SECOND NINE WEEKS
NOTES PACKET
COLLIER
CHEMISTRY PRE-AP
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UNIT 5
CHEMICAL NAMING &
BALANCING
Chapter 6, 15.1, 16.1
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NOMENCLATURE:
 Atoms of _______________ elements combine to form
_________________ that are represented by
________________________________.
 All compounds can be _________________ in such a way
that others can understand the make-up of the
compound.
 Chemical formulas are represented by using the
______________ symbols with ________________ that
show the number of each element in a given compound
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Chemical Bonding
Molecules – the smallest electrically neutral unit of a
substance.
1. _________________ are the building blocks of substances.
2. Atoms of most elements will _______________________
different compounds (except ____________________)
3. Molecules are made up of ______________________ that
act as one unit.
Ex:
4. Molecular compounds can be called _________________
or ______________________ molecules.
5. Molecular Compounds are composed of molecules that
are ___________________________________________.
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Ions & Ionic Compounds
1. ___________ are atoms or a group of atoms that
have ____________________________________
2. A loss of electrons gives the atom a positive charge;
therefore a gain of an electron would give the atom a
negative charge.
3. Atoms of ___________________ elements tend to
___________ electrons to become ______________
and ________________ tend to _________________
electrons to become ____________________.
(+) ions also called _____________.
Ex:
(-) ions = ________________.
Ex:
4. Ionic compounds consist of a _______________
bonded to a ____________. The charge of the
compound is ___________________.
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OXIDATION NUMBER
 This is a ______________ or _______________ number
assigned to an atom according to a set of rules
 The oxidation number of a ___________________ ion is
equal to its ionic charge
Ex. Br- = _______
Mg +2 = __________
 The oxidation number of ____________ in a compound is
_________ except in a metal hydride where it is ______
 The oxidation number of _____________ is ___________
except in ____________ where it is __________
 The oxidation number of the atom in its _______________
form is _________
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Electrons in Ionic Bonding
Valence Electrons --
1. Determines the _______________ properties of an
element.
2. The __________ determines the number of valence
electrons
3. Electron Dot Structures – ___________________ that
show only valence electrons.
Octet rule -- atoms tend to achieve the ___________________
_____________________________.
**Atoms will gain or lose electrons to achieve this.
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Ionic Bonds -- the ____________ of attraction between
___________________.
(____________ &__________________)
Covalent Bonding -- a bond in which two atoms _________ __
a pair of electrons. (________________________)
1. Structural formula H – H
2. Sharing of electrons occur if atoms involved
achieve the _________________ rule.
3. _______________ and ______________ Covalent
Bonding means that more than one pair of
electrons are shared.
Example:
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Chemical Compounds
Chemical Formula’s -- Shows the ______________ and
_______________ of atoms in a compound.
- _______________ show the number of atoms. Ex: H2O
• Ionic compounds consist of a ____________ ion with a
________________ ion in the lowest whole number ratio.
Ex:
• Molecular compounds typically consist of
_____________________ covalently bonded together
Ionic Charges
• ________________________ ions consist of only 1 ion.
• _________________ ions consist of tightly bound atoms
that behave as a single unit, that has either lost or gained
electrons to become (+) or (-). Ex: NO3-, SO4-2
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Balancing Formulas
Writing Formulas
Ionic Compounds
• Criss Cross Method
1. Write __________________ and ______________ for
each ion present.
Compare charges:
 If charges equal _________ then simply write symbols
together.
 If charges don’t equal 0 then _____________________
and write them as a _________________.
Criss-Cross Method
Al+3 + O-2 → Al2O3
Examples:
Sodium Chloride:
Aluminum Sulfide:
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Binary Molecular Compounds Consists of two nonmetals
1. Write the symbols of the ___________ elements,
change ________________ to subscripts.
Prefixes
Mono =
Hexa =
Di =
Hepta =
Tri =
Octa =
Tetra =
Nona =
Penta =
Deca =
Acids
1. Acids consist of a ___________ bonded to an
_____________ (negative ion ).
Step 1:
Step 2:
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Examples:
- Hydrochloric Acid
- Nitric Acid
- Nitrous Acid
Step 3: _________________________________ for this
compound as you would when writing any molecular formula.
Hydro- ; -ic = -ide
-ic = -ate
-ous = -ite
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Naming Compounds
Ionic Compounds
1. Write the ___________ of the ___________ present.
2. Determine which ____________ forms the compound
by looking at the ratio of the atoms and compare to
the ratio of the charges.
3. Remember that the overall charge of the compound
has to ________________________.
4. Example: CuO is it copper I or II?
 The ratio of atoms is 1:1 and the charge on O is –2
so the Cu has to be 2 to have the same ratio. The
name of the compound would be
___________________
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Binary Molecular Compounds:
1. Write the name of the ____________ in the first position
and if it has a subscript change it to a _____________.
2. Write the name of the ______________ element. Change
the ending to ___________, and always use a
_______________ that corresponds with the number of
atoms present in the molecule.
Example:
PO5 =
S2 N =
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Acids
**If the compound starts with an H then it considered to be an acid.
Step 1:
Step 2:
- If the ending is __________, place ____________ in front of
the negative ion name, then drop the –ide and add _________.
- If the ending is ____________, drop the -ate ending and add
_____________.
- If the ending is ______________, drop the ending and add
__________.
Examples:
H2S =
H2SO4 =
Hydro- ; -ic = -ide
H2SO3 =
-ic = -ate
-ous = -ite
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Molecular Weights
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UNIT 6
Chemical Quantities--Chapter 7
Atomic Mass (Weight)
 The ______________ of an ______________ expressed in
_______________.
 Examples:
Carbon –
Sulfur –
Iron –
Mass of a Mole of a Compound
 1st you must _____________________________ of the
compound
 The formula tells you ___________________ of each
element that is present.
 Example:
SO3 –
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 After you know the _________________ then you
calculate the _______________ of the compound.
Molar mass


Find the molecular mass of each compound.
1. Li2S
2. FeCl3
3. Ca(OH)2
4. N2O5
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Percent Composition
% mass of element =
Or
Grams of element
Grams of Compound X 100
% Comp. = Part
X 100
Whole
Practice Problem
 Ethane (C6H6) ---- 222.06 g of C6H6
mole of C = 216.00 g
mole of H = 6.06 g
Find the % Composition of Carbon & Hydrogen in Ethane (C6H6).
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Significant Figures & Scientific Notation
 All the _______________that can be known precisely in a
measurement, plus a __________________________ digit.
 Expression of numbers in the form ___________________
where n is equal to or greater than one and less than 10
H2O –
 How many much Hydrogen is in Water?
 What is Hydrogen’s % Composition in Water?
 What would happen to the % of hydrogen with 2
molecules of Water?
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SIGNIFICANT DIGITS
To determine the number of significant digits in a written number complete one of the
following:
1. Qualifying statement: The numerical value has no decimal place written.
Action to take: Count from the first non-zero digit to the last non-zero digit.
Examples:
1.
18,004 - 5 significant digits
2.
10,040,000– 4 sig digs
3.
10 – 1 sig dig
2.
Qualifying statement: The numerical value contains a decimal place.
Action to take: Count from the first non-zero digit to the end of the number.
Examples:
1.
100.00 – 5 sig digs
2.
1,050. – 4 sig digs
3.
0.000 145 00 – 5 sig digs
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What is a mole?


Avogadro’s number = _________________________________
Standard Temperature and Pressure (STP) –
a. Standard temperature is ____________ or ____________
b. Standard pressure is _______________________________
_______________________________________________.
• 1 mol of any gas at STP = ________________
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Representative Particle
 Refers to the _________________________ in a
substance: usually atoms, molecules, or formula units.
(ions)
 Example:
 Fe is composed of _______________ atoms.
 K is composed of ___________________ atoms.
Example: How many moles of Mg is 1.25 x 1023 atoms of Mg?
Know:


Unknown:



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Practice Problems
How many moles are 2.80 x 1024 atoms of silicon?
How many molecules is 0.360 moles of water?
Amt. in Compounds
 Carbon dioxide has 3 atoms.
 1-carbon 2-oxygen
 Thus one mole of CO2 contains three times Avogadro’s # of
atoms.
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Example: How many atoms are in 2.12 mole of propane
(C3H8)?
Know:



Unknown:

Desired conversion:
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Practice Problems
How many atoms are there in 1.14 mole SO3?
How many moles are there in 4.65 x 10 24 molecules of NO2?
Mole -- gram
How many grams are in 3.32 mole K?
How many grams are in 4.52 x 10-3 moles of C20H42?
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Gram -- Mole
How many moles are in 3.70 x 10-1 grams of Boron?
How many moles are in 27.4 g TiO2?
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Unit 7
Balancing Chemical
Reactions
&
Predicting Products
Ch. 8 & 19.1
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Reactions
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Equations
An _______________________ is a description of a
chemical reaction indicating the
___________________________, the
_____________________ and a _____________ of their
quantities.
REMEMBER:
__________________ → _________________
_______________ (reactants) = ___________________
(products)
Coefficients --
The reaction is said to go to ______________________
when no reactants remain when the reaction has stopped.
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The reaction may be ______________________. When this
happens, a complete cycling of events exists. It is a
continuous ongoing process.
The Law of Conservation of Mass: in reactions atoms are
neither _______________________ nor
___________________ only ______________________
The Law of Conservation of Mass
__________________________________________.
What is the relationship between the mass of the reactants
and the mass of the products?
How are chemical reactions balanced?
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_________________ are used to separate reactants or
products on the same side of the arrow.
Catalyst:
2H2 + O2 → 2H2O
1. What are the reactants?
2. What are the products?
3. What are the coefficients for the whole reaction?
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Rules for Balancing Equations:
1. Determine the ____________________ for all reactants and
products. Indicate physical states in parentheses.
2. Write the formulas for the __________________
and _____________________________ of the arrow. Then
separate two or more formulas with + signs.
3. __________________________________of each element for
both sides of the equation.
4. ______________ the elements one at a time by using
________________.
5. Check each atom to be sure that the equation is balanced.
6. Make sure the _______________ are in the
______________________________that balances.
Diatomic Elements


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Classifying reactions
Synthesis (Combination) reaction:
 The ____________ can be elements or compounds.
 The _______________ will always be a compound.
A + B → AB
Decomposition reactions:
• The __________ is always a ______________.
• Most decomposition Reactions requires energy such as
light, heat or electricity.
AB → A + B
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Single replacement reactions:
 Single elements changes places with the other ion that is the
same as the single element. ____________________ &
______________________
A + BC → AC + B
Double Replacement:
AB + CD → CB + AD
Combustion Reactions:
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Factors Effecting Reaction Rates:
Temperature:
Concentration:
Catalyst:
Inhibitor:
Neutralization:
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Predicting Products
1. To predict products of any reaction the first step is to
______________________________________ it is.
2. Then apply existing rules for the type of reaction it is:
Activity Series of metals: an invaluable aid to predicting the
products of replacement reactions. It also can be used as an aid in
predicting products of some other reactions.
Going from bottom to top, the metals:





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Metal
Metal Ion
Reactivity
Lithium
Li+
Most Reactive
Potassium
K+
-
Calcium
Ca2+
-
Sodium
Na+
-
Magnesium
Mg2+
-
Aluminum
Al3+
-
Manganese
Mn2+
-
Zinc
Zn2+
-
Chromium
Cr2+,Cr3+
-
Iron
Fe2+,Fe3+
-
Lead
Pb2+
-
Copper
Cu2+
-
Mercury
Hg2+
-
Silver
Ag+
-
Platinum
Pt2+
-
Gold
Au+,Au3+
Least Reactive
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Synthesis (Combination): ______________ +
___________________ →
 Combine the 2 elements into a compound
Practice:
1. Aluminum + Oxygen →
2. Chlorine + Potassium →
3. Chromium + Bromine →
4. Oxygen + Hydrogen →
5. Calcium + Iodine →
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Single replacement: _____________ + _______________ 
 If the element forms a + ion replace the + ion of the
compound, and if it is a - ion the replace the – ion of the
compound.
Practice:
1. Aluminum + Magnesium oxide →
2. Copper (II) sulfide + chlorine →
3. Fluorine + Nickel bromide →
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Double Replacement: ______________ + _____________ 
 Switch the position of the (+) ions.
Practice:
1. Aluminum sulfite + Calcium acetate→
2. Tin (IV) phosphate + Ammonium dichromate →
3. Magnesium chloride + Chromium peroxide →
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Decomposition Reactions
Metallic Chlorates:
_____________ ________ + ________
 Substitute the specific + ion for the “M” and follow the word
equation.
Metallic Carbonates:
____________  _______ + ________
 Substitute the specific + ion for the “M” and follow the word
equation.
Metallic Hydroxides:
___________  _______ + _________
 Substitute the specific + ion for the “ M “ and follow the
word equation.
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Oxy Acids:

Example:
Simple Decomposition:
__________  ______ + ______
 Split the compound into is simpler parts.
Practice
1. Potassium sulfide 
2. Zinc oxide 
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3. Zinc carbonate →
4. Sodium carbonate →
5. Calcium hydroxide →
6. Aluminum hydroxide →
7. Potassium chlorate →
8. Calcium chlorate →
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