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Prentice Hall Physical Science
Chapter 4 Atomic Structure
4.1 Studying Atoms
A. Ancient Greek Models of Atoms
- Democritus believed all matter was made
of extremely small particles that could not be
divided
- he called the particles atoms (from the
Greek atomos meaning indivisible)
- he thought there were different types of
atoms with different properties (liquids had
round, smooth atoms and solids were rough
and prickly)
B. Dalton’s Atomic Theory
- he proved atoms exist by discovering that
compounds have a fixed composition
-
Dalton’s Theory:
-
all elements are composed of atoms
•all atoms of the same element have the same mass, and atoms of
different elements have different masses
•compounds contain atoms of more than one element
•in a particular compound, atoms of different element always
combine in the same way
-
he pictured atoms as solid spheres
- in time it was found that not all of this
was completely correct
C. Thomson’s Model of the Atom
- some objects have a positive or negative
charge and objects with the same charge repel
and objects with opposite charges attract
- Thomson used an electric current to learn
more about atoms
- he passed an electric current through a
gas and a glowing beam appeared
- the beam was repelled by a negatively
charged metal plate (like charges repel and
opposite charges attract)
- Thomson concluded that the beam most
be made of negatively charged particles,
smaller than atoms.
electrically neutral, he developed a model in
which negative charges were evenly scattered
throughout an atom filled with a positively
charged pudding like material
- called the “Plum Pudding” model (or
chocolate chip cookie dough)
D. Rutherford’s Atomic Theory
- Rutherford discovered positively charged,
fast moving particles called alpha particles
- he shot some of these particles a piece of
gold foil, thinking they would go straight
through
- when he actually performed the
experiment, he found that the positively
charged alpha particles were greatly deflected
by the gold (figure 7, p. 104)
- he determined that the alpha particles
must have come very close to another
positively charged particle causing them to be
repelled and put off course
- this led him to believe that the positive
charge in an atom is not evenly distributed,
but concentrated in a very small central area
he called the nucleus
- his model of the atom is a small,
positively charged nucleus surrounded by
negatively charged particle
Prentice Hall Physical Science
Chapter 4 Atomic Structure
4.2 The Structure of the Atom
A. Properties of Subatomic Particles
-
Protons (p+)
-
positively charged (1+)
-
found in the nucleus
-
mass of 1 amu (atomic mass unit)
-
Neutrons (no)
-
no charge (neutrally charged or neutral)
-
found in the nucleus
-
mass of 1 amu
-
Electrons (e-)
-
negatively charged (1-)
- found in the electron cloud outside the
nucleus
-
mass of 1/1836 amu
B. Atomic Number, Mass Number, and
Atomic Mass
-
Atomic Number
-
it is the number of protons in an atom
- it is the WHOLE number in each square
of the periodic table
- all atoms of the same element have the
same atomic number
•since atoms are electrically neutral, the number of protons
(positive charges) they have is equal to the number of electrons
(negative charges) they have
•it is the number of protons in an atom that determine what
element the atom is
-
Mass Number
- it is the number of protons PLUS the
number of neutrons in an atom
-
mass number = #protons + #neutrons
-
#neutrons = mass number - #protons
-
NOT on the periodic table
-
Atomic Mass
•because protons are so small, it is hard to find their mass using
traditional units, so scientists developed the atomic mass unit (amu) to
describe the mass of a proton
•1 amu = 1/12 the mass of a carbon-12 atom
•every atoms of one element does NOT have the same number of
neutrons so every atom in one element does not have the same mass
•atomic mass is the weighted average of all of the possible atoms
in an element
•example: chlorine comes in two forms: chlorine-35 and
chlorine-37
• 75% is chlorine-35, 25% is chlorine-37
C. Isotopes
- atoms of the same element with the same
number of protons, but different numbers of
neutrons and different mass numbers
- example:
oxygen-18
oxygen-16, oxygen-17, and
Prentice Hall Physical Science
Chapter 4 Atomic Structure
4.3 Modern Atomic Theory
A. Bohr’s Model of the Atom
-
also called the planetary model
- it focused on the arrangement of the
electrons in the atom
-
Energy Levels
•Bohr said that electrons move with constant speed in fixed orbits
(energy levels) around the nucleus
•each electron in an atom has a specific amount of energy which is
equal to the energy of one of the energy levels in the atom
•an electron must be in an energy level, it cannot be between
•energy levels closest to the nucleus have the lowest amount of
energy and the ones frthest away has the most energy
•if the atom loses or gains energy, the electrons lose or gain energy
too and must, therefore, change energy levels to match their energy
•if the electron gains energy, the electron moves up to a higher E
level
- electrons want to be in the lowest E state
possible
B. Electron Cloud Model
- scientists now know that electrons don’t
move like planets; they are MUCH less
predictable
- in this model, scientists state only where
electrons probably are
- the electron cloud represents where the
electrons probably are outside the nucleus;
the cloud is denser where the probability of
finding an electron is greater
-
example: fan blades spinning
C. Atomic Orbitals
- the electron cloud contains orbitals which
are the regions around the nucleus where an
electron is likely to be found
-
each orbital can only hold 2 electrons
- the lowest E level contains only 1 orbital,
and can hold only 2 electrons
-
E level 2 = 4 orbitals and 8 electrons
-
E level 3 = 9 orbitals, 18 electrons
-
E level 4 = 16 orbitals, 32 electrons
D. Electron Configurations
- the arrangement of electrons in the
orbitals of an atom
- the most stable is the one with the
electrons in the lowest possible E levels
-
this is the ground state