Zumdahl • Zumdahl • DeCoste

Zumdahl • Zumdahl • DeCoste
World of CHEMISTRY

Chapter 11 Modern Atomic Theory

Review Trends in the periodic table:
certain elements grouped together because they behave similarly there is great similarities within groups, but the differences in behavior between groups is what we will be studying. Copyright © Houghton Mifflin Company

Objective: To describe Rutherford’s model of the atom
11.1 Rutherford’s Atom Objective: To describe Rutherford’s model of the atom Copyright © Houghton Mifflin Company

Review of the composition of an atom
Nucleus Proton (+) Orbital's Electrons (-) Copyright © Houghton Mifflin Company

Figure 11.1: The Rutherford atom.
Rutherford’s model shows a very small nucleus with electrons arranged in such a way that he believed they revolved around the nucleus (i.e. like the solar system) What Rutherford could not explain was why the negative electrons aren’t attracted into the positive protons, causing the atom to collapse. Copyright © Houghton Mifflin Company

Objective: To explore the nature of electromagnetic radiation.
11.2 Energy and Light Objective: To explore the nature of electromagnetic radiation. Copyright © Houghton Mifflin Company

Electromagnetic radiation
The energy transmitted from one place to another by light. Types of electromagnetic radiation X-Rays The white light from a light bulb Microwaves used to cook Radio waves that transmit sound Copyright © Houghton Mifflin Company

3 Components of Waves Wavelengths Frequency speed
Copyright © Houghton Mifflin Company

Figure 11.2: A seagull floating on the ocean moves up and down as waves pass.
Notice that the gull just moves up and down with the motion of the waves, it does not move forward. Copyright © Houghton Mifflin Company

A wavelength (λ) is the distance between two consecutive wave peaks.

Figure 11.3: The wavelength of a wave.

Frequency (nu ν ) Frequency indicates how many wave peaks pass a certain point per given time period. i.e how many times does the sea gull go up and down per minute. Copyright © Houghton Mifflin Company

Light (electromagnetic radiation) Travels in Waves
Different forms of electromagnetic radiation travel at different wavelengths Refer pg 325, fig. 11.4 Sunlight: ultraviolet radiation and visible light Glowing coals transmits heat via infrared radiation. Microwave radiation: water molecules absorb the microwave and this energy is transferred to other molecules by collision (KE). Copyright © Houghton Mifflin Company

Light = waves that carry energy
Photons – tiny energy packets that travel through space in a beam of light. Copyright © Houghton Mifflin Company

Figure 11.5: Electromagnetic radiation.
A beam of light can be pictured in two ways: as a wave and as a stream of individual packets of energy called photons Copyright © Houghton Mifflin Company

Figure 11.6: Photons of red and blue light.
The photons of red light (long wavelengths) carries less energy than a photon of blue light (shorter wavelengths). Copyright © Houghton Mifflin Company

11.3 Emission of Energy by Atoms
Objective: To see how atoms emit light Video – flame test Copyright © Houghton Mifflin Company

The heat from the flame causes atoms to absorb energy.
The Colors of Flames Result from Atoms Releasing Energy by Emitting Visible Light The heat from the flame causes atoms to absorb energy. The electrons become excited. Energy is released in the form of light The energy is carried by photons Energy change = energy carried by the photon High energy photons = short-wavelengths Low-energy photons = long-wavelengths Copyright © Houghton Mifflin Company

Figure 11.6: Photons of red and blue light.
The photons of red light (long wavelengths) carries less energy than a photon of blue light (shorter wavelengths). Copyright © Houghton Mifflin Company

Figure 11.8: An excited lithium atom emitting a photon of red light to drop to a lower energy state.

What is the difference between the frequency of a wave and its speed?
Focus Questions 11.1 – 11.3 What was wrong with Rutherford’s model of the atom? Why did it need to be modified What is the difference between the frequency of a wave and its speed? What is the relationship between the wavelength of light and the energy of a photon? Copyright © Houghton Mifflin Company

4. How can red light improve germination and production in tomato crops?

11.4 The Energy Levels of Hydrogen
Objective: To understand how the emission spectrum of hydrogen demonstrates the quantized nature of energy. Copyright © Houghton Mifflin Company

11.4 The Energy Levels of Hydrogen
Review: different wavelengths of light carry different amounts of energy per photon An atom with excess energy is said to be in an excited state. An excited atoms can release some or all of its excess energy by emitting a photon. Copyright © Houghton Mifflin Company

Figure 11.9: A sample of H atoms receives energy from an external source.

Figure 11.9: The excited atoms release energy by emitting photons.

Figure 11.10: An excited H atom returns to a lower energy level.
Notice: the energy contained in the photon corresponds to the change in energy that the atom experiences in going from the excited state to the lower state. Copyright © Houghton Mifflin Company

Figure 11.11: Colors and wavelengths of photons in the visible region.
In a sample of Hydrogen atoms a lot of energy is put into the system, and the following visible light is emitted. That is, only certain types of photons are produced. What does this tell you? Copyright © Houghton Mifflin Company

Energy levels of Hydrogen are Quantized
Because only certain photons are emitted, we know that only certain energy changes are occurring. Copyright © Houghton Mifflin Company

Figure 11.12: The color of the photon emitted depends on the energy change that produces it.

Figure 11.13: Each photon emitted corresponds to a particular energy change.
Therefore Hydrogen atoms must have certain discrete energy levels. Copyright © Houghton Mifflin Company

The energy of all atoms are Quantized
Each atom has its own set of discrete energy levels. Copyright © Houghton Mifflin Company

Figure 11.15: The difference between continuous and quantized energy levels.

11.5 The Bohr Model of the Atom
Objective: To learn about Bohr’s model of the hydrogen atom Copyright © Houghton Mifflin Company

Figure 11.17: The Bohr model of the hydrogen atom.
Bohr theorized that the atom consisted of a small nucleus with electrons orbiting around it like the sun and the planets. The electrons could jump to different orbits by absorbing or emitting a photon. Copyright © Houghton Mifflin Company

Bohr’s model is only consistent with Hydrogen
Current theory of atomic structure do not support the Bohr model. Electrons do not move around the nucleus in circular orbits like planets orbiting the sun. Copyright © Houghton Mifflin Company

11.6 The Wave Mechanical Model of the Atom
Objective: To understand how electron’s position is represented in the wave mechanical model Copyright © Houghton Mifflin Company

The Wave Mechanical Model of the Atom
Louis Victor de Broglie and Erwin Schrodinger: Since light has both wave and particle characteristics……the electron might also exhibit both of these characteristics. Copyright © Houghton Mifflin Company

Figure 11.18: A representation of the photo of the firefly experiment.

Figure 11.19: The orbital that describes the hydrogen electron in its lowest possible energy state.

1. Figure 11. 11 contains four colored lines
1. Figure contains four colored lines. You already know that hydrogen only has one electron. How can we get four lines from one electron? The lines produced in an emissions spectrum represent thousands of atoms of the element in question. Each line represents a single electron transition. Copyright © Houghton Mifflin Company

2. What is wrong with the Bohr model of the atom. 3
2. What is wrong with the Bohr model of the atom? 3. How does the wave mechanical model differ from Bohr’s model? Copyright © Houghton Mifflin Company

Figure 11.20: The hydrogen 1s orbital.
Because the edges of an orbital are fuzzy, an orbital does not have an exactly defined size. The sphere that contains 90% of the total electron probability, that is the electrons spend 90% of the time inside this surface and 10% of the time somewhere outside this surface. Copyright © Houghton Mifflin Company

If hydrogen can absorb energy to transfer electron to a higher energy state (excited),
the wave mechanical model states that each energy state corresponds to a different orbital with different shapes. Copyright © Houghton Mifflin Company

Focus Questions 11.4 – 11.6 Figure contains four colored lines. You already know that hydrogen has only one electron. How can we get four lines from one electron? What is wrong with the Bohr model of the atom? How does the wave mechanical model of the atom differ from Bohr’s model? Copyright © Houghton Mifflin Company

Figure 11.21: The first four principle energy levels in the hydrogen atom.

Figure 11.22: How principal levels can be divided into sublevels.

These sublevels contain spaces (orbitals) for the electrons.
Like the inverted triangle, the principle energy levels in the hydrogen atom contain subatomic particles. These sublevels contain spaces (orbitals) for the electrons. Copyright © Houghton Mifflin Company

Figure 11.20: The hydrogen 1s orbital.

Figure 11.23: Principal level 2 shown divided into the 2s and 2p sublevels.

Figure 11.25: The three 2p orbital‘s are not spherical, but are lobes.

Figure 11.26: Diagram of principal energy levels 1 and 2.

Summary of Orbital Labels
1. The number tells the principal energy level 2. The letter tells the shape S = sphere P = two lobes x, y, z Copyright © Houghton Mifflin Company

As the level number increases , the average distance of the electron in that orbital from the nucleus increases. When hydrogen electron is in its 1s (ground state) it spends most of its time closer to the nucleus than when it occupies 2s orbital (excited state). Copyright © Houghton Mifflin Company

Orbitals can be thought of as potential space for electrons.
Hydrogen orbital’s Orbitals can be thought of as potential space for electrons. Hydrogen with its one electron can only occupy a single orbital at a time But the other orbitals are available should the electron be excited. Copyright © Houghton Mifflin Company

Figure 11.27: Relative sizes of the spherical 1s, 2s, and 3s orbital's of hydrogen.

Level 1 = 1s Level 2 = 2s; 2px 2py 2pz Level 3 = 3s; 3p; 3d
3dzy 3dxz 3dxy 3dx2- y2 & 3dz2 Level 4 = 4s; 4p; 4d; 4f Where the f sublevel has four compartments Copyright © Houghton Mifflin Company

Figure 11.28: The shapes and labels of the five 3d orbital's.

11.8 The Wave Mechanical Model: Further Development
Objectives: To review the energy levels and orbitals of the wave mechanical model of the atom To learn about electron spin Copyright © Houghton Mifflin Company

Each electron appears to spin on its axis.
Electron Spin Each electron appears to spin on its axis. Two electrons in the same orbital must have opposite spins. Pauli Exclusion Principal An atomic orbital can hold a maximum of two electrons, and those two electron must have opposite spin. Copyright © Houghton Mifflin Company

Principal Components of the Wave Mechanical Model of the Atom
1. Atoms have a series of energy levels called principal energy levels, which are designated by whole numbers symbolized by n 2. The energy of the levels increase as the value of n increases. 3. Each principal energy level contains one or more types of orbitals called sublevels. Copyright © Houghton Mifflin Company

Principal Components of the Wave Mechanical Model of the Atom cont.,
4. The number of sublevels corresponds to the principal energy level 1s (1) 2s(1) 2p(3) 3s (1) 3p (3) 3d(5) 4s(1) 4p(3) 4d(5) 4f(7) (note – numbers in parentheses indicates orbitals, not energy levels) Copyright © Houghton Mifflin Company

5. The value of n is always used to label the orbitals followed by the letter that tells the shape of the orbital. 6. An orbital can be empty or it can contain one or two electrons, but never more than two. 7. The shape of an orbital does not indicate the details of electron movement. It indicates the probability distribution of electron movement. Copyright © Houghton Mifflin Company

Refer to pg. 338 for practice

Focus Question 11.7 – 11.8 What is the difference between an orbit and an orbital in atomic theory? Draw Figure and fill in the different types of sublevels for each principal energy level. Tell how many orbitals are found in each type of sublevels: s, p, d, f. What is the Pauli exclusion principal and how does it help us determine where an electron is found within the atom? Copyright © Houghton Mifflin Company

11.9 Electron Arrangement in the First Eighteen Atoms on the Periodic Table
Objectives: To understand how the principal energy levels fill with electrons in atoms beyond hydrogen To learn about valence electrons and core electrons Copyright © Houghton Mifflin Company

11.10 Electron Configuration and the Periodic Table
Objective: To learn about the electron configuration of atoms with Z greater than 18 Copyright © Houghton Mifflin Company

Figure 11.30: Partial electron configurations for the elements potassium through krypton.

Figure 11.31: Orbital's being filled for elements in various parts of the periodic table.

11.11 Atomic Properties and the Periodic Table
Objective: To understand the general trends in atomic properties in the periodic table. Copyright © Houghton Mifflin Company

Figure 11.35: Classification of elements as metals, nonmetals, and matalloids.

Figure 11.36: Relative atomic sizes for selected atoms.

Write the orbital diagram for the elements listed below.
Focus Questions Write the orbital diagram for the elements listed below. Copyright © Houghton Mifflin Company

2. what is the difference between a valence electron and the core electron? Select an element from row 3 and label both using its electron configuration. 3. Elements in vertical columns (families) show similar chemical behavior. How are their electron configurations similar? Copyright © Houghton Mifflin Company

4. Using their positions on the periodic table, write the valence-electron configuration for the following elements. 5. What chemical properties distinguishes a metal from non-metal. 6. How can the property in question 5 be explained using ionization energy trends? Copyright © Houghton Mifflin Company

Figure 11.14: Continuous and discrete energy levels.

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