ATOMS AND MOLES
Chapter 3
Unit Essential Questions:
1) How have atoms been studied
and understood throughout time?
2) How do we handle the very
small sizes of atoms in
calculations?
Lesson Essential Question:
1) How do laws in chemistry
support the existence of atoms?
Section 1: Substances Are Made of Atoms
First, a little history!
 As early as 400 BC, a few people believed in
atoms.

 Democritus
– Greek philosopher, first to develop
idea of the atom.
Had no evidence, only ideas.
 Theory of the atom has changed over time to
become what we have today.

Section 1 Cont.
Laws support atomic theory.
 Developed from observations of compounds
(how they are made up, how they react).

 Recall
what a compound is!
 Atoms of two or more elements chemically combined.

Three laws were developed.
Introduction #1



It turns out that atoms and compounds are very similar to
ingredients and food. How atoms come together to form
compounds is similar to how ingredients come together to
form food.
Think of a recipe for cooking or baking something and
link this process to the law of conservation of mass.
The amount of ingredients you put into a recipe should
equal the amount of food that comes out (but usually in
a different form). No food is lost or gained (ideally)!
Law of Conservation of Mass
Antoine Lavoisier 1782
 Mass cannot be created or destroyed in
chemical or physical changes, only rearranged.

S + O2  SO2
 32 g + 32 g  64 g
 Example:
Introduction #2





If you are baking a cake can you put together any
amount/number of ingredients and always get out a
cake?
No! So what can you say about amounts of ingredients
needed?
The amounts are definite- only certain amounts of
ingredients can make a cake.
Same is true for other recipes- given specific amounts of
ingredients, only certain outcomes (food) are possible.
Atoms and compounds are the same! Atoms come
together in definite amounts to form certain compounds.
Law of Definite Proportions
Atoms form compounds in specific, well defined
proportions.
 Proposed by Joseph Proust in 1797.

 Example:
 Always:
ethylene glycol (antifreeze)
51% oxygen, 39% carbon and 10% hydrogen
 C2H6O2
 Table
salt = sodium chloride
 61%
 NaCl
chlorine, 39% sodium
Introduction #3







If you are making a hamburger at McDonalds, how
many hamburgers and bun slices do you use?
One hamburger and two bun slices: ratio is 1:2.
If you are making a Big Mac at McDonalds, how many
hamburgers and bun slices do you use?
Two hamburgers and three bun slices: ratio is 2:3.
Based on this information, can hamburgers only be found
as having one hamburger and two bun slices?
No! Other combinations are possible!
This is also true when atoms of certain elements combine:
2H2 + O2  2H2O
H 2 + O 2  H 2O 2
Law of Multiple Proportions




Two elements can combine to form two or more different
compounds.
If the mass of the first element is held constant, the ratio
of masses of the second element is always a small whole
number ratio.
Name of
Mass O
Mass N
Example: Compound Description Formula
nitrogen
monoxide
colorless gas
NO
16.00 g
14.01 g
nitrogen
dioxide
poisonous
brown gas
NO2
32.00 g
14.01 g
What is the ratio by mass of O atoms? 1:2
Dalton’s Atomic Theory

1808- John Dalton combined ideas of others to
form an atomic theory.
 Took
Greek idea of atoms and turned it into a
theory that could be tested.
 Used the three laws previously discussed.

Was not 100% correct – we’ll look at new
evidence that disproves some of his theory in
Section 2.
Dalton’s 5 Principles of Atomic Theory
1. All matter is composed of extremely small particles
called atoms, which cannot be, created, destroyed, or
subdivided.
2. Atoms of a given element are identical in their
physical and chemical properties.
3. Atoms of different elements differ in their physical
and chemical properties.
4. Atoms of different elements combine in simple, whole
number ratios to form compounds.
5. In chemical reactions, atoms are combined,
separated, or rearranged but never created or
destroyed.
Dalton’s Atomic Theory- Which are Still
True Today?
1. All matter is composed atoms, which cannot be
subdivided, created, or destroyed. FALSE
2. Atoms of a given element are identical in their
physical and chemical properties. FALSE
3. Atoms of different elements differ in their physical
and chemical properties.
4. Atoms of different elements combine in simple,
whole number ratios to form compounds.
5. In chemical reactions, atoms are combined,
separated, or rearranged but never created or
destroyed.
Lesson Essential Question:
What are the components of the
atom and why are they
important?
Research Topics





Group 1: Thomson’s discovery and how he
discovered it
Group 2: Thomson’s model of the atom- provide
explanation & picture
Group 3: Rutherford’s discovery & how he
discovered it
Group 4: Rutherford’s model of the atom- provide
explanation & picture
Group 5: Chadwick’s discovery
Section 2: Structure of Atoms



Subatomic particles were discovered after Dalton’s
theory.
The three we will discuss:
 Electron
 Proton
 Neutron
Others exist (quarks- make up neutrons and protons,
leptons- make up electrons), but they are normally
discussed in physics.
Electron

Discovered by JJ Thomson in the mid 1800’s.
 Studying
electricity, not atoms, using a cathode ray tube.
 Pumped
all air out of glass tube.
 Applied voltage to two metal plates, called electrodes.
 Anode
– connected to positive terminal
 Cathode – connected to negative terminal
 Glowing
beam came out of cathode toward anode.
 No
atoms were inside the tube, so the beam must have come
from the atoms in the cathode.
 Beam was made of electrons.
 Ray
came from cathode- negatively charged.
 Used
in older TVs, computer monitors, and radar displays.
Further testing…

Used an electric field in addition to the magnetic
field to deflect the ray.
 Further
proof the beam was negatively charged.
What About Mass?
How could Thomson test the beam to see if it
had mass? What can objects that have mass do
that objects without mass can’t do?
 Move things!
 So, Thomson placed a small paddle wheel in
the beam’s path.

 Wheel
turned when hit by the beam. What does
this tell you?
 The beam consists of particles that have mass!
Further testing…
Now that Thomson reached a conclusion, what
should he do next?
 Verify the results… many times!

 Retested
using the same metal AND other metals.
Both gave the same results.
 Why was it important to test other metals and not
just the same metal?
 To show that electrons must exist in all atoms, not
just the atoms in the first metal he tested.

Electrons- subatomic particles that have a
negative charge.
Plum Pudding

How would you describe the locations of raisins in
the plum pudding? Assume it’s the same inside.
Plum Pudding Model

Thomson proposed the “plum-pudding” model
of the atom.
 Electrons

are embedded in a positive ball.
How did he know there should be a positive
charge?
Further Searching




Imagine you were throwing a ball at a ‘special’
wall, and 98% of the balls you threw at the wall
went through it while the remaining 2% bounced off
of it in various directions.
What can you conclude about the composition of
this ‘special’ wall?
Imagine the ball you were throwing had a positive
charge on it. What would this tell you about any
charges present in the wall?
Very similar to the gold foil experiment!
Searching for a Positive
Subatomic Particle


Atoms are neutral, so positive particles must exist in
atoms to balance out negative electrons.
Gold foil experiment: conducted by Ernest
Rutherford, student of Thomson’s (1909).
 Alpha
particles (positive charge) directed at gold foil.
 Most particles went through the foil.
 Atoms must be mostly empty space.
 Some particles were deflected.
 There must be some concentrated positive area in
atoms.
Nucleus & Proton

Rutherford developed the idea of the nucleus.
 Nucleus
– atom’s positive central region, location of
protons (and neutrons).
 Protons- positive subatomic particle in the nucleus.
 Mass is 2,000 times greater than an electron.
 The nucleus is only 1/10,000 of the radius of the
whole atom.
 If the nucleus was the size of a marble, the entire
atom would be the size of a football stadium!
 Measure atom’s radius in picometers (pm) =
10-12m.
Rutherford’s Model of the Atom

Rutherford’s experiments did not support
Thomson’s Plum Pudding model.
Developed the Planetary model- electrons look
like planets orbiting the sun.
 Let’s visit an up-close picture of the gold foil.

Neutron

Discovered 30 years after the proton was found
by James Chadwick in 1932.
 Several
people made observations about the neutron
before Chadwick.
 In studying a powerful beam, Chadwick was not able to
deflect it with magnetic or electric fields.
 Concluded the particles in the beam must be neutral in
charge.
 Neutrons – subatomic particles found in the nucleus and
have no electric charge.
Subatomic Particle Summary
Name
Symbol
electron
e-
proton
P+
neutron
n
Common
charge
notation
Mass (kg)
-1.602 x
10-19
-1
9.109 x 10-31
+1.602 x
10-19
+1
1.673 x 10-27
0
1.675 x 10-27
Actual Charge
(C)
0
Stability of Nuclei

How do protons stay together in the nucleus?
 All
protons are positively charged- why don’t they push
each other out of the nucleus?

Even though protons do repel one another in the
nucleus, neutrons help hold them together.
 Neutrons
provide attractive forces without being subject to
repulsive charge-based forces.
Atomic Number

Number of protons is unique to each element.
 Can
be used to identify elements.
 Example:
atomic number = 1 = 1 proton = hydrogen
 How many protons does carbon have? 6
 What element has 80 protons? Hg (mercury)
The number of protons is the atomic number.
 When atoms are neutral (no net charge),
atomic number (number of protons) must equal
number of electrons.

 For
a neutral atom: p+ = e- = atomic #
Mass Number

Total number of subatomic particles in the nucleus.
# = # of n + # of p+
 Mass #’s are not unique (isotopes – will explain later).
 Mass

Can be used to find number of neutrons.
n

= mass # - atomic #
Same as: mass # - #p+
Examples:
 Hydrogen
can have a mass # = 1, 2, or 3 but atomic
number is always 1.
 So, number of neutrons H atoms can have = 0,1,2
Nuclear Symbols
Representation using symbols with numbers to
identify the atomic number and/or mass number.
 One method is name-mass #.

 Ex:

Carbon-12 or Carbon-14
Another method is called ‘nuclear symbol’.
A
 Example: ZX
A
= mass number ; Z = atomic number; X = element
symbol
12
 6C
 Remember that the bottom number is always the
same for any element, the top number can vary.
 Because changing #p+ changes the element!
Nuclear Symbols & Ions

Not all atoms are neutral! Some have a charge.
 Atoms
with a charge are called an ion.
 Ex: O-2 and Na+


The superscript gives the charge of the ion.
Only # of e- are changed to produce ions. # of
p+ stays the same!
charge = e- lost
Na+: 11e- - 1e- = 10e Subtract the charge number from the # of e Negative charge = e- gained O-2 : 8e- + 2e- = 10e Add the charge number to the # of e Positive
Isotopes

Isotope – an atom with the same number of
protons, but a different number of neutrons
(and therefore a different mass #).
 Identified
using the two methods outlined on the
previous slides.
 Examples:
4
and 32He
 copper-63 and copper-65
 2He
Lesson Essential Question:
How can we describe the location
of an electron in an atom?
Bohr Model

Developed after Rutherford’s planetary
model.
 Recall
that Rutherford’s model showed that
electrons orbit the positively charged nucleus:

Problem: why don’t electrons crash into the
nucleus?
 Rutherford’s
model could not answer this question.
Bohr Model

Bohr proposed that electrons can only orbit the
nucleus at certain energy levels- they cannot exist
anywhere in between.
 Similar
to the rungs on a ladder- you can’t step in
between them, there’s nowhere to put your foot down!

Thus, electrons don’t crash into the positive nucleus
because there is no energy level (orbit) for them to
exist in.
Electrons and Light

Einstein (1905) proposed light had properties of
particles in addition to wave properties.
 Photoelectric effect: a certain amount of
energy is needed to remove an electron from
a piece of metal when struck with light.
 If light only acted as waves, any frequency
would eventually have enough energy to
remove an electron. But this was not seen!
 Only certain frequencies with certain energies
could remove an electron.
Electrons as Particles and Waves
DeBroglie (1924) pointed out that electrons act as
waves as well as particles.
 Quantum model of the atom used orbitals – regions
where electrons are likely to be found.

 Also
called
electron clouds.
 No sharp boundaries.
 Uses probability.
Match the model to the scientist/theory

Thomson (Plum Pudding)

Bohr

Rutherford (Planetary)

Quantum
Light Emission

Electrons have a certain energy level where they
are located: ground state = low energy.
 Higher
state = excited state = higher energy.
 When they are removed or moved their energy
changes.
 The
difference in energy is usually released as light.
 Each element can give a unique line-emission spectrum.
 Bohr
developed an equation to calculate the energy of
each electron.
 This
led to many people accepting his model of the atom.
Electron Excitation
Electromagnetic Spectrum
Warm-Up Question
What is the purpose of an address? Why
is each component necessary?
 Examine the following address:
430 New Schaefferstown Road
Bernville, PA
 Order the components of the address from
most general to most specific.
 PA, Bernville, New Schaefferstown Rd., 430

Quantum Numbers
Like an address, electrons can be identified by
where they “reside” in an atom.
 The 4 parts to this ‘address’ are called quantum
numbers.
 Each quantum number further pinpoints an
electron.
 In other words, quantum numbers separate
electrons from one another- they let you tell
them apart.

Principal Quantum Number

Principal quantum #, n, tells the energy level.
n
can only be positive integers (1, 2, 3, …)
 The larger the n value, the farther the e- is from the
nucleus, and the greater its energy.
 In terms of an address, this would be like the state- it gives
you the most general idea of where an electron is.
n = 1 (first energy level)
n = 2 (second energy level)
Angular Momentum Quantum Number

Angular momentum quantum #, l, tells the
sublevel/orbital.
 Can
be zero or any positive whole number.
 Each sublevel has a different shape.
 Letters
l =0
l =1
l =2
l =3
designate shapes of different l values:
s
p
p
d
s
d
f
Angular Momentum Quantum
Number Continued



In terms of an address, this would be like the city.
Cities give us a more specific area to look for someone
within a state.
Sublevels further specify where an electron is within an
energy level.
Magnetic Quantum Number

Magnetic quantum #, m, tells the orbital’s
orientation in space.
 Each
orbital shape can have different numbers of
orientations.
 More orientations = more orbitals present in a sublevel.
 s has 1 orbital because it only has 1 possible orientation.
 p has 3 orbitals because it has 3 possible orientations.
 d has 5 orbitals, and f has 7.
 Notice
the pattern- the number of orbitals (or
orientations) increases by 2 for each sublevel.
A Further Look at Orientations
Magnetic Quantum Number
In terms of an address, this would be like the
street.
 The street that a person lives on allows us to
further isolate where that person is located within
a city.
 Orientations of orbitals further specify where an
electron is within a sublevel.

Spin Quantum Number

Spin quantum number, s, tells the electron’s
orientation within an orbital.
 Can
only have 2 values, which are symbolized in 3 ways:
 +1/2
and -1/2

and
 clockwise and counterclockwise




Note: any orbital can only hold up to TWO electrons!
In an address, this would be like the street number.
Once we have the street number of a person, we know
exactly where to find them.
The same goes for an electron when we know the spin.
Another Look at Spin

The up and down arrows will come into play when
we learn to write electron configurations.
Thinking Ahead- Electron
Configurations
Can 2 e- in the same orbital have the same spin
( or )? (This would be like two homes on the
same street having the same house number.) Why or
why not?
 Do you think e- would prefer to be closer to the
nucleus at a lower energy level or farther away
from the nucleus at a higher energy level?

 Hint:
think about charges and energy involved.
Thinking Ahead- Electron
Configurations

If given the choice, do you think 2 e- would rather
be paired together in the same orbital or be alone
in different orbitals?
 Hint:
think about charges.
Electron Configuration Rules

When determining the placement of electrons
in an atom, three rules must be followed.
 Pauli
Exclusion Principle – no two electrons in the
same atom can have the same four quantum
numbers.
 Just like no two places can have the same
address!
 aufbau principle – electrons fill up the lowest
energy level first.
Electron Configuration Rules Cont.

Hund’s rule – One electron must occupy each
orbital before pairing.
 Electrons
are negatively charged and repel each
other, so they spread out as much as possible.
 Think of this as the “movie theater” rule.
Electron Configurations
Row numbers = energy levels (n)
Blocks = shapes
1
2
3
4
5
6
7
s
(n)
d
(n-1)
p
(n)
f (n-2)
Remember: atomic # = #e- if there’s no charge!
Types of Electron Configurations

Full
 All

electrons are written out.
Example: Write the electron configuration for an
atom of N.
 First,
determine the atomic number.
 This tells you how many electrons N has.
7
 Then write the electron configuration using the
periodic table.
 1s22s22p3
 Check yourself! The superscripts should add up to the
# e- (atomic #): 2+2+3 = 7
Types of Electron Configurations

Abbreviated Cont.
 Find
the noble gas you’ll need to use:
(1) Go up one row from the element you’re writing the
abbreviated configuration for.
(2) Go all the way to the right-most column on the
periodic table. This element is the noble gas you’ll use.
Write the symbol for this element in brackets [ ].
 Determine how many electrons the noble gas has, and
count up to this many using the orbital filling diagram.
Where you end is implied by the noble gas in brackets.
 Pick up with your configuration as you normally would.
Types of Electron Configurations

Orbital Diagram
 May
use full or abbreviated AND involves the use
of horizontal lines (___) for each orbital.
 Arrows are used with each line to show electrons.
 Recall that up ( ) and down ( ) arrows are used
to show different spins! (spin quantum number
‘s’).
Lesson Essential Question:
How do we count and work
with large quantities of atoms?
Section 4: Counting Atoms
 Atomic
mass units: created just to measure
masses of atoms because they’re so small.
 Abbreviated
amu.
 Daltons (Da) can also be used.
 Use the values on the periodic table.
The Mole & Avogadro’s Number

Mole – SI base unit – amount of substance.
 Like

a counting unit.
 1 dozen = 12 eggs
Number of particles in one mole is called Avogadro’s
number.

1mole = 6.022 x 1023 particles
 Extremely large: 602,200,000,000,000,000,000,000!
 Atoms, molecules, etc. can be used for labels instead of
particles.
 Named after work that Amadeo Avogadro completed.
Example





If you have a dozen people and a dozen cars, what
do they have in common?
They both have 12!
Since the numbers are the same does that mean that
their masses are the same?
No! The heavier items will have a greater mass!
This is the same for atoms! If you have a mole of
carbon and a mole of gold they both have
6.022 x 1023 atoms. But a gold atom is heavier
than a carbon atom, so the mole of gold has a
larger mass!
The Mole & Molar Mass

To convert between moles and grams, we use
molar mass (conversion factor).
 Molar mass – mass in grams of one mole of an
element.
 Units = grams/mole or g/mol
 Example: Cu = 63.55 amu = 63.55
g/mol
 We will round all atomic masses or molar
masses to 2 places after the decimal point.