1. CHEMICAL
BOND
( PRESENTATION )

4. BOND FORMATION
 exothermic process
E
N Reactants
E
R Energy
G released
Y
Products

5. BREAKING BONDS
 Endothermic reaction
 energy must be put into the bond in order to break it
E
N Products
E
R Energy
G Absorbed
Y Reactants

6. BOND STRENGTH
 Strong, STABLE bonds require lots of energy to be
formed or broken
 weak bonds require little E

7. TWO MAJOR TYPES OF BONDING
Ionic Bonding
 forms ionic compounds
 transfer of e-
Covalent Bonding
 forms molecules
 sharing e-

8. ONE MINOR TYPE OF BONDING
 Metallic bonding
Occurs between like atoms of a metal
in the free state
Valence e- are mobile (move freely
among all metal atoms)
Positive ions in a sea of electrons
 Metallic characteristics
 High mp temps, ductile, malleable, shiny
 Hard substances
 Good conductors of heat and electricity as
(s) and (l)

9. It’s the mobile
electrons that enable
me tals to conduct
-
electricity!!!!!!

10. IONIC BONDING
 electrons are transferred between valence shells of
atoms
 ionic compounds are
made of ions
NOT MOLECULES
• ionic compounds are called Salts or
Crystals

11. IONIC BONDING
 Always formed between metals and non-metals
+ -
[METALS ] [NON-METALS ]
Lost e- Gained e-

12. IONIC BONDING
 Electronegativity difference > 2.0
 Look up e-neg of the atoms in the bond and subtract
NaCl
CaCl2
 Compounds with polyatomic ions
NaNO3

14. PROPERTIES OF IONIC COMPOUNDS
SALTS
 hard solid @ 22 C
o
Crystals
 high mp temperatures
 nonconductors of electricity in solid phase
 good conductors in liquid phase or dissolved in
water (aq)

15. COVALENT BONDING
Pairs of e- are shared
molecules
between non-metal
atoms
electronegativity difference <
2.0
forms polyatomic ions

16. MOLECULAR
SUBSTANCES
Covalent
bonding
Low m.p. temp and b.p.
temps
relatively soft solids as
compared to ionic compounds
nonconductors of electricity
in any phase

17. COVALENT, IONIC, METALLIC
BONDING?
NO
2
NH + • CO
sodium 4
Aluminu
hydride • Co
Hg m
phosphate
H S
2 KH
sulfate Also study
KCl your
HF characteristics!

18. DRAWING IONIC COMPOUNDS
USING LEWIS DOT STRUCTURES
• Symbol represents the KERNEL of the
atom (nucleus and inner e-)
• dots represent valence e-

19. NACL
 This is the finished Lewis Dot Structure
-
[Na] [ Cl ]
How did we get here?
+

20.  Step 1 after checking that it is IONIC
Determine which atom will be the +ion
Determine which atom will be the - ion
 Step 2
Write the symbol for the + ion first.
 NO DOTS
Draw the e- dot diagram for the – ion
 COMPLETE outer shell
 Step 3
Enclose both in brackets and show each
charge

21. DRAW THE LEWIS DIAGRAMS
LiF
MgO
CaCl
2
K S
2

22. DRAWING MOLECULES USING
LEWIS DOT STRUCTURES
 Symbol represents the KERNEL of the atom
(nucleus and inner e-)
 dots represent valence e-

23. Always remember atoms
are trying to complete
their outer shell!
The number of electrons the atoms needs is the
total number of bonds they can make.
Ex. … H? O? F? N? Cl? C?
one two one three one four

24. METHANE CH4
 This is the finished Lewis dot structure
How did we get here?

25.  Step 1
count total valence e- involved
 Step 2
connect the central atom (usually the
first in the formula) to the others
with single bonds
 Step 3
complete valence shells of outer
atoms
 Step 4
add any extra e- to central atom
IF the central atom has 8 valence e-

26. SOMETIMES . . .
 You only have two atoms, so there is no central
atom, but follow the same rules.
 Check & Share to make sure all the atoms are
“happy”.
Cl2 Br2 H2 O2 N2 HCl

27.  DOUBLE bond
atoms that share two e- pairs (4 e-)
O O
 TRIPLE bond
atoms that share three e- pairs (6 e-)
N N

28. DRAW LEWIS DOT STRUCTURES
You may represent valence electrons from different
atoms with the following symbols x, ,
CO2
NH3

29. DIAGRAM FOR
POLYATOMIC IONS
 Count all valence e- needed for covalent bonding
 Add or subtract other electrons based on the
charge
REMEMBER!
A positive charge means it
LOST electrons!!!!!

30. DRAW POLYATOMICS
 Ammonium
 Sulfate

31. TYPES OF COVALENT
BONDS
NON-Polar bonds
Electrons shared evenly in the
bond
E-neg difference is zero
Between identical atoms
Diatomic molecules

32. TYPES OF COVALENT
BONDS
Polar bond
Electrons unevenly shared
E-neg difference greater than
zero but
less than 2.0
closer to 2.0 more polar
more “ionic character”

33. POLARITY
WHICH IS LEAST AND WHICH IS
MOST?
HCl
CH
4
CO a.k.a.
2
“ionic character”
NH
3
N
2
HF

34. NON-POLAR MOLECULES
Sometimes the bonds within
a molecule are polar and yet
the molecule is non-polar
because its shape is
symmetrical. H
H C H
Draw Lewis dot first and
see if equal on all sides
H

35. POLAR MOLECULES (A.K.A.
DIPOLES)
Not equal on all sides
Polarbond between 2 atoms
makes a polar molecule
asymmetrical shape of molecule

36. δ+
H Cl δ -

37. SPACE FILLING MODEL
“ELECTRON-CLOUD” MODEL
δ+ H Cl -
δ

38. WATER IS ASYMMETRICAL
δ+ δ+
H H
O
-
δ

39. WATER IS A BENT MOLECULE
H H H
H
O

40. W-A-T-E-R
as bent as it can be!
Water’s polar MOLECULE!
Water’s polar MOLECULE!
The H is positive
The O is not - not - not - not

41. MAKING SENSE OF THE POLAR
NON-POLAR THING
BONDS MOLECULES
Non-polar Non-polar Polar
Polar Symmetrical Asymmetrical
Identical
Different

42. IONIC BONDS ….
Ionic bonds are so
polar that the electrons arenot shared
but transferred between atoms forming ions!!!!!!

43. 4 SHAPES OF
MOLECULES

44. LINEAR (STRAIGHT LINE)
Ball and stick
model
Space filling
model

45. BENT
Ball and stick
model
Space filling
model

46. TRIGONAL PYRAMID
Ball and stick Space filling
model model

47. TETRAHEDRAL
Ball and stick Space filling
model model

48. INTERMOLECULAR ATTRACTIONS
 Attractions between
molecules
 van der Waals
forces
 Weak attractive
forces between
non-polar
molecules
 Hydrogen
“bonding”
 Strong
attraction
between special

49. VAN DER WAALS
 Non-polar molecules can exist in liquid and solid
phases
because van der Waals forces keep the molecules
attracted to each other
 Exist between CO2, CH4, CCl4, CF4, diatomics and
monoatomics

50. VAN DER WAALS PERIODICITY
 increase with molecular mass.
 Greater van der Waals force?
 F2 Cl2 Br2 I2
 increase
with closer distance between
molecules
Decreases when particles are farther away

51. HYDROGEN “BONDING”
 Strong polar
attraction
Like magnets
 Occurs ONLY
between H of one
molecule and N,
O, F of another
H “bond”

52. WHY DOES H “BONDING” OCCUR?
 Nitrogen, Oxygen and Fluorine
 small atoms with strong nuclear charges
 powerful atoms
 very high electronegativities

53. INTERMOLECULAR FORCES
DICTATE CHEMICAL PROPERTIES
 Strong intermolecular forces cause high b.p., m.p.
and slow evaporation (low vapor pressure) of a
substance.

54. WHICH SUBSTANCE HAS THE
HIGHEST BOILING POINT?
 HF
 NH Fluorine has the highest e-neg,
3
 H 2O SO
HF will experience the
 WHY? strongest H bonding and ∴
needs the most energy to
weaken the i.m.f. and boil

55. THE UNUSUAL PROPERTIES OF
WATER
 Unusually
high
boiling
point
 Compared
to other
compounds
in Group 16

56. DENSITY????

57. H2O(S) IS LESS DENSE THAN H2O(L)
 The hydrogen bonding in water(l)
molecules is random. The molecules
are closely packed.
 Thehydrogen bonding in water(s)
molecules has a specific open lattice
pattern. The molecules are farther
apart.

58. THE END