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Unit 1 Notes
Chemical Reactions
Part 1
Grade 11 Chemistry
Unit 1: Chemical Reactions Part 1
The Atom
All atoms, except hydrogen, are made of 3 basic particles: protons,
neutrons and electrons. Each element has a unique number of protons,
which is indicated by its atomic number.
The outermost shell is called the valance shell.
valence shell are called the valence electrons.
The electrons in the
The atoms of elements in Period 1 have one shell. This shell contains a
maximum of 2 electrons.
The atoms of elements in Period 2 have two shells. The valence shell
contains a maximum of 8 electrons.
The atoms of elements in Period 3 have three shells. The valence shell
contains a maximum of 8 electrons.
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The Periodic Table
The periodic table arranges atoms according to their properties. The periodic
table below shows the names of several groups we will be referring to
throughout this course.
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Ionic Compounds
An ion is a charged particle. An ion is formed when a neutral atom gains or
loses electrons. An atom will gain or lose electrons according to the number
of valence electrons in its valence shell. For example, atoms with one
valence electron in its valence shell will have a tendency to lose that electron
and become positively charged. Positively charged ions are called cations,
and negatively charged ions are called anions. Atoms gain or lose
electrons so that their outermost shell is stable (contains 8 electrons).
Ionic compounds are formed when two or more oppositely charged ions
are attracted to each other. This attraction is called a chemical bond. An
ionic bond is formed when a negatively charged ion is attracted to a
positively charged ion.
Ions combine together so that their charges add up to zero.
Ionic compounds are usually made of metal and non-metal ions.
Examples of Ionic Compounds
NaCl – sodium chloride
Fe2O3 – iron oxide
CuSO4 – copper sulfate
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Chemical Formulas
A chemical formula is a shorthand method to represent compounds that uses
the elements’ symbols and subscripts. The chemical formula gives the
following information:


The different elements in the compound
The number of atoms of each element in the compound
Example:
H 2O
Ca3(PO4)2
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Naming Binary Ionic Formulas
A binary compound contains 2 different kinds of elements. There can be
more than one atom of each element in a binary compound. Binary ionic
compounds usually contain one kind of metal ion combined with one kind of
non-metal ion. Metal ions have positive charges and non-metal ions
have negative charges.
When naming an ionic compound from its formula, follow the rules below:
1. The cation (positive ion/metal) is named first, followed by the anion
(negative ion/non-metal).
2. Write the full name of the metallic elements (positive ion).
3. Write the name of the non-metallic element (negative ion) and change
the ending to “-ide”.
Example
Write the name of NaCl.
Step 1: Name the first element.
Step 2: Name the second element and change the ending to “-ide”.
The name of the compound is __________________________.
Example
Write the name of Mg3P2
Step 1: Name the first element.
Step 2: Name the second element and change the ending to “-ide”.
The name of the compound is __________________________.
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Writing Binary Ionic Formulas
The following must occur, when writing the formula for ionic compounds.
1. The formula must have the cation first, followed by the anion.
2. The sum of the charges of the ions must be zero. That is, the number
of positive charges must equal the number of negative charges.
3. You may NOT change the charge of the ions to make the ion charges
equal zero.
The “Criss-Cross” Method
1. Write the ions and their charges side by side.
2. “Criss-cross” the charges. That is, make the number of the charge of
one ion the subscript of the other ion (omitting the + or – sign).
Remember we do not write the number one as a subscript.
3. Reduce all subscripts to their simplest form, if necessary.
Example:
Write the formula for aluminum oxide.
Step 1: Write the ions and their charges.
Step 2: “Criss-cross” the charges. Make the number of the charge of one ion
the subscript of the other ion. Omit the + or – sign.
Example:
Write the formula for barium fluoride
Step 1: Write the ions and their charges.
Step 2: “Criss-cross” the charges. Make the number of the charge of one ion
the subscript of the other ion. Omit the + or – sign.
Note: The charge on the fluoride ion is 1-. Since IUPAC rules do not write
the number one as a subscript, we leave the barium without a subscript.
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Polyatomic Ions
Some ions are composed of several atoms joined covalently. These are
called polyatomic ions (poly = many).
Refer to the Periodic Chart of Polyatomic Ions for a list of ions.
The charge for polyatomic ions is for the whole group of atoms not just for
the atom written last. DO NOT change the subscripts of polyatomic
ions; if you change the subscripts you change the identity of these
ions.
When indicating the presence of more than one polyatomic ion in a
compound, we use parenthesis around the polyatomic ion, followed by its
subscript. For example, the compound Al(C2H3O2)3 has one aluminum ion
(Al) and 3 acetate ions (C2H3O2). Placing the acetate ion in parenthesis and
following it with the subscript 3 indicates there are 3 acetate ions.
Example
Write the name for KNO3
Step 1: Identify the cation.
Step 2: Identify the anion.
Step 3: Write the name of the cation first, followed by the anion.
Example
Write the name of Hg2Cl2
Step 1: Identify the cation
Step 2: Identify the anion.
Step 3: Write the name of the cation first, followed by the anion.
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Example
Write the name of Na3PO4.
Step 1: Identify the cation
Step 2: Identify the anion.
Step 3: Write the name of the cation first, followed by the anion.
Example
Write the name of NH4SCN.
Step 1: Identify the cation
Step 2: Identify the anion.
Step 3: Write the name of the cation first, followed by the anion.
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Stock Naming System
Several transition metals (group 3-12) have more than one possible
charge. They are often referred to as multivalent. For example,
Ion
Possible Charge
Copper
Iron
Cobalt
Chromium
Lead
Tin
1+,
2+,
2+,
2+,
2+,
2+,
2+
3+
3+
3+
4+
4+
In 1919, Alfred Stock (1876-1946), a German chemist, suggested using
numbers to indicate the charge on the ions. Prior to this the ions were given
different names based upon their charge. The Cu+ ion was called cuprous
and the Cu2+ ion was called cupric. However, the Fe2+ ion was ferrous and
the Fe3+ ion was ferric. Since the charges were not always the same, the
“-ic” and “-ous” suffixes caused some confusion. Today, the Stock naming
system uses Roman numerals following the metal ion’s name to indicate the
ion’s charge.
Example:
Copper (I) = Cu+
Copper (II) = Cu2+
Iron (II) = Fe2+
Iron (III) = Fe3+
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Writing the Chemical Formula
Example:
Write the formula for iron (III) chloride.
Step 1: Write out the ions.
Step 2: Criss-cross the charges.
Example:
Write the formula for lead (IV) sulfide.
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Writing the Chemical Name
We name in a very similar manner as those ions with a single ion charge,
except we must determine the charge on the metal ion.
Example:
Write the name for CoBr2
Example:
Write the name for MnO2
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Covalent Compounds
Naming Covalent Compounds
Non-metals tend to combine chemically by sharing electron pairs. These
bonds are known as covalent bonds. Neutral compounds made of atoms
joined covalently are called molecular or covalent compounds.
We name covalent compounds differently than ionic compounds. We must
indicate the number of each element by adding a prefix in front of the
element’s name.
Subscript
Prefix
one
two
three
four
five
six
seven
eight
nine
Ten
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To name covalent compounds, follow these rules:
1. Name the first element in full using a prefix only when there are two or
more of that element. That is, omit "mono" if only one of that element
is in the compound. e.g. NO is nitrogen monoxide, but N2O is
dinitrogen monoxide.
2. Name the second element and end in "-ide". Use prefixes to indicate
the number of that element (including mono).
3. Write the name of the compound writing the substance found more to
the left on the periodic table first.
Note: There are two exceptions to the naming rules. Here the common
names for the compounds are used:
H2O = water
NH3 = ammonia
CH4 = methane
Example:
Write the name for CO2
Step 1: Name the first atom with prefixes.
Step 2: Name the second element using prefixes and end in “-ide”.
Step 3: Write the name of the compound writing the substances found more
to the left on the periodic table first.
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Example:
Write the name for N2O4.
Example:
Write the name for PCl5.
Example:
Write the name for NH3.
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Writing Covalent Compound Formula
Writing formulas for covalent compounds involves the following rules:
1. Write the symbol for the first element followed by the subscript
indicated by the prefix.
2. Write the symbol of the second element followed by the subscript
indicated by its prefix.
DO NOT REDUCE THE SUBSCRIPTS!!!
Example:
Write the formula for dinitrogen monoxide.
Example:
Write the formula for sulphur hexafluoride.
Example:
Write the formula for phosphorus tribromide.
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Diatomic Molecules
Some molecules do not exist as single atoms. These elements exist as pairs
of atoms joined covalently, called diatomic molecules. The elements that
exist as diatomic molecules are hydrogen (H2), oxygen (O2), fluorine (F2),
chlorine (Cl2), bromine (Br2), iodine (I2), and nitrogen (N2).
Here is something to help you remember the diatomic molecules:
I Have No Bright Or Clever Friends
I= iodine
H = hydrogen
N = nitrogen
Br = bromine
O = oxygen
Cl – chlorine
F = fluorine
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Grade 11 Chemistry
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Chemical Equations
A chemical equation indicates the substances reacting and the
substances produced in a chemical reaction. A chemical reaction also shows
the ratio in which these substances react or are produced.
A word equation can describe a chemical reaction.
Example:
Hydrogen gas and oxygen gas react to form (yield) water vapour.
A chemical equation can also show heat changes that occur.
Endothermic reactions is when the reaction absorbs energy. The energy is
used in the reaction, so energy is a reactant.
2KClO3 + heat
2KCl + 3O2
Exothermic reactions is when the reaction releases energy. The energy (or
heat) is a product.
2H2(g) + O2(g)
2H2O(g) + energy
Recall:
Symbol
+
Meaning
(s)
(l)
(g)
(aq)
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Grade 11 Chemistry
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Balancing Chemical Reactions
The Law of Conservation of Mass (or Matter) states that:
In any chemical reaction, matter cannot be created or destroyed.
This means the mass will not change during a chemical reaction.
In other words, in a chemical reaction, mass of reactants is equal to
mass of products.
The amount of matter (mass) on both sides of the equation needs to be the
same. As a result of this law, all chemical equations must be balanced.
To balance atoms, we insert coefficients rather than changing subscripts.
You will eventually develop a method of balancing that works best for you.
Until then, you can use the following suggestions.
1. If the reaction is a word equation, write out the formulas of each
reactant and product.
2. Write each element underneath the equation and keep a tally of the
number of atoms of each element.
3. Use coefficients to balance metals first, then non-metals.
4. Leave single elements and diatomic molecules to balance last.
5. If possible, reduce the coefficients to the lowest whole number ratio.
Multiply fractions, if present, by the denominator to make all
coefficients whole numbers.
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Example:
Balance the equation
C 3 H8
+
O2
+
CaCl2
CO2
+
H2O
Example:
Balance the equation
Al2(SO4)3
AlCl3
20
+
CaSO4
Grade 11 Chemistry
Unit 1: Chemical Reactions Part 1
Types of Chemical Reactions
Combustion Reactions
A combustion reaction is a reaction in which oxygen reacts with a
hydrocarbon to produce carbon dioxide, water, and a lot of heat.
General Formula: hydrocarbon + O2
CO2 + H2O
Example:
CH4(g) (methane) + O2(g)
CO2(g) + H2O(g)
Synthesis Reactions
A synthesis reaction involves the combining of smaller atoms/molecules
into larger, more complex molecules. If only two different atoms appear on
the reactant side, then the reaction must be synthesis.
General Formula:
A+B
AB
Example:
Water and dinitrogen pentoxide gas react to produce aqueous hydrogen
nitrate.
Decomposition Reactions
A decomposition reaction involves the splitting of large molecules into
smaller molecules or elements.
General Formula:
AB
A+B
Example:
Solid nickel (II) hydroxide decomposes to produce solid nickel (II) oxide and
water.
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Grade 11 Chemistry
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Single Displacement (Replacement) Reactions
Single displacement reactions are chemical changes that occur when an
element replaces a less active element in a compound, leaving the replaced
element alone.
Refer to the Activity Series handout
General Formula:
A + BC
B + AC
Example:
Fluorine gas will react with sodium bromide in an aqueous solution to
produce sodium fluoride and liquid bromine.
Double Displacement (Replacement) Reactions
Double displacement reactions occur when elements in different
compounds displace each other or exchange places. Generally, the reaction
occurs in an aqueous system between two ionic compounds.
General Formula:
AB + CD
AD + CB
Example:
When aqueous lithium iodide and aqueous silver nitrate react, they will
produce solid silver iodide and aqueous lithium nitrate.
Acid-Base Reactions
When an acid and base react together, the reaction is known as a
neutralization reaction. The products will always be water and a salt.
General Formula:
ACID + BASE
WATER + SALT
Example:
When a solution of aqueous hydrochloric acid and solid potassium hydroxide
react, water and aqueous potassium chloride are formed.
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Grade 11 Chemistry
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Predicting Reaction Products
When determining what the products will be in a chemical reaction, use the
following information to help you.
Synthesis Reaction
Reactants: Generally two elements or two compounds (where at least one
compound is a molecular compound).
Probable Products: A single compound
Decomposition Reaction
Reactants: Generally a single binary compound or a compound with a
polyatomic ion.
Probable Products: Two elements (for a binary compound), or two or more
elements and/or compounds (for a compound with a polyatomic ion).
Combustion
Reactants: Oxygen and a compound of C, H, (O)
Probable Products: CO2 and H2O
Single Replacement Reaction
Reactants: An element and a compound
Probable Products: a different element and a new compound.
Double Replacement Reaction
Reactants: Two ionic compounds
Probable Products: Two new compounds
Determining the Physical States
To determine if a product is a solid or not, use a table of solubility. The
following are hints to help determine the physical state.
(s) – most metals, precipitates
(l) – mercury, bromine, water
(g) – noble gases, diatomic molecules (except bromine), ammonia
(aq) – substance is in a water based solution (use solubility chart)
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Isotopes
The number of neutrons in each atom varies, even between atoms of the
same element. For example, potassium can exist as three different atoms.
All three atoms contain 19 protons, but one potassium atom has 20
neutrons, another 21 neutrons and yet another has 22 neutrons.
Atoms that have the same number of protons but differ in their number of
neutrons are called isotopes.
If different isotopes have different numbers of neutrons, they will have
different masses. The mass number of an atom is the sum of the
protons and neutrons found in the nucleus of that atom.
If we look at the potassium isotopes above, the isotope containing 19
protons and 20 neutrons will have a mass number of 39. We call this
isotope potassium-39.
The isotope that has 19 protons and 21 neutrons will have a mass
number of 40 and is called potassium-40.
Chemists have designed a symbol for each isotope that includes the
element’s symbol (X), its atomic number (Z) and its mass number (A).
The symbol for potassium-39 would be:
The symbol for potassium-40 would be:
The symbol for uranium-238 would be:
The symbol for uranium-239 would be:
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Grade 11 Chemistry
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Atomic Mass
The masses of individual atoms are expressed as atomic mass units
(amu) or µ. The atomic mass unit is defined as 1/12 the mass of a
carbon-12 atom. A proton or a neutron has mass equal to approximately
one atomic mass unit.
In many cases the amount of each isotope in the sample, or its relative
abundance, can be determined using a mass spectrometer. The relative
abundance of an isotope is the percent of each isotope found in an
average sample of the element.
You have noticed that the atomic mass shown for each element on a
periodic table is rarely a whole number. This is because it is actually an
average mass of all isotopes of that element.
How to Calculate Average Atomic Mass
To determine the average atomic mass, you first need to determine what the
mass contribution is for the isotope.
mass contribution = (mass)(relative abundance)
Once all the mass contributions have been determined, you simply add up
the numbers to find your average atomic mass.
Example:
Find the atomic mass of magnesium, using the information provided:
Relative Abundance of Stable Magnesium Isotopes
Isotope
Relative
Atomic Mass
Mass
Abundance (%)
(amu)
Contribution
Mg-24
78.70
23.98504
Mg-25
10.13
24.98584
Mg-26
11.17
25.98259
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Grade 11 Chemistry
Unit 1: Chemical Reactions Part 1
Calculating the Mass of Compounds
Using the Periodic Table, you can find the atomic mass of any element. You
can also use this information to determine the molecular mass, or formula
mass of any compound.
The mass of molecular (covalent) compounds is referred to as molecular
mass. The mass of an ionic compound is referred to as formula mass.
Example:
Calculate the molecular or formula mass of the following substances:
Mg
H2O
C12H22O11
NaCl
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The Mole
Formula and molecular mass deal with individual atoms and molecules.
Chemists do not work with amounts of individual atoms or molecules.
These particles are far too small to see or mass. Balances or scales do not
measure mass in terms of atomic mass units.
Chemists need a practical unit that relates mass, in grams, with the
number of particles. We will use the word particle as a generic term
meaning atoms, molecules, ions, or any other entity.
What number of atoms or molecules would be large enough? 1 dozen?
1 gross (144)? 1 million?
In the early 1800’s, a scientist named Lorenzo Romano Amedeo Carlo
Avogadro devised a way of measuring the masses of very small particles
of matter.
Avogadro suggested that is you took exactly 12.00 g of the carbon-12
atom you would have a very large number of carbon atoms. Just as
words like “dozen”, “gross” or “couple” have numerical meaning, he
needed a word to describe how many atoms there would be in 12.00 g of
carbon-12.
This word is now known as the “mole” (abbreviated mol). A mole is just
a number! 1 Mole = 6.02 x 1023 atoms
Just like, one dozen atoms would be 12 atoms or one gross of atoms
would be 144 atoms, one mole of atoms is 6.02 x 1023 atoms, known
as Avogadro’s number and abbreviated N or NA.
The mole is the central unit in most calculations;
it is at the heart of chemistry.
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Grade 11 Chemistry
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The Green Pea Analogy
The Green Pea Analogy is intended to give you an idea of exactly how
much 6.02 x 1023 atoms of something actually is.
 One hundred (102) green peas have a volume of about 20 cm3 or
20 mL in a graduated cylinder.
 One million (106) green peas have enough volume to fill up an
ordinary refrigerator.
 One billion (109) green peas can fill an average 3 bedroom house.
 One trillion (1012) green peas can fill 1000 average homes.
 One quadrillion (1015) green peas can fill all the buildings in a
larger city such as Edmonton.
 Imagine all of Alberta covered one metre deep in green peas. Now
you got one quintillion (1018) green peas.
 Imagine that all the continents are now covered 1 metre deep with
green peas. Now you have one sextillion (1021).
 If you froze all the oceans and totally covered the whole earth with
green peas and had 250 planets just like earth, all covered with
green peas, you would now have a mole’s worth!
 250 000 planets like earth, all covered one metre deep in green
peas would give you a cotillion (1027), the number of atoms in your
body!
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Grade 11 Chemistry
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Molar Mass
Recall, Avogadro’s Number relates the number of particles to mass. By
definition, one mole of carbon atoms has a mass of 12.0000 g. If the mass
of one mole of any atom is its atomic mass in grams, then
one mole of aluminum has a mass of 27.0 g
one mole of silver atoms has a mass of ____________g
one mole of sodium atoms has a mass of __________g
One mole of any compound has a mass equal to its formula/molecular mass,
in grams. Water has a molecular mass of 18.0 amu. The mass of one mole
of water molecules is ________________ g.
The mass of one mole of a substance is called the molar mass (M). The
units for molar mass are grams per mole (g/mol), so the molar mass of
water is 18.02 g/mol.
Find the molar masses for each of the following:
Example 1: CaSO4(s)
Ca = 1 x _________________ = ________________
S = 1 x _________________ = _________________
O = 4 x _________________ = _________________
CaSO4(s) = ___________________
Example 2: Al(OH)3(s)
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Example 3: ammonium dichromate
Example 4: Na2H2PO4(s)
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Converting Mass to Moles
The units for molar mass are g/mol. This means the mass, in grams, of one
mole of a substance or mathematically,
molar mass =
mass in grams
1 mole
To determine the number of moles in a sample, we will use the following
steps:
Step 1: Find the molar mass
Step 2: Use the molar mass t calculate the number of moles.
Example:
How many moles of carbon atoms in 24.0 g of carbon?
Step 1: Find the molar mass.
Step 2: Use the molar mass to calculate the number of moles.
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Example:
How many moles of water molecules in 81.0 g of water?
Step 1: Find the molar mass.
Step 2: Use the molar mass to calculate the number of moles.
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Converting Moles to Mass
If one mole of carbon has a mass of 12.0 g, then 2 moles of carbon will have
a mass of ______________ g, 3 moles will have a mass of ___________ g,
etc.
If we need a certain number of moles of a substance, we can calculate the
mass of the substance by multiplying the number of moles of that substance
by its molar mass.
When converting from moles to mass,
Step 1: Determine the molar mass of the substance.
Step 2: Multiply so units cancel to yield grams.
Example
What is the mass of 2.50 moles of gold?
Step 1: Determine the molar mass of the substance.
Step 2: Multiply so units cancel to yield grams.
Example
What is the mass of 1.20 x 10-5 moles of carbon tetrachloride?
Step 1: Determine the molar mass of the substance.
Step 2: Multiply so units cancel to yield grams.
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Converting Moles to Particles
One mole of anything contains 6.02 x 1023 individual entities, or
Avogadro’s Number, NA, of that substance. In this course, the generic
term “particles” refers to any individual entity like atoms, molecules, ions,
etc. If one mole of particles is 6.02 x 1023 particles, then
2 moles of particles is 2 x 6.02 x 1023 = 1.20 x 1024 individual particles, or
2 x NA
3 moles of particles is 3 x 6.02 x 1023 = 1.81 x 1024 individual particles, or
3 x NA
The ratios we will use:

To find the number of particles
6.02 x 1023 particles
1 mole

To find the number of moles
_____1 mole_____
6.02 X 1023 particles
Example:
How many atoms in 25.0 moles of copper atoms?
The type of atom, molecule, ion, etc. is not significant in this calculation.
Use the correct ratio to cancel moles.
Example:
How many molecules of water in 1.50 x 10-5 moles of water?
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Example:
How many moles are there in 5.00 x 1025 particles?
Choose the correct ratio and multiply by the given amount:
Example:
How many moles of atoms are in 5 atoms of zinc?
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Grade 11 Chemistry
Unit 1: Chemical Reactions Part 1
Converting Mass to Particles
The mole allows the conversion between mass and number of particles. The
flow chart below may help to visualize the process for the conversion.
To convert mass to the number of particles:
Step 1: Convert mass to moles.
Step 2: Convert moles to number of particles.
Example
How many molecules of water are in a 10.0 g sample of water?
Step 1: Convert mass to moles.
Do not round at this point. Keep the number in your calculator.
Step 2: Convert moles to number of particles.
or, we can put the two equations together. Notice how all units except
molecules (particles) cancel each other.
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Example:
How many molecules are in 25.0 g of sodium chloride?
Example:
How many molecules are in 32.57 g of magnesium oxide?
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Grade 11 Chemistry
Unit 1: Chemical Reactions Part 1
Converting Particles to Mass
From the flowchart
we can determine the mass of a certain number of particles of a substance.
To find the mass of a given number of particles:
Step 1: Convert the number of particles to the number of moles.
Step 2: Convert the number of moles of the substance to mass.
Example:
What is the mass of 3.01 x 1023 molecules of water?
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Example:
What is the mass of a single atom of carbon?
Example:
What is the mass of a 4.52x1022 molecules of nitrogen monoxide?
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