Chemical Bonds
Standard 2: chapter 7, 8, 9
Vocabulary
leave enough space for definition AND an example
1. Metallic bond
2. Alloy
3. Ionic bond
4. Cation
5. Anion
6. Crystal
7. Covalent bond
8. Polar covalent bond
9. Diatomic molecule
Siddall.
10. Electron dot structure
Chemistry.
Standard 2d: Intermolecular forces
Solids:
• Particles are
strongly attracted
to each other
• High melting point
• Particles vibrate in
place
• fixed shape
• fixed volume
Liquids:
• Particles are
weakly attracted
to each other
• Low melting point
• Particles move
around each
other freely
• no fixed shape
• fixed volume
Gases:
• Particles have no
attraction to each
other
• Extremely low
melting point
• low boiling point
• Particles move
rapidly and
randomly
• no fixed shape
• no fixed volume
study question 1
• If a substance has no fixed shape it could
be: _________ or _________
• If a substance has no fixed shape and no
fixed volume it would be: ___________
Physical state: The state of a material
depends on the balance between:
1. the kinetic energy of the particles.
2. the attractions between particles
•
•
•
Kinetic energy > attractions = gas
Kinetic energy < attractions = liquid
Kinetic energy << attractions = solid
study question 2
• If the kinetic energy of particles in a
substance is much greater than the forces
between particles, the substance is a
__________
Phase changes.
• Melting a solid or evaporating a liquid
requires energy to overcome the
forces holding the particles together.
• Freezing a liquid or condensing a gas is
caused by removing energy so attractive
forces between particles dominate
Condensing
liquid
solid
evaporating
gas
Energy added/absorbed
intermolecular attractions are
overcome
melting
freezing
Energy released/removed
intermolecular attractions take
over
Physical state
study question 3
Which processes occur when energy is
removed from a substance?
Honors Only:
Volatility = degree of change.
• A substance with high volatility will change easily
from solid to liquid or liquid to gas. For example:
o carbon dioxide, oxygen are very volatile compounds
(gas at room temperature)
o Water is less volatile (liquid at room temperature)
o Iron, salt are considered non-volatile compounds
(solids at room temperature)
Study question 4
•
List the following compounds as
‘extremely volatile’, ‘somewhat volatile’ or
‘non-volatile’. Explain each choice.
1. Methane (natural gas)
2. Alcohol
3. Calcium carbonate (rocks)
Standard 2a: Types of Bonds
Type of
atoms
Metallic
bond
Metals
Ionic bond
Metals &
non-metals
Covalent
bond
Non-metals
Electrons
Shared
between
atoms
Transferred
from one
atom to
another
Shared
between
atoms
Solid/
liquid/gas
Solid
Solid
Gas, liquid or
solid
study question 5
• What type of bonds are formed when nonmetal atoms share electrons?
Bonding in Metals
Metal atoms share valence electrons.
 Atoms are very close together
ஃ Metals are solid compounds
 Electrons move around (sea of electrons)
ஃ Metals conduct electricity
ஃ Metals are malleable.
 ex: lead
Alloy: mixture of
different metals with
specific properties
superior to individual
metals.
e.x. steel frame
construction.
study question 6
• Explain why metals are solid and why
they conduct electricity.
Standard 2c: Ionic bonds.
An Ionic bond is formed between metal
and non-metal atoms
• Each atom gains or loses electrons in
order to form an octet
• An ion is a charged particle
• A cation = positive ion = a Metal.
o e.x. Na+, Ca2+, Al3+
• An anion = negative ion = a Non-metal
o e.x. Cl-, S2-, P3-
study question 7
For the compound KF:
• Which atom is the
cation?
• Which atom is an
anion?
Crystal Lattice Structure
• All ionic compounds form a crystal
lattice structure
o formed by a very large network of
electrostatic attractions (positive and
negative ions attracted to each other).
o Lattice energy: The energy needed to
break the electrostatic attractions holding
the lattice structure.
• Ionic compounds are always solid
because of the strong electrostatic
attractions between ions
Crystal Lattice Structure.
+
+
+
+
+
study question 8
Why do ionic compounds form crystal
lattice structures?
Properties of ionic compounds
Electrostatic attractions are very strong
therefore ionic compounds:
1. are solids at room temperature.
2. have very high melting points.
study question 9
Which of the following are solid at room
temperature?
• CaO
• CO
• NO2
• Na2O
naming ionic compounds
e.x. Na2O
1. Use cation name
2. modify anion name (ide)
3. Do NOT use prefixes
Name = sodium oxide
More examples:
CaF2 = calcium fluoride
K2O = potassium oxide
study question 10
Name the following
1. MgCl2
2. Al2O3
3. NaBr
Weird things
• Polyatomic ions: act as one charged
particle in an ionic bond
• Anion names are not modified for
polyatomic ions
 Example: NH4OH
= Ammonium hydroxide
 Example: Al(NO3)3
= aluminum nitrate
study question 11
Name the following
1. NaOH
2. K2SO4
3. Mg(NO3)2
Writing ionic formulas from names
• example:
 Sodium hydroxide (Na+ and OH-)
 Charges must cancel out = NaOH
 example:
 Magnesium hydroxide (Mg2+ and OH-)
For each Mg2+ there must be 2 x OH- =
Mg(OH)2
note: use parenthesis only when showing
more than one polyatomic ion
study question 12
Write formulas for the following
compounds:
1. Aluminum hydroxide
2. Potassium oxide
3. Magnesium nitrate
Standard 2b: covalent bonds
Covalent (molecular) compounds:
• Formed when non-metal atoms bond. O C
O
• Bonds between atoms are strong
• But many covalent compounds are liquids or
gases because molecules are not strongly
attracted to each other
• ex: H2O, CO2
• Properties:
– many covalent molecules have very low melting
points and high volatility
– Many covalent molecules are gases or liquids
H
O H
study question 13
Identify the covalent compounds:
1. CO2
2. CaO
3. MgCl2
4. CCl4
Naming covalent compounds.
•
e.x. CO2 = 1 carbon + 2 oxygen
• Name = carbon dioxide
1. Modify name of second atom (ide).
2. Add pre-fix to indicate number of
atoms.
Prefixes
1.
2.
3.
4.
5.
mono
di
tri
tetra
penta
6. hexa
7. hepta
8. octa
9. nona
10.deca
• Examples:
 CCl4 =
 Carbon tetrachloride.
 N 2O 3 =
 Dinitrogen trioxide.
• Exception: The ‘mono’ prefix is usually
omitted from the first atom
 NO = nitrogen monoxide
Diatomic molecules.
= molecules formed from 2 atoms







H2
N2
O2
F2
Cl2
Br2
I2
hydrogen
nitrogen
oxygen
fluorine
chlorine
bromine
iodine
NOTE:
Nitrogen = N2
N2 is a molecule
N = a nitrogen atom
News flash:
you must
know these
study question 14
Name the following:
1. CO
2. CO2
3. Cl2
4. NO
5. N2O
Standard 2e: Lewis Dot Diagrams
diagrams show:
 Chemical symbol
 Valence electrons
• Each atom has 4 valence electron
orbitals (one s orbital and 3 p orbitals)
• Each orbital can hold 2 electrons.
• Electrons like to be alone.
• electrons pair up if necessary.
Examples of Lewis Dot diagrams
e.x. Nitrogen atom
•
•• N •
•
Electron orbitals
e.x. Sulfur atom
•
•• S •
••
study question 15
1. Draw a Lewis Dot Diagram for an oxygen
atom
2. Draw the Lewis Dot Diagram for a
chlorine atom
Creating an Octet
Non-metal atoms form covalent bonds in
order to share electrons and create an octet.
Example: H2
• Each hydrogen has one electron.
 Each hydrogen needs two electrons (like He).
•
H •H
H H
Covalent bond = 2 shared electrons
(show in between atoms)
study question 16
 Draw the Lewis dot diagram
for a chlorine molecule (Cl2)
e.x. Oxygen molecule (O2).
• Oxygen atom:
 needs 2 more electrons (to have an octet)
 forms 2 bonds (using 2 unpaired electrons)
••
•O
•
••
•
•
•• O •
•
••
••
••
••
O •• O
Double bond
One bond
•The oxygen molecule still has the
same total number of electrons
•But each atom ‘thinks’ it has an
‘octet’.
study question 17
1. Draw the Lewis dot diagram
for N2
Rules for Dot Diagrams:
1. Count total number of valence electrons
for all atoms.
2. Determine number of bonds needed for
each atom.
3. Allocate unpaired electrons to bonds.
4. Allocate unshared (paired) electrons to
orbitals so each atom has an octet.
5. Re-count total number of electrons in
diagram.
e.x.: CH4 (methane molecule)
• Carbon
• Hydrogen
 has 4 electrons
 has 1 electron
 needs 4 electrons
 needs 1 electron
 forms 4 bonds.
 forms 1 bond.
•
•
•
C
•
•H •H
•H •H
Total number of electrons = 8
Single bond.
H
•
•
H •• C •• H
••
H
octet
Helium electron
configuration
• The total number of electrons did not
change.
• Each atom ‘thinks’ it has an octet.
study question 18
• Draw the correct Lewis Dot
Diagram for H2O
Danger!
HONORS STUDENTS ONLY
BEYOND THIS POINT.
VESPR Theory
Valence Shell Electron Pair Repulsion
• An atom with no unshared electrons forms four
bonds:
4 single bonds = tetrahedral
2 single bonds & 1 double bond = trigonal planar
2 double bonds = linear
1 single bond & 1 triple bond = tetrahedral
Study question 19
• Draw the lewis dot diagram for CF4 and
determine the shape of the molecule
• An atom with 1 unshared pairs of electrons
forms 3 bonds:
– The electron pair takes up the space of an
orbital
3 single bonds = trigonal pyramidal
NOT trigonal planar
• An atom with 2 unshared pairs of electrons
forms 2 bonds:
2 single bonds = bent
NOT linear
Study question 20
• Draw the lewis dot diagram for water and
determine the shape of the molecule
2g: electronegativity and bonds.
Polar Covalent Bonds: Formed when atoms share
electrons unequally.
• Electrons have a negative charge
• Atoms with high electronegativity attract the
electrons in a bond. This creates a dipole within a
molecule.
Dipole = charge difference
• Polar molecules (molecules with a dipole) are
attracted to each other like magnets.
• Molecules that are extremely polar are usually
liquids because the molecules are close together
• Non-polar molecules are usually gases
• e.x. Hydrogen fluoride
Slightly
positive
••
H F
••
δ-
••
••
δ+
Slightly
negative
Arrow points to negative part of molecule
tail shows ‘+’ for positive
study question 21
• Draw the Lewis dot diagram for HCl and
use symbols to indicate the dipole
Calculating polar or non-polar bonds.
• Non-polar bond: difference in
electronegativity < 0.5
• Polar bond: difference in en ≥ 0.5
• Ionic bond: difference in en > 2
 Example: H-F
4.0 – 2.1 = 1.9 = polar bond
 Example: H-C
2.5 – 2.1 = 0.4 = non-polar bond
study question 22
•
Label the following molecules polar or
non-polar
1. CO
2. CN
3. O2
Hydrogen bonding (this is NOT a bond)
• A special type of intermolecular attraction
• Hydrogen bond = intermolecular attraction
between a hydrogen nucleus in one
molecule and a very electronegative atom*
in another molecule.
*very electronegative atoms = O, N, F, Cl
• Hydrogen bonding results in
low volatility
• For water there are some
unique properties:
 Lower density of solid water
 High boiling point
study question 23
• Which of the following molecules might
hydrogen bond? CH4, CH2O, HF, H2O, H2S
Van der Waals forces
• Intermolecular forces
• Exist between all types of atoms/molecules
• Electrons orbiting atoms become unevenly
dispersed creating a temporary dipole
• Also called London Dispersion Forces
Study question 24
• Which atoms would probably create
stronger Van der Waals forces? Small
atoms or large atoms. Why?
• Would Van der Waals forces make a liquid
more volatile or less volatile?