Bonding and Naming
Bonding and Naming
1. What is a chemical bond?
2. How do atoms bond with each other?
3. How does the type of bonding affect properties of compounds?
•
•
•
•
•
•
How can all matter in the universe exist from only 92 elements?
Why can you dissolve salt in water but not melt it easily?
Why do the carbon atoms in graphite slide off my pencil lead while
diamonds are forever?
What is the role of carbon in the molecular diversity of life?
How can two toxic and violently reactive elements combine to
produce table salt (essential for life)?
Why is sodium lauryl sulfate found in my shampoo)?
Basics
• Bond
– something that binds, attaches, or restrains
• Chemical Bond
– the force that binds atoms to each other
• Why?
– to Achieve Stability
(Noble Gas Configuration – s2p6)
Octet Rule
• Valence electrons
– Electrons in the outermost energy level of an
atom
• Rules
– Octet—Atoms bond to achieve 8 e- in outer
energy level
– Duet—H bonds to achieve 2 e- in outer
energy level
Lewis Structures
(Electron Dot Diagrams)
• Using X to mean any element, draw Lewis
Structures for elements in Groups
1,2,13,14,15,16,17,18
Electron Dot Diagrams
Group 18
1
2
13
14
15
16
17
Application of Octet Rule to Ionic
Compounds
• Note: The element’s symbol represents the core
electrons.
• Dots represent the valence electrons

Na
1s22s22p6
3s1
Application of Octet Rule
• Ionic Compounds
– Metals go to cations
• Na  Na+ + 1 e-
– Nonmetals go to anions
• Cl + 1 e-  Cl-
• Formation of NaCl
-
[ ] [ ]
Na
+
Cl
Practice
• Draw Lewis Structures for the following
compounds:
• magnesium oxide
• calcium fluoride
• aluminum oxide
Answers
2+
Mg
2-
[O]
Answers
-
2+
F Ca
F
Answers
[O]
[O]
[O]
2-
[Al]
Al
[]
3+
2-
3+
2-
Types of Chemical Bonds
• Three basic types of
bonds
– Ionic
• Electrostatic attraction
between ions
– Covalent and covalent
network
• Shared electrons
– Metallic
• Metal atoms bonded to
several other metal
atoms
Types of Chemical Bonds
• Type of bond  properties of compound
• We will begin with ionic bonding
Ionic Compounds
• also called salts
• held together by attraction between (+) and
(-) charges in a lattice
(electrostatic attraction)
• these charged particles are formed by
donating (and receiving) electrons
• have neutral charge overall: the number of
+ charges = number of – charges
• strength of the bond is amplified through the
crystal structure
Ionic Compounds
• repeating 3-D structure in crystal lattice =
unit cell
Ionic Compounds
• Energy considerations
– Forming a crystal lattice  loss of a lot of PE
• ↑ stability
– “lattice energy” = E released when 1 mol of an
ionic compound is formed from gaseous ions.
Making NaCl
http://www.youtube.com/watch?v=2mzDwgy
k6QM
Making NaCl
Lattice Energy
Definition:
lattice energy = energy released when a crystal
containing 1 mole of an ionic compound is formed
from gaseous ions.
• Related to charge on the ion and distance
between nuclei
 higher charge  stronger pull
 smaller means closer together  stronger pull
Lattice Energies of Some Ionic
Compounds
Compound
Lattice
Energy
(kJ/mol)
Compound
Lattice
Energy
(kJ/mol)
KI
-632
KF
-808
KBr
-671
AgCl
-910
RbF
-774
NaF
-910
NaI
-682
LiF
-1030
NaBr
-732
SrCl2
-2142
NaCl
-769
MgO
-3795
Lattice Energies of Some Ionic
Compounds
• What happens as the ionic radius
increases? (e.g. Check out the Na
compounds or the F compounds)
• What happens as the amount of charge on
a single ion increases? Both ions? (look
at bottom right hand corner of table)
Lattice Energies of Some Ionic
Compounds
• What happens as the ionic radius
increases?
Lattice energy decreases – ions are
further away from each other.
• What happens as the amount of charge on
a single ion increases? Both ions?
Lattice energy increases, especially if both
ions have multiple charges.
Lattice Energies of Some Ionic
Compounds
• Predict whether the lattice energy of CsCl
is larger or smaller than that of KCl.
• Predict whether the lattice energy of Na2O
will be larger of smaller than that of MgO.
• What do you think will be the effect of
lattice energy on melting point?
Lattice Energies of Some Ionic
Compounds
• Predict whether the lattice energy of CsCl is larger or smaller than
that of KCl.
Smaller – Cs+ is larger than K+. (CsCl: -657 kJ/mol, KCl: -701
kJ/mol)
• Predict whether the lattice energy of Na2O will be larger of smaller
than that of MgO.
Smaller – Na+ is 1+, while Mg2+ has double the charge. (Na2O: 2481 kJ/mol)
• What do you think will be the effect of lattice energy on melting
point?
The higher the lattice energy, the more energy it takes to melt it,
therefore melting point increases as lattice energy increases.
Characteristics
•
•
•
•
•
solid at room temperature
hard, brittle
high melting point, boiling point
usually soluble in water
conduct electricity when melted or
dissolved in water (charges can move)
• do not conduct electricity when solid
Dissolving NaCl in Water
Naming Chemical Compounds
Terms for describing how atoms
bond
• Chemical Formula
– Relative #’s of different elements using symbols and
subscripts
– C2H4
• Empirical Formula
– Chemical Formula Reduced to lowest whole #’s of
each element
– C2H4→CH2
• Formula Unit
– Lowest whole number ratios of Ionic Compounds
– NaCl, MgBr2, etc.
Terms for describing how atoms
bond
• Molecular Formula
– Chemical formula of one molecule
– e.g. C2H4
• Structural Formula of Molecule
– Shows how atoms are bonded together in a
molecule H
H
C
H
C
H
Naming Ions
• Monatomic
– 1 atom + or –
• Polyatomic
– Many atoms covalently bonded with an overall + or –
charge (especially anions)
Naming Cations
1. Monatomic Cations
– Use Atomic Symbol + Charge
e.g.
Na+ = sodium Ion
(element’s name + ion)
Naming Cations
2. Polyatomic cations (Memorize these!)
NH4+
Hg22+
ammonium ion
mercury(I) ion
How Do We Know Charges?
• Apply Octet Rule (Groups 1, 2, 16, 17)
• Memorize Ion List
• Use Roman numberals to name elements
with >1 charge
– Copper(I) Ion =Cu+
– Copper(II) Ion=Cu2+
Naming Anions
• monatomic anion:
anions provide the second part of the
name of an ionic compound.
e.g. Cl-
chloride ion
use suffix “ide” + “ion”
Can you name the rest of the monatomic
anions?
Naming Polyatmic Anions
• Polyatomic (with oxygens)
– ite
– ate
NO2- Nitrite
NO3- Nitrate
SO32- Sulfite
SO42- Sulfate
• ClO-=Hypochlorite=least oxygens
• ClO2-=Chlorite
• ClO3-=Chlorate
• ClO4-=Perchlorate=most Oxygens
Naming Polyatmic Anions
• If the anion contains hydrogen, prefix
name with “hydrogen”
– HCO3- = hydrogen carbonate
– HSO4- = hydrogen sulfate
– HS= hydrogen sulfide
Let’s try all the phosphates that contain
hydrogen:
Naming Binary Ionic Compounds
• cation +
Na+ +
anion
Cl-
• NaCl = sodium chloride
Naming Binary Ionic Compounds
• Steps
– Write Ions
Pb4+
O2– Balance Charges (LCM) 4+
2x2– Empirical Formula/Formula Unit PbO2
– Name
Lead (IV) Oxide
Practice
• Calcium
CaCl2 Chloride
• Iron
FeO(II) Oxide
Sulfide
• Iron
Fe2S(III)
3
Ionic Compounds Containing
Polyatomic Ions
• Few positive, many negative polyatomic ions
• Use Atomic Symbols + Charge
– polyatomic cation
• NH4+ ammonium ion
• Hg22+ mercury(I) ion
– polyatomic anion
• OH- hydroxide ion
• CO32- carbonate ion
Polyatomic Ion Formulae
• Write ions
– NH4+
CO32-
• Balance Charges (LCM)
– 2 x 1+
2-
• Write Formula with lowest whole # ratios
– (NH4)2CO3 ammonium carbonate
– Show >1 Polyatomic ion with parentheses
Polyatomic Ion Formulae
• Practice:
aluminum sulfate
magnesium hydroxide
copper(II) acetate
Polyatomic Ion Formulae
• Practice:
aluminum sulfate Al2(SO4)3
magnesium hydroxide Mg(OH)2
copper(II) acetate
Cu(C2H3O2)2
Oxidation #’s
• Oxidation #=Charge of an Ion
– K+ = oxidation # of +1
– O2- = oxidation # of -2
• We use oxidation numbers to figure out
the formulas of ionic compounds.
• The sum of oxidation numbers for the
formulas of an ionic compounds must = 0.
Metallic Bonds
• held together by the attraction of free-floating
valence electrons (-) for the positive ions in a
lattice structure: “electron-sea”
Metallic Bonds
1. What is a regular, repeating three-dimensional arrangement of
atoms called?
2. Do the separate electrons that are shown belong exclusively to a
single atom? What word is used to describe such electrons?
3. Are the electrons shown the only ones actually present? Explain.
4. Why are the central atoms shown as positively charged?
5. How does the number of separate electrons shown for the group
1A metal atoms compare to the number of atoms?
Explain why in terms of valence electrons.
Metallic Bonds - KEY
1.
What is a regular, repeating three-dimensional arrangement of atoms
called? Crystal lattice
2.
Do the separate electrons that are shown belong exclusively to a single
atom? What word is used to describe such electrons?
no, they are delocalized
3.
Are the electrons shown the only ones actually present? Explain.
No, they are valence electrons from the metal atoms.
4.
Why are the central atoms shown as positively charged?
The delocalized (valence) electrons come from the neutral atoms, thus
leaving the atoms with a positive charge – i.e. cations.
5.
How does the number of separate electrons shown for the group 1A metal
atoms compare to the number of atoms?
Explain why in terms of valence electrons. They are equal – have only one
valence electron.
Metallic Bonds
6. How does the number of separate electrons shown for the group 2A
metal atoms compare to the number of atoms?
7. What holds the metal atoms together in such an arrangement?
8. What term is used to describe this model of metallic bonding?
9. How well do metals tend to conduct electricity? How does the
model of metallic bonding account for that property?
10. Do metals tend to be brittle, or are they malleable and ductile?
How does the model of metallic bonding account for that property?
Metallic Bonds - KEY
6.
How does the number of separate electrons shown for the group 2A metal atoms
compare to the number of atoms?
twice as many electrons
7.
What holds the metal atoms together in such an arrangement?
Delocalized electrons are simultaneously attracted to > 1 metal cation.
8.
What term is used to describe this model of metallic bonding?
electron sea
9.
How well do metals tend to conduct electricity? How does the model of metallic
bonding account for that property?
Metals tend to conduct electricity well. The delocalized electrons are not held
strongly by individual atoms and are thus able to move easily throughout the metal.
10. Do metals tend to be brittle, or are they malleable and ductile?
How does the model of metallic bonding account for that property?
malleable and ductile. The delocalized electrons are able to move around the
positive metal core atoms and keep the crystal from breaking.
Characteristics of Metals
• Characteristics of Metals
– high m.p. and b.p.
– Good conductors of electricity and heat
conductors in the solid state
– malleable, ductile
– shiny, reflective, usually gray (or grey)
Characteristics of Metals
As a metal is struck by a hammer, the atoms slide through
the electron sea to new positions while continuing to
maintain their connections to each other.
Warmup – Naming Ionic
Compounds – binary, monatomic
1.Name the following ionic compounds, and
write the pairs of ions that make up these
compounds:
a) KCl
b) SrO
c) PbF2
Warmup – Naming Ionic
Compounds – binary, monatomic
2.Write the chemical formula for the
following ionic compounds, and write the
ions that make up these compounds:
a) lithium bromide
b) barium sulfide
c) chromium (III) oxide
Warmup – Naming Ionic
Compounds, polyatomic ions
3. Write the chemical formula for the
following ionic compounds, and write the
ions that make up these compounds:
a) cesium oxalate
b) calcium hydroxide
c) potassium sulfate
d) ammonium hydrogen phosphate
Warmup – Naming Ionic
Compounds
4. Write the chemical formula for the
compounds made up of the following
ions, then name the compounds:
a)
Cs+, O2b)
Pb2+, Brc)
Fe3+, Fd)
Na+, CO3-
Ionic Bonding
Electron Dot Diagrams
5. How many of each ion will you need to form
each of the following ionic compounds?
Draw electron dot diagrams to demonstrate the
ionic bonding of each compound.
a) lithium and bromine
b) potassium and sulfur
c) aluminum and chlorine
d) gallium and oxygen
e) bismuth (V) and sulfur
Alloys
• Mixture of elements that has metallic
properties (solid solutions)
• Substitutional alloys—atoms of similar sizes
– Brass (Cu +Zn)
– Bronze (Cu + Sn + Pb)
– Pewter (Sn + Sb + Pb)
• Interstitial—much smaller atoms fill spaces
between larger atoms
– Carbon Steel
Alloys
• It is more difficult for layers of atoms to move over
each other, especially in interstitial alloys.
http://www.frankswebspace.org.uk/ScienceAndMaths/chemistry/alloys.htm
Review
• Ionic compounds – crystal lattice of
oppositely-charged ions, held together by
electrostatic charges
• Metals – lattice of positively-charged ions
in a sea of electrons
Covalent Bonding
• Held together by
shared electrons
(covalent bond)
• Both electrons spend
time around each
nucleus but spend
most of their time in
the middle
Diatomic Elements
• Memorize: N O F Cl Br I H
• all exist diatomically in nature (more
stable): N2 O2 F2 Cl2 Br2 I2 H2
• 7- rule: Go to Element #7, travel across to
Group 7A, then down in the shape of a 7.
Another way to remember them:
• Hairogens: H2 (H), N2 and O2 (air), F2, Cl2,
Br2, I2 (halogens)
Diatomic Molecules
• 7 Elements bond with themselves
(increased stability)
– H→H2
– N→N2
– O→O2
– F→F2
– Cl→Cl2
– Br→Br2
– I→I2
Covalent Bonding
• Characteristics of covalently-Bonded
compounds (molecules)
– relatively low m.p. and b.p.
– do not conduct electricity under any
circumstances
– generally not soluble in water, but soluble in
alcohol
Covalent Network Solids
• form covalent bonds in all directions 
continuous network of strong covalent
bonds  no individual molecules
• extremely hard, very high m.p. and b.p.
• nonvolatile, insoluble in all solvents
• brittle, nonconductors of heat and
electricity
• diamond (C), quartz (SiO2)
Compare structures of
networks/lattices:
a) Covalent network (quartz, SiO2)
b) Salt (NaCl)
c) Metal (Cu)
Chemical Bonds
• Metallic ‘bond’
– Metals and Alloys
– ‘Sea’ of electrons
• Ionic bond
– Ionic Compounds or Salts
– Metal + Non-metal: NaCl, MgSO4
– Electrons exchanged between
atoms
• Covalent bond
– Molecules and covalent network
solids
– Non-metals: H2O, CH4
– Electrons shared among atoms
Formation of a Covalent Bond
• Covalent Bond
– When orbitals from two atoms overlap
– 2 electrons of opposite spin in the overlap
– As the amount of overlap ↑, the energy of the
interaction ↓
• When minimum energy is reached, bonding
distance occurs
• Attraction and repulsion of electrons and nuclei are
exactly balanced
– At some distance, nuclei repel, increasing
energy again
Formation of a Covalent Bond
Energy Considerations
• What does “stable” mean? Changes
that lower potential energy are favored.
• In covalent bonds:
• Shared electrons  loss of PE
 stability
• Bond energy = energy required to break a
chemical bond and form neutral atoms
Relationship between bond
length and bond energy in
molecules
• Bond length = average distance between
two bonded atoms (distance of minimum
potential energy)
• As Ebond ↑ , Lengthbond , because the
closer the atoms are, the more attraction
between nuclei and electron clouds.
\ harder to separate.
Naming Molecular Compounds
• 2 systems – prefixes, vs. oxidation
numbers (Honors/AP)
• See p. 248 in textbook for list of prefixes to
memorize
Naming Molecular Compounds
A. Prefixes, roots, suffixes
1. Begin with element with lowest electronegativity.
Nitrogen
2. Add appropriate prefix (unless it is mono-). Dinitrogen
3. End second element with "ide" (as for ionic
compounds…). Oxide
4. Use appropriate prefix for second element. tetroxide
 dinitrogen tetroxide
 N2O4
Naming Molecular Compounds
Practice:
CCl4
CO
CO2
As2S3
P2O5
P4O10
Naming Molecular Compounds
Practice:
CCl4 carbon tetrachloride
CO
carbon monoxide
CO2
carbon dioxide
As2S3 diarsenic trisulfide
P2O5 diphosphorus pentoxide
P4O10 tetraphosphorus decoxide
Oxidation Numbers
(p. 2 of Honors Supplement)
Oxidation number = the charge on an ion
e.g. K+ has an oxidation number of 1+
O2- has an oxidation number of 2We use oxidation numbers to figure out the
formulas of ionic compounds.
The sum of oxidation numbers for the formulas
of an ionic compounds must = 0.
Oxidation Numbers
We can also use oxidation numbers for
molecular compounds, by pretending the
atoms are ions = “apparent charge”.
Oxidation Numbers
The oxidation number of:
• an element in the uncombined state is 0.
• a monatomic ion equals the charge on the ion.
• hydrogen is generally +1; in hydrides, -1.
• oxygen is generally -2; in peroxides, -1.
Oxidation Numbers
The oxidation number of:
• the more electronegative element in a binary covalent compound is
negative, while that of the other element is positive.
• elements other than oxygen and hydrogen in a neutral compound is
such that the sum of the oxidation numbers for all atoms in the
compound is 0.
• elements other than oxygen and hydrogen in a polyatomic ion is
such that the sum of the oxidation numbers for all atoms in the ion
equals the charge on the ion.
Use these rules to assign oxidation
numbers to each element in each of the
given formulas
e.g. H2O H: 2 x +1, O: -2
N2
0
Assign oxidation numbers to each
element:
1. Cl2
2. Cl3. Na
4. Na+
5. KCl
6. H2S
7. CaO
8. H2SO4
Assign oxidation numbers to each
element:
9. NO310. Cr2O7211. NH4Cl
12. NH3
13. NO2
14. CaH2
(calcium hydride)
15. Na2O2
(sodium peroxide)
Assign oxidation numbers to each
element: KEY
1. Cl2
0
2. Cl-
1-
3. Na
0
4. Na+
1+
5. KCl
K: 1+, Cl: 1-
6. H2S
H: 2 x 1+, S: 2-
7. CaO
Ca: 2+, O: 2-
8. H2SO4
H: 2 x 1+, S: 6+, O: 4 x 2-
Assign oxidation numbers to each
element: KEY
9. NO3-
N: 5+, O: 3 x 2-
10. Cr2O72-
Cr: 2 x 6+, O: 7 x 2-
11. NH4Cl
N: 5+, H: 4 x 1-, Cl: 1-
12. NH3
N: 3+, H: 3 x 1-
13. NO2
N: 4+, O: 2 x 2-
14. CaH2
(calcium hydride)
15. Na2O2
(sodium peroxide)
Ca: 2+, H: 2 x 1-
Na: 2 x 1+, O: 2 x 1-
Oxidation Numbers - warmup
• Write the oxidation numbers for all elements in
the following compounds:
P2O5
NO2NO3HNO3
K2CrO4
K2Cr2O7
Oxidation Numbers - warmup
• Write the oxidation numbers for all elements in
the following compounds:
P2O5 P: 2 x 5+, O: 5 x 2NO2- N: 3+, O: 2 x 2NO3- N: 5+, O: 3 x 2HNO3 H: 1+, N: 5+, 3 x 2K2CrO4 K: 2 x 1+, Cr: 6+, O: 4 x 2K2Cr2O7 K: 2 x 1+, Cr: 6+, O: 7 x 2-
Using Oxidation Numbers to
Name Molecular Compounds
1. Name the following molecular
compounds using oxidation numbers:
a) CCl4 carbon (IV) chloride
b) CO
c) CO2
d) P2O5
e) PCl5
f) SO2
Using Oxidation Numbers to
Name Molecular Compounds
1. Name the following molecular
compounds using oxidation numbers:
a) CCl4 carbon (IV) chloride
b) CO
carbon (II) oxide
c) CO2
carbon (IV) oxide
d) P2O5 phosphorus (V) oxide
e) PCl5 phosphorus (V) chloride
f) SO2
sulfur (IV) oxide
Using Oxidation Numbers to
Name Molecular Compounds
2.
Write formulas for the following
molecular compounds:
a) carbon(IV) iodide
Cl4
b) sulfur(VI) oxide
c) nitrogen(IV) oxide
d) arsenic(III) sulfide
e) phosphorus (III) fluoride
Using Oxidation Numbers to
Name Molecular Compounds
2.
Write formulas for the following
molecular compounds:
a) carbon(IV) iodide
Cl4
b) sulfur(VI) oxide
SO3
c) nitrogen(IV) oxide
NO2
d) arsenic(III) sulfide
As2S3
e) phosphorus (III) fluoride
PF3
Compounds That Become Acids
When Dissolved in Water
• General Formula: HX
H+ Xmonatomic
or
polyatomic
anion
Compounds That Become Acids
When Dissolved in Water
Three Rules:
1. When X ends in “ide”
(e.g. chloride, cyanide)
 “hydro_______ ic acid”
e.g. hydrochloric acid,
hydrocyanic acid
Compounds That Become Acids
When Dissolved in Water
Three Rules:
2. When X ends in “ite”
(e.g. chlorite, sulfite)
 “______ous acid”
e.g. chlorous acid
sulfurous acid
Compounds That Become Acids
When Dissolved in Water
Three Rules:
3. When X ends in “ate”
(e.g. chlorate, sulfate)
 “______ ic acid”
e.g. chloric acid
sulfuric acid
Compounds That Become Acids
When Dissolved in Water
Your turn:
HBr
HNO2
HNO3
Compounds That Become Acids
When Dissolved in Water
Your turn:
HBr
hydrobromic acid
HNO2
nitrous acid
HNO3
nitric acid
Warmup
Write formulas for the following molecular
compounds:
• sulfur trioxide
• phosphorus pentachloride
• nitrogen dioxide
• tetraphosphorus decoxide
• oxygen difluoride
Warmup
• Write Lewis structures (dot diagrams) for the
following elements:
• carbon
• hydrogen
• fluorine
• sulfur
• nitrogen
• oxygen
• phosphorus
• bromine
Molecules and Lewis Structures
Lewis structures show
• All atoms in the molecule
• How atoms are connected (# of bonds)
• Any unshared electron pairs (lone pairs)
Making Lewis Structures
1. Find total # of valence e-’s for ALL atoms
(for ions, consider the charge).
2. Write atom symbols, beginning with the
central atom:
a) Carbon in center – CH4
b) Most electropositive in center – SO42c) Non metal (other than H or O) in center –
H2PO4d) Hydrogen and oxygen are usually on the
outside
Making Lewis Structures
3. Add valence e’-s to each atom,
beginning with the bonding electron for
each atom added to the central atom.
4. Shared e- pairs represent bonds and
are counted in the valences of both
elements.
5. Adjust so that each element has 8
valence electrons (2 for H), note
exceptions…
Lewis Structure Practice
(Electron Dot Diagrams)
Molecules
CH3Br
BrI
H2S
PH3
Ions
ClOSO42H2PO4NH4+
Hydrocarbons
C2H6
C2H4
C2H2
C6H6
Warmup – naming ionic
compounds
How many of each ion is in each of the
following compounds?
a) AlBr3
b) PbCl4
c) RbNO3
d) MgSO4
e) K3PO4
Warmup – naming ionic
compounds
How many of each ion is in each of the
following compounds?
a) AlBr3 1 Al3+, 3 Brb) PbCl4 1 Pb4+, 4 Clc) RbNO3 1 Rb+, 1 NO3d) MgSO4 1 Mg2+, 1 SO42e) K3PO4 3 K+, 1 PO43-
Warmup – naming ionic
compounds
• Write the formula for each of the following
compounds:
a) ammonium phosphate
b) cesium oxide
c) copper(I) fluoride
d) silver nitride
e) beryllium nitrate
Warmup – naming ionic
compounds
Write the formula for each of the following
compounds:
a) ammonium phosphate (NH4)3PO4
b) cesium oxide Cs2O
c) copper(I) fluoride CuF
d) silver nitride Ag3N
e) beryllium nitrate Be(NO3)2
Exceptions to the Octet Rule
Central atom has
< 8 valence e-’s
Central atom
has > 8 e-’s
BeF2
BF3
PF5
SF6
XeF4
Answers
e
e
e
ee
C 4
Br 7
H 1
H 1
H +1
14 e-
CH3Br
Answers
e
Cl 7
O6e
+ 1e
14 eClO
Answers
7
e
Br + 7
e
I
14 e-
BrI
Answers
6
6
6
6
6
+2
e
ee
ee
e-
2SO4
Warmup – Lewis Structures
Draw Lewis Structures for the following
molecular compounds:
• carbon tetrachloride
• sulfite ion
• ammonia (NH3)
• water
• chlorine (Cl2)
• oxygen (O2)
• nitrogen (N2)
Molecular Geometry
From 2D Lewis Structures to 3D:
Valence Shell Electron Pair Repulsion (VSEPR Theory)
The influence of Unshared Pairs
on Geometry
When describing the shape of a molecule, consider
the arrangement of the atoms,
which is influenced by the unshared (lone) pairs
Molecular Geometry
Terminology
• Electron domain = region where electron pairs reside
= #bonded pairs + #lone pairs
• Bonding domain = bonded pair
• Nonbonding domain = lone pair
• Annotate the diagrams in your notes with these terms
VSEPR Theory – Determining
Molecular Shapes
2-3 electron domains
4 electron domains
5 electron domains
5 electron domains (Honors)
6 electron domains
Back to Activity
• Observe the models, then draw and name the
shapes in table B
Note: double and triple bonds = 1 lone pair
• Next, you will build models of the compounds in
table A, then draw and name the structures
• We will discuss polarity later this week.
Hybridization (Honors)
• Atomic orbitals are mixed  new, identical
hybrid orbitals
• helps to explain VSEPR
• # of hybrid orbitals = # atomic orbitals
mixed, including lone pairs
Hybridization
Hybrid Orbital Summary
Electron-domain geometry must be known before
hybridization is assigned.
• To assign hybridization:
• Draw a Lewis structure.
• Assign the electron-domain geometry using VSEPR
theory.
• Specify the hybridization required to accommodate the
electron pairs based on their geometric
arrangement.
• Name the geometry by the positions of the atoms.
Warmup
1. What is meant by the term “polar”?
Use an example that makes sense to you.
2. Using the PT of theElectronegativities, calculate the
difference in electronegativities between
a) Na and Cl
b) Br and Cl
c) B and Cl
3. Identify the bonds in Q2 as ionic, polar covalent or
nonpolar covalent.
Warmup
1. What is meant by the term “polar”?
Use an example that makes sense to you.
2. Using the PT of theElectronegativities, calculate the
difference in electronegativities between
a) Na and Cl (2.1) ionic
b) Br and Cl (0.2) NP covalent
c) B and Cl
(1.0) P covalent
3. Identify the bonds in Q2 as ionic, polar covalent or
nonpolar covalent.
Ionic, polar covalent or
nonpolar covalent?
• The nature of the bond between any two
atoms is determined by the difference in
their electronegativities (see chart on ho)
• The greater the difference, the more ionic
the bond (on a continuum)
ionic
4.0
covalent
0
0.4
1.7
polar (P)
NP
Ionic range: 4.0 - 1.8
Cl – Na: 3.0 - 0.9 = 2.1 (ionic)
Polar covalent range: 1.7 - 0.5
Cl – Al: 3.0 - 1.5 = 1.5 (polar covalent)
Nonpolar covalent range: 0.4 - 0.0
Cl – Br: 3.0 - 2.8 = 0.2 (nonpolar covalent)
ionic
4.0
covalent
Cl - Na
0
0.4
1.7
polar (P)
NP
Cl - Al
Cl - Br
Ionic, polar covalent or
nonpolar covalent?
• ionic bonds - the less electronegative
atom donates 1 or more e-’s to the more
electronegative atom
• covalent bonds - e-’s from both atoms are
shared
o Polar covalent bonds - e-’s are shared, but
are not shared equally between two atoms
o nonpolar covalent bonds- e-’s are shared
equally between two atoms
Polar vs. Non-polar:
a matter of symmetry
Determining Molecule Polarity
1. Draw electron dot diagram (Lewis structure) of
the molecule, let lines represent bonds
2. Compare the electronegativities of each of the
bonded atom pairs
3. Determine whether each bond is polar (P) or
nonpolar (NP)
4. Draw an arrow parallel to each bond directed
towards the more electronegative atom
5. Is the molecule symmetrical? (If you cut it
through the xy, xz and yz planes, would it split
into mirror images that are identical?)
•
•
Yes - molecule is NP
No - molecule is P
Determining Molecule Polarity
Examples
HCl
CCl4
CH3Cl
NH3
BF3
H2O
Back to Activity
• Last column of Data Table B
• Last two columns of Data Table A
Sigma Bonds and Pi Bonds
(Honors) canceled this year
• Sigma (s) bonds: electron density lies on
the axis between the nuclei.
• All single bonds are s bonds.
Sigma Bonds and Pi Bonds
(Honors) canceled this year
• What about overlap in multiple bonds?
• Pi (p) bonds: electron density lies above and
below the plane of the nuclei.
• A double bond consists of one s bond and
one p bond.
• A triple bond has one s bond and two p
bonds (above and below the plane of the nuclei;
in front of and behind the plane of the nuclei).
Sigma Bonds and Pi Bonds
canceled this year
Formation of a Double Bond
canceled this year
Formation of a Double Bond
canceled this year
Formation of a Double Bond
canceled this year
Triple Bond
canceled this year
Warm Up
For each of the following three compounds:
CBr4, CH2O, and CCl2F2
a)Draw the Lewis structures
b)Name the shape
c)Calculate the polarity of each bond
d)Predict the polarity of each molecule
e) (Honors only) What is the hybridization of
the central atom?
Intermolecular forces
(van der Waals forces)
• Dispersion (London) forces
• Dipole force
• Hydrogen bonding
1. Dispersion (London) forces
• named after Fritz London, 1900-1954)
a. weakest intermolecular force
b. results from the constant motion of electrons
 uneven distribution of electrons at any particular moment:
“temporary dipole” which may  dipole in nearby molecule.
c. acts on all molecules all the time
d. only intermolecular force acting among noble gas atoms
and nonpolar molecules
1. Dispersion (London) forces
e. with number of electrons:
note m.p., b.p.
e.g. halogens
F2, Cl2 gases at room T
Br2
liquid at room T
(more e-’s than F2 and Cl2)
I2
solid at room T (most e-’s)
1. Dispersion (London) Forces
2. Dipole Force
(polar molecules)
a. the attraction between two polar
molecules:
(-) end of one polar molecule attracts the
(+) end of another polar molecule
b. more polar  stronger dipole force
c. closer together  stronger dipole force
2. Dipole Force
(polar molecules)
3.
Hydrogen bonding
a. always involves H
usually involves O, F or N (small, high
electronegativity)
b. strongest intermolecular force
How strong? 5% of the strength of a
covalent bond
c.  higher b.p. and higher viscosity
e.g. H2O
3.
Hydrogen bonding
Comparing Bond Types
Properties
Description of
bond
General
appearance
Malleable vs.
brittle
Conduct
electricity?
Under what
circumstances
?
m.p., b.p.
State of
matter at
room temp?
Ionic: salts
Metallic
Covalent:
molecules
Covalent
network
Comparing Bond Types
Properties
Ionic: salts
Metallic
Covalent:
molecules
Covalent
network
Description of
bond
Electrostatic
attraction
Sea of
electrons
Shared
electrons
Shared
electrons in
network
General
appearance
crystalline
Shiny, prob.
gray
Range of
colors
Range of
colors
Malleable vs.
brittle
brittle
malleable
brittle
Prob. brittle
Conduct
electricity?
yes
yes
no
no
Under what
circumstances
?
dissolved in
H2O or
molten
all
N/A
N/A
m.p., b.p.
High
High
Low
Very high
State of
matter at
Always solid
Mainly solid
Gas/solid
solid
Comparing ALL bond types:
Which is stronger?
covalent network > metallic > ionic > covalent (molecules) > H bond > dipole > dispersion
e.g. Compare melting points:
SiO2 > Fe > NaCl > C12H22O11 > H2O >
HCl
>
H2
sand > iron > salt > sugar > ice > hydrogen chloride > hydrogen gas