PowerPoint Lectures
to accompany
Physical Science, 6e
Chapter 8
Atoms and Periodic
Properties
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
The origins of quantum physics
Atomic structure discovered
Ancient Greeks
• Democritus (460-362 BC) - indivisible
particles called “atoms”
• Prevailing argument (Plato and Aristotle) matter is continuously and infinitely divisible
John Dalton (early 1800’s) - reintroduced
atomic theory to explain chemical reactions
Dalton’s atomic theory
1. All matter = indivisible atoms
2. An element is made up of identical atoms
3. Different elements have atoms with different
masses
4. Chemical compounds are made of atoms in
specific integer ratios
5. Atoms are neither created nor destroyed in
chemical reactions
Discovery of the electron
J. J. Thomson (late
1800’s)
• Performed cathode ray
experiments
• Discovered negatively
charged electron
• Measured electron’s
charge-to-mass ratio
• Identified electron as a
fundamental particle
• What appears to be visible light coming through the slit in this vacuum
tube is produced by cathode ray particles striking a detecting screen.
You know it is not light, however, since the beam can be pulled or
pushed away by a magnet and since it is attracted to a positively
charged metal plate. These are not the properties of light, so cathode
rays must be something other than light.
Electron charge and mass
Robert Millikan
(~1906)
• Studied charged oil
droplets in an electric
field
• Charge on droplets =
multiples of electron
charge
• Charge + Thomson’s
result gave electron
mass
Early models of the atom
• Dalton - atoms indivisible
• Thomson and Millikan experiments
– Electron mass very small, no measurable volume
– What is the nature of an atom’s positive charge?
• Thomson’s “Plum pudding” model
– Electrons embedded in blob of positively charged
matter like “raisins in plum pudding”
The nucleus
Ernest Rutherford (1907)
• Scattered alpha particles off
gold foil
• Most passed through without
significant deflection
• A few scattered at large
angles
• Conclusion: an atom’s
positive charge resides in a
small, massive nucleus
• Later: positive charges =
protons
• James Chadwick (1932):
also neutral neutrons in the
nucleus
• From measurements of alpha particle scattering,
Rutherford estimated the radius of an atom to be
100,000 times greater than the radius of the nucleus.
This ratio is comparable to that of the (A) thickness of
a dime to the (B) length of football field.
– In 1917 Rutherford broke up the nucleus of the
nitrogen atom by bombarding it with alpha particles
and was able to identify a particle with a positive
charge called a proton.
• He also thought that there were neutral particles in the
nucleus called neutrons. They were later identified by
James Chadwick
• The atom has a tiny, massive nucleus made up of protons
and neutrons.
• Negatively charged electrons, whose charge balances
the charge on the protons, move around the nucleus at a
distance of about 100,000 times the radius of the nucleus.
• atomic number is the number of protons in the nucleus.
All atoms of the same element have the same atomic
number.
Every element has a distinctive atomic number that
identifies it
The nuclear atom
• Atomic number
– Number of protons in
nucleus
– Elements distinguished
by atomic number
– 113 elements identified
– Number of protons =
number of electrons in
neutral atoms
• Isotopes
– Same number of protons;
different number of
neutrons
Atomic symbols and masses
Mass number
• Number of protons +
neutrons
Atomic mass units (u)
• 1/12 of carbon-12
isotope mass
Atomic weight
• Atomic mass of an
element, averaged over
naturally occurring
isotopes
Classical “atoms”
Predictions of classical theory
•
•
•
•
•
Electrons orbit the nucleus
Curved path = acceleration
Accelerated charges radiate
Electrons lose energy and spiral into nucleus
Atoms cannot exist!
Experiment - atoms do exist
 New theory needed
The quantum concept
• Max Planck (1900)
– Introduced quantized
energy
• Einstein (1905)
– Light made up of
quantized photons
• Higher frequency
photons = more
energetic photons
Atomic spectra
Blackbody radiation
• Continuous radiation
distribution
• Depends on temperature of
radiating object
• Characteristic of solids,
liquids and dense gases
Line spectrum
• Emission at characteristic
frequencies
• Diffuse matter: incandescent
gases
• Illustration: Balmer series of
hydrogen lines
Bohr’s theory
Three rules:
1. Electrons only exist in
certain allowed orbits
2. Within an orbit, the
electron does not
radiate
3. Radiation is emitted or
absorbed when
changing orbits
Quantum theory of the atom
• Lowest energy state =
“ground state”
• Higher states = “excited
states”
• Photon energy equals
difference in state
energies
• Hydrogen atom
example
– Energy levels
– Line spectra
• These fluorescent lights emit light as electrons of mercury
atoms inside the tube gain energy from the electric current. As
soon as they can, the electrons drop back to their lower-energy
orbit, emitting photons with ultraviolet frequencies. Ultraviolet
radiation strikes the fluorescent chemical coating inside the
tube, stimulating the emission of visible light.
Quantum mechanics
• Bohr theory only modeled the line spectrum
of H
• Further experiments established waveparticle duality of light and matter
– Young’s two slit experiment produced interference
patterns for both photons and electrons
Matter waves
Louis de Broglie
(1923)
• Postulated matter
waves
• Wavelength related to
momentum
• Matter waves in atoms
are standing waves
Wave mechanics
• Developed by Erwin Schrodinger
• Treats atoms as three dimensional systems
of waves
• Contains successful ideas of Bohr model and
much more
• Describes hydrogen atom and many electron
atoms
• Forms our fundamental understanding of
chemistry
• Wave Mechanics
– Electrons do emit light in certain wavelengths based
on their energy levels (orbital radius)
– Since waves spread out from the electron, the wave
mechanic model predicts an area where an electron
would be found, and not a specific place where it
would be found.
The quantum mechanics
model
• Quantum numbers
specify electronic
quantum states
• Visualization - wave
functions and
probability
distributions
• Electrons
delocalized
Electronic quantum numbers
in atoms
1. Principle quantum number,
n
– Energy level
– Average distance from
nucleus
2. Angular momentum
quantum number, l
– Spatial distribution
– Labeled s, p, d, f, g, h, …
3. Magnetic quantum number
– Spatial orientation of orbit
4. Spin quantum number
– Electron spin orientation
• The Quantum Mechanics Model
– Quantum mechanics describes the energy levels
of an electron wave with four quantum numbers.
•
•
•
•
distance from nucleus (n)
energy sublevel (l)
orientation in space. (m)
direction of spin (s)
– Principal quantum number (n)
• Describes main energy level of the electron in terms of its
distance from the nucleus.
• n = 1, 2, 3, 4, 5, 6, 7
– Angular momentum quantum number ( l )
• Defines energy sublevels within the main energy
levels
• s, p, d, or f designating the type of orbital and
also the orbital shape.
• The Heisenberg Uncertainty Principle states
that you cannot measure the momentum and
exact position of an electron at the same time.
–What you can measure is the probability that
an electron will be found in a certain area,
called an orbital.
• (A)An electron distribution sketch representing
probability regions where an electron is most likely to
be found. (B) A boundary surface, or contour, that
encloses about 90 percent of the electron distribution
shown in (A). This three-dimensional space around
the nucleus, where there is the greatest probability of
finding an electron, is called an orbital.
– Magnetic quantum number ( m)
• Defines the orientation in space of the orbitals
relative to a magnetic field.
• The s orbital has one orientation
• The p sublevel can have 3 orientations
• The d sublevel can have 5 orientations
• The f sublevel can have 7 orientations.
• (A) A contour representation of an s orbital. (B)
A contour representation of a p orbital.
– Spin quantum number ( s )
• Describes the direction of spin of an electron in
its orbit.
• Electrons occur in pairs and each of the
orientations for a sublevel can have one electron
pair.
• Experimental
evidence
supports the
concept that
electrons can
be considered
to spin one way
or the other as
they move
about an orbital
under an
external
magnetic field.
– Pauli Exclusion Principle
• No two electrons can have the same set of quantum
numbers.
• At least one of the quantum numbers must differ.
Electron configuration
• Arrangement of
electrons into atomic
orbitals
• Principle, angular
momentum and
magnetic quantum
numbers specify an
orbital
• Specifies atom’s
quantum state
• Pauli exclusion principle
– Each electron has
unique quantum
numbers
– Maximum of two
electrons per orbital
(electron spin
up/down)
• Chemical properties
determined by
electronic structure
Writing electron configurations
• Electrons fill available
orbitals in order of
increasing energy
• Shell capacities
–
–
–
–
s=2
p=6
d = 10
f = 14
• Example: strontium (38
electrons)
1s 22s 22p 63s 23p 64s 23d 104p 65s 2
• Electron Configuration
– This is a shorthand designation for electron
orientation.
– The lowest possible energy level is n=1.
• If one electron already occupies this energy level, a
second can only occupy it if it has a different spin
quantum number.
– Electron configurations tells you the quantum
numbers of the electron.
–
Energy sublevel

Principle quantum number 1s2  two electrons
• There are three possible
orientations of the p orbital,
and these are called px, py,
and pz. Each orbital can hold
two electrons, so a total of
six electrons are possible in
the three orientations; thus
the notation p6.
Periodic chemical properties
• Understood in terms of
electron configuration
• Electrons in outer orbits
determine chemical
properties
• Summarized in the
periodic table
• Rows = periods
• Columns = families or
groups
–
–
–
–
Alkali metals (IA)
Alkaline earths (IIA)
Halogens (VIIA)
Noble gases (VIIIA)
• A-group families = main
group or representative
elements
• B-group = transition
elements or metals
The periodic table
Metals, nonmetals and
semiconductors
• Noble gases - filled shells,
inert
• 1-2-3 outer electrons
– Lose to become positive
ions
– Metals
• 5-7 outer electrons
– Tend to gain electrons
and form negative ions
– Nonmetals
• Semiconductors intermediate between
metals and nonmetals
•Elements and the
Periodic Table
• Classification is arranging items into groups or
categories according to some criteria.
• Classifying helps organize your thinking &
reveals patterns that might go unnoticed.
• The act of classifying creates a pattern that
helps you recognize and understand the
behavior of fish, chemicals, or any matter in
your surroundings.
• These fish, for
example, are
classified as
salmon because
they live in the
northern Pacific
Ocean, have
pinkish colored
flesh, and
characteristicall
y swim from salt
to fresh water to
spawn.
• Classifying Matter
• Matter is usually defined as anything that has
mass and occupies space.
• Metals and Nonmetals
– A metal had the following properties.
•
•
•
•
•
Metallic luster
High heat and electrical conductivity.
Malleability, able to be rolled or pounded into a thin sheet.
Ductile, can be pulled into a wire.
all metals are solids at room temp. except mercury,
which is
is a liquid
• Most matter can be
classified as metals or
nonmetals according to
physical properties.
Aluminum, for example, is a
lightweight kind of matter
that can be melted and
rolled into a thin sheet or
pulled into a wire. Here you
see aluminum pop cans
that have been compressed
into1,600 lb bales for
recycling, destined to again
be formed into new pop
cans, aluminum foil, or
perhaps aluminum wire.
– A nonmetal has the following properties
• No metallic luster
• Poor conductor of heat and electricity.
• When it is a solid it is brittle so it cannot be
pounded or pulled into a wire.
• Occur as solids, liquids, or gases at room
temperature
*Chemistry is the science concerned with the study
of
the composition, structure, and properties of
substances and the transformations they
undergo.
• Solids, Liquids, and Gases
– Gases have no defined shape or defined volume
• Low density
– Liquids flow and can be poured from one container
to another
• Indefinite shape and takes on the shape of the container.
– Solids have a definite volume
• Have a definite shape.
• (A)A gas dispenses throughout a container, taking the
shape and volume of the container. (B) A liquid takes
the shape of the container but retains its own volume.
(C) A solid retains its own shape and volume.
• Mixtures and Pure Substances
– A mixture has unlike parts and a composition that
varies from sample to sample. Can be separated
into parts by
physical means, involving physical changes. Can
be
separated into pure substances.
– A heterogeneous mixture has physically distinct
parts with different properties.
– A homogeneous mixture is the same throughout
the sample
– Pure substances are substances with a fixed
composition
• A classification scheme for matter.
– A physical change is a change that does not alter
the identity of the matter.
– A chemical change is a change that does alter the
identity of the matter.
– A compound is a pure substance that can be
decomposed by a chemical change into simpler
substances with a fixed mass ratio (There are
millions of different compounds)
– An element is a pure substance which cannot be
broken down into anything simpler by either
physical or chemical means. (There are 115 known
elements)
• Sugar (A) is a compound that can be easily
decomposed to simpler substances by heating.
(B) One of the simpler substances is the black
element carbon, which cannot be further
decomposed by chemical or physical means.
• Elements
• Names of Elements
– The first 103 elements have internationally
accepted names, which are derived from:
• The compound or substance in which the element was
discovered
• An unusual or identifying property of the element
• Places, cities, and countries
• Famous scientists
• Greek mythology
• Astronomical objects.
• Here are some of the symbols Dalton used for
atoms of elements and molecules of
compounds. He probably used a circle for each
because, like the ancient Greeks, he thought of
atoms as tiny, round hard indivisible spheres.
– Chemical Symbols
• There are about a dozen common elements that have a
single capitalized letter for their symbol
• The rest, that have permanent names have two letters.
– the first is capitalized and the second is lower case
– the lower case letter is either the second letter in
the name, or the letter of a strong consonant heard
when the name of the element is spoken
• Some elements have symbols from their Latin names.
• Ten of the elements have symbols from their Latin names
and one element with a symbol from a German name.
• a symbol both identifies a specific element and
represents an atom of that element
• The elements of aluminum, Iron, Oxygen, and Silicon make up
about 88 percent of the earth's solid surface. Water on the
surface and in the air as clouds and fog is made up of hydrogen
and oxygen. The air is 99 percent nitrogen and oxygen.
Hydrogen, oxygen, and carbon make up 97 percent of a
person. Thus almost everything you see in this picture us made
up of just six elements.
– Symbols and Atomic Structure
• A molecule is a particle that is composed of two
or more atoms held together by a chemical bond.
• Isotopes are atoms of an element with identical
chemical properties, but different masses due to
a difference in the number of neutrons. (Frederic
Soddy)
• The atomic weight of an element is the
weighted average of all the masses of the
isotopes.
–an isotopes contribution is determined by its
relative abundance.
– Aston provided evidence of isotopes with the
mass spectrograph.
• A mass spectrum of chlorine from a mass
spectrometer. Note that that two separate masses of
chlorine atoms are present, and their abundance can
be measured from the signal intensity. The greater the
signal intensity, the more abundant the isotope.
• The atomic mass of an element is the mass of the
element compared to an isotope of carbon Carbon 12.
– Carbon 12 is assigned an atomic mass of 12.00 u.
• The number of protons and neutrons in an atom is its
mass number.
• Atomic numbers are whole numbers, identify the element
(number of protons, and electrons)
• Mass numbers are whole numbers, used to identify
isotopes.( closest whole number to atomic mass)
• The atomic mass of an isotope is not a whole number,
with exception of Carbon-12.
• Atomic weights are not whole numbers, due to being
a weighted average of the masses of the isotopes.
• The Periodic Law
• Dmitri Mendeleev gave us a functional scheme
with which to classify elements.
– Mendeleev’s scheme was based on chemical
properties of the elements and their atomic weights.
– Meyer’s scheme was based on physical properties
of the element and their atomic weights.
– It was noticed that the chemical properties of
elements increased in a periodic manner.
– The periodicity of the elements was demonstrated
by Medeleev when he used the table to predict the
occurrence and chemical properties of elements
which had not yet been discovered. Mendeleev is
given credit for developing the periodic table, due to
these dramatic
predictions.
• Mendeleev left blank
spaces in his table when
the properties of the
elements above and below
did not seem to match. The
existence of unknown
elements was predicted by
Mendeleev on the basis of
the blank spaces. When
the unknown elements
were discovered, it was
found that Mendeleev had
closely predicted the
properties of the elements
as well as their discovery.
• The Periodic Law
– Similar physical and chemical properties recur
periodically when the elements are listed in order of
increasing atomic number.
– the number of electrons around the nucleus is the
essential basis for the modern periodic table.
This fact was made known after the work of
Rutherford
And Moseley on the atomic nucleus.
• The Modern Periodic Table
• Introduction
– The periodic table is made up of rows of elements and
columns.
– The arrangement has a meaning about atomic
structure
and about chemical behavior. For the table to be
meaningful, one must understand the code of this structure.
– An element is identified by its chemical symbol.
– The number above the symbol is the atomic number
– The number below the symbol is the atomic weight of
the element.
– A row is called a period
– A column is called a family
• (A) Periods of the periodic table, and (B)
families of the periodic table.
• Periodic Patterns
– The chemical behavior of elements is determined
by its electron configuration
– Energy levels are quantized so roughly correspond
to layers of electrons around the nucleus. (resemble
the
layers of an onion)
– A shell is all the electrons with the same value of n.
• n is a row in the periodic table.
– The first three periods contain just A families. Each
period begins with a single electron in a new outer
electron shell.
– Each period ends with a completely filled outer shell
that has the maximum number of electrons for that
shell.
– The number identifying the A families identifies the
number of electrons in the outer shell, except
helium
– The outer shell electrons are responsible for
chemical reactions. Elements in the same family
have the same
number of outer shell electrons; so they will have
similar chemical properties.
– Group A elements are called representative
elements
– Group B elements are called transition elements.
• Chemical Families
– IA are called alkali metals because the react with
water to from an alkaline solution, very soft metals.
(except H2)
– Group IIA are called the alkali earth metals
because they are reactive, but not as reactive as
Group IA.
• They are also soft metals, though not as soft as alkali
metals
– Group VIIA are the halogens
• These need only one electron to fill their outer shell
• They are very reactive.( disinfectants, bleach, headlights)
– Group VIIIA are the noble gases as they have
completely filled outer shells
• They are almost non reactive.
• Four chemical families of the
periodic table: the alkali
metals (IA), the alkaline earth
metals (IIA), halogens (VII),
and the noble gases (VIIIA).
• Metals, Nonmetals, and Semiconductors
– Chemical behavior is determined by the outer
electrons.
• These are called valence electrons
– These outer shell electrons are represented using
electron dot diagrams.
– The noble gases have completely filled outer shells
and are therefore stable.
• All other elements react so as to fill their outer shell and
become more stable.
• Electron dot notation for the representative
elements.
– When an atom or molecule gain or loses an
electron it becomes an ion.
• A cation has lost an electron and therefore
has a positive charge
• An anion has gained an electron and
therefore has a negative charge.
• (A) Metals lose their outer electrons to acquire a noble
gas structure and become positive ions. (B)
Nonmetals gain electrons to acquire an outer noble
gas structure and become negative ions.
– Elements with 5 to 7 electrons in their outer shell tend to
gain electrons to fill their outer shell and become anions.
• These are the nonmetals which always tend to gain
electrons.
– Semiconductors (metalloids) occur at the dividing line
between metals and nonmetals. (silicon, germanium,
– Elements with 1, 2, or 3 electrons in their outer shell tend to
lose electrons to fill their outer shell and become cations.
• These are the metals which always tend to lose
electrons.
• about 80% of all the elements are metals
arsenic-conduct electric current under certain conditions)
Transition elements are all metals. They are located in B
Group families. Form horizontal rows of elements with
similar properties. Have variable charges. (Cu+1, Cu+2)
• The location of metals, nonmetals, and
semiconductors in the periodic table.
• The periodic table of the elements.