Download The Periodic Table - River Dell Regional School District

Periodic Table
I. History
A. Stanislao Cannizzaro (1826-1910)
(Italian Chemist)
1. reliable method to measure atomic
masses
B. Johann Dobereiner (1780-1949)
(German Scientist)
John Newlands (1837-1898 - English)
1. related atomic mass to properties
2. Newland’s Law of Octaves
John Newlands – Law of Octaves
John Newland’s Law of Octaves
C. Lothar Meyer (1835-1895 - German)
1. properties of elements show a
repetitive pattern when they are
arranged by atomic mass
D.Dimitri Mendeleev (1834-1907- Russian)
(father of modern periodic table)
1. published system used today (1869)
2. elements arranged by increasing mass
3. left spaces for elements not yet
discovered - predicted properties
http://video.pbs.org/video/2365538111/
Dimitri Mendelev
Mendeleev’s Table
His table re-organized
video
E. Henry Mosley (1887-1915) English
1.Arrange elements
by increasing atomic
number – this led to
the-periodic law
2. Periodic Law properties of elements
are periodic functions
of their atomic number
II. Arrangement of Elements
A. Periodic Table – arrangement of
elements in order of increasing
atomic number so that elements with
similar properties are in the same
column
1. period – horizontal row (7)
2. group(family)- vertical columns (1-18)
3. periodicity – reoccurrence of similar
properties of elements in groups
C. Special Groups on the Periodic Table
Group # and Name
1 - Alkali Metals
2 - Alkaline Earth
Metals
15 - Nitrogen
Family
16 - Oxygen
Family
17 - Halogens
18 - Noble Gases
Metals – Metalloids - Nonmetals
• www.ptable.com
I. History of the Atomic Theory
F. Modern Atomic Theory
1. All matter is made up of small particles
called atoms.
2. Atoms of the same element have the same
chemical properties while atoms of different
elements have different properties
(isotopes)
3. Not all atoms of an element have the same
mass, but they all have a definite average
mass which is characteristic.
(isotopes)
I. History of the Atomic Theory
4. Atoms of different elements combine to
form compounds and each element in
the compound loses its characteristic
properties.
5. Atoms cannot be subdivided by chemical
or physical changes.
(nuclear reactions)
I. History of the Atomic Theorytime line
1803
1897
1909
1913
1935
Today
solid
particle
electron
proton
e- orbit
nucleus
neutron
Quantum
Atom
theory
Dalton
Thomson
Rutherford
Bohr
Chadwick
Schrodinger
and others
II. History of the Atomic Structure
A. J.J. Thomson (1887)
– Cathode Ray Tube
– Discovered matter contained negative
charge
– Electron
– e-
Thomson’s Model
• Found the electron
• Said the atom was
like plum pudding
• A bunch of positive
stuff (pudding), with
the electrons
suspended (plums)
II. History of the Atomic Structure
B. Robert Milikan (1909)
– Oil Drop Experiment
– Discovered mass and actual charge of
electron
– Mass is 1/1840 the mass of a hydrogen
atom
– e – has a mass of 9.11 x 10-28 g
II. History of the Atomic Structure
• So, at this point we know:
– Dalton’s Atomic Theory
– Electrons are negatively charged
– The mass of an electron is very small
II. History of the Atomic Structure
• But:
– Atoms are neutral, so there must be a positive
charge.
– Electron’s mass is so small…there must be
more to an atom.
II. History of the Atomic Structure
C. Ernest Rutherford (1909)
–
–
–
–
Gold Foil Experiment
Discovered the proton
p+
When alpha (+2) particles hit screen, the
screen lights up
– P.S.
•
#p+ = atomic number
Lead
block
Uranium
Florescent
Screen
Gold Foil
• He expected:
– The alpha particles to pass straight through
He thought this would happen:
Because
Because, he thought the mass of the positive
charge was evenly distributed in the atom
Here is what he observed:
• So, he noticed:
– Most positive alpha particles pass right
through
– Only a few were deflected
• He reasoned:
– If a + particle hit a + point on the foil, it was
repelled and deflected
• He concluded:
– Atom is mostly empty space
– Has a small, dense positive center
II. History of the Atomic Structure
• At this point in 1909, we know:
– p+ = 1.67 x 10-24 g
– e- = 9.11 x 10-28 g
– The charges are balance!
• But,
– How are the electrons arranged?
– There is still mass that is unaccounted for
II. History of the Atomic Structure
D. Niels Bohr (1913)
– Electrons orbit nucleus in
predictable paths
– Much more on him later
II. History of the Atomic Structure
E. Chadwick (1935)
– Discovers neutron in
nucleus
– Neutron is neutral
– n0
– Mass is 1.67 x 10-24 g
II. History of the Atomic Structure
• Charges balanced
• Mass accounted for
• But today, we subscribe to the Quantum
Atom Theory to describe the atomic
structure
II. History of the Atomic Structure
F. Quantum Atom Theory
• The atom is mostly empty
space
• Two regions:
– Nucleus- protons and
neutrons
– Electron cloud- region where
you have a 90% chance of
finding an electron
III. Subatomic Particles
Relative Actual
mass (g)
Name Symbol Charge mass
Electron
e-
-1
Proton
p+
+1
1amu 1.67 x 10-24
Neutron
n0
0
1amu
0
9.11 x 10-28
1.67 x 10-24
Positive-negative
attraction electric
force between
protons in one
atom and
electrons in
another atom
hold atoms
together 
chemical bond
A.
III. Subatomic Particles
Atomic number
–
The number of protons in the nucleus of an atom
–
Identifies the element
–
No two elements have the same atomic number
B. Mass number
–
The number of protons plus neutrons in the nucleus of an atom
–
Mass number is very close to the mass of an atom in amu (atomic
mass units)
–
Two atoms with the same atomic number but different mass number
are called isotopes
–
(mass number) – (atomic number) = #n 0 (number of neutrons)
C. Electrons and Ions
–
For neutral atoms, #e- = #p+
–
If there are more electrons, a negative ion forms
–
If there are less electrons, a positive ion forms
–
For now, we will work only with neutral atoms
III. Subatomic Particles
C. Electrons and Ions
– For neutral atoms, #e- = #p+
– If there are more electrons, a negative ion
forms
– If there are less electrons, a positive ion
forms
– For now, we will work only with neutral
atoms
III. Subatomic Particles
• You can never change the number of
protons and have the same element
• If you change the number of neutrons in
an atom, you get
– An isotope
• If you change the number of electrons in
an atom, you get
– An ion
III. Subatomic Particles
D. Notation
1. Nuclear Notation- how we depict isotopes
 contains the symbol of the element, the mass
number, and the atomic number
Mass
number
# of P +N
Atomic
number
# of P
X
III. Subatomic Particles
D. Notation
1. Nuclear Notatioin
23
Na
11



How many protons?
How many neutrons?
How many electrons?
III. Subatomic Particles
2. Hyphenation Notation
– Symbol or name of element – mass number
Fluorine-19
Protons? Neutrons? Electrons?
C-12
Protons? Neutrons? Electrons?
III. Subatomic Particles
E. Average Atomic Mass
– Measured in grams (for a lot)
– Measured in amu (for a few)
Atomic Mass
• Mass of an atom
• Too small to measure in grams
• Use relative mass (amu)
• Almost the same as mass number
• Standard: 1 amu is defined as 1/12 the
mass of one C-12 atom
III. Subatomic Particles
E. Average Atomic Mass
– Weighted average mass of all known
isotopes
– Weighted means that the frequency of an
isotope is considered
– blackboard
• 1. Horizontal rows are called periods
– A. There are 7 periods
• 2. Vertical columns are called
groups.
– a. Elements are placed in columns by
similar properties.
• i. Also called families
1A
• 3. The elements in the A groups
are called the
2A
representative elements
3A 4A 5A 6A 7A
8A
0
4. The group B are called the
transition elements
a. These are called the inner
transition elements and they
belong here- lanthanide actinide
series
5. Group 1A are the alkali metals
6. Group 2A are the alkaline earth metals
7. Group 7A is called the Halogens
8. Group 8A are the noble gases
• The part of the atom another atom sees is
the electron cloud.
• More importantly the outside orbitals.
• The orbitals fill up in a regular pattern.
• The outside orbital electron configuration
repeats.
• The properties of atoms repeat.
Classification of the elements
A. Periodic Table- arrangements of elements
so that elements with similar properties are in
the same column
1. Valence electrons- electrons that
exist on the highest principal energy level of an
atom
a. these are the e- that can be
lost or gained in formation of a compound
• 2. energy level and period
a. the energy level of valence electrons is
indicated by which period it is found in.
Ie: period 4 elements have valence electrons on
the 4th energy level
Lets work through a few.
• 3. Valence electrons and group number
– a. Representative element’s group number
and number of valence electrons it contains
also are related.
– Ie: Na has one e- on valence shell
– Lets work on some
D. Periodic Table Showing s,p,d,f Blocks