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Atomic Theory
Chapter 4
History
Democritus
• 400BC
• Greek philosopher
• Thought all matter
was made up of
atoms (atomos),
which are the basic,
indivisible particles
of matter
Aristotle
• Believed all matter
was continuous-did
not believe in atoms
• His opinion was
accepted for nearly
2000 years
Late 1700’s
Scientists agreed that
1. Most natural materials are mixtures of pure
substances.
2. Pure substances are elements or compounds.
3. Law of constant composition-a compound
always has the same composition.
John Dalton
• English schoolteacher
who tied all three laws
together in his atomic
theory
• Dalton turned
Democritus’s idea
into a scientific theory
that could be tested
by experimentation.
Dalton’s Atomic Theory
1.
2.
3.
4.
5.
Elements are made of tiny particles called
atoms.
All atoms of a given element are identical.
The atoms of an element are different from
those of any other element.
Atoms of one element combine with others
to form compounds.
Atoms are indivisible in chemical processes.
A chemical reaction changes the way the
atoms are grouped together.
Joseph John Thomson
• Cathode-ray
experiment: showed
that the atoms of
any element can be
made to emit tiny
negative particles
• Determined the
charge ratio of
electrons
William Thomson
• Plum pudding
model-a bunch of
positive stuff with
the electrons
scattered
throughout.
Rutherford, Geiger, Marsden-nucleus
• Gold foil
experiment, which
led to the discovery
of the nucleus.
• Like bullets through
a tissue
Lead
block
Uranium
Florescent
Screen
Gold Foil
What he expected
Because, he thought the mass was
evenly distributed in the atom.
What he got
+
Atomic Structure
Protons
• The number of protons in an atom determines
the element’s identity
• Nuclear forces hold the nuclear particles
together
• The atomic number equals the number of
protons
Electrons
• Are very small.
• If the nucleus is a grape, the electrons would be
about one mile away.
• Have a negative charge
• The arrangements of electrons determines the
element’s chemical properties.
Neutrons
• Mass number= protons+neutrons
• Neutrons=mass number-atomic number
• Isotope-atoms that have the same number of
protons and electrons but different numbers of
neutrons (disproves point 2 of Dalton’s theory)
• Nuclide-any isotope of any element
Table 2.1 The Mass and Charge of
the Electron, Proton, and Neutron
PracticeGive the protons, neutrons, and
electrons for each
•
•
•
•
Mercury
Sodium
Carbon
13C
6
Answers
•
•
•
•
Mercury 80p, 80e, 121n
Sodium 11p, 11e, 12n
Carbon 6p, 6e, 6n
13C
6p, 6e, 7n
6
Ions
• An ion is formed when we remove or add an
electron to a neutral atom.
• Cation-a positive ion
• Anion-a negative ion
Ion Example
Regular sodium has 11 electrons, 11 protons, and
12 neutrons. If we take away 1 electron, it would
have 10 electrons (-), and 11 protons (+) so the
charge would be +1. (Neutrons would stay the
same).
Periodic Table
Organization
Groups/families-vertical columns
Periods-horizontal rows
Conductors Metals
Lose electrons (cations +)
Malleable and ductile
Nonmetals
Brittle
Gain electrons (anions -)
Covalent bonds
Semi-metals or Metalloids
Alkali Metals
Alkaline Earth Metals
Halogens
Transition metals
Rare Earth Metals
(Inner transition metals)
Noble Gases
Periodic Table
Label the following on your
periodic table
Group
1 and 11
2 and 12
13
14
Charge
+1
+2
+3
+-4
Group Charge
15
-3
16
-2
17
-1
18
none
+1 +2
-4 -3 -2 -1
+3
+1 +2
Trends
Notice that metals tend to give up
electrons while nonmetals tend to
gain electrons.
Ionic Compounds
• Contains a metal and a nonmetal (causes ions
that is why it is called ionic)
• The net charge of an ionic compound has to be
zero.
Ionic Compound Examples
• Sodium chloride Na +1 Cl -1 NaCl
• Magnesium chloride Mg +2 Cl -1 MgCl2