1. Chemical Bonds and Shapes
Chapter 7 and 8

2. Review electron shells
Atomic number = number of Electrons in neutral atoms
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

3. 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.
The electrons in the outer energy shells are called VALENCE ELECTRONS
Only focuses on s and p orbitals

4. 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

5. 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 (looks familiar?)

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

9. Formation of Ions from Metals -Cations
Ionic compounds result when metals react with nonmetals
Metals lose electrons to match the number of valence electrons of their nearest noble gas (lower in atomic number)
Positive ions form when the number of electrons are less than the number of protons (more positives than negatives)
Group 1A metals → ion 1+
Group 2A metals → ion 2+
Group 3A metals → ion 3+

10. Some Typical Ions with Positive Charges (Cations)
Group 1A Group 2A Group 3A
H Mg Al3+
Li Ca2+
Na Sr2+
K Ba2+

11. Learning Check: Aluminum
A. Number of valence electrons aluminum
1) 1 e ) 2 e- 3) 3 e-
B. Change in electrons for octet
1) lose 3e ) gain 3 e ) gain 5 e-
C. Ionic charge of aluminum
1) ) ) 3+

12. Solution A. Number of valence electrons in aluminum 3) 3 e-
B. Change in electrons for octet
1) lose 3e-
C. Ionic charge of aluminum
3) 3+

13. Learning Check Give the ionic charge for each of the following:
A. 12 p+ and 10 e-
1) 0 2) 2+ 3) 2-
B. 50p+ and 46 e-
1) 2+ 2) 4+ 3) 4-
C. 15 p+ and 18e-
2) ) 3- 3) 5-

14. Ions from Nonmetal Ions Anions
In ionic compounds, nonmetals in 5, 6, and 7 gain electrons from metals
Nonmetal add electrons to achieve the octet arrangement
Nonmetal ionic charge:
3-, 2-, or 1-

15. Chemical bonds: an attempt to fill electron shells
Ionic bonds –
Covalent bonds –
Metallic bonds

16. IONIC BOND bond formed between two ions by the transfer of electrons

17. Ionic Bonds Atoms gain or lose electrons to form charged ions.
+/- attraction creates electrostatic attraction to form bond
Recognize by Metal with a Nonmetal (M/NM)
Ex: NaCl AlBr3 MgSO4
Ionic Compounds are also called salts

18. Ionic Bond
Cations are attracted to the anions and combine in ratios that make a neutral charge
More examples; NaCl, CaCl2, K2O

19. F K In an IONIC bond, electrons are lost or gained,
resulting in the formation of IONS
in ionic compounds. F K

20. K F

21. K F

22. K F

23. K F

24. K F

25. K F

26. + _ K F

27. K F _ + The compound potassium fluoride
consists of potassium (K+) ions
and fluoride (F-) ions

28. K F _ + The ionic bond is the attraction between the positive K+ ion
and the negative F- ion

29. Ionic Compounds Properties
Most are crystalline solids at room temp.
High melting and boiling points
They are electrolytes (conduct e- in water)

31. What is an ionic bond?
What are properties of ionic bonds?

32. Polyatomic Ion Formulas
Concept:
Polyatomic ions are groups of atoms that behave as one unit.
*Some ions have more that one atom but their overall charge can be determined.
Examples: (SO4) = S +6 and O -8 = -2
(NO3)= N +5 and O -6 = -1
(NH4) = N -3 and H +4 = +1

33. Polyatomic Ion Formulas
These group ions, polyatomic ions, are treated like single ions in formulas, but must have parentheses when more than one is used in a formula.
Examples: Ca(NO3) Ga2(SO4)3
(NH4)2O

34. COVALENT BOND bond formed by the sharing of electrons

35. In covalent bonding,
atoms still want to achieve
a noble gas configuration
(the octet rule).

36. But rather than losing or gaining electrons,
In covalent bonding,
atoms still want to achieve
a noble gas configuration
(the octet rule).
But rather than losing or gaining
electrons,
atoms now share an electron pair.

37. The shared electron pair is called a bonding pair
In covalent bonding,
atoms still want to achieve
a noble gas configuration
(the octet rule).
But rather than losing or gaining
electrons,
atoms now share an electron pair.
The shared electron pair is called a bonding pair

38. Covalent or Molecular Bonds
Atoms share v.e. to obtain octet rule.
Recognize by Nonmetal and Nonmetal (NM/NM)
Ex: CO NH H2O
Organic molecules are Covalent
Ex: C6H12O6 (sugars), amino acids

39. Covalent or Molecular Properties
Tend to be liquids or gases at room temp.
Lower melting/boiling point than Ionic
Non-electrolytes (doesn’t conduct e- in water)
Some elements tend to be Diatomic
Diatomic elements come in pairs when not in compound
-N2, O2, F2, Cl2, Br2, I2, H2
BrINClHOF

40. What are Diatomics?
Diatomic elements: elements that only exist when paired with itself.
Ex: H2
There are 7 diatomics you MUST memorize
BrINClHOF- Magic 7
Br2, I2, N2, Cl2, H2, O2, F2

41. Bonds in all the polyatomic ions and diatomics are covalent bonds

42. Chlorine
forms a
covalent
bond
with
itself
Cl2

43. Cl Cl do to achieve an octet? What’s the solution – what can they
Neither atom will give up an electron –chlorine is highly electronegative.
What’s the solution – what can they
do to achieve an octet?

44. How
will
two
chlorine
atoms
react? Cl Cl

45. Cl Cl

46. Cl Cl
octet

47. Cl Cl This is the bonding pair circle the electrons for
each atom that completes
their octets

48. Single bonds are abbreviatedCl Cl
Single bonds are abbreviated
with a dash
circle the electrons for
each atom that completes
their octets

49. O2
Oxygen is also one of the diatomic molecules

50. O
How will two oxygen atoms bond?

51. O

52. Both electron pairs are shared.

53. O O
two bonding pairs,
making a double bond

54. This is the oxygen molecule,=
This is the oxygen molecule, O2

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

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

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

58. Ionic, Non-Polar or Polar?
You can use the difference in electronegativies (ΔEN) to distinguish between the types of bonds:
Ionic- ΔEN 1.9 and above
Covalent- ΔEN below 1.9
Polar
Non-Polar

59. METALLIC BOND bond found in metals; holds metal atoms together very strongly

60. Metallic Bonds
Metallic bonds explain why metals have the properties they have (shiny, etc)
Bonds can occur within an element or Alloy
Alloy is a mixture of 2 or more metals
Stainless steel, bronze, brass, etc.
In metallic bonds v.e. are mobile (like a “sea” of e-. V.E. can float between atoms.
Al Brass Bronze

61. Metallic Bond, A Sea of Electrons

62. Metals Form Alloys Metals do not combine with other metals.
They form alloys which is a solution of a metal in a metal.
Examples are steel (iron and carbon), brass (copper and zinc), bronze (copper and usually tin or other metals) and pewter (tin with copper, antimony, or bismuth).

63. Chemical Bonding and VSEPR

64. Attractions between molecules
Intermolecular forces are weaker than ionic or covalent bonds
Van der Waals forces
Dipole interactions: attraction between 2 or more polar molecules
Dispersion forces: weakest IMF, caused by the motion of electrons
Hydrogen bonds: force in which a hydrogen covalently bonded to a highly electronegative atom is also weakly bonded to an unshared electron pair of another EN atom
Relatively strong IMF
Occurs because hydrogen’s nucleus is electron deficient when bonded

65. The Shapes of Molecules
The shape of a molecule has an important bearing on its reactivity and behavior.
The shape of a molecule depends a number of factors. These include:
Atoms forming the bonds
Bond distance
Bond angles

66. Valence Shell Electron Pair Repulsion
Valence Shell Electron Pair Repulsion (VSEPR) theory can be used to predict the geometric shapes of molecules.
VSEPR revolves around the principle that electrons repel each other (like repels like).
One can predict the shape of a molecule by finding a pattern where electron pairs are as far from each other as possible.

67. Bonding Electrons and Lone Pairs
In a molecule some of the valence electrons are shared between atoms to form covalent bonds. These are called bonding electrons.
Other valence electrons may not be shared with other atoms. These are called non-bonding electrons or they are often referred to as lone pairs.

68. VSEPR*
In all covalent molecules electrons will tend to stay as far away from each other as possible
The shape of a molecule therefore depends on:
the number of regions of electron density it has on its central atom,
whether these are bonding or non-bonding electrons.

69. Lewis Dot Structures
Lewis Dot structures are used to represent the valence electrons of atoms in covalent molecules
Dots are used to represent only the valence electrons.
Dots are written between symbols to represent bonding electrons

70. VSEPR Predicting Shapes

71. VSEPR: Predicting the shape
Once the dot structure has been established, the shape of the molecule will follow one of basic shapes depending on:
The number of regions of electron density around the central atom
The number of regions of electron density that are occupied by bonding electrons

72. VSEPR: Predicting the shape
The number of regions of electron density around the central atom determines the electron skeleton
The number of regions of electron density that are occupied by bonding electrons and hence other atoms determines the actual shape

73. *Basic Molecular shapes
The most common shapes of molecules are shown at the right

74. Linear Molecules
Linear molecules have only two regions of electron density.

75. Angular or Bent
Angular or bent molecules have at least 3 regions of electron density, but only two are occupied
Bond angles tend to be about 105⁰

76. Triangular Plane*
Triangular planar molecules have three regions of electron density.
All are occupied by other atoms

77. Tetrahedron*
Tetrahedral molecules have four regions of electron density.
All are occupied by other atoms

78. Trigonal Bipyramid*
A few molecules have expanded valence shells around the central atom. Hence there are five pairs of valence electrons. The structure of such molecules with five pairs around one is called trigonal bipyramid.

79. Octahedron*
A few molecules have valence shells around the central atom that are expanded to as many as six pairs or twelve electrons. These shapes are known as octahedrons