Download Electrons and Periodic Trends Summary

I.
Electrons
and
Period
Trends
Summary
Properties
of
Light
A. Wave
Description
of
Light
1. Electromagnetic
radiation
is
a
form
of
energy
that
exhibits
wavelike
behavior
as
it
travels
through
space.
2. Electromagnetic
spectrum
is
all
forms
of
electromagnetic
radiation
put
together.
http://lot.astro.utoronto.ca/images/spectrum.png
II.
III.
3. Wavelength
the
distance
between
corresponding
points
on
adjacent
waves.
Photoelectric
Effect
A. The
emission
of
electrons
from
a
metal
when
light
shines
on
the
metal.
B. Quantum:
the
minimum
quantity
of
energy
that
can
be
lost
or
gained
by
an
atom.
C. Photon:
is
a
particle
of
electromagnetic
radiation
having
zero
mass
and
carrying
a
quantum
energy.
Bright
Line
Emission
A. Ground
state:
the
lowest
energy
state
of
a
atom.
B. Excited
state:
a
state
in
which
the
atom
has
a
higher
potential
energy
than
it
has
in
its
ground
state.
C. Line‐emission
spectrum:
Bands
of
light
that
correspond
to
an
elements
identity.
D. Continuous
spectrum:
the
emission
of
a
continuous
range
of
frequencies
of
electromagnetic
radiation.
IV.
Bohr
Model
of
the
Atom
A. Niels
Bohr
proposed
a
model
of
the
hydrogen
atom
based
on
its
line‐emission
spectrum.
(http://www.splung.com/nuclear/images/atoms10.gif)
V.
Quantum
Model
of
the
Atom
A. Quantum
Theory
(Heisenberg
and
Schrodinger)
describes
mathematically
the
wave
properties
of
electrons
and
other
very
small
particles.
(http://farm1.static.flickr.com/59/169563400_cf75d3ccf9.jpg)
B. Quantum
Number:
specifies
the
properties
of
atomic
orbitals
and
the
properties
of
electrons
in
orbitals.
C. Principle
Quantum
Number
(n)
the
main
energy
level
occupied
by
the
electrons
(1,2,3,4,5,6,7
‐‐one
for
each
period
on
the
periodic
table)
http://image.tutorvista.com/content/atomic-structure/relative-energy-levels-hydrogen-and-hydrogen-like-atoms.gif
VI.
D. Magnetic Quantum Number: the orientation of the orbital
around the nucleus (s, px, py, pz, dx2-y2, dxy, dyz, dxz, dz, f etc…)
E. Spin Quantum Number: + or – the two fundamental spin states of
an electron in orbit.
F. General rules that hold true:
IF
n = the number of the energy level (1-8)
THEN n = the number of orbital types in that energy level
n2
=
the
total
number
of
orbitals
in
that
energy
level
2n2
=
the
total
number
of
electrons
that
can
exist
in
the
energy
level.
Electron
Configurations
A. Aufbau
Principle:
an
electron
occupies
the
lowest
–energy
orbital
that
can
receive
it.
B. Pauli
Exclusion
Principle:
no
two
electrons
in
the
orbital
can
have
the
same
spin
(they
must
have
opposite
spins).
C. Hund’s
rule:
when
filling
an
energy
level
all
orbitals
MUST
get
one
electron
before
any
orbital
can
have
two
electrons.
D. Electron
Dot
Notation:
Dots
are
used
to
denote
the
outermost
shell
of
electrons.
The
dots
are
placed
according
to
which
orbital
electrons
they
represent.
3
4
7
X
2
1
5
6
8
E. Electron
Configuration:
1.
Lists
the
energy
level,
the
orbital
type
and
the
number
of
electrons.
EXAMPLE:
3p3
this
means
that
in
the
p
orbitals
of
the
third
energy
level
there
are
three
electrons.
2.
When
listing
a
complete
configuration
notation,
group
the
energy
levels
together.
Even
though
the
order
of
filling
for
an
atom
may
be:
1s2
2s2
2p6
3s2
3p6
4s2
3d6
We
group
our
energy
levels
and
show
it
as
follows:
1s2
2s2
2p6
3s2
3p6
3d6
4s2
The
sum
of
the
superscripts
will
always
be
the
number
of
electrons
in
the
atom
(atomic
number).
F.
Orbital
Notation:
This
is
very
similar
to
configuration.
We
label
the
orbitals
by
1s,
2s,
etc,
but
we
show
the
electrons
by
slashes
or
arrows.
VII.
History
of
the
Periodic
Table
A. Dmitri
Mendeleev:
1. Organized
elements
based
on
properties.
2. Left
empty
spaces
for
“new”
elements
B. Henry
Moseley:
1. Arranged
elements
by
their
increasing
atomic
charge
(number
of
protons)
C. Periodic
Law:
The
physical
and
chemical
properties
of
the
elements
are
periodic
functions
of
their
atomic
numbers.
1. Elements
with
the
same
properties
fall
in
the
same
column.
VIII. Groups/Families
(Columns)
A. Noble
gases:
chemically
unreactive,
full
outer
shell
(octet
rule)
B. Alkali
metals:
Group
1
C. Alkaline
Earth
Metals:
Group
2
D. Transition
Metals
(Groups
3‐12)
E. Halogens:
(Group
17)
IX.
Periods
(Rows)
A. The
Lanthanides
(58‐71)—fill
the
f
block
B. The
Actinides
(90‐103)—fill
f
block
X.
Trends
A. Atomic
Radius:
as
one
half
the
distance
between
the
nuclei
of
identical
atoms
that
are
bonded
together.
B. Ionization
Energy:
the
energy
required
to
remove
one
electron
from
a
neutral
atom
of
an
element.
1. Second
Ionization
Energy:
the
addition
energy
need
to
remove
a
second
electron.
C. Electron
Affinity:
the
energy
CHANGE
that
occurs
when
an
electron
is
acquired
by
a
neutral
atom.
D. Ionic
Radii:
The
size
of
the
charged
atom
1. Cation:
+
charge
2. Anion:
‐
charge
E. Valence
Electrons:
the
electrons
available
to
be
lost,
gained
or
shared
in
the
formation
of
chemical
bonds.
F. Electronegativity:
the
measure
of
the
ability
of
an
atom/element
to
attract/steal
electrons
from
another
element/atom.
http://www.meta‐synthesis.com/webbook/24_complexity/periodic_trends.jpg
http://www.chemistry.wustl.edu/~courses/genchem/Tutorials/LED/images/periodic.jpg