Chapter 4 Atoms
Section 1
Development of Atomic Theory
BR. Who came up with the first
theory of atoms?
• Objectives• 1. Give an example of how new scientific
data can cause an existing scientific
explanation to be supported, rejected or
revised
• 2. Evaluate selected theories based on
supporting scientific evidence.
Objectives (cont.)
• 3. Cite evidence that scientific
investigations are conducted for many
reasons.
• 4. Identify scientific evidence that has
caused modifications in previously
accepted theories.
• GLE’s:
Democritus
• Over 2000 years ago
• Democritus
• Universe made of indivisible units called
atoms.
• Atomos- unable to be cut or divided
• Did not have evidence to support theory
Dalton
• 1808 John Dalton revised atomic theory
• Atoms could not be divided
• All atoms of a given element were exactly
alike; and atoms of different elements
could join to form compounds.
• Based on experimental evidence
Dalton (cont.)
• Law of definite proportions- a chemical
compound always contains the same
elements in exactly the same proportions
by weight or mass. (supported Dalton’s
theory)
• Foundation of modern atomic theory
• Could not explain all experimental
evidence
Thomson
• 1897 experiment suggested that atoms
were not indivisible
• He was experimenting with electricity
• studying cathode rays not atoms
• Cathode ray tube experiment suggested
that cathode rays were made of negatively
charged particles that came from inside
the atoms
Thomson (cont.)
• Revealed that atoms could be divided
• Discovered electrons, negative particles
• New model- electrons spread through-out
the atom ( plum pudding model)
Mass and positive charge evenly distributed
Electrons scattered through out
Rutherford
• Found Thomson’s model needing revising
• Proposed that most of the mass of the
atom was in the center
• Conducted Gold-foil experiment where
most particles passed straight through
• Some particles were deflected
• Some particles came straight back
• Not what he expected
Rutherford (cont.)
• Discovered the nucleus
• Nucleus was very small
• Electrons orbit the nucleus (like sun and
planets)
• Led to new model of the atom
Section 2- Structure of Atom
• BR. Rutherford’s Gold foil experiment led
to the discovery of what?
• Objectives:
• 1. Identify the 3 subatomic particles by
location, charge, and relative mass
• 2. Describe the results of loss or gain of
electrons on charges of atoms
Objectives (cont.)
• 3. Identify valence electrons in first 20
elements.
GLE’S:
Atom
• Three subatomic particles compared by
mass, charge, and location in the atom.
• Copy chart on page 119
• Nucleus- small, dense, center of atom
• Atoms are Neutral.
Atom (cont.)
• Nucleus is made of
1. protons- + charged particle
2. neutron- neutral or no charge particle
Protons and neutrons are almost equal in
size and mass.
Electrons
• Move in a dense cloud (fan blades
moving) outside the nucleus
• Very tiny--- 1837 electrons = 1 neutron or
proton
• Negatively charged
• Exact location cannot be determined.
Speed and direction cannot be determined
• Located by shading; shaded region is
orbital: darker shading better chance to
find
Protons
• Each element has a unique number of
protons
• Elements are identified by the number of
protons that they have in the nucleus of an
atom
• Atoms are neutral because they have the
same number of p and e. They cancel
each other out.
Ions
• Atoms that have lost or gained electrons.
• Lose electrons become positive.
• Gain electrons become negative
• If atoms lose or gain electrons they are not
atoms, but IONS.
Atoms
• Atoms are held together by an electric
force
+ and – charges attract each other by an
electric force
This attraction is what holds the atom
together just like the attractive force
between solids and liquids.
Atoms and Elements
• Atoms of different elements have unique
structures.
• Because atoms have different structures, they
have different properties.
• Atoms of the same element can vary in structure
also.
• Atoms of each element have the same number
of protons, but different numbers of neutrons.
Atomic number
• Atomic number equals the number of
protons.
• Since atoms are neutral, it also equals the
number of electrons.
• Neutral atom--
+=-
Atomic numbers go from 1 to 116
Mass Number
• Equals the total number of subatomic
particles in the NUCLEUS of the atom.
• Nucleus contains p and n.
• Mass number is equal to p + n.
• Neutrons vary so mass number can vary
for the same element.
• See figure 3 on page 121
Isotopes
• Isotope has same atomic number but
different number of neutrons.
• Isotopes have same atomic number but
different mass numbers.
• Isotopes have the same number of
protons but different number of neutrons
Isotopes (cont.)
• Some isotopes are more common than
others.
• Radioisotopes- unstable isotopes that emit
radiation and decay into other isotopes.
They continue to decay until they reach a
stable isotope.
They decay at a fixed rate. (fraction of a
second to millions of years)
Isotopes (cont.)
• Since isotopes have the same number of
protons and electrons, they have the same
chemical properties.
• Isotopes have different masses.
• Isotopes of an element vary in mass
because their numbers of neutrons differ.
Isotopes of Water
• Water has 3 isotopes. Each isotope has 1
proton and 1 electron.
• Protium- 1 proton and no neutrons (mass
number of 1
• Deuterium-1 proton and 1 neutron (mass
number of 2)
• Tritium- 1 proton and 2 neutrons mass number
of 3)
Calculating Neutrons
• Isotopes are written
35
•
Cl
17
mass number ( p+n)
symbol
atomic number (p)
Mass number – atomic number = neutrons
Atomic mass
• Atoms are expressed in unified atomic
mass units because the mass is so small.
• Unified atomic mass = equal to 1/12th of
the mass of a carbon-12 atom.
• Also called atomic mass unit
Atomic Mass
• Average atomic mass for an element is a
weighted average.
• More common isotopes have more effect
than less common isotopes of the
element.
Mole
• Mole- collection of a very large number of
particles.
602,213,670,000,000,000,000,000 particles
Written 6.02 x 1023 particles and is called
Avogadro’s number
Avogadro’s number= the number of atoms in
12 grams of carbon-12. (Popcorn kernels covering
the US 310 miles tall)
Molar mass
• Molar mass- the mass in grams of 1 mole
of a substance
1 mole C12 = 12 g
Converting between
Moles and Grams
Amount x molar mass of element = Mass(g Moles
1 mole of element
3 moles x 32.07 g S
= 96.21 g S
S
1 mole S
96.27 g S
x
1 mole S = 3 mol S
32.21 g S
Work problem 1 a-d on page 126
Compounds have Molar mass
• Add all molar masses in compounds and
then work the same way.
H2O x (1.01 g H x 2) + 16 g O
1 mole H2O
1 mol H20 = 18.02 g
=
18.01 g/mol
H2O
Section 3
Modern Atomic Theory
• BR List the 3 subatomic particles and give their
location, relative mass, and charge.
• Objectives:
• Describe the results of the loss or gain of
electrons on the charges of atoms.
• Identify valence electrons in first 20 elements.
• Draw Bohr models of 1st 20 elements
• GLE’S:
Modern Model of Atom
• Electrons are found only in certain energy
levels. NOT between levels
• Location of electrons can not be predicted
precisely
• Bohr- electrons can be in only certain
energy levels.
• Bohr energy level related to electron’s
path around the nucleus
Modern Model of Atom (cont.)
• Electrons must gain energy to move to a
higher energy level.
• Electrons must lose energy to move to a
lower energy level.
• 1925 Bohr’s model revised.
Modern Atomic Theory (cont.)
• Old out – not like sun and planets
• New in –
• Electrons behave more like waves on a
vibrating string than like particles.
Energy levels and Electrons
• Many energy levels for electron to occupy
• The number of energy levels that are filled
in an atom depends on the number of
electrons.
• Valence electrons- those electrons in outer
most energy level
• Valence electrons determine the chemical
properties of the atom.
Energy levels
• Maximum electrons in energy level
•
1st = 2
•
2nd= 8
•
3rd= 18
•
4th= 32
• Must fill 1st and 2nd energy level before
going to the 3rd energy level.
Energy levels (cont.)
• There are 4 types of orbitals.
• Orbitals are s, p, d, f.
• Orbitals determine the number of electrons
that each level can hold.
Electron Jumping
• Electrons jump between energy levels
when an atom gains or loses energy.
• Lowest energy level called ground state.
• Excited state-gains energy it moves to
another level
How Electrons Move
• Electrons gain energy by absorbing a photon
and move to a higher energy level
• Photon- particle of light; each have different
energies
• The electron may fall back to previous energy
level when it releases a photon.
• Photons determine which level the electron will
• jump to.
Light• Photons determine which level the
electron will jump to.
• Atoms absorb or emit light at certain
wavelengths.
• Energy of photon is related to the
wavelength of the light.
• High energy photons= short wavelengths
• Low energy photons= long wavelengths
Atomic Fingerprint
• Because of each element’s unique atomic
structure, the wavelengths emitted depend
on the particular element.
• Each element emits its own characteristic
color. Neon= red blue=copper
sodium= yellow strontium= red
orange = calcium green= barium