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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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
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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)
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Why do molecules attract each other?
electrostatic attraction between… permanent dipoles (in polar molecules) temporary diploes
A dipole is a overall charge imbalance in a molecule.Which of the following molecules are polar?
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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. E.g. He b.p 4 K : Xe b.p. 165 K.
Only effective over short range so the molecule “area” is also important.
E.g: Pentane, C5H12, b.p. 309 K
Dimethylpropane, (CH3)4C b.p. 283 K
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Is a molecule polar?A polar molecule Has polar covalent bonds.
Look at the difference in electronegativity (FONClBrISCH)
AND Unsymmetrical shape according to charge
distribution.
Otherwise it will be a non-polar molecule.
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Molecular polarity
Images: http://en.wikipedia.org/wiki/Molecular_polarity
HF
H2O
NH3
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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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Polar and non-polar liquids are immiscible
Image: http://en.wikipedia.org/wiki/Petroleum
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Hydrogen bonding H bonded to a highly electronegative element eg
F, O or N proton strongly attracts electronegative element
in another molecule important in water
Image: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force
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Hydrogen bond
In molecules that contain Hydrogen bonded to Oxygen, Nitrogen or Fluorine (high electronegativity and non-bonding electron pair).
Interaction of the non-bonding electron pair in one molecule and hydrogen (with high positive charge) in another molecule.
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Examples H2O b.p.=100oC H2S b.p.= -61oC
NH3 b.p.= -33 oC PH3 b.p.= -88oC
C3H8 b.p. CH3CHO C2H5OH
b.p. 20 oC 42 oC 78 oC
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Examples H2O b.p.=100oC H2S b.p.= -61oC
NH3 b.p.= -33 oC PH3 b.p.= -88oC
C3H8 b.p. CH3CHO C2H5OH
b.p. 20 oC 42 oC 78 oC
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Ice
Image: http://en.wikipedia.org/wiki/Ice
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Trends in physical 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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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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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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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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Diamond
Hard, colourless, insulator Tetrahedral, giant structure Covalent bond => sp3 orbitals.
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Graphite
Slippery, black, conductor Layers of fused six-membered rings. Each carbon
surrounded by three others in a planar trigonal arrangement => sp2 + p-orbital
The p-orbital is perpendicular to the layer and give close packed p-orbitals
stabilise the layers Delocalisation of electrons => electrical
conductivity
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Fullerene, C60
Spherical molecule. Looks like a football. 12 pentagons and 20 hexagons.
Bonds: C60 –hydration C60H60
(C2H4 + H2 C2H6 ; 1 H2 / double bond)
Each carbon has a double bond
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Silicon
Metalloid, Semiconductors, non-metallic structure Similar structure as diamond.
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Silicon dioxide
SO2 Silica, giant structure similar to diamond
Silicates, SiO4, tetrahedrical, silicon-oxygen single
bond
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Physical properties Melting points (impurities lower the melting point) Boiling points Volatility (how easy a compound will convert to
gas) Electrical conductivity Solubility
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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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Ionic salts Typical properties
Hard, brittle, Conduct electricity in solution or melted. High melting points => Strong bonds Hydration of Ion in Water solution
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Metallic bond Metals have low electronegativity. The atoms are packed close together in a lattice. The valence electrons are delocalised among all
atoms. The valence electron have no “home” The atoms can be seen as positive ions in a see of
electrons that keep them together.
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This can explain the metallic properties Electrical conductivity: electrons float around. If
you put in one, one will fall out.
Malleability (smidbarhet) and Ductility (sträckbarhet): if the atom is pushed from its location the electron will follow. The bond is between the ion and the electrons not between the ions.
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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