Download Compounds and Equations

Compounds and
Equations
Part One
Interpreting Chemical
Formulas…
Chemical formulas give us important
information about the chemical
compound that we are studying or
working with in a laboratory setting.
These formulas are part of the “recipe”
for important chemical reactions.
New Terminology
Empirical Formulas: Useful when talking about ionic
compounds. The empirical formula is the simplest
whole number ratio of the elements in an ionic
compound.
We use these for ionic compounds because ionic
compounds do not form separate molecules. Instead,
it’s just a ratio of alternating positive and negative
charges.
The subscripts tell you how many of each ion there are
in a single unit or how many moles of ions there are in
1 mole of that sample.
Example: CaCl2: The ratio of Calcium to Chlorine
ions is 1:2
This means if there are 50 Ca+2 ions, there will
be 100 Cl- ions.
Mole: It is useful when counting objects, to
define them in terms of a collective unit. For
example, we use a dozen to mean 12 items. This
makes it easier to count a large amount of
items.
It’s the same way with particles, we cannot
possibly count all the molecules or atoms in a
substance. Instead we use a unit called the mole
1 mole = 6.02 x 1023 particles of anything
1 mole is the number of atoms that weigh their
atomic mass in grams.
For example, the atomic mass for sodium is 23.0
grams. We can say that 1 mole of sodium atoms
(6.02x1023 atoms) would have a mass of 23.0
grams.
This is known as Avogadro’s number
Avogadro’s number is a constant, and does not change.
However, the mass of 1 mole of a substance WILL
change depending on the substance, because the
composition changes.
For instance, while we know that a dozen always means
12 items… A dozen bricks has a much greater mass than
a dozen feathers.
The substance you are working with determines the
mass of a “dozen”.
New Terminology
Molecular Formula: This tells us how many
atoms of each element we would need to
form the molecule.
Remember, molecules are formed through
covalent bonding and done exclusively with
non-metals.
Molecular formulas are not useful when
studying ionic compounds but are most
appropriate when studying molecular
compounds.
For example, to form a molecule of methane, it
takes 1 carbon atom and 4 hydrogen atoms giving
a molecular formula of CH4. Note, this is also
methane’s empirical formula because it cannot be
simplified.
To form a molecule of glucose it takes 6 C, 12 H, 6
O, giving a molecular formula of C6H12O6. This is
not the empirical formula of glucose, but it tells us
how to build that molecule.
The empirical formula is CH2O… However, this
does not tell us how to build glucose.
New Terminology
Coefficient: tells you how many
moles of a given compound you
have. Useful when balancing
equations – in the future.
Review: Interpreting a
Chemical Formula
Empirical
Formula
Fe2O3
K2SO4
Al2(SO4)3
# of Ions of
Each Element
# of moles of
each element
Total # of
moles of ions
Review: Interpreting a
Chemical Formula
Empirical
Formula
Moles of Each
Element
Formula with
Coefficient
CaCl2
3 CaCl2
Al(NO3)3
2 Al(NO3)3
(NH4)3PO4
5 (NH4)3PO4
Moles of each
ion
Gram Formula Mass
Just like a dozen bricks has more mass than a
dozen feathers…
The molar mass aka the Gram Formula Mass
aka the Molecular mass is the mass of 1 mole
of a substance.
The units are g/mol which translates into the
number of grams in one mole of the substance.
Gram Formula Mass
How to calculate the GFM:
1. Count the atoms of each element in the
compound
2. Look up the mass of each element on
the periodic table, round to the tenths
place.
3. Multiply the number of atoms by the
mass
4. Add up each mass.
Gram Formula Mass
Examples
What is the molar mass of Cu?
What is the molar mass of KCl?
Calculate the GFM of Na3PO4:
Calculate the GFM of Al2(SO4)3 :
Mole to Gram Conversions
When chemical reactions take place, we can think about
them as if they were recipes. These recipes are based on
mole amount of each ingredient.
However, we don’t have balances that measure moles, we
always measure in grams. Therefore, we need to be able
to convert between grams and moles of our ingredients so
that way we can ensure the reaction runs properly.
Two types of problems:
1. Determine how many moles there are in a certain mass
2. Determine the mass of a certain number of moles.
Mole to Gram Conversions
Anytime you are talking about grams and moles, you
must calculate the GFM of the compound.
To solve: there is a formula on the back of the reference
table:
Formula:
You can also use unit cancelling ( factor label method) to
solve these problems. See guided practice below.
Mole to Gram Conversions
Helpful hint:
• If it’s more than 1 mol, your answer must
be greater than the GFM
• If it’s less than 1 mol, your answer must be
less than the GFM.
Mole to Gram Conversions
Problem 1: Determine how many moles….
Example 1: Calculate the number of moles in a
500 gram sample of MgCl2.
Formula Method:
Factor Label Method:
Mole to Gram Conversions
Example 2: How many moles of Sr(NO3)2 in
75.5 grams of strontium nitrate?
Formula:
Factor Label Method:
Mole to Gram Conversions
Example 3: How many grams are
present in 3.4 moles of NH4Cl?
Formula:
Factor Label Method:
Mole to Gram Conversions
Example 4: How many grams are present in
4.5 moles of potassium oxide?
Formula for potassium oxide:
Formula:
Factor Label Method:
Percent Composition
PERCENT COMPOSITION AND MOLECULAR FORMULAS:
Percent Composition: The proportion by mass of each
element in a compound.
This can be done experimentally or using the formula..
Either way, there is a formula on Table T of your reference
table.
Formula:
Percent Compostion
Example: Determining percent composition experimentally:
A compound containing nitrogen and oxygen has a mass of 80.0
grams. Experiments show that the 80.0 grams is made up of
56.0 grams of Oxygen and 24.0 grams of Nitrogen. What is %
composition by mass of each element?
• %O
%N
Percent Composition
Example: Determine using the formula… You need
the periodic table for this:
Step 1: Calculate the GFM of the whole compound
Step 2: Plug into formula
Percent Composition
• What is the % composition by mass of each element
in SiO2?
• What is the % composition by mass of each element
in H3PO4 ?
• What is the percent composition of nitrogen in
NH4NO3?
Percent Composition of a
Hydrate
Hydrate: Ionic solids with water trapped within the crystal
lattice. The amount of water trapped in the solid is expressed
in this format:
Ionic Compound • n H2O
Example: CuSO4 • 5 H2O
This can be named as Copper II Sulfate Pentahydrate. The 5
water molecules are trapped within the copper sulfate crystal
lattice.
Anhydrate: An ionic solid with no water trapped inside. Can be
formed by heating a hydrate until all of the water is evaporated.
Percent Composition of a
Hydrate
PERCENT COMPOSTION OF A HYDRATE: Can also be done
experimentally or using the formula.
Example 1: Experimental determination:
• Subtract the mass of the anhydrate (without water, after
heating) from the mass of the hydrate (with water, before
heating).
• This is the mass of the water that evaporated.
• Divide the mass of the water by the mass of the hydrate.
Percent Composition of a
Hydrate
1. A 10.40 gram sample of a hydrated crystal is heated to a
constant mass of 8.72 grams. What is the percent
composition of water in this hydrate?
2. An empty evaporating dish is found to have a mass of 29.993
grams. A sample of the hydrated crystal is placed into the
evaporating dish, the combined dish 39.486 g. The evaporating
dish is heated and removed from the flame when it reaches a
constant mass of 38.378 g. what is the percent by mass of water
in the hydrate?
Percent Composition of a
Hydrate
Example 2: Determining using the formula:
Calculate the GFM of the hydrate, including the water.
Plug into the formula on Table T.
1. What is the percent by mass of water in Na2CO3 • 10 H2O
2. What is the percent by mass of water in BaCl2 • 2 H2O?
Empirical and Molecular
Formulas
Molecular Formulas: Empirical formulas give you the simplest
whole number ratio. While this is fine for ionic compounds, they
do not tell the full story for molecular compounds.
Molecular compounds are a multiple of an empirical formula:
Example: Glucose
Divide
Molecular Formula
C6H12O6
Empirical Formula
CH2O
Multiply
Empirical and Molecular
Formulas
It is easy to arrive at a compounds empirical formula, you
simply divide each of the subscripts by the greatest common
factor to reduce them to the simplest whole number ratio.
Examples: Determine the empirical formula:
Molecular Formula
N2O4
C4H8
B4H10
P2O5
C2H6
C3H9
Empirical Formula
Empirical and Molecular
Formulas
TO DETERMINE THE MOLECULAR FORMULA:
You cannot determine the molecular formula, you will need the
empirical formula and you will need to know what to multiply
the empirical formula by to get the molecular formula.
Steps:
1. Determine the GFM of the empirical formula
2. Divide the molecular mass by the empirical mass.
* This should be a whole number. If it is not, you did
something wrong*
3. Multiply the empirical formula by the whole number.
Empirical and Molecular
Formulas
Example: The empirical formula for ethylene is CH2. Find the
molecular formula if the molecular mass is 28.1 g/mol.
Example: What is the molecular formula of a compound that has
an empirical formula of NO2 and a molecular mass of 92.0
g/mol?
Example: a compound has an empirical formula of HCO2 and a
molecular mass of 90 grams per mole. What is the molecular
formula of this compound?
Compounds and
Equations
Part Two
How is the Law of Conservation of mass
applied during chemical changes?
Chemical Equations use chemical formulas and physical
states like solid (s), liquid (l), gases (g), and aqueous (aq) to
describe a chemical reaction.
These equations give us information about what is
reacting, what is being formed and how much or the
number of moles of each substance and even the energy
that might be absorbed or released.
All chemical equations follow the same basic format:
Reactants  Products
How is the Law of Conservation of mass
applied during chemical changes?
For example: Write the equation for the following
reaction: Zinc reacts with hydrogen chloride to
produce hydrogen gas and zinc chloride.
Eq: Zn(s) + HCl (aq)  H2 (g) + ZnCl2 (aq)
In the equation above, you may notice
something not right. As written, this
equation does not obey the Law of
Conservation of Mass.
How is the Law of Conservation of mass
applied during chemical changes?
Law of Conservation of Mass can be thought of as
the number of atoms of each element on the left
must be equal to the number of atoms of each
element on the right.
In order to “balance” this equation, we use
numbers known as coefficients. Coefficients are
interpreted to mean the moles of each substance
How is the Law of Conservation of mass
applied during chemical changes?
Balance the equation in order to follow the Law
of Conservation of Mass:
Eq: Zn + 2 HCl  H2 + ZnCl2
This can be read as 1 mole of zinc reacts with
2 moles of hydrogen chloride to produce 1
mole of hydrogen gas and 1 mole of zinc
chloride. These numbers give us a ratio of
reactants to products.
How is the Law of Conservation of mass
applied during chemical changes?
Some guidelines for writing equations and balancing equations:
1. Make sure that the formulas are written properly. Criss cross for ionic compounds,
follow the prefixes for molecular compounds.
2. Make sure that you have reactants on the left hand side and products on the right
hand side. Remember “yields” means the same as “produces” and is usually where the
arrow should be drawn.
3. Evaluate the equation: what elements are not balanced?
4. Balance in pencil!
a. Start with an element that is on both sides
b. Keep polyatomic ions together when possible (balance as a unit)
c. Leave lone elements until the very end
d. Never, ever change the formula.
How is the Law of Conservation of mass
applied during chemical changes?
Examples: Balance the following
___ Al + ____ CuCl2  ____ Cu + ____ AlCl3
____ H2 + ____ O2  ____ H2O
____ N2 + ____ H2  _____ NH3
____ BaCl2 + _____ AgNO3  ____ Ba(NO3)2 + _____ AgCl
____ FeCl3 + _____ Pb(NO3)2  _____ Fe(NO3)3 + _____ PbCl2
How is the Law of Conservation of mass
applied during chemical changes?
How is the Law of Conservation of mass
applied during chemical changes?
Examples: Write the chemical equations given the following
word equations and balance:
1. Copper reacts with oxygen gas to produce copper (I) oxide.
Word equation:
____________________________________________
Chemical equation:
_____________________________________________
How is the Law of Conservation of mass
applied during chemical changes?
2. Lead(II)Nitrate reacts with sodium phosphate to yield
lead(II)phosphate and sodium nitrate.
Word Eq:
_________________________________________________
Chemical Eq:
______________________________________________
How is the Law of Conservation of mass
applied during chemical changes?
Examples: Use the following equations and your understanding
of the Law of Conservation of Mass to answer the following:
1. If 35.0 grams of nitrogen gas are reacted with hydrogen gas to
produce 42.5 grams of ammonia gas, how many grams of
hydrogen gas were reacted?
Equation: ________________________________________
How is the Law of Conservation of mass
applied during chemical changes?
2. How many grams of aluminum are formed when 45.0
grams of aluminum oxide are decomposed into aluminum
and 21.0 grams of oxygen?
Equation: _______________________________________
These are called “Missing Mass Calculations: To abide by the
Law of Conservation of Mass:
The mass on the reactants side and products side must be
equal to each other.
Guiding Question 2: What are the five
major types of chemical reactions?
Reaction Type 1: Synthesis
In its most basic form: two elements combine to form a
compound. Today, our definition is expanded a bit, where two
substances combine to form one product.
If an equation fits the following format, it can be classified as a
synthesis reaction:
Format:
A + B  AB
or
A + B2  AB2
*The second example shows that it does not matter what the
formula of the compound or the individual substances might be.
Sickly Bird
+ Worms
2 reactants  1 product
= Healthy Bird
Guiding Question 2: What are the five
major types of chemical reactions?
Reaction Type 2: Decomposition
This can be thought of as the opposite of
synthesis. A compound is broken down into
2 simpler substances.
Format: AB  A + B
or AB2  A + B2
Baby Turtle in Egg
Egg shell
1 reactant  2 or more products
+ Turtle
Guiding Question 2: What are the five
major types of chemical reactions?
Reaction Type 3: Single Replacement
This is identified by: one element and a compound on
EACH side.
Two options:
Option 1: Metal + Compound: In these reactions, the
metal replaces the cation in the compound.
Format:
A + BC  AC + B
Derek Jeter
+
Tony Romo
and Ms. 
Hannon
Derek
Jeter and
Ms.
Hannon
One element replaces another in a compound
+
Tony Romo
+
+
+
+
Hey! She’s mine! I’m
stronger than you,
I’m taking her back!
+
Finally, all is right
in the world.
+
Guiding Question 2: What are the five
major types of chemical reactions?
Example: Li + CaCl2  LiCl + Ca
This occurs because Li is more reactive than
Calcium. This is determined by using Table J
in your reference tables.
This type of reaction can only occur if the
metal that is REPLACING the other is more
active and therefore HIGHER on Table J.
Guiding Question 2: What are the five
major types of chemical reactions?
Example: Predict whether the following
reactions can happen. If yes: then predict
the products:
No Reaction: Au is lower on
Au + NaCl 
Table J
Zn +
2 CuNO3  2 Cu + Zn(NO3)2
Guiding Question 2: What are the five
major types of chemical reactions?
Option 2 : Nonmetal + Compound: The non-metal
replaces the anion in the compound. These can also
be predicted by using Table J.
Format:
A + BC  BA + C
Example:
F2 + ZnCl2  ZnF2 + Cl2
This occurs because fluorine is more reactive than
chlorine. This is shown on Table J as F2 is higher than
Cl2.
Guiding Question 2: What are the five
major types of chemical reactions?
Reaction Type 4: Double Replacement:
Two aqueous compounds swap positive ions
to form two new products.
Format:
AB + CD  AD + CB
+
Mickey and Minnie
Donald and Daisy
Donald and Minnie
+
Mickey and Daisy
Guiding Question 2: What are the five
major types of chemical reactions?
Once AB is dissolved, the ions will separate as the
polar water molecules pull them apart. The same
thing will occur between CD. The A,B,C,D ions are all
able to move freely in solution and find new partners.
In some cases, the new partnership is such as strong
attraction that the water molecules cannot pull them
apart. This will form an insoluble (water can’t pull
apart) precipitate (more to come in our Solutions
unit). These are abbreviated with an (s) after the
compound formula.
Guiding Question 2: What are the five
major types of chemical reactions?
Example: K2CrO4 (aq) + Ba(NO3)2 (aq)
 2 KNO3 (aq) + BaCrO4 (s)
Example: NaCl (aq) + AgNO3 (aq) 
AgCl(s) + NaNO3(aq)
Guiding Question 2: What are the five
major types of chemical reactions?
Reaction Type 5: Combustion:
This is the most specific type of reaction as it has a
predetermined set of reactants and products.
Format:
CxHy + O2  CO2 + H20
These are highly exothermic reactions… See Table I. Combustion
reaction are described as a hydrocarbon burns (which means
reacts with oxygen) to form carbon dioxide and water.
This is the same reaction that occurs in car engines, known as
combustion engines. The fuel that you purchase is a
hydrocarbon!
Guiding Question 2: What are the five
major types of chemical reactions?