Download Chemistry Topic III – The Atom

III – The Atom
Chemistry Student Notes
Chemistry Topic III – The Atom
PRE-TEST QUESTIONS
1. Convert 43 cm into meters.
2. Classify sugar water.
3. Define matter.
4. Define element.
5. What is the chemical symbol for silver?
Disclaimer: Please note that these early models we will be learning are now known to be wrong.
However, having said that, you do need to understand these early models of the atom to see how
it has evolved into our current understanding.
I. Early Concepts of the Atom
A. How divisible is matter?
1. Around 400 B.C., there was a debate in ancient Greece over what matter was made of
and how _________________ it was.
2. The Greek philosopher, Democritus (460 B.C. – 370 B.C.) believed that matter was made
of basic _________________ that were indivisible and indestructible.
a. He named these particles _________________ or _________________, from the
Greek meaning “uncuttable” or “indivisible.”
3. Aristotle (384 B.C. – 322 B.C.), another Greek philosopher, disagreed with Democritus
and believed matter was _________________ and infinitely _________________.
a. In fact, he proposed matter was made of four elements: earth, air, fire and water.
4. However, these were merely _________________, and neither philosopher explained
_________________ _________________.
a. As a result of not being able to provide ___________________ _________________
for either idea, neither idea ever truly took hold for about 2,000 years.
B. What did we know early on?
1. Before atomic theory was proposed, there were a few laws that scientists knew were true,
but could not explain.
2. The Law of Conservation of Mass states that for any chemical reaction, mass is neither
_________________ nor _________________.
3. The Law of Definite Proportions (aka Law of Constant Composition) states that a
chemical compound always contains the _________________ elements in exactly the
_________________ proportions by mass.
a. e.g., pure glucose sugar (C6H12O6), if broken into its constituent elements, is always
composed of _________________ carbon, _________________ oxygen and
_________________ hydrogen, by mass, of each.
b. e.g., table salt (NaCl) is always composed of _________________ by mass sodium
and _________________ by mass chlorine. For example…
i. 2000 tons of salt is made of _________________ tons of sodium and
_________________ tons of chlorine.
ii. 100 grams of salt is made of _________________ grams of sodium and
_________________ grams of chlorine.
iii. 1 mg of salt is made of _________________ mg of sodium and
_________________ mg of chlorine.
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TOPIC III Notes, Page 1
III – The Atom
Chemistry Student Notes
4. The Law of Multiple Proportions states that whenever two elements form more than
one compound, the different masses of one element that combine with the same mass of
the other element are in the ratio of small _________________ numbers.
a. e.g., Water’s formula is ______ and Hydrogen peroxide’s formula is __________.
b. Water is 2 hydrogen atoms for 1 oxygen atom.
c. Hydrogen peroxide is 2 hydrogen atoms for 2 oxygen atoms.
d. The ratio between the oxygen atoms is 2:1 (_________________ numbers)
C. What did Dalton think up?
Draw Dalton’s Atomic Model
1. John Dalton (1766 – 1844) was an English
chemist and schoolteacher.
2. He used _________________________ methods
to help explain the laws listed above to turn
Democritus’s idea of atoms into a scientific
theory.
a. Considering the laws above, the ratios in the
law of multiple proportions were always
_________________ numbers and the law of
definite proportions were always the same.
b. From this, he believed that matter was made
of some basic unit, since the ratios were
always WHOLE and the same.
3. Dalton’s Atomic Theory (c. 1803) was the first theory to relate chemical changes to
events at the atomic level.
a. All elements are composed of tiny indivisible particles called _________________.
b. Atoms of the same element are _________________. The atoms of any one element
are different from those of any other element.
c. Atoms cannot be _________________, _________________ or _________________
in physical or chemical reactions.
d. Atoms of different elements can physically mix together or can chemically combine
in simple whole-number ratios to form compounds.
e. Chemical reactions occur when atoms are ____________________,
____________________ or ______________________. Atoms of one element,
however, are never changed into atoms of another element as a result of a chemical
reaction.
4. Today, 200 years later, there are a few exceptions known. (Note: we will be learning
more about these exceptions later.)
a. _________________ are atoms of the same element but with a different mass
(violates #2 above)
b. There are parts to the atom like protons, neutrons, electrons, and even further into
quarks (violates #3 above). However, it remains true the atom is the smallest part of
an element that retains its properties.
c. We do know atoms can be transmuted into other elements via _________________
reactions (violates #5 above).
D. How big are atoms?
1. If you took a sample of an element and divided it in half, over and over again, you would
eventually get a piece that could not be divided anymore, but would still have the
properties of the element.
a. This is called the _________________.
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TOPIC III Notes, Page 2
III – The Atom
Chemistry Student Notes
2. Atoms are extremely small and cannot be seen by the unaided eye or even with the light
microscopes used in Biology to look at cells.
a. Despite this small size, individual atoms are observable with instruments such as
scanning tunneling microscopes (STM for short) [1981] that use an extremely small
needle to show the atom.
II. Refining the Atomic Model
A. Was the atom really the smallest form of matter?
1. Much of Dalton’s Atomic Theory of Matter is still accepted today, except we now know
about isotopes and that atoms are divisible.
2. There are three subatomic particles: _________________, _________________ and
_________________.
B. What were cathode ray tubes?
1. William Crookes constructed glass tubes, with most of the air vacuumed out. When an
electric current was passed through the tube, a beam originated from the
_________________ electrode, called the cathode.
a. These were called ______________ __________ (they originated from the cathode).
b. It was discovered that if you brought _________________ fields near the beam, it
would deflect.
c. If you brought a positively-charged magnet near it, it attracted the beam.
i. Thus, because _________________ _________________, we knew the beam was
composed of _________________-charged particles.
d. Scientists also knew that if you put a paddle-wheel in the tube, it moved when the
cathode ray struck it.
i. This showed the beam was made of _________________, not _______________.
C. What did cathode rays tell us?
1. English physicist, Sir Joseph John (“J. J.”) Thomson (1856-1940) experimented with the
cathode rays and discovered the electron in 1897.
2. In the cathode ray tube (CRT), most of the gas was vacuumed out so that it was under
very low pressure.
a. Metal electrodes were metal disks placed at each end of the glass tubes
b. Anodes are the _________________ charged disks.
c. Cathodes are the _________________ charged disks.
3. Draw this “default” setting:
4. Thomson noted that you could pass a _________________ _________________ field by
the cathode ray and get the same results as if you had passed a magnet by it.
5. The beam went towards the _________________ end or was repelled by the
_________________ end.
a. This meant the beam had a _________________ charge because
_________________ _________________.
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III – The Atom
Chemistry Student Notes
b. He called these particles “_________________,” but later scientists preferred the
term “_________________” suggested by George Stoney prior to Thomson’s
experiment.
c. Draw Thomson’s result:
6. Thomson did further experiments and found the rays were the _________________ no
matter what metal the cathode was made from or the gas inside the tube.
a. This proved that these negatively-charged particles were present in
_________________ atoms of _________________ elements.
b. They also had the _________________ mass-to-charge ratio.
7. Thomson constructed a cathode ray with a fluorescent screen at one end so he could
calculate the amount of deflection of the ray due to the strength of the electric and
magnetic fields.
a. He found the charge-to-mass ratio of the electron to be ____________________
coulombs per gram.
b. The coulomb (C) is the SI unit of electrical
Draw Thomson’s Atomic Model
charge.
8. Thomson proposed a model of the atom.
a. The atom was _________________ overall, so
Thomson knew if it contained these
_________________ electrons, there had to be a
_________________ part of the atom to balance
the charge.
b. He proposed the idea that the atom is a relatively
large positive sphere with these negative
electrons interspersed throughout it, like raisins
in a plum pudding, as he described it.
c. His model is called the plum-pudding model.
D. Robert Millikan (1868-1953) was an American physicist.
1. In 1909, he succeeded in measuring the charge of an electron in his famous “oil-drop”
experiment.
2. In a cast-iron pot, small drops of oil were _________________ with extra electrons.
3. These oil drops were then dropped between two electrically charged plates.
4. Millikan monitored the drops, measuring how the voltage on the plates affected their rate
of fall.
5. From this data, he calculated the charges on the drops. His experiments showed that the
charges were always integral multipliers of _____________________, which he deduced
was the charge of a single electron.
6. Using Thomson’s charge-to-mass ratio, he was able to calculate the mass…
© Hendley
TOPIC III Notes, Page 4
III – The Atom
III.
Chemistry Student Notes
The Nuclear Atom
A. Atomic models so far
1. Dalton proposed evidence for the atom and knew only that atoms were the basic particles
of matter.
a. He thought they looked like tiny indivisible spheres.
2. Thomson knew the atom was neutral overall, and that a part of the atom was negative, so
he proposed the plum-pudding model.
a. The atom is a positive sphere with negative electrons dispersed throughout it.
B. Rutherford
1. Lord Ernest Rutherford (1871-1937), a former student of Thomson, set out in 1911 to
evaluate the atomic model of the day: Thomson’s plum-pudding model.
2. Rutherford and his co-workers used heavy radioactive particles called
_________________ (α) particles.
a. _________________ particles are essentially helium atoms that have lost their two
electrons, such that they have a ______ charge.
3. The _________________ particles were shot through lead barriers to focus the beams to
a sheet of _________________ foil, which had been flattened so thin that it was only a
few atoms thick.
4. A specially coated film was placed around the gold foil so that it would light up when
struck by the alpha particles, to show where they had gone.
5. Rutherford’s hypothesis: The alpha particles should go _________________ through the
gold foil, and be deflected only _________________, if at _________________.
a. Draw the hypothetical result:
6. The Result: Rutherford found that most of the alpha particles _________________,
indeed, go _________________ through the gold foil.
a. However, he noted that a few alpha particles _________________ off at large (more
than 90 degrees) angles (1 in _________________)
b. Draw the result:
C. What did Rutherford conclude about the atom?
1. Most of the particles went straight through the gold foil.
a. This showed that most of the atom is _________________ _________________.
2. The alpha particles have a lot of mass, and they bounced off.
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III – The Atom
Chemistry Student Notes
a. This showed what they were bouncing off had a lot of _________________.
3. The alpha particles were positively charged and bounced off something.
a. Whatever it was they bounced off must also have been _________________-charged
to deflect them.
4. The alpha particles were very rarely (about 1 in
Draw Rutherford’s Atomic Model
80,000) deflected.
a. This meant whatever they were bouncing off was
very _________________.
D. What explained all this?
1. Rutherford concluded there was a part of the atom
that had most of the _________________ and all of
the _________________ charge concentrated in a
very small _________________.
2. He called this region the _________________.
a. The _________________ is the tiny central core
of the atom comprised of _________________
and _________________.
3. He proposed that the atom has a positive _________________ around which the
electrons circled.
a. The _________________ and _________________ are located in the nucleus.
b. The electrons are distributed around the nucleus and occupy almost all the
_________________ of the atom.
c. This model is still incomplete, but better.
E. What about protons and neutrons?
1. _________________ are positively-charged particles found in the nucleus of the atom.
a. This dates back to Eugen Goldstein (1850-1930) and the discovery of canal rays,
which were positively-charged rays that went in the opposite direction of cathode
rays.
b. Each proton has a mass about _________________ times that of an electron.
2. James Chadwick (1891-1974) was a student of Rutherford’s and discovered a particle
that had a _________________ charge.
a. He named this particle the _________________.
b. Neutrons have no charge, but their mass is nearly ____________ to that of the proton.
Particle
© Hendley
Symbol
Properties of Subatomic Particles
Relative
Location
Relative Mass
Charge
inside the
(mass of
atom
proton = 1)
Actual Mass
(g)
TOPIC III Notes, Page 6
III – The Atom
Chemistry Student Notes
IV. Distinguishing Among Atoms
A. Atomic Number
1. The atomic number of an element is
the number of _________________
in the nucleus of an atom of that
element.
2. This number is _________________
for each element.
3. On the Periodic Table, it is the
_________________ number and is
always a whole number.
a. Usually written at the top.
B. Mass Number
1. Most
of
an
atom’s
_________________ is found in the _________________ and _________________ that
comprise its nucleus.
2. The mass number is the total number of protons and neutrons in the atom’s nucleus.
a. It is also called the atomic _________________ and atomic _________________,
among other names.
3. On the Periodic Table, it is the larger number, usually expressed as a weighted average in
the form of a decimal, usually written at the bottom.
C. How can we determine the number of subatomic particles?
1. Each element is defined by the number of protons in its nucleus.
a. The number of protons is equal to the _________________ _________________.
2. Because all atoms are neutral (for now), the amount of negative charge must balance the
amount of positive charge in the atom.
a. The number of electrons = the number of _________________.
3. Most of the mass comes from the protons and neutrons and is found in the mass number.
a. Thus, the number of neutrons is the difference between the mass number and the
atomic number.
b. Number of neutrons = _________________ Number – _________________ Number
4. For example, Lithium has an atomic number of ___ and a mass number of ________
which rounds off to ____.
a. Lithium has ____ protons.
b. Lithium has ____ electrons.
c. Lithium has ____ neutrons.
EXAMPLE III-01: Determining Subatomic Particles
Element
Symbol
Atomic
Mass
Number
Number
Hydrogen
1
Boron
Protons
Neutrons
Electrons
5
Gallium
Gold
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III – The Atom
Chemistry Student Notes
D. Shorthand Notations
1. Nuclear Notation
a. The element’s symbol is written with the mass number on its left at the top and the
atomic number at the bottom.
b. Z – _________________ Number
c. A – _________________ Number (Rounded to a whole number)
2. Hyphen Notation is when the name of an element is written first, then hyphenated with
the mass number attached (rounded to a whole number).
a. The atomic number is not given because it is a constant, giving the identity of the
element.
b. For example, Oxygen with a mass of 16 is written ______________________.
c. This also helps distinguish between _________________ of elements.
EXAMPLE III-02: Shorthand Notations
Element
Symbol
Atomic
Number
Hydrogen
1
Boron
Mass
Number
Nuclear
Notation
Hyphen
Notation
5
Gallium
Gold
E. Isotopes
1. Isotopes are atoms that have the _________________ number of protons as other atoms
of the same element, but _________________ numbers of neutrons.
a. Because isotopes of an element have different numbers of neutrons, they also have
different _________________ numbers.
2. Despite the difference in mass, isotopes of the same element are chemically alike because
they have ________________ numbers of ________________ and ________________.
a. Preview: __________ are the subatomic particles responsible for chemical reactivity.
3. For example, there are three major isotopes of hydrogen. They all have 1 proton and 1
electron, but a different number of neutrons.
a. The most common isotope is called _________________, or generally, just
‘hydrogen.’ It has no neutrons and is written Hydrogen-1.
b. The second isotope has one neutron and a mass number of 2. It is called
_________________ and is written Hydrogen-2.
c. The third isotope has two neutrons and a mass number of 3. It is called
_________________ and is written Hydrogen-3.
F. Atomic Mass Units
1. The mass of protons and neutrons are very small.
a. In comparison, the mass of the electron is negligible.
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TOPIC III Notes, Page 8
III – The Atom
Chemistry Student Notes
2. Atomic masses in grams are too small to be practical, so relative masses are used instead.
a. Historically, hydrogen was used since it was the smallest atom known.
b. Since humans are carbon-based life-forms, we use _________________.
c. Carbon-12 is assigned to be exactly 12 atomic mass units (amu’s).
3. An atomic mass unit (amu) is defined as one-twelfth of the mass of a carbon-12 atom.
a. Carbon-12 has 6 protons and 6 neutrons in its nucleus. The electrons barely make a
difference, so each proton and neutron is roughly 1 amu.
G. Average Atomic Mass
1. Most elements in nature occur in mixed amounts of isotopes.
a. e.g., 99.985% of all hydrogen is hydrogen-1.
2. The average atomic mass depicted on the Periodic Table, then, is a _________________
average of all isotopes in nature.
a. To calculate the atomic mass of an element, multiply the mass of each isotope by its
natural abundance, expressed as a decimal, and then add the products.
3. Carbon has mostly two isotopes in nature.
a. Carbon-12 has a natural abundance of 98.89% while carbon-13 has a natural
abundance of 1.11%.
b. The mass of carbon-12 is 12.000 amu; the mass of carbon-13 is 13.003 amu.
c. Avg atomic mass =
EXAMPLE III-3: Average Atomic Mass of Chlorine
Chlorine has two isotopes: Chlorine-35 (mass of 34.969) which has a natural abundance of
75.4% and chlorine-37 (mass of 36.966) which has a natural abundance of 24.6%. Determine the
average atomic mass of chlorine.
V. The Periodic Table – A Preview
A. A Preview
1. The Periodic Table allows you to easily compare properties of one element (or group of
elements) to those of another element (or group of elements).
2. Each square in the Periodic Table generally shows the symbol, name, atomic number and
mass number of each element.
3. _________________ are the horizontal rows in the periodic table.
a. A period has either 2, 8, 18, or 32 elements in its period.
4. _________________, or _____________, are the vertical columns in the periodic table.
a. Elements are placed in the same family because they have similar properties, like
reactivity.
i. e.g., Group 11 (Copper, Silver, Gold) are called the Coinage metals.
b. There are ____ groups in the periodic table.
5. Elements on the _________________ side of the P.T. are _________________.
6. Elements on the _________________ side of the P.T. are _________________.
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TOPIC III Notes, Page 9