UNIT 4: ELECTRONS IN ATOMS
1. All the early research on the atom focused on the nucleus. The models lacked detail about how the electrons occupy the space surrounding the nucleus.
2. Rutherford’s model did not explain how the electrons were arranged in space. Nor did it address the question of why the negatively charged electrons are not pulled into the nucleus.
3. Chemists found that it did not account for the differences in chemical behavior among the various elements.
4. Scientists observed that certain elements emitted light when heated in a flame. When they analyzed the light, they found that an element’s chemical behavior is related to the arrangement of the electrons in the atom.
Wave Nature of Light (pp. 324 - 325)
1. ______is a form of energy that exhibits wavelike behavior as it travels through space.
2. Waves have three primary characteristics:
______- distance between equivalent points on a continuous wave (usually crests).
______- number of waves that pass a given point per second.
______- the wave’s height from the origin to either a crest or a trough.
3. The three characteristics are related to one another through the equation:
4. All electromagnetic waves travel at the same speed, but they have different wavelengths and frequencies. This gives the various waves different properties.
5. Visible light is just a small portion of the overall electromagnetic spectrum.
6. The electromagnetic spectrum encompasses all forms of electromagnetic radiation with the only differences in the types of radiation being their frequencies and wavelengths.
7. The energy of the radiation increases with increasing frequency.
***** Determine the wavelength of a microwave that has a frequency of 3.44 x 109 Hz.
***** Determine the frequency of green light. Its wavelength is 4.90 x 10-7m.
Particle Nature of Light
1. Considering light as a wave does explain much of its everyday behavior.
2. It fails to describe some important aspects of light’s interactions with matter.
3. The wave model cannot explain why heated objects emit only certain frequencies of light at a given temperature.
4. It cannot explain why some metals emit electrons when colored light of a specific frequency shines on them.
5. Either another model or a revision of the current model of light was needed.
The Quantum Concept
1. Hot objects emit a glowing light.
2. Temperature is a measure of the average kinetic energy of its particles. As a metal gets hotter, it possesses a greater amount of energy and emits different colors of light.
3. The colors correspond to different frequencies and wavelengths. The wave model could not explain this.
4. Max Planck studied the phenomenon and it led him to a startling conclusion: matter can gain or lose energy only in small, specific amounts called ______.
5. A ______is the minimum amount of energy that can be gained or lost by an atom.
6. Planck demonstrated a mathematical relationship between the energy and the frequency of the emitted radiation:
7. According to this theory, for a given frequency, matter can emit or absorb energy only in whole-number multiples of hν.
8. Matter can have only certain amounts of energy – quantities of energy between these values do not exist.
The Photoelectric Effect
1. In the ______electrons, called photoelectrons, are emitted from a metal’s surface when light of a certain frequency shines on the surface.
2. No matter how intense or how long it shines, light with a frequency below that required will not cause electrons to be ejected from the surface of the metal.
3. Once the correct frequency is reached, even a dim light will cause the electron to be ejected.
4. Einstein proposed that electromagnetic radiation has both wavelike and particle like natures.
5. While a beam of light has many wavelike characteristics, it can also be thought of as a stream of tiny particles, or bundles of energy, called photons.
6. A ______is a particle of electromagnetic radiation with no mass that carries a quantum of energy.
7. The energy of a photon is given by:
***** Determine the energy of a photon from the violet portion of the rainbow if it has a frequency of
7.23 x 1014 s-1.
Atomic Emission Spectra (pp. 328 - 330)
1. The ______of an element is the set of frequencies of the electromagnetic spectrum emitted y atoms of the element.
2. The element’s atomic emission spectrum consists of several individual lines of color, not a continuous range of colors as seen in the visible spectrum.
3. Each element’s atomic emission spectrum is unique and can be used to determine if that element is part of an unknown compound.
4. The fact that only certain colors appear in an element’s atomic emission spectrum means that only certain specific frequencies of light are emitted. Because those emitted frequencies of light are related to specific energy levels, it can be concluded that only photons having certain specific energies are emitted.
5. Scientists were still confused. They had not expected line spectra. They had expected to observe the emission of a continuous series of colors and energies as excited electrons lost energy and spiraled toward the nucleus. But, the electrons weren’t doing this.
QUANTUM THEORY AND THE ATOM
Bohr Model of the Atom (pp.330 - 331)
1. Bohr attempted to explain why the emission spectra of the hydrogen atom were not continuous.
2. He proposed, using Planck’s and Einstein’s concepts, that the hydrogen atom has only certain allowable energy states for its electron.
3. He suggested that the single electron in a hydrogen atom moves around the nucleus in only certain allowed circular orbits.
4. The smaller the electron’s orbit, the lower the atom’s energy state, or energy level. The larger the electron’s orbit, the higher the atom’s energy state, or energy level.
5. Bohr assigned a quantum number, _____, to each orbit. The higher the value of n corresponds to a higher energy level.
Hydrogen’s Line Spectrum
1. Bohr suggested that when hydrogen is in the ground state, the electron is in the n = 1 orbit. In this state the atom does not radiate energy.
2. If energy is added from an outside source, the electron moves to a higher – energy orbit. The atom is now in an excited state.
3. Since nature likes being in the lowest energy state possible, the atom will return to the ground state. The electron will return to the n = 1 orbit and in the process emit a photon of energy.
4. Since only certain atomic energies are possible, only certain frequencies of electromagnetic radiation can be emitted.
5. Bohr’s model explained the hydrogen atom, but failed to explain the emission spectrum of any other element.
6. Bohr’s model did not fully account for the chemical behavior of atoms.
7. Also, unlike the models developed to this point in time, evidence was mounting that electrons do not move around the nucleus in circular orbits.
THE QUANTUM MECHANICAL MODEL OF THE ATOM
(pp. 331 - 332)
Electrons as Waves
1. Louis DeBroglie thought that Bohr’s quantized electron orbits had characteristics similar to those of waves.
2. His question: If waves can have particle like behavior, could the opposite also be true? Can particles like electron behave like waves?
3. If an electron has wavelike motion and is restricted to circular orbits of fixed radius, the electron is allowed only certain possible wavelengths, frequencies and energies.
4. Experiments that followed proved that electrons and other moving particles do indeed have wave characteristics.
The Heisenberg Uncertainty Principle
1. Heisenberg concluded that it is impossible to make any measurement on an object without disturbing the object.
2. If one tries to locate an electron by shining a light on it, the electron will interact with the photon. This changes both the wavelength of the photon and the position and velocity of the electron.
3. The ______states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.
4. Schrodinger later derived an equation that treated the hydrogen’s electron as a wave.
5. This model for the hydrogen atom seemed to apply equally well to atoms of the other elements.
6. The atomic model in which electrons are treated as waves is commonly known as the quantum mechanical model of the atom.
7. While it limits as electron’s energy to certain values, it makes no attempt to describe the electron’s path around the nucleus.
8. Schrodinger’s wave equation is very complex. Each solution is known as a wave function. The wave function is related to the probability of finding the electron within a particular volume of space around the nucleus.
9. The wave function predicts a three-dimensional region around the nucleus called an ______that describes the electron’s probable location.
10. The atomic orbital can be thought of as a fuzzy cloud in which the density of the cloud at a given point is proportional to the probability of finding the electron at that point.
Hydrogen’s Atomic Orbitals (pp. 333 - 337)
1. Because the boundary of an atomic orbital is fuzzy, the orbital does not have an exactly defined size.
2. Chemists draw an orbital’s surface to contain 90% of the electron’s total probability distribution.
3. The Bohr model assigns quantum numbers to electron orbits. Similarly, the quantum mechanical model assigns ______.
4. The principal quantum number indicates the relative sizes and energies of atomic orbital. As n increases, the orbital becomes larger, the electron spends more time farther from the nucleus and the atom’s energy level increases.
5. Therefore n specifies the atom’s major energy levels, called ______.
6. The lowest energy level is ______, the atom is in its ground state.
ELECTRON CONFIGURATIONS (pp. 337 - 347)
1. The arrangement of the electrons in an atom is called the atom’s ______.
2. Low energy systems are more stable than high-energy systems so electrons tend to assume the arrangement that gives the atom the lowest possible energy.
3. Three rules, or principles, define how electrons can be arranged in an atom” orbitals:
4. The ______states that each electron occupies the lowest energy orbital available.
5. The ______states that a maximum of two electrons may occupy a single atomic orbital, but only if the electrons have opposite spins.
6. ______states that single electrons with the same spin must occupy each equal energy orbital before additional electrons with opposite spins can occupy the same orbitals.
LABELING ELECTRONS IN ATOMS
1. In quantum theory, each electron in an atom is assigned a set of four quantum numbers.
2. Three of the numbers act like coordinates to describe the location while the fourth describes the orientation in the orbital.
3. The first number is the ______This number describes the energy level that the electron occupies. It is assigned a positive integer starting with one:
4. The larger the value of n, the further the electron is from the nucleus and the more energy it will have.
5. Quantum numbers are used to describe shapes of atomic orbitals. The second quantum number is designated by the letter ______.
6. While “l” does have numerical values, the shapes are often described using letters. This translates as:
lLetter
8. The lowest energy orbital is the s-orbital. It is closest to the nucleus and shaped like a sphere. There is one s-orbital in each energy level.
9. The next higher orbital is the p-orbital. There are three p-orbitals at each energy level. These are shaped like dumbbells and oriented along the x, y, and z-axes.
10. Orbitals designated _____ and _____ have more complex shapes and still higher energy.
11. A third quantum number, the magnetic quantum number, is designated ______. The value of ml tells one the electron’s position by designating the spatial orientation of the orbital that the electron occupies.
12. Electrons can have the same value of n and l, but not necessarily ml. These electrons are said to be in the same SUBLEVEL of the atom.
13. The 2p-sublevel contains a maximum of three orbital: 2px, 2py and 2pz. These orbital are oriented along the x, y and z-axes.
14. To describe the motion of the electron, there is a fourth quantum number, ______. This is the spin quantum number.
15. This number labels the orientation of the electron. Electrons in an orbital spin in the opposite direction. These directions are +1/2 and –1/2.
16. According to the Pauli exclusion principle, no two electrons can have an identical set of four quantum numbers.
Electron Configurations
1. An electron configuration is an abbreviated form of the orbital diagram
2. This diagram will help you sequence the proper filling of the orbitals.
3. This energy level diagram helps explain why the orbitals fill the way they do.
How to Write an Electron Configuration
1. Locate the element whose electron configuration you wish to write in the periodic table.
2. Fill the orbital in the proper order with electrons.
3. Check that the total number of electrons in the electron configuration equals the atomic number.
Electron Configurations – Noble Gases
1. Electron configurations can be abbreviated to save time and space.
2. By using the noble gas that appears just prior to the element and then adding the remaining electrons, one can write the electron configuration.
***** Write the electron configuration for nickel.
***** Write the abbreviated electron configuration for Silicon.
Orbital Diagrams
1. Orbital diagrams are used to show how electrons are distributed within sublevels and to show the direction of spin.
2. Each orbital is represented by a box and each electron by an arrow.
3. The direction of spin is represented by the direction of the arrow.
***** Construct the orbital diagram for oxygen and aluminum.
.
4. Electrons are arranged in accordance with Hund’s Rule. This rule states that orbital of equal energy are each occupied by one electron before any pairing occurs by adding a second electron. In this way, repulsion is minimized.
5. All electrons in singly occupied orbital must have the same spin.
6. When two electrons occupy an orbital they must have the opposite spin.
***** Construct an orbital diagram for iron.
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