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Early Atomic Structure
Dalton through Rutherford
Chemistry I
Owen J. Roberts High School
Daryl Walmer
A Prelude to Atomic Theory
• Atomic theory was based on three scientific
laws which were developed at the end of the
1700’s.
– Law of Conservation of Matter
– Law of Definite Proportion
– Law of Multiple Proportions
Existence of Atoms Inferred
• These were all mathematical laws, and
therefore the existence of matter as
randomized chunks combined in any old
manner did not make sense mathematically.
• These laws made sense (mathematically) if
matter was thought be composed of smaller
pieces that united in specific ways.
• John Dalton inferred the existence of atomsalthough he never saw one.
Dalton’s Atomic Theory 1803
• All matter is composed of atoms, that are
indivisible and indestructible.
• Atoms of different elements are different
• Compounds result from atoms combining in
specific simple whole number ratios.
• During a chemical reactions atoms are not
created or destroyed they only change bonding
partners.
• Atoms of a given element are the same in terms
of structure and fundamental properties.
Dalton’s Model
• Scientific Model is a visual representation to
aid in explaining a theory, law or hypothesis.
– Lock and key model for enzymes
• Dalton envisioned atoms as small, tiny, solid
spheres.
• Billiard Ball Model- Dalton’s Model
Atoms
• All elements are represented and composed
of atoms.
• The atoms that are currently present in the
universe have been here since the beginning
of the universe.
• The number of atoms of a given element are
for the most part fixed.
• Atoms become recycled in the process of life
and death.
Electricity gets attention of Scientists
• As the 1800’s progressed- scientists became
fascinated with electricity.
• Numerous experiments using electricity were
conducted.
• Cathode ray tube was device used to study
electricity
Simple Cathode Ray Tube
• http://www.chem.uiuc.edu/clcwebsite/cathod
e.html
What occurs in CRT
• Tube is near vacuum – means very little gas
• Cathode produces negative charge.
• Gas particles bump into cathode and pick up
negative charge
• Negative charge particles drawn toward
positive anode
• Opposite charges attract.
Cathode Ray is beam of Particles
• Maltese Cross
• Paddlewheel
Cathode Ray is Negative
• Opposite Charges attract
• Charge plates and magnets- bend beam
Electron Discovered- 1896
• J. J Thomson
• With cathode ray tube
• Experiments done using different metals for
cathode and anode.
• Experiments were done using different gases
in the tube.
• Why did Thomson change the electrode and
the gas?
Thomson determines
Charge to Mass Ratio on the Electron
• Measured charge to mass ratio for a stream of
electrons using a cathode ray tube apparatus
• This ratio is 1.76 x 108 coulombs/grams
• A coulomb, C, unit of electrical charge
equivalent to 6.24 x 1018 charges
Millikan’s Oil Drop Experiment
Millikan’s Oil Drop Experiment
• Result of oil drop experiment
– Charge on an electron 1.60 x 10-19 coulombs of
negative charge
• From this a calculation is performed to
determine mass on electron
•
1.60 𝑥 10−19 𝐶
x
1𝑔
1.76 𝑥 108 𝐶
= 9.10 x 10-28 g
• Current accepted 9.10939 x 10-28 g
Discovery of Proton
• Since atoms were neutral- once negative
particle discovered must be positive particle
• Positive particle harder to discover
• Positive particle much harder to make– not
many made- beam of positive charge very
faint hard to see
• Formation of positive particle
Canal Ray Used to See Proton
• Modification made to cathode ray tube to find
positive particle.
• Fritted disk cathode
Canal Ray tube Up Close
Discovery of Proton
• Goldstein designed canal ray tube
• Wien discovered charge to mass ratio on
proton
• Many scientists were involved in proton work,
Ernest Rutherford is generally credited with
naming the particle.
E/m ratio on Proton
• Charge to mass ratio proton = 9.62 x 104 C/g
• The charge on a proton is 1.60 x 10-19 C
• How can one account for the difference in the
charge to mass ratio of an electron and a
proton?
• (remember e/m electron 1.76 x 108 C/g)
Compare and Contrast
Plum Pudding Model of Atom
•
•
•
•
Description
Around 1900
J.J. Thomson
How does this
contradict Dalton’s
model?
• Chocolate chip cookie
Rutherford’s Gold Foil Experiment
• Around 1910
• See independent reading on Rutherford’s Gold
foil Experiment
• Complete the Cornell Sheet based on the
independent reading.
Neutron
• James Chadwick -1932
• Bombarded Beryllium foil with alpha particles
Protons, Neutrons and Electrons in
Atoms
• The number of protons determines the identity
of an element.
– Each element has unique number of protons
• Magnesium 12 protons, aluminum 13 protons
– Atoms of different elements must have different
number of protons
• Atomic Number (Z) number of protons in an
atom.
• How many protons are in:
– Na
Ba
Zn
F
Kr
Mass Number
• Protons and neutrons are roughly equal in
mass.
• Atomic mass scale assigns
– Proton a mass of 1 amu
– Neutron a mass of 1 amu
• Mass number is the number of the nucleons(electrons not factored into an atom’s massWhy?)
• Mass number (A) = neutrons + protons
Practice
• Quick Notes
•
•
•
•
•
•
Atomic # is Z
Mass # is A
# protons = Z
# neutrons = A – Z
# electrons (neutral atom) = Z
A = protons + neutrons
• Find number of p+,
n and e• Cl mass # 35
• X having an atomic #
of 20 and a mass #
of 41
• Pb mass # 209
• If Z = 28 and A = 57
Isotopes
• Isotopes are atoms of the same element which differ in mass
– This mass difference is due to possession of differing
numbers of neutrons
– Nuclear charge cannot change – what if it does?
• Isotopes of an element have different mass numbers
– Mass number is equal to the sum of major nucleons
• A nucleon refers to any particle in the nucleus of an atom
• Some isotopes are stable while others are radioactive
(unstable)
– This stability may be predicted at times by looking at the
proton to neutron ratio or possible one of the “magic
numbers”, or other trends
Isotopes
• Naming isotopes
– The term nuclide is used to refer to a specific atom of an
isotope
– Element name - mass number
– Carbon -12 and carbon- 14 two different isotopes or nuclides of
carbon
• Determine number of protons, neutrons, and electrons in each of
these carbon nuclides.
• The isotope composition of an element is the same no matter
where that element is found. Some isotopes tend to exist in higher
percentages than others.
• For example:
– Chlorine has two common isotopes
• Chlorine – 35 and chlorine – 37
• 75.77% of chlorine is chlorine-35
• Therefore what is the percentage of chlorine-37?
Isotopes of Hydrogen and Isotopic
Notation
• Isotope notation has
the general form
• Isotopes (nuclides ) of
hydrogen:
Practice
• Write the isotope
notation for the
following :
• Atom with 12 protons,
and 13 neutrons
• Atom with 83 protons
and 210 neutrons
• Atom with 34 protons
and 79 neutrons
• Determine # of p+, n
and e•
•
•
40
K
19
235
U
92
20
Ne
10
Atomic Mass
• Mass Number represents the number of nucleons in
one nuclide (atom) of an element or one kind of
isotope of an element.
• In nature, all of the isotopes
– exist mixed together with one another
– Have the same chemical properties
– Very expensive to isolate
• Atomic mass represents the average mass (weighted)
of the mass numbers of all the isotopes• Atomic mass would be the mass of the element as it
exists when extracted from the earth or the seas.
Average Atomic Mass
The average atomic mass may be determined by calculation from any analysis
yielding the actual masses of the isotopes and some relative proportion of
their presence; an actual number of nuclides or as percentages of nuclides. A
rough estimate may be obtained if the mass numbers are utilized
e.g. Two isotopes of chlorine are found to occur naturally; chlorine – 37 (37Cl)
with a mass of 36.9659 amu and representing 24.23% of all chlorine and
chlorine – 35 (35Cl) with a mass of 34.9689 amu and representing 75.77% of
all chlorine. What is chlorine’s average atomic mass?
Ave. At. Mass = [(% x isotope mass) + (% x isotope mass) + …..]
Total %
Ave. mass Cl = (.2423 x 36.9659 u) + (.7577 x 34.9689 u)
0.2423 + 0.7577
OR (24.23% x 36.9659 u) + (75.77 % x 34.9689 u)
24.23 % + 75.77 %
Mass Spectrograph