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Bonding. Elements form bonds in order to reach the lowest potential energy state. Electrons are transferred or shared in order to obey the ‘rule of eight’. Three types of bonds. Ionic Elements transfer electrons from a metal to a nonmetal in order to form an octet in the valence level

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Presentation Transcript
slide1

Bonding

Elements form bonds in order to reach the lowest potential energy state. Electrons are transferred or shared in order to obey the ‘rule of eight’.

slide2

Three types of bonds

Ionic

Elements transfer electrons from a metal to a nonmetal in order to form an octet in the valence level

Covalent

Two nonmetals share electrons in order to form an octet in the valence level

Metallic

Delocalized electrons ‘float’ in a sea of electrons among metal cations

slide3

Ionic bonds

metal + nonmetal

lose e- to form cation gain e- to form anion

Electrons are transferred until the net charge is zero

Example: aluminum oxide Al2O3

O

Al Total charge

O 2 x 3+ = +6

Al 3 x 2- = -6

O 0

••

2-

••

3+

••

••

2-

••

3+

••

••

2-

••

properties of ionic compounds
Properties of Ionic Compounds

Crystalline solid because ions close together

Held together by electrostatic forces (positive-negative attraction between cation and anion)

High melting point

High Boiling point

Soluble in water forms electrolytes

When dissolved in water conducts electricity

Do not conduct in solid state

Brittle

slide5

Electronegativity-ability to attract electrons

Higher the electronegativity value, stronger the attraction

The difference (absolute value; no negative numbers ) between electronegativity values can be used to predict the type of bond formed by elements.

≤ 0.2 nonpolar covalent bond

< 0.2 but ≤ 1.7 polar covalent bond

> 1.7 ionic bond

CO2     C---O   2.5 – 3.5  =  1.0   polar covalent bond 

O2 O = O 3.5 – 3.5 = 0 nonpolar covalent bond

bond facts
Bond facts
  • Number of covalent bonds that an atom can form = number of valence electrons
  • Atoms will position themselves so as to achieve the lowest possible energy
  • Distance where energy is minimum is the bond length
  • Bond will form if the energy of the aggregate is lower than that of the separated atoms

covalent ionic

0polarity increases as

electronegativity

difference increases

slide7

polar covalent bonds—electrons are shared unequally

The element that has the higher electronegativity value will pull electrons closer

This will set up a partial positive (d+) and partial negative (d-) charge. 

These partial charges are called dipoles and can be indicated using this symbol         

The arrow points toward the atom with the higher electronegativity

CO2     C O  or C O

2.5 – 3.5  =  1.0   

d+

d -

slide8

nonpolar covalent bonds-electrons are shared equally

no dipoles

Nonpolar covalent bonds are usually between identical atoms or elements located very close on the periodic table.

Analogy:

Imagine a Tug-O-War between two defensive linemen on a football team.  Each pulls with a strong force and no one wins.  Tug-O-War between two horseracing jockeys.  Each pulls with a weak force and no one wins.  Now, what if the two defensive linemen pull against the two horse-racing jockeys, the football players pull with a stronger force and win the Tug-O-War. 

Same thing can happen in a molecule. If one nucleus has a stronger attraction for the electrons, the shared pair of e- is pulled closer to that nucleus.

properties of covalent compounds
Properties of covalent compounds

Solid, liquid, gas

Low melting point < 300oC

Low boiling point

High to low solubility (variable)( less soluble than ionic)

Poor conductor- nonelectrolytes

Held together by overlap of orbitals

bond strength and bond length
Bond strength and bond length

Single bond (sigma bond) 

longer

weaker

end to end overlap

Double or triple bond (pi bond) 

shorter

stronger O C O

side to side overlap

First bond formed is always a sigma bond, second or third bond would be a pi bond

1 

1 

1 

1 

hybridization blending of orbitals
Hybridization(blending of orbitals)

2 bonding sites sp

3 bonding sites sp2

4 bonding sites sp3

5 bonding sites sp3d

6 bonding sites sp3d2

7 bonding sites sp3d3

lone pair on the central atom = bonding site

double or triple bond = one bonding site

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/hybrv18.swf

slide12

Polar and Nonpolar Molecules

Polar molecules generally result from polar bonds, but the shape of the molecule influences properties.

      CO2 C       O 2.5 - 3.5 = 1.0 polar covalent bond

Draw the electron dot structure        O C O

The linear molecule with dipoles equal and opposite will cause the net pulling force to equal zero, so this is a nonpolar molecule.

CH4               C H   2.5 - 2.1  =  .4                          polar covalent bond

Draw the electron-dot structure                      H

                                                            H         C         H

                                                                        H

the tetrahedral molecule with dipoles equal and opposite will cause the net pulling force to equal zero, so this is a nonpolar molecule.

slide13

CHCl3           C H   2.5 - 2.1  =  .4 polar covalent bond

                     C Cl  2.5 - 3.0  =  .5 polar covalent bond

Draw the electron dot structure                      H

                                                            Cl        C         Cl

                                                                       Cl

The tetrahedral molecule with dipoles that are unequal will cause the net pulling force to not equal zero, so this is a polar molecule.

Polar molecules

Lone pair will make a molecule polar

Central atom bonded to different atoms

slide14

Solubility and polarity

Polar molecules and ionic substances dissolve in polar solvents and nonpolar molecules dissolve in nonpolar solvents.  

Remember…. “Like dissolves like.”

You can predict the properties of the molecule by trying to dissolve a substance in the polar solvent, water or the nonpolar solvent, hexane.

geometry of molecules
Geometry of Molecules

Valence Shell Electron Pair RepulsionVSEPR

Valence electrons will repel each other strongly and cause the peripheral atoms to move as far from each other as possible

This repulsion will determine the shape of the molecule

Every central atom will hybridize

slide16

G

D

Example:

CO2

linear

All atoms in a line

Bond angle of 180o

bent

Usually three atoms that could be linear but instead will bend

Bond angle 104.5o

One or two lone pairs (unbonded electrons)

Always polar because of lone pair

Example:

H2O

slide17

A

L

trigonal planar

Atoms in same plane but in a triangle arrangement

Bond angle 120o

Usually a double bond

Tetrahedral

Central atom with 4 other atoms arranged around this atom

Bond angle 109o

Example:

NO3-

Example:

CH4

slide18

F

K

Trigonal pyramidal

One central atoms with three other atoms and a lone pair arranged in a pyramid shape

Bond angle 107o

Always polar because of lone pair

Trigonal bipyramidal

Central atom with 5 other atoms arranged around it

Exception to the octet rule

Bond angles 120o (trigonal planar middle) 180o and 90o

Example:

NH3

Example:

PCl5

slide19

H

Octahedral

Central atom with 6 other atoms around it

Exception to the octet rule

Example:

SF6

intermolecular forces van der waals forces attractive forces between molecules
Intermolecular forcesvan der Waals forcesattractive forces between molecules
  • London dispersion forces
    • Polar and nonpolar molecules
    • Strength increases with increased number of electrons
    • Strength increases with more mass
    • Induced dipole (results from collision with other molecules- ‘squishy cloud’)
  • Dipole-dipole forces
    • Only polar molecules
    • strong attraction between the partial positive charge and partial negative charge
  • Hydrogen bonding ( NOT a bond)
    • Polar molecules with H—F, H—O, H—N bonds
slide21
IMFs

Predict high boiling point, strong surface tension, high vapor pressure for compounds with strong intermolecular forces

Predict low boiling point, splattering, fast evaporation for compounds with weak intermolecular forces

LDF’s are weakest type of IMF

Dipole-dipole IMFs are strong

Hydrogen bonds are the strongest type of IMF

LDF's  increase with increasing mass

LDF’s increase with increased number of electrons

dipole interactions

d+d-

d+d-

d+d-

d+d-

d+d-

d+d-

d+d-

Dipole Interactions

d+d-

hydrogen bonding

H

O

O

H

H

O

H

H

H

H

O

H

H

H

H

O

O

O

H

H

H

Hydrogen bonding
slide24

Metallic Bonding

Results from the attraction between metal atoms and a surrounding sea of mobile electrons.

Valence electrons are lost to form cations.  These are delocalized, which means they do not belong to any one atom but move freely throughout the entire metal.

Metallic bonding is not directional but uniform throughout the solid.  One group of atoms can slide past another group of atoms amid the electron sea without breaking any attractions. 

Explains metallic properties like electrical and thermal conductivity, malleability, ductility, and luster.

slide25

+

+

+

+

+

+

+

+

+

+

+

+

Metal bond

‘sea of electrons’