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Chapter #3
Chapter 3A
ATOMS:
The Building
Blocks of
Matter
West Valley
High School
General
Chemistry
• Atoms: The Building Blocks of
Mr.
Mata
Matter
Standard 2a
• To compare and contrast
the atomic number and
atomic mass.
Essential Question
• Describe how atomic
theory developed to give us
our modern view of the
atom.
3-1 Early Atomic Theory
• Atoms – smallest particle of matter.
• Democritus (400 B.C) stated world was
made of atoms. (atomos = “indivisible”).
• Antoine Lavoisier (1800’s) discovered
mass didn’t change after chemical rxn.
• Proposed “matter can be changed, but
it cannot be created or destroyed “
(Law of Conservation of Mass).
Dalton’s Atomic Theory
1. All matter composed of atoms.
2. Atoms of same element identical in size, mass,
properties; atoms of diff elements diff. in size,
mass, properties.
3. Atoms can’t be subdivided, created, destroyed.
4. Atoms of different elements combine in wholenumber ratios to form chemical compounds.
5. In chemical reactions, atoms are combined,
separated, or rearranged.
Modern Atomic Theory
• All matter is composed of atoms.
• Atoms of any one element differ in
properties from atoms of another element.
• Element’s average mass unique to element.
• Atoms cannot be subdivided, created, or
destroyed in ordinary chemical rxns.
• Changes CAN occur in nuclear rxns!
Section 3-2
• Atom- smallest particle of element
that retains chemical properties of
element.
• Nucleus- positively charged, dense
central portion of the atom; contains
nearly all mass (~ 99.7%).
Subatomic Particles
Electrons e-
Negative charged particles. Found
in electron clouds. Thomson (1897)
Protons
Positively charged particles. Found
in the nucleus. Rutherford (1918)
p+
Neutrons N
No charge. Found in the nucleus.
Chadwick (1932)
The Atomic Scale
• Most of the mass of the atom is in
the nucleus (protons and neutrons)
• Electrons are found outside of the
nucleus (the electron cloud).
• Most volume of atom=empty space.
Scientist
Famous Scientist
Experiment
Name
Conclusion
JJ Thomson
Cathode Ray
Electron (-)
Robert Millikan
Oil Drop
Mass of a single
electron.
Ernest Rutherford Gold foil
Proton (+)
______________ _______________ _____________
James Chadwick
Beryllium Detection Neutron (0)
Discovery of the Electron
• J.J. Thomson (1897) used a cathode ray
tube to deduce the presence of a negatively
charged particle.
• Discovered the – electron particle.
• (1906)Thomson
awarded the Nobel
Prize in chemistry for
discovery of the
electron.
• The atom could be
broken down into
smaller particles.
Thomson’s Atomic Model
• Thomson believed electrons were like
plums in a + charged “pudding”.
• He called it the “plum pudding” model.
Rutherford’s Gold Foil
Experiment
• Alpha particles are helium nuclei.
• Particles fired at a thin sheet of gold foil.
• Particle hits on screen (film) are detected.
Rutherford’s Findings
• Most particles passed
right through screen.
• Few particles deflected.
• VERY FEW were greatly
deflected.
• Conclusions:
• The nucleus is small.
• The nucleus is dense.
• The nucleus is + charged.
Section 3-3
• Atomic number:
number of protons in
the nucleus of atom.
• # p(+) = # e(-)
6
C
Carbon
12.011
• Atomic Mass
• number of protons & neutrons in
the nucleus.
• Atomic Mass=protons + neutrons
• Atomic Mass C = 12.011
• Mass number = rounded atomic
mass
• Mass Number C = 12
Nuclear Symbols
Atomic
Mass
Mass
number
(p+ +(p+n)+ n)
Atomic number
(# of p+)
235
Element
symbol
92
235
92
U
Hyphen Notation
Sodium-23
(23 is the atomic mass)
Sooo… 23- 11 (atomic #) = 12 for the
# of neutrons.
Atomic number of 11 is the # of
protons (11) and electrons(11).