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Chapter 4: Atoms
4.1 The Development of Atomic Theory
Key Ideas:
(1) Who came up with the first theory of atoms?
(2) What did Dalton add to the atomic theory?
(3) How did Thompson discover the electron?
(4) What is Rutherford’s atomic model?
Key Terms: electron , proton, nucleus, law of definite proportions
Why it Matters:
The electrons stripped from atoms are used in television tubes to help create the
images on a television screen.
The Beginnings of Atomic Theory:
(1) In 4th BC, the Greek philosopher Democritus suggested that the universe was
made of indivisible units.
(2) He called the units atoms from the Greek work “atomos” meaning unable to
be cut or divided.
(3) Democritus did not have evidence for his atomic theory.
Dalton’s Atomic Theory:
(1) In 1808, an English schoolteacher named John Dalton proposed that all
atoms of a given element were exactly alike, and atoms of different elements
could join to form compounds.
(2) Dalton used experimental evidence to support his theory.
(3) According to the law of definite proportions, a chemical compound always
contains the same elements in exactly the same proportions by weight and
mass.
(4) Dalton’s theory could not explain all of the experimental evidence.
Thomson’s Model of the Atom:
(1) In 1897, J.J. Thomson, a British scientist was experimenting with electricity
and revealed that atoms could be divided into smaller parts.
(2) Thomson’s cathode ray tube experiment suggested that cathode rays were
made of negatively charged particles that came from inside the atoms.
(3) The cathode ray tube experiment included a electrical current applied across
two metal plates at the end of the vacuum tube called a cathode (neg.
charge) and a anode (pos. charge). Thomson used magnets to determine the
beam of electrons must have come from the negative cathode.
(4) Thomson’s model was called the plum-pudding model because the believed
that electrons were spread throughout the atom, just as blueberries in a
muffin.
Rutherford’s Model of the Atom:
(1) Ernest Rutherford, a British scientist found that Thomson’s model needed
revising.
(2) Rutherford proposed that most of the mass of the atom was concentrated at
the atom’s center.
(3) Rutherford proved this to be true with his gold-foil experiment.
(4) Rutherford’s experiment suggested that an atom’s positive charge is
concentrated in the center of the atom in a dense core called the nucleus.
(5) Rutherford suggested that the nucleus is very small compared to the size of
the atom.
(6) Rutherford’s Model: the negative electrons orbit the positively charged
nucleus in much the same way that planets orbit the sun.
4.2 The Structure of Atoms
Key Ideas:
(1) What is the difference between protons, neutrons, and electrons?
(2) What do atoms of an element have in common with other atoms of the same
element?
(3) Why do isotopes of the same element have different atomic masses?
(4) What unit is used to express atomic mass?
Key Terms: proton, neutron, atomic number, mass number, isotope, unified atomic
mass unit, mole
Why It Matters:
Radioisotopes emit energy when they decay. To diagnose and treat diseases,
doctors use this property of radioisotopes to track where in the body certain atoms
go.
What is an Atom?
(1) Atoms are made up of three main subatomic particles that are distinguished
by mass, charge, and location in the atom.
Particle
Proton
Neutron
Electron
Charge
+1
None
-1
Mass (kg)
1.67 x 10-27
1.67 x 10-27
9.11 x 10-31
Location in Atom
Nucleus
Nucleus
Outside Nucleus
(2) Each element has a unique number of protons identifying that element.
(3) Most atoms do not have an overall charge. This is because most atoms have
equal numbers of positive protons and negative electrons and their charges
exactly cancel.
(4) If an atom loses or gains electrons, it becomes charged. A charged atom is
called an ion.
(5) The electric force holds the atoms together. The electric force between
protons in the nucleus and electrons outside the nucleus holds the atom
together.
Atomic Number and Mass Number
(1) Atoms of each element have the same number of protons, but they can have
different numbers of neutrons.
(2) Atomic Number – the number of protons in the nucleus of an atom.
(3) Mass Number – the sum of the numbers of protons and neutrons in the
nucleus of an atom.
(4) Isotopes – an atom that has the same number of protons as other atoms of
the same element do but that has a different number of neutrons ( and thus a
different atomic mass).
(5) Unstable isotopes, called radioisotopes, emit radiation and decay into other
isotopes. A radioisotope continues to decay until the isotope reaches a
stable form. (decay can vary from a fraction of a second to millions of years)
(6) The number of subatomic particles can be calculated.
a. Protons = Mass Number - Neutrons
b. Neutrons = Mass Number – Protons
c. Mass Number = Protons + Neutrons
Atomic Masses
(1) Because working with such tiny masses is difficult, atomic masses are
usually expressed in unified atomic mass units.
a. A unified atomic mass unit (u) is equal to one-twelth of the mass of a
carbon -12 atom (sometimes called the atomic mass unit, amu)
(2) Average atomic mass is a weighted average. It is the average of all of
elements isotopes.
(3) The mole is useful for counting small particles.
(4) Mole – the SI base unit used to measure the amount of a substance whose
number of particles is the same as the number of atoms of carbon in exactly
12 g of carbon-12.
(5) The mole’s value is written as 6.02 x 1023 and is called Avogadro’s number.
(6) 1 mole is equal to the mass of individual atoms or the molar mass of a
compound. Ex. 1 mole = 55.84 g Fe or 1 mole = 18.02 g H2O.
(7) 1 mole is equal to 22.4 L of gases.
(8) 1 mole is equal to 6.02 x 1023 atoms, ions, molecules or particles.
Conversions in moles, grams, liters, and atoms (molecules, particles)
Mole to gram: Convert 2.50 mol S to grams.
2.50 mol x 32.06 g S = 80.15 g S
1 mol
Steps in solving problem:
(1) List given and unknown values
(2) Identify the conversion factor that converts moles to grams.
(3) Multiply the amount of sulfur by the correct conversion factor.
Gram to Mole: Convert 95 grams of Na to mol.
95 g Na x 1 mol Na = 4.13 mol Na
22.98 g Na
Gram to Atoms: Convert 60 g of Au to atoms.
60 g Au x 1 mol Au x 6.02 x 1023 atoms
196.96 g Au
1 mol
= 1.83 x 1023 atoms
4.3 Modern Atomic Theory
Key Ideas:
(1) What is the modern model of the atom?
(2) How are the energy levels of an atom filled?
(3) What makes an electron jump to a new energy level?
Key Terms: orbital, valence electrons, and photon
Why It Matters:
Modern atomic theory explains why different elements produce different colors in
fireworks.
Modern Models of the Atom
(1) In the modern atomic model, electrons can be found only in certain energy
levels, not between levels. The location of electrons cannot be predicted
precisely.
(2) In 1913, Niels Bohr, a Danish physicist, suggested that the energy of each
electron was related to the electron’s path around the nucleus.
(3) Electrons can gain energy and move to higher energy levels or lose energy
and move to lower energy levels.
(4) By 1935, Bohr’s model no longer explained all aspects of electron behavior.
The new model demonstrated electrons behaving more like waves on a
vibrating string than like particles.
(5) Determining the exact location of an electron in an atom and its speed and
direction is impossible. Scientist can only calculate the chance of finding an
electron in a certain place.
(6) Orbital – a region in an atom where there is a high probability of finding
electrons.
Electron Energy Levels
(1) The number of energy levels that are filled in an atom depends on the
number of electrons.
(2) Valence electrons – electrons that are found in the outermost shell of an
atom and that determines the atom’s chemical properties.
(3) There are 4 types of orbitals: s, p, d, and f
a. S orbital – the simplest of the orbitals. S orbital has the lowest energy
and can only hold 2 electrons. (shaped like a sphere)
b. P orbital – each p orbital can hold two electrons, so the three p
orbitals can hold 6 electrons. (shaped like a dumbbell)
c. D and F orbitals – are more complex. 5 possible d orbitals and 7
possible f orbitals.
d. Orbitals determine the number of electrons that each level can hold.
Electron Transitions
(1)
(2)
(3)
(4)
(5)
Electrons jump between energy levels when an atom gains or loses energy.
The lowest state of energy of an electron is called the ground state.
As electrons gains energy, it moves to an excited state.
Electrons gains energy by absorbing a particle of light called a photon.
Electrons lose energy and fall back to a lower level releasing a photon.

A photon is defined as a unit or quantum of light.
Atoms absorb or emit light at certain wavelengths.
(1) The energy of photons is related to the wavelength of the light.
(2) High-energy photons have short wavelengths and low-energy photons have
long wavelengths.
(3) Because atoms gain or lose energy in certain amounts, they absorb or emit
only certain wavelengths.
(4) Each element has a unique atomic structure and the wavelengths emitted
depend on the particular element.
(5) Wavelengths serve as a atomic fingerprint for identification of elements.