Download - 1 Covalent bonding Molecules & Structures. 2 What do you already know about Covalent bonding?
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Covalent bonding
Molecules & Structures
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What do you already know about Covalent bonding?
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What kind of bonding occurs when a metal and a nonmetal transfer electrons?– Ionic bonding
What is made when two metals just mix and don’t react? – An alloy
What do two nonmetals or some metalloids with nonmetals form when they bond together?– Covalent bond
Bonding Review
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Covalent bonding makes molecules Molecules are
– Specific atoms, usually nonmetals, joined by sharing electrons
Two major kinds of molecules: Molecular compound
– Sharing e- between different elements– Example: CH4
Diatomic molecules– Sharing e- between two of the same atom– These atoms occur naturally as compounds b/c
they are more stable that way– Examples:
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Diatomic elementsThere are 7 elements that always form
molecules
H2 , N2 , O2 , F2 , Cl2 , Br2 and I2
1 + 6 pattern on the periodic table
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1 and 6
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Properties of Molecular CompoundsTend to have low melting and boiling
pointsHave a molecular formula which
shows type and number of atoms in a molecule– Not necessarily the lowest ratio of
elementsEx: C6H12O6 or H2OThe molecular formula doesn’t tell you
how bonded atoms are arranged
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How does H2 form?The nuclei repel
++
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How does H2 form?
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The nuclei repel
But they are attracted to electrons
They share the electrons
So the bond forms when the attractive forces balance the repulsive forces
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How do we show bonding?Lewis structures
– Use electron-dot diagrams to show how electrons are arranged in molecules
– Ex: H2
These are called structural formulas
– They show what atoms are bonded together and the type(s) of bond(s).
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Covalent bondsNonmetals hold onto their valence
electrons.Need noble gas configuration to be stable
– Usually with 8 valence electronsGet it by sharing valence electrons with
each other.
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Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals
F F
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Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals
F F8 Valence electrons
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Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals
F F8 Valence electrons
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Single Covalent Bonds: A sharing of two valence electrons. Ex: H2
Double Covalent Bonds: A sharing of four valence electrons Ex: O2
Triple Covalent Bonds: A sharing of six valence electrons Ex: N2
As the number of shared electrons pairs increase, bond length decreases
The shorter the bond length, the stronger the bond
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Drawing Structural Formulas We use Lewis Structures to show this:
1. Predict the location of atoms- Terminal atoms will be an end atom on the structure
because they can only form one bond - Ex: H & F ALWAYS
- The central atom is usually the one that is less electronegative
2. Make sure you have the correct number of valence e- for all of your atoms
– For polyatomic ions, add one e- for each negative charge & subtract one e- for each positive charge
3. Start bonding by creating single bonds between the central atom(s) and each of the terminal atoms
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4. Make sure to share electrons between atoms as needed to get an octet on EACH atom.
– SOME EXCEPTIONS: H wants only 2 e-, Be is OK with only 4 e-, and B is OK with only 6 e-
– Remember C, N, O, & S can form double or triple bonds with the same element or another element if cannot get an octet with only single bonds
5. Redraw with lines for each shared pair of electrons (covalent bonds). Remember to enclose a polyatomic ion in brackets and indicate the overall charge on the ion.
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Some things to think about:
Too few e- for octets? Consider a double or triple bond.
Too many electrons? Can an atom have an expanded octet? Any atom in rows 3-7 of the periodic table can have an expanded octet. –Why? B/c of empty d-orbitals–The atom with the expanded
octet is usually the central atom.
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Examples:PH3
H2S
CCl4
SiO2
NO3-
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Practice Draw Lewis structures for the following:
PCl3
CH2O
C3H6
SO42-
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Practice on Molecules with More than One Central AtomC2H2
CH3COOH
H2O2
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Dealing with Exceptions to the Octet Rule
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Resonance Structures When more than one correct structure can be written
for a molecule or ion.– Usually happens with molecules or polyatomic ions
that have both double and single bonds. They only differ in the position of the electron pairs
NOT the atom positions.
EX: NO2-
Which one is the true structure? Does it go back and forth? Double bonds are shorter than single
In NO2- all the bonds are the same length
The actual molecule behaves as if it only has one structure.
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Let’s Draw Some Resonance Structures:SO2
SO32-
O3
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Coordinate Covalent BondWhen one atom donates both electrons
in a covalent bond.
Ex: Carbon monoxide
OC
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Coordinate Covalent Bond When one atom donates both electrons
in a covalent bond. Ex: Carbon monoxide
OC
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Coordinate Covalent Bond When one atom donates both electrons
in a covalent bond. Ex: Carbon monoxide
OCOC
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Another Example:The ammonium ion (NH4
+), which is formed from the combination of ammonia and a H+ ion.
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Molecules with odd numbers of electrons
Molecules when valence e- are counted cannot form octets around each atom
Examples: NO2, ClO2 and NO
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Don’t forget Expanded OctetsThe central atoms contain more than 8
valence electrons
Ex: XeF4
Ex: SF6
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Bond Dissociation Energy The energy required to break a bond C - H + 393 kJ C + H You ALWAYS have to add energy to break bonds (it
will always be a positive number) Double bonds have larger bond dissociation energies
than single Triple bonds are even larger
– Examples of bond dissociation energy between different types of carbon-carbon bonds
– C-C 347 kJ– C=C 657 kJ– C≡C 908 kJ
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Bond Dissociation Energy The larger the bond energy, the harder it is to
break the bond Large bond energies make chemicals less
reactive In chemical reactions, bonds in the reactants
are broken and new bonds are formed to make products.
The total energy change of a chemical reaction is determined from the energy of the bonds broken and formed.– This is where Endothermic and Exothermic
reactions come from
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Let’s Do An Example: First balance the equation for the
combustion of methane:
CH4 + O2 → CO2 + H2O
Draw the Lewis structures of the reactants and products
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Refer to the structures and add up the energy released from forming all the bonds in the products. (For example: if 2 moles of water are formed, that’s 4 O-H bonds)
Add up the bond energies for the moles of the reactants.
Subtract the energy used to break all the bonds in the reactants from the energy used to make the bonds in the products. This is how much energy released in the combustion reaction.
Is the reaction endothermic or exothermic?
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Bond Polarity
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Electronegativity and Polarity Covalent bonds are the sharing of electrons
between atoms. The amount of sharing can change
depending on how strongly an atom holds onto its electrons.
We use the periodic table and values of electronegativity to determine how strongly an atom will pull electrons in a bond.
The electronegativity of the atoms was assigned by Linus Pauling when he studied the bonding abilities of atoms in molecules.
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When the atoms in a bond are the same, the electrons are shared equally.
Also when there is little difference in electronegativity, the electrons are essentially shared equally
These are considered nonpolar covalent bonds.
When two different atoms are connected, the electrons may not be shared equally.
This is a polar covalent bond.
How do we measure how strong the atoms pull on electrons?
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Electronegativity A measure of how strongly the atoms attract
electrons in a bond. The bigger the electronegativity difference the
more polar the bond. Use Figure 9-15 Pg. 263 in text to get
electronegativities of atoms. We use general guidelines to determine if a
bond is polar, nonpolar or ionic.– Chemical bonds between different
elements are never completely ionic or covalent
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Use the following differences in electronegativities of bonded atoms as general guidelines for bond polarity:
0.0 - 0.4 nonpolar covalent bond
0.5 - 1.7 polar covalent bond
>1.7 Ionic
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How to show when a bond is polar
Isn’t a whole charge just a partial charge
means a partially positive
means a partially negative
The Cl pulls harder on the electrons
The electrons spend more time near the Cl
H Cl
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For each pair of elements, calculate the electronegativity difference and label the
bond type (polar covalent, nonpolar covalent, or ionic).
H, Cl
H, S
S, Cl
Na, F
Cl, Br
Al, Br
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Examples with compounds Let’s determine if the bonds in the following
compounds are polar, nonpolar, or ionic. You will need to show your calculations!
HCl
CH4
HSF
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Molecular Shapes
Some Theories
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VSEPR Theory Valence Shell Electron Pair Repulsion.
Predicts three dimensional geometry of molecules based on the number of pairs of valence electrons both bonded and unbonded.
Name tells you the theory.
Valence shell - outside electrons.
Electron Pair repulsion - electron pairs try to get as far away as possible.
Can determine both the angles of bonds
AND the shape of molecules
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VSEPRThe bond angle is the angle formed by two
terminal atoms
Both shared pair e- and lone pair e- repel each other
– Lone pair e- repel more than shared e-
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# of e- domains
Shared e- pairs
Lone pair e-
Molecular shape
Bond
Angle
Example
2 2 0 Linear 180o BeH2
3 3 0 Trigonal planar
120o BH3
4 4 0 Tetrahedral 109.5o CH4
4 3 1 Trigonal pyramidal
107.3o NH3
3 or 4 2 1 or 2 Bent 104.5o H2O
5 5 0 Trigonal bipyramidal
90o/ 120o
PF5
6 6 0 Octahedral 90o SF6
http://intro.chem.okstate.edu/1314F00/Lecture/Chapter10/VSEPR.html
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Examples of how we get the molecular shapes
Single bonds fill all atoms.
There are 4 pairs of electrons pushing away.
The furthest they can get away is 109.5º.
C HH
H
H
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4 atoms bondedBasic shape is
tetrahedral.
A pyramid with a triangular base.
Same basic shape for everything with 4 pairs. CH H
H
H109.5º
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3 bonded - 1 lone pair
N HH
H
NH HH
<109.5º
Still basic tetrahedral but you can’t see the electron pair.
Shape is calledtrigonal pyramidal.
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2 bonded - 2 lone pair
OH
H
O HH
<109.5º
Still basic tetrahedral but you can’t see the 2 lone pair.
Shape is calledbent.
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3 atoms no lone pair
CH
HO
The farthest you can the electron pair apart is 120º.
Shape is flat and called trigonal planar.
Will require 1 double bond
C
H
H O
120º
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2 atoms no lone pairWith three atoms the farthest they can
get apart is 180º.
Shape called linear.
Will require 2 double bonds or one triple bond
C OO180º
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Try to determine the shapes of these molecules:
SiH4
PF3
HBr
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How do we get the shapes of these?
C2H2
CH3COOH
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Molecular Orbitals (MO)The overlap of atomic orbitals from
separate atoms makes molecular orbitals
Each molecular orbital has room for two electrons
Two types of MO
– Sigma ( σ ) between atoms
– Pi ( π ) above and below atoms
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Sigma bonding orbitals From s orbitals on separate atoms
+ +
s orbital s orbital
+ ++ +
Sigma bondingmolecular orbital
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Sigma bonding orbitals From p orbitals on separate atoms
p orbital p orbital
Sigma bondingmolecular orbital
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Pi bonding orbitalsP orbitals on separate atoms
Pi bondingmolecular orbital
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Sigma & Pi BondsSigma bonds ( occur from overlap of orbitals between the atomsPi bond ( bond) occur between p orbitals. above and below atomsAll single bonds are
bondsDouble bond is 1
and 1 bond Triple bond is 1
and 2 bonds
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Hybrid Orbitals
Combines bonding with geometry
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Hybridization The mixing of different atomic orbitals to form the same
type of hybrid orbitals.
All the hybrid orbitals that form are identical.
Each hybrid orbital contains one electron that it can share with another atom.
The number of atomic orbitals mixed to form the hybrid orbital equals the total number of pairs of electrons (double and triple bonds get treated as though they are one pair of electrons)
Lone pairs on the central atom also occupy hybrid orbitals.
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Types of Hybridization sp3 -1 s and 3 p orbitals mix to form 4 sp3
orbitals.
EX: CH4, NH3, H2O
sp2 -1 s and 2 p orbitals mix to form 3 sp2 orbitals leaving 1 regular p orbital.
EX: BH3, AlCl3, C2H4
sp -1 s and 1 p orbitals mix to form 2 sp orbitals leaving 2 regular p orbitals.
Ex: BeCl2, CO2, O2, N2
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sp3 hybridizaiton
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sp2 hybridization
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Where is the p orbital?Perpendicular
The overlap of orbitals makes a sigma bond ( bond)
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CCH
H
H
H
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sp hybridizationwhen two things come off
one s and one p hybridize
linear
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sp hybridizationend up with two lobes 180º
apart.
p orbitals are at right angles
makes room for two bonds and two sigma bonds.
a triple bond or two double bonds
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CO2
C can make two and two O can make one and one
CCOO OO
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N2
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N2
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sp3d1 s, 3 p, and 1 d orbitals mix together
making 5 sp3d hybrid orbitals
Ex: PCl5
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sp3d2
1 s, 3 p, and 2 d orbitals mix together making 6 sp3d2 hybrid orbitals
Ex: SF6
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Molecular Polarity
How to show if the entire molecule is polar or not.
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Molecular PolarityMolecules are either nonpolar or polar,
depending on the location and nature of the covalent bonds.
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Nonpolar Molecules
There is symmetry with regard to the distribution of electrons.
Determine the shape! If there is an electronegative atom on
one part of the molecule and one that “balances” it on another part, then the molecule is nonpolar. If not, it is a polar molecule
Ex: CH4 and CCl4 and CH4Cl
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Polar Molecules Molecules with a partially positive end and a
partially negative end Symmetry can not cancel out the effects of
the polar bonds. (There is no “balancing” of electronegative atoms on another part of the molecule)
Must determine shape first.Examples: H2O and NF3
For each molecule, draw the Lewis For each molecule, draw the Lewis structure, predict the shape and structure, predict the shape and bond angle, and identify as polar or bond angle, and identify as polar or nonpolar. nonpolar. Br2
HCN
C2H2
NH4+
H2S
PF3
CH2O
MgO
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“Symmetrical” shapes are those without lone pair on central atom – Tetrahedral– Trigonal planar– Linear
The molecule will be nonpolar if all the atoms are the same or have low differences in electronegativities
Shapes with lone pair on central atom are not symmetrical
Can be polar even with the same atom bonded to the central atom
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Is it a polar or nonpolar molecule?
HF
H2O
NH3
CBr4
CO2
CH3Cl
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Properties of Molecules Most have LOW melting & boiling points
tend to be gases and liquids at room temperature
Ex: CO2, NH3, H2O
Polar and Nonpolar molecules have a little bit different properties due to the partial charges.
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H - F+ -
H - F
+-H - F+
-
H - F
+-
H - F +-
H - F+-
H - F
+-
H - F
+-
+-
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Properties of Solid MoleculesTwo kinds of crystals:
– Molecular solids – molecules held together by attractive forces
• Ex: BI3, Dry Ice, sugar
– Network solids- atoms held together by bonds
• One big molecule (diamond, graphite)• High melting & boiling points, brittle, extremely hard
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Graphite Diamond
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Intermolecular Forces
What holds molecules to each other
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Intermolecular ForcesThey are what make solid and liquid
molecular compounds possible.
The weakest are called van der Waal’s forces - there are two kinds
– Dispersion forces
– Dipole Interactions
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Dispersion ForceDepends only on the number of
electrons in the molecule
Bigger molecules more electrons
More electrons stronger forces
• F2 is a gas
• Br2 is a liquid
• I2 is a solid
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Dispersion force
H H H HH H H H
+ -
H H H H
+ - +
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Dipole interactionsOccur when polar molecules are
attracted to each other.
Slightly stronger than dispersion forces.
Opposites attract but not completely hooked like in ionic solids.
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Dipole interactionsOccur when polar molecules are
attracted to each other.
Slightly stronger than dispersion forces.
Opposites attract but not completely hooked like in ionic solids.
H F
H F
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+- +-
+-
+
-+
-
+-
+-+-
+-
+-
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Hydrogen bondingAre the attractive force caused by
hydrogen bonded to F, O, or N.F, O, and N are very electronegative so
it is a very strong dipole.They are small, so molecules can get
close togetherThe hydrogen partially share with the
lone pair in the molecule next to it.The strongest of the intermolecular
forces.
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Hydrogen Bonding
HH
O+ -
+
H HO+-
+
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Hydrogen bonding
HH
O H HO
HH
O
H
H
OH
HO
H
HO HH
O
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Video lessonWater, a polar molecule, on YouTube:
https://www.youtube.com/watch?v=iOOvX0jmhJ4
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ReviewIonic and Covalent Compounds
Practice Quiz and Graphics: http://www.elmhurst.edu/~chm/vchembook/145Areview.html
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Internet resources Molecular polarity:
http://www.elmhurst.edu/~chm/vchembook/210polarity.html
Polar covalent compounds: http://www.elmhurst.edu/~chm/vchembook/152Apolar.html
Nonpolar covalent compounds: http://www.elmhurst.edu/~chm/vchembook/150Anpcovalent.html
Ionic compounds: http://www.elmhurst.edu/~chm/vchembook/143Aioniccpds.html
Compare Ionic, Polar, and Nonpolar Bonds: http://www.elmhurst.edu/~chm/vchembook/153Acompare.html