Name: Electrons in Atoms Chemical Periodicity Chapters 13 and 14

Transcription

1 Name: Electrons in Atoms Chemical Periodicity Chapters 13 and 14 1

2 Chapter 13 Electrons in Atoms We need to further develop our understanding of atomic structure to help us understand how atoms bond to form compounds: 13.1 Models of the Atom Briefly describe how the following people viewed the atom: A. John Dalton (Page 361) B. JJ Thomson (Page 361) C. Ernest Rutherford (Page 362) D. Niels Bohr (Page 362) In 1926, the Austrian physicist Erwin Schrodinger took atomic models one step further. He has given us the modern description of the electrons in atoms, called the. Describe this model: (bottom paragraph page 363 through top paragraph page 364) Atomic Orbitals Page 364 a. What is a principal energy level? b. Within each principal energy level, the electrons occupy energy. 2

3 c. The number of sublevels within each principal energy level is the as the principal quantum number. d. What are atomic orbitals? e. What are the four letters used to denote the atomic orbitals? f. What shape do the first 3 atomic orbitals take? (draw them below) S shape P shape D shape (f shape clouds are too complicated to draw ) The S orbital has spatial orientation, therefore can hold electrons. The P orbitals have spatial orientations, therefore can hold electrons. The D orbitals have spatial orientations, therefore can hold electrons. The F orbitals have spatial orientations, therefore can hold electrons Electron Arrangement in Atoms Page 367 There are 3 general rules that guide us when writing electron configurations. They are: 1. Aufbau principle: 2. Pauli Exclusion Principle: 3. Hund s Rule: 3

4 There is a helpful tool that makes writing electron configurations easier: The Electron Pyramid Draw the Electron Configuration Pyramid that helps us diagram atomic structures: Use these rules to write electron configurations and orbital diagrams for the following atoms: a. Lithium b. Fluorine c. Rubidium Example 2: After writing the electron configuration for each of the following, indicate how many unpaired electrons the atom has: a. Phosphorus b. Nickel c. carbon d. argon 4

5 Exceptional Electron Configurations Chromium and Copper have electron arrangements that completely fill the d sublevel, leaving their 4s partially filled. This is a much more stable arrangement. Write the correct, (exception to the rules) configuration for a. Chromium b. Copper 13.3 Physics and the Quantum Mechanical Model (Page 372) This section studies the electron as a property of light. Electrons travel as waves and are made of particles called photons. According to the wave model, light consists of. This form of energy includes o o o o o o o o 5

6 Every element emits light when it is excited by the passage of electric discharge through its gas or vapor. The atoms first absorb energy, then lose the energy as they emit light. They are said to move from their (lowest energy level) to an (higher energy level.) When the electron falls back to its lower energy, it emits a of energy, and can be seen in the visible spectrum. Passing the light emitted by an element through a prism gives the of the element. Because each atom has a unique electron arrangement, each atom emits a unique wavelength during this process. The wavelength falls within the visible spectrum. Section Review: 1. How many electrons are in the highest occupied energy level of these atoms? a. Barium b. Sodium c. aluminum d. oxygen 2. Write electron configurations and orbital diagrams for the elements that are identified only by these atomic numbers. a. 15 b. 12 c. 9 d Give the symbol for the atom that corresponds to each electron configuration. a. 1s 2 2s 2 2p 6 3s 2 3p 6 b. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 7 5s 1 c. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 1 6s 2 6

7 Kernel Structures The kernel is a structure used to shorten an electron configuration. A kernel is an inert gas symbol in brackets that stands in place of all of the filled orbitals contained in the inert gas. Example: Using a kernel structure, write the electron configuration for the following elements: a. Sodium b. Germanium c. Silver d. Tellurium (Honors Only) The Quantum Concept and the Photoelectric Effect Electrons travel as waves around the nucleus of an atom. Let s review the concept of wave mechanics. Define the following terms: amplitude wavelength frequency hertz The frequency and wavelength of all waves, including light and thus electrons, are inversely related. Show this relationship mathematically: What is the value for the speed of light constant? Example Problems: 1. Calculate the wavelength of the yellow light emitted by a sodium lamp if the frequency of the radiation is 5.1 E 14 Hz. 2. What is the wavelength of radiation with a frequency of 1.5 E 13 Hz? Does this radiation have a longer or shorter wavelength than red light? 7

8 3. What frequency is radiation with a wavelength of 5.00 E -6 cm? In what region of the electromagnetic spectrum is this radiation? (Honors Only) Photoelectric Effect By studying black body radiation, German physicist Max Planck described mathematically that the amount of radiant energy (E) absorbed or emitted by a body is proportional to the frequency of the radiation. The constant (h) is known as Planck s constant, which has a value of. In 1905, Albert Einstein proposed that light could be described as quanta of energy that behave as if they were particles he called. In the, metals eject electrons when light shines on them. This revolutionary idea has allowed scientists to develop photoelectric cells to generate electricity, such as used in solar panels. Example Problems: 1. Calculate the energy (in Joules) of a quantum of radiant energy (the energy of a single photon) with a frequency of 5.00 E 15 Hz. 2. What is the energy of a photon of microwave radiation with a frequency of 3.2 E 11 Hz? What is the energy of one mole of photons at this same frequency? 3. The threshold photoelectric effect in tungsten is produced by light of wavelength 260 nanometers (10 E -9). Give the energy of a photon of this light in joules. 8

9 (Honors Only) Quantum Mechanics and Matter Waves In 1924 Louis De Broglie derived an equation that described the wavelength of a moving particle, such as an electron. If the mass of an electron is 9.11 E -28 grams and moving nearly at the speed of light, an electron has a wavelength of about 2 E -10 cm. De Broglie s prediction that matter would exhibit both wave and particle properties is summarized in the following two statements: Examples: 1. What is the wavelength of an electron moving with a speed of 5.97 E 6 m/s? (The mass of an electron is 9.11 E -28 grams). (h = 6.63 E -34 J-s) (1 Joule = 1 kg-m 2 /s 2 ) 2. The electron microscope has been widely used to obtain highly magnified images of biological and other types of materials. When an electron is accelerated through a particular potential field, it attains a speed of 9.38 E 6 m/s. What is the characteristic wavelength of this electron? 9

10 (Honors Only) Orbitals and Quantum Numbers The quantum mechanical model uses a set of 4 quantum numbers to describe a single electron s position within an atom. This is analogous to our homes having a single unique address to represent its location. 1. The Principle Quantum number, n: 2. The azimuthal quantum number, l: Value of l Letter used 3. The magnetic quantum number, ml : 4. The spin magnetic quantum number, m s Examples: 1. For n = 4, what are the possible values of l? 2. For l = 2, what are the possible values of m l 3. How many possible values for l and ml are there when a. n = 3 b. n = 5 4. Give the numerical values of n and l corresponding to each of the following designations: a. 3p b. 2s c. 4f d. 5d 10

11 (ALL) Chapter 14: Chemical Periodicity 14.1 Classification of the Elements (Page 391) A. Classifying Elements by Electron Configuration a. Of the three major subatomic particles, the plays the most significant role in determining the physical and chemical properties of an element. b. The arrangement of elements in the depends on these properties. c. Elements can be classified into four categories according to their electron configurations: i. The Noble Gases: Write the electron configurations for the first four noble gases here: Helium: Neon Argon Krypton ii. The representative elements: Write the electron configurations for: Lithium Sodium Potassium Carbon Silicon Germanium 11

12 iii. The transition metals: Write the electron configurations for the following: Manganese: Zinc Zirconium iv. The inner transition metals: B. There is a pattern to the periodic table. The Periodic Table can be divided into sections, or blocks, that correspond to the sub levels that are filled with electrons: Examine the diagram below: 12

13 Practice: Use the Periodic Table to write the electron configurations of these elements: a. Nitrogen b. Nickel c. Vanadium Practice: What are the symbols for all the elements that have the following outer configurations? a. s 2 b. s 2 p 5 c. s 2 d Periodic Trends (Page 398) Trends in Atomic Size Define atomic radius: Describe the group trends as related to atomic radius: Describe the periodic trends as related to atomic radius: Trends in Ionization Energy Define an ion: Define ionization energy: Define the term first ionization energy : Define the term second ionization energy : Describe the group trends as related to ionization energy: Describe the periodic trends as related to ionization energy: 13

14 Trends in Ionic Size The atoms of elements have low ionization energies. They form ions easily. By contrast, the atoms of elements readily form ions. Describe the trends as related to ionic size: Positive ions are always than the neutral atoms from which they form. o. Negative ions are always than the neutral atoms from which they form. o This is because the is less for an increased number of electrons. Trends in Electronegativity Define electronegativity: Electronegativity generally as you move down a group. As you go across a period from left to right, the electronegativity of the representative elements. The electronegativity of cesium, the least electronegative element, is. The electronegativity of fluorine, the most electronegative element, is. Electronegativity values help predict the type of bonding that can exist between atoms in compounds, either or bonds. Summary of Periodic Trends Upon examining Figure 14.6 on page 406, copy the figure into your notes below. Be sure to include all periodic trends. 14

15 Section Review: 1. For which of these properties does lithium have a larger value than potassium? a. first ionization energy b. atomic radius c. electronegativity d. ionic radius 2. Arrange these elements in order of decreasing atomic size: sulfur, chlorine, aluminum, and sodium. Does this demonstrate a group trend or a periodic trend? 3. How does the ionic radius of a typical anion compare with the radius for the corresponding neutral atom? 4. Which element in each pair has the larger ionization energy? a. sodium, potassium b. magnesium, phosphorus c. lithium, boron d. magnesium, strontium e. cesium, aluminum 5. Indicate which element in each pair has the greater atomic radius: a. sodium, lithium b. strontium, magnesium c. carbon, germanium d. selenium, oxygen 6. Arrange the following elements in order of increasing ionization energy: a. Be, Mg, Sr b. Bi, Cs, Ba c. Na, Al, S 15

16 7. Why is there a large increase between the first and second ionization energies of the alkali metals? 8. Which particle has the larger radius in each atom/ion pair? a. Na, Na + b. S, S -2 c. I, I -1 d. Al, Al How does the ionic radius of a typical metallic atom compare with its atomic radius? 10. Which element is more electronegative? a. Cl, F b. C, N c. Mg, Ne d. As, Ca 16

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