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Transcript
Atomic Structure
Atomic Structure-The BIG Picture
Discovery of the components of the atom and subsequent
modeling of the atomic structure led to explosive advances in
chemistry, medicine, and energy
Chemistry
• The nature of the chemical
bond
• New molecule synthesis
• Predictions about reactivity
• Information about how
reactions
work
• Electronics / computer
development
• New analytical (measuring)
methods
• Emergence of the field of
Nuclear Chemistry
ATOMIC STRUCTURE
Energy
Medicine
• Isotope tracers
• New drugs
• Cancer treatments
• New cell screening
methods
• Nuclear fission
• Nuclear fusion
• Power plants
• Understanding of the
nature of the sun,
planets, stars, etc
• Weapons
A History Lesson
• Not only does history help you become an
educated person, it helps you understand
current theories if you see how they
developed…
A History Lesson
• When you learn this stuff, try to put yourself in
the time of the person making the discovery.
– What was it like back then?
– How did society influence thinking?
– Would you have helped to make this
discovery?
– Do you think another problem of that day was
more important?
A History Lesson
• People were thinking about
the ATOM at a time when
indoor plumbing and
electricity were not
available, priorities to work
on, or imminently possible!
A History Lesson
1. Thousands of years ago, a Greek
philosopher Democritus speculated if
a piece of matter were divided in half
enough times, you’d finally find the
smallest piece that could not be subdivided any further
A History Lesson
a. This little piece of matter would have
the same properties as the big piece.
b. It was called the ATOM.
c. He also believed:
– matter could not be created, destroyed, or
further divided
– matter is mostly empty space
A History Lesson
• So what made the atom of one type of matter
different than another?
• The Greek philosophers thought SHAPE
(geometry) was the key difference.
• The Greeks were great practitioners of
Geometry…
• I wonder if THAT influenced their thinking on
the ATOM?
A History Lesson
Let’s mark this for later
Democritus: Shape
A History Lesson- 1800 years
later…
2. Dalton’s (1766-1844) experimentation on
matter led him to believe:
a. All atoms were spherical in shape but
differed from one another by mass.
• Mass was a big deal in Dalton’s time.
• Maybe it influenced his thinking?
A History Lesson
2. Dalton’s (1766-1844) experimentation on matter led him
to believe:
b. All matter is composed of atoms
c. All the atoms for a given element were identical*.
Atoms of a specific element are different from atoms
of another element.
d. Atoms cannot be created, divided* or destroyed.
e. Atoms could combine to make compounds only in
whole number ratios
f. In a chemical reaction, atoms are separated,
combined or rearranged.
*later shown to be incorrect (Make sure you marked them!)
A History Lesson
• These statements have come to be known
as Dalton’s Atomic Theory.
• Dalton was right about the MASS part of
his theory and how they combine to make
compounds
*Let’s mark this for later: Dalton: mass
A History Lesson
• Near the turn of the 20th century, evidence
began to emerge that suggested charged
subatomic parts made up the atom…
A History Lesson
3. Thomson’s (1897) experimentation led
to his discovery of a negativelycharged subatomic particle!
a. The negatively charged particle is called
the electron.
b. Discovered while studying electricity with a
cathode ray tube
Cathode Ray Tube (CRT)
A History Lesson
• Thomson inferred from his data that the
ATOM must be composed of a
combination of (+) charged “matrix” with
(–) charged particles (electrons)
dispersed in it
• like raisins or plums in a pudding
c. Developed the “plum pudding” model
of the atom
Progression of the Atomic Model…Discovery of the electron
“Plums”
or
Thomson’s “Plum Pudding”
Model
• He discovered the electron
(link) in 1897 before the
nucleus was discovered
• Later discoveries invalidated
this model
“Pudding”
J.J. Thomson in
Philosophical Magazine,
1904
“... the atoms of the elements
consist of a number of
negatively electrified
corpuscles enclosed in a
sphere of uniform positive
electrification, ... “
• By the way, the cathode ray tube, which
Thomson used to generate and study electrons
evolved into the television set (or at least
became the central component in TVs).
Lets mark this for later:
Thomson: electron
A History Lesson- Refining the Atom
4. Millikan (1909) used Thomson’s work to
determine:
a. the electron charge (-)
b. the mass of an electron is 9.11x10-28
grams
c. his experiment was the oil drop
experiment
Millikan’s Oil Drop Experiment
Millikan later determined the mass and charge of the electron
A History Lesson- Refining the Atomic
Model
5. Rutherford (1911) proved:
a. The (+) in an atom was not spread out but
concentrated in a central location called
the nucleus.
b. The volume of an atom is mostly empty
space!
Gold Foil Experiment
Gold Foil Experiment:
The Explanation
A History Lesson- Refining the Atomic
Model
c. Rutherford’s gold foil experiment:
i. He bombarded thin, gold foil with heavy,
positively (+) charged He (alpha) particles and
most passed through the foil.
ii. Occasionally, the particle would bounce back (as
if it were a tennis ball hitting a brick wall!)
iii. Rutherford assumed the deflected particles hit a
dense, gold nucleus and could not pass because
of the large mass of the gold nucleus and its (+)
charge
Plum
Pudding
vs.
Gold Foil
Experiment
A History Lesson- Refining the Atomic Model
Rutherford’s famous gold foil experiment:
• showed that the positive charge of the atom MUST be
concentrated in a tiny, yet heavy volume he called the nucleus
• almost ALL of the mass of the atom is in the nucleus
• very light electrons surround this nucleus
• the volume that an atom occupies is mostly empty space
Gold foil
animation
If a nucleus were as big as you are wide, the edge of its atom (outermost
electron orbital) would be over a mile away!
About 1.25 miles
.
• *Let’s mark this for later: Rutherford:
proton, nucleus
Click 
A History LessonRefining the Atomic Model
6. Bohr (1913)- suggested electrons must
move around in well-defined orbits or
energy levels
a. His experiments suggested that electrons
reside at different energy levels
because it took more (or less) energy to
knock them loose from an atom
*Lets mark this for later: Bohr: planetary
orbit of the electrons around the
nucleus
A History Lesson- Refining the Atomic
Model
7. Chadwick (1932) discovered the
neutron
a. The neutron is a particle with no charge but
about the same mass as a proton.
Chadwick’s experiment
• This history lesson gets us closer to the
modern atomic model.
• We still need to understand how the
electrons behave.
• BUT we’ll do that in the next unit…
II. The Subatomic Particles
A. Protons (p+)
1.Positively-charged subatomic
particle
2. Contained in the nucleus
3. Confirmed by Rutherford
4. Mass: 1.67 x 10-24 g (1840x more
massive than an electron!)
II. The Subatomic Particles
B. Neutron (n0)
1.Not charged subatomic particle
2. Contained in the nucleus
3. Discovered by Chadwick
4. Mass: 1.67 x 10-24 g
(same as proton)
II. The Subatomic Particles
C. Electrons (e-)
1. Negatively charged subatomic
particle (the charge is equal and
opposite to the charge of the proton)
2. Surrounding the nucleus
3. Mass: 9.11 x 10-28 gram
4. Tiny mass but occupies the majority of
the volume of the atom
II. The Subatomic Particles
C. Electrons (e-)
5. Each electron has an “electronic
address”
• Each resides in a well defined
energy level some distance from
the nucleus. The further from the
nucleus, the higher the energy
level.
6. Responsible for chemical bonding
III. Current Model of the Atom
• Spherically-shaped
• Small, dense positively-charged nucleus
surrounded by a cloud of negativelycharged electrons
• Most of the atom is empty space
• >99% of mass is in the nucleus
• Very small (there are 6.5 x1021 atoms in a
drop of water)
• Nucleus is held together by strong nuclear
forces
• THE STRONG NUCLEAR FORCE is the
name given to the attractive force that
holds protons and neutrons together
in the nucleus.
• If you think it is weird that likecharged protons can get together in
an atom’s nucleus without flying
apart, I don’t blame you!
• Don’t worry about the theory of how
this force works.
• Just remember, the STRONG
NUCLEAR FORCE is able to overcome
like-charge repulsion and hence
atomic nuclei are quite stable.
Here’s the atom so far:
Nucleus
PROTONS
NEUTRONS
RESPONSIBLE
FOR MASS,
IDENTITY OF THE
ELEMENT
valence shell
RESPONSIBLE
FOR CHEMICAL
BONDING
core electrons
RESPONSIBLE FOR NUCLEAR
SHIELDING
IV. The Subatomic Particles
A. Atomic Number
1. Protons are responsible for the nuclear
charge
2. # protons = atomic number
+
+
NUCLEUS
+
+
+
+
PROTONS:
RESPONSIBLE FOR NUCLEAR
CHARGE & = ATOMIC NUMBER
IV. The Subatomic Particles
3. This is the big (or top) number shown on the
periodic table
4. # of protons identifies the element
If for some reason the number of protons
changes (like a nuclear reaction), the
ELEMENT CHANGES!
THE NUCLEUS
+
+
+
NUCLEUS
+
+
+
PROTONS:
RESPONSIBLE FOR NUCLEAR
CHARGE & = ATOMIC NUMBER
NEUTRONS
“STRONG NUCLEAR FORCE” HOLDS THE NUCLEUS TOGETHER. IT OVERCOMES THE
REPULSIVE FORCE OF “LIKE CHARGES” (REMEMBER COULOMB’S LAW)
IV. The Subatomic Particles
Practice
1. Determine the number of protons in:
a. Fluorine
9
b. Magnesium
12
2. Identify the element:
Zinc
a. 30 protons
Chlorine
b. 17 protons
Sodium
c. 11 p+
Hydrogen
d. 1 p+
IV. The Subatomic Particles
• Remember, an atom is the smallest particle of an
element that retains the identity of that
element.
• When the number of protons (+) and the
number of electrons (-) are the same, the atom
is neutral (has no net charge).
• An ion is a charged atom or group of atoms
bonded together.
• An ion can be positive or negative.
IONS ARE CHARGED ATOMS.
IF AN ATOM GAINS ELECTRONS SO THAT IT HAS MORE ELECTRONS
THAN PROTONS, IT IS A NEGATIVELY CHARGED ATOM CALLED AN
ANION
a negative ion
IF AN ATOM LOSES ELECTRONS SO THAT IT HAS FEWER ELECTRONS
THAN PROTONS, IT IS A POSITIVELY CHARGED ATOM CALLED AN
electron
+
Ca ion
Atom about to
become a cation
Atom about to
become an anion
electron
Neutral Beryllium
2+
Be
Beryllium Ion
+
-
+
-
+
-
+
-
Neutral Nitrogen
Nitrogen Ion
“Nitride Ion”
+
-
+
-
+
-
+
-
+
-
+
-
+
-
3N
For a neutral atom: p+ = eCharge of an atom or ion: p+ - e- = Charge of Ion
Atoms are always neutral
Ions have a charge
Example
What are the charges of:
Lithium has 3 p+ and 3 e-
3–3=0
Lithium has 3 p+ and 2 e-
3 – 2 = +1
Oxygen has 8 p+ and 8 e-
8–8=0
Oxygen has 8 p+ and 10 e-
8 – 10 = -2
Charges = Oxidation Numbers
IV. The Subatomic Particles
B. Mass Number
1. Both neutrons and protons are
responsible for nearly all the atom’s
mass
2. # protons + # neutrons = mass
number
Ex: An oxygen atom has 8 protons and 8
neutrons and has a mass number =
16
__
Practice with Atomic number and Mass number
# of protons + # neutrons = mass number
A carbon atom with 6 protons and 6 neutrons
12
has a mass number = ____
# of protons = atomic number
The atomic number of carbon
6
is ___.
Number of electrons will equal the number of protons for an atom with NO NET
CHARGE
Ex 1:
How many p+, e- and n0 are in an atom of
Neon with a mass # 22?
10 p+
• Neon’s atomic number is 10  ___
• Mass number = protons + neutrons
22
= 10
+ neutrons
12 = neutrons
• If this atom is electrically10
neutral,
protons = electrons  ___ e-
Ex 2:
Determine which element has a mass # of
23 and contains 12 n0.
Mass number = protons + neutrons
23
= protons + 12
11 = protons
p+ = atomic number
 the element with 11 protons is
Na (sodium)
_______________
Practice
atomic
#
mass # # of
p+
# of
no
# of
e-
Atomic
mass
symbol
7
14
7
7
7
14.007
N
9
19
9
10
9
18.998
F
19
39
19
20
19
39.098
K
27
59
27
32
27
58.933
Co
Practice
2. If 2 protons were removed from the
nucleus of an oxygen atom, what nucleus
remains?
O: 8 p+
- 2 p+
Carbon
6 p+  ________
IV. The Subatomic Particles
C. Isotopes
1. All atoms of an element must have the
same number of protons
2. BUT they may have a different number
of neutrons
3. All atoms of an element must have the
same atomic # but can have a different
mass #.
4. These are called isotopes.
isotopes
For Isotopes
Parts of the Atom
Isotopes of Lithium
Lithium – 6
Mass # 
Atomic # 
6
3
Li
Lithium - 7
7
3
Li
ISOTOPES
Atoms in the same element with different MASS NUMBER but identical ATOMIC NUMBER.
NUCLEI OF ATOMS
IN THE SAME ELEMENT
+
+
+
+
+
UNRAVEL
+
+
+
ISOTOPE SYMBOL
12
+
6
+
6 PROTONS
6 NEUTRONS
C
ISOTOPES
Atoms in the same element with different MASS NUMBER but identical ATOMIC NUMBER.
NUCLEI OF ATOMS
IN THE SAME ELEMENT
+
+
UNRAVEL
+
+
+
+
+
ISOTOPE SYMBOL
12
+
+
6
+
C
6 PROTONS
6 NEUTRONS
+
+
+
+
+
+
13
+
+
6
+
+
6 PROTONS
7 NEUTRONS
C
A
Z
X
X –ELEMENT SYMBOL
A-MASS NUMBER
Z-ATOMIC NUMBER
Weighted Averages
• The atomic mass of one atom is TINY if
reported in kilograms or even grams.
• Atomic mass is reported in atomic mass
units (amu).
• 1 amu is about the mass of one proton
or neutron (about 1.67 x 10-24)
• Generally, atoms of a given element will
have 1 isotope in high abundance and
several others in much smaller numbers.
• To calculate the average atomic mass, we
use a WEIGHTED AVERAGE that accounts
for the abundance of each isotope.
Ex 1:
• Carbon-12 makes up 98.89% of naturallyoccurring carbon. Carbon-13 makes up
1.11% of naturally occurring carbon. Use
this information to determine the average
atomic mass of carbon.
Average atomic mass =
(12 X 0.9889) + (13 X 0.0111) =
12.0111 amu
Ex 2:
• Chromium has 4 naturally-occurring
isotopes. Their abundance is as follows:
Cr-50 – 4.35%, Cr-52 – 83.79%, Cr-53 –
9.50%, and Cr-54 – 2.36%. Determine the
average atomic mass for chromium.
(50 X .0435) + (52 X .8379) + (53 X .0950) + (54 X .0236) =
52.0552 amu
Practice:
1. The element copper is found to contain 69.1%
of copper-63 and 30.9% of copper 65.
Calculate the average atomic mass of copper.
63.618 amu
2. Gallium occurs in nature as a mixture of two
isotopes. They are gallium-69 with an
60.108% abundance and gallium-71 with a
39.892% abundance. Calculate the atomic
mass of gallium.
69.7978 amu
3. Neon has 2 isotopes: neon-20 and neon22. Use the information from the periodic
table to determine which occurs in greater
abundance.
Neon – 20
Because the atomic
mass on the periodic
table rounds to 20 and
not 22
4. Use the table below to calculate the
atomic mass of Element X and then
identify it.
Isotope
% Abundance
16X
99.762
17X
.038
18X
.2
(16 X .99762) + (17 X .00038) + (18 X .002) =
16.0044 amu
Oxygen