Download The Modern Periodic Table

Prior Knowledge:
An electron configuration tells us where the electrons are in an atom.
Example:
Ca: 1s22s22p63s23p64s2
↓
2
1s
↓
Energy level the
orbital are in
How many e- in the orbital
Shape of the orbital
Valence Electrons
(Outermost electrons/Why we learn e- configurations)
Notice that in electron configurations, the last two s & p orbitals
are in the furthest energy level from the nucleus.
If you add up the electrons in the outer most s & p orbitals you get
a number between 1-8
These electrons are “Valence Electrons” (e-’s in the outer-most
e.l.)
Valence Electrons:
- are responsible for chemical bonding
- All of the other electrons are called “core electrons”.
- Help tell us see some physical and chemical properties
- Helps predict the charges on ions
Valence Electrons:
Counting Valence Electrons
Group A # = number of valence electrons
(only exception Helium = 2 e-’s)
Examples: Ca = 2 e-’s
Nitrogen = 5 e-’s
Argon = 8 e-’s
Si: 1s22s22p63s23p2
outer most energy level: 3
s & p orbital e-’s: 2 and 2
Valence electrons: 2+2= 4
Cl:
outer most energy level:
s & p orbital e-’s:
Valence electrons:
Octet Rule (and ions)
The s & p orbitals allow for a total of 8 electrons. When the
outer-most s & p are full (8e-) the atom is stable.
-atoms will gain or lose electrons in order to reach a “full
octet”
(this makes the atoms an ion)
Predictable Ions (Groups 1A-8A):
1. Locate the nearest noble gas
2. Count how many “places” it is away
3. If you count backwards to the left you have a + charge
if you count forward, you have a negative charge
Practice Problems: How many electrons are gained or lost when
forming an ion from the following elements?
a) Magnesium: ____ (gained or lost) b) Iodine: ____ (gained or
lost)
c) Gallium:____ (gained or lost)
d) Boron:____ (gained or lost)
Practice Problems: How many electrons are
gained or lost when forming an ion from the
following elements?
1. Magnesium: ____ (gained or lost)
2. Iodine: ____ (gained or lost)
3. Gallium:____ (gained or lost)
4. Boron:____ (gained or lost)
Expressing Valence Electrons
Electron dot structures represent electrons as dots located
around the symbol of the element in the pattern shown below.
4
7
3
8
X
6
Examples: Nitrogen =
N
1
5
2
Hydrogen =
(important exception.... Helium = He )
H
Valence Electrons On The Periodic Table
1
8
2
3 4 5 6 7
The Modern Periodic Table
Dmitri Mendeleev constructed the 1st periodic table
Mendeleev:
-left blank spaces for “missing elements”.
Later when
these elements were discovered, the gaps could be
filled.
- arranged the elements in columns and rows
according to their properties. Elements with similar
properties were in the same horizontal row.
-was able to accurately predict the properties of the
missing elements based on the properties of the
elements in similar rows.
-
ordered the elements by increasing atomic mass
The Modern Periodic Table
- In 1913, Henry Moseley determined the atomic
number, (# of p+), of the elements.
- He then arranged the elements in the
periodic table by increasing atomic number.
- This switched the position of some elements.
This is how the modern periodic table is
arranged today.
-In 1942 Glenn T. Seaborg rearranged the
periodic table to include the 5f row. He won a
Nobel Prize after publishing his new periodic
table.
Features of the Modern Periodic Table
Periodic repetition of chemical and physical
properties of the elements when they are
arranged by increasing atomic number is called
periodic law.
The Modern Periodic Table
Columns of elements are called groups.
Rows of elements are called periods.
Elements in groups 1,2, and 13-18 possess a
wide variety of chemical and physical properties
and are called the representative elements.
Elements in groups 3-12 are known as the
transition metals.
The Modern Periodic Table
Elements are classified as metals, non-metals,
and metalloids.
Metals are elements that are generally
• shiny when smooth and clean
• solid at room temperature
• good conductors of heat and electricity
• Malleable
The Modern Periodic Table
Alkali metals are all the elements in group 1
except hydrogen, and are very reactive.
Alkaline earth metals are in group 2, and are also
highly reactive.
The transition elements are divided into transition
metals and inner transition metals.
The two sets of inner transition metals are called
the lanthanide series and actinide series and
are located at the bottom of the periodic table.
The Modern Periodic Table
Non-metals are elements that are generally:
• gases or brittle
• dull-looking solids
• poor conductors of heat and electricity.
Group 17 is composed of highly reactive elements
called halogens.
Group 18 gases are extremely unreactive and
commonly called noble gases.
The Modern Periodic Table
Metalloids have physical and chemical properties
of both metals and non-metals:
• Semiconductors
-(which can be used in computer chips)
Examples: Silicon & Germanium
Figure 11.35: Classification of elements as
metals, nonmetals, and metalloids.
Parts of the Periodic Table
Trends on the Periodic Table
The Periodic Table is organized so that it
follows many trends:
Trends you have
learned:
Trends you will learn in the next few
slides:
•Increasing Atomic
Number
•Increasing valence
electrons across row
•Groups have similar
properties
•Atomic Radius
•Atomic Radius vs Ion radius
•Ionization Energy
•Electronegativity
Trends: Atomic Radius
Atomic size is a periodic trend influenced by
electron configuration.
Atomic radius is the
distance from the
nucleus to the
outer most electrons.
Atomic Radius
There is a general decrease in atomic radius from left to
right:
• caused by increasing positive charge in the nucleus.
• As you add more electrons and protons in the same
energy level, it causes a stronger attractions which
shrinks the e- cloud.
As you move down a group the size generally increases:
• Valence electrons are not shielded from the increasing
nuclear charge because no additional electrons come
between the nucleus and the valence electrons.
Atomic Radius
Atomic radius generally increases as you move down a group.
The outermost orbital size increases down a group, making the
atom larger.
Figure 11.36: Relative atomic sizes for selected atoms.
Ionic Radius
An ion is an atom or bonded group of atoms with
a positive or negative charge.
When atoms lose electrons and form positively
charged ions, they always become smaller for
two reasons:
1.The loss of a valence electron can leave an
empty outer orbital resulting in a small radius.
2.Electrostatic repulsion decreases allowing the
electrons to be pulled closer to the radius.
Ionic Radius
When atoms gain electrons, they can become larger,
because the addition of an electron increases
electrostatic repulsion.
Ionic Radius
Both positive and negative ions increase in size
moving down a group.
Ionization Energy
Ionization energy is the energy required to remove the
outer most electron in an atom.
•The energy required to remove the first electron
is called the first ionization energy.
•As you move down a group less energy is
required to remove electrons. Therefore,
ionization energy decreases.
It is easier to remove an electron if it is farther away
from the nucleus. These e-’s are not as attracted to the
nucleus. In general, the larger the atom, the less
attracted it is to its e-’s.
Ionization Energy
The ionization at which the large increase in
energy occurs is related to the number of
valence electrons.
• First ionization energy increases from
left to right across a period.
• First ionization energy decreases down
a group because atomic size increases
and less energy is required to remove
an electron farther from the nucleus.
Ionization Energy
Ionization Energy
Ionization Energy Discrepancies
When an orbital is half full the ionization energy
increases a little bit.
When an orbital is full the ionization energy
increases
When the energy level is full the ionization
energy is the highest
Ionization Energy
The octet rule states that atoms tend to gain, lose
or share electrons in order to acquire a full set
of eight valence electrons.
The octet rule is useful for predicting what types
of ions an element is likely to form.
Electronegativity
The electronegativity of an element indicates its relative
ability to attract electrons in a chemical bond.
Electronegativity decreases down a group and increases
left to right across a period.
As we move down a group, electronegativity decreases.
-This occurs because the electron is farther away
from the nucleus.
As we move across a period, electronegativity
increases.
-This is because the atoms get smaller moving
across the period and the electrons are more strongly
attracted to the nucleus