Post on 19-Jan-2016
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Chemical Bonds: The Formation of Compounds from Atoms
Dr. Bixler-Zalesinsky
PERIODIC TRENDS IN ATOMIC PROPERTIES
11.1
Metals and Nonmetals (review)
Ionization energy is the energy required to remove an electron; corresponds to their charge
Atomic Radii increase going down a group and decrease across a period
LEWIS DOT DIAGRAMSValence Electrons
Valence Electrons & Per. Table
Lewis Structures of an atom shows the valence electrons (ones involved in bonding)
Octet Rule
• Every atom aspires to have eight electrons in its outermost shell (2 s electrons and 6 p electrons just like the noble gases)
• They must borrow (covalent molecules), release or accept (ionic compounds) electrons to get to the eight.
Types of bonding
BONDINGIons and
Ionic bonding occurs between metals (cations) and nonmetals (anions)
The nonmetal accepts the electron(s) and the metal donates the electron(s) ionic bond is the attractive between oppositely charged ions
Form large crystals; our formulas are the smallest whole number ratios not
the true number of atoms
The charges must cancel out and equal zero to form stable compounds If you have a +2 ion then you need either two -1 ions or one -2 ion
Link to Video clip on Ionic Bonding (1:39)
COVALENT V. IONIC
Bonding
11.5 Covalent Bonding: Sharing Electrons
• Covalent bonding occurs between two nonmetal atoms
• Electrons are shared between two atoms
11.7 Lewis Structures of Molecules
1. Find the number of valence electrons for each element in the structure
2. Multiply the number of valence electrons times the number of atoms you have of that element
3. Determine which element can make the most bonds and put it in the center and attach the other elements to it
4. Make each atom have 8 valence electrons around it. 5. Add up the number of electrons you used in the
structure. This number must match the total number of electrons you started with
H2.1Li1.0
Na1.5
He
B2.0
C2.5
N3.0
O3.5
F4.0 Ne
Na 0.9
Mg1.2
Al1.5
Si 1.8
P2.1
S2.5
Cl3.0
Ar
K08
Ca1.0
Br2.8
Kr
Rb0.8
Sr1.0
Se2.4
Cs0.7
Ba0.9
I 2.5
Fr0.7
Ra0.9
Xe
Rn
Number of Valence Electrons:
1 2 3 4 5 6 7 8
3 4 3 2 1 0 (number of bonds each can make) S S/D/T S/T S/D S 0 (types of bonds
s=single, D= double, T= triple)
Page 151 in textbook
Ex. Write the Lewis Dot Diagrams for the following molecules• I2• H2O
• FCl
• CF4
• NBr3
Molecule (Covalent) Nomenclature• Naming: These binary inorganic
molecules are named using prefixes like mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-,
• The first element only gets a prefix if there is more than one of it; otherwise, the element name remains the same.
• The second element ALWAYS gets a prefix and the ending changes to –ide.
• Ex. CO is carbon monoxide (two words not capitalized)
HW p. 164 # 2 a – f Write the question and answer!
Multiple Bonds
Double and Triple Bonds Double and Triple Bonds Knowing when NOT to use them is as important as understanding when to use them!
Multiple Bonds
• Some times using the correct number of electrons will not give you a full octet. When this happens:
• 1st double check your math and counting
• 2nd see if the atoms involved can make a double or triple bond
Double Bonds
• O, S, and C can make double bonds with each other than themselves but no others!
• Double bond is 4 electrons in a bond
• Symbolized by an = sign
• Take a look at CO2
Triple Bonds
• P, N, and C can make a triple bond with each other or themselves
• A triple bond is 6 electrons in a bond
• The symbol for a triple bond is =
• Let’s try N2
Molecular Geometry
VSEPR Theory and Application
Structural Formula
• Shows how elements of a molecule are connected to each other
VSEPR • V = valence• S = shell• E = electron• P = pair• R = repulsion• Electrons will arrange themselves as far
apart from one another as possible • Unbonded pairs take up more room than
bonded ones
3-D
Hybridized orbitals, shapes, and decision tree
Linear Shape
Bent Shape
Trigonal planar
Pyramidal
Tetrahedral
VSEPR Video Review (3:21)
sp hybridization
Path to hybridization
Sp3 hybridization
HYBRIDIZATIONVideo Review (1:36)
VSEPR Theory of Molecular Geometry # of atoms
Central Atom
Shape Bond angle
Example
VSEPR Theory of Molecular Geometry # of atoms
Central atom
Shape Bond Angle
(in
degrees)
Example
2 None Linear 180 HF
3 Any Linear 180 CO2
3 S or O Bent 105 H2O
4 B Trig. Planar
120 BCl3
4 (3-D) P or N Pyramidal 107 NH3
5 (3-D) C or Si Tetrahedral
109.5 CH4
Shape Decision Tree
Polar Covalent v.
Nonpolar Covalent
Polar and Nonpolar Covalent Bonds• If they are shared equally they are said to be nonpolar
bonds if they are not equally shared they are said to be polar bonds
• Sharing of electrons has to do with the pull of one element compared to the other element sharing the electron pair. This pulling is called electronegativity (eneg) which increases across the period and up the group
• The larger the electronegativity the greater the time the electrons spend with the more electronegative atom giving it a slightly positive charge and because of this imbalance we call this a polar molecule
• If the eneg difference lies between 0.5 to 1.6 it is a polar bond
Polar or Nonpolar Molecules1. Determine the shape of the molecule2. Determine how many polar bonds
there are in the molecule3. If there are NO polar bonds the
molecule must be NONpolar.4. If there is exactly one polar bond, the
molecule is polar. 5. If there is more than one polar bond,
follow the chart below.
Molecules with more than one polar bond (assuming polarity is equal)
Shape # of polar bonds
Molecular polarity
Linear 2 Nonpolar
Bent 2 Polar
Trig Planar 2 Polar
Trig Planar 3 Nonpolar
Pyramidal 2 Polar
Pyramidal 3 Polar
Tetrahedral 2 Nonpolar
Tetrahedral 3 Polar
Tetrahedral 4 Nonpolar
VSEPR Theory of Molecular Geometry # of atoms
Central Atom
Shape Bond angle
Example
Polarity of Molecules with more than one polar bondShape # of polar
bondsMolecular Polarity