Transcript Periodic table development and trends

VCE CHEMISTRY-UNIT 1
PERIODIC TABLE
PART 1
DEVELOPMENT OF THE PERIODIC TABLE
MENDELEEV &
THE PERIODIC TABLE
Mendeleev discovered that the chemical
properties of elements follow a periodic pattern as
their relative atomic weights increase.
This pattern is called periodicity.
He noticed as the atomic weights increased
across a period, the chemical properties gradually
changed.
Development of the Periodic table
In 1869 he created the periodic table by:
• arranging elements with similar chemical
properties into vertical columns (groups)
• arranging the elements in order of
increasing relative atomic mass (periods)
Development of the Periodic table
Mendeleev left gaps in the periodic table where he
believed there should be elements that hadn’t
yet been discovered.
He was able to predict their chemical properties
by the periodic pattern.
Three predicted elements (Gallium,
Germanium, and Scandium) were discovered
20 years later.
MENDELEEV’S
PERIODIC TABLE
Today’s table is ordered according to atomic
number and electron configuration.
HOWEVER.. since chemical properties result
from electron configuration the modern table
and Mendeleev’s table are the same
Mendeleev’s Periodic Law
The properties of elements vary periodically with
their atomic weights.
William Ramsay and the
Noble Gases
• Ramsay was first to isolate the noble gases
from the atmosphere (helium argon, neon,
krypton, and xenon)
• Contributed to the discovery that helium is a
product of the atomic disintegration of
radium.
GLENN SEABORG AND THE
TRANSURANIUM ELEMENTS
Seabourg discovered many radioactive elements such as:
•californium,
•einsteinium,
•fermium,
•mendelevium
•nobelium
•plutonium
•americium
•curium
•berkelium
Generally called the transuranium elements.
Task
Write a short summary in your books of the
contribution to the development of the
Periodic Table by :
•
Mendeleev
•
Ramsay
•
Seaborg
VCE CHEMISTRY-UNIT 1
PERIODIC TABLE
Part 2: Understanding
the Trends in the
periodic table
Periodic Properties of the
Elements
• We now know that the periodic properties are
due to the electronic structure of atoms.
• Electronic structure explains the observed
trends in
– Atomic size
– Metallic/non-metallic character
– Chemical reactivity
Al
Zn Ga
Cd In Sn
RIM vs RAM
• RIM (Ir ) uses the mass spectrometer to
calculate the mass of the isotope
• RAM (Ar ) is an average of the masses of
all of the isotopes of that element naturally
found in any mixture (p.55-56 of text)
Group I elements
• Group I elements (minus hydrogen) are
called the Alkali metals.
• They are very reactive elements.
Eg: they reactive violently with water,
increasingly down a group
Hydrogen is placed in Group I as it has 1
electron in it’s outer shell but does not
share similar properties
Group II elements
• Group II elements are called Alkaline Earth
metals.
• They are not as reactive as Group I
elements.
• They tend to form bases when mixed with
H 2O
Electron Shells
• The inner, completed electron shells are
called core shells. The electrons in these
shells are called core electrons.
• The outermost electron shell is called the
valence shell. The electrons that sit in this
shell are called valence electrons.
Core Charge
• The higher the core charge, the greater
the ‘pull’ of the valence electrons to the
nucleus – this decreases the radius
• A greater ‘pull’ draws the cloud of
electrons further in, thus reducing the
diameter.
Atomic Size
The radius of an
atom is found
from the distance
between nuclei in
a molecule
http://www.mhhe.com/physs
ci/chemistry/essentialchemis
try/flash/atomic4.swf
ATOMIC SIZE -DOWN A
GROUP
EXPLANATION
• Number of shells ……………. Down a
group
• Electrons are ……….…... away from
the nucleus
• Size of atoms therefore …………..
down a group
Ionization Energy
Ionization energy - The minimum
amount of energy required to remove
the outermost electron
The greater the amount of electrons in the
shell, the harder it is to remove one
• Shells are described in terms of the
energy required to remove an electron
Electronegativity
Definition
The ability of an atom to attract electrons.
• The more electrons in the valence shell,
the greater the ability to attract more
electrons.
Metals and Nonmetals
Metals are characterized by low ionization energy
and high electrical conductivity; Non-metals by
high ionization energy and poor electrical
conductivity.