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ATOMIC STRUCTURE AND
THE PERIODIC TABLE
DEVELOPMENT OF THE ATOMIC THEORY
THE ATOM
• Democritus: around the 5th Century B.C. proposed that matter
is composed of tiny, indestructable units called “atoms.”
• Atoms themselves do not change
• Move around and combine with other atoms
• Properties of an atom are determined by the shape of the atom
• Aristotle –320B.C.
– believed in 4 Essences Theory, that all
matter was made of fire, air, earth and water.
JOHN DALTON
• 1820’s
DALTON’S ATOMIC THEORY
• 1. An element is composed of tiny indivisible indestructible
particles called atoms
• 2. All atoms of an element are identical and have the same
properties
• 3. Atoms of different elements combine to form compounds
• 4. Compounds contain atoms in small whole number ratios
• 5. Atoms can combine in more than one ratio to form different
compounds
Atomic structure
THOMSON EXPERIMENT
1897, Joseph John Thomson:
Determined the charge: Mass ratio of cathode rays
(discovered electrons).
© 2013 Pearson Education, Inc.
8
MODERN CATHODE RAY TUBES
Television
Computer Monitor
Cathode ray tubes pass electricity
through a gas that is contained at a
very low pressure.
J. J.’s Model of the atom
–Developed the “Plum Pudding” model of the atom
Negatively
charged electrons
Positively charged
material
• cathode ray tube
• Student
of JJ Thomson
•1911 he described a small,
heavy nucleus with electrons in
orbit around the nucleus
•1919 – discovered a proton
and determined that it exists in
the nucleus of an atom
•Protons are positively charged
RUTHERFORD’S GOLD FOIL EXPERIMENT
NIELS BOHR
•
•
•
•
1913
Put electrons into energy levels or orbits
Electrons with lower energy are closest to nucleus
Nuclear Model of the Atom
• http://www.youtube.com/watch?v=ecsgC1w5p5I&feature=related
LOUIS DEBROGLIE AND IRWIN SCHRODINGER
• early 1900’s
• Found that electrons can shift from one energy
level to another.
• Because of electron movement they described
electrons as 3-D waves.
• This became known as the Quantum
Mechanical or Cloud Model of the atom
QUANTUM MECHANICAL MODEL
• Based on idea of electron clouds – based on
statistical probability of finding an electron in a
certain position – mathematical model
• Regions where electrons are likely to be found are
called atomic orbitals and are labeled s,p,d,f
ATOMIC MODELS THROUGH TIME
• A boy and his atom how its made
• a boy and his atom
THE TWO MAIN PORTIONS OF
THE ATOM ARE…….
Nucleus
Electron cloud
ATOMIC STRUCTURE
3 Main Subatomic Particles
-Protons
-Neutrons
-Electrons
PROTONS
• Located in nucleus
• Have a positive charge
• Have a mass of 1 a.m.u. (atomic
mass unit)
• Made of 3 smaller parts called
quarks
NEUTRONS
•
•
•
•
Located in the nucleus
Are neutral – no charge
Have a mass of 1 a.m.u.
Made of 3 smaller parts called
quarks
Isotopes
• Atoms with the same # of protons,
but different # of neutrons.
• Ex.
H-1
Protium
H-2
Deuterium
H-3
Tritium
-Isotopes have different Mass
Numbers but same atomic number.
ELECTRONS
• Located in electron cloud outside
of the nucleus
• Arranged in energy levels
• Have a negative charge
• Extremely small mass.
• Have a diameter of 1 x 10-18m –
(about 10 million times smaller than
an atom)
ELECTRONS…..WHERE ARE THEY IN THE
CLOUD?
REMEMBER THE BOHR MODEL….
• Understanding Electron Arrangements (e- configurations)
is the key to understanding the chemical properties of
elements.
• As scientists began to determine these configurations,
they realized that energy and stability play a huge part in
how electrons are configured.
• We, in chemistry, need to gain a basic understanding in
order to proceed.
PRINCIPAL ENERGY LEVEL
SUBLEVEL
ORBITAL
QUANTUM NUMBERS
• Used to describe the region of space with the highest
probability of finding an electron.
• Principal Energy level (n)
• n = 1,2,3,4,5,6,7
• Sublevels correspond to an orbital of a different shape
• Each energy level has an equal number of sublevels
• Denoted by letters (s,p,d,f)
• Orbital
• Describes highest probability
• Contained within sublevels
• Shapes describe probability
RULES FOR ELECTRON CONFIGURATIONS
• Aufbau Principle: electrons occupy the orbitals of lowest
energy first.
• Pauli Exclusion Principle: An atomic orbital may describe at
most two e- with opposite spin.
• Hund’s Rule: When occupying a sublevel where all orbitals
have equal energy, orbitals must all fill with one electron
before any orbital contains two electrons.
ENERGY LEVELS
n=1
n=2
n=3
n=4
TYPES OF ORBITALS
• The most probable area to find these electrons
takes on a shape
• So far, we have 4 shapes. They are named s, p, d,
and f. (sharp, principal, diffuse, fundamental)
• No more than 2 e- assigned to an orbital – one spins
clockwise, one spins counterclockwise
ELECTRON CONFIGURATIONS
A list of all the electrons in an atom (or ion)
• We need electron configurations so that we can
determine the number of electrons in the outermost
energy level. These are called valence electrons.
• The number of valence electrons determines how
many and what this atom (or ion) can bond to in
order to make a molecule
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.
ELECTRON CONFIGURATION
1
1s
# of electrons
row # (on periodic table)
in that subshell
Also known as shell #
(principal # n)
possibilities are 1-7
subshell
7 rows
possibilities are
s, p, d, or f
4 subshells
group # = # valence (outside) e-
1A
1
2A
8A
3A 4A 5A 6A 7A
2
Row 3
=
4
# shells
5
6
3B 4B 5B 6B 7B 8B 8B 8B 1B 2B
s
d
p
7
6
7
f
SUBSHELLS D AND F ARE “SPECIAL”
group # = # valence e-
1A
period # = # e- shells
1
2A
3A 4A 5A 6A 7A
2
3
3B 4B 5B 6B 7B 8B 8B 8B 1B 2B
4
3d
5
4d
5d
6d
6
7
8A
d
6 4f
7 5f
f
SHORTHAND NOTATION
Step 1 Find the closest noble gas to the atom (or
ion), WITHOUT GOING OVER the number of
electrons in the atom (or ion). Write the noble gas
in brackets [ ].
Step 2: Find where to resume by finding the next
energy level.
Step 3: Resume the configuration until it’s finished.
THE PERIODIC TABLE
ORGANIZING THE ELEMENTS
• Cu, Ag, Au are three of the oldest known elements.
• By the year 1700, only 13 elements had been identified and
isolated.
• From 1765-1775, five new elements including H, N, and O had
been isolated.
• As soon as elements were (are) identified, scientists begin to look
for similarities and classify them.
CLASSIFYING ELEMENTS
• JW Dobereiner published a classification system
in 1829.
• Elements grouped in triads (3 elements with similar
properties)
• Couldn’t group all elements using that system
DEVELOPMENT OF THE PERIODIC
TABLE
• Mendeleev- arranged elements by increasing
atomic mass
• Saw pattern in valence #’s and properties
• Predicted properties of “missing” elements
• Some seemed out of order
MENDELEEV’S TABLE
THE PERIODIC LAW
• In developing the Periodic Table, Mendeleev noticed problems.
• Atomic mass order didn’t always match periodic trends.
• The table was eventually rearranged according to atomic
number.
• Mosley – arranged elements by increasing
atomic number
• Out of order elements, now fell into correct
order
PERIODIC LAW:
• When elements are arranged in order of
increasing atomic number, there is a
of their physical and
chemical properties.
•
OTHER PERIODIC TABLES
• periodic tables
Parts of the Periodic Table
Columns = Groups or Families
-Have same # of valence electrons
-Have similar properties
Rows
= Periods or Series
-Have same number of energy levels
METALS, NONMETALS AND METALLOIDS
Properties of Metals
•
•
•
•
•
•
•
•
Malleable
Ductile
Electron donors
Conductors of heat and electricity
Solids at room temp. except Hg
High melting points
Silver, gray or white (except Cu and Au)
Higher densities
Properties of Nonmetals
•
•
•
•
•
•
•
•
Non-malleable
Non-ductile
Electron acceptors
Insulators
Low melting points
Variety of colors
Solids, liquids and gases
Lower density
Metalloids
7 - Boron, Silicon, Germanium, Arsenic,
Antimony, Tellurium, Astatine
-Have some prop. of metals
and nonmetals
-most common is silicon
-used in semiconductors in
computer chips
WHAT FAMILY HAS…..
a metalloid and 4 non metals?
A metalloid and 4 metals?
3 nonmetals, 1 metalloid and 1 metal?
4 nonmetals, 2 gases and a liquid and 2
solids?
• a gas and 4 solids?
• a nonmetal, 2 metalloids and 2 metals?
•
•
•
•
ELEMENTS CAN BE SORTED INTO…
• noble gases
• representative elements
• transition elements
• …based on their
electron
configurations
• inner
transition
elements
TRANSITION ELEMENTS
• 2 Types
• Transition metals
- Inner transition metals
TRANSITION ELEMENTS (METALS)
• In “d” sublevel (“B” or # 3-12)
Ex. Copper, silver, gold and Iron
Since “s” sublevel fills before “d” and “p” after, they
usually have two valence electrons, but there are
exceptions.
Transition Metals
INNER TRANSITION ELEMENTS
(METALS)
•In “f” sublevel (not numbered)
•Also known as “Rare Earth Elements”
Since “s” sublevel fills before “f” and “d
and “p”after, they usually have two
valence electrons, but there are
exceptions.
Inner Transition Metals
THE PERIODIC TABLE SHOWS
TRENDS OR PATTERNS IN..
1. Atomic Number
2. Valence Electrons
3. Atom Size and Ion Size
4. Ionization Energy
5. Electronegativity
6. Chemical Reactivity
7. Electron Affinity
Atomic Size
•Size Increases going down a group or
family.
•Because…
–adding more energy levels
–electrons added are further from the nucleus,
there is less attraction. This is due to additional
energy levels and the shielding effect. Each
additional energy level “shields” the electrons
from being pulled in toward the nucleus.
Size decreases across a period
• Because….
– more protons, increases the + charge
– Each added electron feels a greater and greater +
charge because the protons are pulling in the same
direction, whereas the electrons are scattered.
Large
Small
Which is Bigger?
• Na or K ?
• Na or Mg ?
• Al or I ?
Remember… All atoms want to
achieve a FULL outside energy level
• That means last “s” and “p” must be full
with 8 valence electrons
• Hydrogen and Helium are the exceptions
( only have a “s” sublevel so they are full
with 2 electrons)
IONS
• An atom or group of atoms that has a
positive or negative charge due to losing or
gaining electrons.
• Indicated with a + or – after the symbol
showing the charge on the ion.
• Ex. Na
1+
, Mg
2+
, Cl1-
IONS ARE FORMED WHEN
ELECTRONS ARE TRANSFERRED
BETWEEN ATOMS
• Atoms want to have a filled outer energy level
• Generally if an atom has….
• 1,2 or 3 valence electrons – the atom will give away it’s
electrons (cations)
• 5,6,or 7 valence electrons – the atom will take
electrons (anions)
METALS
• Metals tend to lose lelectrons
• Ex: Sodium loses one: there are now more
protons (11) than electrons (10), and thus a
positively charged particle is formed
• Positive ions = “cation”
•The charge is written as a number followed
by a plus sign: Na1+
• •Now named a sodium ion
NONMETALS
• Nonmetals tend to
• Ex: Chlorine will gain one electron- protons (17) no longer
outnumber the electrons (18)
• So now has a charge of Cl-1
• Known as a chlor
ion
•Cations are positive
•Anions are negative
Try Some Ions!
• Write the longhand notation for these:
F1Li1+
Mg2+
• Write the shorthand notation for
these:
Br1Ba2+
Al3+
Ionization Energy (IE)
IE = the amount of energy required to
remove an electron from an atom (in
the gas phase).
Trends in Ionization Energy
•IE decreases DOWN a group or family
•Because size increases (Shielding
Effect)
•The further away from nucleus,
the less the electrons feel the
attraction of the protons therefore
easier to remove
Trends in Ionization Energy
•General trend IE increases across a period
•because the positive charge increases,
holding the electrons tighter.
•It dips down after filling “s’ orbital and
after 1 electron is put in each “p” orbital
1ST
2ND
3RD
H 1312
He 2731
5247
Li 520
7297
11810
Be 900
1757
14840
B 800
2430
3569
C 1086
2352
4619
N 1402
2857
4577
O 1314
3391
5301
F
1681
3375
6045
Ne 2080
3963
6276
Trends in Ionization Energy
I. E. generally decreases
I. E. generally increases
5. Electronegativity, 
 is a measure of the ability of an
atom in a molecule to attract
electrons to itself.
Concept proposed by
Linus Pauling
1901-1994
Electronegativity
•In a group: Atoms with fewer energy
levels can attract electrons better
(less shielding). So, electronegativity
decreases DOWN a group (FAMILY) of
elements.
•In a period: More protons, while the
energy levels are the same, means
atoms can better attract electrons.
So, electronegativity increases RIGHT
in a period of elements.
Electronegativity
increases
decreases
Which is more electronegative?
•F or Cl ?
•Na or K ?
•Sn or I ?