Download atoms

On Earth, matter
usually can be found
as a solid, liquid, or
gas. BEC & Plasma
are usually found in
space.
What’s Matter?
Volume & Mass
Classifications of Matter
Substances
always the same composition (makeup)
Elements (1 kind of atom) 92 Natural About 120
Total see Periodic Table
Examples: Gold (Au), Lead (Pb), Mercury (Hg)
* 8 Elements make up 98% of crust – Al, Fe, Ca,
Na, K, Mg, Si 27%,O 47%
Molecules (2 or more kinds of atoms chemically
bonded that act as a unit)
Examples: Hydrogen (H), Bromine (Br), Sugar
Compounds (2 or more elements chemically
joined in a specific combination)
Examples: Water (H2o),Carbon Dioxide CO2
Mixtures
Matter that varies in composition. 2 or more
substances that are blended but not bonded.
Heterogeneous Mixtures - not evenly mixed
Examples: Trail Mix, Granite, Smoke
Homogeneous Mixtures - evenly mixed/no
bond!
Examples: Brass, Natural Gas, Windex
Solutions – not bonded mixture
Composition
Changes in
composition
Properties of
parts
Solutions
Substances
evenly mixed
together
Compounds
Atoms bonded
together in the
same
combination
Solution still the
Changes
same but relative
composition
substances differ.
makes a new
compound with
new properties.
Substances keep Properties of the
their own
compound are
properties when
different than
mixed.
atoms that make
it up.
The Greeks named these particles atoms, a
term that means ―cannot be divided.‖
• Early philosophers
didn’t try to prove their
theories by doing
experiments as
scientists now do.
• Their theories were the
result of reasoning,
debating, and
discussion—not of
evidence of proof.
400
B.C. - Democritus thought matter
could not be divided indefinitely.
• This led to the idea of atoms in a void.
fire
Democritus
earth
air
water
• 350 B.C - Aristotle modified an earlier
theory that matter was made of four
“elements”: earth, fire, water, air.
• Aristotle was wrong. However, his
Aristotle theory persisted for 2000 years.
• During the eighteenth century, scientists,
especially the French, began debating the
existence of atoms once more. Newton
proposed held together by force.
• They found that certain substances couldn’t
be broken down into simpler substances.
• Scientists came to realize that all matter is
made up of elements.
• An element is matter made of 1 kind of atom.
Dalton’s Concept
1800 -Dalton an English schoolteacher
proposed a modern atomic model
based on experimentation not on
pure reason.
Dalton pictured an atom as a hard
sphere that was the same throughout.
•
•
•
•
All matter is made of atoms.
Atoms of an element are identical.
Each element has different atoms.
Atoms of different elements combine in
constant ratios to form compounds.
• Atoms are rearranged in reactions.
• His ideas account for the law of conservation of
mass (atoms are neither created nor destroyed)
and the law of constant composition (elements
combine in fixed ratios).
Scientific Evidence
• In 1870, the English scientist William
Crookes did experiments with a glass tube
that had almost all the air removed from it.
• The glass tube had two pieces of metal
called electrodes sealed inside.
• The electrodes were connected to a battery
by wires. Crookes’s tube is known as a
cathode-ray tube, or CRT.
• An electrode is a
piece of metal that
can conduct
electricity.
• One electrode, called
the anode, has a
positive charge. The
other, called the
cathode, has a
negative charge.
• Crookes hypothesized that the green glow in the
tube was caused by cathode rays.
Models of the Atom
Cathode Rays
• Many scientists were not convinced that
the cathode rays were streams of particles.
• In 1897, J.J. Thomson, an English physicist,
tried to clear up the confusion.
• He placed a magnet beside the tube from
Crookes’s experiments.
• Thomson concluded that the beam must be
made up of (-) negatively charged particles
of matter that came from the cathode.
Models of the Atom
Thomson’s Model
• These negatively charged particles are now
called electrons. Electrons were 1st proposed by
proposed in 1874 by G. Johnstone Stoney
Thomson also inferred that electrons are a part
of every kind of atom.
• If atoms contain one or more negatively
charged particles, then all matter, which is
made of atoms, should be negatively charged
as well.
Thomson’s Atomic Model
• The negatively charged electrons were spread
evenly among the positive charge.
• The negatively charged electrons and the
unknown positive charge would then neutralize
each other in the atom.
• The atom is neutral.
Thomson’s Atomic Model
• Later discovered not all atoms are neutral.
#Electrons within an element can vary.
• More positive charge than negative electrons
= overall positive charge.
• More negative electrons = overall negative charge.
• Rutherford wanted to see what would happen
when they fired fast-moving, positively charged
bits of matter, called alpha particles +, at a thin
film (400nm) of a metal such as gold surrounded
by a fluorescent screen.
Models of the Atom
Rutherford’s Results Fail!
• His prediction -speeding alpha
particles would pass right
through the foil and hit the
screen on the other side.
• Rutherford reasoned - the thin,
gold film did not contain enough
matter to stop the speeding alpha
particle or change its path.
Rutherford was shocked when his
students Hans Geiger & Ernest
Marsden rushed in to tell him that
some alpha particles were veering
off at large angles.
The Model Fails
Positively charged alpha particles moving with such
high speed that it would take a large positive
charge to cause them to bounce back.
The Proton
• The actual results did not
fit this model, so
Rutherford proposed a
new one.
• He hypothesized – most of the
mass of the atom and its positive
charge is in the nucleus.
Models of the Atom
The Proton
• In 1920 scientists identified
the positive (+) charges in
the nucleus as protons.
• Rutherford’s new model of
the atom fits the
experimental data.
• Most alpha particles could
move through the foil with
little or no interference.
Electricity is called “cathode rays” when passed
through an evacuated tube.
These rays have a small mass and are negative.
Thompson noted that these negative subatomic
particles were a fundamental part of all atoms.
Dalton’s “Billiard ball” model (1800-1900)
Atoms are solid and indivisible.
Thompson “Plum pudding” model (1900)
Negative electrons in a positive framework.
Rutherford - Atoms are mostly empty
space.
Negative electrons orbit a positive nucleus.
The Neutron
• According to Rutherford’s model, the only
other particle in the atom was the proton.
• He proposed that another particle must be in
the nucleus to account for the extra mass.
• The mass of most atoms is at least twice as
great as the mass of its protons.
• The particle, which was later call the
neutron, would have the same mass as a
proton and be electrically neutral.
Models of the Atom
The Neutron
• 20 Years Later in 1932
-The model of the atom
was revised again to
include the newly
discovered neutrons in
the nucleus by James
Chadwick.
• The nuclear atom has
a tiny nucleus
tightly packed with
positively charged protons
and neutral neutrons.
Bohr Models
• Then, electrons would travel in
orbits around the nucleus.
• A physicist named Niels Bohr even
calculated energy levels for the hydrogen
atom.
• However, scientists soon learned that electrons
are in constant, unpredictable motion and
can’t be described try & explain the orbits.
Electrons orbit the nucleus in “shells”
Electrons can be bumped up to a higher shell if
hit by an electron or a photon of light.
The Electron Cloud Model
• The electrons are more
likely to be close to the
nucleus. They are
attracted to the positive
charges of the proton.
• Electrons travel in a
region surrounding the
nucleus, called the
electron cloud.
Identifying Numbers
• The atomic number of an element is the
number of protons in the nucleus.
• Atoms of an element are identified by the
number of protons because this number
never changes without changing the identify
of the element.
Mass Number
• The total masses of the protons and neutrons in an
atom make up most of the mass of an atom. The
mass number isotope is neutrons plus protons.
• You can find the # of neutrons in an isotope by
subtracting the atomic # from the mass #.
• Electron 1/1,800 of a proton. (We don’t calculate!)
Number of Neutrons
• These 3 kinds of carbon atoms are called
isotopes. Isotopes are atoms of the same
element that have different numbers of neutrons,
but the same # of protons.
•
Radioactive Decay
• Many atomic nuclei are stable when they
have about the same number of protons
and neutrons. Some nuclei are unstable
because they have too many or too few
neutrons. This is especially true for
heavier elements such as uranium and
plutonium.
• The release of nuclear particles and energy
is called radioactive decay.
• In these nuclei, repulsion builds up. The
nucleus must release a particle to become
stable.
Radioactive Decay
• When the particles that are ejected from a
nucleus include protons, the atomic number of
the nucleus changes. When this happens, one
element changes into another (Transmutation).
• A smoke detector makes use of
radioactive decay. It contains
americium-241, undergoes
transmutation by ejecting
energy and an alpha particle.
Radioactive Decay
• The fast-moving alpha particles enable the air to
conduct an electric current. As long as the electric
current is flowing, the smoke detector is silent.
• The alarm is
triggered when
the flow of
electric current
is interrupted by
smoke entering
the detector.
Changed Identity Alpha Particles
• When americium expels an alpha particle,
it’s no longer americium, it becomes the
element that has 93 protons, neptunium.
Loss of Beta Particles
• During this different kind of transmutation, a
neutron becomes unstable and splits into an
electron and a proton. The proton, however,
remains in the nucleus.
• The electron, or beta particle, is released with a
large amount of energy, so the atomic # of the
element that results is greater by one..
• A beta particle is a high-energy electron that
comes from the nucleus, not from the electron
cloud.
•
Rate of Decay
• Radioactive decay is random.
• The rate of decay of a nucleus is measured
by its half-life.
• The half-life of a radioactive isotope is the
amount of time it takes for half of a sample
of the element to decay.
The radioactive decay of
unstable atoms is steady,
unaffected by conditions
(weather, pressure,
magnetic or electric
fields, and even chemical
reactions.
Radioactive
Decay
• Carbon-14 is used to determine the age of
dead animals, plants, and humans.
• In a living organism, the amount of carbon14 remains in constant balance with the
levels of the isotope in the atmosphere or
ocean.
• This balance occurs because living
organisms take in and release carbon.
• Geologists examine the decay of
uranium & potassium argon to age rocks.
Uses of Radioactive Isotopes
• Tracer elements are used to diagnose disease
and to study environmental conditions.
• The radioactive isotope is introduced into
a living system such as a person, animal,
or plant.
• It then is followed by a device that detects
radiation while it decays.
• The isotopes chosen for medical purposes have short
half-lives, which allows them to be used without the
risk of exposing living organisms to prolonged
radiation. The isotope iodine-131 has been used to
diagnose problems with the thyroid. Also used to detect
cancer, digestion problems, & circulation.
Environmental Uses
• Tracers such as phosphorus-32 are injected
into the root system of a plant.
• A detector then is used to see how the plant
uses phosphorus to grow and reproduce.
• Radioisotopes also can be placed in
pesticides or the water cycle and followed to
see the impact to the ecosystem.

1789 Antoine Lavoisier grouped elements into
categories of metals, nonmetals, gases, & Earth.

Before 1830 55 elements were known to early man &
used to make weapons and jewelry. Examples: Gold,
Silver, Iron, Copper, Tin

1869 Dimitri Mendeleev – Periodic Table – The elements
were ordered by atomic mass & arranged according to
similarities in properties. Not all elements were known so
he left blank spaces for later discovery.

Early 20th Century Henry Mosely - Rearranged the table
by atomic number rather than mass.
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Organized by increasing atomic number
Periods (rows) are labeled 1-7 (# of electrons)
18 Columns of elements contains a group (family) of
elements that have similar properties. (Properties
Repeat -Periodic Law) The number of an “A” group
element matches the number of valence electrons for
an element in that group. ( Valence electrons are the
electrons that are in the highest occupied energy level
of an atom. Valence electrons are the ones involved in
chemical bonding.)
Divided into Zones.
- Representative Elements 1&2 & 13-18
*Metals *Metalloids *Nonmetals
- Transition Elements 3-12
*Lanthanide Series & *Actinide Series

1 or 2 letter abbreviation - may be for the
Latin name, discoverer’s name, or place
Examples: O = Oxygen Sc = Scandium

3 letter abbreviation for New Names – Must
be approved by the IUPAC International
Union of Pure and Applied Chemistry
Keys on the Periodic Table usually show the
information represented in the periodic table.
Atomic Number
Element name
Symbol
Often Atomic Mass number
(amu – atomic mass unit 1/12 of carbon -12)
Sometimes state of matter at room temp.
(Solid, Liquid, or Gas)
Sometimes Natural (1-92) or Synthetic (93+)
Using the Periodic Table
Step 1 - Gather Information
Go to the Periodic Table of Elements. Use the
Table of Elements to find your element's
atomic number and atomic weight.
The atomic number usually the number
located in the upper left corner and the atomic
weight is the number located on the bottom,
as in this example for krypton:
Step 2 - The Number of Protons is...
The atomic number is the number of protons in an
atom of an element. In our example, krypton's is 36.
This tells us that an atom of krypton has 36 protons in
its nucleus.
Every atom of krypton contains 36 protons. If an
atom doesn't have 36 protons, it can't be an atom of
krypton. Adding or removing protons from the
nucleus of an atom creates a different element. For
example, removing one proton from an atom of
krypton creates an atom of bromine.
Step 3 - The Number of Electrons is...
Atoms have no overall electrical charge (balanced).
Atoms must have equal numbers of protons and electrons.
In our example, an atom of krypton must contain 36
electrons since it contains 36 protons.
An atom can gain or lose electrons, becoming what is
known as an ion. An ion is nothing more than an
electrically charged atom. Adding or removing electrons
from an atom doesn’t change it, just its net charge.
Step 4 - The Number of Neutrons is...
The atomic mass the total number of particles in an atom's
nucleus. The atomic mass is a weighted average of all of the
naturally occurring isotopes of an element relative to the
mass of carbon-12.
All you really need to find is something called the mass
number. To find the mass number, round the atomic mass
to the nearest whole number. Krypton's mass number is 84
since its atomic mass, 83.80, rounds up to 84.
The mass number is a count of the number of particles in
an atom's nucleus. Remember that the nucleus is made up
of protons and neutrons.
http://videos.howstuffworks.com/hsw/14741-simplyscience-the-history-of-the-periodic-table-video.htm
http://videos.howstuffworks.com/hsw/5966chemistry-of-life-atoms-and-elements-video.htm
In Summary...For any element:
Number of Protons = Atomic Number
Number of Electrons = Number of Protons =
Atomic Number
Mass Number = (Number of Protons) +
(Number of Neutrons)
Number of Neutrons = Mass Number Atomic Number
The # of electrons increases from L to R
across a period on the table.
 The last element in each period has 8 in its
outer energy level so it is complete.
 Halogens have 7 in their outer shells.
 Noble gases have a stable electron
configuration. That’s why they don’t
combine! Each has 8 valence electrons,
with the exception of helium, which has 2.

The nucleus containing protons & neutrons
is surrounded by the electron cloud.
 The # & arrangement of electrons
determines the chemical & physical
properties.
 More electrons can fit into orbitals farther
away from the nucleus. (see overhead)
 When the highest occupied energy level of
an atom is filled with electrons, the atom is
stable and not likely to react.

Ions are atoms that are no longer neutral.
 + & - attract each other to form an ionic
bond
 An ionic bond is formed when a cation &
anion bond. A nonmetal & metal bond!
 Metals form positive (+) ions (Cations) by
losing electrons
 Nonmetals form negative (-) (Anions) by
gaining electrons

 Ionization
Energy – amount of
energy used to remove an
electron.
 Ionization Energy – increases from
L to R across the period.
 Ionization Energy – decreases
from the top of a group to the
bottom.
Covalent Bond – 2 atoms sharing a pair
of valence electrons. Ex. P 166
 Covalent Bond – 2 nonmetal elements
 Single bond – 1 pair of electrons

Double bond – 2 pairs of electrons
 Triple bond – 3 pairs of electrons

Covalent bond – Sharing of electrons
 Molecule is a neutral (no charge) group of
atoms joined by a covalent bond.
 Diatomic means 2 atoms (7 elements – are
only covalent – Br, H, N,O,F,I,Cl
 Remember - Brighties Have NO Fights In Class
 L side of periodic table has a greater
attraction for electrons. Top of the group has
greater attraction.

Ionic Compounds – can be shown by a
chemical formula. Shows the ratio (3:2)
of the compounds.
 Ionic Crystals – shape determined by the
arrangement of ions in its lattice.
 Ionic Crystals – classified by shape
 Have a high MP & good conductors.
 Properties explained by strong
attractions among ions within a crystal
lattice.
 Ex. P
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Polar covalent bonding
– electrons unevenly
shared
Polar bonds have
opposite sides like a
magnet + / Polar covalent bonding
means the atom with
the greater charge has
a partial negative
charge. Greek symbols
are used to show this.


Nonpolar covalent
bonding means the
atoms share the
electrons equally.
(Diatomic molecule
do this).
Like atoms form
nonpolar bonds.
There are forces of attraction between
molecules. (Ex. Water surface tension)
 Polar molecules have stronger
attractions than nonpolar molecules.
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Symbols used to represent atoms & compounds.
Elements use 1,2, & 3 letters.
Compounds are symbols & numbers.
In formulas the small # after the letter is a subscript
This shows the number of atoms in the molecule.
A superscript may show the charge if an ion is
formed.
A chemical formula tells which elements are
present & how much of the atom.