Download Chemistry Unit IV – The Electron

IV – The Electron
Chemistry Student Notes
Chemistry Unit IV – The Electron
PRE-TEST QUESTIONS
1. What is the chemical symbol for Tungsten?
2. Define Atomic Number.
3. How many neutrons does silver have?
4. How many electrons does silver have?
5. Who is credited with discovering the electron?
I. Light and Waves
A. What is light?
1. Sir Isaac Newton (1642 – 1727) tried to explain what was known about the behavior of
light by assuming light consists of __________________.
a. Random Fact: Newton is considered the father of Classical Physics, or Newtonian
Physics. He studied the motion and forces affecting large objects in our everyday
world, like the motion of projectiles, forces that affect motion and so on.
2. However, by the year 1900, there was enough experimental evidence to convince
scientists that light consists of __________________ instead of particles.
B. What are waves?
1. There are several measurable characteristics.
a. The __________________ __________________ is a straight line around which the
wave oscillates. Kind of the ‘mid-point’, if you will.
b. The __________________ is the high point on a wave.
c. The __________________ is the low point on a wave.
d. The __________________ of a wave is the wave’s height from zero (rest position) to
the crest.
e. The __________________ ______, represented by λ (the Greek letter lambda), is the
__________________ between crests or any two identical points on two waves.
f. The __________________ ______, represent by ν (the Greek letter nu), is the number
of wave cycles to pass a given point per unit of time. (Sometimes represented by f
also)
i. The units of frequency are usually cycler per __________________.
ii. The SI unit of cycles per second is called __________________ ______.
2. Draw a wave and label its parts
C. What is the relationship between wavelength and light?
1.
2. c = speed of light, __________________
3. The wavelength and frequency of light are _______________ proportional to each other.
a. As λ increases, the ν __________________.
b. As λ decreases, the ν __________________.
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Unit IV Notes, Page 1
IV – The Electron
Chemistry Student Notes
D. What does light consist of?
1. According to the wave model, light consists of electromagnetic __________________.
2. Electromagnetic radiation includes __________________ waves,
__________________, __________________ light, __________________ light,
__________________, __________________, and __________________ rays.
a. Draw the EMR in order. Indicate wavelength, frequency and energy.
i. Note that each of these travel at the speed of light, __________________
b. Note that visible (white) light is a very small range of the entire EMR spectrum and
consists of all the colors we can see.
i. When light passes through a prism, the slower moving red slows down more than
the faster-moving blue, which results in a splitting of light into a spectrum.
ii. A __________________ is a range of wavelengths of EMR (such as all the colors
of visible light).
iii. The colors of light can be remembered by the mnemonic:
EXAMPLE IV-01: Wavelength from frequency
Calculate the wavelength of the light emitted if the frequency of the wave is 4.70 × 1014 Hz.
II.
The Experiments That Led to the Quantum Theory
A. The Photoelectric Effect
1. At the time of Rutherford, electrons were pictured as __________________, and light
was pictured as __________________.
a. But! This doesn’t explain all the properties of __________________!
2. When light was shined onto a piece of metal, it ejected __________________ in
something referred to as the photoelectric effect.
a. Scientists noted the light had to be a
__________________ frequency for the
electrons to be ejected.
i. This was called the ______________
_______________.
b. Any light _____________ that frequency,
the electrons would not be ejected.
c. The theory was that the light of any
frequency could do this, though. Scientists
wanted to know why was there a
________________ frequency required?
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Unit IV Notes, Page 2
IV – The Electron
Chemistry Student Notes
B. Max Planck and Quanta
1. Max Planck, in 1900, noticed that when a hot object __________________ energy, it
does not do so continuously.
a. Waves, however, do give off energy _____________________.
2. Since the energy was not emitted continuously, Planck called these “packets of energy”
__________________.
a. A __________________ is a finite, minimum quantity of energy that can be gained or
lost by an atom.
b. This is the beginnings of “quantum theory.”
c. Since quanta are the __________________ units of energy possible, energy exists
only in __________________ of these quanta.
3. Planck proposed the following relationship: E = hν
a. Energy = __________________ (measured in __________________)
b. h = __________________ constant, __________________
c. ν = __________________ of radiation emitted
4. You do not need to worry with solving for energy, but you DO need to understand that
there is a __________________ relationship between frequency and energy.
a. A high frequency correlates to a __________________ energy.
b. A low frequency correlates to a __________________ energy.
C. Einstein and Photons
1. In 1905, __________________ expanded on this idea of quanta.
2. He proposed the radical idea of a dual wave-particle nature for light.
a. While light exhibits many wave-like properties, it can also be thought of as a stream
of particles he called photons.
3. A __________________ is a quantum of light – a particle of radiation.
a. Has __________________ rest mass and carries a quantum of energy.
b.
4. According to this theory proposed by Einstein, radiation is absorbed and emitted only in
whole numbers of photons.
5. Different metals are attracted at different strengths to their electrons. Thus, different
metals require different minimum energies to eject their electrons.
a. This explained the photoelectric effect and gave credibility to quantum theory.
b. PREVIEW: In the next Unit on the Periodic Table, we will be discussing this. It is
called “Ionization Energy”.
6. Using quantum theory, Einstein was able to explain what classical physics could not.
a. A photon is a “packet of light” or a __________________ of energy.
b. A photon will knock out an electron which has the __________________ amount of
energy.
ACTIVITY IV-01: Spectroscopy
Turn to the Spectroscopy sheet in your notes and observe the emission spectrum for
the various elements.
D. Hydrogen Atom Line Emission Spectrum
1. A __________________ tube was filled with Hydrogen gas under low pressure and then
an electric current was applied to it.
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Unit IV Notes, Page 3
IV – The Electron
Chemistry Student Notes
2. A __________________ glow was emitted from the tube.
3. When this light was passed through a prism, it split into very distinct wavelengths.
4. Scientists had expected a __________________ spectrum, but had only seen very
__________________ wavelengths.
a. WHY???
E. Atomic Spectra
1. Light emitted by atoms consists of a mixture of only specific frequencies.
a. Thus, it forms very specific lines which are called an __________________
__________________ __________________.
b. The atomic emission spectrum for each element is like a person’s fingerprint – unique
and helps identify the element.
c. This is how scientists can determine what elements are in the stars in space.
d. But why only specific frequencies?
F. Summarizing…
1. What are the properties of waves and particles and how are they alike and different?
WAVES
III.
PARTICLES
Electrons in Atoms
A. Review of Atomic Models
1. Dalton proposed that matter was made of basic particles he called __________________.
2. Thomson discovered the electron. He proposed the atom was a positive sphere with
negative __________________ dispersed throughout.
3. Rutherford discovered the nucleus. He proposed the atom had a dense, positive
__________________ that occupied very little space. Most of the space of the atom was
occupied by electrons, moving around the nucleus like planets around the sun (only not in
such specific paths).
a. However, this model does not explain the __________________ properties of
elements.
b. e.g., Why do metals and other objects change __________________ when heated?
B. The Bohr Model
1. Niels Bohr (1885-1962) was a young Danish physicist and a student of Rutherford.
2. He took these new ideas about how the energy of an atom changes when it absorbs or
emits light to help devise a new model of the atom.
a. He used the simplest element, __________________, and its emission spectrum.
3. He proposed that electrons are found only in specific circular paths, or
__________________, around the nucleus.
a. Sometimes referred to as the __________________ model or __________________
__________________ model because the electrons orbit the nucleus in specific paths
much like the planets orbit the sun.
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Unit IV Notes, Page 4
IV – The Electron
Chemistry Student Notes
Draw Bohr’s Model of the Atom
4. Each possible electron orbit in Bohr’s model has a fixed
energy, called ________________ ________________.
a. Like the rungs of a ladder, but they’re not spaced
evenly between one another.
b. The closer they are to the nucleus, the farther
________________ they are. The farther from the
nucleus, the ________________ they are together.
c. In order to move from one level to the next, must
move just the right distance.
5. A ________________ of energy, then, is the amount of
energy required to move an electron from one energy
level to another energy level.
a. “Quantum leap”
6. Note that because the energy levels get ________________ as they get farther from the
nucleus, the amount of energy an electron gains or loses in the atom is not always the
________________.
C. Absorption and Emission
1. Each orbit of the atom has a specific energy.
a. In order to change energy levels, the electron must ________________ or
________________ energy equal to the ______________________.
2. The ________________ ________________ is the ________________ possible energy
of an electron described by quantum mechanics.
3. Absorption
a. Electrons can gain an amount (quanta) of energy equal to the difference between
energy levels and “leap” from a lower energy level (near the nucleus) to a higher
energy level.
b. This is called an ________________ ________________.
c. Draw Absorption…
4. Emission
a. The excited state is unstable for an electron, so when the electron goes from the
excited state back down to the ground state, it has to ________________ that excess
energy it acquired in going from ground state to excited state in the first place.
b. This energy is given off in the form of electromagnetic radiation (______________).
c. Draw Emission…
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Unit IV Notes, Page 5
IV – The Electron
Chemistry Student Notes
D. Back to the Hydrogen spectrum
1. Bohr used the frequencies of the colors of light emitted by the hydrogen atom to calculate
the __________________________ of the light emitted.
2. Why did we see only certain colors of light? Shouldn’t there be many more possibilities?
a. We can’t ________________ all the light emitted by the hydrogen atom.
3. Lines in the _______________ region (which we _____ see) are called the Lyman series.
a. Excited states go back to ground state n = ___.
4. Lines in the ________________ light region (we _____ see) are called the Balmer series.
a. Excited states go back to ground state n = ___.
5. Lines in the ________________ region (we ________ see) are called the Paschen series.
a. Excited states go back to ground state n = ___.
6. Spectral lines for the transitions from higher energy levels to n = 4 and n = 5 also exist.
a. VERY close together in energy (because the energy levels are much closer together)
b. There is an upper limit to the frequency of light that can be emitted, so an electron
with enough energy completely escapes the atom.
E. The Bohr Model Diagram
1. Please note the wavelength is given in nanometers (nm).
2. To convert wavelength from nm into m, simply tack on “________________”.
3. Notice the various regions.
a. All electrons that go to n = 1 are in the ________________ region of EMR.
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Unit IV Notes, Page 6
IV – The Electron
Chemistry Student Notes
b. All electrons that go to n = 2 are in the ________________ region of EMR.
c. All electrons that go to n = 3 are in the ________________ region of EMR.
4. Notice the lines are ______________ for Ultraviolet, thus there is ____________ energy.
a. Likewise, IR lines are ______________, so they have the ______________ energy.
5. Consider an electron falling from n = ___ to n = ___ energy level.
a. What is its wavelength?
b. What color is this light?
F. Bohr’s Model: The Result
1. Bohr’s model was great at describing ________________!
2. But, anything with more than one electron, it was ________________.
3. Thus, the quantum model of the atom would be born.
IV. Electrons as Waves
A. Another way to think about electrons
1. To scientists of the early 20th century, Bohr’s model contradicted common sense.
2. Why did Hydrogen’s electrons exist only in certain orbits with _____________ energies?
3. Why not limitless orbits, with slightly ___________________ energies?
4. At the time, it was known that light could behave both as a ______ and a ____________.
5. French scientist Louis de Broglie asked the same question about electrons.
a. Could they have a _____________ nature too?
6. De Broglie pointed out that Bohr’s electron orbits were similar to the behavior of waves.
a. For example, a wave confined to a space is known to have only specific
___________________.
b. So, you can think of electrons as ___________________ confined to the space
around the ___________________ within the ___________________.
c. Electrons, then, can exist only at certain ___________________.
d. E = hν relates the frequency to ___________________.
e. This is the ___________________ that Bohr had defined!
B. Proof for electrons as waves
1. Waves can be bent.
a. ___________________ can be bent!
2. Diffraction is the bending of a wave as it passes the edge of an object, such as the edge
of atoms in a crystal.
a. ___________________ can be diffracted!
3. Waves can interfere with one another when they overlap.
a. ___________________ can interfere with one another!
4. Davisson and Germer at Bell Labs in New Jersey performed an experiment showing
electrons reflecting off metal in curious patterns, just like X-Rays do!
a. ___________________ must be like waves!
b. Here was the proof for De Broglie’s theory of electrons behaving as waves!
c. Nobel Prizes for De Broglie and Davisson (and George Paget Thomson).
5. The answers to De Broglie’s equation show that all moving objects have wavelike
behavior, but the mass of the object being measured must be very ___________________
in order for its wavelength to be large enough to observe.
6. Classical (Newtonian) mechanics adequately describes the motions of bodies much
___________________ than atoms, while quantum mechanics describes the motions of
subatomic particles and atoms as waves.
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Unit IV Notes, Page 7
IV – The Electron
Chemistry Student Notes
C. Heisenberg
1. German physicist Werner Heisenberg explained another feature of quantum mechanics.
2. The Heisenberg Uncertainty Principle states that it is ___________________ to know
exactly both the ___________________ and the ___________________ of a particle at
the same time.
a. Concerns only small, microscopic particles such as electrons and not macroscopic
objects like cars or baseballs.
3. In order to see objects, we must bounce light off of them. For electrons, when the light
hits them, because they’re so tiny, it affects their ___________________. It is
unpredictable how this affects the motion, so we can never be certain about both its
momentum and position at the same time.
D. Schrödinger
1. In 1926, Austrian physicist Erwin Schrödinger created an equation that used electrons as
___________________.
2. Solutions to the Schrödinger wave equations are known as wave functions.
3. The quantum mechanical model is the modern description of the electrons in atoms that
originates from Schrödinger’s wave equation.
a. It only gives the ___________________ of finding the electrons.
E. The Quantum Mechanical Model
1. Like the Bohr model, the quantum mechanical model of the atom restricts the energy of
electrons to certain values.
2. Unlike the Bohr model, the quantum mechanical does not involve an exact path the
electron follows. Instead, it occupies a region of space.
3. The quantum mechanical model determines the allowed energies an electron can have
and how likely it is to find the electron within various ____________ around the nucleus.
4. The probability of finding the electron within the atom can be represented by a couple of
different models:
a. A probability cloud, where the _____________ areas
show where electrons are more likely to be found.
b. An enclosed, solid sphere that shows it is somewhere
___________________ that sphere.
V.
Quantum Numbers
A. Defining the Electrons
1. Within each energy level, the Schrödinger equation also leads to a mathematical
expression, called an atomic ___________________, describing the probability of
finding an electron at various locations around the nucleus.
2. An ___________________ ___________________ is often thought of as a region of
space in which there is a _________ probability of finding an electron.
3. Each electron is defined mathematically by four numbers called Quantum Numbers.
a. The ___________________ Quantum Number
b. The _______________ _______________ (or _______________) Quantum Number
c. The ___________________ Quantum Number
d. The ___________________ Quantum Number
B. The Principal Quantum Number
1. Represented by the symbol ____.
2. Corresponds to the main energy ________________ (________________) of the atom.
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Unit IV Notes, Page 8
IV – The Electron
Chemistry Student Notes
a. The maximum we deal with is usually ____.
b. Within each energy level, there can be multiple sublevels, orbitals and electrons.
3. Values of n are positive whole numbers: ___________________
4. The distance from the nucleus ___________________ as n increases.
5. Total number of orbitals possible = _____ (generally not to exceed _____)
C. The Angular Momentum (Azimuthal) Quantum Number
1. Represented by the symbol ____.
2. Corresponds to an energy ___________________ (or ___________________).
3. Gives the ___________________ of the orbital.
Value of l
Orbital
Designation
Origin Name
(Don’t have to know)
Mnemonic
4. Each energy sublevel corresponds to an orbital of a different ___________________,
which describes where the electron is likely to be found.
5. The value of l is equal to ____________, not to exceed ____.
a. If n = 1, then l can be _______.
b. If n = 2, then l can be _______ and _______.
c. If n = 3, then l can be _______, _______, and _______.
d. If n = 4, then l can be _______, _______, _______ and _______.
6. The number of sublevels possible is ___________________ to the principal quantum
number, not to exceed ___________________.
a. This is because they are only four: ___________________
7. Sublevels are designated by using the principal quantum number (energy level) first, then
the letter of the shape.
a. The s orbital in the first energy level is called _____.
b. The s orbital in the second energy level is called ______.
c. The p orbitals in the second energy level is called ______.
D. The Magnetic Quantum Number
1. Represented by the symbol _____.
2. Corresponds to the ___________________ of the orbital within the sublevel.
a. Gives the ___________________ along the axes.
3. The mathematical value for the Magnetic Q.N. is equal to a range of whole integers from
______ to _______, including the number _____.
Orbital
l=
Possible ml values
Number of orbitals
designation
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Unit IV Notes, Page 9
IV – The Electron
Chemistry Student Notes
4. p orbitals are designated by the axis they are on.
a. ______, ______, and _______, are the three designations for p orbitals.
5. The number of total orbitals within each energy level = ______.
E. The Spin Quantum Number
1. Represented by the symbol ______.
2. Corresponds to the ___________________ of the spin of the electron.
3. Only ___________________ electrons exist within each orbital.
4. Only two values for Spin Q.N.: ___________________.
F. Energy Levels
1. There are ___________________ energy levels.
a. Within those energy levels, there are up to ___________________ types of sublevels.
i. Within those sublevels, there are ___, ___, ___, or ___ orbitals.
(a) Within those orbitals, there can be only ___ electrons.
2. So…
a. The first main energy level can hold ____ electrons.
i. ___ orbital has 1 orbital, which can hold ____ electrons.
b. The second main energy level can hold ____ electrons.
i. ___ orbital has 1 orbital, which can hold ____ electrons.
ii. ___ orbital has 3 orbitals, which can hold ____ electrons.
c. The third main energy level can hold ____ electrons.
i. ___ orbital has 1 orbital, which can hold ____ electrons.
ii. ___ orbital has 3 orbitals, which can hold ____ electrons.
iii. ___ orbital has 5 orbitals, which can hold ____ electrons.
d. The fourth main energy level can hold ____ electrons.
i. ___ orbital has 1 orbital, which can hold ____ electrons.
ii. ___ orbital has 3 orbitals, which can hold ____ electrons.
iii. ___ orbital has 5 orbitals, which can hold ____ electrons.
iv. ___ orbital has 7 orbitals, which can hold ____ electrons.
3. The above can be more easily determined by remembering that the maximum number of
electrons each energy level can hold = ______.
a. The number of total orbitals within each energy level = ____.
VI. Electron Configuration Rules
A. Introduction
1. Most things in nature tend to prefer the ___________________ amount of energy as this
is most ___________________.
2. The electron configuration is the arrangement of electrons of an atom in its
___________________ state into various orbitals around the nuclei of atoms.
3. Three rules—the Aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell
you how to find the electron configurations of atoms.
B. The Aufbau Principle
1. The ___________________ ___________________ is the rule that electrons occupy the
orbitals of ___________________ ___________________ ___________________.
a. These are ___________________ to the nucleus.
2. Up until 3p, the orbitals are filled in order of number and letter, but then the energies get
close together, so the filling order is not as clear.
a. Within an energy level, s is always the lowest energy subshell, then p, then d, then f.
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Unit IV Notes, Page 10
IV – The Electron
Chemistry Student Notes
b. The problem comes when you compare energy levels.
3. Copy the filling chart:
C. The Pauli Exclusion Principle
1. According to the _________________ ___________________ ___________________,
an atomic orbital may describe, at most, ___________________ electrons and that no
two electrons can share the same set of ______ quantum numbers within the same atom.
2. To occupy the same orbital, two electrons must have ___________________ spins.
3. In orbital notation, different electron spins are represented by different arrows:
D. Hund’s Rule
1. ___________________ ___________________ states that electrons occupy orbitals of
degenerate (the ___________________) energy in a way that makes the number of
electrons with the same spin direction as ___________________ as possible.
2. Orbitals of ___________________ energy are occupied by one electron each
___________________ any one orbital receives a ___________________ electron, and
all electrons in singly occupied orbitals must have the ___________________ spin.
3. Each of the three p orbitals gets one electron each before any one gets a second electron.
4. Each of these unpaired p electrons will have the same spin, either +½ or – ½.
(a)
(b)
(c)
VII. Electron Configurations
A. Electron Configurations
1. There are three ways to express electron configurations.
B. Orbital Notation
1. The most visual notation.
2. Each orbital is represented by a horizontal ___________________ (or sometimes boxes).
3. The energy level and orbital designation are usually written under the line or grouped
together if it has more than one orbital (p, d, f)
4. Electrons are represented by ___________________. Up or down differentiates spin.
5. Hydrogen:
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Unit IV Notes, Page 11
IV – The Electron
Chemistry Student Notes
6. Helium:
7. Lithium:
8. Beryllium:
9. Boron:
10. Carbon:
C. Electron Configuration Notation
1. Most commonly used method.
2. Does not display each individual orbital.
3. Uses the energy ___________________ and ___________________ designation, with a
___________________ to indicate how many _________________ are in each sublevel.
a. s, with one orbital, can hold ____ electrons.
b. p, with three orbitals, can hold ____ electrons.
c. d, with five orbitals, can hold ______ electrons.
d. f, with seven orbitals, can hold ______ electrons.
4. Hydrogen:
5. Helium:
6. Lithium:
7. Beryllium:
8. Boron:
9. Carbon:
D. Noble Gas Notation
1. Noble Gases are found in Group 18 of the P.T.
2. Remember: Noble gases are unique in that their outer energy level of ____ and ____ are
full of electrons, so they are ___________________ and ___________________.
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Unit IV Notes, Page 12
IV – The Electron
Chemistry Student Notes
3. Thus, if an element occurs after a noble gas (reading the P.T. in order of Atomic
Number), then the Noble Gas can be substituted into the electron configuration notation
to represent inner energy levels that are already filled.
4. This helps to focus attention on the ___________________ energy levels, because they
are the ones that are ___________________ to how atoms ___________________.
5. Hydrogen:
6. Helium:
7. Lithium:
8. Beryllium:
9. Boron:
10. Carbon:
11. Neon:
12. Sodium:
E. Exceptional Electron Configurations
1. Copper is a special case in filling electron configurations.
a. It should fill ___________________
b. However, it is more stable to have a _________ d than a ________ s, so the d takes 1
electron from the 4s to fill the 3d orbital.
c. Copper, instead, fills: ___________________
d. This looks like it breaks the Aufbau Principle, but it is actually more stable this way!
2. Why is this?
a. ___________________ energy sublevels are ___________________ stable than
partially filled sublevels.
b. Some actual electron configurations differ from those assigned using the Aufbau
principle because half-filled sublevels are not as stable as filled sublevels but they are
more stable than other configurations.
3. Chromium is another good example.
a. It is in Group 6, so it seems it should fill: ___________________
b. However, it is more stable to have half of a filled d orbital than to have the 4s full
(remember, s orbitals are fairly low energy compared to the others).
c. Chromium, instead, fills: ___________________
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Unit IV Notes, Page 13
IV – The Electron
Chemistry Student Notes
Spectroscopy
Name: _____________________
Date: ___________ Period: ____
Directions: Place lines in the emission spectrum boxes below corresponding to what you see using the glasses or
spectrscopes. Put general lines for each of the colors. Remember, Red is on the right with 700 nm and violet is on
the left, with 400 nm (380 nm). Try and use colored pencils or crayons to sketch what you see.
Continuous Spectrum (look at the fluorescent lights in the room)
↑ 400 nm (Violet)
↑ 500 nm
↑ 600 nm
700 nm (red) ↑
↑ 500 nm
↑ 600 nm
700 nm (red) ↑
↑ 500 nm
↑ 600 nm
700 nm (red) ↑
↑ 500 nm
↑ 600 nm
700 nm (red) ↑
↑ 500 nm
↑ 600 nm
700 nm (red) ↑
↑ 500 nm
↑ 600 nm
700 nm (red) ↑
Gas ___________________
↑ 400 nm (Violet)
Gas ___________________
↑ 400 nm (Violet)
Gas ___________________
↑ 400 nm (Violet)
Gas ___________________
↑ 400 nm (Violet)
Gas ___________________
↑ 400 nm (Violet)
4.
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Unit IV Notes, Page 14