This is one of the chapters you must read…. chapter 6…bonding
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Transcript of This is one of the chapters you must read…. chapter 6…bonding
This is one of the chapters you must
read….chapter 6…bonding
Student will learn:
1. three types of bonding
ionic, covalent, metallic
2. two categories of bonding
polar, non-polar
3. how to draw Lewis structure
4. how to calculate electronegative
5. bond characteristics
6. VSPER theory
Chemical Bonds ch6 p.161
?What holds chemicals together? Chemical Bonds:
electrical attraction between +nuclei and
–valence electron of different atoms
• By bonding together the atoms are more stable, and have a lower level of energy arrangement
3 types of Bonding:
l. Ionic Bonding: lose or gain –e
metal + Non-metal = ionic bonding
= makes ions
2. Covalent Bonding: share –e
Non-metal + non-metal=covalent bonds
= makes molecules
3. Metallic Bonding: -e flow free in a sea of –e
Transition metals
Pick the bonding
NaCl, CH4, HCl, K2S, FeSO4, LiF, H20, Cu, Zn, Mg(OH)2
2 catagories for the bonding
polar non-polarUnequal equal
attraction for electrons balanced attraction
Ionic bonding (Metals+nonMetals) is always polar
Covalent (nonMetals +nonMetals) maybe either polar/nonpolar
2 ways to figure out
Draw Lewis structure calculate and use chart
Lewis Structure for iodomethane CH3I1. Lewis Dot for each element
C H I2. Arrange to form skeleton
If a carbon then always in middleIf no carbon then least electronegative atom
in middleHydrogen never in middle
See that it is lopsided…… polar covalent molecule
How about individual bonds?
Lewis for ammonia NH3
1. Draw Lewis dot for each element
NH
2. Do skeleton : hydrogen never in middle
Does it look lopsided….
polar covalent molecule
Lewis for formaldehyde CH20l. Lewis dot for each C H O
2. Do skeleton: carbon always in middle
Notice left out –e….move to make a double bond
Notice lopsided: polar……..covalent
Single bonds
Double bonds
Triple bonds
Lewis struture for : CCl4
l. Draw Lewis dot for each: C CL
2. Draw skeleton: carbon always in middle
Does it look lopsided?........No…..
non-polar covalent molecule What about each bond?
Use table on p.151 and chart on page 162
Calculating Polarity of Bonds: using differences in electronegativity
Remember: electronegativity = ability to gain electrons
Bonding is rarely purely ionic or covalent………most of time somewhere in between
Use table on p.151 and chart on page 162 (overhead 31)
Subtract the two electronegativity numbers then ? is it less than 1.7
= polar covalent?
Calculate bond type and polarity
KCl, MgCl2, H2, H2S Cs2S, SCL2,
Comparing Characteristics Ionic bonds (vs) Covalent bonds
Metals + nonmetalsGain or lose electrons so….+ or – ends ……very polarWill form a crystalline lattice,
look on page 177 (model)
Stronger bondsMost are solidsHigher melting pointHigher boiling pointMany dissolve in water, +ion and -ion break apart in water so will conductelectricity in water. Some do not dissolve becausethe pull between the charges are greater thanthe attraction of H2O molecule
Hard but brittle----why?A shift of one row of ions causes a large build up of repulsive forces. And --do not like-- so if one layer moves that forces the other layers to move so they are brittle.
Non-metals + Non-metals
Share electrons
Exist as individual molecules
Weaker bonds
Most are gases, some liquids
Very low melting point
Very low boiling point
Will evaporate at room temperature
Overhead 70
Ionic compounds form
Crystalline lattice
How ionic compounds dissolve
Why Ionic Compounds are brittle
Comparing Characteristics Ionic bonds (vs) Covalent bonds
Metals + nonmetalsGain or lose electrons so….+ or – ends ……very polarWill form a crystalline lattice,
look on page 177 (model)
Stronger bondsMost are solidsHigher melting pointHigher boiling pointMany dissolve in water, +ion and -ion break apart in water so will conductelectricity in water. Some do not dissolve becausethe pull between the charges are greater thanthe attraction of H2O molecule
Hard but brittle----why?A shift of one row of ions causes a large build up of repulsive forces. And --do not like-- so if one layer moves that forces the other layers to move so they are brittle.
Non-metals + Non-metals
Share electrons
Exist as individual molecules
Weaker bonds
Most are gases, some liquids
Very low melting point
Very low boiling point
Will evaporate at room temperature
Overhead 70
Why most covalents are liquids or gases and evaporate easy.
Metallic Bonding p.181
Transitional Metals: vacant outer p orbitals
because filling up d orbitals first.
4s2, 3d10, 4p… they overlap
This overlapping lets –e roam freely about the metal network of empty atomic orbitals.
These mobile –e form a sea of electrons which are packed in a lattice form.
Overhead 68
Characteristics of Metallic bonding “Cu,Au,Ag, Fe”l. Conduct electricity
Conduct heat :::::: due to the “sea of electrons” ability to move freely
2. Reflect light, Shiny, Polish :::::: Contain many orbitals (d10)
separated by extremely small energy differences, metals can absorb a wide range of light frequencies. This absorption of light energy accounts for the ability to reflect light and be shiny. …..p. 181
3. Malleable: hammer into a thin sheet.::::::possible because the metallic bonding is same in all directions throughout the
solid because of “sea of electrons”
Ductile: ability to be drawn into a thin wire.
::::::Because the metallic bonding is same in all directions throughout the solid because of
“sea of electrons”
VSEPR THEORY
Valence Shell Electron Pair Repulsion
Theory : replusion between Valence Shell Electrons Pairs surrounding an atom causes these sets to be oriented as “far apart as possible”.
“AS FAR APART AS POSSIBLE”
Lewis dot, VSEPR TO PREDICT GEOMETRY OF MOLECULE,
Intermolecular force: the attraction between molecules
3 types: dipole-dipole
hydrogen bonding
London dispersion forces
Dipole-dipole: strongest intermolecular force
created when by equal but opposite charges of the molecule come within close distance of each other.
Hydrogen bonding: hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a near-by molecule.
“this is why water expands when frozen”
London Dispersion force: results from the constant motion of electrons and the creation of instantaneous dipoles.
Because London dispersion force depends on the motion of electrons, the strength increases with increasing atomic masses.