Content Benchmark P.8.A.2
Students know elements can be arranged in the periodic table which shows repeating patterns
that group elements with similar properties. E/S
Development of the Periodic Table
French scientist Antoine Laurent Lavoisier (1743–1794) constructed a table which consisted of
33 known elements. The 33 elements, known during that time, were arranged in four categories;
gases, metals, nonmetals, and earths. As time progressed towards the 1800s, along came the
discovery of new elements. Scientific instruments which allowed for the increase of the new
elements included the introduction of electricity (used to breakdown the compounds into their
components) and the innovation of the spectrometer (used to identify isolated elements).
To learn more about Antoine Lavoisier, see
http://cti.itc.virginia.edu/~meg3c/classes/tcc313/200Rprojs/lavoisier2/home.html
As the discovery of elements and compounds increased over time, so did learning their
properties. Due to the overwhelming amount of elements being discovered, chemists needed a
method or tool which would assist in learning the elemental properties. In 1860, chemists came
to an agreement upon a method that would help organize and arrange the elements. This was
noted to be by their atomic masses and elemental properties.
John Newlands, an English chemist (1837-1898), developed a directorial system for elements. It
was in his studies that he noticed a repeating pattern of properties for every eighth element when
arranged according to their increasing atomic masses, a term coined as periodic. Newlands
named the observations as the law of octaves, similar to musical notes repeating every eighth
tone. Although his law was unaccepted in the scientific community for its analogy to music, as
well as its inability to be applied for all known elements, it was nonetheless shown to be correct
years later.
For more information about John Newlands and the Law of Octaves, see
http://www.wou.edu/las/physci/ch412/perhist.htm
Father of the Periodic Table
Dmitri Mendeelev, a Russian chemist (1834-1907), published a methodical scheme reflecting the
relationship between atomic mass and elemental properties between 1868 and 1870. As elements
were arranged in increasing atomic mass, a repeating pattern in elemental properties every eighth
element was also observed by Dmitri. He developed a table which organized the elements by
increasing atomic mass into columns with similar elemental properties. Mendeleev’s periodic
table became broadly accepted for its ability to predict the existence and properties of
undiscovered elements left as blank spaces on the periodic table, which would later to be found
such as scandium, gallium, and germanium. Dmitri Mendeleev is accredited as being the father
of the periodic table.
To learn more about Dmitri Mendeelev, see
http://www.aip.org/history/curie/periodic.htm
Henry Moseley
Although Mendeelev’s table was accepted by the scientific community, it was not completely
correct. As elements and their atomic masses became more accurately discovered, several
elements were noticeably misplaced. By arranging the elements according to their increasing
atomic masses, several elements were placed in groups of conflicting elemental properties. In
1913, English chemist Henry Moseley (1887-1915) observed this error which piloted him to the
discovery of the atomic number. According to Moseley, arranging the elements in increasing
order of atomic number (number of protons) resulted in a consistent periodic pattern of elemental
properties. A periodic repetition of physical and chemical properties of elements arranged by
increasing atomic number is called the periodic law.
For more information about the Henry Moseley, see
http://corrosion-doctors.org/Periodic/Periodic-Moseley.htm
The Modern Periodic Table
The modern periodic table consists of boxes, each containing the element’s name, symbol,
atomic number, and atomic mass.
Oxygen
Atomic Number
8
Element Symbol
O
Element Name
State of Matter
Atomic Mass
15.999
Figure 1. A typical box from the periodic table illustrates the element’s
name, symbol, atomic number, atomic mass, and state of matter.
The elements on a periodic table are arranged from left to right, top to bottom, in order of
increasing atomic number. The boxes are arranged in order of increasing atomic number in a
series of vertical columns, called groups or families, and horizontal rows, called periods.
Elements that have similar characteristics are arranged in vertical columns and horizontal rows,
which help to define their characteristics. Elements located in groups all share in the same
number of electrons in their outer orbital. Elements located in periods all share in the same
number of atomic orbitals.
Figure 2. Illustration of Groups versus Periods on a Periodic Table
(From http://www.chem4kids.com/files/elem_pertable.html)
There are a total of seven periods. Each group is numbered 1 through 18. The elements found in
groups 1, 2, 13, and 18 all possess a broad range of chemical and physical properties. It is for this
reason that they are referred to as the main group, or representative elements. Elements located
in groups 3 through 12 are referred to as the transition elements. Elements on the periodic table
may be classified as metals, nonmetals, and metalloids.
Figure 3. Periodic Table of Elements
(From http://www.elementsdatabase.com/)
Metals
The elements that are generally shiny and smooth, solid at room temperature, and classified to be
good conductors of heat and electricity are called metals. Most metals are malleable and ductile,
allowing them to be crushed into thin sheets or drawn into wires. In general, metals lose
electrons in forming compounds.
Alkali Metals
With the exception of hydrogen, all of the elements on the left side of the periodic table are
metals. Group 1 elements (except for hydrogen) are called alkali metals. Since alkali metals are
highly reactive, they usually exist in nature as compounds with other elements; never as free
elements. Alkali metals are silver colored, soft metals with only one valence electron in its
valence level reflecting in low ionization of energy. Thus, alkali metals easily lose electrons. As
a result, alkali metals are well-known for their violent reactions when placed in water. As one
moves down the group, these reactions become increasingly more violent.
Figure 4. Alkali Metals
(From http://chem4kids.com/files/elem_alkalimetal.html)
To learn more about Alkali Metals, visit
http://www.chemicalelements.com/groups/alkali.html
Alkaline Earth Metals
Elements found in group 2 are as well highly reactive and similar to group 1 elements. Alkaline
earth metals are silver colored, soft metals with two valence electrons in its valence level, and are
basic. In comparison to Alkali metals, group 2 elements have higher densities and melting points.
As a result of their higher ionization of energy, they are not as reactive as alkali metals.
However, they are nonetheless still not found in nature in its elemental state. As one moves down
the group, reactivity increases among the elements in the alkaline earth family.
Figure 5. Alkaline Earth Metals
(From http://chem4kids.com/files/elem_alkalineearth.html)
To learn more about Alkaline Earth Metals, visit
http://chem4kids.com/files/elem_alkalineearth.html.
Transition and Inner Transition Metals
Transition elements are divided into transition metals and inner transition metals. Groups 3
through 12 make up the transition metals, whilst the two rows located along the bottom of the
periodic table make up the inner transition metals, more specifically the lanthanide series and
actinide series. Many of these metals play an important role in living organisms and materials.
Transition metals are located in the middle of the periodic table. Properties may vary from family
to family as in their oxidation states. However, in general, most of the metals have high densities
and melting points, silvery-blue at room temperature with the exceptions of copper, iron, and
gold, and may act as a catalyst in chemical reactions.
Figure 6. Transition and Inner Transition Metals
(From http://chem4kids.com/files/elem_transmetal.html)
To learn more about Transition Metals, visit
http://www.chemguide.co.uk/inorganic/transition/features.html
Inner transition metals, which make up the two series Lanthanides and Actinides, are able to at
ground state occupy f-orbitals. Elements in the lanthanide series are silver colored, soft metals
with the capability to lose 3 electrons, and are somewhat less reactive than alkaline earth metals
which allows for limited use structurally. They discolor easily in air and react slowly with water.
Though there is little commercial employment of these metals, their main use is in making steel
alloys.
Figure 7. Transition and Inner Transition Metals
(From http://chem4kids.com/files/elem_transmetal.html)
To learn more about the Lanthanide Series, visit
http://chemistry.about.com/od/elementgroups/a/lanthanides.htm
On the other hand, the actinide series are of importance because of their radioactivity. All
isotopes of these elements are denser, have higher melting and boiling points than the alkaline
earth metals, but similar to the lanthanide series, their reactivity allows for limited use
structurally. Nonetheless, all elements after uranium may be produced by nuclear bombardment
reactions.
Figure 8. Actinide Series
(From http://chem4kids.com/files/elem_transmetal.html)
To learn more about the Actinide Series, visit
http://chemistry.about.com/od/elementgroups/a/actinides.htm
Metalloids
Metalloids are also known as semimetals; having both physical and chemical properties of both
metals and nonmetals. Elements of this classification may be shiny or dull, and typically conduct
heat and electricity better than nonmetals.
Elements in the boron group (group 13) consist of boron as the only metalloid in this family.
Aluminum, gallium, indium, and thallium are all metals. These elements have three electrons in
their outer valence levels. The most important element in this group is aluminum, which is the
third most abundant element in the earth’s crust.
Figure 9. Semimetals
(From http://nobel.scas.bcit.ca/chem0010/unit4/4.3.3_property_semimetals.htm)
Nonmetals
Nonmetals occupy the upper right side of the periodic table. Nonmetals are generally gases and
include groups such as the carbon group (group 14), the nitrogen group (group 15), the oxygen
group (group 16), the halogens (group 17), and the noble gases (group 18). Characteristics shared
amongst nonmetals in comparison to metals reflect solid in form, dull and brittle, lower densities,
lower melting and boiling points, high electronegativty, poor conductors of heat and electricity,
and are more acidic. In general, nonmetals gain electrons in forming compounds.
Elements in the carbon group (group 14) consist of carbon as the only nonmetal, silicon and
germanium as metalloids, tin and lead as metals. All elements have four electrons in their outer
valence levels. Aside from carbon being the only nonmetal, elements in the group will often
share their electrons. As one moves down the group, the tendency to lose electrons increases as
the size of the atom increases with increasing atomic number.
Elements in the nitrogen group (group 15) consist of nitrogen and phosphorus as nonmetals,
arsenic and antimony as metalloids, and bismuth as a metal. Except for nitrogen, all elements in
the group are found to be solid at room temperature and have five electrons in their outer valence
levels. Out of all the elements in this group, nitrogen is considered the most important. At room
temperature, it is a gas, making up nearly 80 percent of the Earth’s atmosphere.
Elements in the oxygen group (group 16) consist of oxygen, sulfur, and selenium as nonmetals,
and tellurium and polonium as metalloids. All elements in group 16 have six electrons in their
outer valence levels. In this group, oxygen and sulfur are considered to be most important.
Oxygen makes up about 23 percent of the mass of air, 89 percent of the mass of water, and about
46 percent of the mass of the crustal rocks of Earth. Oxygen (O2) is an odorless, tasteless, and
colorless gas whereas sulfur has a foul odor and found as a bright yellow solid. The largest
incorporation of sulfur use occurs in the production of sulfuric acid such as lead in storage
batteries.
Elements in the halogen group consist of fluorine, chlorine, bromine, and iodine as nonmetals
and astatine as a metalloid. All elements in group 17 have seven electrons in their outer valence
levels. Group 17 is the only group to contain elements in all three states of matter at room
temperature. Most metals that react with halogens will form compounds called salts.
Elements in the noble gases consist of helium, neon, argon, krypton, xenon, and radon all as
nonmetals. At room temperature, they are all odorless, colorless, gases and are considered to be
the least reactive elements. Noble gases neither lose nor gain electrons. Elements in group 18
have their outer shell completely filled, allowing them little ability to partake in chemical
reactions.
Trends of the Periodic Table
In addition to the elements being organized based on family and their similar properties, periodic
trends for elements also include the element’s size (atomic radius), their ability to lose or gain
electrons (ionic radius), ionization energy, and electronegativity. Periodic trends reflect a general
tendency for elements to change in a predictable way as one travels down a group or across a
period.
Atomic Radius
Atomic radius is a periodic trend influenced by the element’s electron configuration. Within each
atom, there is a nucleus that is surrounded by an electron cloud. The size of the atom is
dependent upon how far away its valence electrons are from the nucleus. Thus, reflecting the
varying distances of space between the atom and its neighboring atom. For example, if the
valence electrons are close to the nucleus, the atom will be small. If they are far away from the
nucleus, the atom will be fairly large.
Figure 10. Illustration of Atomic Radius
(From http://www.chemguide.co.uk/atoms/properties/atradius.html)
Periods and Atomic Radius
Trends within periods tend to decrease in atomic radius when moving from left to right across
the periodic table. The change in atomic radius is a direct result of the nucleus’s positive charge
and movement across on the same principal energy level within a period. Regardless of whether
one moves left to right or right to left, the principal energy level for the atoms within the row
remain identical. As one moves across a period, there are no additional electrons added to the
outer valence electrons and the nucleus. The valence electrons are not protected from the
increased nuclear charge caused by the additional protons in the nucleus. Therefore, the
increased nuclear charge pulls the outermost electrons in closer to the nucleus, making the atom
smaller.
Groups and Atomic Radius
In general, atomic size increases moving down a group. The nuclear charge on the atom
increases as well as the number of electrons added to the orbitals. This is analogous to its
increase in principal energy levels. Though the increased nuclear charge on an atom in a period
becomes smaller, moving left to right; however, in a group, the atom’s radius is the exact
opposite. The atom’s outermost orbitals increase in size together with the principal energy level
as one moves down a group. Because of this increase, the atom’s size becomes larger due to the
increased space between the outer electrons and the nucleus. This increase of distance or space is
what negates the pull of the increased nuclear charge.
For more information about Trends in Atomic Radius, see
http://intro.chem.okstate.edu/1314F00/Lecture/Chapter7/Lec111300.html
Ionic Radius
Atoms are able to gain or lose one or more electrons in order to form ions. Ions are atoms that
have gained a net charge; positive or negative. Atoms become smaller when they have lost
electrons and formed positively charged ions. Electrons lost are usually the atom’s outer valence
electrons which then leave an empty orbital resulting in a smaller radius. In addition to the loss
of valence electrons, the electrostatic repulsion between the lesser amount of electrons and
positively charged nucleus declines, permitting them to be drawn closer to the nucleus.
However, when atoms gain electrons and form negatively charged ions, they become larger.
Electrons gained increases the electrostatic repulsion between the additional outer electrons,
pushing them to move farther apart from the nucleus.
Figure 11. Illustration of Ionic Radius
(From http://www.chemguide.co.uk/atoms/properties/atradius.html)
Periods and Ionic Radius
The ionic radius of most elements, moving left to right across a periodic table, will gradually
decrease in positive ions; thus, decreasing the size of the atom’s ionic radius. Yet, moving right
to left on the periodic table, until reaching group 15, the larger negative ions will begin to slowly
decrease.
Groups and Ionic Radius
An atom moving down a group results in a gradual increase in its ionic radius. This is attributed
to the ion’s outer electrons in orbitals matching the higher principal energy levels. Both positive
and negative ions increase in ionic radii moving down a group.
Figure 12. Atomic Radius versus Ionic Radius
(From http://intro.chem.okstate.edu/1314F97/Chapter8/Ionic%20Radii2.GIF)
For more information about Ionic Radius, see
http://www.chemcool.com/regents/periodictable/aim3.htm
Ionization Energy
Ionization of energy is the amount of energy needed to remove an electron from a gaseous atom.
The energy necessary to remove the first electron from an atom is called the first ionization of
energy. High ionization of energy values specifies that the atom has a strong pull on its electrons.
An atom reflecting a strong hold on its electrons is less likely to form positive ions. Similarly,
atoms with values of low ionization of energy will lose their outer electrons readily and as
expected form positive ions.
Periods and Ionization of Energy
With each additional removal of an electron, the amount of energy needed to remove one
electron from an atom increases for each successive ionization. In a period or row, when moving
across from left to right, ionization of energies increases. The increased nuclear charge of each
consecutive element produces an increased pull on the valence electrons.
Groups and Ionization of Energy
In a group or family, when moving down, first ionization of energy decreases. This is due to the
atomic size of the atom increasing as one travels down a group. As the atom becomes larger, less
and less energy is required to remove its outer valence electrons away from the nucleus.
Figure 13. Ionization of Energy
(From http://grandinetti.org/Teaching/Chem121/Lectures/PeriodicTrends/index.html)
For more information about Ionization Energy, see
http://dl.clackamas.cc.or.us/ch104-06/periodic.htm
Electronegativity
Electronegativity usually increases from left to right across a period and decreases moving down
a group. The electronegativity of an element indicates the atom’s ability to attract shared
electrons in a chemical bond. Values for electronegativity are expressed on a numerical scale of
4.0 or less. In a chemical bond, atoms with the higher electronegativity value will attract the
bond’s shared electrons. An atom’s electronegativity is correlated to its ionization of energy.
Both trends for ionization of energy and electronegativity are similar; both increase moving from
left to right across a period and decrease moving down a group.
Figure 14. Electronegativity on a Periodic Table
(From http://www.green-planet-solar-energy.com/periodic-table-picture.html)
Figure 15. Electronegativity Trend
(From www.arbuiso.com/diary/trendsdiary.pdf)
For more information about Electronegativity, see
http://www.chemguide.co.uk/atoms/bonding/electroneg.html
Summary of Properties of Metals, Nonmetals, and Metalloids







Metals
silver or white in color, lustrous (shiny)
malleable and ductile
good conductor of heat and electricity
high melting and boiling points
high densities
many react with acids
all solids except, mercury (Hg)







Nonmetals
dull in appearance
brittle if solid
poor conductor of heat and electricity
low melting and boiling points
low densities
does not react with acids
exist in all three states of matter
* Metalloids are combined properties of both metals and nonmetals.
Summary of Periodic Trends
Periodic Trends
Atomic Radius
Moving Left to Right across a
Period (Horizontal Row)
 decreases; atoms become
smaller due to an increase
in nuclear charge
 more protons equal greater
nuclear charge
Ionization of
Energy

increases; strong nuclear
charge due to atomic radius
Electronegativity

increases; tend to gain
electrons
Going down across a Group
(Vertical Column)
 increases; atoms become
larger due to a decrease in
nuclear charge
 atoms tend to lose
electrons; nuclear charge
is less
 decreases; weak nuclear
charge due to atomic
radius
 decreases; tend to lose
electrons
* Ions

Negative ions are much larger than their neutral atoms. The atomic size of negative ions
is larger due to their ability to gain one or more electrons in order to fill their orbitals.
The nuclear charge, hold on the electrons, is lessened.

Positive ions are much smaller than their neutral atoms. The atomic size of positive ions
is smaller due to their ability to lose one or more outer valence electrons. The nuclear
charge is increased and electrons are pulled inwards toward the nucleus.
Figure 16. Periodic Trends
(From http://www.tutorvista.com/content/science/science-ii/periodic-classification-elements/trends.php)
For more information about Periodic Trends, see
http://www.sciencegeek.net/Chemistry/chempdfs/PeriodicTrendsOrganizer.pdf
Properties of Matter
Density
Density is defined as the mass of an object divided by its volume; expressed in grams per cubic
centimeter. Using density as a property of matter, allows one to identify a substance. For
example, imagine both boxes in the picture below have identical volume and mass, which box
would weigh more?
Figure 17. Schematic of Density
(From http://www.nyu.edu/pages/mathmol/textbook/density.html)
Since the mass of each ball and the volume of each box are the same, the box that is filled with
more balls, weighs more. Thus, it has more mass per unit of volume; greater density. Using this
example, allows one to decipher between the characteristics and properties of any substance.
Density increases as one moves down a group which is related to its increase in atomic size.
Moving down a group considerably increases the atomic radius size due to the number of protons
and neutrons occupying the nucleus. The increased space between the outer electrons and the
nucleus increases the distance or space negates the pull of the increased nuclear charge for the
new energy levels added moving down a group. This increase in atomic radius constructs each
atom down a group to become more massive in its density. So, as one moves down a group, the
volume of the atom increases as well. How much mass an atom can hold, depends of their
volume, and their volume depends on their atomic radius. Additionally, mass of an atom
increases moving down a group.
Conductivity
Conductivity is the ability to carry an electrical current when electrons are free to move.
Conductivity of an atom is typically determined by its ionization energy (how easily electrons
can be removed) and by its electronegativity (atom’s ability to attract electrons in a chemical
bond).
Metals versus Nonmetals
Conductivity differs between metals and nonmetals. Conductivity of a metal atom decreases as
one travels from left to right across a period. However, it increases as one goes down a group. In
a period, when moving across from left to right, the amount of energy required to remove one
electron from an atom increases for each following ionization. When moving down a group, as
the atom becomes larger, less and less energy is required to remove its outer valence electrons
away from the nucleus, thus making it easier to be taken away.
Conductivity of a nonmetal atom is the exact flip-flop of the reactivity of metal atoms. As one
travels from left to right across a period, the conductivity of the nonmetal atom increases. As one
travels down a group, its conductivity decreases. Therefore, the farther right and up one goes on
the periodic table, the higher the electronegativity becomes resulting in a more dynamic
exchange of electrons for nonmetal atoms.
Metals
Period atomic number increases
ionization energy increases
conductivity decreases
Conductivity
Nonmetals
atomic number increases
electronegativity increases
conductivity increases
Group atomic number increases
ionization energy decreases
conductivity increases
atomic number increases
electronegativity decreases
conductivity decreases
Magnetism
Magnetochemistry is the study of magnetic properties. When studying periodic trends,
magnetochemistry may be classified into three distinct categories: diamagnetic, paramagnetic,
and ferromagnetic.
According to Eric Weisstein’s World of Physics
(http://scienceworld.wolfram.com/physics/Diamagnetism.html), diamagnetism is a weak
repulsion in a magnetic field generated by the current of the orbiting electron. Though all
materials possess diamagnetism, some materials may exhibit stronger inherent magnetic
properties; diamagnetism being the weakest out of the three categories. Diamagnetic atoms are
not attracted to a magnetic field because all of its electrons are paired up.
When a material has been placed in between magnetic poles and becomes attracted into the
magnetic field, it is called paramagnetism which arises from electron spins. Paramagnetism may
occur when a material has unpaired electrons in its orbitals. For example, the electron
configuration of cobalt, Co is [Ar] 4s2 3d7, and 3 of the 7 d electrons are unpaired.
Magnetic behavior experienced in everyday life is known as ferromagnetism, which is
commonly viewed as the typical form of magnetism. Ferromagnetic effects sometimes produce
magnetizations of magnitude to be larger than diamagnetic or paramagnetic effects. This may
occur when materials of ferromagnetic properties, such as iron, in an external magnetic field
become attracted and remain attracted even after the material has left the external magnetic field.
For example, iron, Fe, will form magnetic domains within the magnetic field when its atoms
align themselves to increase their ferromagnetic properties.
Similarities in properties of elements of the periodic table are a result of similar valence shell
electron configurations. Since atomic size determines outermost valence shell electrons, as the
atomic radius increases, so does the magnetic properties of the atom. Although, the principal
energy is the same for elements moving across a period, as the atomic radius decreases, the
attraction between atoms increases due to their increase in atomic mass number of protons and
electrons.
Solubility
Solubility is the amount of solute that dissolves in a given amount of solvent at a given
temperature to form a saturated solution. Solute is the substance that is being dissolved in a
solvent to form a solution whereas a solvent is the substance that is doing the dissolving in a
solution. For example, sugar dissolving in water, the sugar is the solute whereas the water is the
solvent.
Reactivity of an element’s solubility as a periodic trend is dependent upon its atomic radius. As
one moves down a group, atomic radius begins to increase, as a result, its solubility will begin to
increase as well due to the space it occupies based on atomic size. The larger the atomic radius,
the less it will fit in a given amount of solution, therefore making it more difficult to dissolve.
In addition, as one moves down and across on the periodic table, atoms will begin to share
electrons no longer separate when placed in a solution. This is a result of the atom’s ionic radius
and their ionic behavior. As one moves across a period, more electrons are transferred or pulled
off of metals, allowing for a greater net charge. For nonmetals, more of their protons will pull or
gain electrons. This effect is the same for elements moving down a group.
Content Benchmark P.8.A.2
Students know elements can be arranged in the periodic table which shows repeating patterns
that group elements with similar properties. E/S
Common misconceptions associated with this benchmark
1. Students may have the misconception that all atoms of the same element are identical.
Since Dalton was only able to probe the chemical behavior of materials, he assumed all atoms of
the same element were identical. Recall that neutrons, along with protons, make up the nucleus.
Since neutrons have no electrical charge, they do not affect the chemical behavior. What Dalton
did not know is that there can be different numbers of neutrons in atoms of the same element;
that is, the atoms are not all identical. Atoms of the same element with different numbers of
neutrons are called isotopes.
To learn more about isotopes, go to
http://www.electricianeducation.com/theory/atomic_particles.htm
2. Students may have misconceptions that electrons orbit the nucleus in fixed paths.
Atomic orbitals are the specific regions in space around the nucleus of an atom in which the
electrons are most likely to be found. Atomic orbitals are entirely described by their quantum
numbers. Each electron has a set of four quantum numbers, which specify it completely; no two
electrons in the same atom can have the same four quantum numbers; example of the Pauli
Exclusion Principle. Niel Bohr suggested that electrons are confined to specific shells which
have fixed energy levels. However, because of the Heisenberg Uncertainty Principle, the
locations of these shells about a nucleus cannot be rigidly defined. Thus, the atomic orbitals are
diffuse regions in space around the nucleus of an atom in which the electrons are most likely to
be found.
To learn more about this misconception, go to
http://www2.ucdsb.on.ca/tiss/stretton/CHEM1/elecon5.html
3. Students may confuse atomic number with numbers of valence electrons.
The atomic number of an element in the periodic system; where the elements are arranged in
order of increasing number of protons in the nucleus, represents the number of protons, which is
always equal to the number of electrons in the neutral atom, is also the atomic number.
The valence electrons are the electrons in the last shell or energy level of an atom. They show a
repeating or periodic pattern. For example, as you move across a period, you increase in number.
Then when you start the new period, the number drops back down to one and starts increasing
again.
To learn more about atomic number and valence electrons, go to
http://misterguch.brinkster.net/subatomic%20particles.html
4. Students may have misconceptions that the periodic table is unstructured.
The periodic table is an array of the elements in increasing atomic number or atomic mass,
grouped according to chemical property. The layout of the periodic table demonstrates recurring
periodic or chemical properties. Elements are listed in order of increasing atomic number. Rows
are arranged so that elements with similar properties fall into the same vertical columns (groups).
According to quantum mechanical theories of electron configuration within atoms, each
horizontal row (period) in the table corresponded to the filling of a quantum shell of electrons.
There are progressively longer periods further down the table, grouping the elements into s-, p-,
d- and f-blocks to reflect their electron configuration.
To learn more about how the periodic table is structured, go to
http://www.chemicalelements.com/
Content Benchmark P.8.A.2
Students know elements can be arranged in the periodic table which shows repeating patterns
that group elements with similar properties. E/S
Sample Test Questions
Questions and Answers to come in separate file
Content Benchmark P.8.A.2
Students know elements can be arranged in the periodic table which shows repeating patterns
that group elements with similar properties. E/S
Answers to Sample Test Questions
Questions and Answers to come in separate file
Content Benchmark P.8.A.2
Students know elements can be arranged in the periodic table which shows repeating patterns
that group elements with similar properties. E/S
Intervention Strategies and Resources
The following is a list of intervention strategies and resources that will facilitate student
understanding of this benchmark.
1. Periodic Table of Elements
Teacher Domain is a user-friendly site for all educators to use in the classroom. This particular
activity comes in both flash and adobe type media. The overview journeys through the periodic
table of the elements which holds a wealth of information about our material world. This
interactive periodic table helps unlock some of the information the periodic table contains,
establishing clear relationships among some elements and illustrating the electron configurations
responsible for each element's chemical properties.
To access the interactive activity, go to.
http://www.teachersdomain.org/resources/phy03/sci/phys/matter/ptable/index.html.
Another helpful websites, in addition to the above activities, is the Periodic Table of Comic
Books, found at http://www.uky.edu/Projects/Chemcomics/.
2. Periodic Table Trends
Ponder Chemistry Lesson index provides several activities online for students exploring the
periodic table as a web quest. The Periodic Table Trends exercise is designed for students to
discover the periodic trends of certain physical properties of elements related to their position on
the Periodic Table of Elements. In this activity, students will look at a few physical properties of
elements and examine how those properties are related to their position on the Periodic Table.
They will analyze the data found on the Periodic Table sites to answer the questions listed below.
To access the interactive activity, go to
http://www.lynchburg.net/hhs/chemistry/trends/.
Helpful websites, in addition to the above activities, include the Visual Elements Periodic Table,
found at http://www.chemsoc.org/viselements/pages/pertable_fla.htm.
And the Environmental Chemistry of Elements, found at
http://environmentalchemistry.com/yogi/periodic/#Chemical%20elements%20sorted%20by.
3. Creative Chemistry Flame Tests
Creative Chemistry is a site which provides students with a hands-on experience understanding
chemical properties of elements. The site provides a student, teacher, and technician guide in
adobe format. The flame test is used to visually determine the identity of an unknown metal of an
ionic salt based on the characteristic color the salt turns the flame of a Bunsen burner. When
solutions of metals are heated in a Bunsen burner flame, they give off characteristic colors. For
example, sodium makes the flame turn bright orange – this is the same orange color made by
sodium street lamps and many fireworks. The student can be an analytical chemist or forensic
scientist in this activity, by finding out what color flames different metals make, and working out
the identity of some unknown metal solutions. This activity could certainly be done as a
demonstration.
To access this activity, go to
http://www.creative-chemistry.org.uk/activities/flametests.htm.
And PDF format at
http://www.creative-chemistry.org.uk/activities/documents/flametests.pdf.
Helpful websites that may accompany the above activities, include the Visual Representation of
a Flame Test (requires Adobe Flash Player)
http://www.metacafe.com/watch/658019/flame_test_of_the_elements_experiment/.
And the Flame Test at http://www.chemcool.com/regents/atomicconcepts/aim3.htm.
4. The Periodic Table Challenge
Rob Toreki designed this interactive online activity for students to challenge themselves as to
how well they can complete the blank periodic table under a certain amount of time. The activity
automatically maintains their scores and resets when the time runs out. The periodic table
resembles the periodic table from http://www.chemicool.com/. This is a great activity for
students’ challenge themselves regarding the arrangement of elements in the periodic table.
To access this activity, go to
http://www.ilpi.com/genchem/instantquiz.html.
For the Challenging Classic Version, go to
http://www.ilpi.com/genchem/periodicquiz.html.
5. The Particle Adventure: Fundamentals of Matter and Force
The Particle Adventure is a constantly evolving educational project sponsored by the Particle
Data Group of the Lawrence Berkeley National Laboratory (LBNL). In an online multimedia
presentation students will take a unique tour of the atom, be introduced to the electron, proton
and neutron, and then dive into the nucleus of the atom to explore parts that most of us never
heard about. This is a great introductory activity to the atom and its atomic structure. At the end
of each module, students are posed a trivia question and allowed to view the answer in a separate
window.
To begin the particle adventure, go to
http://particleadventure.org/.
6. Atomic Musical Chairs
Middleschoolscience.com a website developed by Marc Bonem from the Science and Arts
Academy provides middle school science teachers a collection of activities for the science
classroom. Atomic musical chairs will allow the students to identify the different parts of an
atom, determine the atomic number, atomic mass, and the number of protons, neutrons and
electrons for each atom, realize that electrons are not static, but always moving, know the
relationship between the number of electrons to the type of atom, differentiate between ions and
isotopes, recognize that atoms have a numeric relationship in the periodic table, and understand
the connection between energy levels and valence electrons to the shape of periodic table.
To access the lesson plan and answer key, go to
http://www.middleschoolscience.com/atomicmusicalchairs.pdf.