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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 Theory1. 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 reaction2. 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 atomsYES!
4. atoms combine in certain whole-number ratiosYES! 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
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• 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.
Gold Foil Animation
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 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 (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 todetermine the charge on the electron independently of itsmass. Then using Thompson’s charge-to-mass ratio, hewas able to calculate the mass of the electron.
e = 1.602 10 x 10-19 coulombe/m = 1.7588 x 108 coulomb/gramm = 9.1091 x 10-28 gram
Goldstein - Conducted “positive” ray experiments thatlead to the identification of the proton. The chargewas found to be identical to that of the electron andthe 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 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 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
• 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.
The atom is mostlyempty space
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
• 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
+
–
• 11 electrons• 11 negative charges
• 11 positive charges• 11 protons
Atoms are neutral because the numbers of protons and electrons are equal - the opposite charges cancel.
IonsA 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+ chargeIf 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 I- O-2
Na Ca I O compared to
PREDICTING ION CHARGES
In general
• metals lose electrons ---> cations
• nonmetals gain electrons ---> anions
Charges on Common Ions-1-2-3
+1
+2
By losing or gaining e-, atom has same number of e-’s as nearest Group 8A atom.
-/+4+3
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
Z
10
5B
A
Z
10
5B
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-to-
charge 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 X 234
X235
X238
X
92 93 92 92
Learning Check
Which of the following represent isotopes of the same element? Which element? The red ones are isotopes of Uranium
234 X 234
X235
X238
X
92 93 92 92
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+ 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 is1) 14 2) 16 3) 34
C. The element is1) Si 2) Ca 3) Se
D. Another isotope of this element is1) 34X 2) 34X 3) 36X
16 14 14
Atomic Symbols: Nuclide Notation
Nuclide: atomic species determined by nuclear
contents
Show the name of the element, a hyphen, and
the mass number in hyphen notation
sodium-23
Show the mass number and atomic number in
nuclear symbol frommass number 23 Na
atomic number 11
Nuclide notation: p+, charge, and average atomic mass
37
Mass number (protons + neutrons)
Cl17Atomic number (number of protons)
A-Z =20number of neutrons
As atoms have no charge, the number of electrons is the same as the number of protons. This atom has 17 electrons.
Nuclide notation – ions
23Mass number Na+
11Atomic number
number of neutrons=
1+ charge means 1 electron less than the number of protons. This atom has 10 electrons.
Nuclide notation –ions
16Mass number (protons + neutrons) O2–
8Atomic number (number of protons)
number of neutrons= 2– charge means 2 electrons more than the number of protons. This atom has 10 electrons.
Learning Check
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- ___________
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 7
B. 12X 14X 6 6
C. 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 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 grelative 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 Science
10 70%
Final 30 70%
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 • 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
10B
11B
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
