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Transcript
The Periodic Table

J.W. Dobereiner
The elements in the triad has similar chemical properties.
 Several elements could be classified into groups of three
called triads.


Jar Newlands

Arranged the elements in order of increasing mass.
The Periodic Table

Dmitri Mendeleev
Organized the elements into the Periodic Table
 He started by sorting elements by increasing mass.
 He saw a repetition of pattern.
 Same column → same properties
 Predicted the existence of new elements


Henry Moseley

The atomic number is based on the amount of positive
charges in the nucleus.
The Periodic Law
o
When elements are arranged in order of
increasing atomic number, their physical and
chemical properties show a periodic pattern.
o
Elements with similar properties are in vertical
columns called Groups or Families.
o
Horizontal rows are called Periods.
Labeling
European notation
 American system
 IUPAC
(Inter. Union of Pure
& Applied Chemistry)





Roman numerals
IA, IIA, IIIA
1-18 (no letters)
The elements in a group have similar properties
because they have valence electrons in similar
configurations
Valence Electrons

Elements in the same column contain the same
number of valence electrons (electrons in their
outer-shell orbitals).
Atomic Radius

Measure the atomic radius from the center of the nucleus to
the outermost electron.
1)
Atom size increases going down a group.

Reason: The principal quantum number of the outermost (valence)
electrons increases.
Atomic size decreases going across the period.
2)

Reason: Same principal quantum number going across the period.
More protons are added going across the period. The protons have a
stronger pull on the electrons. The strong attractive forces between the
protons and the outermost (valence) electrons shrinks the orbitals and
makes the atoms smaller. Effective nuclear charge increases (which
decreases shielding).
Atomic Radius

http://www.shodor.org/chemviz/ionization/students/background.html
Ionic Size

An ion is created when an atom gains or loses an electron.
1)
LOSE AN ELECTRON (create a positive ion)


2)
GAIN AN ELECTRON (create a negative ion)
Size becomes larger

Reason: There are a greater number of electrons. There is a
greater repulsion force between electrons.
***Elements in a group form ions of the same charge***
Left side of periodic table forms POSITIVE ions.
Right side of the periodic table form NEGATIVE ions.



Size becomes smaller
Reason: Loss of an electron vacates outer orbitals and
reduces the repulsive forces between electrons.
Ionization Energy
This is the energy needed to remove the outermost
electron of an atom.
Li(g) → Li+(g)+ e- ionization energy 8.64 x 10-19 J/atom
1)
HIGH ionization energy means the atom hold onto the
electron tightly.
2)
LOW ionization energy means the atom holds onto the
electron loosely.

Since an atom is very small, scientists use a larger unit
of measure called the mole. Therefore, ionization
energy is measured in J/mol

Ionization Energy Trend
1)
Ionization energy decreases as you move down a group.
Reason: the electrons being removed are, in general, farther from the
nucleus. As "n" increases, the size of the orbital increases, and the
electron is easier to remove.
1)
Ionization energy increases as you move from left to right on the
periodic table.
Reason: electrons added in the same principal quantum level do not
completely shield the increasing nuclear charge caused by the added
protons. The electrons in the same principal quantum level are generally
more strongly bound when moving left to right across the periodic
table
(NOTE: This trend is the opposite of the atomic radius trend)
Ionization Energy

http://www.shodor.org/chemviz/ionization/students/background.html
Successive IE’s

The energy required to remove a second
electron from an atom is called its second IE,
and so on…
Electron Affinity (attraction)

This is the energy change that occurs when a
gaseous atom gains an extra electron.
Ne(g) + e- → Ne-1(g)


electron affinity= 29 kJ/mole
If the electron affinity is a negative number, the
atom releases energy.
Normally, non-metals have a more negative
electron affinity than metals. The exception is
the noble gases.
Electron Affinity


Electron affinity generally becomes increasingly
negative moving from left to right.
(exception: the addition of an electron to a
noble gas would require the electron to reside in
a new, higher-energy subshell. Occupying a
higher-energy subshell is energetically
unfavorable, so the electron affinity is positive,
meaning that the ion will not form)
Electron Affinity

Electron affinity does not change greatly as we move down a
group. Electron affinity should become more positive (less
energy released).
Reason: Moving down a group the average distance
between the added electron and the nucleus steadily increases,
causing the electron-nucleus attraction to decrease. The orbital
that holds the outermost electron is increasingly spread out,
however, proceeding down the group, reduces the electronelectron repulsions. A lower electron-nucleus attraction is thus
counterbalanced by lower electron-electron repulsions.

http://www.shodor.org/chemviz/ionization/students/background.html
Electronegativity
o
This is the ability for an atom to attract an
electron in a chemical bond.
o
The scale ranges from 0.7 to 4.0
o
There are no units for this number.
METALS/NON-METALS/METALLOIDS

The more an element exhibits the physical and
chemical properties of metals, the greater its
metallic character.

Metallic character generally increases going
down a column and decreases going from left to
right across a period
Metals
Non-Metals
Have a shiny luster; various
Do not have a luster; various
colors, although most are silvery colors
Solids are malleable and ductile
Solids are usually brittle; some
are hard, and some are soft
Good conductors of heat and
electricity
Poor conductors of heat and
electricity
Most metal oxides are ionic
solids that are basic
Most non-metallic oxides are
molecular substances that form
acidic solutions
Tend for form cations in
aqueous solutions
Tend to form anions or
oxyanions in aqueous solution
METALS

Metals conduct heat and electricity. They are
malleable (can be pounded into thin sheets) and
ductile (can be drawn into wire).

Metals are solids at room temperature except Hg
(which is a liquid)
METALS

Metals tend to have low ionization energies and
are consequently oxidized (lose electrons) when
they undergo chemical reaction.

Many transition metals have the ability to form
more than one positive ion.
Reactions with Metals



Metal oxide + water → metal hydroxide
Metal oxide + acid → salt + water
These reactions will be helpful to answer your
Reaction Prediction questions.
Non-Metals



Non-metals are not lustrous and are poor
conductors of heat and electricity.
Non-metals commonly gain enough electrons to
fill their outer p sub-shell completely, giving a
noble gas electron configuration.
Compounds composed entirely of nonmetals
are called molecular substances.
Non-Metal Reactions



Nonmetal oxide + water → acid
Nonmetal oxide + base → salt + water
These reactions will be helpful to answer your
Reaction Prediction questions.
Metalloids

Metalloids have properties between those of
metals and nonmetals.

Metalloids are also called semi-metals.
Group Names





Know the group (family) names:
Alkali metals
Alkaline earth metals
Halogens
Noble gases
Alkali Metals






Soft
Metallic luster
High thermal and electrical conductivity
Low densities
Low melting points
Exist in nature as compounds
Alkaline Earth Metals




Solids
Harder than alkali metals
More dense than alkali metals
Higher melting points than alkali metals
Homework Problems

4, 5, 9, 13-17, 23, 24, 27, 28, 32, 33, 35, 38, 40,
39, 41, 43, 45, 54

Do 40 before 39