Transcript The Periodic Table

Unit 6 – The Periodic Table
Origins of the Periodic Table
By the year 1700, only 13 elements had
been identified
 Scientific discovery led to a higher rate
of element discovery
 A logical organization of elements was
needed for all the new elements

Early Organization

J.W. Dobereiner (1829) organized
elements in triads
 Triad – three elements with similar
properties (ex: Cl, Br, I)

J.R. Newlands (1864) organized
elements in octaves
 Octave – repeating group of 8 elements
Mendeleev
Dmitri Mendeleev (1869) arranged
elements according to their
properties
 Mendeleev noticed that when the
elements were arranged in order
of increasing atomic mass, there
was a repeating pattern to their
properties
 This is known as Periodicity
 Mendeleev left some spaces on
his table blank, but was able to
predict the properties of the
unknown elements

Mendeleev’s Periodic Table
Moseley
Mendeleev’s table was imperfect
– Te and I had to be reversed
 Henry Moseley (1913) arranged
elements according to atomic
number
 The periodic repetition of
chemical and physical properties
when elements are arranged by
atomic number is known as the
Periodic Law

Modern Periodic Table
The modern periodic table consists of
Rows and Columns
 Rows 
 Horizontal
 Also known as Periods
 Numbered 1-7

Columns  Vertical
 Also known as Groups and Families
 Numbered 1-18
Classifying Elements
The elements on the periodic table can
be simply classified by groups
 Groups 1,2,13-18 (1A-8A) are known as
the Representative Elements

Classifying Elements

Groups of representative elements have
the same valence electrons and
Oxidation State

Oxidation State is how many electrons are
gained or lost by an atom in a chemical
reaction
 Lost Electrons = Positive Oxidation State
 Gained Electrons = Negative Oxidation State

Think of Oxidation State as the charge of
the ion
Driving Force
 Full
Energy Levels require lots of
energy to remove their electrons.
 Noble Gases have full orbitals.
 Atoms behave in ways to achieve
noble gas configuration
Classifying Elements

Groups 3-12 (3B-2B) , as well as the
lanthanide and actinide series are
known as Transition Metals
Metals
The most common class of elements is
Metals
 Metals become cations

 What is a cation? How are they formed?
○ Positively charged atom/positive oxidation
state - Lose electrons

Metals are generally solid (except Hg),
conductive of heat and electricity,
malleable, ductile, and shiny
Alkali Metals
Group 1 elements are known as
Alkali Metals
 Alkali metals include Li, Na, K,
Rb, Cs, Fr
 Alkali metals are generally dull,
soft, and reactive – rarely found
as free elements

Alkali Metals

Write the noble gas configuration for each
Alkali Metal

[He]2s1

[Ne]3s1

[Ar]4s1

[Kr]5s1

[Xe]6s1

[Rn]7s1
How many valence
electrons do all Alkali
Metals have?
One
What is the
oxidation state of
Alkali Metals?
+1
Alkaline Earth Metals
Group 2 elements are known as
Alkaline Earth Metals
 Alkaline earth metals include Be,
Mg, Ca, Sr, Ba, and Ra
 Alkaline earth metals are harder,
denser, and stronger than alkali
metals
 Less reactive than alkali metals,
but still rarely found as free
elements

Alkaline Earth Metals

Write the noble gas configuration for each
Alkaline Earth Metal

[He]2s2

[Ne]3s2

[Ar]4s2

[Kr]5s2

[Xe]6s2

[Rn]7s2
How many valence
electrons do all
Alkaline Earth Metals
have?
Two
What is the oxidation
state of Alkaline
Earth Metals?
+2
Transition Metals
Elements in groups 3-12 (3B-2B) are
known as Transition Metals
 Transition metals include Mn, Fe, Ag,
Au, Mo, etc.
 Transition metals fill in the d orbital and
often have multiple oxidation states
 Lanthanide and Actinide Series
elements fill in the f orbitals – known as
inner transition elements

Metalloids
Elements that border
the staircase on the
periodic table are
known as Metalloids
 Metalloids include: B,
Si, Ge, As, Sb, Te, Po,
At
 Metalloids have
properties of both
metals and nonmetals

Nonmetals
Nonmetals are found to the right of the
staircase on the periodic table
 Nonmetals generally become anions

 What is an Anion? How are they formed?
○ Negatively charged atom/oxidation state Gain electrons
Nonmetals are often gases or dull, brittle
solids
 Nonmetals generally show poor
conductivity, ductility, and malleability

Halogens
Group 17 elements are known as
Halogens
 Halogens include F, Cl, Br, and I
 Halogens are the most reactive
nonmetals – often found in
compounds

Halogens

Write the noble gas configuration
for each Halogen

[He]2s22p5

[Ne] 3s23p5

[Ar] 4s23d104p5

[Kr] 5s24d105p5
How many valence
electrons do all
Halogens have?
Oxidation State?
Seven / -1
Why are the Halogens the
most reactive non-metals?
They are 1 electron
short of having an
octet.
Noble Gases
Elements in group 18 are known
as Noble Gases
 Noble Gases include He, Ne, Ar,
Kr, Xe, Rn
 Noble gases are extremely
unreactive

Noble Gases

Write the electron configuration for each
Noble Gas

1s2

[He]2s22p6

[Ne]3s23p6

[Ar]4s23d104p6

[Kr]5s24d105p6

[Xe]6s25d106p6
How many valence
electrons do all
Noble Gases have?
Eight
Why are Noble Gases
so unreactive?
They contain a full octet
– atoms gain/lose
electrons to achieve
noble gas notation
Other Groups

All other groups can be identified by the
top most element in that group.
 Ex: Group 15 can be called the Nitrogen





Group
Oxidation State: -3
Q: What is another name for Group 16?
A: Oxygen group
Q: Oxidation State
A: -2
Periodic Trends

The elements on the periodic table show
repeating trends related to electron
configuration
What is the trend for Oxidation
State?
Atomic Radius
The Atomic Radius is ½ the
distance between nuclei of
bonded atoms from the same
element
 Atomic radius decreases from
left to right across a period
 Atomic radius increases from
top to bottom in a period

Why?
Not changing energy level, but increasing
nuclear force (more positive charge in
nucleus)
Ionization Energy

If an atom is becoming an ion, it is
gaining or losing electrons in an effort to
have an octet (8 valence electrons)
Ionization Energy

The energy required to remove an
electron from an atom is called
Ionization Energy
Ionization Energy
1st Ionization Energy- energy required to
remove 1st electron from an atom
 2nd Ionization Energy- energy required to
remove 2nd electron from an atom

 2nd Ionization Energy is ALWAYS higher than
the 1st

3rd Ionization Energy- energy required to
remove 3rd electron from an atom
 3rd Ionization Energy is ALWAYS higher than
the 1st or 2nd
Ionization Energy
IE Decreases as you move down a
group
 Why?
 Electron is further away

Ionization Energy
IE Increases as you move across a
period
 Why?
 You are in the same energy level but
have more nuclear charge

Ionization Energy
 Full
Energy Levels require lots of
energy to remove their electrons.
 Noble Gases have full orbitals.
 Atoms behave in ways to achieve
noble gas configuration.
Ionization Energy
Write the electron configuration for Be
 1s22s2
 How many valence electrons does Be
have?
2
 Why is the ionization energy low?
 It is easier for Be to lose those 2 valence
electrons than it is to gain 6. Therefore, it
has a low ionization energy.

Ionization Energy
Move across the period. Write the electron
configuration for F.
 1s22s22p5
 How many valence electrons does F have?
7
 Why is the ionization energy high?
 It is easier for F to gain 1 valence
electron than is it for it to lose 7.
Therefore, its’ ionization energy (energy
to lose an electron) is high

Ionization Energy
Electron Affinity
Electron affinity is the energy change
associated with adding an electron to a
gaseous atom.
 Easiest to add to group 7A (halogens).
 Why?

 Gets them to full octet.
Increase from left to right: atoms become
smaller, with greater nuclear charge.
 Decrease as we go down a group.

Ionic Size
Cations are smaller than the atoms
from which they form (less electrons)
 Anions are larger than the atoms from
which they form (more electrons)

Ionic Size
 Across
the period, nuclear charge
increases so they get smaller.
 Energy level changes between
anions and cations.
Li1+
B3+
Be2+
C4+
N3-
O2-
F1-
Electronegativity
Electronegativity is the ability for an
atom to attract electrons in a compound
 Electronegativity increases from left to
right in a period
 Electronegativity decreases from top to
bottom in a group

Electronegativity
We do not consider noble gases when
talking about electronegativity because
they do not bond.
 What is the most electronegative
element?

 Fluorine
Electronegativity
Write the electron configuration for Li
 1s22s1
 How many valence electrons does Li
have?
1
 Why is the electronegativity low??
 It is easier for Li to lose 1 valence
electrons than it is to gain 7. It has a low
electronegativity because it would be
difficult for Li to attract 7 electrons

Electronegativity



Move across the period. Write the electron
configuration for O.
 1s22s22p4
How many valence electrons does O have?
6
Why is the electronegativity high?
 It is easier for O to gain 2 valence electrons
than is it for it to lose 6. Electronegativity is
high because it can gain electrons more
easily than it can lose them.
Electronegativity
Ionization energy, Electronegativity,
and Electron Affinity INCREASE
Atomic size increases,
Ionic size increases