Transcript Chapter 20 Oxidation-Reduction Reactions

Chemistry Warm Up
Some Dimensional Analysis Review.
PLEASE SHOW YOUR WORK USING CONVERSION
FACTORS AND DIMENSIONAL ANALYSIS
1. If 6.02 x 1023 atoms of carbon have a mass of 12.0 grams,
what the mass of 1.51 x 1023 atoms of carbon. Hint: set up
the equality that you know. Make two conversion factors and use one to
solve the problem. Check your work using dimensional analysis.
2. How many atoms are there in sample of carbon that weighs
36.0grams?
3. What is the mass of a sample containing 1.204x1022 atoms.
ATOMIC MASS V ATOMIC MASS NUMBER
Almost all carbon is one of three isotopes.
1.Write the isotope notation for carbon-14:
2.Write the isotope notation for carbon 12:
3.What is the mass number of carbon 13?
4. Use your periodic table to determine the atomic
mass of carbon. Write it here.
5. How many neutrons does carbon have?
6.Learning objective here! Explain the difference
between atomic mass and mass number of an
element in terms of isotopes, mass, protons, and
neutrons.
•
3.3 Electrons and Light
California State Science Standards Chemistry
1. The periodic table displays the elements in increasing
atomic number and shows how periodicity of the physical and
chemical properties of the elements relates to atomic structure.
As a basis for understanding this concept:
g.* Students know how to relate the position of an element in
the periodic table to its quantum electron configuration and to
its reactivity with other elements in the table.
i.* Students know the experimental basis for the development
of the quantum theory of atomic structure and the historical
importance of the Bohr model of the atom.
3.3 Electrons and Light
California State Science Standards Chemistry
1. The periodic table displays the elements in increasing
atomic number and shows how periodicity of the physical
and chemical properties of the elements relates to atomic
structure. As a basis for understanding this concept:
2. i.* I know the experimental basis for the
development of the quantum theory of
atomic structure and the historical
importance of the Bohr model of the
atom.
Quick Review:
Models of the Atom
Dalton- Indivisible Atom
J.J.Thomson
discovers subatomic
particle
“Plum pudding,”
model
Rutherford’s Gold Foil
Experiment
Alpha particles shot at
a thin piece of gold
foil did not pass right
through with slight
deflection. Instead,
most passed
straight through.
Some bounced right
back!
Rutherford’s Gold Foil
Experiment
Rutherford concluded that
•Most of the atom is empty space
•All of the positive charge
and almost all of the mass is
Concentrated in the tiny core, “nucleus.”
composed of protons and neutrons.
An idea of the size:
Atom = football stadium
Nucleus = marble
5.3 Physics and the Quantum
Mechanical Model
Or, “How do they get all those colors of neon lights?”
Wave Terminology
Amplitude = height of wave
Wavelength = distance between crests
Frequency = number of crests to pass a point
per unit of time
Light waves
Amplitude = height of wave
Wavelength = distance between crests
Frequency = number of crests to pass a point
per unit of time
For light, the product of frequency and
wavelength = speed of light, c
Frequency • Wavelength = 3.00 x 108
So, as the frequency of light increases, the
wavelength decreases
Electromagnetic
SpectrumWavelength of Light p140
Visible light is only part of the
electromagnetic spectrum:
Atomic Spectra
When atoms absorb energy,
they move to higher energy (excited) levels.
When electrons return to the lower energy
level, or ground state, they emit light
Each energy level produces a certain frequency
of light resulting in an emission spectrum
Atomic Spectra
Emission spectra are like a fingerprint of the
element
We know what stars are made of by comparing
their emission spectra to that of elements we
find on earth
Explanation of Atomic Spectra
Emission spectra like a fingerprint of the
element
We know what stars are made of by comparing
their emission spectra to that of elements we
find on earth
Development of Atomic Models
Rutherford’s Model
l
l
l
l
Dense central Nucleus
Electrons orbit like planets
Atom mostly empty space
Does not explain chemical
behavior of atoms
The Bohr Model
l
l
l
Electrons orbit the nucleus
Specific circular orbits
Quantum =
energy to move
from one level
to another
Note Taking Strategies
WHEN YOU ARE
DONE, YOU NEED TO
BE ABLE TO EXPLAIN
THE CONNECTION
BETWEEN LIGHT
AND ELECTRONS.
We will
Take notes on section 3.3 Electrons and Light ppg92-95
Using an outlining strategy
Use phrases in RED as main topic headings.
Use phrases in BLUE as subheadings.
Important: illustrations are often more important than the
text.
study each illustration for meaning.
sketch the illustration into your notes if it is pertinent
Include vocabulary with meaning.
If you don’t understand a term, don’t just copy its
definition.
READ FOR UNDERSTANDING.
Note Taking Strategies
We will
Take notes on section 3.3 Electrons and Light ppg92-95
Using an outlining strategy
Use phrases in RED as main topic headings.
Use phrases in BLUE as subheadings.
Important: illustrations are often more important than the text.
study each illustration for meaning.
sketch the illustration into your notes if it is pertinent
Include vocabulary with meaning.
If you don’t understand a term, don’t just copy its definition.
READ FOR UNDERSTANDING.
WHEN YOU ARE DONE, YOU NEED TO BE ABLE TO
EXPLAIN THE CONNECTION BETWEEN LIGHT AND
ELECTRONS.
The Bohr Model
Energy level like rungs of the ladder
The electron cannot exist between energy
levels, just like you can’t stand between
rungs on a ladder
A quantum of energy is the amount of
energy required to move an electron from
one energy level to another
The Bohr Model
Energy level of an
electron analogous to
the rungs of a ladder
But, the rungs on this ladder
are not evenly spaced!
Title your paper:
Atomic Emission Spectra
Goal: to understand what emission spectra are.
We will make a table in which we draw and describe
several emission spectra.
1. Put on the diffraction glasses and look outside
Find a rainbow. Draw and describe the spectrum you
see.
Source
Sunlight
Fluorescent
bulb
Drawing
Description
Title your paper:
Atomic Emission Spectra
Goal: to understand what emission spectra are.
We will make a table in which we draw and describe several emission
spectra.
1. Put on the diffraction glasses and look outside
Find a rainbow. Draw and describe the spectrum you see.
2. Repeat with fluorescent lights on in the room.
3. Repeat with neon, helium, and five other elements
Source
Sunlight
Fluorescent
bulb
Drawing
Description
Quantum Mechanical Model
Energy quantized; comes in chunks.
A quantum is the amount of energy
needed to move from one energy level to
another.
Since the energy of an atom is never “in
between” there must be a quantum leap
in energy.
1926 Erwin Schrodinger equation
described the energy and position of
electrons in an atom
Quantum Mechanical Model
•Things that are very small
behave differently from things
big enough to see.
•The quantum mechanical model
is a mathematical solution
•It is not like anything you can
see.
Quantum Mechanical Model
•Has energy levels for electrons.
•Orbits are not circular.
•It can only tell us the
probability of finding an
electron a certain distance
from the nucleus.
Atomic Orbitals
•Energy levels (n=1, n=2…)
•Energy sublevels = different
shapes
•The first energy level
has one sublevel:
1s orbital -spherical
Atomic Orbitals
•The second energy level has
two sublevels, 2s
and 2p
There are 3 p-orbitals
Atomic Orbitals
•The third energy level has three
sublevels, 3s
3p
And 5 3d orbitals
py
Atomic Orbitals
•The forth energy level has four
sublevels, 4s
4p
4d orbitals
And seven 4f orbitals
Atomic Orbitals
The principal quantum number
(energy level) equals the
number sublevels
5.2 Electron Arrangement in
Atoms
Electron Configuration
Electrons and nucleus interact
to produce most stable
arrangement=
Lowest energy configuration
3 rules:
Aufbau Principle Electrons fill
the lowest energy orbitals first
Hydrogen
has 1
electron
1
1s
3 rules:
Pauli Exclusion Principal- two
electrons per orbital (one spin
up, one spin down)
Boron has 5
electrons
2
2
1s 2s
1
2p
3 rules:
Hund’s rule- In orbitals with
equal energy levels, arrange
spin to maximize electrons with
22s22p3
1s
the same spin
Nitrogen has 7 electrons
Hund’s Rule: Separate
the three 2p elecrons into
the three available 2p
orbitals to maximize the
electrons with the same
spin.
Conceptual Problem
Electron Configuration for Phosphorus (atomic # = 15)
1s2 2s2 2p6 3s2 3p3
Practice Problem
Electron Configuration for Carbon (atomic number = 6)
1s2 2s2 2p2
Practice Problem
Electron Configuration for Argon (atomic # = 18)
1s2 2s2 2p6 3s2 3p6
Practice Problem
Electron Configuration for Nickel (atomic # = 28)
1s2 2s2 2p6 3s2 3p6 3d8 4s2
Practice Problem
Electron Configuration for Boron (atomic # = 5)
1s2 2s2 2p1
How
many
unpair
ed
electro
ns?
1
Practice Problem
Electron Configuration for Silicon (atomic # = 14)
1s2 2s2 2p6 3s2 3p2
How
many
unpair
ed
electro
ns?
2
Exceptions to the Aubau Rule
Copper atomic number=29
1s2 2s2 2p6 3s2 3p6 3d9 4s2
This is
the
expected
electron
configura
tion
Exceptions to the Aubau Rule
Copper atomic number=29
1s2 2s2 2p6 3s2 3p6 3d10 4s1
HalfThis is
filled
and
the
filled
actual
sublevels
electron
are
more
configura
stable,
tion.if it
even
means
stealing an
electron
from a
nearby
sublevel
Exceptions to the Aubau Rule
Chromium atomic number=24
1s2 2s2 2p6 3s2 3p6 3d4 4s2
This is
the
expected
electron
configura
tion
Exceptions to the Aubau Rule
Chromium atomic number=24
1s2 2s2 2p6 3s2 3p6 3d5 4s1
HalfThis is
filled
and
the
filled
actual
sublevels
electron
are
more
configura
stable,
tion.if it
even
means
stealing an
electron
from a
nearby
sublevel