, gain three electrons
 Group 6A - share, gain two electronsGroup 6A - share, gain two electrons
 Group 7A - gain one electronGroup 7A - gain one electron
 Group 8 - do not react, noble gasesGroup 8 - do not react, noble gases

60. 60
Think!Think!
 Why is it possible to bend metals but notWhy is it possible to bend metals but not
ionic crystals?ionic crystals?
 In an ionic compound, ions of like chargeIn an ionic compound, ions of like charge
do not have mobile electrons asdo not have mobile electrons as
insulation. When forced into contact byinsulation. When forced into contact by
physical stress, the ions of like chargephysical stress, the ions of like charge
repel, causing the crystal to shatter.repel, causing the crystal to shatter.

61. 6161
AttractionsAttractions
betweenbetween
MoleculesMoleculesName and describe the weakName and describe the weak
attractive forces that hold groups ofattractive forces that hold groups of
molecules together.molecules together.

62. 62
Van der Waals ForcesVan der Waals Forces
 Weaker than either an ionic or covalentWeaker than either an ionic or covalent
bondbond
 Responsible for determining whether aResponsible for determining whether a
molecular compound is a gas, liquid, ormolecular compound is a gas, liquid, or
solid at a given temperaturesolid at a given temperature
 Two types: dispersion forces and dipoleTwo types: dispersion forces and dipole
interactionsinteractions
 Dispersion – caused by motion ofDispersion – caused by motion of
electrons; dispersion generally increaseselectrons; dispersion generally increases
as the number of electrons increasesas the number of electrons increases

63. 63
 Example of dispersion forces: (Refer toExample of dispersion forces: (Refer to
Group 7A) F and Cl are gases at STP; BrGroup 7A) F and Cl are gases at STP; Br
is a liquid at STP, and I is a solid at STPis a liquid at STP, and I is a solid at STP
 Dipole interactions – electrostaticDipole interactions – electrostatic
attractions between oppositely chargedattractions between oppositely charged
regions (Example: water)regions (Example: water)

64. 64
Hydrogen BondsHydrogen Bonds
 Strongest of the intermolecular forcesStrongest of the intermolecular forces
 Important in determining the properties ofImportant in determining the properties of
water and biological molecules such aswater and biological molecules such as
proteinsproteins
 Has only about 5% of the strength of anHas only about 5% of the strength of an
average covalent bondaverage covalent bond