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.

•   4 elements:
•   Earth
•   Wind
•   Fire
•   Water
•   Everything was different combinations of these
•   Matter could be endlessly divided
•   This idea didn’t really change until1661!

•   Matter is not infinitely divisible
•   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/
PlainTextHistories.asp?
historyid=ac20#ixzz1UvX6le4i •   “Atomos”
100 Greek Drachma,
1967

§   Aristotle 384 BC – 322 BC
•   Originally opposed the idea of atoms, then
•   Added hot/cold or moist/dry 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 different balances of these atoms
•   Changing the balance could change matter
•   ex: what we know as copper changed to gold

§   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 –
§   Lightening is electricity

•   John Dalton
(1766-1844) proposed an atomic theory
•   While this theory was not completely correct, it revolutionized how chemists looked at matter and brought about chemistry as we know it today instead of alchemy

1.
Elements are made of very small indivisible particles called atoms.
2.
All atoms of a given element are identical
(all hydrogen atoms are identical)
.
3.
The atoms of an element are different than the atoms of another element
(hydrogen is different than helium)
.
4.
Atoms of one element can combine with the atoms of another element to make compounds. A given compound should have the same relative numbers and types of atoms.
5.
Atoms are indivisible in chemical processes … they are not created or destroyed just reorganized.

Problems with Dalton’s Atomic Theory?
1. matter is composed of indivisible particles
Atoms Can Be Divided, but only in a nuclear reaction
2. all atoms of a particular element are identical
Does Not Account for Isotopes (atoms of the same element but a different mass due to a different number of neutrons)!
3. different elements have different atoms
YES!
4. atoms combine in certain whole-number ratios
YES! Called the Law of Definite Proportions
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 for nuclear reactions that can change atoms of one element to a different element

•   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
(1903)
•   Proposed the Plum
Pudding Model, but didn’t name it
•   Electrons embedded in a positive, spherical cloud

JJ Thomson (1904)
•   Discovered electrons (1897)
•   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 (1904)

Cathode Ray Tube

Cathode ray tube

Gold Foil Animation
•   alpha
( α )
particles: positively charged particles directed at thin metal foil
•   most particles made it through → empty space
•   others were deflected back
→ since alpha particles are positive, they had to bounce off of something positive
So … there is a dense positive charge (nucleus) that the electrons move around.

Rutherford’s experiment led to the nuclear view of the atom (1909/ published 1911)
(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 15inch 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 (1932)
•   Worked with Ernest
Rutherford
•   Proved the existence of the neutron.
•   same mass as a proton, but with zero charge
•   its mass was about 0.1% more than the proton's.

JJ Thomson (1912)
•   Determined isotopes of atoms exist (1912)
•   Used anode rays
•   Found Ne deflected in two different paths using what we now call mass spectroscopy

Millikan ’ s  Experiment

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

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 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.

Niels Bohr (1885-1962)
•   Bohr Model or the Solar System Model
•   Niels Bohr in 1913 introduced his model of the hydrogen atom.
•   Electrons circle the nucleus in orbits, which are also called energy levels.
•   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 energy..

Glenn Seaborg
(1912-1999 )
•   Discovered 8 new elements.
•   Only living person for whom an element was named.

The atom is mostly empty space
•   protons and neutrons in the nucleus.
•   the number of electrons is equal to the number of protons.
•   electrons in space around the nucleus.
•   extremely small.
•   One teaspoon of water has 3 times as many atoms as the Atlantic Ocean has teaspoons of water.

•   Protons (p + )
•   positive (+) electrical charge
•   mass = 1.672623 x 10 -24 g
•   relative mass = 1.007 atomic mass units (amu)
•   but we can round to 1
•   Electrons (e )
•   negative (-) electrical charge
•   relative mass = 0.0005 amu
•   but we can round to 0
•   Neutrons (n o )
•   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.

can also be further broken down into Fundamental Particles
•   Quarks
•   component of protons & neutrons
•   6 types
•   Up, down
•   Spin, charm
•   Top, bottom
He
•   3 quarks = 1 proton or 1 neutron

•   Electrons are electrons
•   They are not made from quarks
•   Which is why they weigh so much less than p + or n o
•   Classified as a lepton

Subatomic Particles
More information at http://www.lns.cornell.edu/~nbm/NBM_INTRO_TO_HEP1.htm

All atoms of the same element have the same number of protons in the nucleus,
13
Al
26.981
Atomic number
Atom symbol
AVERAGE Atomic Mass

+
–

§   A charged atom because of a gain or loss of electrons.
§   If an atom is neutral, the # of p + = # of e -
§   If it has lost 1 e , the atom has a 1+ charge
§   If it has gained 1 e , the atom has a 1- charge

•   Taking away electrons from an atom gives a
CATION with a positive charge
•   Adding electrons to an atom gives an
ANION with a negative charge.
•   Atoms may gain or lose more than 1 e -
•   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
Na Ca I O compared to

In general
•   metals lose electrons ---> cations
•   nonmetals gain electrons ---> anions

-2 -1
+1
+ 3
-/+ 4 -3
+2
By losing or gaining e-, atom has same number of e-’s as nearest Group 8A atom.

•   C atom with 6 protons and 6 neutrons is the mass standard
•   = 12 atomic mass units
•   Mass Number (A)
•   =(# protons) + (# neutrons) A
•   NOT on the periodic table … (that is the
AVERAGE atomic mass on the table)
Z
•   Ex: A boron atom can have
A = 5 p + 5 n = 10 amu
10
B
5

Atomic Math
On periodic table- but not all PTs look exactly like this set up, but they have the same information

•   John Dalton stipulated that all atoms of a particular element were identical
•   Their atomic numbers were the same, and also their #’s of neutrons 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:
•   91% of the atoms had one mass
•   The remaining atoms were 10% heavier
•   All of the atoms had 10 protons, however some had more neutrons

•   atoms with the same number of protons (Z) but a different number of neutrons
•   same element, different atomic mass number (A)

1H (hydrogen):
2H (Deuterium):
3H (Tritium):
A=1 Z=1
A=2 Z=1
A=3 Z=1

Bone scans with radioactive technetium-99.

The tritium 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.

Which of the following represent isotopes of the same element? Which element?
234
X
234
X
235
X
238
92 93 92 92
X

Which of the following represent isotopes of the same element? Which element? The red ones are isotopes of Uranium
234
X
234
X
235
X
238
92 93 92 92
X

•   Atomic number (Z)
•   the number of protons in the nucleus
•   gives the element’s identity
•   (Atomic) Mass Number (A)
•   sum of the protons and neutrons for a given isotope of an element
•   Atomic Mass (also called Atomic Weight)
•   Weighted average mass of the atoms (accounts for all the isotopes) is average atomic mass

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 +
19
16 O -2
8
41 Ca +2
20
#p + ______ ______ _______
#n o ______ ______ _______
#e ______ ______ _______

Learning Check – Counting
State the number of protons, neutrons, and electrons in each of these ions.
39 K +
19
16 O -2
8
41 Ca +2
20
#p + 19 8 20
#n o 20 8 21
#e 18 10 18

Learning Check – Counting
Naturally occurring carbon consists of three isotopes,
12 C, 13 C, and 14 C. State the number of protons, neutrons, and electrons in each of these carbon atoms.
12 C
6
13 C
6
14 C
6
#p + _______ _______ _______
#n o _______ _______ _______
#e _______ _______ _______

12 C
6
13 C
6
#p + 6 6
#n o 6
#e 6
7
6
14 C
6
6
8
6

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) 34 X 2) 34 X 3) 36 X
16 14 14

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) 34 X 2) 34 X 3) 36 X
16 14 14

l   Nuclide: atomic species determined by nuclear contents l   Show the name of the element, a hyphen, and the mass number in hyphen notation sodium-23 l   Show the mass number and atomic number in nuclear symbol from mass number
23 Na
atomic number 11

+
Cl

Na +
less

O 2–
more

Write the nuclear symbol form for the following atoms or ions:
A. 8 p + , 8 n, 8 e - ___________
B. 17p + , 20n, 17e - ___________
C. 47p + , 60 n, 46 e ___________

Solution
A. 8 p + , 8 n, 8 e -
16
O
8
B. 17p + , 20 n, 17e -
37
Cl
17
C. 47p + , 60 n, 46 e -
107
Ag +
47

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. 15 X 15 X
8 7
B. 12 X
6
14 X
6
C. 15 X
7
16 X
8

Solution
B. 12 X
6
14 X
6
Both nuclear symbols represent isotopes of carbon with six protons each, but one has 6 neutrons and the other has 8.
C. 15 X
7
16
8
X
An atom of nitrogen (7) and an atom of oxygen (8) each have 8 neutrons.

•   We are used to calculating #’s of p + , n o and e using whole numbers; however on the
Periodic Table we often see a decimal number à Why?
•   Atomic Mass (on the Periodic Table)
•   The average of the isotopic masses, weighted according to the naturally occurring abundances of the isotopes of the element
•   In a weighted average we must assign greater importance – give greater weight – to the quantity that occurs more frequently

•   The atomic mass for each element on the periodic table reflects the relative abundance of each isotope in nature.
•   The mass on the periodic table is NOT the atomic mass number (A)

•   Atomic mass unit (amu) is the unit for relative atomic masses of the elements
•   1 amu =1/12 the mass of C-12 isotope.
•   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 (n o )
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
/100
•   To calculate your overall average, we use a weighted average instead of a simple average since different tasks are worth more
•   For example:
(30/100 x 80)
+ (30/100 x 75)
+ (10/100 x 70)
+ (30/100 x 70)
= 74.5%
Exams 30
Course work
30
Applied
Science
10
Final 30
Your mark
80%
75%
70%
70%

•   You add up (fractional abundance X mass) for each isotope to get the weighted average
•   Fractional abundance = natural abundance/100
•   Ex: If something has 3 isotopes:
(fractional abundance) isotope 1
X (mass) isotope 1
+
(fractional abundance) isotope 2
X (mass) isotope 2
+
(fractional abundance) isotope 3
X (mass) isotope 3
= average atomic mass

•   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
+
(.6917) x (62.93)
(.3083) x (64.93)
= 63.55 amu

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

Answer to Learning Check
(.00337) x (35.97)
+ (.00063) x (37.96)
+ (.996) x (39.96)
= 39.95amu

•   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

•   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
•   (.9051x 20) + (.0027x21) + (.0922x22)= 20.19 amu

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

10 B
11 B
•   Boron is 20% 10 B and 80% 11 B. That is, 11 B is 80 percent abundant on earth.
•   For boron, atomic weight=
= 0.20 (10 amu) + 0.80 (11 amu) = 10.8 amu

•   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?

Answer Check Part 1
•   if 2 isotopes, then the total is 100%. assume one is x%
(x), the other is automatically 100-x%, (1-x)
•   x(34.97) + (1-x)(36.97) = 35.453

Answer Check Part 2
•   x(34.97) + (1-x)(36.97)=35.453
•   Solve for x
•   34.97x+36.97-36.97x=35.453
•   -2x+36.97=35.453
•   -2x=-1.517
•   x=.7585
•   1-x=.2415

Answer Check Part 3
•   Therefore Cl-35 has a % abundance of 75.85% and Cl-37 has a % abundance of 24.15%

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

Answer Check
•   The atomic mass of nitrogen is 14.00674amu
•   14.00674 = p(14.003074) + (1 -p)(15.000108)
14.00674 = 14.003074p + 15.000108 - 15.000108p
0.997034p = -0.993368
•   p = 0.9963 = 99.63% (N14)
1 - p = 0.0037 = 0.37% (N15)