ib dp1 chemistry bonding
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IB DP1 Chemistry Bonding. What makes atoms join together to make compounds?. Topic 4: Bonding (12.5 hours). 4.1 Ionic bonding 4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged ions. 4.1.2 Describe how ions can be formed as a result of electron transfer. - PowerPoint PPT PresentationTRANSCRIPT
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IB DP1 ChemistryBonding
What makes atoms join together to make compounds?
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Topic 4: Bonding (12.5 hours)
4.1 Ionic bonding4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged ions.4.1.2 Describe how ions can be formed as a result of electron transfer.4.1.3 Deduce which ions will be formed when elements in groups 1, 2 and 3 lose electrons.4.1.4 Deduce which ions will be formed when elements in groups 5, 6 and 7 gain electrons.4.1.5 State that transition elements can form more than one ion.4.1.6 Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values.4.1.7 State the formula of common polyatomic ions formed by non- metals in periods 2 and 3.4.1.8 Describe the lattice structure of ionic compounds.4.2 Covalent bonding4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei.4.2.2 Describe how the covalent bond is formed as a result of electron sharing.4.2.3 Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom.4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength.4.2.5 Predict whether a compound of two elements would be covalent from the position of the elements in the periodic table or
from their electronegativity values.4.2.6 Predict the relative polarity of bonds from electronegativity values4.2.7 Predict the shape and bond angles for species with four, three and two negative charge centres on the central atom using the valence shell electron pair repulsion theory (VSEPR).4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities.4.2.9 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C60 fullerene).4.2.10 Describe the structure of and bonding in silicon and silicon dioxide.4.3 Intermolecular forces4.3.1 Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how theyarise from the structural features of molecules.4.3.2 Describe and explain how intermolecular forces affect the boiling points of substances.4.4 Metallic bonding4.4.1 Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons.4.4.2 Explain the electrical conductivity and malleability of metals.4.5 Physical properties4.5.1 Compare and explain the properties of substances resulting from different types of bonding.
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Ionic Bonding
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Crystals: 7 ‘perfect’ crystal shapes
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Halite- rock salt- sodium chloride
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Sodium chloride is an ionic compound with ions arranged in a lattice
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Ionscharged particles with electrostatic attraction between them
Na+ Cl-
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Sodium and chloride ions formed when electrons transfer
Na + Cl Na+ + Cl-
2,8,1 2,8,7 2,8 2,8,8
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Ions Group 1: H+, Li+, Na+, K+, Rb+, Cs+, Fr+
Group 2: Be2+, Mg2+, Ca2+, Sr2+, Ba2+
Group 3?/13: B3+, Al3+, Ga3+
Group 6?/16: O2-, S2-, Group 7?/17: F-, Cl-, Br-, I-
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Which is the smallest ion?
Na+
Al+3
Cl-P3-
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Two or more electrons can be transferred
Different sized atoms give different mineral structures as they pack in a different way
Hexagonal Beryl crystal; Image Wikipedia
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What is the formula of iron (III) oxide?
Fe2OFeOFe3O2Fe2O3
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Polyatomic ions: charge distributed over more than one atom
For example phosphate, PO4-
3
can be found in products of reactions of phosphoric acid
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Some common polyatomic ions Nitrate NO3
-
Hydroxide OH- Sulphate SO4
2-
Carbonate CO32-
Hydrogen carbonate HCO3-
(Bicarbonate)
Phosphate PO43-
Ammonium NH4+
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Common Anions Common Name Formula Alternative
name Simple Anions Chloride Cl− Fluoride F− Bromide Br− Oxide O2− Polyatomic anions Carbonate CO3
2- Hydrogen carbonate
HCO3− bicarbonate
Hydroxide OH− Nitrate NO3
2- Phosphate PO4
3- Sulfate SO4
2- Anions from Organic Acids Ethanoate CH3COO− acetate Methanoate HCOO− formate Ethandioate C2O4
−2 oxalate Cyanide CN-
Common Cations Common Name Formula Alternative
name Simple Cations Aluminium Al3+ Calcium Ca2+ Copper(II) Cu2+ cupric Hydrogen H+ Iron(II) Fe2+ ferrous Iron(III) Fe3+ ferric Magnesium Mg2+ Mercury(II) Hg2+ mercuric Potassium K+ kalic Silver Ag+ Sodium Na+ natric Polyatomic Cations Ammonium NH4
+ Hydronium H3O+
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Careful with... name of atom can change when ion is formed
chlorine atom (Cl) chloride ion (Cl-)
-ate is often a polyatomic ion with oxygen eg sulphate, phosphate, etc.
different ions often have similar names... nitrate NO3
-
nitrite NO2-
nitride N-3
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What is the formula of ammonium sulphate? NH4SO4 (NH4)2SO4 NH4(SO4)2 SO4(NH4)2
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d-block (transition elements) can have variable valencies
Mn2+ manganese(II)Mn3+ manganese(III)Mn4+ manganese(IV)Ni2+ nickel(II)/nickelousNi3+ nickel(III)/nickelicPb2+ lead(II)/plumbousPb4+ lead(IV)/plumbic
Cr2+ chromium(II)/chromousCr3+ chromium(III)/chromicCu1+ copper(I)/cuprousCu2+ copper(II)/cupricFe2+ iron(II)/ferrousFe3+ iron(III)/ferricHg2+ mercury(I)/mercurous
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Covalent bonding
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Define electronegativity
Electronegativity is the tendency of an atom to attract electrons towards itself. The atoms with higher values attract electrons more strongly.
Highest flourine (and rest of groups 7,6,5)FONClBrISCHWikipedia table
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How ionic is an ionic compound? bigger difference in electronegativity more ionic (‘ionic’ usually De-neg> 1.8 difference) usually metal + non-metal
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Which aluminium compounds will be ionic?atom Al F O Cl Brelectronegativity
1.5 4.0 3.5 3.0 2.8
Formula of aluminium compound
De-neg ‘Ionic’ or ‘covalent’?
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‘Sharing’ electrons De-neg < 1,7covalent bonding forms molecules
Often between non-metals
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Covalent bond formation- valence electrons
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2, 4 or 6 electrons? Single bond: the two atoms share two electrons
(1 pair) Double bond: the two atoms share four
electrons (2 pairs) Triple bond: the two atoms share six electrons (3
pairs)
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Lewis structures (dot structures) show valence electrons in pairs as dots, crosses or lines
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skeletal formula for complex organic molecules
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Condensed formulapropanol CH3CH2CH2OH
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Coordinate covalent bond (dative bond)
both electrons in the bond from the same atomonce formed, is the same as any other covalent bond
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Bond lengths and Bond strengths
As the number of shared electrons increases (single to triple) the bond lengths shortens and the bond energy increase
Bond Bond type Lengths (pm) Energy (kJ/mol)
CC Single 154 347
CC Double 134 614
CC Triple 120 839
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Which bond has the highest bond polarity, δ
H-HCl-ClAl-FAl-Br
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Non-polar covalent bond
In, H2 the two electrons in the bond are shared equally between the two hydrogen atoms. H-H De-neg =0. The electron distribution is symmetrical.
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Polar covalent bond If two different atoms form a covalent bond there
will be a difference in De-neg.
The atom with highest electronegativity will have the electrons closer; they don’t share equally.
Unsymmetrical electron distribution.
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Bonds100% Covalent bond Polar covalent bond Ionic bond % ionic character of a bond: 0-90%
(there are no 100% ionic compounds)
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Molecular shapes
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What shape are molecules? VSEPR theory (Valence shell electron pair
repulsion) pairs of electrons repel and sit as far away as
possible from each other double and triple bonds count as a pair
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VSEPR: electron repulsion molecular shape
Structure of molecule given by pairs of electrons arranging around an atom to be as far apart as possible
non-bonded pairs repel more than bonded pairs double and triple bonds count as one
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Build molecules from plasticine and straws bond: 3cm length of straw atom: 1cm diameter plasticine ball unbonded pair of electrons 1cm straw length
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Number of charge centres
Name of shape Bond angles (s)
Example
2 linear 180 BeCl23 trigonal planar 120 BF3
4 tetrahedral 109.5 CH4
5 trigonal bipyramidal
90, 120, 180
6 octahedral 90, 180
Shapes of simple molecules
http://en.wikipedia.org/wiki/Phosphorus_pentafluoridehttp://en.wikipedia.org/wiki/Sulphur_hexafluoridehttp://en.wikipedia.org/wiki/Boron_triflouride
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Methane, Water and Ammonia
greater repulsion between non-bonding pairssmaller bond angles than predicted
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VSEPR for 5- and 6-negative charge centres
Some molecules have expanded valence shell with five or six negative centres.
Shapes similar to trigonal bipyramid and octahedron.
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Trigonal bipyramidTwo types of electron rich region:
Equatorial: 3 bonds with 120o between. Axial: 2 bonds with 180o between
Equatorial to Axial: 90o. Non-bonding orbitals always occupy equatorial
positions. Many of the compounds that form trigonal bipyramids
and octahedrons are fluorides because only highly electronegative atoms can increase the number of valence electrons. Fluoride is also quite small, but there is not enough space for larger ions.
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Octahedron All positions are equal. 90o angle If two non-bonding orbitals: they take place
opposite each other plane square shape.
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Examples PCl5 SF6
XeF4
PF6-
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Intermolecular forcesWhy do molecules stick together to form liquids and solids?
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Intermolecular forces hold molecules together, affecting physical properties
Melting and boiling points Strength Flexibility Viscosity Deflection in electric field Volatility (how easy a compound will convert to gas) Electrical conductivity Solubility
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Intermolecular forcesHydrogen bond strongDipole-dipole weakervan der Waal’s forces weakest
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Why do molecules attract each other to make liquids and gases?
Intermolecular forces: electrostatic attraction between permanent dipoles (polar molecules) permanent dipole and a temporary dipole
(induced polarity) temporary diploes (induced polarity)A dipole is a overall charge imbalance in a molecule.
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Induced dipoles in all molecules (van der Waal’s forces)
Image: http://www.uwec.edu/boulteje/Boulter103Notes/11December.htm
Movements in electron cloud Temporary dipoles.
Temporary dipole in one molecule can induce a temporary dipole in another.
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van der Waals forces The strength increases with molar mass of the
molecule. He b.p 4KXe b.p. 165K.
Only effective over short range so the molecule “area” is also important.
Pentane, C5H12, b.p. 309K
Dimethylpropane, (CH3)4C b.p. 283K
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Trends in physical properties
melting point /C boiling point /CFlourine -220 -188Chlorine -102 -34Bromine -7 59Iodine 114 184Astatine 302 337
Plot one graph showing melting point and boiling point (in Kelvin) against molar mass for the halogensDescribe the pattern (2 sentences)Explain the pattern (2 sentences)
Data: http://en.wikipedia.org/wiki/Halogen
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Is a molecule polar?A polar molecule has polar covalent bonds.
Is there a difference in electronegativity? (FONClBrISCH)
AND has an asymmetric shape according to charge
distribution.
Otherwise it is a non-polar molecule.
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Molecular polarity
Images: http://en.wikipedia.org/wiki/Molecular_polarity
HF
H2O
NH3
http://phet.colorado.edu/en/simulation/molecule-polarity
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Molecular polarity http://phet.colorado.edu
/en/simulation/molecule-polarity
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Dipole-dipole
Electrostatic attraction between molecules with permanent dipoles.Stronger than vdW.Hydrogen chloride M= 36,5 g/mol b.p. 188 KFluorine M= 38 g/mol b.p. 85K
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Induced dipole
Image: http://www.uwec.edu/boulteje/Boulter103Notes/11December.htm
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Hydrogen bonding H bonded to a highly electronegative element- F,
O or N proton unbonded pair important in water
Image: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force
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ExamplesH2O b.p. 373K H2S b.p 212K
NH3 b.p. 240K PH3 b.p 185K
C3H8 bp20 oC
CH3CHO bp42 oC
C2H5OH bp78 oC
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Ice
Image: http://en.wikipedia.org/wiki/Ice
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Polar and non-polar liquids are immiscible
Image: http://en.wikipedia.org/wiki/Petroleum
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Allotropes
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Allotropes: different structural forms of the same element
http://catalog.flatworldknowledge.com/bookhub/4309?e=averill_1.0-ch18_s04
OxygenO2 diatomic oxygenO3 ozone
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Allotropes of Carbon
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C allotropes: Diamond Hard, colourless, insulator Tetrahedral, giant structure Covalent bonds sp3
orbitals.
Image: http://en.wikipedia.org/wiki/Diamond
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C allotropes: Graphite Slippery, black conductor Layers of fused six-
membered rings Each carbon surrounded by 3
others in a trigonal planar arrangement sp2 + p-orbital
p-orbital perpendicular to layers and gives close-packed p-orbitals
Delocalized electrons electrical conductor
Image: http://en.wikipedia.org/wiki/Diamond
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C: Allotropes: Fullerene, C60
Spherical molecule 12 pentagons and 20
hexagons.
Image:http://en.wikipedia.org/wiki/Fullerine
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C allotropes: Carbon nanotubes
Image:http://en.wikipedia.org/wiki/Fullerine
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Silicon solid at room temperature high melting and boiling points of 1414 and 3265
°C conducts heat well grey color and a metallic luster strong, very brittle crystallizes in a diamond cubic crystal structure
Images: http://en.wikipedia.org/wiki/Silicon
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Silicon dioxide SO2 Silica giant structure
similar to diamond silicates, SiO4,
tetrahedral, silicon-oxygen single bond
Image: http://www.green-planet-solar-energy.com/silicon-element-facts.html
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Metals
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Metallic bond Metals have low
electronegativity. Atoms packed into a lattice. Valence electrons
delocalised valence electrons have no
“home” atoms positive ions in a sea
of electrons that keep them together.
Image: http://www.bbc.co.uk/schools/gcsebitesize/science/add_ocr_gateway/periodic_table/metalsrev2.shtml
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Patterns in bonding and properties
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How strong are the forces between molecules?
Bond type Dissociation energy (kJ/mol)
Covalent 1600Hydrogen bonds 50–70Permanent dipoles 2–8Induced dipoles <4
Data: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force
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Ionic salts Hard, brittle, Conduct electricity in solution or melted. High melting points and boiling points Ions hydrated in aqueous solution
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Covalent molecular compounds low mp and bp poor conductors of electricty and heat
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Metallic properties Electrical conductivity: electrons float around. Put
one in, one is pushed out.
Malleability and Ductility: if the atom moves, the electron follows. Bond is between ion and electrons, not between ions.
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Summary of properties Structure typeProperty
GiantMetallic
GiantIonic
GiantCovalent
MolecularCovalent
Hardness and malleability
Variable hard-ness, malleable rather than brittle
Hard and brittle Hard and brittle Usually soft and malleable unless hydrogen bonded
Melting and boiling points
Variable dep. On No of valence e-
High Very High Low
Electrical and thermal conductivity
Good in all states
Not as solids, conduct in (aq) or (l)
No No
Solubility
Insoluble, except as alloys
In Water mostly Insoluble Often more soluble in other than water except if H-bonded
Examples Iron, copper NaCl, Na2SO4 Diamond,SiO2 (Sand)
CO2, Cl2, ethanol, sugar
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Investigation
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Investigate a physical property of a mixture related to intermolecular forces
Quantitative independent variable (cause)
Quantitative dependent variable (effect) viscosity, deflection by charged object, or other physical property
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Links Ionic bonding
http://www.teachersdomain.org/asset/lsps07_int_ionicbonding/
Covalent bonding http://www.teachersdomain.org/asset/lsps07_int_covalentbond/
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Polarity links
http://phet.colorado.edu/en/simulation/molecule-polarity
Viscosity http://www.youtube.com/watch?v=3KU_skfdZVQ
States of matter http://phet.colorado.edu/en/simulation/states-of-matter
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Polarity links http://phet.colorado.edu/en/simulation/molecule-polarity http
://antoine.frostburg.edu/chem/senese/101/liquids/faq/h-bonding-vs-london-forces.shtml
States of matter http://phet.colorado.edu/en/simulation/states-of-matter http://employees.oneonta.edu/viningwj/modules/
CI_dipoleinduced_dipole_forces_13_5a.html Notes: http://www.uwec.edu/boulteje/Boulter103Notes/11December.htm Snowflakes: http://www.its.caltech.edu/~atomic/snowcrystals/class/
class.htm Ice crystals http://www.edinformatics.com/interactive_molecules/
ice.htm
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Links http://phet.colorado.edu/en/simulation/molecule-
shapes http://en.wikipedia.org/wiki/
Phosphorus_pentafluoride http://en.wikipedia.org/wiki/Sulphur_hexafluoride http://en.wikipedia.org/wiki/Boron_triflouride
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Teaching notes