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Ch. 4: Atoms/Atomic
Theory
Atoms
• Definition
- the smallest particle that has the
properties of an element, basic unit of
matter
• 119 distinct atoms as of 1999, form
elements
Atomic Symbols
- each element has its own name,
accompanied by a symbol
- usually one/two letters (first one is
always capitalized)
ex. Iron: Fe
- Fe represents 1 atom of iron,
- 2Fe represents 2 atoms of iron etc…
- can also be written as Fe2
History

Democritus
 named the most basic
particle
 atom- means “indivisible”

Aristotle
 didn’t believe in atoms
 thought matter was
continuous
History
 by


1700s, all chemists agreed
on the existence of atoms
that atoms combined to make
compounds
 Still
did not agree on whether
elements combined in the same ratio
when making a compound
Law of Conservation of Mass
 mass
is neither created or destroyed
during regular chemical or physical
changes
Law of Definite Proportions
 any
amount of a compound contains
the same element in the same
proportions by mass
No matter
where the
copper
carbonate is
used, it still
has the same
composition
Law of Multiple Proportions
applies when 2 or more elements combine
to make more than one type of compound
 the mass ratios of the second element
simplify to small whole numbers

Dalton’s Atomic Theory
1.
2.
3.
4.
5.
All mass is made of atoms
Atoms of same element have the same
size, mass, and properties
Atoms can’t be subdivided, created or
destroyed
Atoms of different elements combine in
whole number ratios to make compounds
In chemical reactions, atoms are
combined, separated, and rearranged.
Modern Atomic Theory

Some parts of Dalton’s theory were wrong:



atoms are divisible into smaller particles
(subatomic particles)
atoms of the same element can have different
masses (isotopes)
Most important parts of atomic theory:


all matter is made of atoms
atoms of different elements have different
properties
Structure of Atom

Nucleus:



contains protons
and neutrons
takes up very little
space
Electron Cloud:


contains electrons
takes up most of
space
Subatomic Particles
 includes



all particles inside atom
proton
electron
neutron
 charge
on protons and electrons are
equal but opposite
 to make an atom neutral, need equal
numbers of protons and electrons
Subatomic Particles
 number
of protons identifies the
atom as a certain element
 protons and neutrons are about
same size
 electrons are much smaller
Comparing Subatomic
Particles
Discovery of Electron
resulted from scientists passing electric
current through gases to test conductivity
 used cathode-ray tubes
 noticed that when current was passed
through a glow (or “ray”) was produced

Discovery of
Electron
Noted Qualities of Ray Produced:
1. existed- there was a shadow on the
glass when an object was placed
inside
2. had mass- the paddle wheel placed
inside, moved from one end to the
other so something must have been
“pushing” it
Discovery of Electron
3.
4.
negatively charged- the rays behaved the
same way around a magnetic field as a
conducting wire
negatively charged- were repelled by a
negatively charged object
Discovery of Electron

All of these led scientists to
believe there were
negatively charged particles
inside the cathode ray
Discovery of Electron
 J.J.
Thomson (English 1897) did more
experiments to actually make the
discovery
 he found ratio of charge of this
particle to this mass of the particle
 since the ratio stayed constant for any
metal that contained it, it must be the
same in all of the metals
Are electrons the only particles?
 since
atoms are neutral, something
must balance the negative charge
 since an atom’s mass is so much
larger than the mass of its electrons,
there must be other matter inside an
atom
Discovery of Nucleus
Rutherford discovered the nucleus by
shooting alpha particles (have positive
charge) at a very thin piece of gold foil
 he predicted that the particles would go
right through the foil at some small angle

Discovery of Nucleus
Discovery of Nucleus
 some
particles (1/8000) bounced back
from the foil
 this meant there must be a “powerful
force” in the foil to hit particle back
Predicted Results
Actual Results
Discovery of Nucleus
Characteristics of
“Powerful Force”:
1. dense- since it was strong
enough to deflect particle
2. small- only 1/8000 hit the
force dead on and bounced
back
3. positively charged- since
there was a repulsion
between force and alpha
particles
Atomic Math and
Isotopes
Atomic Number
number of protons is the atomic #
 It is the identity of an element.
 All atoms of the same element have the same
atomic number
 located above the symbol in the periodic table
 order of the elements in the periodic table

Isotopes
 atoms
of the same element with
different numbers of neutrons
 most elements exist as a mixture
of isotopes
B. Isotopes
© Addison-Wesley Publishing Company, Inc.
 What
do the Carbon isotopes below
have in common? What is different
about them?
Mass Number For Isotopes
sum of particles in nucleus
 Mass number for isotope = #p + #n
Hydrogen isotopes have special names:




protium
deuterium
tritium
Designating Isotopes
 Hyphen


notation:
Name - mass number
ex. Carbon – 13
 Nuclear
pn
p
Symbol notation:
Symbol
Ex : 136C
Examples
1.
7 protons, 8 neutrons
Nitrogen-15
2.
15
7
N
17 electrons, 19 neutrons
Chlorine- 36
36
17
Cl
Examples
3.
Z=5, 6 neutrons
Boron- 11
3.
11
5
B
A=75, 42 neutrons
Arsenic- 75
75
33
As
Ch. 4 Atoms
Average Atomic Mass
Review
Subatomic Particles
ATOM
NUCLEUS
ELECTRONS
PROTONS
NEUTRONS
POSITIVE
CHARGE
NEUTRAL
CHARGE
NEGATIVE CHARGE
equal in a
Atomic
Most Number
of the atom’s mass.
neutral atom
equals the # of...
Relative Atomic Mass
 since
masses of atoms are so small, it
is more convenient to use relative
atomic masses instead of real masses
 to set up a scale, we have to pick one
atom to be the standard
 since 1961, the carbon-12 nuclide is
the standard and is assigned a mass of
exactly 12 amu
Relative Atomic Mass
 atomic
mass unit (amu)- one is exactly
1/12th of the mass of a carbon-12
atom
 mass
of proton= 1.007276 amu
 mass of neutron= 1.008665 amu
 mass of electron= 0.0005486 amu
Average Atomic Mass
 weighted
relative atomic masses of the
isotopes of each element
 each isotope has a known natural
occurrence (percentage of that
elements’ atoms)
Calculating Average Atomic Mass
 Naturally


occurring copper consists of:
69.71% copper-63 (62.929598 amu)
30.83% copper-65 (64.927793 amu)
(0.6971 x 62.929598)+(0.3083 x 64.927793)
=63.88 amu
Calculating Average Atomic Mass
 An
element has three main isotopes
with the following percent occurances:



#1: 19.99244 amu, 90.51%
#2: 20.99395 amu, 0.27%
#3: 21.99138 amu, 9.22%
 Find
the average atomic mass and
determine the element.
Calculating Average Atomic
Mass
(19.99244x90.51)  (20.99395x0.27)  (21.99138  9.22)
100
 20.17945amu
Neon