I. Elements, Compounds and
Mixtures
What is Chemistry?

The science of matter
 How atoms interact with each other
 Physical and chemical properties of substances

What is Matter?
 Anything that has mass and takes up space (has
volume)
○ EVERYTHING is made of matter.
 If it doesn’t have mass
and volume, it’s not there
A. Mixtures
Mixture - Two or more substances, combined in
varying proportions - each retaining its own
specific properties.
 The components of a mixture can be separated
by physical means, i.e. without the making and
breaking of chemical bonds.


Examples: Air, milk, wood, salt water, ink, soda, and
concrete.
1. Heterogeneous Mixture

Mixture in which the properties and
composition are not uniform throughout the
sample.

Examples: milk, wood, and concrete.
2. Homogeneous Mixture
Mixture in which the properties and
composition are uniform throughout the
sample.
 Liquid mixtures are termed solutions.

Examples: air and table salt thoroughly
dissolved in water.
B. Pure Substance

A substance with constant composition.
 Crystals, substance made of just one type of
molecule or element

Can be classified as either an element or as a
compound.

Examples: Table salt (sodium chloride, NaCl),
sugar (sucrose C12H22O11)
water (H2O),
iron (Fe),
copper (Cu)
and
oxygen (O2)
1. Element

A substance that cannot be separated into two
or more substances by ordinary chemical (or
physical) means.
 We use the term ordinary chemical means to
exclude nuclear reactions.

Elements are composed of only one kind of
atom.

Elements are found on the Periodic Table.

Examples: Iron (Fe), copper (Cu), and oxygen (O2).
2. Compounds
A substance that contains two or
more elements, in definite proportion
by weight.
 Compounds are composed of more
than one kind of atom bonded
together.
 The term molecule is often used for
the smallest unit of a compound that
still retains all of the properties of the
compound.


Examples: Table salt (sodium chloride, NaCl),
sugar (sucrose, C12H22O11), and water (H2O).

Separate Elements
X and Y

Mixture of Elements
X and Y

Compounds of
Elements X and Y
Which is it?
Mixture
Element
Compound
Element, Compound
or Mixture ?
Element, Compound
or Mixture ?
Element
Element, Compound
or Mixture ?
Element, Compound
or Mixture ?
Element
Element, Compound
or Mixture ?
Element, Compound
or Mixture ?
Mixture
Element, Compound
or Mixture ?
Element, Compound
or Mixture ?
Mixture
Element, Compound
or Mixture ?
Element, Compound
or Mixture ?
Compound
Element, Compound
or Mixture ?
Element, Compound
or Mixture ?
Compound
Element, Compound
or Mixture ?
Element, Compound
or Mixture ?
A mixture of a
compound and
an element
QOD

VOCAB
Classify the
 Compound – two
following as an
more elements
element, compound
BONDED together
or mixture
in definite
(homogeneous or
proportions
hetergenous)
 Lemonade Homogenous
 Iron Element (Fe)
 Carbon Dioxide Compound
C. Physical & Chemical Properties of Matter
1. Physical Properties:

Physical properties can be
observed or measured without
changing the composition of
matter. Physical properties are
used to observe and describe
matter.

Physical properties include:
• appearance
• texture
• color
• odor
• melting point
• boiling point
• density
• solubility
• polarity and many others.
2. Chemical Properties:

Chemical properties of matter
describes its "potential" to
undergo some chemical change
or reaction by virtue of its
composition. What elements,
electrons, and bonding are
present to give the potential for
chemical change.

For example hydrogen has the
potential to ignite and explode given
the right conditions. This is a
chemical property.

Metals in general have they
chemical property of reacting with
an acid. Zinc reacts with
hydrochloric acid to produce
hydrogen gas. This is a chemical
property.
D. Physical & Chemical Change of Matter
1. Physical Change:
2. Chemical Change:

Changes in matter
that do not alter the
matter itself.

Changes that do alter
the identity of a
substance.

EX:
 Size
 Shape
 Phase:
Solidliquidgas

EX:
(freezing, melting, boiling)
 Iron rusting:
4Fe(s)+3O2(g) 2Fe2O3
 Wood burning
 Copper turning to brass.
3. Observing Chemical Changes

Watch for the following to establish
that a chemical rxn has taken place:
 Precipitate (solid formed from
solutions).
 Emitted gas.
 Color change.
 Energy change
(hotter/colder/emit light)
*all chemical rxs have a temp
change*
Last slide
QOD

Which of the
following does NOT
demonstrate a
physical change?
VOCAB

Which of the
following does NOT
demonstrates a
chemical change?
 Evaporation
 Melting
 Cooking
 Flammability
 Hardness
 Solubility
 Luster
 Rotting
A. History

Democritus, 400 B.C. proposed that the world
was made up of two things:
 1. Empty space
 2. Small particles he called atoms (which
is tiny indivisible particles)
○ His Views were not supported

Isaac Newton and Robert Boyle in the 17th
century published articles stating their belief in
atomic structure.

John Dalton in the
early 1800s
offered the first
logical quantitative
explanation of
atomic structure.
B. Dalton’s Atomic Theory





1. All matter is composed of tiny particles
called atoms. *These particles could not be
broken down into smaller substances.
*2. Atoms of an element were exactly alike
and atoms of different elements were unalike.
3. Atoms combine in simple ratios to form
compounds. (law of definite and multiple
proportions)
4. Atoms can not be destroyed (only
rearranged during a chemical reaction)
video
C. Research and Revisions of Dalton’s
Atomic Theory
1. Plum Pudding Model
using the Cathode
Ray Tube
 JJ Thomson proposed
the Plum Pudding
Model in which the
negative electrons
were held in place by
a random scattering of
positive charges.
C. Research and Revisions of Dalton’s
Atomic Theory

Thomson work with the cathode ray tube led to
the discovery that cathode rays consisted of
electrons. By exposing the ray to a magnetic field
and measuring the bending of the ray, he was able
to calculate the ratio of an electron’s charge to its
mass.

http://science.jrank.org/pages/627/Atomic-Models.html
Cathode Ray Tube:
High voltage electricity is passed into the cathode (negative end).
A ray is generated toward the anode (positive end). When a
magnet is placed near the ray, the negative end of magnet would
cause the ray to bend in the opposite direction and the positive
end would bend the ray towards the magnet. Video
C. Research and Revisions of Dalton’s
Atomic Theory
Using data from Thomson,
Robert Millikan obtain the
first accurate measurement
of the electron charge
 Millikan was able to
calculate mass of an
electron

C. Research and Revision of Dalton’s
Atomic Theory
2. Rutherford’s Gold
Experiment
 Rutherford predicted
the presence of
neutrons in the
nucleus.
Most of the alpha particles are passing
through the foil. A few are slightly deflected
while even fewer are greatly deflected.
Results of Gold Foil Experiment

1. The atom contains a tiny dense center
called the nucleus
 the volume is about 1/10 trillionth the volume of the
atom
2. The nucleus is essentially the entire mass
of the atom
 3. The nucleus is positively charged

 the amount of positive charge of the nucleus balances
the negative charge of the electrons

4. The electrons move around in the empty
space of the atom surrounding the nucleus
D. Atomic Theory Early Laws

Law of Conservation of Mass = matter cannot
be created nor destroyed, only chemically
altered. A + B  AB.

Law of Definite Proportions = specific
substances always contain elements in the
same ratio by mass.
EX: mass of sodium to the mass of chlorine
in salt is always the same.
D. Atomic Theory Early Laws

Law of Multiple Proportions = ratio of masses
of one element that combine with a constant
mass of another element can be expressed in
small whole numbers.

EX: Tin (II) oxide = SnO = 1:1 ratio
Water = H20 = 2:1 ratio
Last slide
QOD

Which of the
following scientists is
NOT responsible for
contributing to the
discovery of atomic
structure?
 Thomson
 Einstein
 Dalton
 Rutherford
VOCAB
 Law
of the
Conservation of
Matter– Matter is
neither created nor
destroyed, only
chemically altered
From studying the history of the atomic
theory… We have concluded that atoms are
composed of protons, neutrons & electrons.
But what is the internal structure of atoms?
Atomic Structure

Atoms are divisible:
 1. Electrons = negatively charged
particles (Millikan and Thomson).
 Smallest of the subatomic particles
(e-) (has almost no mass)
 Found on the outside of the central
mass (nucleus)
Atomic Structure
 2. Protons = positively charged
subatomic particles (Thomson and Rutherford)
 Slightly smaller in mass than the
neutron.
 Found in the dense central mass
called nucleus.
Atomic Structure
 3. Neutrons = neutral, no charge
(Rutherford).
 Largest of the subatomic particles.
 Found in the dense central mass
called the nucleus.
Atomic Mass of Subparticles
Electrons = 0.000549 amu
 Protons = 1.0073 amu
 Neutrons = 1.0087 amu

Where can we find information on the number
of protons, neutrons & electrons of atoms?
THE PERIODIC TABLE
6
Atomic Number (Z)
C
Element Symbol
Carbon
Element Name
12.0107
Atomic Mass
52
Symbols of Elements

Symbols are used to represent elements in
the periodic table
 They are 1 or 2 letter abbreviations
 Capitalize the first letter only
Examples:
C carbon
N nitrogen
F fluorine
O oxygen
Co
Ca
Br
Mg
cobalt
calcium
bromine
magnesium
53
Determining the # of protons:
The number of protons in the nucleus
is known as the atomic number (Z).
Determining the # of neutrons:
Isotopes
 An
isotope: atoms of an element that have
the same numbers of protons but differ in
the number of neutrons.
 Isotopes
of an element will have different
number of neutrons.
 Isotopes of an element will have different
mass numbers due to more mass of the
increased number of neutrons
Isotope Symbols

Examples of isotope Symbols
mass number
atomic number
sodium-23
Na-23
23 Na
11
25 Na
11
sodium-25
Na-25
Example: Isotopes of carbon
C-12 has 6 p+ and 6 N0
C-14 has 6 p+ and 8 N0
56
Atomic Structure - Isotopes

Isotopes of an element always have the
same number of protons (atomic number)
but different mass numbers due to more
mass of the increased number of neutrons

Example: Isotopes of carbon
 C-12 has 6 p+ and 6 N0
 C-14 has 6 p+ and 8 N0
 Video
To Calculate P+, E-, and NO

Atomic number and protons
always equal one another.

Number of protons equals
the number of electrons (for
now). (p = e)

If you subtract the number of
protons from the mass
number= number of neutrons
(n = mass – p)

Ex:
P+ = 17
E- = 17
No = 35-17=18
Atomic Mass and Mass Number
Atomic Mass
 Mass of protons and
neutrons.
 If in decimal form is
an average of the
mass of all isotopes.

Mass Number
 Mass of protons and
neutrons.
 Always expressed as a
whole number.
 Can obtain by rounding
atomic mass.

19
19
K
K
39.0983
39
Now try:
1.
2.
3.
4.
5.
I
Kr
Ca
Na
Fe
Atomic Mass
Mass Number
126.90
83.80
40.08
22.98
55.85
127
84
40
23
56
Determining the # of electrons:
Remember…
Protons are Positively charged (p+)
 Electrons are Negatively charged (e-)


Electroneutrality means that an atom has
equal number of positive and negative charges.
 So… protons = electrons

Atoms are electrically neutral… but they can
gain or lose electrons to become ions.
61
Calculating Electrons
Ions - an atom or group of atoms that has an electric
charge because it has lost or gained electrons
 Cation- an ion that has a
positive charge.
○ More protons than
electrons.
○ Gave away an electron
 Example: Li+
Atomic # 3
Protons = 3
Electrons = 3-1 = 2
 Anion- an ion that has a
negative charge.
○ More electrons than
protons.
○ Accepted an electron
 Example: As3-
Atomic # 33
Protons = 33
Electrons = 33+3 = 36
*remember… do the opposite of what the charge is!
62
REVIEW
How to calculate subatomic particles:
P = atomic number
N = mass # - p
E = protons charge
Total # of subatomic particles = p + n + e
Last slide
63
QOD
 Find
how many
neutrons are in iron
(Fe) if it has a mass
of 55.
 mass = p + n
 n = mass - p
 Atomic # = p = 26
 n = 55-26 = 29
VOCAB
– Atomic
Number, which is
the number of
protons in the
nucleus of an
atom.
Z
Atomic Mass
Atomic mass is the
weighted average mass of
all the atomic masses of the
isotopes of that atom.
19
K
39.0983
Calculating Average Atomic Mass
About 75.5 % of the chlorine found in
nature is Cl -35 (17 protons, 18 neutrons)
and about 24.5% of the chlorine found in
nature is Cl-37 (17 protons, 20 neutrons) .
(mass # x %) + (mass # x %) + … =
Cl-35
Cl-37
(35 x 75.5%) + (37 x 24.5%) =
26.4 g
+
9.07 g
= 35.5 g
QOD
 Which
atom
contains exactly
15 protons?
 Phosphorus-32
 Sulfur - 32
 Oxygen-15
 Nitrogen-15
VOCAB
 Alpha
particle
The type of particle
that passed through
an electrical field in
the Gold Foil
Experiment proving
that the nucleus of
an atom is positively
charged.
Atomic Models
Since the atom is
too small to be seen
even with the most
powerful
microscopes,
scientists rely upon
models to help us to
understand the
atom.
Believe it or not this is a
microscope. Even with the
world’s best microscopes we
cannot clearly see the
structure or behavior of the
atom.
Atomic Models
Scientists create
models to help them
to visualize complex
properties, structures
or behaviors. Since
the atom is so small,
scientists must gather
Indirect Evidence to
develop their models.
This is a model of a very
complex molecule made of
many different kinds of atoms.
Each colored ball represents an
atom of a different element.
Indirect Evidence
Indirect Evidence is evidence gathered
without being able to directly observe the
object. The Atomic - Molecular Theory of
Matter is based upon a vast amount of
indirect evidence gathered over a long
period of time. Just like pieces being added
to a puzzle, each new bit of information
gives us a better understanding of atoms.
Atomic Models:
 Old
version = Bohr’s
Also known as the
planetary atomic model
 Describes electron paths
as perfect orbits with
definite diameters
 Good for a visual

New version =
Quantum Theory
 Most accepted
 Diagrams electrons of a
atom based on
probability of location at
any one time

Bohr’s model:
Nucleus is in the center of an atom(like the
sun) and the electrons orbit the nucleus
similar to the planets.
 Orbits are called shells

 1st shell = 2 electrons
 2nd shell = 8 electrons
 3rd shell = 18 electrons
 4th shell = 32 electrons
Last slide
QOD
 What
is the
approximate
mass of an
electron?
0.000549 amu
VOCAB

WHICH IS A TRUE
STATEMENT?
 Compounds can be
broken down
(decomposed) by
chemical means
 Compounds can be
decomposed by
physical means
Study of how light interacts with matter
Quantum Theory:
 To
better the description of the atomic
structure, atoms were exposed to energy
(heat) which made the electrons go into what
is called the excited state (normal = ground
state).
 When electrons returned to ground state they
emitted energy in the form of light.
Quantum Theory:
 This
method of study is called
spectroscopy (spectrum)
 Visible light = part of the
electromagnetic spectrum
between 400-700 nm
Electromagnetic Spectrum
Quantum Theory:

Electromagnetic
Spectrum
 From crest to crest =
frequency which is
measured in hertz.
 This therefore can be
used to identify
elements (absorption of
energy and color emitted
is a fingerprint of an
element)
Kind of like wearing your team colors.



Continuous spectrum of white light
When you pass sunlight through a
prism, you get a continuous spectrum of
colors like a rainbow.
Line-Emission Spectrum
However, when light from Hydrogen &
Helium gases were passed through a
prism, they found a dark background
with discrete lines.
WHY? This lead to the quantum theory.
H
He
Quantum Theory:
A scientist, Bohr suggested
that electrons must exist in
Electron Orbitals (shows the
most probable area to find an
electron of a certain energy.)
So whenever an excited
hydrogen atom falls to its
ground state or lower energy
level, it emits a photon of
light, which means that
energy levels must be fixed.

Video
Quantum Theory: Electron Configuration
 Electrons
(e-) of atoms
are the basis for every
chemical reaction.
 In quantum theory,
electrons exist in
orbitals based on
probabilities and these
orbitals are arranged
within energy levels.
Notice… these orbitals
look different from Bohr’s.
This diagram is more
correct.
Quantum Theory: Electron Configuration
Quantum Numbers
 Quantum
numbers
specify the
properties of atomic
orbitals and the
properties of
electrons in those
orbitals
 We will define these
numbers & letters.
Example of Quantum #:
2
3s
Quantum Theory: Electron Configuration
Principle Quantum
Number (n)
 Is
equal to the number of
the energy level (n).
 The principle quantum #
corresponds to the energy
levels 1-7 which is the
period number (row) on
the periodic table.
Example of Quantum #:
2
3s
Blocks and Sublevels
P
E
R
I
1
3
O
4
D
6
S
1-7
8-85
d (n-1)
2
5
7
4
5
Quantum Theory: Electron Configuration
Maximum e- in Energy Levels

The maximum number of e- in
any one level is given by the equation 2n2
N=4, 32e
Calculate the maximum
number of electrons that can
occupy the 4th principal
quantum number (period 4).
 Solve: Use 2n2
2(4)2
32 electrons total

N=3, 18e
N=2, 8e
N=1, 2e
Quantum Theory: Electron Configuration
Sublevels and Orbitals
An energy level in made up of
many energy states called
sublevels.
 The number of sublevels for
each energy level is equal to the
value of the principal quantum
number.

EX: one sublevel in energy level one (period 1)
two sublevels in level two (period 2)
three sublevels in level three (period 3)
*now lets find out what those sublevels are
called…
Example of Quantum #:
2
3s
Quantum Theory: Electron Configuration
Sublevels and Orbitals

There are 4 sublevels:
s
p
d
f
Energy levels and sublevels work together to
form an e- cloud.
 e- are repelled by one another and move as
far apart as possible.
 e- clouds take on characteristic shapes
called orbitals.

Sublevels and Orbitals (notice the shapes)
Orbital Shapes
s orbitals are
spherical.
This diagram
represents an s
orbital.
d orbitals
contains 5
possible
orbital
shapes.
p orbitals are
“dumbbell”
shaped.
This diagram
represents 1 of the
3 types of p
orbitals.
f orbitals
contain 7
possible
orbital
shapes.
Electrons & Orbitals
Example of Quantum #:
 Orbitals
overlap and
change shape
as electrons
are added.
 Each orbital
can only hold
2 electrons.
2
3s
Electrons and Orbitals
(count 2 electrons per orbit)
Orbitals, and Electrons per Sublevel
Principal Quantum
Number (n)
Sublevel
# of
Orbitals
# of Electrons per
Orbital
1
s
1
2
2
s
p
2
6
3
s p
d
s p
d f
1
3
1 3
5
1 3
5 7
4
2
6
10
2 6
10 14
QOD
 The
principal
quantum number
corresponds to the:
•Energy Levels
•Periods on the
periodic table
VOCAB

Which statement is
true: The
characteristic brightline spectrum (color)
of an element is
produced when its
electrons…
 Move to an excited
higher energy state
 Return to a lower
ground energy
state
Distribution of Electrons
Atoms are electronically neutral. (for now)
 There is an electron for every proton in the
nucleus.
 The larger the atom, the larger the electron
cloud.
1. Pauli Exclusion Principle: only two e- can
occupy the same orbital due to
the opposite electronic spin .

Electron Filling Diagram
•Sublevels and
orbitals are filled as
indicated in the
diagram.
•Example:
1s2 2s2 2p6 3s2 3p6
4s2 …
energy level
Notice… they don’t go in order !
Sublevel
orbital
# electrons in
the orbital
Label your blank periodic
table.
Read it “like a book”
WRITE the Electron Configuration
Now try:
1. C
2. Kr
3. Ca
4. Fe
5. Hg
DRAW the Electron Configuration

Carbon has 6 e- (same as protons)

Start with lowest energy level and place
one electron in each orbital. Spins must
be in same direction within orbitals of the
same energy level.

If there are remaining e-, pair up singles
in same energy level before moving to
next highest energy level.
Electron Configuration
1s
2s
2p
Carbon’s electron config. is:
1s2 2s2 2p2
Superscripts total the number of
electrons
2+2+2=6
*Notice that you can write the electron configuration
based on the orbital diagram.
*When asked to draw or diagram, use arrow
configuration.
Last
*When asked to write, use 1s2,2s2… configuration. slide
QOD
 What
is the total
number of electrons
that can be held in
the third principal
energy level?
2n2
18
VOCAB

Quantum Theory of
atomic structure states
all except:
 Electrons orbit the
nucleus in perfect paths
 Electrons form clouds
based on probability of
location
 Electron clouds form
characteristic shapes
due to repelling of
negative charges
 Electrons occupy the
lowest energy levels
before moving into
higher energy levels
Electron Configuration – Noble Gas
Configuration
Electron Configuration demonstrates a
periodic trend, so you can write shorthand
electron configuration using the electron
configuration of the noble gases in Group 18
of the periodic table.
 Noble gases have stable configurations.

Noble Gas Configuration
When writing shorthand econfig for an element, refer
to the noble gas in the
energy level (period) just
above the element.
 Write the symbol of the
noble gas in brackets.
 Write out the remaining econfig based on the energy
filling diagram.

Electron
Configuration
Na = 1s22s22p63s1
Al = 1s22s22p63s23p1
Ne = 1s22s22p6
Shorthand Electron
Configuration
Na = [Ne] 3s1
Al = [Ne] 3s23p1
Noble Gas Configuration
EX: Na
Step 1: Na is in period 3 so refer to the
noble gas in period 2 which is Neon.
Step 2: Write Ne in brackets. [Ne]
Step 3: Now write remaining electrons in
standard form. 3s1.
Step 4: Combine. [Ne]3s1
Noble Gas Configuration
EX: Br
Step 1: Br is in period 4 so refer to noble gas
from period 3 which is Argon.
Step 2: Write in brackets. [Ar]
Step 3: Write remaining electrons.
4s23d104p5
Step 4: Combine to form: [Ar] 4s23d104p5
*Check your work: Add the number of electrons
from the noble gas (18) to the subscripts of
the remaining e-config (17). 18+17=35 which
is the electrons for Br.
Nobel Gas Configuration
Now try:
1. I
2. Kr
3. Na
4. Cu
Electron Configuration with Ions

When we write the electron configuration of
a positive ion, we remove one electron for
each positive charge:
Na
1s2 2s2 2p6 3s1

→
→
Na+
1s2 2s2 2p6
When we write the electron configuration of
a negative ion, we add one electron for
each negative charge:
O
1s2 2s2 2p4
→
→
O21s2 2s2 2p6
Electron Configuration with Ions
Now try:
1. Ca+2
2. Fe-3
QOD
VOCAB
 What
element
has completely
filled 3p orbitals?
 Which
of the
following is the
correct name for
Ca+1?
 Calcium isotope
 Calcium
Argon (Ar)
1s2
2s2
2p6
3s2
3p6
 Calcium ion
 Calcium with extra
electrons
Label your blank periodic table.
Read it “like a book”
*
**
*
**
S - Block
1
1
D - Block
2
6
P - Block
1
1
2
3
4
5
2
2
3
1
4
3
5
4
6
7
2
3
4
5
6
7
8
9
3
10
4
5
1
1
6
5
F - Block
6
1
2
3
4
5
6
7
8
9
10 11 12 13 14
4
4
5
5
Last
slide