Matter & The Atom
Matter
Anything that takes up space and has mass
Can be classified as solid, liquid, gas or plasma
Is it matter?
What is not matter?
ENERGY, HEAT, LIGHT,
ELECTROMAGNETIC WAVES,
MAGNETIC FIELDS, IDEAS, ETC.
Properties of Matter
Describe the characteristics and behavior
of matter, including the changes that
matter undergoes
Observing Matter
 Macroscopic Observations: Observations made with the
5 senses
 Microscopic Observations: Observations made with a
microscope
 Submicroscopic Observations: Observations of
substances so small they cannot even be seen with a
microscope
Macroscopic
Microscopic
Submicroscopic
 Qualitative
Observation:
Describes the
properties of a
substance
 Quantitative
Observation:
An observation
that involves a
numerical
value.
Physical Properties
What are the physical properties
represented in the image above?
Chemical Properties
MIXTURES
Two or more elements physically combined.
How can you tell something is a mixture?
It can be physically separated into its parts.
Heterogeneous Mixtures
• The prefix “hetero” means “different”
• A mixture with different compositions throughout
• You can see each phase (part) of the mixture
Homogeneous Mixtures
• The prefix “homo-” means “the same”
• A mixture that is the same throughout
• You cannot see the phases (parts) of
the mixture.
HETEROGENEOUS
OR
HOMOGENEOUS?
Solutions
• Solute: The substance being dissolved in a
solution
• Solvent: The substance that dissolves the solute
• Aqueous Solution: A solution in which water is
the solvent
What is the solute?
What is the solvent?
82% Fe
18% Cr
ALLOY
NAME OF ALLOY
% MAKE UP
Stainless Steel
73-79% Fe
14-18% Cr
7-9% Ni
Sterling Silver
92.5% Ag
7.5% Cu
18-karat white gold
75% Au
12.5% Ag
12.5% Cu
14 karat gold
58% Au
14-28% Ag
14-28% Cu
EXAMPLE
Methods to Separate Mixtures
• Filtration: Separates a solid from a liquid
Separating…
• Magnet: Separates Fe, Co, or Ni
Separating…
• Distillation: Separates two or more liquids
with different boiling points.
Separating…
• Crystallization: Separates crystalline solids
from a saturated liquid
Separating…
• Chromatography: Separates different types
of liquids
PURE SUBSTANCE

Matter with the same fixed composition
and properties
– First Type of Pure Substance
 Element
– The Periodic Table: A chart that lists the chemical name
and chemical symbol for each element
– Chemical Symbol: A shorthand abbreviation for the
name of an element
– You can tell a substance is an element because it is on
the periodic table
-Can you separate an element? No
Aluminum = ___
Gold = ____ Tin = ____
PURE SUBSTANCE

Matter with the same fixed composition and properties
– Second Type of Pure Substance
 Compound
– Chemical Formula: A combination of chemical
symbols that show what elements make up a
compound and the number of atoms of each
element
 Subscript: A number written to the lower
right of an element symbol to indicate the
number of atoms of that
– How do you know if a substance is a
compound? If it is 1 thing only—and it is not
on the periodic table.
– Can you separate a compound? Yes—by
chemically decomposing it.
NaH2CO3
Mg(OH)2
Decomposing a Compound
Electrolysis
– “To tear apart with electricity”
– The process in which electrical energy causes
a non-spontaneous chemical reaction to occur
 May break a compound apart into its elements
 Electrolysis of PbBr2 & ZnCl2
 Electrolysis of Water
THE GREEK PHILOSOPHERS
• 250 B.C.
• Four Fundamental Elements: Earth,
Wind, Water, and Fire made up
everything in the world
DEMOCRITUS
• 450 B.C.
• Seashell experiment led to development
of the idea of an indivisible piece of matter
called “atomos”
•Atom: The building block
of matter
•Problem: No experimental
data to back his concept
•Aristotle: Discredited
Democritus
DEMOCRITUS’ ATOM
Law of Conservation of Mass
•Antoine Lavoisier--1782 (Mercury & Oxygen
Experiment)
•Mass cannot be created nor destroyed in a chemical
reaction
•The mass of the reactants must equal the mass of the
products in a chemical reaction
Law of Definite Proportions
• Joseph Louis Proust-1799
• In a pure compound, the elements combine
in definite proportions to one another
according to mass
– Water is always 2 Hydrogen : 1 Oxygen
Law of Definite Proportions
John Dalton-1803
·Elements are made of tiny particles called
atoms
·All atoms of a given element are identical
·Atoms of a given element are different then
every other element
·Atoms of one element can combine with
atoms of another element to form compounds
·Atoms are indivisible and indestructible.
Atoms can only be rearranged in chemical
reactions--not created, divided, nor destroyed.
DALTON’S ATOM
Benjamin Franklin-1700’s
• Kite experiment:
– Objects have 1 of
2 electric charges
• Called them + &
• Like charges
repel
• Opposite
charges attract
Michael Faraday--1839
• Suggested that atoms contain particles that
have electrical charge
– Electricity (elektron, Greek word for amber)
• The flow of electrons in a substance
– Static: Stationary
– Static Electricity: Electrical charges not in motion (socks
out of a dryer)
J.J. THOMSON--1897
CATHODE RAY TUBE: Evacuated glass tube in which a stream of electrons
emitted by a cathode strikes a fluorescent material, causing it to glow
CATHODE: The electrode that brings electrons to the ions or atoms in a
solution.
TELEVISIONS ARE CATHODE
RAY TUBES
THOMSON’S PLUM-PUDDING MODEL
OF THE ATOM
Thomson measured the degree to which a magnetic field and
an electric field deflected the cathode ray. Since the field was
attracted to the positive charge, he knew it must contain a
negative charge. By doing this he discovered the electron.
Electron’s
(negatively charged
particles) are
embedded in a ball
of positive charge.
Henri Becquerel--1896
• Accidentally placed uranium on unexposed
photographic film
– Found an image had been produced on the film
– Discovered that uranium exhibits radioactivity
Radioactivity
• Radioactivity is the spontaneous emission of
radiation from an element
• Marie & Pierre Curie were awarded the Noble
Peace Prize, along with Becquerel, for the
discovery of radioactivity
• The Curies isolated two other radioactive
elements—radium and polonium
• Elements with atomic numbers greater than 83 are
radioactive
Robert Millikan—1909
Oil Drop Experiment
Using this experiment, Millikan determined the
charge of the electron
Ernest Rutherford
Alpha particle (α)—a particle with a +2 charge
Beta particles ()—high-speed electrons
Gamma radiation ()—not composed of particles
RUTHERFORD’S GOLD FOIL
EXPERIMENT--1911
If Thomson’s model was correct, positive α
particles would all go straight through the atom.
However, Rutherford’s Gold Foil Experiment
proved this to be untrue. Instead, every once in a
while the α particle was repelled. Since α particles
are positive, that meant there was a small, positive
part of the atom.
RUTHERFORD’S NUCLEAR MODEL OF THE ATOM
--
Center of the atom
-- subatomic particle
with no charge
located in the nucleus
-- subatomic particle
with a positive charge
located in the nucleus
Concept: The atom is made of mostly empty space,
containing electrons, surrounding a small, dense,
positively charged nucleus.
James Chadwick—1930’s
Scanning Tunneling Microscope
(STM)
STM IMAGES
Nickel
Iron on Copper
Platinum
WHO RECEIVED THE NOBEL
PRIZE FOR STM IN 1986?
• Gerd Binnig and Heinrich Rohrer of the
IBM Research Laboratory
– 1981-Invented the STM which formed images
of individual atoms
ATOM MANIPULATION
Iron on Copper
Carbon
Monoxide Man
(on Platinum)
Xenon on Nickel
Particle Accelerator--FermiLab
What’s smaller than a proton?
Other particles: quarks, gluons, mesons, muons &
other exotic particles—no immediate chemical
impact
ATOMIC SIZE
A typical atom is 0.000000001 meter across
or 1 billionth of a meter
A quark is 0.000000000000000001 meter
ATOMIC SIZE
If the atom was the same size as the distance between the
Earth and the Moon then:
-the nucleus would be the length of a golf course
-a proton would be about the size of a football field
-a quark would be about the size of a golf ball
SUBATOMIC PARTICLES
Symbol
Charge
(relative)
Location
Mass
(relative)
Electron
e-
-1
Outside
nucleus
0 amu
Proton
p+
+1
Inside
nucleus
1 amu
Neutron
n0
0
Inside
nucleus
1 amu
Amu = atomic mass units= 1/12 the mass of a carbon-12 atom
Atomic Number
= # of protons
In an ATOM also
= # of _________________
because: an ATOM is electrically neutral
so: the positive = the negative
A lithium atom with 3
protons & 3 electrons has
no overall charge.
A boron atom with 5 protons
& 5 electrons has no overall
charge.
3 protons + 3 electrons =
5 protons + 5 electrons =
(+3) + (-3) = 0
(+5) + (-5) = 0
protons
electrons
Atomic Number
• The atomic number is the number that is always used to
identify an element
– Each element has a unique # of protons and the
number of protons of a particular element can never
change
• The elements are arranged by atomic number on the
periodic table
Example
Element
Name
Element
Symbol
Atomic
Number
# of
protons
10
Ag
62
# of
electrons
ISOTOPES
Isotopes: Different atoms of the same element with the
same number of protons but a different number of
neutrons
• What’s this mean? There are different types of the
same element
• Isotopes of the same element are: chemically alike,
because they have the same number of protons &
electrons
90% of the universe, H2O
D2O (moderator in nuclear reactors to
slow down neutrons)
Radioactive, produced in nuclear
reactors, fission bombs
Carbon Isotopes
MASS NUMBER
Mass Number: The sum of the isotope’s number
of protons and neutrons
-The mass number is the mass of just 1 of the
element’s isotopes
-The mass number is always a whole number
Mass # must be calculated, it cannot be found on the
periodic table!
Mass Number = (# of p+) + (# of no)
Neutrons =
APE
T
R
L
O
O
E
M
T
C
I = O & T
C
N
R
S
O
N
N
U
S
M
B
E
R
MAN
A
S
S
N =
U
M
B
E
R
T
E
O
U
M
T
I + R
C
O
N
N
S
U
M
B
E
R
Examples
What is the mass number of an atom with 17
protons and 20 neutrons?
How many neutrons does a carbon atom with
a mass number of 14 have?
ISOTOPES
Isotopic Symbol:
Mass #
Atomic #
Symbol
Naming an Isotope:
Element Name-Mass #
EXAMPLES
Write the isotopic symbols for the element with:
21 protons, 24 neutrons
53 protons, 74 neutrons
How many protons, electrons and neutrons are in
an atom of 3215P?
EXAMPLE OF AN ISOTOPE
Isotope
Carbon-12
Carbon-13
Isotopic Mass #
Symbol
Atomic #
# of
Protons
# of
Neutrons
ISOTOPES
Atomic Mass Unit: Unit for atomic mass
• 1 amu = 1/12 (mass of carbon-12 atom)
= 1.66 x 10-24 g
Atomic Mass: a weighted average of all of
the isotopes of an element
• Different from mass number
• Is found on the periodic table
• the atomic mass value will be closest to the mass of
the isotope that is most abundant
ATOMIC MASS
To calculate atomic mass:
Atomic Mass = [(% abundance isotope 1) x (mass of
isotope 1)] + [(% abundance isotope 2) x (mass of isotope
2) ] + …
Example: Calculate the atomic mass of carbon if carbon-12 has a percent
abundance of 98.89% and carbon-13 has a percent abundance of 1.11%.