Download Chapter 3: Atoms and the periodic table

CHAPTER 3:
ATOMS AND THE PERIODIC
TABLE
Section 2:
A Guided Tour of the Periodic Table
Goals/Objectives

After completing the lesson, students will be able to
...
 Relate
the organization of the periodic table to the
arrangement of electrons within an atom;
 Explain why some atoms gain or lose electrons to form
ions;
 Determine how many protons, neutrons, and electrons
an isotope has, given its symbol, atomic number, and
mass number;
 Describe how the abundance of isotopes affects an
element’s average atomic mass.
Organization of the Periodic Table


The periodic table groups
similar elements together,
which makes it easier to
predict the properties of
the element based on
where it is located on the
table.
Elements are represented
by their symbols and are
arranged in an order
based on the number of
protons found in the
nucleus of its atoms.
Organization of the Periodic Table


Elements are listed in this order, because the
periodic law states that properties will occur in a
regular pattern.
Periodic Law—Properties of elements tend to
change in a regular pattern when elements are
arranged in order of increasing atomic number, or
number of protons in their atoms.
Turn to page 78-79 in the textbook.
Using the Periodic Table to Determine
Electronic Arrangement


Periods—A horizontal row of elements in the periodic table.
As you move left to right in the periodic table, not only do the
number of protons increase by one, but so do the electrons.



Protons = positive charge
Electrons = negative charge
Neutrons = neutral charge (no charge)



One way to remember that there are the same number of protons and
electrons is the fact that elements on the periodic table are neutral (their
positive charges cancel out with the same number of negative charges).
Atoms have “energy levels” where the electrons are found.
Valence electrons—An electron in the outermost energy level of
an atom.

Valence electrons determine the chemical properties of atoms.
Elements in the Same Group Have
Similar Properties


Group (Family)—A vertical column of elements in
the periodic table.
Elements in groups have the same number of
valence electrons, so they have similar properties.
 THEY
ARE NOT EXACTLY THE SAME . . . Just similar.
 They
have different number of protons in their nuclei and
different number of electrons in their filled inner energy
levels
Some Atoms Form Ions





Ionization—The process of adding electrons
to or removing electrons from an atom or
group of atoms.
Some energy levels are only partially
filled, which makes it easier for them to
gain or lose electrons in order to have a full
outer energy level.
If an atom gains or loses an electron, it no
longer has the same number of electrons as
it does protons.
Because the charges do not cancel
completely like before, the ion that forms
has a net electric charge.
Ion—An atom or group of atoms that has
lost or gained one or more electrons and
therefore has a net electric charge.
Ionization
Cation


In elements with high reactivity,
electrons are easily shared.
Cation—An ion with a positive
charge.


This indicates that an electron
(negative charge) was “given up,”
or removed from this atom.
When an element gives up an
electron, there is a positive charge
and is indicated with putting a
“positive” (plus sign) after the
chemical symbol.

Example: If lithium (Li) gave up an
electron, it would be written Li+
Anion

Anion—An ion with a negative charge.



This indicates that an electron was
“gained” to help complete the
outermost energy level.
This usually occurs within elements that
it’s easier to “fill up” the energy level
as opposed to “giving up” electrons in
order to accomplish the same goal
(stable energy levels).
When an element gains an electron,
there is a negative charge and is
indicated with putting a “negative”
(minus sign) after the chemical symbol.

Example: If Fluorine (F) gains an
electron, it would be written F-
How Do the Structures of Atoms Differ?


Atoms of different elements have their own unique
structures and, because of this, they have different
properties.
Atomic Number—The number of protons in the
nucleus of an atom.
 Since
atoms have the same number of electrons as the
protons, the atomic number indicates the number of
electrons found in the atom as well.

The atomic number for a given element never
changes.
How Do the Structures of Atoms Differ?




Mass Number—The total number of protons and
neutrons in the nucleus of an atom.
This number does not include the number of
electrons in the atom, because protons and neutrons
provide most of the atom’s mass.
Although atoms of an element always have the
same atomic number, they can have different mass
numbers.
This occurs when there is a different number of
neutrons in the nucleus.
Isotopes


Isotopes—Any atom having
the same number of protons
but different numbers of
neutrons.
If you want to refer to a
certain isotope, you write it
like this: AXZ.




X = Chemical Symbol
Z = Atomic Number
A = The number of protons and
neutrons combined.
In a normal hydrogen element, it
would be written 1H1
Isotopes (Q & A)

How many isotopes can one element have? Can an
atom have just any number of neutrons?
 No;
there are "preferred" combinations of neutrons and
protons, at which the forces holding nuclei together
seem to balance best.
 Light elements tend to have about as many neutrons as
protons; heavy elements apparently need more
neutrons than protons in order to stick together.
 Atoms with a few too many neutrons, or not quite
enough, can sometimes exist for a while, but they're
unstable.
Isotopes (Q &A)

I'm not sure what you mean by "unstable." Do atoms
just fall apart if they don't have the right number of
neutrons?
 Well,
yes, in a way.
 Unstable atoms are radioactive: their nuclei change or
decay by spitting out radiation, in the form of particles
or electromagnetic waves.
Calculating the Number of Neutrons in
an Atom

In order to calculate the
number of neutrons in
the isotopes, subtract
their atomic numbers
(found in the periodic
table) from their given
mass numbers.

Example: Carbon-14
Subtract their atomic
number (6) from the mass
number (14).
 14 – 6 = 8

The Mass of an Atom


The mass of a single
atom is very small,
which causes it to be
quite hard to work
with.
Atomic masses are
usually expressed in
atomic mass units
(amu).


Atomic Mass Unit—A
quantity equal to onetwelfth of the mass of
a carbon-12 atom.
Average Atomic
Mass—The weighted
average of the masses
of all naturally
occurring isotopes of
an element.
Summary



Elements are arranged in
order of increasing atomic
number so that elements with
similar properties are in the
same column, or group.
Elements in the same group
have the same number of
valence electrons.
Reactive atoms may gain or
lose valence electrons to form
ions.




An atom’s atomic number is its
number of protons.
An atom’s mass number is its
total number of subatomic
particles in the nucleus.
Isotopes of an element have
different numbers of neutrons,
and therefore have different
masses.
An element’s average atomic
mass is a weighted average
of the masses of its naturally
occurring isotopes.