Part 1: Writing Chemical Compounds
Chemical Formula
 short hand method of indicating a ratio of atoms in a
compound.
 identifies the atoms in a compound
 the less electronegative element is listed first in the
formula
 for covalent compounds, the formula tells you how
many atoms of each type in the molecule.
What does this look like?
Valence or Oxidation Numbers
 describes how many electrons from an atom are used
in bond formation
 if its an ionic bond


indicates how many electrons are donated ( + )
indicates how many electrons are received ( - )
 if it’s a covalent bond it indicates how many electrons
are contributed to the bond as if the electrons were
completely removed or gained.
Oxidation Numbers
 An oxidation number is the charge an atom would
have if the electron pair that is shared between two
atoms belonged entirely to the more electronegative
atom.
Rules for Oxidation Numbers
The following rules will help to assign oxidation numbers. The rules are listed in Priority
Sequence.
 free atoms (0)
Ex. Al(s)
 atoms bound to eachother. (0) Ex. Cl2, H2 .....
 monoatomic ions have the same oxidation number as the charge number
Ex. Cl1- (-1), Mg2+ (+2), .......
 F is always (-1)
 O is almost always (-2) -except in peroxides
 H is almost always (+1) -except for metallic hydrides
 Assign oxidation numbers to the most electronegative atom first
Rules Continued
Group I (+1)
Group II (+2)
Al (+3)
Group VII (-1)
Group VI (-2)
Group V (-3)
When forming binary compounds with metals.
nonmetals only
In general, if two atoms form an ionic bond, the valence tells
you the charges on the ions that are formed. If a covalent
bond is formed, the valence tells you how many electrons
the atoms contribute to the covalent bond.
Writing Chemical Formulas Using
Valences
Use the Zero Sum Rule (for neutral compounds only)
 algebraic sum of oxidation numbers is zero
 for charged polyatomic ions, the sum of oxidation numbers
equals the charge of the ion.
 Example: KF
 Each potassium ion has a charge of +1 and each fluorine ion
has a charge of -1. Because there is one of each ion in the
formula, the sum of the valences is zero.
Polyatomic Ions
 ions made of more than two atoms
 are charged molecules
 very strong covalent bonds keep these ions together
and react as a single inseparable ion.
Need to Memorize...
Ammonium NH4+
Chlorate ClO3-
Acetate C2H3O2-
Sulfate
Hydroxide OH-
Phosphate
Nitrate NO3-
Cyanide
Carbonate
CO32-
SO42PO43CN-
Writing Chemical Formulas
Writing Chemical Formulas
Writing Chemical Formulas
Part 2: Naming Chemical Compounds
Chemical Nomenclature
 Each chemical compound has been given a name.
Some have a trivial or common name, H2O is known
as water.
 These names have been used for centuries. However all
chemicals have been given a name using Systematic
Nomenclature.
 uniquely describes it s compound
 the name is derived from its chemical formula
 from a name a formula can be determined

ex. NaCl common name: salt; (sodium chloride)
H2O common name: water (hydrogen oxide)
Positive Monoatomic Ions
 the name of the positive monoatomic ion is the same as the element
name

Ex. Li+ -lithium
Stock System:
if the metal has more than one oxidation number, a Roman Numeral is
used: Sn2+ -Tin(II)
Sn4+ Tin(IV)
Old Method:
the old name along with the suffix -ous or -ic is used to indicate the
oxidation number.
Sn2+ -stannous
Sn4+ -stannic
Memory Aid:
-ous -indicates the lower oxidation number
-ic -indicates the higher oxidation number
Negative Monoatomic Ions
 add -ide to the end of the element name.
 Ex. chlorine (Cl) becomes chloride (Cl- )
Binary Compounds Containing a Metal and
a Nonmetal
 consist of only 2 different elements
 the element with the more positive oxidation number is
written first.
 Ex. NaCl and not ClNa
Binary Compounds Containing Hydrogen
and Another Element
 compounds of hydrogen and nonmetals from group
VI and VII are called:
 hydrogen _______ide

ex. HBr is call hydrogen bromide
 compounds of hydrogen and a metal, the metal is
written first and are called:
 ________________ hydride

ex. NaH is called sodium hydride
Binary Compounds Containing 2
Nonmetals
 this method uses prefixes to
indicate the number of
nonmetal elements present
__________name _______name
ex. As2S3 is called diarsenic
trisulfide
NO is called nitrogen monoxide
N2O5 is called dinitrogen pentoxide
Naming Chemical Compounds Containing More
than Two Elements
Polyatomic ion + nonmetal ion
 name the positive portion first then the negative ion
 negative monoatonic ions end in –ide
 ex. NH4Cl is called ammonium chloride
Cation + polyatomic ion
 oxygen containing polyatomic ions end in either –ate or –ite.
 an –ite polyatomic ion contains one less oxygen than does an-ate
polyatomic ion.
(Memory aide: memorize the –ate form of ions)
 ex. SO42- is sulfate and . SO32- is sulfite
Group VII Polyatomic Ions have
Additional Forms
 ClO ClO2 ClO3-
 ClO4-
hypochlorite
chlorite
chlorate
perchlorate
Nomenclature of Acids
 Acid molecules contain hydrogen atoms that are easily removed
when dissolved in water. These ionizable hydrogen atoms appear
at the beginning of the chemical formula.
Binary Acids
 hydrogen and a nonmetal element from group VI or VII is called:
hydrogen ______ide.
 when this compound is dissolved in water, it becomes an acid
called:
hydro________ic acid
ex. HCl (g) + H2O (l) →
hydrogen chloride
HCl (aq)
hydrochloric acid
Oxyacids
 hydrogen and a negative polyatomic ion is called:
hydrogen polyatomic ion
ex. H2SO4 (g) hydrogen sulfate
H2SO3 (g) hydrogen sulfite
 when these compounds are dissolved in water are called:
ex.
H2SO4 (aq) sulfuric acid
H2SO3 (aq) sulfurous acid
Note: In acid form –ate polyatomic ions become -ic acid
-ite polyatomic ions become -ous acid
Proof for existence of
intermolecular forces
Gases have large distances between their particles, so
short range forces between molecules are insignificant.
ON COOLING AND COMPRESSING
These weak forces become significant as the distances are
reduced and their magnitude can now make a difference.
So gases liquefy. So (GL) i.e. Liquefy
ON further COOLING AND COMPRESSING
The distances are further reduced.
Forces become more significant and hence the liquid
solidifies.
So (LS ) i.e. solidifies
Types of Intermolecular forces
Van der waals’ forces
• Arise due to random
movement of electrons
leading to the
formation of
instantaneous dipole
and hence induced
dipoles in molecules
• Their strength depends
upon the molar mass of
the molecule.
• These forces are
effective over a short
range.
• They are dependent on
the surface area of the
molecule
Dipole-dipole forces
• These occur due to
electrostatic attraction
between molecules with
permanent dipoles.
• They are significantly
stronger than van der
waals’ forces in
molecules of a similar
size.
• Molecule will have not
just VVF but in addition
to them the DPDP
forces also.
Hydrogen bonding
• This occurs in
molecules that contain
H bonded to N/O/F
• The non bonding epair on these N/O/F
atoms interacts with the
H atom that carries a
high δ+ character coz its
bonded to another of
these small very electroatoms.
• It’s the strongest of all
the intermolecular
forces.
Examples to explain Van der waals’ forces
Boiling
points of
noble gases
Boiling
points of
Alkanes
Boiling
points of
Halogens
• He – 4K
• Xe – 165 K
• More atomic
mass hence
more no. of
electrons
• Methane111 K
• Hexane 341 K
• As molar mass
increases the
VVF also
increase
• As molar mass
of halogens
increases from
F2 to I2 the
boiling points
also increase
Boiling
points of
straight
chain &
branched
Alkanes of
same molar
mass
• n pentane BP
309 K
• neo pentane
BP 283 K
Plastics and
polymers
They have
very high
Molar mass
and very
high surface
area
Examples to explain dipole-dipole forces
Hδ+— Clδ- ||||||||||||Hδ+— Clδelectrostatic attraction
• The HCl molecule is polar
and has slight negative and
slight positive centres on it.
• This is a permanent DIPOLE
and it interacts with other
dipoles also and this
electrostatic attraction will
be DP-DP forces.
Comparing the boiling
points of non polar noble
gases and polar hydrogen
halides of same molar masses
• Boiling points of Non-polar Ar,
Kr and Xe will be lower than
those of Polar hydrogen halides
like HCl, HBr and HI
• Because these are permanent
dipoles they will have DP-DP
forces along with VV forces
Hydrogen bonding
 Hδ+— Fδ- ---------------Hδ+— Fδ
hydrogen bond
 Considerably stronger than other intermolecular
forces.
 Affects the physical properties of the compounds in
which it exists.
Examples of H-Bonding
Comparing BP of Hydrogen peroxide, Fluorine
and Hydrogen Chloride (Mr ~ 34-36)
• Hydrogen peroxide
• Hydrogen Chloride
• Fluorine
431 K
188 K
85 K
(polar, hydrogen bonded)
(polar)
(non polar)
Comparing BP of Propane, Ethanal and
Ethanol (Mr~ 44-46)
• Propane C3H8
only VVF
BP is 231 K
• Ethanal CH3CO-H VVF, polar
BP is 294 K
• Ethanol CH3CH2OH VVF, polar, Intermolecular H bonding—352
K
Collision Theory
A reaction will only be successful if it has the correct
orientation and energy.
Baseball Model: Baseball bat is Reactant A, baseball is
Reactant B. The reaction is successful if the batter hits
a homerun!
The scenarios…
1) Pitcher throws a fast ball, batter swings and misses.
2) Pitcher throws off-speed and the batter just makes
contact with the ball. It’s a foul.
3) Pitcher throws a curve ball, batter swings and just
hits the ball sending it to right field.
4) Pitcher throws a fast ball, batter swings and the ball
goes flying high in the air. HOMERUN!
When did the reaction occur?
Try with molecules
Reaction: H2 + I2  2 HI
+

H2
I2
HI
How would you orient the molecules to get an effective
collision capable of producing two hydrogen iodide
molecules?
Types of Reactions
There are five types of chemical reactions we will
talk about:
•
1.
2.
3.
4.
5.
•
Synthesis reactions
Decomposition reactions
Single displacement reactions
Double Displacement reactions
Combustion reactions
You need to be able to identify the type of reaction
and predict the product(s)
Steps to Writing Reactions
Some steps for doing reactions
•
1.
2.
3.
Identify the type of reaction
Predict the product(s) using the type of reaction as a
model
Balance it
Don’t forget about the diatomic elements! (BrINClHOF) For
example, Oxygen is O2 as an element.
In a compound, it can’t be a diatomic element because it’s not
an element anymore, it’s a compound!
1. Synthesis reactions
• Synthesis reactions occur when two substances (generally
elements) combine and form a compound. (Sometimes
these are called combination or addition reactions.)
reactant + reactant  1 product
• Basically: A + B  AB
• Example: 2H2 + O2  2H2O
• Example: C + O2  CO2
Practice
• Predict the products. Write and balance the following
synthesis reaction equations.
• Sodium metal reacts with chlorine gas
2 Na(s) + Cl2(g)  2 NaCl (s)
• Solid Magnesium reacts with fluorine gas
Mg(s) + F2(g)  MgF2 (s)
• Aluminum metal reacts with fluorine gas
2 Al(s) + 3 F2(g)  2 AlF3 (s)
2. Decomposition Reactions
• Decomposition reactions occur when a compound
•
•
•
•
breaks up into the elements or in a few to simpler
compounds
1 Reactant  Product + Product
In general: AB  A + B
Example: 2 H2O  2H2 + O2
Example: 2 HgO  2Hg + O2
Decomposition Exceptions
• Carbonates and chlorates are special case
decomposition reactions that do not go to the
elements.
• Carbonates (CO32-) decompose to carbon dioxide and a
metal oxide
•
Example: CaCO3  CO2 + CaO
• Chlorates (ClO3-) decompose to oxygen gas and a metal
chloride
•
Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2
• There are other special cases, but we will not explore
those in Chemistry 11
Practice
• Predict the products. Then, write and balance the
following decomposition reaction equations:
• Solid Lead (IV) oxide decomposes
PbO2(s)  Pb (s) + O2 (g)
• Aluminum nitride decomposes
2 AlN(s)  2 Al (s) + N2(g)
Practice
Identify the type of reaction for each of the
following synthesis or decomposition
reactions, and write the balanced equation:
Nitrogen monoxide
N2(g) + O2(g) 
BaCO3(s) 
Co(s)+ S(s) 
(make Co be +3)
NH3(g) + H2CO3(aq) 
NI3(s) 
3. Single Displacement Reactions
• Single Displacement Reactions occur when one element
replaces another in a compound.
• A metal can replace a metal (+) OR
a nonmetal can replace a nonmetal (-).
• element + compound product + product
A + BC  AC + B (if A is a metal) OR
A + BC  BA + C (if A is a nonmetal)
(remember the cation always goes first!)
When H2O splits into ions, it splits into
H+ and OH- (not H+ and O-2 !!)
Single Replacement Reactions
• Write and balance the following single replacement
reaction equation:
• Zinc metal reacts with aqueous hydrochloric acid
Zn(s) + HCl(aq)  ZnCl2 + H2(g)
Note: Zinc replaces the2hydrogen ion in the reaction
Single Replacement Reactions
• Sodium chloride solid reacts with fluorine gas
NaCl(s) + F2(g)  NaF(s) + Cl2(g)
Note that fluorine replaces chlorine in the compound
• Aluminum metal reacts with aqueous copper (II) nitrate
Al(s)+ Cu(NO3)2(aq)
4. Double Replacement Reactions
• Double Replacement Reactions occur when a metal
replaces a metal in a compound and a nonmetal replaces a
nonmetal in a compound
• Compound + compound  product + product
• AB + CD  AD + CB
Double Replacement Reactions
• Think about it like “foil”ing in algebra, first and last ions go
together + inside ions go together
• Example:
AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq)
• Another example:
K2SO4(aq) + Ba(NO3)2(aq)  KNO3(aq) + BaSO4(s)
•
Practice
Predict the products. Balance the equation
HCl(aq) + AgNO3(aq) 
2. CaCl2(aq) + Na3PO4(aq) 
3. Pb(NO3)2(aq) + BaCl2(aq) 
4. FeCl3(aq) + NaOH(aq) 
5. H2SO4(aq) + NaOH(aq) 
6. KOH(aq) + CuSO4(aq) 
1.
5. Combustion Reactions
• Combustion reactions
occur when a hydrocarbon
reacts with oxygen gas.
• This is also called burning!!!
In order to burn something
you need the 3 things in the
“fire triangle”:
1) A Fuel (hydrocarbon)
2) Oxygen to burn it with
3) Something to ignite the
reaction (spark)
Combustion Reactions
• In general:
CxHy + O2  CO2 + H2O
• Products in combustion are
ALWAYS carbon dioxide and
water. (although incomplete
burning does cause some byproducts like carbon monoxide)
• Combustion is used to heat
homes and run automobiles
(octane, as in gasoline, is C8H18)
Combustion Reactions
Edgar Allen Poe’s drooping eyes and
mouth are potential signs of CO
poisoning.
Combustion
• Example
•
C5H12 + 8 O2  5 CO2 +6 H2O
• Write the products and balance the following
combustion reaction:
•
C10H22 + O2 
Mixed Practice
•
1.
2.
3.
4.
5.
State the type, predict the products, and balance the
following reactions:
BaCl2 + H2SO4 
C6H12 + O2 
Zn + CuSO4 
Cs + Br2 
FeCO3 
Total Ionic Equations
 Once you write the molecular equation (synthesis,
decomposition, etc.), you should check for
reactants and products that are soluble or
insoluble.
 We usually assume the reaction is in water
 We can use a solubility table to tell us what
compounds dissolve in water.
 If the compound is soluble (does dissolve in
water), then splits the compound into its
component ions
 If the compound is insoluble (does NOT dissolve
in water), then it remains as a compound
Solubility Table
Solubilities Not on the Table!
 Gases only slightly dissolve in water
 Strong acids and bases dissolve in water
 Hydrochloric, Hydrobromic, Hydroiodic, Nitric,
Sulfuric, Perchloric Acids
 Group I hydroxides (should be on your chart anyway)
 Water slightly dissolves in water! (H+ and OH-)
 For the homework… SrSO4 is insoluble; BeI2 and
the products are soluble
 There are other tables and rules that cover more
compounds than your table!
Total Ionic Equations
Molecular Equation:
K2CrO4 + Pb(NO3)2 
Soluble
Soluble
PbCrO4 + 2 KNO3
Insoluble
Total Ionic Equation:
2 K+ + CrO4 -2 + Pb+2 + 2 NO3- 
PbCrO4 (s) + 2 K+ + 2 NO3-
Soluble
Net Ionic Equations
 These are the same as total ionic equations, but you
should cancel out ions that appear on BOTH sides of
the equation
Total Ionic Equation:
2 K+ + CrO4 -2 + Pb+2 + 2 NO3- 
PbCrO4 (s) + 2 K+ + 2 NO3Net Ionic Equation:
CrO4 -2 + Pb+2  PbCrO4 (s)
Net
Ionic
Equations
 Try this one! Write the molecular, total ionic, and net
ionic equations for this reaction: Silver nitrate reacts
with Lead (II) Chloride in hot water.
Molecular:
Total Ionic:
Net Ionic: