The Structure of Atoms and Mole
Theory
Element vs. Compound
• Element: is a pure chemical substance made of one type of atom. They are
classified as metals, nonmetals or semimetals. For example: Na is a metal
in elemental form. Cl forms Cl2, a diatomic gas in elemental form.
– No charge (neutral state)
• Compound: is made up of a combination of two or more elements (more
specifically atoms from these elements). For example: An ionic compound
is made of a metal and a nonmetal (NaCl).
– Charge: Na is 1+ and Cl is 1‒
chemical formula
• Chemical formula:
1) Molecular: C12H22O11 (sucrose aka table sugar)
2) Empirical: NaCl (table salt)
2 HgO
2 Hg
chemical equation
+ O2
Periodic Trends
• Metals: shiny, silvery, soft
(malleable), good conductors of
heat/electricity, react violently in
water, all solid (except Hg).
• Nonmetals: not silvery (some
colored), brittle, poor conductors,
some solid, liquid (Br) and gas.
• Semimetals (metalloids): Have
properties that cross between metals
and nonmetals. Partially conduct
electricity (rather poor), silvery in
appearance. All are solid. B (Boron),
Si (Silicon), Ge (Germanium), As
(Arsenic), Sb (Antimony), Te
(Tellurium), At (Astatine)
Group 1A: known as the alkali metals
Li (Lithium), Na (Sodium), K (Potassium), Rb (Rubidium), Cs
(Cesium), Fr (Francium)
When part of an ionic compound, these cations have a 1+ oxidation
state (An electron configuration that is full is preferred)
Group 2A: known as the alkaline earth metals
Be (Beryllium), Mg (Magnesium), Ca (Calcium), Sr (Strontium), Ba
(Barium), Ra (Radium)
When part of an ionic compound, these cations have a 2+ oxidation
state (An electron configuration that is full is preferred)
Group 7A: known as the halogens
F (Fluorine), Cl (Chlorine), Br (Bromine), I (Iodine), At (Astatine)
When part of an ionic compound, these anions have a 1‒ oxidation
state (An electron configuration that is full is preferred)
Group 8A: known as the noble gases
He (Helium), Ne (Neon), Ar (Argon), Kr (Krypton), Xe (Xenon), Rn
(Radon)
Are rather inert and only under special circumstances do they form
compounds.
Law of Conservation of Mass: Mass is neither created nor
destroyed in chemical reactions.
Aqueous solutions of mercury(II) nitrate and potassium iodide will react to form a
precipitate of mercury(II) iodide and aqueous potassium iodide.
3.25 g + 3.32 g = 6.57 g
Hg(NO3)2(aq) + 2 KI(aq)
HgI2(s) + 2 KNO3(aq)
4.55 g + 2.02 g = 6.57 g
Evolution of Atomic Theory
https://www.youtube.com/watch?v=UDIprICe9kg
Atomic Theory Fill in the Blank Worksheet
Dalton’s Atomic Theory
• Elements are made up of tiny particles called atoms.
• Each element is characterized by the mass of its atoms.
Atoms of the same element have the same mass, but
atoms of different elements have different masses.
• The chemical combination of elements to make different
chemical compounds occurs when atoms bond together in
small whole-number ratios.
• Chemical reactions only rearrange how atoms are
combined in chemical compounds; the atoms themselves
don’t change.
The Structure of Atoms: Electrons
Cathode-Ray Tubes: J. J. Thomson (1856–1940) proposed that cathode rays
must consist of tiny, negatively charged particles, which we now call electrons.
Thomson was only able to measure the charge to mass ratio of an electron
The Structure of Atoms: Electrons
Robert Millikan’s Oil Drop Experiment
Was able to measure the mass
of an electron and ultimately
the charge (from Thomson)
The Structure of Atoms: Protons and
Neutrons
Rutherford’s Scattering Experiment
The Structure of Atoms: Protons and
Neutrons
Atomic Nucleus: When Ernest Rutherford (1871–1937)
directed a beam of alpha (α) particles at a thin gold foil, he
found that almost all the particles passed through the foil
undeflected. A very small number, however (about 1 in
every 20,000), were deflected at an angle, and a few
actually bounced back toward the particle source.
Rutherford explained his results by proposing that a metal
atom must be almost entirely empty space and have its
mass concentrated in a tiny central core that he called the
nucleus.
The Structure of Atoms: Protons and
Neutrons
The mass of the atom is primarily in
the nucleus.
The charge of the proton is opposite in
sign but equal to that of the electron.
Song about Atomic Theory
https://www.youtube.com/watch?v=07yDiELe83Y
Can you create a song based on what you learned? Create a song
using the people of atomic theory and a one liner about their
contributions. One that rhymes and will help you remember the
pioneers of atomic theory. You can get in groups of 2-3 to help. Be
prepared to at least read what you have written. Feel free to use
the fill in the blank handout.
Atomic Number and Mass Number
Atomic weight is also known as the average
atomic mass. The Mass Number comes from
this value unless given. Round to whole
number.
Some Definitions
Atomic Number: number of protons in an atom
(different elements have a different number of
protons). This number defines the atom.
Mass Number: Protons + Neutrons
Isotopes: atoms of the same element that have
different numbers of neutrons.
Ion: atom that has gained or lost 1 or more
electrons.
Ex: Na+ has a charge of +1 because it has lost 1 electron.
Ex: O-2 has a charge of -2 because it has gained 2 electrons.
Determining # of protons, neutrons and
electrons.
Look up the element in the periodic table. Write its
average atomic mass and atomic number down.
Round its average atomic mass off to the nearest
whole number to make it the mass number.
Protons: atomic number
Neutrons: Mass number – atomic number
Electrons: atomic number – charge (no charge written
down, the charge is zero)
*if the element has a number after the symbol and a
hyphen, use this as the mass number. It is an
isotope.
Isotope Notation (If Given)
nitrogen-14 or N-14
mass number
14
7
N
7 protons
7 electrons
7 neutrons
N
7 protons
7 electrons
8 neutrons
atomic number
nitrogen-15 or N-15
mass number
15
7
atomic number
Isotope Notation (If Not Given)
If you were just given N and asked to determine the number of electrons, protons, and
neutrons.
Look at periodic table.
Top left is atomic number = number of protons = 7
Bottom is average atomic mass.
The mass number will be 14.
So, 14 = number of protons + number of neutrons
14 = 7 + number of neutrons
Number of neutrons = 7
If there is no charge, then the number of protons equals
the number of electrons. So, there are 7 electrons.
Isotope Notation (If Not Given)
What if you were given N-3 and asked to determine the number of electrons, protons,
and neutrons.
Look at periodic table.
Top left is atomic number = number of protons = 7
Bottom is average atomic mass.
The mass number will be 14.
So, 14 = number of protons + number of neutrons
14 = 7 + number of neutrons
Number of neutrons = 7
A charge of -3 is given (this is an ion).
Number of electrons = atomic number – charge
Number of electrons = 7 – (-3) = 10
Red phosphorus is an allotrope of phosphorus that can be
made by heating white phosphorus to temperatures above
240 oC. The most common isotope of phosphorus can be
represented by
How many protons,
electrons, and neutrons does
this isotope have?
a) 16, 16, 15
e) 16, 16, 1
b) 15, 15, 1
c) 15, 15, 16
d) 16, 16, 31
The most stable isotope of silver is given by
How many protons, neutrons, and electrons does it have?
p,
n, e
a) 107, 47, 47
b)
60, 47, 60
c)
47, 47, 60
d)
47, 60, 47
e)
47, 107, 60
Example Problem #1
1. Determine the number of protons, neutrons
and electrons in each.
A. Ba
B. P-3
C. Na+
D. U-235
Average Atomic Mass
The mass number on the periodic table is a reflection of
the masses of each type of isotope of an element and
the percent of that element present as each isotope.
Avg. At. Mass = (%1)(mass1) + (%2)(mass2) + (%3)(mass3)…..
Neutrons and protons are most of the mass. The number
of protons defines the atom. Isotopes are of the same
atom, just differ in number of neutrons. So, a change in
mass only really reflects changes in the number of
neutrons.
Average Atomic Mass
Why is the average atomic mass of the element carbon 12.011 u?
carbon-12:
98.89% natural abundance
12 u
carbon-13:
1.11% natural abundance
13.0034 u
You have to convert percentages to decimal. Divide percent by 100 or move decimal two
places to the left.
Avg. At. Mass = (0.9889)(12 u) + (0.0111)(13 u)
= 11.867 u + 0.144 u
= 12.011 u
Example #2
An element (X) has two isotopes. Isotope
X-6 (7.5% of 6.015u) and
X-7 (92.5% of 7.016u).
What is the average atomic mass of element X?
What is the element?
Example #3
Chlorine has two isotopes. It is 75.77% of Cl-35
which has a mass of 34.969u and 24.230% of
Cl-37 which has as mass of 36.966u. What is
the average atomic mass of chlorine?
Example #4
Silicon has three isotopes shown in the table
below:
Isotope Mass (u)
Abundance(%)
Si-28
27.98
92.21%
Si-29
28.92
4.70%
Si-30
29.97
3.09%
What is the average atomic mass of silicon?
What the heck is a mole?
Nope, not a little animal.
A mole is a word that means a number.
The word “dozen” means 12
The word “mole” means 6.022 x 1023
(this is called Avogadro’s number)
You can have a mole of atoms, a mole of molecules or a
mole of pennies (that would be awesome!!)
Mole concept
The mass number on the periodic table
represents the mass of one mole of each
element.
1 mole C = 12.01g = 6.022 x 1023 atoms
1 mole of Fe = 55.85g = 6.022 x 1023 atoms
etc.
Converting between grams-moles and
atoms
Write down the number, the unit, and the chemical that
is given.
Make a bracket with a line.
Place the unit you currently have on the bottom, the unit
you need to switch to on the top.
In the brackets: next the unit “mole” place a 1, next to
the unit “grams” place the number on the periodic
table, next to the unit “atoms” place Avogadro’s
number.
Put the number on the left in the calculator first.
Multiply by the top number, divide by the bottom
number.
Example #5
A. How many moles are there in a
4,000,000atom sample of copper?
B. How many grams are in a 0.037mole sample
of iron?
C. How many atoms of sodium are in a 3.08g
sample of sodium?
D. How many grams of aluminum will a
3 x 1020 atom same weigh?
E. How many moles are in a 4g sample of
potassium?
Molar Mass or Molecular Weight
There are about 118 elements (so far) but there are
millions of compounds. To perform grams mole
conversions with compounds we have to calculate
the mass of one mole.
Subscript of the first element (atomic mass of the first
element) + subscript of the second element (atomic
mass of the second element)…..
If there is a poly atomic ion with parenthesis and
another subscript you need to multiply the
subscripts.
Example #6
Determine the molar mass of C6H12O6.
6(12.01) + 12(1.01) + 6(16) = 180.18g/mol
Round the atomic masses to two places past the
decimal.
Example #7
Determine the molecular weight of Cr(NO3)3.
1 Cr, 3x1 = 3 N, 3 x 3 = 9 O
1(52) + 3(14.01) + 9(16) = 238.03g/mol
Example #8
(let’s try some together)
Find the molar mass (molecular weight) for:
a. NaOH
b. CaCO3
c. AlPO4
d. C12H22O11
e. (NH4)2SO4
f. K2C2O4
Gram-mole-molecule conversions
Write down the number the unit and the chemical
that is given.
Make a bracket with a line. Place the unit you
currently have on the bottom, the unit you need
to switch to on the top.
Inside the Bracket: next the unit “mole” place a 1,
next to the unit “grams” place the number from
finding molar mass, next to the unit “molecule”
place Avogadro’s number.
Put the number on the left in the calculator first.
Multiply by the top number, divide by the bottom
number.
Example #9
A. How many moles of NaCl are there in a 75g sample?
B. How many molecules of Al2O3 are present in a
0.0059mole sample?
C. How many molecules of Fe2(CO3)3 are there in a
0.045g sample?
D. What mass will 1,000,000 molecules of (NH4)2SO4
have?
E. What mass will a 0.59mole sample of K2C2O4
have?
Percent composition
Determine the molecular weight.
Place the total mass of one element over the
molecular weight x 100%.
Do this for each element. Percent composition is
based on mass.
Example #10
Determine the percent composition of
Fe2(CO3)3.
Determine the percent composition of C2H6.
Determine the percent composition of
C12H22O11.
Empirical & Molecular Formulas
Empirical: The formula with the simplest whole number
ratio of elements in the compound.
Molecular: The formula with the exact numbers of
each type of atom present in the molecule.
CH2O
empirical
C6H12O6
molecular
Finding the empirical formula
If percents are given, assume you have the same
number of grams.
Do a gram mole conversion for each element
present in the molecule.
Identify the element with the smallest number
of moles. Divide all the moles by that number.
The answer will give you the subscripts for each
element.
What if you don’t get a whole number
when you divide?
You decide what fraction the decimal
represents. Multiply all the numbers by the
denominator of the fraction.
0.5 = ½
0.2 = 1/5
0.33 = 1/3
0.4 = 2/5
0.67 = 2/3
0.6 = 3/5
0.25 = ¼
0.8 = 4/5
0.75 = 3/4 if the decimal is <0.2 or >0.8 just
round off to a whole number.
Example #11
A compound is analyzed and determined to be
27.37% sodium, 1.20% hydrogen, 14.30%
carbon and 57.14% oxygen. Determine the
empirical formula for this compound.
Example #12
A compound is analyzed and determined to
contain 4.864g of carbon, 0.816g of hydrogen
and 4.320g of oxygen. Determine the
empirical formula of this compound.
Example #13
A compound is analyzed and determined to
contain 0.873g of phosphorus and 1.127g of
oxygen. What is the empirical formula?
Example #14
A compound is analyzed and the percent
composition is determined. The compound is
89.92% carbon and 10.08% hydrogen. What is
the empirical formula of the compound?
Hydrates
Hydrates are compounds with waters attached.
You can determine the formula for a hydrate
in the same way you find the empirical
formula.
Convert the grams of the anhydrous compound
and the grams of water to moles.
Divide both by the smaller number of moles.
This will give you the number that goes in front
of water.
Example: MgSO4 • 7H2O
Example #15
A hydrate is found to have the following percent
composition:
48.8% magnesium sulfate (MgSO4) and 51.2%
water (H2O).
When 11.75g of a hydrate are heated to drive all
of the waters of hydration off, 9.25g of
anhydrous cobalt(II)chloride (CoCl2) remain.
Determining Molecular Formulas
Find the Empirical formula just like before.
Determine the mass of the empirical formula in
the same way you found the molar mass of a
compound.
Place the molecular weight given in the problem
over the empirical weight.
Multiply each subscript by the answer to the
division of molecular weight by empirical
weight.
Example #16
A compound is analyzed and determined to be
39.993% carbon, 6.727% hydrogen and
53.280% oxygen. The molecular weight of the
compound is 180.18g/mol. Determine the
empirical and molecular formulas for the
compound.
Example #17
A substance is analyzed and determined to be
40.68% carbon, 5.08% hydrogen and 54.24%
oxygen and has a molar mass of 118.1g/mol.
Determine the empirical and molecular
formulas for this substance.
Example #18
A compound contains 0.856g of carbon and
0.144g of hydrogen. Its molecular weight is
56.12g/mol. What are the empirical and
molecular formulas for this compound?