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Modern Chemistry
Chapter 3
Atoms:
the building block of matter
1
Chapter Vocabulary
Law of conservation
of mass
Law of definite
proportions
Law of multiple
proportions
Atom
Nuclear forces
Atomic number
Isotope
Mass number
nuclide
Atomic mass unit
Average atomic
mass
Mole
Avogadro’s number
Molar mass
2
Atomic Number
• Atoms are composed of identical
protons, neutrons, and electrons
– How then are atoms of one element
different from another element?
• Elements are different because they
contain different numbers of PROTONS
• The “atomic number” of an element is
the number of protons in the nucleus
• # protons in an atom = # electrons
Composition of The Atomic Nucleus
• Nuclei contain protons and neutrons
• The nuclei is neutral because the number of
protons equal the number of electrons
• Each element has a different number of
protons in their nucleus
– The number of protons determines the
atom’s identity
• Nuclear forces hold protons & neutrons
together
Chapter 3 Section 2 The Structure
of the Atom pages 72-76
4
p. 76
Properties of Subatomic Particles
Chapter 3 Section 2 The Structure
of the Atom pages 72-76
5
Discovery of the Atomic Nucleus
Relative size of the nucleus
Chapter 3 Section 2 The Structure
of the Atom pages 72-76
6
Isotopes
• Atoms of the same element that have
different masses
• Isotopes of hydrogen
– Protium 1p+ 0n0
– Deuterium 1p+ 1n0
– Tritium 1p+ 2n0
• Isotopes do not differ significantly in
their chemical behavior
Chapter 3 Section 3 Counting
Atoms pages 77-87
7
Mass Numbers
• Mass numbers = # of p+ + # of n0
of a specific isotope
• Examples
– Protium 1p+ + 0n0 = 1
– Deuterium 1p+ + 1n0 = 2
– Tritium 1p+ + 2n0 = 3
Chapter 3 Section 3 Counting
Atoms pages 77-87
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Designating Isotopes
• Hyphen notation
– name of element – mass number
– Hydrogen – 3
• Nuclear symbol
mass number
atomic number
Chapter 3 Section 3 Counting
Atoms pages 77-87
9
Number of neutrons in an atom
neutrons = mass number – atomic number
How many p+, e- and n0 are there in an
atom of chlorine-37?
17 p+ 17e- 20n0 (37-17)
Nuclide – a general term for a specific
isotope of an element
Chapter 3 Section 3 Counting
Atoms pages 77-87
10
Atoms:
From Philosophical Idea
to
Scientific Theory
Chapter 3 Section 1 Atoms: Ideas
to Theory pages 67-71
11
Foundation of Chemical Atomic Theory
• Law of Conservation of Mass
– Mass is neither created or destroyed
during ordinary chemical reactions or
physical changes
Chapter 3 Section 1 Atoms: Ideas
to Theory pages 67-71
12
p. 69*
Law of Conservation of Mass Image
Chapter 3 Section 1 Atoms: Ideas
to Theory pages 67-71
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p. 69*
Law of Conservation of Mass Image
Chapter 3 Section 1 Atoms: Ideas
to Theory pages 67-71
14
The atom – Philosophical Idea to Scientific Theory
• The Greek philosopher Democritus (460
B.C. – 370 B.C.) was among the first to
suggest the existence of atoms (from
the Greek word “atomos”)
– He believed that atoms were indivisible and
indestructible
– His ideas did agree with later scientific
theory, but did not explain chemical
behavior, and was not based on the
scientific method – but just philosophy
Dalton’s Atomic Theory (experiment based!)
John Dalton
(1766 – 1844)
1) All elements are composed of
tiny indivisible particles called
atoms
2) Atoms of the same element are
identical. Atoms of any one
element are different from
those of any other element.
3) Atoms of different elements combine in
simple whole-number ratios to form
chemical compounds
4) In chemical reactions, atoms are combined,
separated, or rearranged – but never
changed into atoms of another element.
Sizing up the Atom
 Elements are able to be subdivided into
smaller and smaller particles – these are
the atoms, and they still have properties
of that element
If you could line up 100,000,000
copper atoms in a single file, they
would be approximately 1 cm long
Despite their small size, individual
atoms are observable with instruments
such as scanning tunneling (electron)
microscopes
Structure of the Nuclear Atom
• One change to Dalton’s atomic
theory is that atoms are divisible
into subatomic particles:
–Electrons, protons, and neutrons are
examples of these fundamental
particles
–There are many other types of
particles, but we will study these three
Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray
tube to deduce the presence of a negatively
charged particle: the electron
Mass of the Electron
Mass of the
electron is
9.11 x 10-28 g
The oil drop apparatus
1916 – Robert Millikan determines the mass
of the electron: 1/1840 the mass of a
hydrogen atom; has one unit of negative
charge
Conclusions from the Study of
the Electron:
a) Cathode rays have identical properties
regardless of the element used to
produce them. All elements must contain
identically charged electrons.
b) Atoms are neutral, so there must be
positive particles in the atom to balance
the negative charge of the electrons
c) Electrons have so little mass that atoms
must contain other particles that account
for most of the mass
Thomson’s Atomic Model
J. J. Thomson
Thomson believed that the electrons
were like plums embedded in a
positively charged “pudding,” thus it
was called the “plum pudding” model.
Ernest Rutherford’s
Gold Foil Experiment - 1911
Alpha particles are helium nuclei The alpha particles were fired at a thin
sheet of gold foil
 Particles that hit on the detecting
screen (film) are recorded

p. 75
Gold Foil Experiment Image
Chapter 3 Section 2 The Structure
of the Atom pages 72-76
24
Relative Atomic Mass
• One atom, carbon-12, is set as a
standard
• All masses are expressed in relation to
this standard
• 1 atomic mass unit = 1/12 the mass of
a carbon-12 atom
Chapter 3 Section 3 Counting
Atoms pages 77-87
25
Relative Atomic Mass
• Examples
– Hydrogen – 1 = 1.007825 amu
– Oxygen – 16 = 15.994915 amu
– Magnesium – 24 = 23.985042 amu
• p+ = 1.007276 amu, n0 = 1.008665
amu, e- = 0.0005486 amu
• Relative mass and mass number are
close in value but not the same
Chapter 3 Section 3 Counting
Atoms pages 77-87
26
Average Atomic Mass
• The weighted average of the atomic
masses of the naturally occurring
isotopes of an element
• Example
– Copper
Cu-63: .6915 x 62.93 amu = 43.52
Cu-65: .3085 x 64.93 amu = 20.03
63.55 amu
Chapter 3 Section 3 Counting
Atoms pages 77-87
27
The Mole
• An amount of a substance that contains
as many particles as there are atoms in
exactly 12 g carbon-12.
• Similar to a dozen or a pair or a gross
• 6.022 x 1023 carbon-12 atoms = 12
grams of carbon-12
• Avogadro’s number = 6.022 x 1023
particles
Chapter 3 Section 3 Counting
Atoms pages 77-87
28
Molar mass
• The mass of one mole of a pure
substance
• Unit = g/mol
Chapter 3 Section 3 Counting
Atoms pages 77-87
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Gram-Mole Conversions
• The conversion factor for gram-mole
conversion is molar mass.
g
OR
mol
mol
g
• What is the mass, in grams, of 3.50
moles of Cu?
– 222 grams Cu
Chapter 3 Section 3 Counting
Atoms pages 77-87
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Conversions Image
p. 84
Chapter 3 Section 3 Counting
Atoms pages 77-87
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