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Chapter 3

Democritus (400 B.C.) suggested matter was
made of tiny particles which he called Atoms
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Law of Conservation of Mass
o Mass is neither created nor destroyed in
chemical or physical changes
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Law of Definite Proportions
o A compound always contains the same
elements in the exactly the same proportions

Law of Multiple Proportions
o If the same two elements form two different
compounds, the masses of the second element
that combine with a certain mass of the first
element is always a ratio of small whole
numbers
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Matter is composed of small particles called
Atoms
Atoms of the same element are the same;
atoms of different elements are different
Atoms cannot be created, divided, or
destroyed
Atoms of different elements combine to form
compounds
Atoms are rearranged in chemical reactions
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Atoms are made up of smaller particles
Atoms of the same element have different
numbers of neutrons (different masses)
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Atom
o
Smallest part of an element that has the properties
of that element
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Joseph John Thomson (1897) showed that a
cathode ray consists of negative particles
called electrons
Robert Millikan (1909) confirmed that
electrons have a negative charge and
determined the mass of an electron
Since atoms are neutral, it was determined
that atoms have positive particles called
protons

Electron
o Negative Charge (-)
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Proton
o Positive Charge (+)
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Ernest Rutherford (1911) discovered the
nucleus when he bombarded a piece of gold
foil with alpha particles.
The nucleus contains most of the mass in the
atom but occupies only a very small amount
of space in the atom.
Contains Protons (+) and neutrons (0), which
are neutral.
The number of protons determines the
atom’s identity.
Particle
Electron
Symbol
e-
Charge
-
Location
Surrounds
Nucleus
Proton
p
+
Nucleus
Neutron
n
0
Nucleus
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Atomic Number = # of Protons

# of Protons = # of Electrons

Mass = # of Protons + # of Neutrons
Atomic
Number
Mass
Number
7 14
9
19
19 39
27 59
6 14
Protons
Neutrons
Electrons
7
7
9 10
19 20
27 32
6
8
7
9
19
27
6
Symbol
N
F
K
Co
C
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Isotopes
o Atoms of the same element that have different
numbers of neutrons
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Isotopes have different masses but react the
same way
Mass number
Atomic number
Or it can be written
Uranium - 235
Isotopes of Carbon
Carbon - 12
Carbon - 14
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Abbreviated to amu
Defined as exactly 1/12 of the mass of
carbon – 12
Unit used to measure the mass of atoms
Mass (amu)
Proton 1 (approx.)
Neutron 1 (approx.)
Electron 0 (actually 1/1836 amu)
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Atomic Mass
o The weighted average of the masses of the isotopes
of that element
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Most elements occur as two or more isotopes
Atomic mass is found on the periodic table
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Calculating Atomic Mass
If the element consists of Isotope A and
Isotope B
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An element has two isotopes. The isotope
with a mass of 10.012 amu has a relative
abundance of 19.91%. The isotope with a
mass of 11.009 has a relative abundance of
80.09% Calculate the atomic mass and name
the element.
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An Element has two isotopes. One has a
relative abundance of 69.2% and a mass of
62.93. The other isotope has a relative
abundance of 30.8% and a mass of 64.93.
Calculate the atomic mass
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SI unit which measures the amount of a
substance.
1 mol = 6.022 x 1023 particles
Most elements consist of particles called
atoms
6.022 x 1023 Avogadro’s number
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How many moles are in 1.20 x 1025
phosphorus atom?
How many atoms are in 0.750 moles Zn?
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Molar Mass
o Mass of the one mole of substance
Atomic Mass
Molar Mass
(1 Atom)
(6.02 x 1023 atoms)
Carbon
12.01 amu
12.01 g
Copper
63.55 amu
63.55 g
Magnesium
24.31 amu
24.31 g
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How many moles are in 336 grams of carbon?
How many grams are in 2.38 moles of
magnesium?
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How many atoms are in 12 grams of sodium?
How many grams are in 9.6 x 1023 atoms of
potassium?