September 16, 2013
- When the bell rings voices are at zero, working on the DO NOW
( All electronics away!!)
- Pick up daily handouts
- Turn in – (LATE M&Ms Lab Report, LATE Beanium ClEvR, LATE
Anticipation guides, Lab safety contracts)
- Bohr Article (HW last night) is the next page in your NB – DO
NOT TURN IT IN
- HAVE OUT YOUR REFERENCE TABLE
- Start on the Do Now in your Do Now form
DO NOW 9/17/13
Complete the following decay reactions
For the following find Protons, Neutrons, Electrons, and charge
1. 9038Sr
2. 9038Sr+2(think of it as “to the +2)
Do Now Review
DO NOW 9/17/13
Homework
Lab Report for Chemistry in a bag Due Friday
9/20/2013 ***
( Cover Letter, Introduction, and Hypothesis)
Study for Nuclear Chemistry Test ( tomorrow )
Graded***
Upcoming Dates
•
•
•
•
•
•
9/18/2013 Wednesday Nuclear Chemistry Test
9/20/2013 Mole Conversion test
9/20/2013 Chemistry in a bag report due
9/25/2013 Empirical Formula Test
9/27/2013 Electron / Periodic Properties Test
9/27/2013 Periodic Table Wanted Poster Due
Tutoring
• Tuesday 2:30-3:30 in room 529
• Thursday 2:30-3:30 in room 710
Test 2 Review: Atomic Structure
•
•
•
•
Isotopes (atomic mass vs. atomic number)
Finding Protons, Neutrons, and Electrons
Cathode Ray (Vision Learning Assignment)
Percent Abundance
Percent Abundance
• Mg has 3 naturally occurring isotopes, Mg-24
(78.99%), Mg-25 (10%), and Mg-26 (11.01%).
Calculate the average atomic mass from Mg.
•
•
•
24Mg=
23.985042amu, 78.99%
25Mg= 24.985837 amu, 10.00%
26Mg= 25.982593, 11.01%
Objective 9/17/2013
1. SWBAT articulate that electromagnetic radiation is
given off as photons
2. SWBAT communicate inverse relationship between
wavelength and frequency, and the direct
relationship between energy and frequency.
3. SWBAT Describe the wave/particle duality of
electrons
4. SWBAT utilize the Bohr Model for Hydrogen Atom
and the Electromagnetic spectrum diagrams from
the reference table to relate color, frequency, and
wavelength of the light emitted to the energy of the
photon.
Modern Atomic Theory
(a.k.a. the electron chapter!)
Bohr Model and Spectrum
Video
• http://educationportal.com/academy/lesson/the-bohrmodel-and-atomic-spectra.html
Electromagnetic radiation.
Electromagnetic Radiation
• Most subatomic particles behave as
PARTICLES and obey the physics of
waves.
Electromagnetic Radiation
• Waves have a frequency
• Use the Greek letter “nu”,
are “cycles per sec”
, for frequency, and units
  = c
• All radiation:
•
where c = velocity of light = 3.00 x 108 m/sec
Electromagnetic Spectrum
Long wavelength --> small frequency
Short wavelength --> high frequency
increasing
frequency
increasing
wavelength
Electromagnetic Spectrum
In increasing energy, ROY G BIV
Excited Gases
& Atomic Structure
Atomic Line Emission Spectra
and Niels Bohr
Bohr’s greatest contribution to
science was in building a simple
model of the atom. It was based
on an understanding of the
LINE EMISSION SPECTRA of
excited atoms.
Niels Bohr
(1885-1962)
Bohr Model
•
•
•
•
•
•
Pick an Element (H)
Find P, N, E
Draw Nucleus
Draw first energy level
Draw in Electrons
Practice; (He) (B) (Ne)
Energy Levels
n=1
n=2
n=3
n=4
n=5
n=6
n=7
Line Emission Spectra
of Excited Atoms
• Excited atoms emit light of only certain
wavelengths
• The wavelengths of emitted light depend on
the element.
Spectrum of
Excited Hydrogen Gas
Line Spectra of Other
Elements
The Electric Pickle
• Excited atoms can emit
light.
• Here the solution in a
pickle is excited
electrically. The Na+ ions
in the pickle juice give
off light characteristic of
that element.
http://www.youtube.com/
watch?v=mh7VHcuaPCg
Atomic Spectra
One view of atomic structure in early 20th
century was that all electron (e-) traveled
about the nucleus in an orbit randomly.
Atomic Spectra and Bohr
Bohr said classical view is wrong.
Need a new theory — now called
QUANTUM or WAVE MECHANICS.
e- can only exist in certain discrete orbits
e- is restricted to QUANTIZED energy state
(quanta = bundles of energy)
Quantum or Wave Mechanics
Schrodinger applied idea of ebehaving as a wave to the problem
of electrons in atoms.
He developed the WAVE
EQUATION
E. Schrodinger
1887-1961
Each solution describes an allowed
energy state of an e-
Heisenberg Uncertainty
Principle
W. Heisenberg
1901-1976
Problem of defining nature of
electrons in atoms solved by
W. Heisenberg.
Cannot simultaneously define
the position and momentum
(= m•v) of an electron.
We define e- energy exactly
but accept limitation that we
do not know exact position.
Arrangement of Electrons in
Atoms
Electrons in atoms are arranged as
LEVELS (n)
SUBLEVELS (l)
ORBITALS (ml)
QUANTUM NUMBERS
The shape, size, and energy of each orbital is a
function of 3 quantum numbers which
describe the location of an electron within an
atom or ion
n (principal) ---> energy level
l (orbital) ---> shape of orbital
ml (magnetic) ---> designates a particular
suborbital
The fourth quantum number is not derived from
the wave function
s (spin) ---> spin of the electron
(clockwise or counterclockwise: ½ or – ½)
QUANTUM NUMBERS
So… if two electrons are in the same place at the same
time, they must be repelling, so at least the spin
quantum number is different!
The Pauli Exclusion Principle says that no two electrons
within an atom (or ion) can have the same four
quantum numbers.
If two electrons are in the same energy level, the same
sublevel, and the same orbital, they must repel.
Think of the 4 quantum numbers as the address of an
electron… Country > State > City > Street
Energy Levels
• Each energy level has a number called
the PRINCIPAL QUANTUM NUMBER, n
• Currently n can be 1 thru 7, because
there are 7 periods on the periodic table
• Periods = Rows
• Groups = Colomns
Energy Levels
n=1
n=2
n=3
n=4
n=5
n=6
n=7
Relative sizes of the spherical 1s, 2s, and
3s orbitals of hydrogen.
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.
• No more than 2 e- assigned to an orbital – one
spins clockwise, one spins counterclockwise
Types of Orbitals (l)
s orbital
p orbital
d orbital
p Orbitals
this is a p sublevel
with 3 orbitals
These are called x, y, and
z
3py
orbital
There is a PLANAR
NODE thru the
nucleus, which is
an area of zero
probability of
finding an electron
p Orbitals
• The three p orbitals lie 90o apart in space
d Orbitals
• d sublevel has 5
orbitals
The shapes and labels of the five 3d
orbitals.
f Orbitals
For l = 3,
---> f sublevel with 7 orbitals
Diagonal Rule
• Must be able to write it for the test! This will
be question #1 ! Without it, you will not get
correct answers !
• The diagonal rule is a memory device that
helps you remember the order of the filling of
the orbitals from lowest energy to highest
energy
• Aufbau principle states that electrons fill from
the lowest possible energy to the highest
energy
Diagonal Rule
Steps:
1s
2s
3s
4s
1.
Write the energy levels top to bottom.
2.
Write the orbitals in s, p, d, f order. Write
the same number of orbitals as the energy
level.
3.
Draw diagonal lines from the top right to the
bottom left.
4.
To get the correct order,
2p
3p
4p
3d
4d
4f
follow the arrows!
By this point, we are past
the current periodic table
so we can stop.
5s
5p
5d
5f
5g?
6s
6p
6d
6f
6g?
6h?
7s
7p
7d
7f
7g?
7h?
7i?
How many electrons can be in a sublevel?
Pauli Exclusion Principle: A maximum of two
electrons can be placed in an orbital.
s orbitals p orbitals d orbitals f orbitals
Number of
orbitals
Number of
electrons
1
3
5
7
2
6
10
14
Electron Configurations
A list of all the electrons in an atom (or ion)
• Must go in order (Aufbau principle)
• 2 electrons per orbital, maximum
• 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 Configurations
4
2p
Energy Level
Number of
electrons in
the sublevel
Sublevel
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6
6s2 4f14… etc.
Let’s Try It!
• Write the electron configuration for the
following elements:
H
Li
N
Ne
K
Zn
Pb
September 18, 2013
- When the bell rings voices are at zero, working
on the DO NOW ( All electronics away!!)
- Pick up daily handouts
- Turn in – (LATE M&Ms Lab Report, LATE
Beanium ClEvR, LATE Anticipation guides, Lab
safety contracts)
- Bohr Article (HW Monday night) is the next
page in your NB – DO NOT TURN IT IN
- HAVE OUT YOUR REFERENCE TABLE
- Start on the Do Now in your Do Now form
DO NOW 9/17/13
Write the electron configuration for Ne and Li
Do Now Review
DO NOW 9/18/13
Write the electron configuration for Ne and Li
Homework
Lab Report for Chemistry in a bag Due Friday
9/20/2013 ***
( Cover Letter, Introduction, and Hypothesis)
Electron Configuration and Orbital Practice
Worksheet
Graded***
Upcoming Dates
•
•
•
•
•
9/20/2013 Mole Conversion test (FRIDAY)
9/20/2013 Chemistry in a bag report due
9/25/2013 Empirical Formula Test
9/27/2013 Electron / Periodic Properties Test
9/27/2013 Periodic Table Wanted Poster Due
Tutoring
• Tuesday 2:30-3:30 in room 529
• Thursday 2:30-3:30 in room 710
Objectives 9/17/2013
1. SWBAT articulate that electromagnetic radiation is
SWBAT
Describe
concepts of excited and
given off
as photons
ground
state
of electrons
the atom:between
2. SWBAT
communicate
inverseinrelationship
1. Gaining
energy
resultsand
in theelectron
wavelength
and frequency,
the direct
moving
from
its ground
to a higher
relationship
between
energy state
and frequency.
level the wave/particle duality of
3. energy
SWBAT Describe
2. When
the electronmoved to a lower
electrons
levelthe
it relases
theforenergy
4. energy
SWBAT utilize
Bohr Model
Hydrogen Atom
and the Electromagnetic
difference
in the two spectrum
levels asdiagrams from
the reference table to
relate color,
frequency, and
electromangetic
radiation
(emissions
wavelength of the light emitted to the energy of the
spectrum
photon.
An excited lithium atom emitting a
photon of red light to drop to a
lower energy state.
An excited H atom returns to a
lower energy level.
Orbitals and the Periodic
Table
• Orbitals grouped in s, p, d, and f orbitals (sharp, proximal,
diffuse, and fundamental)
s orbitals
f orbitals
d orbitals
p orbitals
Shorthand Notation
• A way of abbreviating long electron
configurations
• Since we are only concerned about
the outermost electrons, we can
skip to places we know are
completely full (noble gases), and
then finish the configuration
Shorthand Notation
• Step 1: It’s the Showcase Showdown!
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.
Shorthand Notation
• Chlorine
– Longhand is 1s2 2s2 2p6 3s2 3p5
You can abbreviate the first 10
electrons with a noble gas, Neon.
[Ne] replaces 1s2 2s2 2p6
The next energy level after Neon is 3
So you start at level 3 on the diagonal
rule (all levels start with s) and finish
the configuration by adding 7 more
electrons to bring the total to 17
[Ne] 3s2 3p5
Practice Shorthand Notation
• Write the shorthand notation for each of the
following atoms:
Cl
K
Ca
I
Bi
Valence Electrons
Electrons are divided between core and valence
electrons
B 1s2 2s2 2p1
Core = [He] , valence = 2s2 2p1
Br [Ar] 3d10 4s2 4p5
Core = [Ar] 3d10 , valence = 4s2 4p5
Rules of the Game
No. of valence electrons of a main group atom =
Group number (for A groups)
Atoms like to either empty or fill their outermost
level. Since the outer level contains two s
electrons and six p electrons (d & f are always in
lower levels), the optimum number of electrons
is eight. This is called the octet rule.
Keep an Eye On Those Ions!
• Electrons are lost or gained like they always
are with ions… negative ions have gained
electrons, positive ions have lost electrons
• The electrons that are lost or gained should be
added/removed from the highest energy level
(not the highest orbital in energy!)
Keep an Eye On Those Ions!
• Tin
Atom: [Kr] 5s2 4d10 5p2
Sn+4 ion: [Kr] 4d10
Sn+2 ion: [Kr] 5s2 4d10
Note that the electrons came out of the highest
energy level, not the highest energy orbital!
Keep an Eye On Those Ions!
• Bromine
Atom: [Ar] 4s2 3d10 4p5
Br- ion: [Ar] 4s2 3d10 4p6
Note that the electrons went into the highest
energy level, not the highest energy orbital!
Try Some Ions!
• Write the longhand notation for these:
FLi+
Mg+2
• Write the shorthand notation for these:
BrBa+2
Al+3
Orbital Diagrams
• Graphical representation of an electron
configuration
• One arrow represents one electron
• Shows spin and which orbital within a
sublevel
• Same rules as before (Aufbau principle,
d4 and d9 exceptions, two electrons in
each orbital, etc. etc.)
Orbital Diagrams
• One additional rule: Hund’s Rule
– In orbitals of EQUAL ENERGY (p, d, and
f), place one electron in each orbital
before making any pairs
– All single electrons must spin the same
way
• I nickname this rule the
“Monopoly Rule”
• In Monopoly, you have to build
houses EVENLY. You can not put
2 houses on a property until all
the properties has at least 1
house.
Lithium
Group 1A
Atomic number = 3
1s22s1 ---> 3 total electrons
3p
3s
2p
2s
1s
Carbon
Group 4A
Atomic number = 6
1s2 2s2 2p2 --->
6 total electrons
3p
3s
2p
2s
1s
Here we see for the first time
HUND’S RULE. When
placing electrons in a set of
orbitals having the same
energy, we place them singly
as long as possible.
Lanthanide Element Configurations
4f orbitals used for
Ce - Lu and 5f for
Th - Lr
Draw these orbital diagrams!
• Oxygen (O)
• Chromium (Cr)
• Mercury (Hg)
Ion Configurations
To form anions from elements, add 1 or more efrom the highest sublevel.
P [Ne] 3s2 3p3 + 3e- ---> P3- [Ne] 3s2 3p6 or [Ar]
3p
3p
3s
3s
2p
2p
2s
2s
1s
1s