Name:_______________________________ Assignment #:_______
Investigating the Periodic Table
6.1 & 6.2 Worksheet
In examining the modern periodic table, one notices the chemical elements arranged in groups and periods.
They are classified by their general physical and chemical properties into their groups: metal, metalloids,
and nonmetals, which can be further subdivided. These classifications help chemists understand the known
elements and predict the properties of new manmade elements. Let’s investigate some of these properties
more closely.
1. Use the chart to color in your periodic table and complete the “Group Number” column.
Color
Atomic Numbers
Family Name
Blue
3, 11, 19, 37, 55, 87
Alkali Metals
Red
4, 12, 20, 38, 56, 88
Alkaline Earth Metals
Yellow
21-30, 39-48, 72-80, 104-109
Transition Metals
Pink
13, 31, 49, 50, 81, 82, 83
Other Metals
Purple
5, 14, 32, 33, 51, 52, 84
Metalloids
Green
1, 6, 7, 8, 15, 16, 34
Non-Metals
Brown
9, 17, 35, 53, 85
Halogens
Gray
2, 10, 18, 36, 54, 86
Noble Gases
Orange
57-71, 89-103
Inner Transition Elements
Group Number(s)
2. What do you notice about the elements within a given group? Within a period?
3. The electron configuration of an atom reveals the placement of electrons within the orbitals of the atom
and is a key to chemical behavior. However, it is mostly the outer-shell (or valence) electrons that will
be engaged in bonding and reactivity. We can write a shorthand electron configuration that helps to
highlight these electrons through “the noble gas method.” In this shorthand, write the last noble gas
configuration the element achieved in brackets and then write the electrons that were added after that.
Examples: Sodium is [Ne]3s1 and Bromine is [Ar]4s23d104p5

Write shorthand electron configurations for the following atoms:
o
Lithium
o
Sodium
o
Potassium

What do these electron configurations have in common?

What would you expect about the relative properties of these elements?

What group are these elements located in?

How does their position on the periodic table compare to their electron
configuration?
Name:_______________________________ Assignment #:_______

Write shorthand electron configurations for the following atoms:
o
Fluorine
o
Chlorine
o
Bromine

What do these electron configurations have in common?

What would you expect about the relative properties of these elements?

What group are these elements located in?

How does their position on the periodic table compare to their electron
configuration?
4. Atoms become electrically charged by gaining or losing electrons. The typical number of electrons
gained or lost is related to electron configuration and to position in the periodic table. The goal of an
atom in forming an ion is to become isoelectronic to a noble gas. That means the ion form is trying to
achieve the stable octet (s2p6) outer-shell configuration of a noble gas.


Calcium, element 20, tends to form a 2+ ion, this means it has lost two electrons (remember, we
can’t add protons). Write the electron configuration of a Ca atom and of a Ca2+ ion.
o
Ca
o
Ca2+

Explain why the 2+ ion is the one that tends typically to form

In what group of the periodic table is Ca located?

What ions would the other elements in this group tend to form? Why do you think
that?
Sulfur, element 16, tends to form a 2- ion, this means it has gained two electrons (remember, we
can’t lose protons). Write the electron configuration of a sulfur atom and of a S2- ion.
o
S
o
S2
Explain why the 2- ion is the one that tends typically to form

In what group of the periodic table is S located?

What ions would the other elements in this group tend to form? Why do you think
that?
Name:_______________________________ Assignment #:_______
5. Putting the concepts together: Looking at the periodic table, predict the valance number (number of
electrons in the highest occupied energy level) and possible ion charge of the various groups on the
periodic table.
Group 1
Group 2
Group 13
Group 14
Group 15
Group 16
Group 17
Outershell
Configuration
Valence
Number
Ion Charge
6. Circle the correct response that would make each statement true:
a. Metals would tend to (lose/gain) electrons to form (positive/negative) ions.
b. Nonmetals would tend to (lose/gain) electrons to form (positive/negative) ions.
7. How did Mendeleev organize the periodic table? How did Moseley organize the periodic table?
8. Identify the following elements based on the clues provided:
a. [Ar] 4s23d8
b. A metal in Group 15
c. An element with 5 electrons in the third energy level
d. The halogen found in a liquid state at room temperature
e. Group 14 element in Period 4
f.
The transition metal with the smallest atomic mass
g. The alkali earth metal with the largest atomic number
h. The element in Period 4, Group 17
Group 18
Name:_______________________________ Assignment #:_______
Other things to know about the various families:
Alkali Metals
The alkali metals are very reactive metals that do not occur freely in nature. As with all metals, the alkali metals are
malleable, ductile, and are good conductors of heat and electricity. The alkali metals are softer than most other metals.
Cesium and francium are the most reactive elements in this group. Alkali metals can explode if they are exposed to
water and are stored under oil. The alkali metals have an oxidation number of +1.
Alkaline Earth Metals
The alkaline earth elements are metallic elements found in the second group of the periodic table. All alkaline earth
elements have an oxidation number of +2, making them very reactive. Because of their reactivity, the alkaline metals
are not found free in nature.
Transition Metals
The elements in groups 3 through 12 of the periodic table are called "transition metals" (previously called heavy
metals). As with all metals, the transition elements are both ductile and malleable, and conduct electricity and heat.
The interesting thing about transition metals is that their valence electrons, or the electrons they use to combine with
other elements, are present in more than one shell, so their common ion charge cannot be predicted solely by their
position on the periodic table.
Other Metals
The 7 elements classified as "other metals" are located in groups 13, 14, and 15. While these elements are ductile and
malleable, they are not the same as the transition elements. These elements, unlike the transition elements, do not
exhibit variable oxidation states, and their valence electrons are only present in their outer shell. All of these elements
are solid, have a relatively high density, and are opaque.
Metalloids
Metalloids are the elements found along the stair-step line that distinguishes metals from non-metals. This line is
drawn from between Boron and Aluminum to the border between Polonium and Astatine. The only exception to this
is Aluminum, which is classified under "Other Metals". Metalloids have properties of both metals and non-metals.
Some of the metalloids, such as silicon and germanium, are semi-conductors. This means that they can carry an
electrical charge under special conditions.
Non-metals
Non-metals are the elements in groups 14-16 of the periodic table. Non-metals are not able to conduct electricity or
heat very well. As opposed to metals, non-metallic elements are very brittle, and cannot be rolled into wires or
pounded into sheets. The non-metals exist in two of the three states of matter at room temperature: gases (such as
oxygen) and solids (such as carbon). The non-metals have no metallic luster, and do not reflect light.
Halogens
The halogens are five non-metallic elements found in group 7 of the periodic table. The term "halogen" means "saltformer" and compounds containing halogens are called "salts". All halogens have 7 electrons in their outer shells,
giving them an oxidation number of -1. The halogens exist, at room temperature, in all three states of matter: SolidIodine, Astatine; Liquid- Bromine; Gas- Fluorine, Chlorine.
Noble Gases
The six noble gases are found in group 18 of the periodic table. These elements were considered to be inert gases until
the 1960's, because their oxidation number of 0 (they don’t gain or lost electrons) prevents the noble gases from
forming compounds readily. All noble gases have the maximum number of electrons possible in their outer shell (2 for
Helium, 8 for all others), making them stable.
Inner Transition Elements
The inner transition elements are composed of the lanthanide and actinide series. They used to be called the rare
earth metals because they are difficult to find, not necessarily because they are rare. The lanthanides are so similar in
chemical and physical properties that they are difficult to separate from each other. The actinides are all unstable, and
most do not occur in nature.