Download CP-Chem Ch 3 PowerPoint(Atomic Theory

CHAPTER 3
Atoms
&
Atomic Theory
Brief History of Atomic Theory
• Particle Theory (400 BC)
– A Greek philosopher named Democritus
proposed that matter was composed of atoms
– Democritus coined the term atom which in
Greek meant “indivisible”
• Chemical Reactions (1700’s)
– Scientists determined that a chemical reaction
was the change of one substance into an
entirely new substance
Laws of Matter
• Conservation of Mass
– Matter can not be created or destroyed in a chemical
reaction
• Law of Definite Proportions
– A chemical compound contains the same elements in
the same proportions regardless of the size of the
sample
• Ex: Salt is NaCl,
• it’s composition is always 39.4 % Na & 60.6% Cl
• Law of Multiple Proportions
– Atoms combine in whole # ratios
• Ex: CO2, H2O, C6H12O6
Conservation of Mass
Dalton’s Atomic Theory
• In 1808 an English school teacher proposed an
atomic theory that he created using the laws of
matter and previously known atomic theory
• 1) All matter is composed of atoms
• 2) All atoms of a given element are identical in size,
mass, and other properties
• 3) Atoms can not be divided, created or destroyed
• 4) Atoms of different elements combine in simple whole
# ratios to form chemical compounds
• 5) Chemical reactions cause atoms to combine,
separate, and rearrange
Ex: Fe + O2  Fe2O3
(iron) (oxygen) (iron oxide aka rust)
Modern Atomic Theory
Not all aspects of Dalton’s atomic
theory have proven to be correct.
We now know that:
1) Atoms are divisible into even smaller particles
(subatomic particles)
2) A given element can have atoms with different masses
(isotopes)
Atomic Structure
• Atoms- atoms are the smallest particles of an
element that retain the properties of the element
• Subatomic Particles:
– Protons: positively charged(+), located in the nucleus
of an atom
– Neutrons: neutral particles(o), located in the nucleus of
an atom
– Electrons: negatively charged(-), located in the
electron cloud region surrounding the nucleus
* Photons- particles of light
* Alpha Particles- 2 fused protons
Properties of Subatomic
Particles
Discovery of The Electron
Discovery of the Nucleus
• A scientist named Ernest Rutherford shot
Alpha-Particles at a thin sheet of gold foil
Rutherford’s Experiment
• 99.9% of the alpha-particles went directly
through the gold foil
• It was previously believed that atoms were
solid dense spheres
• The experiment had 3 main conclusions:
– 1) Atoms are made mostly of empty space
– 2) The Nucleus is positively charged
– 3) The Nucleus is Very Small and Very Dense
Gold Foil Experiment
Nuclear Forces
• Nuclear forces hold the positively(+)
charged protons together in the nucleus
• Atomic #- the number of protons in the nucleus in
the atom & the # of electrons surrounding the
nucleus
• Mass #- The total number of protons + neutrons in
the nucleus of an atom
* the mass # that appears on the periodic table is an average of all
masses of all isotopes of a given element
• Isotopes- atoms of the same element with
different numbers of neutrons
– Nuclide: an isotope of an element
Isotopes and Nuclides
Atomic Mass
• The mass of an element
• Atomic masses are based on the Carbon-12 atom
as a standard reference
• Average Atomic Mass- the weighted average
mass of all isotopes of an atom
– Ex: Copper(Cu) has 2 Isotopes: Cu-63 & Cu-65
70% of all copper is Cu-63, 30% is Cu-65
Average Mass = 63.55 amu
Calculating Average Atomic
Masses
To find the average atomic mass you multiply
the % abundance of each isotope by the mass
of each isotope and take the sum of all masses
• Cu-63 = 63 amu x .70 = 44.10 amu
• Cu-65 = 65 amu x .30 = 19.45 amu
+ ------63.55 amu
Section 3.2 Review
1) Why was it important for the Cathode Ray to
pass through a vacuum tube?
2) Compare the 3 sub-atomic particles in terms of
location in the atom, mass, and charge. (you
may use a diagram in your response)
3) Describe one conclusion made by each of the
following scientists that led to the
development of the current atomic theory:
a) Dalton
b) Thompson
c) Rutherford
Atomic Mass & # of Particles
• Moles- moles are a counting unit (dozen = 12)
a mole = 6.02 x 1023
Avogadro's #: The number of particles in one mole,
6.02 x 1023
Molar Mass- the mass of one mole of a substance
(numerically equivalent to atomic mass)
Amedeo Avogadro
“The Man, The Myth, The Legend”
***The Mole Road***
multiply (x)
Mass(g)
Molar
Mass(g/mol)
multiply (x)
Moles
divide (/)
6.02 x
1023
# of
Atoms
divide (/)
• To convert from Moles  Mass
– Multiply the moles by the molar mass of the compound
• To convert from Moles  # of Atoms
– Multiply the moles by 6.02 x 1023
Mole Road Sample
Calculations
END OF
CHAPTER 3
NOTES!!!