Download Atomic Theory

Atomic Structure
Chapter 5
The Atomists:
The first atomic theory
• 460 BCE: Greek
Democritus suggested that
matter is “ composed of
minute, invisible,
indestructible particles of
pure matter which move
about eternally in infinite
empty”
http://www.winneconne.k12.wi.us/middle_school/7th%20Grade/LENZ/history_of_atomic_theory.htm
Development of Modern Atomic
Theory
• In 1782, a French chemist, Antoine Lavoisier
(1743-1794), made measurements of chemical
change in a sealed container.
• He observed that the mass
of reactants in the container
before a chemical reaction
was equal to the mass of the
products after the reaction.
Law of Conservation of Mass
• Lavoisier concluded
that when a chemical
reaction occurs, mass
is neither created nor
destroyed but only
changed.
• Lavoisier’s conclusion
became known as the
law of conservation of
mass.
Dalton’s Atomic Theory
• John Dalton (17661844), an English
schoolteacher and
chemist, studied the
results of experiments
by Lavoisier and many
other scientists.
• Dalton proposed his
atomic theory of matter
in 1803.
Dalton’s Atomic Theory
(The main points)
1. All matter is made up of indivisible particles
called atoms.
2. All atoms of one element are exactly
alike, but are different from atoms of
other elements.
3. Atoms can mix in simple whole number
ratios to form compounds.
4. Chemical reactions happen when atoms
are separated, joined, or rearranged.
The Electron
• Because of Dalton’s atomic theory, most scientists in
the 1800s believed that the atom was like a tiny
solid ball that could not be broken up into parts.
• In 1897, a British
physicist, J.J. Thomson,
discovered that this
solid-ball model was not
accurate.
Cathode-Ray Tube
• Thomson knew that objects with like charges repel
each other, and objects with unlike charges attract
each other
The Electron
• Thomson named the negative particles found
in his experiment electrons.
• Electrons (e-):
– Are the smallest subatomic particles
– Have a charge of -1
Protons
• Atoms aren’t negative, so there must be
something positive in the atom to balance the
electrons.
• Protons (p+):
– Positively charged subatomic particles (+1)
– Mass 1 proton = mass 1840 electrons
• At this point, it seemed that atoms were made up
of equal numbers of electrons and protons
• However, in 1910,
Thomson
discovered that
neon consisted of
atoms of two
different masses
Neutrons
• Because of the discovery of isotopes,
scientists hypothesized that atoms contained
still a third type of particle that explained
these differences in mass.
• Neutrons (n0):
• Neutral charge
• Same mass as a proton
Rutherford’s Gold Foil Experiment
• In 1909, a team of
scientists led by Ernest
Rutherford in England
carried out the first of
several important
experiments that
revealed an arrangement
far different from the
cookie-dough model of
the atom.
The Gold Foil Experiment
The Nuclear Model of the Atom
• To explain the results of the experiment,
Rutherford’s team proposed a new model
of the atom.
• Because most of
the particles passed
through the foil,
they concluded that
the atom is nearly
all empty space.
The Nuclear Model of the Atom
• Because so few particles were deflected, they
proposed that the atom has a small, dense,
positively charged central core, called a
nucleus.
• Nucleus:
• Contains most of the mass of the atom
(protons and neutrons)
• Is very small compared to the rest of the
atom
You know you want to know how
many of these particles are in an
atom, right?
• Atomic Number : # of protons in the atom
• Mass number: # of protons + # neutrons
How to get the info from the
periodic table
1
H
1.0013
Atomic Number:
Number of protons in
the nucleus of the
atom.
In this chapter you will fill out
lots of tables…
Name
Helium-4
Carbon-12
Carbon-13
Protons
Electrons
Neutrons
How many subatomic particles?
• # protons = atomic number (always)
• For a NEUTRAL atom # electrons=# protons
(this will be a bit different in upcoming chapters…)
• # neutrons = mass number - # protons
(always)
In this chapter you will fill out
lots of tables…
Name
Protons
Electrons
Neutrons
Helium-4
Carbon-12
Carbon-13
Carbon-12 and Carbon-13 are isotopes
Isotopes
• Atoms of an element that are chemically
alike but differ in mass are called isotopes of
the element.
So why the difference in mass?
So what is the bottom number
for?
1
H
1.0013
Atomic Number:
Number of protons in
the nucleus of the
atom.
The atomic mass unit (amu)
• Atoms don’t have much mass! (10-23 g)
– This is small, and a bit hard to work with
• Use carbon-12 isotope as a ref. and make
12 carbon atom= 1 amu
• On the periodic table the value is a weighted
average of all isotopes of the element – that
is why it isn’t a whole number (includes
carbon-12, carbon-13, etc).
The Periodic Table
• Chemists kept discovering new
elements and decided they
needed a way to organize
them.
• Dmitri Mendeleev constructed
the 1st periodic table
– Grouped elements according to
similar chemical properties
– He was so smart he left spaces
for elements that would be
discovered later on.
The modern periodic table
• Henry Moseley tweaked the table and
ordered by atomic number rather than atomic
mass.
Metals
Noble Gases
Alkaline Earth Metals
Halogens
Group or family
Period
Transition Metals
Alkali Metals
Nonmetals
Inner Transition Metals