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BONDING: Part 1
Topic 4
•Ionic Bonding
•Covalent bonding
Let’s Review• Valence electrons
– electrons in the highest occupied energy level – always in the s and p orbitals
• normally just a draw a circle to represent these two orbitals
– determines the chemical properties of an element– usually the only electrons used in chemical bonds
Let’s Review• when forming compounds, atoms tend to achieve
the electron configuration of a noble gas (ns2np6)– highest energy level will be filled with 8 electrons the
easiest way possible– atoms of metallic elements (groups 1,2,3) lose
electrons producing cations (positive ions)• Ca becomes Ca2+
– atoms of nonmetallic elements (groups 5,6,7) gain electrons producing anions (negative ions)
• Cl becomes Cl1-
– group 4 can go either way
NeNeNNNaNa FF
NaNa+
OO
OO2-
MgMg
MgMg2+
Cations
Anions
NN3- FF1-
...etc.
As it turns out, atoms bond together for a very simple reason: atoms like
to have full valence shells.
1+ 2+ 3-3+ 4+/- 2- 1- 0
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IB Topic 4: Bonding
Only one group of elements are stable (nonreactive). What is unique about their electron structure?
Filled s & p sublevels
All other elements react in order to achieve this stable electron configuration.
Ionic Bond: Transfer of electrons; metal + nonmetalCovalent Bond: Sharing of electrons; nonmetal +
nonmetal
IONIC BONDINGIdentification
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4.1.1 Describe the ionic bond as the electrostatic attraction between
oppositely charged particles.
Ionic bonding occurs when one or more electrons are transferred from the outer shell of another
atom.
When one atom gives up its electron(s), it becomes a positively charged cation.
When the other atom gains the electron(s), it becomes a negatively charged anion.
OPPOSITES ATTRACT…The ionic bond formed between the two elements is
caused by the attractive force (electrostatic attraction) between the positive and negative
ions.
Get it??
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Na “gives” Cl one electron and now both atoms have a full valence shell (electron configuration of a noble gas)
3.9
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4.1.2 Making ionic bonds
•Sodium gives up an electron to chlorine.
•Sodium now has a full 2nd shell, & chlorine has a full 3rd shell.
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4.1.2 Describe how ions can be formed as a result of electron transfer.
Reaction between Sodium and Chlorine
• Since opposite charges attract, the Na+ and Cl- ions form an ionic bond.
• Formula NaCl (1:1 ratio)
• Name: Sodium chloride
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4.1.2 Describe how ions can be formed as a result of electron transfer.
Ions are formed when atoms gain or lose electrons
Reaction between Sodium and Chlorine
• Sodium configuration: 1s22s22p63s1 or
__ __ __ __ __ __1s 2s 2p 3s
• If Na loses one electron then it would end in 2s22p6 and be stable.
• Then sodium has 11 protons (11+), but only 10 electrons (10-) so it acquires a charge of 1+ and becomes the sodium ion, Na+. In order for it to lose an electron, something has to gain an electron
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4.1.2 Describe how ions can be formed as a result of electron transfer.
Ions are formed when atoms gain or lose electrons
Reaction between Sodium and Chlorine
• Chlorine configuration: 1s22s22p63s23p5 or
__ __ __ __ __ __ __ __ __1s 2s 2p 3s 3p
• If Cl gains one electron then it would end in 3s23p6 and be stable.
• Then chlorine has 17 protons (17+), and 18 electrons (18-) so it acquires a charge of 1- and becomes the chloride ion, Cl-. It will gain the electron from the sodium.
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Li + F Li+ F -
The Ionic Bond
1s22s1 1s22s22p5 1s2 1s22s22p6
[He] [Ne]
Li Li+ + e-
e- + F F -
F -Li+ + Li+ F -
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4.1.2 Describe how ions can be formed as a result of electron transfer.
Ions are formed when atoms gain or lose electrons
Reaction between Magnesium and Chlorine
• Magnesium configuration: 1s22s22p63s2 or
__ __ __ __ __ __1s 2s 2p 3s
• If Mg loses two electrons then it would end in 2s22p6 and be stable.
• Then magnesium has 12 protons (12+), but only 10 electrons (10-) so it acquires a charge of 2+ and becomes the magnesium ion, Mg2+. In order for it to lose two electrons, something has to gain two electrons
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4.1.2 Describe how ions can be formed as a result of electron transfer.
Ions are formed when atoms gain or lose electrons
Reaction between Magnesium and Chlorine
• Chlorine configuration: 1s22s22p63s23p5 or
__ __ __ __ __ __ __ __ __1s 2s 2p 3s 3p
• If Cl gains one electron then it would end in 3s23p6 and be stable.
• Then chlorine has 17 protons (17+), and 18 electrons (18-) so it acquires a charge of 1- and becomes the chloride ion, Cl-.
• Since chlorine can only gain one electron and magnesium gives up two electrons, magnesium requires two chlorine atoms.
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4.1.2 Describe how ions can be formed as a result of electron transfer.
Reaction between Magnesium and
Chlorine
• Since opposite charges attract, the Mg2+ and the 2 Cl- ions form an ionic bond.
• Determine the formula and name of the compound produced.
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4.1.2 Describe how ions can be formed as a result of electron transfer.
Reaction between Magnesium and
Chlorine
• Since opposite charges attract, the Mg2+ and the 2 Cl- ions form an ionic bond.
• Name: Magnesium chloride
• Formula MgCl2
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4.1.2 Describe how ions can be formed as a result of electron transfer.
• Magnesium gives away 2 electrons; one to each chlorine atom• Now Magnesium has a full 2nd shell and both Chlorines have a full
3rd shell
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4.1.2 Describe how ions can be formed as a result of electron transfer.
Reaction between Potassium and Oxygen
Potassium configuration: 1s22s22p63s23p64s1
• Potassium will lose 1 electron and become the potassium ion K+.
Oxygen configuration: 1s22s22p4
• Oxygen will gain 2 electrons and become the oxide ion O2-.
Two potassiums are needed to combine with one oxygen.
Formula: K2O
Name: Potassium oxide
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Reaction between Aluminum and Bromine
Diagram the bonding between Al and Br,
write the formula, and give the name.
4.1.2 Describe how ions can be formed as a result of electron
transfer.
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Reaction between Aluminum and Bromine
Diagram the bonding between Al and Br,
write the formula, and give the name.
Formula: AlBr3
Name: Aluminum bromide
4.1.2 Describe how ions can be formed as a result of electron
transfer.
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Metal + NonmetalIonic compounds are generally composed of a metal combined
with a nonmetal(s)
Give + TakeThe metal gives it’s valence electrons to the nonmetal so it can fill
it’s valence shell.
Cation + AnionIonic compounds are the attraction between oppositely charged
ions
Electronegativity Difference > 1.7Ionic bonds have electronegativity differences greater than 1.7
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.
0.1 – 1.0
1.1 – 1.7
>1.7
0.0 covalent, nonpolar
covalent, slightly polar
covalent, very polar
ionic
electronegativtydifference
probable type of bond
IONIC BONDINGFormulas & Naming
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Writing formulas for ionic compounds
• Write symbol of cation and then anion
• Add subscripts to balance the charges– calcium bromide
• Ca2+ and Br1- is CaBr2
– potassium sulfide
• K+1 and S2- is K2S
– iron(III) oxide
• Fe+3 and O2- is Fe2O3
“Crisscross” method The ionic charge number of each ion is crossed
over and becomes the subscript for the other ion
Reduce to lowest ratio
Naming Ionic CompoundsNaming Ionic Compounds
• 1. Cation first, then anion
• 2. Monatomic cation = name of the element
• Ca2+ = calcium ion
• 3. Monatomic anion = root + -ide
• Cl = chloride
• CaCl2 = calcium chloride
Naming Ionic Compounds
Examples:
NaCl
ZnI2
Al2O3
sodium chloride
zinc iodide
aluminum oxide
Learning Check
Complete the names of the following binary compounds:
Na3N sodium ________________
KBr potassium ________________
Al2O3 aluminum ________________
MgS _________________________
Names of Variable IonsNames of Variable Ions
Transition metals (except Ag, Zn, Cd) REQUIRE Roman Numerals because they can have more than one possible charge.
I II III IV V VI VII 1 2 3 4 5 6 7
FeCl3 (Fe3+) iron (III) chlorideCuCl (Cu+1) copper (I) chlorideSnF4 (Sn4+) tin (IV) fluoridePbCl2 (Pb2+) lead (II) chloride
Fe2S3 (Fe3+)iron (III) sulfide
Naming ionic compounds• Binary Compounds
– cation is written first, followed by the anion with and –ide ending• Cs2O cesium oxide
• SrF2 strontium fluoride
• CuO copper(II) oxide– oxygen is always 2- and therefore copper will be 2+
• Cu2O copper(I) oxide– oxygen is 2- and therefore needed two copper atoms
with 1+ charge
• Ionic compounds with transition metals• indicate charge after metal with Roman numerals
• The overall charge of the compound should = 0
FeCl2
FeCl3
Cr2S3
• Ionic compounds with transition metals• indicate charge after metal with Roman numerals
• The overall charge of the compound should = 0
FeCl2 2 Cl- -2 so Fe is +2 iron(II) chloride
FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride
Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide
– SnF2 ?
• fluorine is always 1- and therefore tin will be 2+
– tin(II) fluoride
– SnS2 ?
• sulfur is always 2- and therefore tin will be 4+
– tin(IV) sulfide
Learning Check
Complete the names of the following binary
compounds with variable metal ions:
FeBr2 iron (_____) bromide
CuCl copper (_____) chloride
SnO2 ___(_____ ) ______________
Fe2O3 ________________________
SnF2 ________________________
Learning Check
Complete the names of the following binary
compounds with variable metal ions:
FeBr2 iron (II) bromide
CuCl copper (I) chloride
SnO2 tin (IV) oxide
Fe2O3 iron (III) oxide
SnF2 tin (II) fluoride
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4.1.5 State that transition elements can form more than
one ion.MOST transition metals REQUIRE Roman
Numerals because they can have more than one possible charge.
BUT there are ALWAYS exceptions to rules
(in chemistry)• These transition metals only form ONE ion:
Ag+1, Zn+2 and Cd+2
Label these on your periodic table!!
NONO33--
nitrate ionnitrate ion
NONO22--
nitrite ionnitrite ion
Polyatomic IonsPolyatomic Ions
Know these!!!Know these!!!NH4 +1 Ammonium
HCO3 –1 Bicarbonate aka hydrogen carbonate
CO3 – 2 Carbonate
Cr2O7 –2 Dichromate
OH –1 Hydroxide
NO3 –1 Nitrate
MnO4 –1 Permanganate
PO4 –3 Phosphate
SO4 –2 Sulfate
Naming Ternary Compounds
Contains at least 3 elementsThere MUST be at least one polyatomic ion
(it helps to circle the ions)Examples:
NaNO3 Sodium nitrate
K2SO4 Potassium sulfate
Al(HCO3)3 Aluminum bicarbonate
or
Aluminum hydrogen carbonate
Learning Check
Match each set with the correct name:
1. Na2CO3
MgSO3
MgSO4
2 . Ca(HCO3)2
CaCO3
Ca3(PO4)2
Learning Check
Match each set with the correct name:
1. Na2CO3 a) magnesium sulfite
MgSO3 b) magnesium sulfate
MgSO4 c) sodium carbonate
2 . Ca(HCO3)2 a) calcium carbonate
CaCO3 b) calcium phosphate
Ca3(PO4)2 c) calcium bicarbonate
Ionic Nomenclature
Writing Formulas
• Overall charge must equal zero.– If charges cancel, just write symbols.– If not, use subscripts to balance charges.
• Use parentheses to show more than one of a particular polyatomic ion.
• Use Roman numerals indicate the ion’s charge when needed (transition metals)
• Don’t show charges in the final formula.
Ternary Ionic Nomenclature
Sodium SulfateNa+ and SO4 -2
Na2SO4
Iron (III) hydroxideFe+3 and OH-
Fe(OH)3
Ammonium carbonateNH4
+ and CO3 –2
(NH4)2CO3
Learning Check
1. aluminum nitrate
a) AlNO3 b) Al(NO)3 c) Al(NO3)3
2. copper(II) nitrate
a) CuNO3 b) Cu(NO3)2 c) Cu2(NO3)
3. Iron (III) hydroxide
a) FeOH b) Fe3OH c) Fe(OH)3
4. Tin(IV) hydroxide
a) Sn(OH)4 b) Sn(OH)2 c) Sn4(OH)
Formula to NameFe(NO3)3
Choose the correct name for the compound
1. Iron trinitrate
2. iron(I) nitrate
3. iron(III) nitrite
4. iron(III) nitrate
5. none of the above
next problemPolyatomic IonsPeriodic Chart
Name to Formulasodium chlorite
Choose the correct formula for the compound
1. NaCl
2. NaClO
3. NaClO2
4. Na(ClO)2
5. none of the above
next problemPrefixesPeriodic Chart
Mixed Practice!
Name the following:
1. Na2O
2. CaCO3
3. PbS2
4. Sn3N2
5. Cu3PO4
6. MgF2
Mixed Up… The Other Way
Write the formula:
1. Copper (II) chlorate
2. Calcium nitride
3. Aluminum carbonate
4. Potassium bromide
5. Barium fluoride
6. Cesium hydroxide
IONIC BONDINGPhysical Properties
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Properties of Ionic Compounds
• OVERARCHING CONCEPT– Anions and cations are strongly attracted to each
other and difficult to separate an ionic compound– Fuse School video:
https://www.youtube.com/watch?v=TxHi5FtMYKk&list=PLW0gavSzhMlReKGMVfUt6YuNQsO0bqSMV&index=44
Properties of Ionic Compounds• Crystalline “lattice” structure
– repeating arrays of cations and anions
– an ionic lattice of alternating positive and negative ions • Each sodium ion is surrounded by up to 6 chloride ions and
each chloride ion is surrounded by up to 6 sodium ions.
Properties of Ionic Compounds• Hard and brittle
– When pressure is applied, ions of like charge will be forced closer to each other
– The electrostatic repulsion can split the crystal
Properties of Ionic Compounds
• High melting and boiling points – High temperatures are required to overcome the
attraction between the cations and anions
– Therefore, it takes a lot of energy to break apart the electrostatic forces allowing them to melt or boil
• Volatility: how easily a substance turns into a gas– very low as electrostatic forces between cations
and anions is very strong– How often have you seen or heard of gaseous
salt??
• Solubility: the ability to dissolve – Most will dissolve in other polar solvents such as water
– Ions dissociate, or separate from each other, when they dissolve, breaking apart the lattice structure
– ions keep their charges in solution
– Solubility of salt in water: http://www.youtube.com/watch?v=xdedxfhcpWo&feature=related
Conductivity• Electrical conductivity: the ability to allow for
the flow of electrons• Substances must possess Freely Moving
Charged Particles (ions or subatomic particles) – This occurs in…
• molten ionic compounds (+ and – ions can move)– http://www.dynamicscience.com.au/tester/solutions/che
mistry/bonding/bonding5.htm
• ionic compounds in aqueous solution (dissolved in water)
– water pulls apart + and – ions and allows them to move
– This does not occur in solid ionic compounds• ions are bound so tightly to each other that there is nowhere
for electrons to flow
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4.1.8 Describe the lattice structure of ionic compounds.
http://ed.ted.com/lessons/how-atoms-bond-george-zaidan-and-charles-morton
https://www.youtube.com/watch?v=lhC42qxk5kQ
COVALENT BONDINGLet’s share
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IB Topic 4: Bonding4.2: Covalent bonding
Essential Idea: Covalent compounds form by the sharing of electrons.
Nature of Science:Looking for trends and discrepancies – compounds
that contain non-metals have different properties from compounds that contain non-metals. (2.5)
Use theories to explain natural phenomena – Lewis introduced a class of compounds which share electrons. Pauling used the idea of electronegativity to explain unequal sharing of electrons (2.2)
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IB Topic 4: Bonding4.2: Covalent bonding
Understandings:1. A covalent bond is formed by the electrostatic
attraction between a shared pair of electrons and the positively charged nuclei.
2. Single, double, and triple covalent bonds involve one, two, and three shared pairs of electrons, respectively.
3. Bond length decreases and bond strength increases as the number of shared electrons increases.
4. Bond polarity results from the difference in electronegativities of the bonded atoms.
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IB Topic 4: Bonding4.2: Covalent bonding
Applications and Skills:1. Deduction of the polar nature of a covalent bond
from electronegativity values
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4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei
Covalent bonds occur between nonmetals since both want to gain electrons.
Atoms share valence electronsthis is still in order to achieve an noble gas
electron configuration (stable and less energy)
exists where groups of atoms (or molecules) share one or more pair/s of electrons
Each hydrogen now has the electron configuration of the nearest noble gas- helium
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4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei
The electrons in the bond are electrostatically attracted by both nuclei,
so that it forms a directional bond between the two atoms.
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4.2.2 Describe how the covalent bond is formed as a result of
electron sharing.Covalent bonds occur between nonmetals since both want to
gain electrons. They share electrons to achieve a stable configuration
Electron Dot Diagrams
Need 1 electron: H F Cl Br I
Need 2 electrons: O S Se Te
Need 3 electrons: N P As
Need 4 electrons: C Si
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4.2 U1. Electron sharing
Key Terms:Lone pairs: electrons on a dot diagram that are
already paired (also called non-bonding pairs)
Shared pairs (Bond pair): electrons that are shared in a covalent bond
H S
H Shared pair
Lone pair
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4.2.2 A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.
Why should two atoms share electrons?
F F+
7e- 7e-
F F
8e- 8e-
F F
F F
Lewis structure of F2
lone pairslone pairs
lone pairslone pairs
single covalent bond
single covalent bond
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4.2 U2. Bond Types
Double covalent bonds
Sharing 2 pair of electrons
Examples: O2
H2CO
Single covalent bonds
Sharing 1 pair of electrons
Examples:H2 H H H H
HCl H Cl H Cl
CCl4
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4.2 U2. Bond Types
Triple covalent bonds
Sharing 3 pair of electrons
Examples:N2
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8e-
H HO+ + OH H O HHor
2e- 2e-
Lewis structure of water
Double bond – two atoms share two pairs of electrons
single covalent bonds
O C O or O C O
double bonds double bonds
Triple bond – two atoms share three pairs of electrons
N N N N
triple bondtriple bond
or
CCl4 - Covalent
C
Cl
Cl
Cl
Cl
HCl - Covalent
H Cl
MgF2 - Ionic
[ F ]2– [Mg]2+
H2O - Covalent
H O H
NH3 - Covalent
H N H
H
NaCl - Ionic
[ Cl ]– [Na] +
OH– - Covalent
O H
H2 - Covalent
H H
HCl - Covalent
H Cl H Cl
CO2 - Covalent
C OO
Na2O - Ionic
[ O ]2– [Na]2+
H N H
H
H N H
H
OO
OO
O2 - Covalent
OO C
II
II
I2 - Covalent
[ O ]32– [Al]2
3+
Al2O3 - Ionic
NH3 - Covalent
OO O
O OO
O3 - Covalent
H C H
H
H
H C H
H
H
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Modeling Covalent Bonds• Molecular model sets: modeling covalent
bonding. Create the 8 molecules listed on the next slide and complete a data table similar to the following:
Formula (given)
Line Diagram Dot Diagram
You will be making 8 molecules!!
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Modeling Covalent BondsRULES:•All bonding sites must have a bond
• Exception: nitrogen only has 3 bonding sites… the 4th should have a “hat”
•Each bond must be attached to two atoms•Molecules will not be circular
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Modeling Covalent Bonds• 1 bond = 2 shared electrons• Red = oxygen• Black = carbon• White = hydrogen• Blue = nitrogen (3 bonding sites + 1 “hat”)• Green & SilverSilver = halogensgens• H2O ● H2
• C2H4 ● O2
• CO2 ● C6H12
• NH3 ● N2
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4.2 U3. Bond lengths and Strength
The more pairs of Electrons that are
shared between two atoms (bonds) in a molecule will
make the attraction between
the atoms • stronger
• and shorter
Lengthnm
Strength(kj mol-1)
C-O 0.143 356
C-C 0.154 348
C=O 0.121 736
C=C 0.134 657
C C 0.120 908
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4.2 U3. Bond lengths and Strength
So, if you compare a single bond to a triple bond, • Which one is stronger?• Which one is shorter?
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4.2 U3. Bond lengths and Strength
Bond Strength
Triple bond > Double Bond > Single Bond
Strongest Medium Weakest
Bond Lengths
Triple bond < Double Bond < Single Bond
Shortest Medium Longest
Know difference in strength and length between single, double, and triple bonds between two
carbon atoms.
Know the bond length difference between the C and O in the carboxyl group
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4.2 U3. Bond lengths and Strength
Let’s practice!
Draw C3H6 (or CH3CHCH2)
• Identify the longest carbon-to-carbon bond• Identify the strongest carbon-to-carbon bond
• Now try: CH3COOH
• Now try: HCOOCH3
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COVALENT BONDINGBOND Polarity
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4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei
Covalent bonds occur between nonmetals since both want to gain electrons.
Atoms share valence electronsthis is still in order to achieve an noble gas
electron configuration (stable and less energy)exists where groups of atoms (or molecules)
share one or more pair/s of electronsSometimes the sharing of electrons between
atoms is unequal, leading to polarity
Polarity• Covalent bonds can either be nonpolar or polar• Shared bonding electrons pairs are sometimes
pulled (as in a “tug-of-war”) between atoms– Equal sharing is non-polar– Unequal sharing is polar
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4.2.6 Predict the relative polarity of bonds from electronegativity values.
Nonpolar Covalent Bond:Nonpolar bonds form when electrons are shared equally
between two atoms with the same electronegativity values
2.2 2.2 – 2.2 = 0
Nonpolar covalent bonds
• Since atoms in the bond have the same electronegativity values, they pull on the shared pair of electrons equally, so no polarity = nonpolar
• Always the case in diatomic molecules
–BrINClHOF meaning…
•Br2 I2 N2 Cl2 H2 O2 F2
Memorize these!!!
Polar Covalent Bonds• Atoms in the bond pull the shared pair of
electrons unequally since they have different electronegativities
• Results in a dipole because it has two poles use the symbol + or – for areas that are slightly
positive or negatively charged
• More electronegative atoms have a greater attraction for electrons• A number is assigned to each element to
quantify its attraction to a pair of electrons in a shared in bond (example- F is 4.0)
• Polar Covalent Bond: Polar bonds form when electrons are shared unequally between two atoms due to a difference in electronegativity
So why are some BONDS polar?
• Atoms with the higher electronegativity give that “side” of the molecule a slightly negative charge (-)
• Atoms on the “other side” with a lower electronegativity therefore have a slightly positive charge (+)
• The greater the difference in electronegativity, the more polar the bond will be
So why are some BONDS polar?
covalent, non-polar
covalent, polar
ionic
BONDING Practice
Arrange the following BONDS from most to least polar:
a) N–F O–F C–F
a) C–F, N–F, O–F
b) C–F N–O Si–F
b) Si–F, C–F, N–O
c) Cl–Cl, B–Cl, S–Cl
c) B–Cl, S–Cl, Cl–Cl
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4.2 U4. Electronegativity
Electronegativity Difference
Nonpolar covalent bonds have electronegativity difference is 0
Polar covalent bonds have electronegativity differences above 0, but less than 1.7Ionic bonds have electronegativity
differences greater than 1.7http://animatedchemistry.org/?p=528
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4.2.6 Predict the relative polarity of bonds from electronegativity values.
Nonpolar vs Polar Covalent Bonds:
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4.2 U4. Electronegativity
Nonmetal + NonmetalCovalent compounds generally occur between
elements that are closer to each other on the periodic table (with the exception of
hydrogen)
Electronegativity DifferenceCovalent bonds have electronegativity differences
from 0 to 1.7. A bond with electronegativity difference of 1.7 is
classified as covalent, NOT ionic.
0.1 – 1.0
1.1 – 1.7
>1.7
0.0 covalent, nonpolar
covalent, slightly polar
covalent, very polar
ionic
electronegativtydifference
probable type of bond
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Nonpolar Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
Increasing difference in electronegativity
4.2.6 Classification of bonds by difference in electronegativity
0 1.7
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Polarity Practice
Let’s get some
practice!!!
COVALENT BONDINGFormulas & Naming
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Covalent Compound Naming
• Covalent compounds are most often called called molecules
Molecular (Covalent) Namingfor two nonmetals
• Prefix System (binary compounds)
1. Less electronegative atom comes first.
2. Add prefixes to indicate # of atoms. Omit mono- prefix on the 1st element. Mono- is REQUIRED on the 2nd element.
3. Change the ending of the 2nd element to -ide.
PREFIXmono-di-tri-tetra-penta-hexa-hepta-octa-nona-deca-
NUMBER123456789
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Molecular Naming Prefixes
Molecular (Covalent) Namingfor two nonmetals
Exceptions to the prefix method
•H2O = water
– not dihydrogen monoxide
•NH3 = ammonia
– not nitrogen trihydride
ELEMENTS THAT EXIST AS DIATOMIC MOLECULES
ELEMENTS THAT EXIST AS DIATOMIC MOLECULES
Remember:
BrINClHOF
These elements only exist as
PAIRS. Note that when they
combine to make compounds, they
are no longer elements so they are no longer in
pairs!
• CCl4
• N2O
• SF6
• carbon tetrachloride
• dinitrogen monoxide
• sulfur hexafluoride
Molecular Naming: Examples
• arsenic trichloride
• dinitrogen pentoxide
• tetraphosphorus decoxide
• AsCl3
• N2O5
• P4O10
Molecular Formula Examples
Practice 1
Fill in the blanks to complete the following names of covalent compounds.
CO carbon ______oxide
CO2 carbon _______________
PCl3 phosphorus _______chloride
CCl4 carbon ________chloride
N2O _____nitrogen _____oxide
NH3 ______________
watch out for that last one…
Practice 1
Fill in the blanks to complete the following names of covalent compounds.
CO carbon monoxide
CO2 carbon dioxide
PCl3 phosphorus trichloride
CCl4 carbon tetrachloride
N2O dinitrogen monoxide
Practice 2
1. P2O5 a) phosphorus oxide
b) phosphorus pentoxide
c) diphosphorus pentoxide
2. Cl2O7 a) dichlorine heptoxide
b) dichlorine oxide
c) chlorine heptoxide
3. Cl2 a) chlorine
b) dichlorine
c) dichloride
Learning Check
1. P2O5 a) phosphorus oxide
b) phosphorus pentoxide
c) diphosphorus pentoxide
2. Cl2O7 a) dichlorine heptoxide
b) dichlorine oxide
c) chlorine heptoxide
3. Cl2 a) chlorine
b) dichlorine
c) dichloride
COVALENT BONDINGPhysical Properties
123
Covalent• Strong intramolecular forces (between atoms)• Weak intermolecular forces (between
molecules)• Usually liquids or gases at room temp or soft
solid (http://www.rsc.org/periodic-table at 25°C)
– strength of polarity determine mp and bp• greater polarity = higher mp and bp
• Polar substances are soluble in water, but nonpolar are not
• Do not conduct electricity
Solubility of Molecules
• “Like dissolves like”– Polar substances tend to dissolve in polar
solvents , such as water– Non-polar substances do not dissolve in
polar substances like water, but do tend to dissolve in non-polar solvents
Conductivity of Molecules• Substances must possess Freely Moving
Charged Particles (ions or subatomic particles) – Since they do not contain ions, they have no means
to carry an electrical current (flow of electrons)
Type of Bonding
Melting
Point
Boiling Point
Volatility
Electrical Conductivit
y
Solubility in Non-
polar Solvent
Solubility in Polar
Solvent
Ionic Bonding
high high low Yes (molten or aqueous)
No Yes (most)
Polar Covalen
t
high varies high No No No
Nonpolar
Covalent
High High high No (except graphite
and graphene)
No No
128
Characteristics of Covalent Bonds
If the type of bonding is known, physical properties can be predicted
Type of Bonding
Melting Point
Boiling Point
Volatility
Electrical Conductivity
Solubility in Non-polar
Solvent
Solubility in
Polar Solvent
Non-polar Low Low High No Yes No
Polar No No Yes
129
4.5.1 Compare and explain the properties of substances resulting from different
types of bonding
If the type of bonding is known, physical properties can be predicted
Type of Bonding
Melting Point
Boiling Point
Volatility
Electrical Conductivity
Solubility in Non-polar
Solvent
Solubility in
Polar Solvent
Metallic Bonding
Yes No No
Covalent No No No
Giant Covalent
High High Low No (except graphite)
No No
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