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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