Download Chapter 2/Unit 2: Matter is Made of Atoms

Chapter 2/Unit 2:
Matter is Made of Atoms
Notes
Chemistry CPA
Unit Goals
• After learning this chapter you will be able to:
Identify “who” discovered each part of the atom
Describe the historic and present models of the atom
Label and describe the function of each part of the
atom (nucleus, proton, neutron, electrons)
Define and identify an isotope of any element
Calculate the average atomic mass of elements as
listed on the periodic table
Describe the connection between light waves and
electron energy
Atomic Theory Timeline
Atomic Models
• This model of the atom
may look familiar to
you.
• This is the Bohr model.
• In this model, the
nucleus is orbited by
electrons, which are in
different energy levels.
• A model uses familiar
ideas to explain
unfamiliar facts
observed in nature.
• A model can be
changed as new
information is collected.
• The atomic model
has changed
throughout the
centuries, starting
in 400 BC, when it
looked like a
billiard ball →
Who are these men?
In this lesson, we’ll learn about
the men whose quests for
knowledge about the
fundamental nature of the
universe helped define our
views.
Democritus
• This is the Greek
philosopher Democritus
who began the search for a
description of matter more
than 2400 years ago.
– He asked: Could matter
be divided into smaller
and smaller pieces
forever,
– Or was there a limit to
the number of times a
piece of matter could be
divided?
400 BC
Atomos
• His theory: Matter could
not be divided into smaller
and smaller pieces forever,
eventually the smallest
possible piece would be
obtained.
• This piece would be
indivisible.
• He named the smallest
piece of matter “atomos,”
meaning “not to be cut.”
This theory was ignored and
forgotten for more than 2000
years!
Why?
 The popular
philosophers of
the time,
Aristotle and
Plato, had a
more
respected, (and
ultimately
wrong) theory.
Aristotle and Plato favored the earth, fire, air and
water approach to the nature of matter. Their
ideas were most believed because of their
popularity as philosophers. The atomos idea was
buried for approximately 2000 years.
Dalton’s Model
In the early 1800s,
the English Chemist
John Dalton
performed a number
of experiments that
eventually led to the
acceptance of the
idea of atoms.
Dalton’s Theory (The Four Postulates)
1. He deduced that all
elements are composed of
atoms. Atoms are indivisible
and indestructible particles.
2. Atoms of the same element
are exactly alike.
3. Atoms of different elements
are different.
4. Compounds are formed by
the joining of atoms of two
or more elements.
.
This theory
became one of
the
foundations of
modern
chemistry.
Thomson’s Plum Pudding Model
In 1897, the
English scientist
J.J. Thomson
provided the first
hint that an atom
is made of even
smaller particles.
Thomson Model
• He proposed a
model of the atom
that is sometimes
called the “Plum
Pudding” model.
• Atoms were made
from a positively
charged substance
with negatively
charged electrons
scattered about, like
raisins in a pudding.
Thomson Model
Thomson studied
the passage of an
electric current
through a gas.
As the current
passed through
the gas, it gave off
rays of negatively
charged particles.
Thomson Model
This surprised
Thomson, because
the atoms of the
gas were
uncharged. Where
had the negative
charges come
from?
Where did
they come
from?
Thomson concluded that the negative
charges came from within the atom.
A particle smaller than an atom had to
exist.
The atom was divisible!
Thomson called the negatively charged
“corpuscles,” today known as
electrons.
Since the gas was known to be neutral,
having no charge, he reasoned that
there must be positively charged
particles in the atom.
But he could never find them.
Rutherford’s Gold Foil Experiment
• In 1908, the English
physicist Ernest
Rutherford was hard
at work on an
experiment that
seemed to have little
to do with
unraveling the
mysteries of the
atomic structure.
• Rutherford’s experiment Involved firing a
stream of tiny positively charged
particles at a thin sheet of gold foil (2000
atoms thick)
– Most of the positively
charged “bullets” passed
right through the gold
atoms in the sheet of gold
foil without changing
course at all.
– Some of the positively
charged “bullets,” however,
did bounce away from the
gold sheet as if they had hit
something solid. He knew
that positive charges repel
positive charges.
• This could only mean that the gold atoms in the
sheet were mostly open space. Atoms were not a
pudding filled with a positively charged material.
• Rutherford concluded that an atom had a small,
dense, positively charged center that repelled his
positively charged “bullets.”
• He called the center of the atom the “nucleus”
• The nucleus is tiny compared to the atom as a whole.
Rutherford
• Rutherford reasoned
that all of an atom’s
positively charged
particles were
contained in the
nucleus. The negatively
charged particles were
scattered outside the
nucleus around the
atom’s edge.
Bohr Model
• In 1913, the Danish
scientist Niels Bohr
proposed an
improvement. In his
model, he placed
each electron in a
specific energy
level.
Bohr Model
• According to Bohr’s
atomic model,
electrons move in
definite orbits
around the nucleus,
much like planets
circle the sun. These
orbits, or energy
levels, are located at
certain distances
from the nucleus.
Wave Model
The Wave Model
• Today’s atomic
model is based on
the principles of
wave mechanics.
• According to the
theory of wave
mechanics, electrons
do not move about
an atom in a definite
path, like the planets
around the sun.
The Wave Model
• In fact, it is impossible to determine the exact
location of an electron. The probable location of an
electron is based on how much energy the electron
has.
• According to the modern atomic model, at atom has
a small positively charged nucleus surrounded by a
large region in which there are enough electrons to
make an atom neutral.
Electron Cloud:
• A space in which electrons
are likely to be found.
• Electrons whirl about the
nucleus billions of times in
one second
• They are not moving around
in random patterns.
• Location of electrons
depends upon how much
energy the electron has.
Electron Cloud:
• Depending on their energy they are locked into a
certain area in the cloud.
• Electrons with the lowest energy are found in the
energy level closest to the nucleus
• Electrons with the highest energy are found in the
outermost energy levels, farther from the nucleus.
Indivisible Electron
Greek
X
Dalton
X
Nucleus
Thomson
X
Rutherford
X
X
Bohr
X
X
Wave
X
X
Orbit
Electron
Cloud
X
X
Atomic Structure
What’s in an Element?
• Each atom can be classified as one of the
elements on the periodic table.
• Each element has the atomic number (which
is the # of protons), the element symbol, the
element name, and the average atomic mass.
The parts of an atom
• Parts of an atom:
– Nucleus
– Electron Cloud
• Subatomic particles:
– Protons (in the nucleus)
– Neutrons (in the nucleus)
– Electrons (outside the
nucleus in the electron cloud)
The Nucleus
• The “core” of the atom
• Protons and neutrons are found in the nucleus of
an atom
• Contains 99.9% of the mass of the atom
• The MASS NUMBER
– is the mass of the nucleus
– a sum of the mass of the protons + the mass of the
neutrons:
Mass number = mass of protons + mass of neutrons
The Atomic Mass Unit
• In an atom, mass is measured in “atomic mass
units”, or “amu”.
• Protons have a mass of 1 amu
• Neutrons have a mass of 1 amu
• Electrons have a mass of 1/1837 of an amu –
so small, we don’t even count it!
• So…which particles give the atom it’s “mass”?
____________________
Protons
• Positively charged
subatomic particle in an
atom’s nucleus
• Gives the atom it’s
“identity”
• Has a mass of 1 amu
• The number of protons is
the atom’s ATOMIC
NUMBER:
#PROTONS = ATOMIC
NUMBER
Neutrons
• A subatomic particle
that has no charge
• Found in the nucleus of
the atom
• Contributes to the mass
of an atom
• Has a mass of 1 amu
• # of neutrons = mass
number - #protons
Electrons
• Are found in the “electron cloud” which is the
space outside of the nucleus.
• Are negatively charged particles
• Have almost no mass
• Are responsible for all chemical reactions and
bonding that happens with other atoms, which
means they give the atom its chemical properties
• In an electrically neutral atom:
# protons = #electrons
More on the nucleus…
• …a nucleus generally has an equal number of
protons and neutrons.
• When a nucleus has a different number of
neutrons than protons, it is called an
“isotope”.
Isotopes
• Atoms of the same element with different mass
numbers.
• Different number of neutrons
• Nuclear symbol:
Mass #
12
Atomic #
6
• Hyphen notation: carbon-12;
carbon-13
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
C
Calculating Mass # Practice
• RECALL:
MASS # = # PROTONS + # NEUTRONS
• Mass numbers are always WHOLE NUMBERS
(they aren’t the decimal numbers found on
the periodic table)
Practice:
Element
# protons
# neutrons
Carbon
6
Carbon
8
Mass #
17
Cl
Isotope Practice
37
• Chlorine-37
– atomic #:
17
– mass #:
37
– # of protons:
17
37
– # of electrons:
17
– # of neutrons:
17
20
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Cl
Isotopes
Isotope notation
• Element name – Mass number
• Examples: Carbon-14, Chlorine - 37
• Practice:
# protons
# neutrons
Carbon - ____
6
Carbon - ____
7
Carbon - ____
8
Mass #
Shorthand Notation
Mass number
SYMBOL OF ELEMENT
Atomic number
Examples: Write the shorthand notation for the
following Isotopes:
Carbon-14 Oxygen-18 Magnesium-25
Average Atomic Mass
• The number on the
periodic table is an
average of all of the
masses of all of the
isotopes that exist in
nature
• Based on percent
abundance of an
isotope’s occurrence in
nature.
Example
Average Atomic Mass Equation
Avg.
=
Atomic
Mass
(mass)(%) + (mass)(%)
100
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Average Atomic Mass Practice Problem
• EX: Calculate the avg. atomic mass of oxygen if its
abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20%
18O.
Avg.
(16)(99.76) + (17)(0.04) + (18)(0.20)
Atomic =
100
Mass
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
=
16.00
amu
Atomic Structure Recap:
• ATOMS
– Differ by number of protons
• IONS
– Differ by number of electrons
• ISOTOPES
– Differ by number of neutrons
Electrons are found in the electron
cloud
The cloud has regions of
space called energy
levels
• The first energy level
holds 2 electrons
• The second energy level
holds 8 electrons.
• The third energy level
holds 18 electrons
Valence Electrons
• Are found furthest from
the nucleus
• Dictate the physical and
chemical properties of an
element
• Use the periodic table to
determine the number of
valence electrons.
• All atoms want 8 valence
electrons
Examples of electron filling
Lewis Dot Diagram
• A way to illustrate the
number of valence
electrons
– Use one dot for each
valence electron
– Place the dot around each
side of the symbol before
pairing the electrons
– The symbol represents the
nucleus plus all the inner
electrons for the element.
Lewis Dot Diagram Practice
H
O
N
F
Ne
Electrons and Light
• Electrons are normally in
the ground state
• When the atom is given
energy the electrons
move to the excited state.
• When the electrons lose
this energy they fall back
to the ground state and
emit (give off) light.
• Each element has a
unique emission
spectrum
Electromagnetic Spectrum
• Electromagnetic Radiation –
– A broad range of energetic emissions
– made up of photons
• Photons – bundles of energy
– Travel like waves
– Move at the speed of light = 3.0 x 108 m/s
– Electromagnetic waves do not require a medium
to move
Parts of the wave
• Amplitude – the height of the wave
• Wavelength – the distance between the two
successive waves
• Frequency – the number of waves that pass a
given reference point per second
Wave Units
• Wavelength = lamda unit is nanometer
• Frequency = nu in units of 1/s or s-1
What is the difference between these
waves?