Download Particles

Ex02
Particles
Atoms, ions, and molecules.
Working with the pieces that define matter.
version 1.5
© Nick DeMello, PhD. 2007-2017
No one had ever seen an atom. The wavelength of visible light is more than 1000 times bigger than an atom,
so that light cannot be used to observe an atom.
However, ESM probe microscopes
can now be used to feel atomic
surfaces and even move individual atoms.
These pictures show 48 iron
atoms on the surface of a copper
crystal being arrange in a circle.
We can now “see” and even “touch”
the atom.
Ex02
Particles
‣ Atoms
‣ Atoms, Molecules & Ions
‣ Compounds & Elements
‣ Experiment
‣ Copper Sulfide
‣ Next Week
‣ Atomic Structure
‣ Molecules
‣ Bonding
‣ Formulas
C6H2(NO2)3CH3
‣ Weighing Atoms
‣ AMU & Moles
‣ Molar Mass
‣ Finding Formulas
1 mol = 6.022 x1023 singles
3
Atoms, Molecules, & Ions
4
‣
A pure substance is made up of identical
particles.
‣
We haven’t talked about how those particles
can differ from particles in a different pure
substance.
‣
Those particles can be made up of different
flavored pieces called atoms.
Atoms, Molecules, & Ions
5
‣
A pure substance is made up of identical
particles.
‣
We haven’t talked about how those particles
can differ from particles in a different pure
substance.
‣
Those particles can be made up of different
flavored pieces called atoms.
‣
Some of the particles are just that, an atom.
‣
Others are collection of atoms bonded into a
single particle, a molecule.
‣
There are also ions. Ions are particles with an
electric charge on them (positive or negative).
Atoms, Molecules, & Ions
Atoms
‣
A pure substance is made up of identical
particles.
‣
We haven’t talked about how those particles
can differ from particles in a different pure
substance.
‣
Those particles can be made up of different
flavored pieces called atoms.
‣
Some of the particles are just that, an atom.
‣
Others are collection of atoms bonded into a
single particle, a molecule.
‣
There are also ions. Ions are particles with an
electric charge on them (positive or negative).
Molecules
Ions
+
6
-2
3+
Elements & Compounds
We use the word element two
ways: it's used to describe the
flavor of an atom and it's also
how we name a substance made
only of that flavor atom.
Particle
Substance
1 kind of atom
Element or Ion
Element
2 or more kinds of atoms
Molecule or Ion
Compound
‣ Atoms are the building blocks of matter.
‣ Each element is made of only one kind of atom.
‣ A compound is made of two or more different kinds of elements.
Monatomic, Diatomic, Polyatomic
Elements
Compounds
Ions
+
Monatomic
Particles
Atoms
Diatomic
Particles
+
Polyatomic
Particles
8
3+
2-
-2
Molecules
Ions
Monatomic, Diatomic, Polyatomic
Elements
Polyatomic
Particles
9
Ions
+
Monatomic
Particles
Diatomic
Particles
Compounds
Binary Compounds and
Binary Ions are made of two
elements (not necessarily two atoms).
+
-2
3+
2-
Part of an Atom
Part of an Atom
Ions vs Atoms
‣ Around the beginning of the 1900’s chemists
discovered some atoms could hold an electrical
charge.
‣ Charges can be positive or negative
‣ Charges can be different sizes
‣ The properties of charged atoms were documented
by Michael Faraday, who named them ions.
‣ Charged atoms move in solution, toward or away
from electrically charged wires.
‣ The word “ion” is greek for wanderer.
‣ Ions that move towards a cathode (neg charged
wire) are positively charged ions.
‣ They’re called cations.
+IONS
CA
‣ Ions that move towards an anode (pos charged wire)
are negatively charged ions.
‣ They’re called anions.
‣ Atoms and ions made from those same atoms have
different properties.
‣ Silver, Ag
‣ Not soluble in water
‣ Not attracted to magnets
1+
‣ Silver Ions, Ag
‣ Soluble in water
‣ Attracted to magnets
12
H
H1+
H1-
Faraday
Ex02
Particles
‣ Atoms
‣ Atoms, Molecules & Ions
‣ Compounds & Elements
‣ Experiment
‣ Copper Sulfide
‣ Next Week
‣ Atomic Structure
‣ Molecules
‣ Bonding
‣ Formulas
C6H2(NO2)3CH3
‣ Weighing Atoms
‣ AMU & Moles
‣ Molar Mass
‣ Finding Formulas
1 mol = 6.022 x1023 singles
13
Chemical Bonding
‣ Bringing elements together to form new compounds.
‣ Not a mixture, but bonding the elements at the
atomic level.
‣ When you mix hydrogen and oxygen, you don’t have
a new substance — no new properties are observed.
‣ Mixtures are useful, but it’s not a reaction.
‣ It’s just stirring up the particles.
‣ When you react hydrogen and oxygen, you have a
new substance — you see new properties.
‣ Water is a compound
‣ won’t burn
‣ liquid at room temperature
‣ causes sodium to burn
‣ The compound forms because the atoms bond to
each other.
‣ All chemical bonding uses electrons to glue atoms to
each other
‣ There are different types of bonding.
‣ metallic
‣ ionic
‣ covalent
14
Bonding Atoms
Iron Pyrite
- does not burn
Iron
- hard to burn
heat & pressure
stir together
- maleable
- burns bright yellow
- not attracted to magnets
Fireworks Additive
Sulfur
- easy to burn
- easy to burn
- burns bright yellow
- burns dull
- cannot be separated mechanically
- has constant properties
- is always 46.6% iron and 53.4 % sulfur
- a pure substance
‣ How did we go from a mixture to a pure substance?
‣ We changed the particles — we created a new substance.
‣ We know a new substance was created because we see properties that didn’t exist before.
‣ Not just more or less of a property that was already there, but something entirely new.
‣ We can isolate a pure substance that did not exist in the original mixture.
‣ A new substance, responsible for the new properties.
15
- lustrous (shiny)
Bonding Atoms
Chemical
Change
Iron
- hard to burn
heat & pressure
stir together
Fireworks Additive
- easy to burn
- burns bright yellow
- burns dull
- cannot be separated mechanically
- has constant properties
- is always 46.6% iron and 53.4 % sulfur
- a pure substance
‣ How did we go from a mixture to a pure substance?
‣ We changed the particles — we created a new substance.
‣ We know a new substance was created because we see properties that didn’t exist before.
‣ Not just more or less of a property that was already there, but something entirely new.
‣ We can isolate a pure substance that did not exist in the original mixture.
‣ A new substance, responsible for the new properties.
16
- lustrous (shiny)
- not attracted to magnets
Physical
Change
- easy to burn
- does not burn
- maleable
- burns bright yellow
Sulfur
Iron Pyrite
Combining (Bonding) Atoms
‣ Metallic Bonding (only metals)
‣ In pure metals (Fe, Au, Co) or alloys (mixtures of metals) electrons
break off and float between the atoms.
‣ These free flowing electrons make metals extremely good conductors
of electricity.
‣ Metal atoms pull on the electrons flowing between them causing the
mass to stick together.
‣ Metallic bonding does not form compounds.
‣ Ionic Bonding (metal with non-metal)
‣ When you mix metals and non-metals electrons break off from metals
and are captured by non-metals.
‣ This creates positively and negatively charged particles.
‣ The particles attract each other, this is an ionic bond.
‣ Ionic bonds are extremely strong.
‣ These ions clump together in simple, large complexes.
‣ Compounds made from ionic bonds are ionic compounds.
‣ Covalent Bonding (only non-metals)
‣ Nonmetals pull on each others electrons.
‣ If neither non-metal pulls hard enough to remove the electron from the other, the two end up sharing a pair of electrons.
‣ The shared electrons are localized between two atoms, creating a
bond between just those two atoms.
‣ This produces discrete microscopic structures built of atoms.
‣ Particles made of covalent bonds are molecules.
‣ Compounds made from covalent bonds are molecular compounds.
Bonding in Compounds
‣
Covalent Bonding (only non-metals)
‣
Nonmetals pull on each others electrons.
‣
The shared electrons are localized
between two atoms, creating a bond
between just those two atoms.
‣
‣
‣
Covalent bonding produces discrete
microscopic structures built of atoms.
Particles made of covalent bonds are
molecules.
Compounds made from covalent bonds
are molecular compounds.
‣ Ionic Bonding (metal with non-metal)
‣ Ionic bonding creates positively and
negatively charged particles.
‣ The particles attract each other, this is
an ionic bond.
‣ These ions clump together in simple,
large complexes.
‣ Compounds made from ionic bonds are
ionic compounds.
Chemical Formulas
‣ We use symbols to represent elements and also to represent atoms
of that element.
Au
‣ You must memorize the symbols of the first 18 elements! (this is
easier than it sounds)
‣ The order of elements goes from the most metal-like element to
the least. Na before C before H before F, etc (we’ll talk more
about this later)
AlBr3
CH4
Cl –
‣ We use subscripts to indicate the number of atoms of that element.
Na+
‣ Subscripts of 1 are omitted.
‣ Omitted subscripts mean 1.
HF
‣ We use superscripts to indicate the net charge (if any) on the
entire particle.
‣ Superscripts of 0 (charge 0) are omitted.
‣ Omitted superscripts are assumed to mean 0 (no charge).
H 2O
Water is a binary compound, it is a
polyatomic molecule composed of
2 hydrogen atoms and 1 oxygen
atom. It has a charge of zero.
19
SO4
Br2
NO2–
Charge
2-
C6H8NO4
PO42–
Sulfate is a binary ion, it is a
polyatomic ion composed of 1 Atom Count
sulfur atom and 4 oxygen atoms.
It has a charge of minus two.
Al3+
Sn
Chemical Formula
‣ We use chemical formulas to describe both
types of compound.
‣ There are three kinds of chemical formulas.
‣ Empirical formulas describe the ratio of
elements in the compound.
‣ Empirical formulas can describe either
molecular or ionic compounds.
Butane
Empirical
Molecular
Structural
20
C2H5
C4H10
Salt
NaCl
does not apply
‣ The smallest whole number ratio of
elements is also called a formula unit.
‣ Molecular formulas describe the number of
atoms in each molecule.
‣ Molecular formulas can only be used to
describe molecular compounds.
‣ Structural formulas graphically describe the
connectivity between atoms.
‣ Structural formals can only be used to
describe molecular compounds.
does not apply
‣ We will talk more about these shortly.
Ex02
Particles
‣ Atoms
‣ Atoms, Molecules & Ions
‣ Compounds & Elements
‣ Experiment
‣ Copper Sulfide
‣ Next Week
‣ Atomic Structure
‣ Molecules
‣ Bonding
‣ Formulas
C6H2(NO2)3CH3
‣ Weighing Atoms
‣ AMU & Moles
‣ Molar Mass
‣ Finding Formulas
1 mol = 6.022 x1023 singles
21
Weighing Atoms
‣ We use different units to measure different sizes of
mass.
‣ A car weighs 4,000 lbs or 2 tons.
‣ If you have many cars it’s easier to work in tons.
‣ Most of the material we use in lab, weighs as much
as a penny. A penny weighs…
‣ 0.0000028 tons
‣ 0.0055 lbs
‣ 2.5 g
‣ Grams are a good unit for many lab preparations.
‣ A hydrogen atom weighs about 1.67 x 10-24 grams.
‣ Grams are not a good unit for talking about the
weights of individual atoms.
‣ We need a unit of weight that is a better fit.
‣ We use AMU (atomic mass units)
22
The AMU
‣ AMU is the atomic mass unit.
‣ Every atom is made of neutrons & protons.
‣ It’s convenient when we’re working on a
molecular scale to have a unit of weight
about the size of a neutron or proton.
‣ We call that unit amu (atomic mass unit).
‣ Most interesting molecules are made of
carbon.
‣ The most common isotope of carbon is made
almost entirely of 6 protons and 6 neutrons.
‣ An amu is defined as: exactly ⅟12 the mass of a Carbon-12 atom
unit
23
12
C
6
‣ 1 amu is measured to be 1.6606 x 10-24 g.
Mole is 6.022 x1023 Singles
1 mol = 6.022
x1023
singles
‣ We will need to jump back and forth between
designing reactions at the AMU scale and doing them in
the lab.
‣ We need a way of quickly jumping back and forth
between atoms and large (gram sized) collections of
atoms.
‣ 6.022 x1023 amu sized things fit in a gram.
‣ So we will design reactions by looking at atoms and
molecules, but we’ll measure out substances in
bunches of 6.022 x1023 atoms and molecules.
‣ A pair is two of something.
‣ A dozen is twelve singles.
‣ A score is 20 singles.
‣ A mole (mol) is 6.022 x1023 singles.
‣ 6.022 x1023 is Avogadro’s Number.
24
Why we work in moles
‣ Working in moles (mol) means if you
know the weight of something in amu,
you know the weight of a mole of that
same thing in grams (g).
‣ Hydrogen atom weighs 1 amu
‣ A mol of hydrogen weighs 1 g
‣ Helium atom weighs 4 amu
‣ A mol of helium weighs 4 g
‣ Copper atom weighs 63.55 amu
‣ A mol of copper weighs 63.55 g
1 mol = 6.022 x1023 singles
25
The Molecular Blueprint
‣ Chemical Formulas Identify Compounds
‣ We use them as shorthand to name of a substance (“Pass me the H2O”)
‣ Chemical Formulas indicate the composition of a substance.
‣ Each element is indicated with it’s symbol.
‣ The a subscript indicates the total number of atoms of
that element.
C6H2(NO2)3CH3
6+1 Carbon Atoms
2+3 Hydrogen Atoms
‣ Subscripts of 1 are omitted.
‣ Omitted subscripts mean 1.
‣ Parenthesis are used to indicate groups of atoms.
‣ Chemical Formulas may contain hints of the connectivity of
the atoms.
3 NO2 Groups
‣ Chemical Formulas may show a CH3 group of atoms and three
NO2 groups of atoms are bonded to a C6H2 group by writing:
3 (3x1) Nitrogen Atoms
6 (3x2) Oxygen Atoms
C6H2(NO2)3CH3
instead of:
26
C7H5N3O6
Problem:
You have 2.85 mols of C6H2(NO2)3CH3 (trinitrotoluene). How many mols of oxygen atoms are in it?
Problem:
You have 2.85 mols of C6H2(NO2)3CH3 (trinitrotoluene). How many atoms of oxygen do you have?
Ex02
Particles
‣ Atoms
‣ Atoms, Molecules & Ions
‣ Compounds & Elements
‣ Experiment
‣ Copper Sulfide
‣ Next Week
‣ Atomic Structure
‣ Molecules
‣ Bonding
‣ Formulas
C6H2(NO2)3CH3
‣ Weighing Atoms
‣ AMU & Moles
‣ Molar Mass
‣ Finding Formulas
1 mol = 6.022 x1023 singles
29
Atomic Weights / Molar Weights
‣
Weights are listed in the periodic table without units.
‣
The weight listed is the average mass of one atom of
each element, in amu.
‣
‣
That means:
1 gram ÷ 1.6606 x 10-24 grams = 6.022 x 1023
1 gram ÷ 1 amu = 1 mol
1 gram = 1 mol x 1 amu
1 mol of anything will weigh in grams, what a single of that anything weighs in amu. ‣
If a cat weighs X amu, a mol of cats weighs X grams.
‣
That means each weight in the periodic table is:
‣ the weight of 1 atom of that element, in amu
‣ the weight of 1 mol of that element, in grams
‣
Reading from the periodic table...
1 H = 1.008 amu
1 mol H = 1.008 g
1 Cu = 63.55 amu
1 mol Cu = 63.55 g
‣ a hydrogen atom (H) weighs 1.008 amu
‣ a mol of hydrogen atoms (H) weigh 1.008 g
‣ a copper atom (Cu) weighs 63.55 amu
‣ a mol of copper atoms (Cu) weighs 63.55 g
.
30
.
New Conversion Factors
You are responsible for these
conversion factors, a periodic table will be
provided.
Avogadro’s Number
1 mol = 6.022 x1023 singles
Atomic Mass
1 copper atom = 63.55 amu
Molar Mass
1 mol copper atoms = 63.55 grams
31
Atomic Mass & Avogadro’s Number
Elements like Copper (Cu)
‣
Molar Mass
1 mol Cu = 63.55 grams
Molar Mass/Atomic Mass grams
(the average mass of atoms of that elements)
‣ We get this from the periodic table
mol
‣ It tell’s us the weight of:
‣ 1 mol of a substance (in grams)
‣ 1 atom of a substance (in amu)
Avogadro’s number
1 mol = 6.022 x 1023
molar scale
molecular scale
‣ Avogadro’s Number
‣ 6.022 x 1023
‣ It’s a measurement
‣ You have to memorize it
molecules
& atoms
‣ It let’s us go from the moles to molecules
Atomic Mass
1 Cu = 63.55 amu
amu
Counting by Weight
1 mol = 6.022 x 1023 singles
How many atoms in exactly 1 mol Copper (Cu)?
How many atoms in 2.53 mol Copper (Cu)?
1 Cu = 63.55 amu
1 mol Cu = 63.55 g
How many mol Cu in 30.5 grams Cu?
How many Cu atoms in 30.5 grams Cu?
.
How many atoms?
A gold ring weighs 1.24 grams. How many atoms of gold are in it?
34
How many grams?
An experiment calls for 4.3 mols of Calcium atoms, how many grams of pure calcium should you weigh out?
35
Weight of 4 atoms?
A phosphorus molecule is composed of 4 atoms of phosphorus. What is it’s weight in AMUs?
36
Molecular Weight/Molar Mass
‣ Molar Mass also applies to molecules and compounds.
‣ We know the atomic weight of elements, what one atoms weighs in amu and
what one mole of atoms weigh in grams.
‣ We can use that information to figure out for compounds what
one molecule weighs or one mole of molecules weigh.
What is the molecular weight of CO2? (in amu)
1 O = 16.00 amu
1 mol O = 16.00 g
.
1 C = 12.01 amu
1 mol C = 12.01 g
What is the molar mass of CO2? (in grams)
What does 2.57 mol of CO2 weigh?
How many moles of CO2 are in 53.256 grams?
.
Molecular Formula & Molar Mass
Molecules like Water (H2O)
‣ ‣
Molecular Formula
grams
Molar Mass
1 mol H2O = 18.02 grams
(& Empirical Formula)
H 2O
‣ It let’s us understand the composition of
molecules.
mol
mol
HO
‣ We can use it as a conversion factor to go
from molecules to how many atoms of any
kind are in that molecule.
2
molar scale
Avogadro’s number
1 mol = 6.022 x 1023
molecular scale
‣ Molar Mass/Molecular Mass
molecules
‣ It relates weight to mols for whole
molecules.
H 2O
Molecular Mass
1 H2O = 18.02 amu
Molecular
Formula
1 H 2O = 2 H
atoms
amu
Problem:
Your experiment requires 4.26 mols of magnesium chloride (MgCl2). What mass of
magnesium chloride do you weigh out for this experiment?
Solution
Problem:
You do an experiment that produces 15.35 grams of nitrogen trioxide (NO3).
How many moles of NO3 were produced?
Solution
Ex02
Particles
‣ Atoms
‣ Atoms, Molecules & Ions
‣ Compounds & Elements
‣ Experiment
‣ Copper Sulfide
‣ Next Week
‣ Atomic Structure
‣ Molecules
‣ Bonding
‣ Formulas
C6H2(NO2)3CH3
‣ Weighing Atoms
‣ AMU & Moles
‣ Molar Mass
‣ Finding Formulas
1 mol = 6.022 x1023 singles
41
Finding the Formula
‣ It’s possible to burn substances by capturing the gases find out the mass of
individual elements in the sample.
‣ If you burn a sample of butane gas, you may find it’s composed of 16.8 grams
carbon and 3.53 grams hydrogen.
‣ From that you can determine the empirical formula of butane.
‣ Empirical Formula is the smallest whole number ratio of atoms of each element in
the substance.
Report:
C 16.8 g
H 3.53 g
42
Burning a sample of butane you find it is composed only of 16.8 g C and 3.53 g H, what is the empirical formula of butane?
Report:
C 16.8 g
H 3.53 g
Chemical Formula
‣ We use chemical formulas to describe both
types of compound.
‣ There are three kinds of chemical formulas.
‣ Empirical formulas describe the ratio of
elements in the compound.
‣ Empirical formulas can describe either
molecular or ionic compounds.
Butane
Empirical
Molecular
Structural
44
C2H5
C4H10
Salt
NaCl
does not apply
‣ The smallest whole number ratio of
elements is also called a formula unit.
‣ Molecular formulas describe the number of
atoms in each molecule.
‣ Molecular formulas can only be used to
describe molecular compounds.
‣ Structural formulas graphically describe the
connectivity between atoms.
‣ Structural formals can only be used to
describe molecular compounds.
does not apply
‣ We will talk more about these shortly.
Finding the Molar Mass
‣ The molar mass can only be determined (without already knowing the
molecular formula) by experiment.
‣ The device we use for this experiment is a mass spectrometer.
‣ The material is atomized and shot through a magnet.
‣ An electron is knocked of one molecule.
‣ By varying the magnetic field you see how much energy is bend the path of
the molecule.
‣ Once you know how much force it takes to move it, you can get the
momentum.
‣ If you know the speed and momentum, you can get the mass.
A mass spectrometer experiment gives us the mass of a molecule like caffeine,
the same way it gave us the mass of an atom like copper.
45
The molar
mass of
Butane is:
58.12 g/mol
Burning a sample of butane you find it is composed only of 16.8 g C and 3.53 g H, what is the molecular formula of butane?
The molar
mass of
Butane is:
58.12 g/mol
Ex02
Particles
‣ Atoms
‣ Atoms, Molecules & Ions
‣ Compounds & Elements
‣ Experiment
‣ Copper Sulfide
‣ Next Week
‣ Atomic Structure
‣ Molecules
‣ Bonding
‣ Formulas
C6H2(NO2)3CH3
‣ Weighing Atoms
‣ AMU & Moles
‣ Molar Mass
‣ Finding Formulas
1 mol = 6.022 x1023 singles
47
Exp 04: Chemical Formula
Your job is to react copper wire (Cu) with sulfur powder (S8) using heat
from a bunsen burner to form a copper sulfide compound.
Then determine and justify the molecular formula of that compound.
CuXSY
Experiment
‣ Support a clean, dry, porcelain crucible and cover
on a clay triangle and dry by heating to a dull red in
a Bunsen burner flame. Allow the crucible and cover
to cool to room temperature and weigh them.
Record the mass to the nearest 0.01 g.
‣ Place 1.5 to 2.0 g of tightly wound copper wire or
copper turnings in the crucible and weigh the
copper, crucible, and lid to the nearest 0.01 g and
record your results. Calculate the mass of copper.
‣ Later you’ll convert this mass of copper to moles of
copper atoms in your proudct.
‣ In the hood, add sufficient sulfur to cover the
copper, place the crucible with cover in place on
the triangle, and heat the crucible gently until
sulfur ceases to burn (blue flame) at the end of the
cover. Do not remove the cover while the crucible is
hot.
‣ Heat the crucible to dull redness for about 5
minutes.
49
Experiment
‣ Allow the crucible to cool to room temperature.
This will take about 10 min. Then weigh with the
cover in place. Record the mass.
‣ Again cover the contents of the crucible with sulfur
and repeat the heating procedure. Allow the
crucible to cool and reweigh it. Record the mass.
‣ If the last two weighings do not agree to within 0.02
g, the chemical reaction between the copper and
sulfur is incomplete.
‣ If so, add more sulfur and repeat the heating and
weighing until constant mass is observed.
‣ Calculate the mass of copper sulfide obtained.
‣ The difference in mass between the copper sulfide
and copper is the mass of sulfur in copper sulfide.
Calculate this mass.
‣ Later, you will cover this mass of sulfur into the moles
of sulfur atoms in your product.
50
Experiment
‣ Sulfur is volatile and flammable. Without knowing the formula of the copper
sulfide formed, we can’t know how much sulfur to use—so we use an excess.
‣ We keep adding more sulfur and heating, consuming all the copper and burning up
any excess sulfur.
‣ In the end all the copper we added should be in the copper sulfide.
‣ All the sulfur should be either in the copper sulfide or have burned away.
51
Experiment
‣ C1: Determine the mass of copper you consumed in
this reaction (+/- .01 grams).
‣ C2: Determine the mass of copper sulfide you
produced (+/- .01 grams).
‣ C3: Determine the mass of sulfur in your copper
sulfide (+/- .01 grams).
‣ C4: Calculate the empirical formula for the copper
sulfide you formed (the ratio of copper to sulfur
atoms).
CuXSY
52
Experiment
‣ C5: Theoretically, two copper sulfides can exist.
One has a molecular weight of 159.157 amu (and
therefore molar mass of 159.157 g/mol), the second
has a molecular weight of 95.611amu (and therefore
molar mass of 95.611 g/mol).
‣ Considering the empirical formula you determined
above, only one of these is the possible product of
today's experiment.
‣ Which is it and why?
‣ C6: Assuming the two copper sulfides above are the
only possible copper sulfides, you now know the
molar mass of the of the compound you produced
today.
‣ Calculate the molecular formula of the copper sulfide
using that molar mass and the empirical formula you
discovered.
‣ The empirical formula is the ratio of copper and
sulfur atoms, the molecular formula is how many of
each atom is a molecule of that molecular weight.
53
Ex02
Particles
‣ Atoms
‣ Atoms, Molecules & Ions
‣ Compounds & Elements
‣ Experiment
‣ Copper Sulfide
‣ Next Week
‣ Atomic Structure
‣ Molecules
‣ Bonding
‣ Formulas
C6H2(NO2)3CH3
‣ Weighing Atoms
‣ AMU & Moles
‣ Molar Mass
‣ Finding Formulas
1 mol = 6.022 x1023 singles
54
Next Meeting
‣ Before next Meeting:
‣ Acquire and bring to class:
‣ Notebook
‣ You will not be turning in notebooks, but this
permanent record of your preparations,
observations and notes will be essential to
success in this class.
‣ Textbook, calculator, pencils
(yes, you can use pen)
‣ Safety Glasses (you cannot participate in the next class without them)
‣ Read and bring a copy of the next experiment
Elemental Analysis (finding the % phosphorus in plant food)
(http://chemskills.com/labs)
‣ Produce and bring to class:
‣ Your pre-lab for exp 03
‣ Your procedure summary for exp 03
‣ Review from your lecture text:
‣ Mole Ratio and Stoichiometry
55
We will start
with a quiz about
the experiment and
reading.
Questions?