Download Mass Number, A

Atomic History and
Structure:
Thales of Miletus (600BC)
•  Noticed what we call static electricity with
amber
•  Things would be attracted to it when
rubbed
•  It was a “magical property”
• The term electron comes from the
Greek word for amber: “elektron”
Kanada (~600-501BC)
•  Indian attributed with first proposing the
idea of atoms (called “parmanu” or “anu”)
•  5 elements
• 
• 
• 
• 
• 
Earth
Fire
Water
Air
Ether
•  Atoms were indestructable and eternal
Empedocles (450BC)
•  4 elements:
•  Earth
•  Wind
•  Fire
•  Water
• Everything was different combinations
of these
•  This idea didn’t really change until1661!
Leucippus (~490 BC)
• Proposed the idea
of atoms
• That two things
exist
• Atoms
• Empty space.
Democritus (420BC)
• 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”
• “Atomos”
From :
http://www.historyworld.net/wrldhis/
PlainTextHistories.asp?
historyid=ac20#ixzz1UvX6le4i
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
Benjamin Franklin (1752)
§  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
§  Words he gave us:
§  battery, conductor, condenser, charge,
discharge, uncharged, negative, minus, plus,
electric shock, and electrician.
J.L. Proust (1794*)
•  Law of constant composition:
• 
A given compound always contains the
same elements in the same proportion
•  In other words…a given compound
always has the same composition,
regardless of where it comes from.
• Ex: H2O is always 89% oxygen and
11% H by mass
*not published or recognized until 1811
Dalton’s Atomic Theory ~1800
• 
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
Dalton’s Atomic Symbols
Dalton’s Atomic Theory
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
Michael Faraday (1832)
§  atoms contain particles with an electric
charge
§  structure of atoms related to electricity
§ The electron was the fundamental
particle of electricity
JJ Berzelius (1779-1848)
•  Came up with how we write chemical
formulas
•  Symbols for elements
•  Subscripts 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
(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
Hantaro Nagaoka (1904)
•  Proposed the planetary
(Saturnian) model of the atom
•  Positive, massive nucleus
•  Electrons bound to the nucleus
via gravity in charged rings
•  Both were confirmed by Rutherford
•  He abandoned the model in 1908
due to errors that were not
confirmed by new studies (charged
rings)
Rutherford’s Gold Foil Experiment
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
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
X-rays give some electrons a charge.
Millikan’s Experiment - 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 par7cles (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 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.
Which brings us to the
modern day view of the
atom….
ATOMIC
STRUCTURE
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.
ATOMIC COMPOSITION
•  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 (no)
• 
• 
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
•  Spin, charm
•  Top, bottom
H
•  3 quarks = 1 proton or 1 neutron
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 number of
protons in the nucleus, Z
13
Al
26.981
Atomic number
Atom symbol
AVERAGE Atomic Mass
Atoms are neutral because the
numbers of protons and electrons are
equal - 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 # 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
IONS
!
•  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 INa
Ca
I
O-2
O
compared to
PREDICTING ION CHARGES
In general
•  metals lose electrons ---> cations
•  nonmetals gain electrons ---> anions
Charges on Common Ions
-/+
+1
3
+
4 -3 -2 -1
+2
By losing or gaining e-, atom has same
number of e-’s as nearest Group 8A atom.
Mass Number, A
•  C atom with 6 protons and 6 neutrons is the
mass standard
•  = 12 atomic mass units
•  Mass Number (A)
•  =(# protons) + (# neutrons)
•  NOT on the periodic table…(that is the
AVERAGE atomic mass on the table)
•  Ex: A boron atom can have
A = 5 p + 5 n = 10 amu
A
10
Z
5
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
•  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
Isotopes
• 
atoms with the same number of protons (Z) but a
different number of neutrons
•  same element, different atomic mass number (A)
1H (hydrogen):
A=1
Z=1
2H (Deuterium):
A=2
Z=1
3H (Tritium):
A=3
Z=1
Isotopes &
Their Uses
Bone scans with
radioactive
technetium-99.
Isotopes & Their Uses
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.
Learning Check
Which of the following represent
isotopes of the same element? Which
element?
234
92
X
234
93
X
235
92
X
238
92
X
Learning Check
Which of the following represent
isotopes of the same element? Which
element? The red ones are isotopes
of Uranium
234
92
X
234
93
X
235
92
X
238
92
X
Atomic Math
•  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
16O -2
8
41Ca +2
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
6
13C
6
14C
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
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
atomic number
23 Na
11
Nuclide notation: p+, charge, and
average atomic mass
Mass number
(protons + neutrons)
37
Atomic number
17
(number of protons)
number of neutrons
A-Z =20
As atoms have no charge, the number
of electrons is the same as the number
of protons. This atom has 17 electrons.
Cl
Nuclide notation – ions
Mass number
Atomic number
23
+
11Na
number of neutrons=
1+ charge means 1 electron less than
the number of protons. This atom
has 10 electrons.
Nuclide notation –ions
Mass number
16
(protons + neutrons)
Atomic number
8
(number of protons)
number of neutrons=
2– charge means 2 electrons more
than the number of protons. This
atom has 10 electrons.
2–
O
Learning Check
Write the nuclear symbol form for the
following atoms or ions:
A. 8 p+, 8 n, 8 eB. 17p+, 20n, 17eC. 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
8
B.
C.
12X
15X
7
14X
6
6
15X
16X
7
8
Isotopes and Average Atomic Mass
•  We are used to calculating #’s of p+, no 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
Isotopes and Atomic Mass
•  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)
AMUs and Atomic Weight
• 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 (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:
(30/100 x 80)
+ (30/100 x 75)
+ (10/100 x 70)
+ (30/100 x 70)
= 74.5%
/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 (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
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
(.6917) x (62.93)
+ (.3083) x (64.93)
= 63.55 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=
= 0.20 (10 amu) + 0.80 (11 amu) = 10.8 amu
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