Download CHAPTER 3 - THE ATOM

CHAPTER 3 – THE ATOM
Read pgs. 107 - 110
I.
History
Democritus – Greek philosopher, 400 B.C., said all matter is made up
of small, indivisible particles he called “atoms” (Greek for
“indivisible”). He wasn’t believed until the late 1700’s
John Dalton – 1803, realized that the Law of Conservation of Mass
and the Law of Definite composition could only be explained if atoms
existed. Wrote Dalton’s Atomic Theory, which was mostly right.
1. Matter is composed extremely small particles called atoms
2. Atoms are indivisible and indestructible
3. Atoms of a given element are identical in size, mass, and chemical
properties.
4. Different atoms combine in simple whole- number ratios to form
Compounds
5. In a chemical reaction, atoms are separated, combined or
rearranged.
Atom – the smallest particle of an element that can exist, either alone
or in combination with other atoms.
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As new info came out, Dalton’s theory was modified into the:
MODERN ATOMIC THEORY
1. All matter is made of atoms.
2. All atoms of a given element have the same chemical properties,
which differ from the chemical properties of other elements.
3. Atoms of a given element may not all have the same mass, but a
sample of the element will have an average mass which is
characteristic of that element.
4. Atoms are not subdivided during physical or chemical rxns.
5. Compounds are formed when 2 or more elements combine; each
atom loses its characteristic properties as a result.
II.
The structure of the Atom
A. Discovery of the electron
--by Thomson in 1897 (pgs. 109 – 110 for background info)
--Robert Millikan – “oil-drop experiment” in 1909 found that
the mass of an e- is very small compared to the mass of the
whole atom. Also discovered that e-‘s have a measureable
amount of negative electric charge.
Q. How could he have measured the mass of an e- when even today we
don’t have a scale that will measure that?
[Hint: How can you find the thickness of a page? (use a ruler)]
--Ernest Rutherford – in 1908- 1909 bombarded thin metal foils
with positively-charged particles. (see pg. 111 - 112 – know his
exp.!)
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Rutherford’s conclusion: The positively-charged particles that
bounced back were repelled by a powerful positive charge which
must occupy a very small area within the atom, since most of the
particles passed through without being affected. (This also
explains the deflected particles.)
He named this area within an atom the nucleus – very small and
extremely dense. Occupied by positively charged particles he
called protons. (He concluded that most of the rest of the atom
was empty space.)
C. Protons – every neutral atom has the same number of protons as
electrons, so the positive and negative charges will balance.
D. Neutrons – discovered in 1932 by Chadwick. Are neutral (no
electrical charge). Mass is virtually identical to the mass of a
proton; both of these particles have a much larger mass than an
electron has (see chart, pg. 111)
E. Fission – in 1939, Lise Meitner and Otto Hahn discovered that
an atom could be subdivided. They called this process “fission”.
(Six years later we dropped a fission bomb on Hiroshima, Japan).
III.
Weighing Atoms
A. Atomic Number (Z) – the number of protons in the nucleus of a
given element’s atoms.
Each element has a different # of protons in their nucleus. That
is what makes elements different from each other.
Elements are placed in order in the Per. Table from 1 118 by
their atomic number.
So: Atomic number of oxygen = __________
# of p in an oxygen atom = __________
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# of e- in an oxygen atom = __________
Atomic number of hydrogen = __________
# of p in a neon atom = __________
# of e- in an iodine atom = __________
B. Mass Number – tells the total number of protons plus neutrons
in the nucleus (doesn’t count e-‘s since they virtually have no
mass).
To find the mass number of an element, round the “atomic
mass” (the decimal number in the P.T.) to the nearest whole
number.
Ex:
Na atomic mass 22.98977, mass number = _________
K atomic mass = __________, mass number = _______
To find number of neutrons in an atom:
# of n = mass# - atomic#
Element
Na
Atomic #
Atomic mass
Mass#
#p
#n
#e
U
Pb
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Fe
Ag
P
F
Ba
Al
Li
C. Isotopes: Atoms of the same element, which have different masses
because they have different numbers of neutrons in the nucleus.
Ex: Most common isotope of H has: 1 proton
0 neutrons
mass number = 1
Other isotopes of H: “Deuterium” 1 proton
1 neutron
mass number = 2
“Tritium”
1 proton
2 neutron
mass number = 3
Two methods of designating isotopes:
1. Hyphen notation
“Name-mass #” ex:
Hydrogen-1
Hydrogen-2
Hydrogen-3
(normal)
(deuterium)
(tritium)
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Ex: the isotope of uranium used in nuclear power plants is
Uranium-235.
Mass # = __________
# of p = ___________
# of n = ___________
2. Nuclear Symbol Method
Mass # 3
H  symbol
(this is Hydrogen-3)
Atomic # 1
Uranium-235 is shown as
Write the nuclear symbol for Chlorine-35 __________
Write the nuclear symbol for Lead-206 ____________
How many protons, neutrons, and electrons are in one atom of
Chlorine-37?
# of p = __________
(atomic number, Z)
# of e = __________
(atomic number, Z)
# of n = __________ (mass # - atomic #)
Practice:
1. How many p, n, and e are in an atom of Bromine-80?
_____ p, _____ n, _____e
2. Write the nuclear symbol for Carbon-13. __________
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3. Write the hyphen-notation for the element that contains 15 p and
15 n. _______________________________________
D. Relative Atomic Mass
The atomic mass of an element = the mass of 1 atom of that
element. Since these atomic masses are so small (tiny fractions
of a gram) they have a designated a new unit:
“Atomic Mass Unit” (u)
1 u = 1/12th the mass of a Carbon-12 atom (everything is compared to
Carbon-12)
So, the mass of 1 atom of carbon-12 is 12 u.
The atomic masses given on the Per. Table are averages for all
isotopes of that element combined, given in “u”.
Mass of 1 atom of He = ____________
Mass of 1 atom of H = ___________
Atomic mass of Mg = ___________
IV.
Counting Atoms
A. The Mole – a counting unit (abbrev. “mol”)
“One Dozen” = 12 of anything
“One Mole” = 6.022 x 1023 of anything
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Ex: one mole of cars = 6.022x1023 cars
one mole of students = 6.022x1023 students
one mole of copper atoms = 6.022x1023 atoms
This number (6.022x1023) is called Avogadro’s Number
How large is it? If every person on earth (5 billion +) counting 1 atom per
second from a 1 mole sample of Cu, it would take 4 million years to count the
atoms.
B. Molar Mass – the mass, in grams, of 1 mole of a pure substance.
Conveniently, this equals the atomic mass from the Per. Table,
with grams as the unit.
Ex: Mass of 1 atom of H = 1.00794 u
Mass of 1 mole of H atoms = 1.00794 g
Mass of 1 mole of He = _______________
Mass of 1 mole of Na = ____________________
Mass of 1 mole of Cu = ____________________
B. Calculating with Moles
**Use factor-label method
**when using molar mass of Avogadro’s Number as a conversion factor, round
it to the sig figs allowed in the problem before using it.
**Moles should end up in every problem.
Ex: What is the mass, in grams, of 3.50 moles of Cu?
3.50 mol Cu x _________g Cu = ________________
mol Cu
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Ex: What is the mass, in grams, of 0.375 moles of K?
Ex: You have 11.9 g of Al. How many moles of Al do you have?
Ex: How many moles of Au are in 3.60 x 10-10 g Au?
Ex: How many moles of Ag are in 3.01 x 1023 atoms of Ag?
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Ex: 2500 atoms Sn = __?__ mol Sn?
Ex: How many atoms of Al are in 2.75 mol Al?
Ex: How many atoms of Zn are in 4.93 mol of Zn?
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Ex: What is the mass in grams of 7.5 x 1015 atoms of Ni?
Ex: How many atoms of sulfur are contained in 4.00 g of sulfur?
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