Chapter 6 electronic structure of atoms

Having introduced the basics of atomic structure and quantum mechanics, we can use our understanding of quantum numbers to determine how atomic orbitals relate to one another. This allows us to determine which orbitals are occupied by electrons in each atom. The specific arrangement of electrons in orbitals of an atom determines many of the chemical properties of that atom. The energy of atomic orbitals increases as the principal quantum number, nincreases.

Figure 1 depicts how these two trends in increasing energy relate. The 1 s orbital at the bottom of the diagram is the orbital with electrons of lowest energy. The energy increases as we move up to the 2 s and then 2 p3 sand 3 p orbitals, showing that the increasing n value has more influence on energy than the increasing l value for small atoms. However, this pattern does not hold for larger atoms.

The 3 d orbital is higher in energy than the 4 s orbital. Such overlaps continue to occur frequently as we move up the chart. Figure 1. Generalized energy-level diagram for atomic orbitals in an atom with two or more electrons not to scale. Electrons in successive atoms on the periodic table tend to fill low-energy orbitals first. Thus, many students find it confusing that, for example, the 5 p orbitals fill immediately after the 4 dand immediately before the 6 s.

The filling order is based on observed experimental results, and has been confirmed by theoretical calculations. As the principal quantum number, nincreases, the size of the orbital increases and the electrons spend more time farther from the nucleus.

Thus, the attraction to the nucleus is weaker and the energy associated with the orbital is higher less stabilized. But this is not the only effect we have to take into account. This phenomenon is called shielding and will be discussed in more detail in the next section. Electrons in orbitals that experience more shielding are less stabilized and thus higher in energy.

For small orbitals 1 s through 3 pthe increase in energy due to n is more significant than the increase due to l ; however, for larger orbitals the two trends are comparable and cannot be simply predicted.

We will discuss methods for remembering the observed order.

chapter 6 electronic structure of atoms

The arrangement of electrons in the orbitals of an atom is called the electron configuration of the atom. We describe an electron configuration with a symbol that contains three pieces of information Figure 2 :. Figure 2. The diagram of an electron configuration specifies the subshell n and l value, with letter symbol and superscript number of electrons. Figure 3. The arrow leads through each subshell in the appropriate filling order for electron configurations.

This chart is straightforward to construct. Simply make a column for all the s orbitals with each n shell on a separate row. Repeat for p, d, and f. Be sure to only include orbitals allowed by the quantum numbers no 1p or 2d, and so forth. Finally, draw diagonal lines from top to bottom as shown. Beginning with hydrogen, and continuing across the periods of the periodic table, we add one proton at a time to the nucleus and one electron to the proper subshell until we have described the electron configurations of all the elements.

Each added electron occupies the subshell of lowest energy available in the order shown in Figure 1subject to the limitations imposed by the allowed quantum numbers according to the Pauli exclusion principle.

Electrons enter higher-energy subshells only after lower-energy subshells have been filled to capacity.Understanding the electronic structure of atoms requires an understanding of the properties of waves and electromagnetic radiation.

A basic knowledge of the electronic structure of atoms requires an understanding of the properties of waves and electromagnetic radiation.

Chapter 6 – The Electronic Structure of Atoms: Part 1 of 10

A wave is a periodic oscillation by which energy is transmitted through space. All waves are periodic, repeating regularly in both space and time. Waves are characterized by several interrelated properties. We say that the allowed energies are quantized that is, their values are restricted to certain quantities. Blackbody radiation is the radiation emitted by hot objects and could not be explained with classical physics.

Max Planck postulated that energy was quantized and may be emitted or absorbed only in integral multiples of a small unit of energy, known as a quantum. Albert Einstein used the quantization of energy to explain the photoelectric effect. There is an intimate connection between the atomic structure of an atom and its spectral characteristics. Most light is polychromatic and contains light of many wavelengths.

Light that has only a single wavelength is monochromatic and is produced by devices called lasers, which use transitions between two atomic energy levels to produce light in a very narrow range of wavelengths.

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Atoms can also absorb light of certain energies, resulting in a transition from the ground state or a lower-energy e. Electrons in a permitted orbit have a specific energy and are said to be in an "allowed" energy state. An electron in an allowed energy state will not radiate energy and therefore will not spiral into the nucleus. All energies given by this equation will be negative.

6.E: Electronic Structure of Atoms (Exercises)

The lower more negative the energy is, the more stable the atom is. Thus, the state in which the electron is removed from the nucleus is the reference, or zero-energy, state of the hydrogen atom. It is important to remember that this zero-energy state is higher in energy than the states with negative energies. Electrons can change from one energy state to another by absorbing or emitting radiant energy.

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Radiant energy must be absorbed for an electron to move to a higher energy state, but is emitted when the electron moves to a lower energy state.

An electron possesses both particle and wave properties. There is a relationship between the motions of electrons in atoms and molecules and their energies that is described by quantum mechanics. Because of wave—particle duality, scientists must deal with the probability of an electron being at a particular point in space. To do so required the development of quantum mechanics, which uses wavefunctions to describe the mathematical relationship between the motion of electrons in atoms and molecules and their energies.

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Defined by the value of 3 quantum numbers; n, l, and m l. The s Orbitals: 1s orbital: most stable, spherically symmetric, figure indicates that the probability decreases as we move away from the nucleus.E: Matter and Measurement Exercises 1. E: Atoms, Molecules, and Ions Exercises 2. E: Stoichiometry Exercises 3. E: Reactions in Aqueous Solution Exercises 4.

E: Thermochemistry Exercises 5. E: Electronic Structure of Atoms Exercises 6. E: Periodic Properties of the Elements Exercises 7. E: Exercises 9. E: Exercises E: Liquids and Intermolecular Forces Exercises E: Properties of Solutions Exercises E: Acid—Base Equilibria Exercises E: Chemistry of the Nonmetals Exercises E: Organic and Biological Chemistry Exercises E: Matter and Measurement Exercises 2. E: Atoms, Molecules, and Ions Exercises 3.

E: Stoichiometry Exercises 4. E: Aqueous Reactions Exercises 5. E: Thermochemistry Exercises 6. E: Electronic Structure Exercises 7. E: Periodic Trends Exercises 8. E: Chemical Bonding Basics Exercises 9.Explain the importance of the quantum theory to our society and you, in particular! Know how to interpret the EM spectrum. No need to memorize, but be able to identify which wave has a longer wavelength or shorter frequency, etc.

Know the "parts" of a wave and how waves are measured wavelength, amplitude, period. Be able to calculate frequency and wavelength. Sample Exercise 6. Explain why energy is quantized and use figure 6. Explain the photoelectric effect and what Einstein learned from Planck.

Explain how Bohr's model helped not fully explained the behavior of electrons. Include limitations of his model. Explain the difference between excited and ground state for electrons. State the Heisenburg uncertainty principle and its implication for chemists.

Explain the results of Schrodinger's experiments not the calculus involved. State the designations of the three quantum numbers n, l and m i.

6: Electronic Structure of Atoms

You will not need to draw the orbitals, just know, in general terms, what they look like. Be able to write electron configurations using long and noble gas core notation. Be able to identify an element by its electron configuration. Purpose: Determine how well you understand the material from chapter 6 sections. Section 6. Purpose: Determine how well you understand the material from chapter 6.

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Chapter 6: Electronic Structure of Atoms.E: Matter and Measurement Exercises 1. E: Atoms, Molecules, and Ions Exercises 2. E: Stoichiometry Exercises 3. E: Reactions in Aqueous Solution Exercises 4. E: Thermochemistry Exercises 5.

E: Electronic Structure of Atoms Exercises 6. E: Periodic Properties of the Elements Exercises 7. E: Exercises 9. E: Exercises E: Liquids and Intermolecular Forces Exercises E: Properties of Solutions Exercises E: Acid—Base Equilibria Exercises E: Chemistry of the Nonmetals Exercises E: Organic and Biological Chemistry Exercises E: Matter and Measurement Exercises 2.

E: Atoms, Molecules, and Ions Exercises 3. E: Stoichiometry Exercises 4. E: Aqueous Reactions Exercises 5. E: Thermochemistry Exercises 6. E: Electronic Structure Exercises 7. E: Periodic Trends Exercises 8. E: Chemical Bonding Basics Exercises 9. E: Bonding Theories Exercises E: Gases Exercises Solids and Modern Materials Exercises E: Kinetics Exercises E: Chemical Equilibrium Exercises After you enable Flash, refresh this page and the presentation should play.

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6.1 Electronic Structure of Atoms

To view this presentation, you'll need to allow Flash. Click to allow Flash After you enable Flash, refresh this page and the presentation should play. View by Category Toggle navigation. Products Sold on our sister site CrystalGraphics. Title: Chapter 6 Electronic Structure of Atoms. Description: Chemistry, The Central Science Tags: atom atoms chapter chemistry electronic structure.

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chapter 6 electronic structure of atoms

Brown H. Eugene LeMay, Jr. Bursten John D. Bookstaver St. Charles Community College St. Peters, MO? The distance between corresponding points on adjacent waves is the wavelength? For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency. Therefore, c??Complete: Journals that are no longer published or that have been combined with another title.

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chapter 6 electronic structure of atoms

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6: Electronic Structure of Atoms

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