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```Unit-4 Atomic Structure and Quantum Theory
1. Bohr was the first to propose that the electrons were located in energy levels. A lower
case “n” is used to denote these principle energy levels (also called principle quantum
numbers). The principle energy levels are numbered, so that the level closest to the
nucleus is labeled n = 1. The next level is labeled n = 2 and so forth. Each principle energy
level had a certain energy value associated with the level. The closer the level was to the
nucleus, the lower the energy of the level. The farther away from the nucleus, the higher
the energy of that level in the atom. As long as the electrons were in these levels, the
electrons do not give off energy. The dark circle below represents the nucleus. The rings
around the nucleus represent the principle energy levels. Number the principle energy
levels starting with the one closest to the nucleus: n = 1, n = 2, n = 3 etc.
2. Electron Configuration and the Periodic Table
Each principle energy level can only hold so many electrons before the level is full. A quick
and easy way to determine the maximum number of electrons (max e-) that a principle
energy level can hold is given by the following:
max e- = 2 n2. First square the
principle energy level number (n) then multiply by 2.
Energy Level (n)
Maximum number of electrons (max e- = 2 n2)
1
2
3
4
5
6
Electrons are arranged around the nucleus by filling up the first principle energy level
(n=1), then the second energy level, etc. This is the electron configuration given on your
periodic table. The number of electrons are listed for each level with a dash between
levels: for oxygen (O) which has a total of 8 electrons, the configuration is 2–6
(2 electrons are located in the first principle energy level and 6 electrons are located in
the second principle energy level. Look up the electron configuration on the periodic table
for the element given and fill in the chart. Ca is done as an example.
Element
Ca
Na
F
B
Al
C
H
n=1
2
n=2
8
n =3
8
n=4
2
3. Drawing Bohr Diagrams of Atoms:
1) A circle is used for the nucleus- the # protons (# p or +) and the # of neutrons (#n)
are placed in the circle.
2) A ring is drawn around the nucleus for each energy level.
3) The electrons for each energy level are placed in pairs symmetrically around the nucleus
For F: atomic # = _____________
atomic mass = _____________
electron configuration: ______________
# p = _________ # n =____________
For Al : atomic # = _____________
atomic mass = _____________
electron configuration: ______________
# p = _________ # n =____________
Going Backwards: Determining the identity of an element from the Bohr diagram:
# p = _____________ # n =______________
atomic # = _____________
atomic mass = # p + # n = ________________
electron configuration:
# p = _____________ # n =______________
atomic # = _____________
atomic mass = # p + # n = ________________
electron configuration:
____________________________________
Isotopic Notation:
____________________________________
Isotopic Notation:
2. Ground and Excited States
The lowest possible energy state that an electron can occupy is called the __________
____________.
This is a very __________ condition.
The principle energy levels, which are occupied match those predicted by the electron
configuration on the periodic table. When electrons gain energy, the electrons move to higher
principle energy levels then they would normally occupy. This unstable situation is called the
_____________ _______________. The electrons will release the absorbed energy, often
seen as the bright line spectrum of the element, and fall back to the ground state.
I) How to tell when energy will be absorbed or released
The Principle Energy Level (n) changes:
If the number of the principle energy level (n) goes up, then energy is _____________
or ______________
n = 1 to n = 3 OR n = 3 to n = 4
If the number of the principle energy level (n) goes down, then energy is
_____________ or ______________
n = 2 to n = 1
OR n = 5 to n = 3
Determine if energy is added/absorbed (+E) or released/emitted (-E) for the following
transitions:
1) n = 1 to n = 2
______________ 6) n = 1 to n = 5
______________
2) n = 4 to n = 3
______________ 7) n = 4 to n = 2
______________
3) n = 2 to n = 1
______________ 8) n = 2 to n = 3
______________
4) n = 2 to n = 4
_____________ 9) n = 3 to n = 6
______________
5) n = 5 to n = 3
______________ 10) n = 6 to n = 2
______________
II) How do you tell the excited and ground state apart from the electron configuration??
Ground State: Matched the predicted electron configuration found on the periodic table.
In other words, it follows the order given
Ground State for Oxygen (O) on PT= 2 – 6
(8 total electrons)
Possible Excited State for Oxygen = 1 – 7 (still 8 total electrons)
The first energy level is not filled before moving into the second energy level.
The KEY here is that the configuration does not MATCH the one on the PT. Another
possible excited state for oxygen: 2 – 5 – 1 (still 8 total electrons)
For the following elements, fill in the chart and determine if the electron configuration is in
the GROUND STATE (GS) or EXCITED STATE (ES).
Element
Symbol
Electron
Configuration
C
2-4
F
1-8
Cl
2-8-7
B
2-1-1-1
Na
2-7-2
O
2-6
S
1-8-7
Zn
2-8-18-2
Br
1-7-17-8-1-1
Protons
Neutrons
Electrons
Principle
Energy
Levels
Ground or
excited
state?
Completely Filled vs. Occupied Principle Energy Levels
Occupied means that there is at least one electron in the Principle Energy Levels
Li: 2 – 1
has 2 occupied Principle Energy Levels
Completely Filled means that each level has its maximum number of electrons which can be
determined by the 2n2 rule.
Li: 2 – 1
has only 1 Completely Filled Principle Energy Level
Principle energy levels
Max number of e-
1
2n2 = 2(
)2 =
2
2n2 = 2(
)2 =
3
2n2 = 2(
)2 =
4
2n2 = 2(
)2 =
5
2n2 = 2(
)2 =
6
2n2 = 2(
)2 =
7
2n2 = 2(
)2 =
(2n2)
a) Copy the electron configuration from the Periodic Table
b) Determine the number of Occupied Principle Energy Levels (PEL)
c) Determine the number of Completely Filled Principle Energy Levels
Element
Electron Configuration
# Occupied PEL
# Completely
Filled PEL
C
Na
O
Cl
He
F
Ne
Si
Zn
Au
3. Bright Line Spectra and Introduction to Light
Every atom has many electron energy levels.
Electrons can drop from the excited state in many different pathways.
Each different drop in energy level gives off light with different amounts of energy, and
therefore wavelength.
The electrons will fall back down in the different ways, giving off different wavelengths of
light.
If you look at the element sample in the flame, it will appear to be glowing with a certain
color. If the light is broken up by a prism (as raindrops can break sunlight up into a rainbow)
and projected onto a white surface, the individual energies of light can be seen as bright
lines of different color.
These colored bright lines are different from element to element, making each spectrum
unique to a particular element.
This process of observing atomic and molecular spectra is called ______________
and the device used to observe the spectra is called a _____________________ .
Example bright-line spectrum for hydrogen (Numbers in parenthesis indicate the drop in
electron energy level that gives off the light, numbers in brackets indicate the wavelength in
nanometers (10-9 m) of the photons making up that bright line):
violet
violet-blue
alpha)
(62)
(52)
[410.2 nm] [434.0 nm]
green
(42)
[486.1 nm]
red (hydrogen(32)
[636.5 nm]
In order for the electron to return to a lower and more stable energy level, the added
energy must be given off. When the electrons return to the lower energy levels this
decreases the PE because the added energy is given off and the colors of the bright line
spectra are seen.
BRIGHT LINE SPECTRUM
BRIGHT LINE SPECTRA are produced when “electrons in the EXCITED STATE” fall back to
lower energy levels of the GROUND STATE. Unlike the continuous spectrum of sunlight, only
certain colors will be present in the BRIGHT LINE SPECTRA. The BRIGHT LINE SPECTRUM is
like a “fingerprint” of the element that produced the spectrum. Like a fingerprint, the BRIGHT
LINE SPECTRA can be used to identify the element. When viewed with a spectroscope, the
individual bands of colors in the BRIGHT LINE SPECTRUM can be seen and the wavelength of
each band determined.
1. Below are the BRIGHT LINE SPECTRA of three elements. From the position of the lines
determine which element is the unknown. (HINT: Match up the lines present in the unknown
with the three known elements.) Unknown element = ____________
Element X
Element Y
Element Z
Unknown
2. Which of the two elements above are present in the BRIGHT LINE SPECTRUM given
below? (HINT: Match up the lines present with the three known elements. Only two
patterns should match perfectly.) ____________ and _____________
Introduction to Light
Visible Light (energy we see with): part of the Electromagnetic Spectrum
1. Two theories to explain light’s behavior:
Waves
Particles of Packets of Energy
There was evidence for both models so the two theories were put together!!
Light: QUANTUM THEORY OF LIGHT
a) packets or bundles of energy called _________________ or ______________
b) travel in wave-like fashion
c) produced when electrons drop from ______________ energy levels to
_____________ energy levels (the greater the drop, the greater the energy
the light has)
2. Properties of Light
Wavelength ( ) - ____________________________________________
Frequency (F) - ______________________________________________
(units: Cycles / second OR Hertz)
Energy (E) - __________________________________
Speed (velocity) – same for all electromagnetic radiation _______________
Relationships:
Frequency and Energy: Type _________________
F ______, E ________ or F ______, E ________
Frequency and Wavelength: Type _________________
F ______,
________ or F ______,
________
Wavelength and Energy: Type _________________
______, E ________ or
______, E ________
3. The Rainbow: A Continuous Spectrum
R
Long
Low F
Low E
O
Y
G
B
I
V
LONG STEM RED ROSES: All “L’s” go together with RED
Short
High F
High E
B. Continuous Spectrum
when radiation from the sunlight passes through a prism, a rainbow – a
spectrum of colors – is seen
the colors are not separated from one another but blend together
due to the overlap of the line spectra of the 67 different elements in
the sun
lightbulb
C. Bright Line Spectrum
when radiation from an excited atom (element) passes through a
prism, the radiation is separated into various wavelengths and colors
Colors are not blended – spectrum is discontinuous – and you observe
lines of color at different locations
Flame
Updating The BoHR MoDeL: QuAnTuM oR WaVe MeChAnIcaL MoDeL
The Bohr Model was very good at explaining many things about atoms and electrons and
how they behaved. Just like your wardrobe needs updating, so did the model of the atom.
A revolution in physics occurred in the early 1900’s when experiments showed that matter,
just like light energy, could have a dual nature…it can act as a particle or a wave.
1. The Wave Mechanical Model
1.__________________________________________________________________
2.__________________________________________________________________
3.__________________________________________________________________
Using complicated math, this math showed that electrons were not moving in definite fixed
orbits like planets but had distinct amounts of energy.
a. These electrons were found in “regions” or “areas” around the
nucleus called orbitals.
b. An orbital is defined as a region in which an electron with a
particular amount of energy is most likely to be found
The probability of finding an electron near the nucleus is _____________
The probability of finding an electron far away from the nucleus is __________
B. Four “Quantum Numbers” are used to describe the location of the electron.
a. Principle Energy Levels (PEL) get divided into sublevels. (PEL are
also called the PRINCIPLE QUANTUM NUMBER. )
b. These sublevels are designated: s, p ,d and f
c. The number of sublevels is determined by the number of the
princple energy level (n): n = # sublevels
d. Within the same principle energy level: the sublevels have different
energies:
s  p  d f
lowest energy

highest energy
e. The sublevels are divided into orbitals (rooms) where electrons are
found. Each oribital can hold 2 electrons.
f. Orbitals are “regions” or “areas” around the nucleus
g. The number of orbitals depends on the sublevel type:
Sublevel
# orbitals
x2=
# of electrons in sublevel
s
x2=
p
x2=
d
x2=
f
x2=
There are two ways to show the Quantum Atom Electron Configuration. One way shows
the number of electrons in the Principle Quantum Level (PEL) and Sublevel type (letter).
The other way, known as BOX DIAGRAMS, show all four quantum numbers by using boxes
for orbitals and arrows for electrons to show the opposite spin of the electrons. Using the
electron configuration from the periodic table, we will do both types.
For Hydrogen: EC form PT is 1
Quantum Atom
Box Diagram
1s1
PEL
1s
# of electrons
SUBLEVEL
PEL
ORBITAL
SUBLEVEL
ARROW SHOW SPIN
RULES TO REMEMBER WHEN DRAWING BOX DIAGRAMS;
* S has one orbital so draw 1 box
* P has 3 orbitals so draw 3 boxes,
* Each “P” box in a sublevel must have an electron before pairing up because each
orbital has the same energy as the other orbitals (Boxes)
* Electrons spin in opposite directions!
Principle Energy Level
Sublevel(s) Present
Orbitals Present
Total # of e per PEL
1
2
3
4
n = _____________________ & n = ___________________________
Element
Elec. Config.
From P T
Quantum Atom
E. C.
Box Diagrams
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
After argon (Ar) the energy levels begin to overlap and things get complicated. We will
not be doing elements above Calcium. However….
The Kernel, Valence Electrons and Lewis (Electron) Dot Diagrams
A very powerful tool for showing how different elements bond together, called Lewis
(Electron) Dot Diagrams, was developed by a chemist named Lewis. Many times in
chemical bonding, only the electrons on the outside or in the highest energy level, called
the VALENCE SHELL, actually become involved in bonding to form compounds. We are
going to look at the Lewis (Electron) Dot Diagrams and elements in this next section.
Valence Shell: Outermost energy level of an element (highest number)
Oxygen: O 2-6 the 2nd principle energy level is the valence shell
Valence Electrons (VE): electrons located in the outermost energy level
- these are the electrons which are the furthest to the right in the electron
configuration found on the PT
- Oxygen: O 2-6 the 6 electrons are the valence electrons
- In a dot diagram, the VE are represented by dots:
Kernel: This is the nucleus and all the other electrons located in the inner energy levels
In a dot diagram, the chemical symbol represents the kernel: O
The valence shell always contains only 4 orbitals. These orbitals are represented by the
four sides around the symbol. The top side represents one orbital which has less energy
than the other three orbitals. The other three orbitals have the same amount of energy.
Frist 2 electrons are always placed first in this top orbital. Going around clockwise, one
electron is placed on each remaining side until you need to pair the electrons up.
Orbital
lower energy
than others
O
Orbitals with
the same energy
O
Orbitals with
the same energy
The Periodic Table:
1) Look up and write down the electron configuration, then circle the valence electrons
2) Try your hand at drawing dot diagrams
Symbol
Electron
Configuration
Li
Be
B
C
N
O
F
Ne
Li
Be
B
C
N
O
F
Ne
Dot Diagram
Guided Practice: Complete the following chart
Symbol
Electron Configuration (Circle
Number of VE
the Valence Electrons (VE)
Dot Diagram
S
Se
Na
K
Si
Ge
Cl
Br
Do you see any patterns in the relationship between the diagrams and the element’s
position on the periodic table?
```