CHAPTER 6 THE PERIODIC TABLE
6.1 Organizing the Elements
6.2 Classifying the Elements
6.3 Periodic Trends
6.1 ORGANIZING THE ELEMENTS
1. How did chemists begin to organize the known elements?
2. How did Mendeleev organize his periodic table?
3. How is the modern periodic table organized?
4. What are three broad classes of elements?
Searching for an Organizing Principle:
1. Only about 13 elements had been identified by 1700, these
included copper, silver, and gold followed by hydrogen, nitrogen,
2. The problem was not only how to isolate the elements but how to
organize them to establish some type of pattern.
3. In 1829, J.W. Dobereiner published a classification system in
which elements were grouped in triads.
4. A triad is any of several sets of three chemically similar
elements, the atomic weight of one of which is approximately
equal to the mean of the atomic weights of the other two.
5. Not all the elements fit into the triad system.
6. Such triads—including chlorine-bromine-iodine, calcium-
strontium-barium, and sulfur-selenium-tellurium
Mendeleev’s Periodic Table
1. In 1869 Mendeleev and then Meyer published nearly identical charts
2. Mendeleev used cards to represent elements and began moving the
cards so that the elements were grouped according to their increasing
3. Mendeleev noticed periodic trends in the elements present and he
predicted properties of missing elements.
4. As predicted elements were isolated it convinced scientists that
Mendeleev’s periodic table was a powerful tool.
The Periodic Law
1. Mendeleev experienced some problems with the organization od
his chart based on his methodology – arrangement by mass.
2. In 1913 Henry Moseley determined the atomic number of each
element based on the protons in the atom.
3. The modern periodic table is arranged in order of increasing
4. There are seven rows or periods on the periodic table.
5. Each period corresponds to a principle energy level.
6. There are more elements in higher number periods because there
are more orbitals in higher energy levels.
7. Elements are also arranged in groups or columns of elements with
8. The properties of elements within a period change as you move
across the periodic table.
9. The pattern of properties that develops in the elements as they
cross the chart repeats itself for each energy level.
10.This periodic repetition of properties is know as periodic law.
11. Periodic law states; “When elements are arranged by increasing
atomic number, there is a periodic repetition of their physical and
Metals, Nonmetals, and Metalloids
1. Each group in the periodic table has three labels
a. In the U.S. Numbers and letters signify the column – 1A or 2B
b. In Europe, Roman numerals and letters are used – VB or IIA
c. The International Union of Pure and Applied Chemistry number
the groups 1 to 18 from left to right as you cross the chart.
2. Elements are can be grouped into three broad classes based on their
a. 80 % of the elements found mostly on the left half of the chart
b. Good conductors of heat and electricity
c. Have a shiny luster or sheen
d. Good ability to reflect light
e. Solid at room temperature except Hg
f. Ductile – can be drawn into wire
g. Malleable – can be hammered into sheets
a. Found in the upper right portion of the periodic table.
b. A greater variation in physical properties
c. Most are gasses are room temperature.
d. Some are solid like Sulfur and Phosphorus
e. Bromine is a dark red liquid
f. Poor conductors except for Carbon
g. Solids tend to be brittle
h. Vary in color
i. Not malleable or ductile
a. Form a stairway from Boron to Astatine
b. Properties reflect those of metals and nonmetals
c. The properties of metalloids can be controlled giving them
many technological applications
Homework – Questions p 160, Handout 6.1, and Journal
6.2 Classifying the Elements
1. What type of information can be displayed in a
2. How can elements be classified based on their electron
Squares in the Periodic Table
1. The periodic table displays the symbols and names of the
elements as well as the atomic number, atomic mass, electron
configuration, crystalline structure and group based on its
2. Groups that are generally color coded include;
a. 1A Alkali Metals
b. 2A Alkali Earth Metals
c. 7A Halogens
d. 1B to 8B the Transition Metals
e. 8A Noble Gases
Electron Configurations in Groups
1. Elements can be sorted into noble gases, representative
elements, transition elements, or inner transition metals based
on their electron configuration
2. Noble Gases
a. Are Group A8 elements
d. The s and p sublevels are completely filled except He which is
full with an s sublevel.
3. The representative Elements
a. Is the portion of the table from 1A to 7A.
b. They are referred to representative elements because they
display a wide range of physical and chemical properties.
c. They include metals, nonmetals, and metalloids
d. Most are solids, some gases, and a liquid
e. The s and p sublevels of the highest occupied energy levels are
f. For any representative element, its group number equals the
number of electrons in the highest occupied level.
1. Elements of the B Group separate the A groups
2. Elements in the B groups provide a connection between the two
sets of representative elements.
3. There are two type of transition elements based on their
a. Transition Metals
b. Inner Transition Metals
4. Transition Metals
a. Are group B elements displayed in the main body of the
b. Examples – Cu, Au, Ag, Fe
c. The highest occupied s sublevel and a nearby d sublevel contain
d. There are some exceptions in writing electron configurations
found in 6B and 1B
5. Inner Transition Metals
a. Appear below the main body of the periodic table
b. the highest occupied s sublevel and the nearby f sublevel
generally contain electrons.
c. Are characterized by f orbitlas
Blocks of Elements
1. If electron configuration and the position of the element are
considered simultaneously, another pattern emerges.
2. The periodic table can be divided into sections or blocks that
correspond to the highest occupied sublevel.
3. S block – 1A and 2A
4. P block – 3A, 4A, 5A, 6A, 7A, 8A
5. D block – Transition elements
6. F block – Inner transition elements
7. D sublevels have a principal energy level that is one less than the
8. F sublevels have a principal energy level that is two less than the
Homework – Question #10 to 15 p. 167, Review sheet 6.2, and
6.3 PERIODIC TRENDS
1. What are the trends among the elements for atomic
2. How do ions form?
3. What are the trends among the elements for the
first ionization energy, ionic size, and
4. What is the underlying cause of periodic trends?
Trends in Atomic Size
1. Molecules –a neutral group of atoms joined together by covalent
a. examples – H2 , O2, N2, F2, Cl2, Br2, I2
2. Atoms in each molecule are identical
3. Distances between the nuclei of molecules can be used to estimate
the size of the atoms
4. Atomic Radius – Is one half the distance between the nuclei of two
atoms of the same element when the atoms are joined.
5. Atomic radius is measured in picometers
6. In general, atomic size increases from top to bottom within a
group and decreases from left to right across a period.
Group trends in atomic size:
1. The atomic radius within a group increases as the atomic number
2. As the atomic number increases in a group so does the charge.
3. As the atomic number increases in a group so does the number
of occupied energy levels
4. How atomic size can be affected;
a. An increase in positive charge draws electrons closer to the
b. An increase in occupied orbitals shields the outer electrons
from being drawn to the nucleus and creates a larger atom
Periodic trend in atomic size:
1. Atomic size decreases from left to right
a. Each element has one more proton and one more electron
b. Electrons are added to the same principal energy level
c. The shielding effect is constant for all the elements in the period
d. The increasing nuclear charge pulls the electrons in the highest
energy level closer to the nucleus and decreases atomic size.
1. Some compounds are composed of particles called ions.
2. An ion is an atom or group of atoms that has a positive or
3. A neutral atom has equal numbers of electrons and protons
4. A cation is a positively charged atom with more protons then
5. Cations are atoms that tend to lose electrons
6. An anion is a negatively charged atom with more electrons then
7. Anions are atoms that tend to gain electrons
Trends in Ionization Energy:
1. Ionization energy – is the energy needed to remove an electron
from an atom.
2. The energy is measured when an element is in its gaseous state.
3. First ionization energy – the energy required to remove the first
electron from the atom.
4. The cation produced by removing one electron is +1
5. First Ionization energy tends to devrease from top to bottom
within a group and increases from left to right across a period.
6. Ionization energy can help predict what ions elements will form.
7. Ionization energy increases as more electrons are removed to
produce cations with +2 or +3 charge.
Group Trends in Ionization energy:
1. In general first ionization energy decreases from top to
bottom within a group
2. As the size of an atom increases, the nucleus has a smaller
effect on the higher level electrons, therefore it requires
less energy to remove the electrons.
Periodic trends in Ionization energy:
1. Ionization energy tends to increase from left to right.
2. Because the nuclear charge increases while the shielding
effect remains constant, there is an increase in the
attraction of the nucleus for an electron. Therefore, there
is an increase in the energy to remove the electrons from
Trends in Ionic Size:
1. In ionic bonds between metals and nonmetals, metals
tend to lose electrons while nonmetals gain electrons.
2. Cations are always smaller than the atoms from which
3. Anions are always larger than the atoms from which they
4. Periodic trend reveals a decrease in size of cations
followed by a decrease in size of anions.
5. Group trends reveal that there is a increase in ionic size
down the group.
Trends in electronegativity:
1. Electronegativity – is the ability of an atom of an element to
attract electrons when the atom is in a compound.
2. Ionization energy is used to calculate electronegativity.
3. Electronegativity can be used to determine the type of bond that
will form during a reaction.
4. Electronegativity values can be represented in a table, however the
noble gases do not appear because they are not very reactive.
5. The values are expressed in units called Paulings.
6. Electronegativity values decrease from top to bottom within a
7. For representative elements, the values tend to increase from left
to right across a period.
8. Electronegativity values among the transition elements are not
9. The least electronegative element is Cesium at 0.7
10.The most electronegative element is Fluorine at 4.0
Homework – Questions 16 to 22 p. 178, Handout 6.3, Journal