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COVALENT COMPOUNDS
TWO TYPES OF BONDS Ionic: Electrons are transferred
Covalent: Electrons are shared Non-polar covalent: equally shared
Polar Covalent: unevenly shared
BOND POLARITY
Which element is the most electronegative?
H F
Fluorine- Has 7 valence e- and wants 8
ability of an atom to attract electrons
REVIEW:WHAT IS ELECTRONEGATIVITY?
H F FH
electron richregion
electron poorregion
e- riche- poor
d+ d-
POLAR BOND :
covalent bond with greater electron density around one of the two atoms
1
2
3 4 5 6 7 8 9 10 11 12
13 14 15 16 17
18
NonpolarCovalentshare e-
Polar Covalentpartial transfer of e-
Ionictransfer e-
Increasing difference in electronegativity
Electronegativity Difference Bond Type
0 to 0.3 Nonpolar Covalent
0.4 to 1.6 Polar Covalent
1.7 Ionic
WHAT TYPE OF BOND IS IT?
Classify the following bonds as ionic, polar covalent,or covalent:
Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic
H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent
Cl – 3.0 N – 3.0 3.0 – 3.0 = 0 Nonpolar Covalent
Cs to Cl
H to S
Cl to N
DO YOU NOTICE A PATTERN FOR THE COMBO OF ELEMENTS THAT ARE IONIC VS COVALENT? Ionic bonds form between:
Covalent bonds form between:
Identify the following as ionic, covalent, or both:
CaCl2 BaSO4
CO2 AlPO4
SO3 H2O
PROPERTIES OF COVALENT COMPOUNDS
Usually soft and squishy
Not soluble in water
Does not conduct electricity
Soluble in organic solvents
Low melting points
Low boiling points
PROPERTIES OF IONIC COMPOUNDS Combination of ions (cation/anion)
Tightly packed solids in a crystal lattice
Hard and Brittle
Usually soluble in water
Conducts electricity when dissolved
High melting points
High boiling points
NAMING COVALENT COMPOUNDS
NAMING COMPOUNDS
Nonmetal – Nonmetal
USE PREFIXES!
1. Change the ending of the second word to -ide
2. No mono on the first word
3. Drop any double vowels
COVALENT PREFIXESNumber of Atoms Prefix
1 Mono-2 Di-3 Tri-4 Tetra-5 Penta-6 Hexa-7 Hepta-8 Octa-9 Nona-
10 Deca-
EXAMPLES
1. CO
2. CO2
3. SO2
4. SO3
5. N2H4
6. N2O3
1. Carbon Monoxide
2. Carbon Dioxide
3. Sulfur Dioxide
4. Sulfur Trioxide
5. Dinitrogen Tetrahydride
6. Dinitrogen Trioxide
EXAMPLES
1. disilicon hexafluoride
2. tricarbon octachloride
3. phosphorus pentabromide
4. nitrogen monoxide
5. selenium difluoride
6. dihydrogen monoxide
1. Si2F6
2. C3Cl8
3. PBr5
4. NO
5. SeF2
6. H2O
EMPIRICAL AND MOLECULAR FORMULAS
Define Empirical Formula:A chemical formula that gives the simplest whole-number ratio of the elements in the formula.
Which of the following is an empirical formula? CO2 C2O4
N2H4 NH2
Define Molecular Formula:A chemical formula that gives the actual number of the elements in the molecular compound. For the following molecular formulas, write the empirical formula:
Molecular: Empirical: C2H4 C6H12O6
C9H21O6N3
LEWIS STRUCTURES
OCTET RULE
Eight electrons in the valence shell (filling s and p orbitals) make an atom STABLE
This is called the octet rule
Bond formation follows the octet rule…Chemical compounds tend to form so that each atom:by gaining, losing, or sharing electrons, has an octet of electrons in its valence energy level.
LEWIS DOT DIAGRAMS
• an electron-configuration notation with only the valence electrons of an element are shown, indicated by dots placed around the element’s symbol.
• tracks the number of valence electrons
• the inner core electrons are not shown
LEWIS DOT PRACTICE
Li N F
Be O Ne
F F
LEWIS STRUCTURES FOR COMPOUNDS The pair of dots between two symbols
represents a shared pair. How many shared pairs does each fluorine have
below?
An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.
F F
LEWIS STRUCTURES
The shared pair of electrons is often replaced by a long dash.
Each dash represents TWO electrons
F F+
7e- 7e-
F F
8e- 8e-
F F
F F
Lewis structure of F2
lone pairslone pairs
lone pairslone pairs
single covalent bond
single covalent bond
WHY SHOULD TWO ATOMS SHARE ELECTRONS?
To get a valence of 8 electrons!
HC
HC
H
H
MULTIPLE COVALENT BONDS double bond:
covalent bond in which two pairs of electrons are shared between two atoms
shown by two side-by-side pairs of dots or by two parallel dashes
MULTIPLE COVALENT BONDS triple bond:
covalent bond in which three pairs of electrons are shared between two atoms
shown by three side-by-side pairs of dots or by three parallel dashes
Bond Type
Bond Length(pm)
C-C 154
CC 133
CC 120
C-N 143
CN 138
CN 116Bond Lengths
Triple bond < Double Bond < Single Bond
LENGTHS OF COVALENT BONDS
BOND LENGTH AND BOND ENERGY
As atomic size increases, bond length increases, and as a result bond energy decreases
As you increase the number of bonds between two atoms, energy increases, while bond length decreases.
BOND LENGTH AND BOND ENERGY EXAMPLES
1. Which bond is greater in length: Br2 or F2?
2. The HF bond is 570 pm, the H2 bond is 436 pm, which bond requires more energy to break?
3. Which bond would require more energy to break C-C single bond or C=C double bond? Which bond is longer?
WRITING LEWIS STRUCTURES
1. Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.
2. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.
3. Complete an octet for all atoms except hydrogen
4. If structure contains too many electrons, form double and triple bonds on central atom as needed.
Step 1 – N is less electronegative than F, put N in center
F N F
F
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
WRITE THE LEWIS STRUCTURE OF NITROGEN TRIFLUORIDE (NF3).
Step 1 – C is less electronegative than O, put C in center
O C O
O
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e-
4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-
2 single bonds (2x2) = 41 double bond = 4
8 lone pairs (8x2) = 16Total = 24
WRITE THE LEWIS STRUCTURE OF THE CARBONATE ION (CO3
2-).
When there are two or more Lewis structures for a single molecule
O C O
O
- -O C O
O
-
-
OCO
O
-
-
What are the resonance structures of the carbonate (CO3
2-) ion?
RESONANCE STRUCTURE:
SOME ELEMENTS DO NOT FOLLOW THE OCTET RULE
H HBe
F B F
F
There can also be expanded octets!
MOLECULAR GEOMETRY
VSEPR THEORY
Lewis Dot Diagrams are 2D but we live in a 3D world.
How are molecules actually arranged??
Follows the Valance Shell Electron Pair Repulsion Theory or VSEPR
AB2 – LINEAR
Number of Surround Atoms Number of Lone Pairs Bond Angle
2 0 180˚
Cl ClBe
AB3 – TRIGONAL PLANAR
Number of Surround Atoms Number of Lone Pairs Bond Angle
3 0 120˚
AB2E1 – BENT
Number of Surround Atoms Number of Lone Pairs Bond Angle
2 1 <120˚
AB4 – TETRAHEDRAL
Number of Surround Atoms Number of Lone Pairs Bond Angle
4 0 109.5˚
AB3E1 – TRIGONAL PYRAMIDAL
Number of Surround Atoms Number of Lone Pairs Bond Angle
3 1 107˚
AB2E2 – BENT
Number of Surround Atoms Number of Lone Pairs Bond Angle
2 2 104.5˚
PREDICTING MOLECULAR GEOMETRY
1. Draw Lewis structure for molecule.
2. Count number of lone pairs on the central atom and number of atoms bonded to the central atom.
3. Use VSEPR to predict the geometry of the molecule.
What are the molecular geometries of SO2 and SF4?
SO O
AB2E
bent
C
F
F
F F
AB4
tetrahedral
INTERMOLECULAR FORCES
Intermolecular forces: attractive forces between molecules.Intramolecular forces: attractive forces within a molecule (the bonds)
intramolecular forces are much stronger than intermolecular forces
Intramolecular Forces
Intramolecular Forces
Intermolecular Forces
DIPOLES
What is a dipole?
A polar molecule
Uneven sharing of electrons so there is a separation of charge
DIPOLE-DIPOLE FORCES
Attraction between two polar molecules
— + — +
HYDROGEN BONDING
Special type of Dipole – Dipole
Attraction between:Hydrogen and Nitrogen/Oxygen/Fluorine
DIPOLE – INDUCED DIPOLE Attraction between one polar and one
nonpolar molecule
— +
— + — +
Electrons shift toward
positive end of dipole
LONDON DISPERSION FORCES Attraction between two nonpolar molecules
— + — +
Electrons become
uneven and form a dipole
STRENGTH OF IMF
Hydrogen Bond
Dipole – Dipole
Dipole – Induced Dipole
London Dispersion Forces
strongest
weakest
Which of the following molecules is polar?H2O, CO2, SO2, and CH4
O HH
dipole momentpolar molecule
SO
O
CO O
no dipole momentnonpolar molecule
dipole momentpolar molecule
C
H
H
HH
no dipole momentnonpolar molecule
SO
O
What type(s) of intermolecular forces exist between each of the following molecules?
HBrHBr is a polar molecule: dipole-dipole forces. There are
also dispersion forces between HBr molecules.
CH4
CH4 is nonpolar: dispersion forces.
SO2
SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.
WHAT DOES IMF EFFECT? Viscosity
Surface Tension
Cohesion/Adhesion
Boiling Point
Stronger IMF Higher Viscosity
VISCOSITY
Measures a fluid’s resistance to flow
Stronger IMF Higher Surface Tension
SURFACE TENSION
result of an imbalance of forces at the surface of a liquid.
BOILING POINT
Point at which liquid particles escape the surface of the liquid into the gas phase
Stronger IMF Higher Boiling Point
Adhesion
Cohesion
ADHESION AND COHESION Cohesion:
intermolecular attraction between like molecules Adhesion:
intermolecular attraction between unlike molecules
