Atomic History and
Structure:
What comes to mind when you think
of the term “atom”?
How do we know what we know about atoms? List any people
you can think of.
Thales of Miletus (________)
• Noticed what we call ________________ with
amber
– Things would be attracted to it when rubbed
– It was a “magical property”
• The term electron _________________
___________________________________
__________________
Kanada (~_________BC)
• Indian attributed with first proposing the
idea of atoms (called “________” or “____”)
• 5 elements
–
–
–
–
–
_______________
_______________
_______________
_______________
_______________
• Atoms were indestructable and eternal
Empedocles (450BC)
• 4 elements:
– _____________
– _____________
– _____________
– _____________
• Everything was different combinations of
these
• This idea didn’t really change until _______!
Leucippus (~_______ BC)
•Proposed the
idea of atoms
•That two things
exist
•__________
•__________
Democritus (_______)
•Student of Leucippus
•Matter is made up of “eternal, indivisible,
indestructible and infinitely small
substances which cling together in different
combinations to form the objects
perceptible to us”
•“_________”
From
:http://www.historyworld.net/wrldhis/PlainTextHistor
ies.asp?historyid=ac20#ixzz1UvX6le4i
100 Greek Drachma,
1967
 Aristotle 384 BC – 322 BC
•Originally opposed the idea of
atoms, then
•Added ____________or
______________ to the four
elements:
•earth (cold and dry)
•air (hot and moist)
•fire (hot and dry)
•water (cold and moist)
•The differences in matter where a
result of ____________________
________________________
•Changing the balance could
change matter
•ex: what we know as
copper changed to gold
Benjamin Franklin (_____________)
 Franklin believed object had 1 of 2 charges (+/-)
 Opposites attract, like charges repel (Coulomb’s Law,
which the Greeks knew a little about)
 Kite experiment (among others):
 Electric charges run from + to –
 ________________________
 Words he gave us:
 ___________________________________________
___________________________________________
_______________________________
J.L. Proust (_____*)
• Law of constant composition:
– ____________________________________
____________________________
– In other words…a given compound always
has the same composition, regardless of
where it comes from.
• Ex: H2O is ______________________
______________________________
*not published or recognized until 1811
Dalton’s Atomic Theory ~____
•
John Dalton (1766-1844)
proposed an atomic
theory
•
While this theory was not
______________
____________________
________________and
brought about chemistry
as we know it today
instead of alchemy
Dalton’s Atomic Symbols
Dalton’s Atomic Theory
Problems with Dalton’s Atomic Theory?
1. matter is composed of indivisible particles
____________________________________________
2. all atoms of a particular element are identical
_______________________________________________
_______________________________________________
______________________________________
3. different elements have different atoms
YES!
4. atoms combine in certain whole-number ratios
YES! Called __________________________________
5. In a chemical reaction, atoms are merely rearranged to
form new compounds; they are not created, destroyed, or
changed into atoms of any other elements.
Yes, except __________________________________
____________________________________________
____________________________________________
Michael Faraday (______)
 atoms contain particles with ___________
______________
 structure of atoms related to electricity
 The electron was the fundamental
________________________________
JJ Berzelius (__________)
• Came up with how we write chemical
formulas
– _____________ for elements
– _______________to indicate numbers of
each element (he used superscripts,
though!)
– Considered one of the fathers of
modern chemistry
• Along with
–John Dalton
–Antoine Lavoisier
–Robert Boyle
Up until the 1900’s….
• Atomic structure was thought about, but
not well known. It took a few more
people to really put things together, and
build off of each other’s knowledge to
come up with what we know today.
• Lord William
Thomson Kelvin
(________)
– Proposed the Plum
Pudding Model, but
______________
• ________________
________________
________________
________________
______
JJ Thomson
• Discovered __________
(_____)
– cathode ray tube
– Called electrons corpuscles
• Name electron came from
George Johnstone Stoney,
who proposed the concept in
1874 and 1881, and the word
came in 1891
• Named the “Plum Pudding”
model of the atom
(________)
Cathode Ray Tube
Hantaro Nagaoka (______)
• Proposed the planetary(Saturnian) model
of the atom
– _______________________
– Electrons bound to the nucleus via
________________________
• Both were _____________ by Rutherford
• He abandoned the model in ______ due to
errors that were not confirmed by new
studies (charged rings)
Rutherford’s Gold Foil Experiment
Gold Foil Animation
– alpha (α) particles: _______
___________directed at thin
metal foil
– most particles made it through
→ _____________
– others were deflected back →
since alpha particles are
positive, they had to bounce off
of something _________
So…there is a dense
__________________________
__________________________
__________________________
__________________
Rutherford’s experiment led to the
nuclear view of the atom (_______/
published _____)
(side note- it was actually Geiger- Marsden Experiment. Scientists Hans G.
and undergraduate Ernest M. worked for Rutherford.)
“It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell
at a piece of tissue paper and it came back and hit you. On consideration, I realized that this scattering backward must be the result of a
single collision, and when I made calculations I saw that it was impossible to get anything of that order of magnitude unless you took a
system in which the greater part of the mass of the atom was concentrated in a minute nucleus. It was then that I had the idea of an atom
with a minute massive center, carrying a charge.[2]”
—Ernest Rutherford
Gold Foil and the Models of the Atom
James Chadwick (
)
• Worked with ___________
______________________.
• Proved the existence of the
________________.
• same mass as a proton, but
with _______________
• its mass was about ______
______ than the proton's.
JJ Thomson
• Determined _____________
__________ (_______)
– Used anode rays
– Found Ne deflected in two
different paths using what
we now call mass
spectroscopy
R. A. Millikan - Measured the charge of the electron
(1909).
In his famous “oil-drop” experiment, Millikan was able to
determine the charge on the electron independently of its
mass. Then using Thompson’s charge-to-mass ratio, he
was able to calculate the mass of the electron.
e = 1.602 10 x 10-19 coulomb
e/m = 1.7588 x 108 coulomb/gram
m = 9.1091 x 10-28 gram
Goldstein - Conducted “positive” ray experiments that
lead to the identification of the proton. The charge
was found to be identical to that of the electron and
the mass was found to be 1.6726 x 10-24 g.
Millikan’s Experiment
X-rays
.
Millikan’s Experiment
- X-rays give some electrons a charge
- Some drops would hover (not fall)
- From the mass of the drop and the charge on
the plates, he calculated the mass of an electron
Millikan oil drop experiment
• Millikan did another experiment to determine
the mass of the –ve particles (electrons). The
experiment used mainly to determine the
magnitude of the electron charge and using
e/m to get m- value.
30
Niels Bohr (1885-1962)
•
Bohr Model or the Solar System Model
– Niels Bohr in ________ introduced his _______
______________________________________
– Electrons _______________________, which
are also called _________________.
– An electron can “jump” from a lower energy
level to a higher one upon absorbing energy,
creating an excited state.
– The concept of energy levels accounts for the
emission of distinct wavelengths of
electromagnetic radiation during flame tests.
Bohr’s Orbit Model (1913)
Electrons occupy
orbitals around the
nucleus according to
their _______.
Glenn Seaborg
(1912-1999 )
• Discovered ___ new
elements.
• Only living person
for whom ______
________________
_____________.
Which brings us to the
modern day view of the
atom….
ATOMIC
STRUCTURE
The atom is mostly
___________________
•protons and neutrons in the _______________.
•the number of electrons is ______________the number of protons.
•electrons in space ______________________.
•extremely small.
•One teaspoon of water has _______________
____________________________________________________
__________________________.
ATOMIC COMPOSITION
• Protons (___)
– positive (+) electrical charge
– mass = 1.672623 x 10-24 g
– relative mass = 1.007 atomic mass units (____)
• but we can round to 1
• Electrons (___)
– negative (-) electrical charge
– relative mass = 0.0005 amu
• but we can round to 0
• Neutrons (___)
– no electrical charge
– mass = 1.009 amu
• but we can round to 1
The following four slides are for
additional information only; you will
not be tested on the fundamental
particles. However, they could appear
as extra credit on a test or quiz.
Subatomic Particles can also be further
broken down into Fundamental Particles
• Quarks
– component of protons & neutrons
– 6 types
• Up, down
• Strange, charm
• Top, bottom
• 3 quarks = 1 proton or 1 neutron
He
Subatomic Particles and Quarks
What about electrons?
• Electrons are
electrons
• They are not made
from quarks
• Which is why
they weigh so
much less than
p+ or no
• Classified as a
lepton
Subatomic Particles
More information at http://www.lns.cornell.edu/~nbm/NBM_INTRO_TO_HEP1.htm
Atomic Number, Z
All atoms of the same element have
the same ____________
__________in the nucleus, ___
13
Al
26.981
Atomic number
Atom symbol
AVERAGE Atomic Mass
Atoms are neutral because the
numbers of _____________________
- the opposite charges cancel.
–
• 11 electrons
• 11 negative charges
+
• 11 protons
• 11 positive charges
Ions
A charged atom because of a gain or loss of electrons.
If an atom is neutral, the __________________
If it has ___________, the atom has a 1+ charge
If it has ___________, the atom has a 1- charge
IONS
• Taking away electrons from an atom gives a
_____________________________
• Adding electrons to an atom gives an _______
______________________________
• Atoms may _____________________
• To tell the difference between an atom and an ion, look
to see if there is a charge in the superscript!
• Examples: Na+ Ca+2 I- O-2 compared to
Na Ca
I
O
PREDICTING ION CHARGES
In general
• metals lose electrons ---> _______________
• nonmetals gain electrons ---> ____________
Charges on Common Ions
By losing or gaining e-, atom has same number of
___________________________.
Mass Number, A
• C atom with 6 protons and 6 neutrons is the mass
standard
– = ____________________________
• Mass Number (A)
– =____________________________
A
10
Z
5
• NOT on the periodic table…(that is the AVERAGE atomic
mass on the table)
• Ex: A boron atom can have
______________________
A =
B
Atomic Math
On periodic table- but not all PTs look exactly like
this set up, but they have the same information
Think Back…
• John Dalton stipulated that all atoms of a
particular element were identical
– ______________________________________
______________________________
• In 1912, J.J. Thomson discovered that this
was not accurate
– In an experiment measuring the mass-tocharge ratios of positive ions in neon gas, he
made a remarkable discovery:
• _________________________________
• _________________________________
• All of the atoms had 10 protons, however some
had ________________
Isotopes
•
atoms with the same number of protons (___) but a
different ___________________________
– same element, different ____________________
1H (___________):
A=1
Z=1
2H (___________):
A=2
Z=1
3H (___________):
A=3
Z=1
Isotopes &
Their Uses
Isotopes & Their Uses
The _____________ content of ground
water is used to discover the source of the
water, for example, in municipal water or the
source of the steam from a volcano.
Learning Check
Which of the following represent isotopes
of the same element? Which element?
234
92
X
234
93
X
235
92
238
X
92
X
Atomic Math: Summary
• Atomic number (Z)
– ________________________________
– ________________________________
• (Atomic) Mass Number (A)
– ______________________________________
______________________________________
• Atomic Mass (also called Atomic Weight)
– _______________________________(accounts for
all the isotopes) is ___________________
Counting Protons, Neutrons, and
Electrons
• Protons: Atomic Number (from periodic table)
• Neutrons: Mass Number minus the number of
protons (mass number is protons and neutrons
because the mass of electrons is negligible)
• Electrons:
– If it’s an atom, the protons and electrons must be
the SAME so that it is has a net charge of zero
(equal numbers of + and -)
– If it does NOT have an equal number of electrons, it
is not an atom, it is an ION. For each negative
charge, add an extra electron. For each positive
charge, subtract an electron (Don’t add a proton!!!
That changes the element!)
Learning Check – Counting
State the number of protons, neutrons, and electrons in
each of these ions.
39 K+
16O -2
41Ca +2
19
8
20
#p+ ______
______
_______
#no ______
______
_______
#e- ______
______
_______
Learning Check – Counting
Naturally occurring carbon consists of three isotopes, 12C,
13C, and 14C. State the number of protons, neutrons, and
electrons in each of these carbon atoms.
12C
13C
14C
6
6
6
#p+ _______
_______
_______
#no _______
_______
_______
#e- _______
_______
_______
Learning Check
An atom has 14 protons and 20 neutrons.
A. Its atomic number is
1) 14
2) 16
3) 34
B. Its mass number is
1) 14
2) 16
3) 34
C. The element is
1) Si
2) Ca
3) Se
D. Another isotope of this element is
1) 34X
2) 34X
3) 36X
16
14
14
Atomic Symbols: Nuclide Notation
 Nuclide_________________________________
 Show the name of the element, a hyphen, and the
mass number in hyphen notation
_______________
 Show the mass number and atomic number in
nuclear symbol from
mass number
atomic number
Nuclide notation: p+, charge, and average
atomic mass
Mass number
(________________)
37
Atomic number
17
(number of _______)
number of ________
A-Z =20
Cl
As atoms have no charge, the number
of electrons is the same as the number
of protons. This atom has ___________.
Nuclide notation – ions
Mass number
Atomic number
23
+
Na
11
number of neutrons=
1+ charge ______________ _____
than the number of protons. This
atom has __________.
Nuclide notation –ions
Mass number
Atomic number
16
2–
O
8
number of neutrons=
___charge means ________________
than the number of protons. This
atom has _____________.
Learning Check
Write the nuclear symbol form for the following
atoms or ions:
A. 8 p+, 8 n, 8 eB. 17p+, 20n, 17e-
C. 47p+, 60 n, 46 e-
___________
___________
___________
Learning Check
1. Which of the following pairs are isotopes of the same
element?
2. In which of the following pairs do both atoms have
8 neutrons?
A.
15X
15X
8
B.
12X 14X
6
C.
7
6
15X 16X
7
8
Isotopes and Average Atomic Mass
• We are used to calculating #’s of p+, no and eusing whole numbers; however on the Periodic
Table we often see a decimal number  Why?
• Atomic Mass (on the Periodic Table)
– The average of the isotopic masses _________
__________________________________________
__________________________________
– In a weighted average we must assign greater
importance – give greater weight – to the quantity
that occurs ______________________
Isotopes and Atomic Mass
• The atomic mass for each element on the
periodic table reflects the ____________
_________________________________ in
nature.
• The mass on the periodic table is ______
_____________________________________
_______________________________
AMUs and Atomic Weight
•________________(____) is the unit for
relative atomic masses of the elements
•1 amu =__________________________
•1 amu = 1.6605x10-24 grams
Protons (p+)
mass = 1.672623 x 10-24 g
relative mass = 1.007 atomic mass units (amu) but we can round to 1*
Electrons (e-)
relative mass = 0.0005 amu but we can round to 0*
Neutrons (no)
mass = 1.009 amu but we can round to 1*
*most times, like now; when we get to nuclear chemistry, we will not be able to!
Comparative Example – Your Grades
• To calculate your overall average,
we use a weighted average
instead of a simple average since
different tasks are worth more
• For example:
/100
Your
mark
Exams
30
80%
Course
work
30
75%
Applied 10
Science
70%
Final
70%
30
To Calculate Average Atomic Mass
• You add up _____________________________for each
isotope to get the weighted average
– Fractional abundance _____________________
• Ex: If something has 3 isotopes:
Example
• Naturally occurring copper exists with the
following abundances:
• 69.17% is Cu-63 w/ atomic mass 62.93 amu
• 30.83% is Cu-65 w/ atomic mass 64.93 amu
Learning Check:
3 Isotopes of Ar occur in nature
• 0.337% as Ar-36, 35.97 amu
• 0.063% Ar-38, 37.96 amu
• 99.6% Ar-40, 39.96 amu
• Calculate the Average Atomic Mass
• In J.J. Thomson’s experiment, he found that the
percent abundances of neon are as follows:
– Neon – 20 = 90.51%
– Neon – 21 = 0.27%
– Neon – 22 = 9.22%
• Calculate the average atomic mass of neon
showing all of your work
If a mass is not specifically given for an
isotope
• Then make the assumption that the mass is
the same as the atomic mass number
– It isn’t exactly correct, but it will be close
AVERAGE
ATOMIC MASS
11B
10B
• Boron is 20% 10B and 80% 11B. That is, 11B is 80
percent abundant on earth.
• For boron, atomic weight=
Calculating & Abundance
• Chlorine has two isotopes: chlorine-35 (mass
34.97 amu) and chlorine-37 (mass 36.97 amu).
• What is the percent abundance of these two
isotopes if chlorine's atomic mass is 35.453?
Problem 1
• The two naturally occurring isotopes of nitrogen are
nitrogen-14, with an atomic mass of 14.003074 amu, and
nitrogen-15, with an atomic mass of 15.000108 amu.
What are the percent natural abundances of these
isotopes?
• The atomic mass of nitrogen is 14.00674amu
End of Chapter