1. Chemical Bonds

2. Chemical BONDING

3. electron shells Atomic number = number of Electrons
Electrons vary in the amount of energy they possess, and they occur at certain energy levels or electron shells.
Electron shells determine how an atom behaves when it encounters other atoms

4. Electrons are placed in shells according to rules:
The 1st shell can hold up to two electrons, and each shell thereafter can hold up to 8 electrons.

5. Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons
C would like to
N would like to
O would like to
Gain 4 electrons
Gain 3 electrons
Gain 2 electrons

6. Why are electrons important?
Elements have different electron configurations
different electron configurations mean different levels of bonding

7. Electron Dot Structures
Symbols of atoms with dots to represent the valence-shell electrons
H He:
          
Li Be  B   C   N   O  : F  :Ne :
       
          
Na Mg  Al  Si  P S :Cl  :Ar :
       

8. Learning Check A. X would be the electron dot formula for
A. X would be the electron dot formula for
1) Na 2) K 3) Al
 
B  X  would be the electron dot formula
1) B 2) N 3) P

9. Chemical Bond
A bond results from the attraction of nuclei for electrons
All atoms trying to achieve a stable octet
IN OTHER WORDS
the p+ in one nucleus are attracted to the e- of another atom
Electronegativity

10. Bond Formation exothermic process ENERGY Reactants Energy released
Products

11. Breaking Bonds Endothermic reaction
energy must be put into the bond in order to break it
ENERGY
Products
Energy Absorbed
Reactants

12. Bond Strength
Strong, STABLE bonds require lots of energy to be formed or broken
weak bonds require little E

13. Is a bond forming or breaking? Strong or weak bond?
Reaction Time
Energy
(KJ)
energy absorbed
endothermic
bond breaking
weak unstable bond
Products
Reactants

14. Is a bond forming or breaking? Strong or weak bond?
Reaction Time
Energy
Energy absorbed
endothermic
bond breaking
strong stable bond
Products
Reactants

15. Is a bond forming or breaking? Strong or weak bond?
Energy
(KJ)
Reaction Time
Energy released
exothermic
bond formation
weak unstable bond
Reactants
Products

16. Is a bond forming or breaking? Strong or weak bond?
Energy
(KJ)
Reaction Time
Reactants
Products
Energy released
exothermic
bond formation
strong stable bond

17. Two Major Types of Bonding
Ionic Bonding
forms ionic compounds
transfer of e-
Covalent Bonding
forms molecules
sharing e-

18. 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)

19. how the electrons do not just stay with one ion
Metallic Bond
A force of attraction between a positively charged metal ion and the electrons in a metal
Many metal ions pass along many electrons
Many properties of metals, such as conductivity, ductility, and malleability, result from the freely moving electrons in the metal
Usually occurs between atoms of metals
Al3+ - - -
Notice
how the electrons do not just stay with one ion - - - - - - - - - - -

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

21. 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

22. [METALS ]+ [NON-METALS ]-
IONic bonding
Always formed between metals and non-metals
[METALS ]+ [NON-METALS ]-
Lost e-
Gained e-

23. 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

25. Ionic Bonds: One Big Greedy Thief Dog!

27. Anion (-) Cation (+)

28. Properties of Ionic Compounds
SALTS
Crystals
hard 22oC
high mp temperatures
nonconductors of electricity in solid phase
good conductors in liquid phase or dissolved in water (aq)

29. Covalent Bonding Pairs of e- are shared between non-metal atoms
molecules
Pairs of e- are shared between non-metal atoms
electronegativity difference < 2.0
forms polyatomic ions

30. Properties of 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

31. when electrons are shared equally
NONPOLAR COVALENT BONDS
when electrons are shared equally
H2 or Cl2

32. 2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons.
Oxygen Atom
Oxygen Atom
Oxygen Molecule (O2)

33. when electrons are shared but shared unequally
POLAR COVALENT BONDS
when electrons are shared but shared unequally
H2O

34. Polar Covalent Bonds: Unevenly matched, but willing to share.

35. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

36. Covalent, Ionic, metallic bonding?
NO2
sodium hydride Hg
H2S
sulfate
NH4+
Aluminum phosphate KH
KCl HF CO Co
Also study
your
characteristics!

37. Drawing ionic compounds using Lewis Dot Structures
Symbol represents the KERNEL of the atom (nucleus and inner e-)
dots represent valence e-

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

39. 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 lewis dot structure for each atom.
- Draw an arrow from the electrons that move to where they move.
Step 3
Write the symbol for the + ion first.
NO DOTS
Draw the e- dot diagram for the – ion
COMPLETE outer shell
Step 4
Enclose both in brackets and show each charge

40. What did the atom of fluorine
say to the atom of
sodium?
You complete me.

41. Draw the Lewis Diagrams
LiF
MgO
CaCl2
K2S

42. Drawing molecules using Lewis Dot Structures
Symbol represents the KERNEL of the atom (nucleus and inner e-)
dots represent valence e-

43. 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

44. Methane CH4 This is the finished Lewis dot structure
How did we get here?

45. 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- surrounding it . . YOU’RE DONE!

46. 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”.
Cl Br H O N HCl

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

48. Draw Lewis Dot Structures
You may represent valence electrons from different atoms with the following symbols x, ,
CO2
NH3

49. Draw the Lewis Dot 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!!!!!

50. Draw Polyatomics
Ammonium
Sulfate

51. Types of Covalent Bonds
NON-Polar bonds
Electrons shared evenly in the bond
E-neg difference is zero
Between identical atoms
Diatomic molecules

52. Types of Covalent Bonds
Polar bond
Electrons unevenly shared
E-neg difference greater than zero but
less than 2.0
closer to more polar
more “ionic character”

53. Place these molecules in order of increasing bond polarity which is least and which is most?
HCl
CH4
CO2
NH3 N2 HF
a.k.a.
“ionic character”

54. Polar molecules (a.k.a. Dipoles)
Not equal on all sides
Polar bond between 2 atoms makes a polar molecule
asymmetrical shape of molecule

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

56. H Cl + -

57. Space filling model “Electron-Cloud” modelH + -

58. Water is asymmetrical + + H O -

59. Water is a bent moleculeH H

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

61. Making sense of the polar non-polar thing
BONDS Non-polar Polar Identical Different
MOLECULES
Non-polar Polar
Symmetrical Asymmetrical

62. IONIC bonds ….
Ionic bonds are so polar that the electrons are not shared but transferred between atoms forming ions!!!!!!

63. Intermolecular attractions
Attractions between molecules
van der Waals forces
Weak attractive forces between non-polar molecules
Hydrogen “bonding”
Strong attraction between special polar molecules

64. 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

65. 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

66. Hydrogen “Bonding” Strong polar attraction
Like magnets
Occurs ONLY between H of one molecule and N, O, F of another
H “bond”

67. H is shared between
2 atoms of OXYGEN or
2 atoms of NITROGEN or
2 atoms of FLUORINE Of 2
different
molecules

68. Why does H “bonding” occur?
Nitrogen, Oxygen and Fluorine
small atoms with strong nuclear charges
powerful atoms
very high electronegativities

69. Intermolecular forces dictate chemical properties
Strong intermolecular forces cause high b.p., m.p. and slow evaporation (low vapor pressure) of a substance.

70. Which substance has the highest boiling point?HF
NH3
H2O
WHY?
Fluorine has the highest e-neg, SO
HF will experience the strongest H bonding and  needs the most energy to weaken the i.m.f. and boil

71. The Unusual Properties of Water
Unusually high boiling point
Compared to other compounds in Group 16

72. Density????

73. H2O(s) is less dense than H2O(l)
The hydrogen bonding in water(l) molecules is random. The molecules are closely packed.
The hydrogen bonding in water(s) molecules has a specific open lattice pattern. The molecules are farther apart.

74. The End