Atomic Theory

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Atomic Theory

• The History of the Atom

2
State Standard

• SC3. Students will use the modern atomic theory
  to explain the characteristics of atoms.
• a. Discriminate between the relative size,
  charge, and position of protons, neutrons, and
  electrons in the atom.
• c. Explain the relationship of the proton number
  to the elements identity.
• d. Explain the relationship of isotopes to the
  relative abundance of atoms of a particular
  element.

3
What do I need to know?

• How did the following people impact the
  development of the atomic model?
• Democritus
• Aristotle
• John Dalton
• JJ Thomson
• Ernest Rutherford
• Robert Millikan
• James Chadwick
• Sir William Crookes
• Niels Bohr

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Democritus

• A contemporary of Aristotle around 400 BC.
• The world was made up of two things
• 1. empty space
• 2. tiny particles called atomos.

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Aristotle

• Matter was continuous. This continuous substance
  was called hyle.
• Believed matter made up of
• Earth, air, water and fire

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• Because of Aristotles reputation, his opinion
  was accepted for the next 2000 years up until the
  17th century.

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Isaac Newton and Robert Boyle

• Lived during the late1600s.
• Spoke out in support of Democritus.
• No experimental evidence of this belief.

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John Dalton

• English schoolmaster and chemist
• Lived in the 1800s
• Called The Father of the Modern Atomic Theory
  because he offered the first rational atomic
  theory based on the scientific evidence of the
  time.

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John Daltons Atomic Theory

• Supported Democritus idea of the atom.
• Was the first atomic theory based on scientific
  evidence.
• Combined ideas of elements with that of atoms
• Published in 1808

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This theory prevailed for another 80 years.
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Scientific Evidence available to Dalton 1

• Antoine Lavoisier
• French Chemist
• Total mass before chemical reaction was the same
  as the total mass after chemical reaction.
• Mass of reactants Mass of products

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Law of Conservation of Mass

• Antoine Lavoisier
• For a closed system, the total mass before
  chemical reaction is equal to the total mass
  after chemical reaction.
• Total mass of the universe is a constant
• Matter can be changed in many ways, but it cannot
  be created or destroyed (for non nuclear
  processes)

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Scientific Evidence Available to Dalton 2

• Joseph Proust
• French Chemist
• Observed that specific substances always contain
  elements in the same ratio by mass.

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Law of Definite Proportions (3)

• Each compound has a specific ratio of elements
• It is a ratio by mass
• Water is always 8 grams of oxygen for each gram
  of hydrogen

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Law of Definite Proportions

• Joseph Proust
• Substances always contain elements in the same
  ratio by mass.
• Example For NaCl, for a 100.0 g sample, it will
  always have 39.3 g Na and 60.7 g Cl
• i.e 39.3 Na and 60.7 Cl

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John Dalton

• Armed with this experimental work, John Dalton
  proposed his atomic theory
• He only had 2 pieces of experimental work when he
  made his proposal although other evidence shortly
  followed (within 3-4 years)
• He kept Democritus idea of particulate matter

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Daltons Atomic Theory (1808)

• An element is composed of extremely small,
  indivisible particles called Atoms
• All of the atoms of a given element have
  identical properties, which differ from the
  properties of the other elements.
• Atoms cannot be created, destroyed, or
  transformed into atoms of another element.
  Experiment 1? Law of Conservation of Mass
• Compounds are formed when atoms of different
  elements combine with each other in small
  whole-number ratios.
• ? Law of Multiple Proportions (Dalton proposal
  not based on experiment but proven later to be
  true by Amadeo Avagadro and J.L. Gay-Lussac)
• The relative numbers and kinds of atoms are
  constant in a given compound. Experiment 2 ? Law
  of Definite Proportions

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Law of Multiple Proportions

• CO or CO2, but not CO1.8

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Law of Multiple Proportions

• if two elements form more than one compound, the
  ratio of the second element that combines with 1
  gram of the first element in each is a simple
  whole number.

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What?

• Water is 8 grams of oxygen per gram of hydrogen.
• Hydrogen Peroxide is 16 grams of oxygen per gram
  of hydrogen.
• 16 to 8 is a 2 to 1 ratio
• True because you have to add a whole atom, you
  cant add a piece of an atom.

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Evidence for Law of Multiple Proportions

• 1811
• Two other scientists studying gaseous reactions
  at constant T and P.
• J.L. Gay-Lussac French Chemist
• Experiment 3 Observed that under constant
  conditions, the ratio of volumes of reacting
  gases and gaseous products were in small whole
  number ratios ? Law of Combining Volumes
• Amadeo Avagadro Italian Chemist
• Experiment 4 Observed that if have a gas at same
  T P, equal volumes of gases have the same
  number of molecules
• ? Avagadros number (mole)

22
How does this confirm LMP?

• Consider this reaction
• Based on Law of Combining Volumes
• 1 vol. Cl 1 vol. H ? 2 vol. HCl
• Avagadro said equal volumes of gases contain the
  same number of particles.
• Does anyone see the problem?

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Confirming LMP

• 1 vol. Cl 1 vol. H ? 2 vol. HCl
• Since equal volumes have same number of
  particles
• 1 Cl 1 H ? 2 HCl
• Now do you see the problem?

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Confirming LMP

• The only way to resolve all of the experimental
  evidence into one coherent understanding was to
  have diatomic particles of hydrogen and chlorine
• 1 Cl2 1 H2 ? 2 HCl
• i.e. the smallest individual particles that made
  up hydrogen and chlorine existed in small
  whole-numbered ratios.

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Subatomic Theory in the 1800s

• Sir William Crookes
• JJ Thomson
• Millikan
• Rutherford

26
State Standard

• SC3. Students will use the modern atomic theory
  to explain the characteristics of atoms.
• a. Discriminate between the relative size,
  charge, and position of protons, neutrons, and
  electrons in the atom.
• c. Explain the relationship of the proton number
  to the elements identity.
• d. Explain the relationship of isotopes to the
  relative abundance of atoms of a particular
  element.

27
Summary of Daltons Atomic Theory

• It was based on experiment
• It explained the concepts of atoms, elements, and
  compounds
• It explained why there were differences b/w the
  properties of the elements.
• Included the three foundational laws of chemistry
  (LCM, LDP, LMP)

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Discovery of the Electron

• 1897
• It required to two separate experiments by two
  different individuals
• JJ Thomson and Robert Millikan

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JJ Thomson and the Cathode Ray Tube (CRT)1897

• JJ Thomson used the CRT to measure the
  charge-to-mass ratio (charge mass) of a kg of
  electrons.
• He won the 1906 Nobel Prize in Physics for this
  work.

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Daltons Atomic Theory

• All matter is made of tiny indivisible particles
  called atoms.
• Atoms of the same element are identical, those of
  different elements are different.
• Atoms of different elements combine in whole
  number ratios to form compounds
• In Chemical reactions, atoms are rearranged,
  combined or separated. No new atoms are created
  or destroyed.

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J. J. Thomson - English physicist 1897

• Used the cathode ray tube
• It is a vacuum tube - all the air has been pumped
  out.
• Couldnt measure the charge directly so he
  determined the ratio of charge to mass.

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Thomsons Experiment
CATHODE
ANODE

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Vacuum tube
Metal Disks
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Thomsons Experiment

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34
Thomsons Experiment

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Thomsons Experiment

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Thomsons Experiment

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• Passing an electric current makes a beam appear
  to move from the negative to the positive end

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Thomsons Experiment

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• Passing an electric current makes a beam appear
  to move from the negative to the positive end

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Thomsons Experiment

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• Passing an electric current makes a beam appear
  to move from the negative to the positive end

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Thomsons Experiment

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• Passing an electric current makes a beam appear
  to move from the negative to the positive end

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Thomsons Experiment

• By adding an electric field

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Thomsons Experiment

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• By adding an electric field

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Thomsons Experiment

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• By adding an electric field

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Thomsons Experiment

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• By adding an electric field

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Thomsons Experiment

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• By adding an electric field

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Thomsons Experiment

-

• By adding an electric field he found that the
  moving pieces were negative

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CRT Expt

• Next, he placed an object in the path of the CR.

Part of the beam was blocked and
Cathode (-)
Anode ()
A shadow was created on the anode.
This proved conclusively that the CR was made of
particles.
47
CRT Expt

• Finally he placed a paddle wheel in the path of
  the CRthe paddle wheel spun.

Cathode (-)
Anode ()
This proved that the particles that made up the
CR had mass.
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CRT Expt Summary

• Thomson showed
• The CR was charged
• The CR had a negative charge
• The CR was made of particles
• The CR particles had mass.

The deflection of the cathode rays gave
evidence for the negatively charged nature of
electrons.
What does it really look like?
49
CRT Expt Summary

• Thomson won the Nobel prize because he measured
  the charge mass ratio of a kg of electrons.
• He measured this to be 1.759 X 1011 C/kg and the
  charge was negative (electrons are negatively
  charged).

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Contributions
PROBLEM?
WHO?
DID WHAT?
Proposed idea of Atoms

• Democritus

What holds atoms together?
Discovered cathode ray (charged particles)
Sir William Crookes
Unable to determine exact charge.
J.J. Thomson
Discovered Electron charge to mass ratio
51
Millikan Oil Drop Experiment

• Conducted by Robert Millikan
• It measured the charge on a single electron.
• He won the 1932 Nobel Prize in Physics for this
  measurement

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HOW DOES IT WORK?
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Millikan Oil Drop Expt
Charged Oil drops
Oil spray
Adjustable voltage
Negatively charged oil drop suspended in air.
microscope
Negatively charged plate
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Millikan Oil Drop Expt

• Once a negatively charged oil drop dropped
  through the hole in the top plate, the voltage
  was switched on and the drop was suspended in mid
  air.
• Once the drop was suspended, the microscope was
  used to measure the size of the oil drop and the
  voltage differences were used to measure the
  charge on the oil drop.

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Millikan Oil Drop Expt

• Millikan was able to measure the charge of a
  single electron with this expt.
• This turned out to be 1.6 X 10-19 C / e-

57
Millikan / Thomson Result

• With the results of these 2 experiments, the mass
  of the electron could then be calculated.
• Thomson 1.759 X 1011 C / kg
• Millikan 1.6 X 10-19 C / e-
• Today, the electron is the standard unit of
  negative charge

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Thomsons Model / Plum Pudding

• Found the electron
• Couldnt find positive Said the atom was like
  plum pudding
• Positive negative charges evenly spread through
  out, with the electrons able to be removed

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Rutherfords experiment(Aware of Thomsons plum
pudding

• Used radioactivity
• Alpha particles - positively charged pieces given
  off by uranium
• Shot them at gold foil which can be made a few
  atoms thick

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• The modern view of the atom was developed by
  Ernest Rutherford (1871-1937).

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He Expected

• The alpha particles to pass through without
  changing direction very much
• Because
• The positive charges were spread out evenly.
  Alone they were not enough to stop the alpha
  particles

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Results of foil experiment if Plum Pudding model
had been correct.
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What he expected
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He thought the mass was evenly distributed in the
atom
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Because, he thought the mass was evenly
distributed in the atom
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What he got
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How he explained it

• Atom is mostly empty
• Small dense, positive center
• Alpha particles are deflected by it if they get
  close enough

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What Actually Happened
71
Rutherfords experiment.
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Plum Pudding Model must be IncorrectYOU MUST
KNOW THISWRITE THE FOLLOWING DOWN

• Concluded
• Atom is mostly empty space
• Center (nucleus) Positive particles and most of
  the atoms mass will be found here.
• Nucleus is DENSE

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Structure of the Atom

• There are two regions
• The nucleus
• With protons and neutrons
• Positive charge
• Almost all the mass
• Electron cloud- Most of the volume of an atom
• The region where the electron can be found

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Modern View

• The atom is mostly empty space
• Two regions
• Nucleus- protons and neutrons
• Electron cloud- region where you might find an
  electron

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Subatomic particles
Actual mass (g)
Relative mass
Name
Symbol
Charge
Electron
e-
-1
1/1840
9.11 x 10-28
Proton
p
1
1
1.67 x 10-24
Neutron
n0
0
1
1.67 x 10-24
76
Contributions
WHO?
DID WHAT?

• Millikan

• Accurately determined
• Charge (-1)
• Mass of electron (1/1840 the mass of Hydrogen
  atom)

Problem couldnt account for all of the mass.
Rutherford

• Concluded
• Discovered Proton
• Nucleus is positive
• Nucleus is dense
• Atom mostly empty space

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Contributions
WHO?
DID WHAT?

• Neils Bohr

Electrons orbit nucleus like planets planetary
model Electrons travel in fixed orbits
James Chadwick (Rutherfords lab assistant)
Discovered Neutron click here for video on
discovery
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Atomic Number

• Moseley found that different elements have
  different numbers of protons.
• In fact, he discovered that elements had
  incrementally different numbers of protons
• Called this number the Atomic Number

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Discovery of the Proton

• Atomic Number tells us
• Number of protons in an atom
• Identity of the element (Au, Ag, etc)

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Implications for Daltons Atomic Theory

• The discovery of the proton, electron, and the
  neutron was the first major contradiction to
  Daltons Atomic Theory
• Daltons Atomic Theory postulated that the atom
  was the smallest piece of matterwe now know this
  not to be true.

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Subatomic Particle Summary (pg 111)

• Particle Actual Charge Relative Charge
• Electron -1.9 X 10-19 C -1
• Proton 1.9 X 10-19 C 1
• Neutron 0 0

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Subatomic Particle Summary KNOW Relative Mass

• Particle Actual Mass Relative Mass
• Electron 9.10939 X 10-31 kg 0
• Proton 1.67262 X 10-27 kg 1
• Neutron 1.67493 X 10-27 kg 1

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Subatomic Particle Summary SO WHAT DO I NEED TO
KNOW?Test on Tue. (10/7)

• You need to remember
• --Who discovered what and how
• (Thomson Millikan, Crookes, Rutherford, Bohr
  Chadwick)
• --The following in this table
• Particle Relative Mass Relative Charge
• Electron 0 -1
• Proton 1 1
• Neutron 1 0

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Symbols

• Contain the symbol of the element, the mass
  number and the atomic number

Mass number
X
Atomic number
85
Mass Number

• Number of protons number of Neutrons

86
Atomic Number

• The number of protons in the nucleus of an atom.

87
Find the number of electrons and protons contain
in an atom of

• A. Arsenic
• B. Gold
• C. Fluorine
• D. Molybdenum
• E. Polonium
• F. Barium

33 p 33 e-
79 p 79 e-
9 p 9 e-
42 p 42 e-
84 p 84 e-
56 p 56 e-
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Identify the atom containing the following number

• A. 34
• B. 5
• C. 31
• D. 61
• E. 94

Selenium
Boron
Gallium
Promethium
Plutonium
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Identify the atom containing the following number
of protons

• A. 74
• B. 20
• C. 49
• D. 70
• E. 93

Tungsten
Calcium
Indium
Ytterbium
Neptunium
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Nuclear Symbol and PEN diagrams

• P protons
• E electrons
• N neutrons
• PEN diagram is nothing more than a diagram that
  tells how many protons, electron, and neutrons
  are in an atom.

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Nuclear Symbols and PEN diagrams

• Mass Number (A) protons neutrons
• Atomic Number (Z) protons
• Nuclear Symbol symbol that designates the
  specific atomic number and mass number of a
  nuclide.

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Nuclear Symbol and names

• 37Cl is the same as chlorine-37
• 15N is the same as nitrogen-15

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PEN Example

• Copper-65
• 65Cu P 29 from PT
• E 29 if neutral, same as Z
• N 36 A - Z
• A Z neutrons

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Your turn class work complete for homework

• Practice Atomic Structure Worksheet1 and 2

95
Isotopes and Atomic Mass

• While scientists were discovering the subatomic
  particles, other scientists were involved with
  isolating and characterizing the naturally
  occurring elements.
• These scientists were having difficulty
  understanding their mass measurements especially
  considering the now widely accepted fact of
  subatomic particles

96
Isotopes and Atomic Mass

• Scientists determined that the atomic mass was
  actually a weighted average of the naturally
  occurring isotopes of an element.
• Isotopes are atoms of the same element that have
  different numbers of neutrons
• This gives atoms of the same element different
  masses and the measured atomic mass of an element
  is then a weighted average of the naturally
  occurring isotopes.

97
Isotopes and Atomic Mass

• Mass Number The number of protons and neutrons
  in an isotope of an atom
• Atomic Mass Weighted average of all of the
  naturally occurring isotopes of an element.

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Whats the standard for the mass of an atom?
99
The Carbon-12 atom was assigned a mass of exactly
12 atomic mass units (amu)
STANDARD IS THE CARBON 12 ATOM
100
ATOMIC MASS UNIT

• amu 1/12 of a carbon 12 atom

Although 1 amu is approximately equal to 1 proton
or 1 neutron - there are differences
101
Subatomic Particles Masses (p.119)KNOW

• Electron 0.000549 0 amu
• Proton 1.007276 1.00 amu
• Neutron 1.008665 1.00 amu

102
Isotopes and Atomic Mass

• What is a weighted average?
• Consider your grade
• 40 Daily 70
• 60 Tests 90
• Avg. 80?

103
Isotopes and Atomic Mass

• Not if its a weighted average. The test grade
  is weighted more heavily so the average should
  reflect this.
• 40 Daily 70 28
• 60 Tests 90 54
• Avg. 82

X .40
X .60
28 54 82
104
Atomic Mass

• Atomic mass is a weighted average.
• To calculate the atomic mass - you need the
  following to be given to you
• relative abundance of the isotopes
• mass of the isotope.

105
Step 1 Convert the relative abundance () into
decimal format.
Steps to determine atomic mass
Step 2 Multiply the decimal from step 1 by the
mass of that isotope.
Step 3 Do this for each isotope.
Step 4 Add the results of all isotope
calculations
This will provide you with the atomic mass.
Lets take a look
106
Atomic Mass Calculation

• Isotope Rel. Abundance Mass (amu)
• 16O 99.759 15.995 amu
• 17O 0.037 16.995 amu
• 18O 0.204 17.999 amu
• At. Mass 0.99759(15.995)
• 0.00037(16.995)
• 0.00204(17.999)
• 15.9995 16.00 amu

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Bellringer (10/6/08)

• Take out Atomic Mass Worksheet
• On a separate sheet of paper answer the
  following
• An ion always contains
• A. equal number of p and e-
• B. equal number of p and n0
• C. Unequal number of p and e-
• D. Unequal number of n0 and e-
• 2. What is the approximate mass of a neutron in
  amus?

108
Lets Review from FridayMeasuring Atomic Mass

• Unit is the Atomic Mass Unit (amu)
• One twelfth the mass of a carbon-12 atom.
• Each isotope has its own atomic mass we need the
  average from percent abundance.

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Naming Isotopes

• Put the mass number after the name of the element
• carbon- 12
• carbon -14
• uranium-235

110

• How do isotopes Hydrogen 1 and Hydrogen -2
  Differ ?

H 2 has one neutron and H 1 has no neutrons
111
Atomic Mass

• are the decimal numbers on the periodic table.
• Is not a whole number because it is an average.

112
Atomic Mass

• Calculate the atomic mass of copper if copper has
  two isotopes. 69.1 has a mass of 62.93 amu and
  the rest has a mass of 64.93 amu.

(.691) (62.93 amu) 43.48 100 - 69.1 30.90
(.3090) (64.93 amu) 20.06 43.48 20.06
63.54 amu Check your work what is the atomic
mass for Copper?
113
Atomic Mass

• Magnesium has three isotopes. 78.99 magnesium 24
  with a mass of 23.9850 amu, 10.00 magnesium 25
  with a mass of 24.9858 amu, and the rest
  magnesium 25 with a mass of 25.9826 amu. What is
  the atomic mass of magnesium?
• If not told otherwise, the mass of the isotope is
  the mass number in amu

114
Need more Practice / information see p.
119-121See top of my web page for all downloads
for this unit.