Transcript The Periodic Table - Mrs Molchany`s Webpage

Ch 7
 The
Periodic Table
Objectives



SWBAT
Define the periodic properties
Apply periodic trends to the periodic table
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 (in a group)

Measure the atomic radius from the center of
the nucleus to the outermost electron.
Atom size increases going down a group.
 Reason: The principal quantum number of
the outermost (valence) electrons increases.
Atomic Radius (in a period)
Atomic size decreases going across the period.

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.
Generally speaking, effective nuclear charge is the
charge felt by the valence electrons after you have
taken into account the number of shielding
electrons that surround the nucleus. Effective
nuclear charge increases (which decreases
shielding).
Atomic Radius

http://www.shodor.org/chemviz/ionization/students/background.html
Ionic Size (when you create an ion)

An ion is created when an atom gains or loses an electron.
1)
LOSE AN ELECTRON (create a positive ion)


2)
Size becomes smaller
Reason: Loss of an electron vacates outer orbitals
and reduces the repulsive forces between electrons.
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.

Ionization Energy

1)
2)

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
HIGH ionization energy means the atom hold
onto the electron tightly.
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.
(NOTE: This trend is the opposite of the atomic radius trend)
Ionization Energy Trend
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
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 electron-electron 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 to form anions or
oxyanions in aqueous solution
Tend for form cations in
aqueous solutions
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.
Try Writing the Reaction
Sodium oxide is reacted with water.
Barium oxide is reacted with HCl.
Reaction Answers
Sodium oxide is reacted with water.
Na2O + H2O  NaOH
Barium oxide is reacted with HCl.
BaO + HCl  BaCl2 + H2O
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.
Try these Reactions.
Carbon dioxide reacts with water.
Tetraphosphorous decaoxide reacts with
water.
Sulfur trioxide reacts with barium hydroxide.
Tetraphosphorous decaoxide reacts with
sodium hydroxide.
Reaction Answers
Nonmetal oxide + water → acid
CO2 (g) + H2O(l) ---> H2CO3 (aq)
P4O10 (s) + 6 H2O(l) ---> 4 H3PO4 (aq)
Nonmetal oxide + base → salt + water
 SO3(g) + Ba(OH)2(aq) → BaSO4(aq) + H2O(l)
 P4O10(s) + 12 NaOH(aq) → 4 Na3PO4(aq) + 6
H2O(l)
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, 11, 13-17, 20, 23, 24, 27, 28, 32, 33, 35,
38, 40, 39, 41, 43, 45, 49, 54

Do 40 before 39