Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Chapter 6: Chemical Bonding
Objectives:
Ionic Bonding
Show ionic bonding of elements in different ways.
1
Covalent Bonding
Show covalent bonding in different ways by using different formulas.
1
Determine the type of bond from electronegativity.
2
Determine whether molecules are polar or nonpolar by using VSEPR theory.
3
Identify and/or explain the hybridization of the carbon atom.
4
Use the hybridization of the carbon atom to show structures of compounds 5
containing carbon.
Identify molecules and their structures in the lab.
6
Identify and/or explain types of intermolecular bonding.
7
Terms for Ionic Bonding:
Chemical Bond
Octet Rule
Ionic Bonding
Ionic Compound
Empirical Formula
Formula Unit
Valence Electrons
Write the orbital notation for phosphorus.
P: 1s
3s
2p
2s
3p
Write the noble gas orbital notation for phosphorus.
P: [Ne]
3s
3p
The noble gas in the series preceding the element is written in brackets. Only the valence electrons are shown with arrows. Only the outermost "s" and "p" orbitals.
Write the noble gas orbital notation for germanium.
Write the noble gas orbital notation for iodine.
Write the noble gas orbital notation for barium.
Write the noble gas orbital notation for selenium.
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Ionic Bonding
Example 1: Sodium and Chlorine
A. Orbital Notation
Na: [Ne]
Octet rule must apply.
Show only valence electrons.
3s
Cl: [Ne] 3s 3p
Each time an electron is moved, place a "+" on the element that loses the electron, and a "­" on the element that gains the electron.
Metals will lose electrons and become positive. Nonmetals will gain electrons and become negative.
B. Electron Dot Notation
C. Ionic Formula
drop the dots
D. Empirical Formula
drop the charges
E. Ratio
F. Chemical Name
In all formulas the positive ion is always written first.
G. Chemical Formula
Ionic compounds form lattice work in crystalline form. Ionic compounds never exist as single ions bonded together. They do form in definite ratios.
nit
U
la
mu
r
Fo
Na Cl
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Example 2: beryllium and fluorine
A. Orbital Notation
Be:
Draw as many atoms as you need for either element.
F:
B. Electron Dot Notation
C. Ionic Formula:
D. Empirical Formula
E. Ratio
F. Chemical Name:
G. Chemical Formula:
Be
F
Write the following examples in your notes. Include the headings for each part. 3. sodium and fluorine
4. potassium and oxygen
5. aluminum and fluorine
6. beryllium and sulfur
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Covalent Bonding:
Terms for Covalent Bonding:
covalent bond
molecule
single covalent bond
double covalent bond
triple covalent bond
diatomic molecule
Lewis structure
molecular formula
hybridization
• Ionic bonds are formed by the transfer of valence electrons.
Covalent bonds are formed by the sharing of valence electrons. • Ionic compounds form crystal structures.
Covalent compounds form molecules.
• Ionic compounds form between a metal and a nonmetal. Covalent compounds form between nonmetals.
• Ionic bonds have charges, "+" and "­".
Covalent bonds have no charges.
• Ionic compounds do not stand alone as individual atoms bonded together.
Molecules can stand alone.
Hydrogen is unique in the way it forms bonds.
1. Hydrogen can lose an electron and become positively charged, and be stable.
2. Hydrogen can gain an electron and become negatively charged, and be stable. 3. Hydrogen can share electrons and be stable.
When hydrogen bonds to another hydrogen atom:
#1
#2
H
1s
H
H
1s
1s
H
1s
#3
H
1s
H
1s
All 3 choices create stability for hydrogen, but in this case only choice three will work.
For extra points write the reason why?
Both hydrogens have the same electronegativity so neither one is more attracted to electrons than the other so they share them.
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Example 1: hydrogen, H2
A. Orbital Notation
H
single covalent bond
H
Never have more than 2 electrons in a box. Boxes must be formed where there are single electrons.
B: Electron Dot Notation:
Show each atom individually, and where it bonds to other atoms.
No subscripts!
No charges!
Show shared paired of electrons between the atoms. Show any other valence electrons around the atoms.
C. Lewis Structure
Replace the shared pair of electrons with a dash. Show all other valence electrons in their proper places. Approximate the shape of the molecule. D. Molecular Formula:
E. Shape of the molecule:
If there are only two atoms in the molecular formula, the shape is linear.
Hydrogen is an example of a diatomic molecule.
There are 7 diatomic molecules:
hydrogen
nitrogen
oxygen
fluorine
chlorine
bromine
iodine
This is considered the free state of these 7 elements. These 7 elements do not exist as single atoms on their own.
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Example 2: oxygen, O2
A. Orbital Notation
double covalent bond
B: Electron Dot Notation:
C. Lewis Formula
Try to space out pairs of electrons equally around the atoms, and try to space the individual atoms as far apart as possible.
D. Molecular Formula:
E. Shape of the molecule:
Show covalent bonding for the following molecules:
3.
4.
5.
6.
7.
8.
9.
10.
chlorine, Cl2
nitrogen, N2
hydrogen chloride, HCl
water, H2O
hydrogen peroxide, H2O2
ammonia, NH3
hydrazine,N2H4
dinitrogen dihydride, N2H2
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Lewis Structures of Molecules
1. Hydrogen
shape:____________
3. Chlorine
shape:____________
5. Hydrogen Chloride
shape:____________
2. Oxygen
shape:____________
4. Nitrogen
shape:____________
6. Water
shape:____________
Extra pairs of unshared e­ will make a molecule bend.
7. Hydrogen Peroxide
shape:____________
8. Ammonia
shape:____________
Looks like a tripod
9. Hydrazine
shape:____________
10. dinitrogen dihydride
shape:____________
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Orbital Notation for Carbon:
__ __ __ __
C: [He] 2s 2p
Hybridization of carbon molecules.
All atoms undergo hybridization when they form covalent molecules.
Carbon needs to make 4 bonds.
Carbon in this form can only make two bonds.
__ __ __ __
C: [He] 2s 2p
By rearranging its orbitals when it bonds, carbon can make the four bonds.
__ __ __ __
C: [He] 2s 2p
These bonds must be equal in strength. The orbitals combine to create one type of orbital to give equal strength bonds.
__ __ __ __
C: [He] 2sp3
Show carbon in its hybridization form for all molecules.
Carbon Molecules:
11. methane, CH4
12. ethane,
C2H6
13. ethyne,
C2H2
14. carbon dioxide, CO2
15. ethyl alcohol, C2H5OH Hint: Bond the two carbons together, then bond the oxygen to one of the carbons, and fill in the hydrogens.
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Structural Formulas of Molecules
11. Methane
shape:____________
13. Ethyne
shape:____________
12. Ethane
shape:____________
14. Carbon Dioxide
shape:____________
15. Ethyl Alcohol
VSEPR Theory ­
Valence Shell Electron Pair Repulsion Theory
Molecules adjust their shape so that electron pairs will push apart as far as possible.
Bonding Pictures
Atoms Video
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Ionic compounds are formed by a metal and a nonmetal.
Covalent compounds are formed by two nonmetals.
Ionic character­ the amount of ionic bonding that exists in a bond between two elements.
100 % Ionic 0% ionic
0 % covalent 100 % covalent
Electronegativity difference between two elements determines the ionic character.
The greater the difference the greater the ionic character.
Three types of bonding:
Ionic bonding ­ there is a complete transfer of charge by the electrons.
Polar covalent ­ one element pulls harder on the electrons than the other element. There is an unequal sharing of electrons.
Pure covalent ­ there is an equal sharing of electrons between atoms.
(nonpolar covalent)
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Electronegativity Table page 161
Electronegativity Difference range:
0 ≤ ED < .4 pure covalent bond
.4 ≤ ED < 1.7 polar covalent bond
F
Fr____
H
H____
N
P____
1.7 ≤ ED ionic bond
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
4.
Cl
K_______
5.
N
Si_______
6.
S
Se_______
7.
Cl
Na_______
8.
P
As_______
9. N
Al_______
Ionic bonding is stronger than covalent. Polar covalent is stronger than pure.
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Types of molecules:
Nonpolar molecules ­ there is an equal distribution of charge within the molecule.
Polar molecules (dipoles) ­ there is an unequal distribution of charge within the molecule. One end of the molecule will be slightly more positive, and one end slightly more negative.
To determine polarity:
1. Are there polar bonds?
No, then the molecule is definitely nonpolar.
Yes, then consider the shape of the molecule.
2. Are the electrons being pulled in one general direction (One end of the molecule)?
Look at the Lewis structure to determine polarity.
No, the molecule is nonpolar.
Yes, then the molecule is polar.
1. hydrogen
2. hydrogen chloride
3. water
4. carbon dioxide 5. methane
6. ammonia
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Intermolecular bonding
Van der Waals forces ­ forces of attraction that exists between molecules.
1. London dispersion forces ­ weakest of the intermolecular forces. These forces are caused by the attraction of opposite charges within molecules.
All molecules have dispersion forces.
H H H Covalent:
Dispersion:
H 2. Polar forces ­ Forces of attraction that exists between polar molecules.
O
O
S S S S Covalent:
Dispersion:
Polar:
3. Hydrogen bonding ­ the strongest of the intermolecular bonding.
It exists between the hydrogen atom of one molecule and highly electronegative atom of another molecule.
H Cl
Cl
H Covalent:
Dispersion:
Polar:
Hydrogen:
H Cl
H Cl
Cl
H
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Chem1 H Chapter 6 Notes Block 4.notebook
April 28, 2017
Chapter 6: Book Assignment
Page 209 1,2,3,4,6,8,10,20,21,25,26,33,37,39,41,
42,44,47,51,52,57
Chemical Bonds Video
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Chem1 H Chapter 6 Notes Block 4.notebook
lithium and sulfur
a. Orbital notation
April 28, 2017
b. electron dot
c. Ionic formula d. ratio
e. empirical formula
f. name
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Chem1 H Chapter 6 Notes Block 4.notebook
carbon tetrachloride
a. Orbital notation
April 28, 2017
b. electron dot
c. Lewis structure d. molecular formula
e. shape
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Attachments
Ionic Bonding.ppt