1. Bonding

2. Calcium Chloride Copper Chloride
The Properties of a Compound are Different from the Properties of the Elements that make up that Compound
Calcium Chloride
Calcium
Chlorine
Copper Chloride
Copper
Chlorine

3. Subscript multiplies to all subscripts in parentheses!
Chemical Formula: Used to represent the kinds and numbers of atoms in a chemical compound.
HC2H3O2
Subscript
Be(ClO2)2
Subscript multiplies to all subscripts in parentheses!
NaHCO3 MgCl2
Al(OH)3 Ca(NO3)2

4. What must happen for elements to react with each other?
Atoms of the elements must collide.
The valence electrons are what actually collide in when elements form compounds.
This is because the valence electrons are the outermost layer.

5. Valence Electrons
The electrons in the outermost energy level of an atom
Responsible for an atom’s chemical properties

6. LEWIS DOT DIAGRAMS
A diagram where dots are placed around the chemical symbol of an element to illustrate the valence electrons

7. Octet Rule
Octet = 8
States that atoms tend to gain, lose or share electrons in order to acquire a full set of valence electrons.
For most elements this is 8 valence electrons
Hydrogen & Helium only require 2 valence electrons
The elements are all trying to be like the Group 18 elements, the noble gases b/c they have 8 valence electrons
Atoms form compounds to become stable!

8. How do atoms obtain a noble gas configuration?
By sharing or transferring electrons with other atoms
Example: Show how Na & Cl can become stable.
The bond between Na & Cl is due to:

9. ANIONS: Have gained electrons and have a negative charge (nonmetals)
If atoms transfer electrons to form a bond they become IONS
ION: AN ATOM THAT HAS AQUIRED A CHARGE BY GAINING OR LOSING ELECTRONS
CA+IONS: Have lost electrons and have a positive charge (metals)
ANIONS: Have gained electrons and have a negative charge (nonmetals)

10. Ionic Compound: Compound formed from ionic bonds
Ionic Bond: The strong attractive force between ions of opposite charge. Occurs when one atom transfers electrons to another atom to become stable

11. EMPIRICAL FORMULA Chemical formula for an ionic compound
Lowest whole number ratio of ions in an ionic compound

12. Types of elements in an ionic compound
An ionic bond is formed between a ______________ and a __________ because:
Ionic bonds are formed between ions of opposite charge
Therefore, ionic bonds form between ____________ and _____________ because metals form cations and nonmetals form anions.
nonmetal
metal
nonmetal
metal

13. How do Ionic Compounds Form?
Electrons are _______________________ in an ionic bond because: one atom is trying to lose electrons to become stable and the other atom is trying to gain electrons to become stable
transferred

14. 3 WAYS TO DETERMINE AN EMPIRICAL FORMULA
Al2O3
Subscript: # written to the lower right of a chemical symbol that shows the number of atoms of that element present in the compound
3 WAYS TO DETERMINE AN EMPIRICAL FORMULA
Use Lewis Dot Diagrams
Use charges
Use the Crisscross Method

15. Draw the Lewis Dot Diagram for each element
DETERMINING THE EMPIRICAL FORMULA BY USING
LEWIS DOT DIAGRAMS TO ILLUSTRATE THE IONIC BOND 1
Draw the Lewis Dot Diagram for each element
Use arrows to show the transfer of electrons between atoms
Determine the number of each element necessary to make each atom

16. PROPERTIES OF IONIC COMPOUNDS
Made of ions
Typically a metal and nonmetal(s)
Crystalline Solids
Hard yet Brittle
Strong Interparticle Forces
High Melting and Boiling Points
Will conduct electricity when molten or dissolved in water

17. INTERPARTICLE FORCES: Forces of attractions between neighboring particles (atoms, ions or molecules)
Strong
Ionic compounds have __________________ interparticle forces due to the strong electrostatic attraction between the cation(s) and anion(s)
Because of these ________________ interparticle forces,
ionic compounds have _____________ melting points.
Strong
high

18. CRYSTAL: A regular, repeating arrangement of ions in an ionic compound
The arrangement of ions in an ionic compound determines the crystal structure
(shape) of the crystal itself.

19. CRYSTAL
CaCl2

20. ELECTROLYTE: A compound that conducts an electric current when it is in an aqueous solution or in a molten state

21. IONIC COMPOUNDS
Ions of opposite charge strongly ___________ each other. Ions of like charge strongly ______________ each other.
As a result of this, how are ions arranged in an ionic compound?
attract
repel
Positive ions tend to be near negative ions and farther from other positive ions.

22. Metallic Bonding: The attraction of free floating valence electrons for the positively charged metal ions. Made of just metal atoms
Pure metals are not simply 1 atom of the metal. They are multiple atoms of that metal bonded by a sea of electrons.

23. Metallic Properties Come From their “Sea of Electrons”
Metallic bonds form from a “Sea of Electrons”: The pool of electrons shared by all the atoms in a metallic substance

24. Malleability: The ability to be hammered into sheets
Metals are malleable because the metal atoms can slide through the electron sea to new positions

25. Ductile: Able to be pulled into wires
Metals are ductile because electrons in the sea of electrons move to allow atoms to align like a wire

26. Conductivity Electricity is caused by moving electrons
Metals conduct because the sea of electrons is free to move and, therefore, free to conduct electricity

27. Interparticle Forces
Metallic bonds have strong interparticle forces. For this reason metals:
Are typically solids
Are crystalline
Metals are arranged in very compact and orderly patterns
Have average to high melting

28. ALLOY: A mixture composed of 2 or more elements, at least 1 is a metal
NAME OF ALLOY
% MAKE UP
EXAMPLE
Stainless Steel
73-79% Fe
14-18% Cr
7-9% Ni
Sterling Silver
92.5% Ag
7.5% Cu
18-karat white gold
75% Au
12.5% Ag
12.5% Cu
14 karat gold
58% Au
14-28% Ag
14-28% Cu

29. Covalent Bonding: Bonding in which electrons are shared
Electrons spend most of their time between the atoms.
The attraction between the nucleus and the shared electrons holds the atoms together.

30. Molecular Compound
Compound formed by covalent bonds (from the sharing of electrons)

31. Covalent Bonds Form molecules instead of crystals.
Molecule: The combination of atoms formed by a covalent bond

32. Types of Elements in a Covalent Bond
Therefore, covalent compounds are formed between 2 or more ____________________ because _____________ are the type of element that want to gain, not lose, electrons.
In a covalent bond, the atoms share electrons because both atoms want to gain electrons.
nonmetals
nonmetals

33. Ionic Compound (Metal & Nonmetal) or Covalent Compound (All Nonmetals)
Which type of bond would form between the following elements?
MgCl2 NI3 AlN
CO2 F2 SnO2

34. How do Covalent Bonds Form?
Two nonmetals share electrons to form a bond

35. Covalent Bonds
Structural Formula: Shows the arrangement of atoms in a molecule or polyatomic ion
Space Filling Model Ball & Stick Model Structural Formula  O H H

36. LEWIS STRUCTURES
Uses dashes and dots to show the bonding arrangement of atoms in a covalent compound
Drawn to model the bonding in a covalent compound.
Based on: Lewis Dot Diagrams
Dashes (): Each dash represents a bond (or 2 shared electrons)
LONE PAIR

37. Single Covalent Bond 1 shared pair of electrons between 2 atoms
Represented by: 
Example:

38. Double Covalent Bond 2 shared pairs of electrons between 2 atoms
Represented by: =
Example:

39. Triple Covalent Bond 3 shared pairs of electrons between 2 atoms
Represented by: ≡
Example:
Lone Pairs (): Represents an unshared pair of electrons

40. Using Lewis Dot Diagrams to Determine the Lewis Structure
Draw the Lewis Dot Diagrams for each atom in the Molecular Formula
Determine how the atoms will SHARE electrons so that all atoms are

41. (N-A)/2 = # of Bonds in Molecule
N = # of electrons needed to make an each element in the compound stable
N = 8 electrons for all elements except H and He which need 2 electrons
A = # of valence electrons each element in the compound has
There are 2 electrons per bond so we divide the # of electrons they need to share by 2 to determine the number of bonds

42. Some General Guidelines
Carbon is typically in the center of the molecule
If possible, keep the molecule symmetrical
Hydrogen & the halogens can only accept one electron and, therefore, tend to be on the perimeter of the molecule
Sometimes, the formula will help you figure out the structure
Check your work: Do all atoms have 8 electrons? Does H have only 2 electrons?

43. (N-A)/2 = # of Bonds in Molecule
NH3 Cl2 CH4

44. (N-A)/2 = # of Bonds in Molecule
H2CO C2H2 NH4+

45. Resonance Structures: Two or more equally valid structures of a molecule or polyatomic ion

46. Exceptions to the Octet Rule
1) Atoms with less than an octet
Boron may not acquire a full octet, it may only obtain 6 electrons--it has only 3 valence electrons to start
Atoms with more than an octet
Especially phosphorus & sulfur
Molecules with Odd Numbers of Electrons
Any molecule that has an odd number of available electrons, especially compounds of Nitrogen

47. Coordinate Covalent Bond
A covalent bond in which one atom contributes both bonding electrons   O C  C O           

48. Properties of Covalent Compounds
Made of molecules
All nonmetals
Often liquids or gasses
Brittle if solid
Weak Interparticle Forces
Low Melting and Boiling Points
Do not conduct electricity when molten or dissolved in water

49. More Properties…
Covalent compounds have ______________ interparticle forces
Because of these _______________ interparticle forces, covalent compounds have low melting and boiling points
Bond Dissociation Energy: The energy needed to break the bond between two covalently bonded atoms
A high bond dissociation energy corresponds to a strong covalent bond
weak
weak

50. Ionic
Covalent
Metallic
Type of Elements
State
Electrons are…
Hard or brittle?
Melting Points
Made of…

52. MOLECULAR GEOMETRY YES!!! NO!!! VSEPR THEORY
Valence Shell Electron Pair Repulsion Theory
In small molecules, the pairs of valence electrons are arranged as far apart from each other as possible
YES!!!
NO!!!

53. Ball-and-Stick Models
Drawings that represent molecular compounds. The balls are used to represent atoms and the sticks are used to represent bonds.

55. VSEPR Theory: Electrons repel (bonds & lone pairs move as far apart as possible)

56. The geometric angle between 2 bonds
Bond Angles
The geometric angle between 2 bonds

58. Other Shapes Octahedral SF6 T-Shaped ICl3 Square Planar XeF4
Trigonal Bypyramidal
PF5

59. EXAMPLES
CO2 PF3 CBr4

60. Bond Length Different pairs of atoms form bonds with different lengths
Not represented in a ball-and-stick model
As you move down a group on the periodic table, the atoms form longer bonds because the atoms become larger
Multiple bonds are shorter than single bonds
The more electrons there are, the more attraction there is for the opposite nuclei
Bonds are “electrical glue”

61. Electronegativity: The ability of an atom to attract electrons when the atom is in a compound
Shows how much an atom will attract valence electrons
The higher the electronegativity the more an atom attracts the electrons towards it

62. EN
The difference between two atoms electronegativities determines the type of bond formed between those two atoms
EN = EN1-EN2
 is the Greek Letter delta and means change
EN is always positive (always subtract the smaller number from the larger number)
Example: Calculate the EN between Cesium and Fluorine

63. You can think of bonding as a tug-of-war for electrons between atoms
The electronegativity of the atom tells you how hard that atom is pulling for the electrons

64. Ionic Bond
When one atom wants the electrons so much more than the other atom that it pulls the electron off the other atom (a tug-of-war where one side wins)
EN > 2.0 (greater than 2.0)

65. Nonpolar Covalent Bond
Like an even tug-of-war—both atoms want the electrons equally so the electrons stay between them
EN < 0.4 (less than 0.4)

66. Polar Covalent
One atom wants the electrons more than the other but not enough to pull it all the way towards itself—like an uneven tug-of-war
EN is between 0.4 & 2.0

67. Electronegativity Difference
EN = EN1 –EN2
Nonpolar
Polar
Ionic
0.0
0.4
2.0

68. Polar Covalent Bonds
Polar bonds have partially charged atoms due to the unequal sharing of electrons
Dipole: A partial charge on an atom.
If the EN on a bond is between 0.4 and 2.0 then that bond has a dipole (partial charge)
This is because one atom wants the electrons more than the other but not enough to cause a complete transfer

69. C-Cl H F H F DIPOLE : Symbol for a dipole or a partial charge
Used to label the partial charges on a polar bond
Two ways to label the dipoles (partial charges) on a polar bond
C-Cl + - H F H F or
2.1
4.0
EN = =0.9 (polar)

70. Images of Polar/Nonpolar

71. To determine the polarity of a molecule:
Does the molecule have only nonpolar bonds?
YES: It is nonpolar
Is the molecule completely symmetrical in all dimensions?
If the molecule contains at least 1 polar bond
and is not completely symmetrical in all
dimensions, then the whole molecule is POLAR

72. Polarity of Carbon Dioxide
CO2

73. Polarity of Ammonia
NH3

74. H2O

75. CH4

76. The Periodicity of Electronegativity

77. Electronegativity Differences & Bonding
The greater the difference between the electronegativities of the bonding atoms, the more unequally those atoms share electrons
The symbol for electronegativity difference: EN

78. Ionic Bonds
EN is > 2.0
This is because one atom wants the electrons much more than the other
Because of the high EN you can assume that the electron on the less electronegative atom is transferred to the other atom

79. Nonpolar Covalent Bonds
EN is < 0.4
This is because both atoms have almost an equal attraction for the electrons
Because of the low EN the atoms are shared equally
These molecules are gases or low-boiling point liquids at room temperature
These bonds are also known as Nonpolar Covalent Bonds

80. Example Label the dipoles on a CCl bond.
Could also use vectors to label a polar bond.
Example: Use a vector to label the dipoles on a C-Cl bond.

82. To determine the polarity of a molecule:
Draw the Lewis Structure of the Molecule with the correct geometry
Use electronegativity values to determine if any of the individual bonds are polar. If so, label the dipoles.
Using vector addition, check to make sure the dipoles do not cancel out.
If at least 1 vector/dipole does not cancel out, the molecule is polar.
If all the vectors/dipoles do cancel out, the molecule is nonpolar
Vectors will only ALL cancel out if the molecule is linear or tetrahedral and completely symmetrical!

83. VECTOR ADDITION
Involves vector addition. = = = =

84. Polarity of Ammonia
NH3

85. H2O

86. Iodine—I2

87. Methanol: CH3OH

88. Dimethyl Ether (CH3)2O

89. EXAMPLES
CO2 CH4 CH2O

90. Properties of Polar Covalent Bonds
Higher Boiling Points than Nonpolar Covalent Bonds
Molecules can “stick” together and form puddles
Hydrophilic: ionic & polar (mix easily with water)
Like dissolves like

91. Large Molecules In a large molecule the polarity oftentimes
helps determine the molecule’s shape.

92. ATTRACTIONS BETWEEN MOLECULES
Responsible for determining whether a
molecular compound is a gas, liquid or
solid at a given temperature.

93. Van der Waals Forces Consist of Dipole Interactions & Dispersion
Dispersion Forces: Caused by the motion of electrons
Weakest of all molecular interactions
Dipole Interactions: Occur when polar molecules are attracted to one another
Similar to but much weaker than ionic bonds

94. Hydrogen Bonds The attraction to a hydrogen atom already
bonded to a strongly electronegative atom
Strongest of intermolecular forces.
Extremely important in determining the properties of water and biological molecules