Atomic Theory Unit 3 Development of the Atomic Theory 1. Where is the mass of the atom concentrated? 2. What is located in the nucleus? 3. What is the negative particle that orbits the nucleus? 4. What is the sum of the protons and neutrons called? 5. With a neutral atom, what two items are equal in number? 6. What is the term for any charged particle? 7. What is the term for the positively charged ion? 8. What is the term for the negatively charged ion? 9. What type of ion is formed when electrons are gained? 10. What type of ion is formed when electrons are lost? 11. How does atomic theory today differ from Dalton s theory? 12. Which model of the atom is based on the solution to the Schrodinger equation? How is this different from the planetary model? Match each term from the experiments of J.J. Thomson with the correct description. 13. anode a. an electrode with a negative charge 14. cathode b. a glowing beam between electrodes 15. cathode ray c. an electrode with a positive charge 16. electron d. a negatively charged particle 17. The diagram shows electrons moving from left to right in a cathode-ray tube. Draw an arrow showing how the path of the electrons will be affected by the placement of the negatively and positively charged electrodes. 18. Indicate the letter of each sentence that is true about atoms, matter and electric charge. a. All atoms have an electric charge. b. Electric charges are carried by particles of matter. H. Cannon, C. Clapper and T. Guillot Klein High School
c. Atoms always lose or gain charges in whole-number multiples of a single basic unit. d. When a given number of positively charged particles combine with an equal number of negatively charged particles, an electrically neutral particle is formed. 19. Indica te the letter next to the number of units of positive charge that remain if a hydrogen atom loses an electron. a. 0 b. 1 c. 2 d. 3 20. The positively charged subatomic particle that remains when a hydrogen atom loses an electron is called a(n). 21. What charge does a neutron carry? 22. Indicate the letter of each sentence that is true about the nuclear theory of atoms suggested by Rutherford s experimental results. a. An atom is mostly empty space. b. All the positive charge of an atom is concentrated in a small central region called the nucleus. c. The nucleus is composed of protons and neutrons. d. The nucleus is large compared with the atom as a whole. e. Nearly all the mass of an atom is in its nucleus. 23. According to Bohr, electrons cannot reside at in the figure below a. point A c. point c b. point b d. point d 24. According to the quantum mechanical model, point D in the above figure represents (a) the fixed position of an electron (b) the farthest position from the nucleus that an electron can be found (c)a position where an electron probably exists (d) a position where an electron cannot exist 25. What does the atomic number of an atom represent? 3-2
Electrons in the atom Select the best possible response. 1.The type of charge on the nucleus is a. negative b. neutral c. positive 2. The number of protons in a neutral atom having 18 electrons is. 3. How many neutrons are in the isotope P-29? A neutral atom of bromine has a mass number of 80. 4. It has protons. 5. It has electrons. 6. It has neutrons. 7. Its nuclear composition is. One isotope of hydrogen has a mass of 3. 8. How many protons does it have? 9. How many neutrons does it have? An atom with a +1 charge has atomic number 19 and mass number 39. 10. It has protons. 11. It has electrons. The anion of oxygen has a -2 charge. 12. How many electrons does it have? 13. The number of protons in a neutral atom having 18 electrons is. An atom has a charge of +1 and has 10 neutrons with a mass number of 21. 14. What is its atomic number? 15. How many electrons does it have? An atom has 10 protons, 8 neutrons, and 12 electrons. 3-3
16. What is its charge? 17. What is its mass number? 18. In forming this ion, the neutral atom gained what? An atom has 15 protons, 16 neutrons, and 16 electrons. 19. What is its mass number? 20. What is the overall charge on this ion? A neutral atom has a mass number of 24 and 11 protons. 21. How many protons does it have? 22. How many neutrons does it have? An atom with a -2 charge has atomic number 16 and mass number 32. 23. It has protons. 24. It has electrons. 25. It has neutrons. Calculate the following: 26. There are 4 isotopes of sulfur, S-32, 95.002%; S-33, 0.76%; S-34, 4.22%; and S-36, 0.014%. What is the average amu of sulfur? 27. There are three common isotopes of chromium, Cr-50, 4.345% abundance; Cr-52, 83.789%; Cr-53, 9.50% abundance; what is the average amu of chromium? 28. Calculate the wavelength of the yellow light emitted by a sodium lamp if the frequency of the radiation is 5.10 x 10 14 Hz (5.10 x 10 14 s -1 ). 29. What frequency is radiation with a wavelength of 5.00 x 10-6 cm? In what region of the electromagnetic spectrum is the radiation? 30. What is the energy of a photon of green light with a frequency of 5.80 x 10 14 s -1? 31. Suppose your favorite AM radio station broadcasts at a frequency of 1150 khz. What is the wavelength in cm of the radiation from the station? 3-4
Atomic Structure Practice Element/Isotope Atomic Mass Protons Neutrons Electrons number number 1) Zinc 2) Aluminum 3) Calcium 4) Sulfur 5) Bromine 6) Gold 7) Silver 8) Platinum 9) Uranium 236 10) Plutonium 246 11) Potassium 39 12) Mercury 201 13) Titanium - 48 14) Titanium - 46 Atomic Structure Exercise I. Use the periodic table to compute the number of electrons. Neutrons and protons in the following: A. Cr B. Cl C. Mg D. Ir E. Si F. Ne II. Moseley used x rays to determine the atomic numbers of the elements identify each of the following elements by name. A. 1 Proton B. 4 Protons C. Protons D. 12 Protons E. 20 Protons F. 30 Protons III. How many protons are in the nucleus of each of the following elements A. Uranium B. Selenium C. Helium D. Magnesium IV. Give the number of neutrons in each of the following isotopes A. Titanium-46 B. Nitrogen-15 C. 34 S 16 D. 85 Cu 29 3-5
The Spectra of Elements Everybody knows that a few drops of soup or milk spilled onto a gas burner will change the blue gas flame into a mixture of colors, predominantly yellow. These colors can be used to identify the elements present in the substance dropped into the flame. We will observe the colors produced by several known substances. The color observed in the flame is the result of atoms of the element absorbing energy from the flame then reemitting it. The energy absorbed has an energy and wavelength in the visible region of the electromagnetic spectrum. What your senses detect as light is actually radiation that is part of a continuum that makes up the electromagnetic spectrum. The electromagnetic spectrum is a continuum of energies ranging from the high-energy gamma and x-rays to low-energy radio and microwaves. Radiation includes any energy emitted in all directions from a single source, not just nuclear decay. Light is one form of energy that can be radiated; - the sun constantly radiates energy into space as its matter is converted into energy. We see this light in a variety of colors, depending on the wavelength of the energy when it reaches us. The waves of light with the longest wavelength (red) are refracted (bent) to a lesser degree than the shorter waves (indigo and violet) as they pass through the atmosphere. The visible range of the spectrum is from about 400 nm to 700 nm. The wavelength () is the distance between successive peaks of a wave. The number of waves that pass a given point in space per second is the frequency (f). The unit is inverse time, sec -1 and is also called hertz (Hz). All light moves through a vacuum at the same rate of speed, about 3 X 10 8 m/s (the speed of light c). Frequency and wavelength are inversely related to each other, as the product of the two equals the speed of light: c = f. Since the electrons of an atom can only absorb certain amounts of energy, the wavelength of the energy emitted after excitement of the atoms is characteristic of the element. When viewed through a spectroscope, the spectrum of energy emitted by a sample of a particular element can be observed. 1. Use the spectroscope to view each of the gas tubes set up in the back of the lab. Each colored band corresponds to a different amount of energy (a different wavelength) of light emitted by an electron in the element. Each element has a unique spectrum based on the amount of energy emitted by its electrons 2. Record the details of the colored lines formed on the scale of your spectroscope and the wavelength of each line. Using the equations we have discussed in class, solve for the frequency and energy of each of the gases observed. If a substance has multiple lined, calculate frequency and energy for each line. Express each answer to three significant digits. Express the energy in ergs. In the next part of our investigation, we will use a procedure known as a flame test to become familiar with the colors produced by several common elements and use these observations to better understand the affect of energy on atoms and ions. 3-6
Procedure: a. Light a Bunsen burner. Make sure that the gas burns with a clear blue flame and has an easily distinguishable inner blue cone. b. Obtain samples of ions as provided by your teacher. These will be metal nitrates, dissolved in a solvent to allow easy distribution of the ion. c. Using the spray bottles, introduce a small amount of each ion (one squirt) into the flame and observe the result. d. After you have identified the color produced by each ion, mix two of the substances and see if you can detect the presence of each of the substances or if the colors mingle so that one cannot be distinguished from the other. e. Obtain an unknown sample and observe the colors produced. Try to identify the element present in the unknown sample. Color Key: Na = Yellow green Sr = Deep red K = Violet Li = Crimson Ca = Yellow-red Ba = Green-yellow Cu = Blue- 1. What is the order of the colors in the spectrum from lowest to highest wavelength? 2. What is the order of metallic ions from lowest to highest wavelength? 3. What correlations if any do you see between the electron configuration of the metallic elements and the energy given off by exciting their electrons Further Analysis In 1913, Neils Bohr hypothesized that the electron in the hydrogen atom is allowed to have only certain amounts of energy. These energy levels would be orbits in which the electron in hydrogen would circle the nucleus of hydrogen. He believed that the electron would move from one energy level to another and would give off light when it jumped from a higher energy level to a lower energy level. The amount of energy would be different for jumps between different levels. 1) The diagram below shows a sketch of some of the possible orbits of the hydrogen electron and their corresponding energy values. If we think of the energy values as being on a number line, which orbit has the greatest value? 3-7
2) When an electron jumps from E 2 to E 1 the amount of energy in the light would be E 2 E 1 ; or (-3.4 ev) (-13.6 ev) which would equal 10.2 ev. What would be the amount of energy in the light when an electron jumps from: a) E 3 to E 4 b) E 4 to E 2 c) E 3 to E 2 d) E 5 to E 2 3) Using Bohr s model, we would assume that the electron would only move between certain orbitals or energy levels. Every possible jump corresponds to light of a different energy. a. How many different energies of light can be emitted from hydrogen when the electron jumps down to E 2 from E 3, E 4, E 5, and E 6? b. How many bands of light did you observe when you viewed the hydrogen tube through the spectrometer? c. How do you think these two observations are related? 4) The colors corresponding to jumps to the E 1 level have higher energies or lower, than those to the E 2. a. As the electron jumped from the E 6 all the way to the E 1 level, how many different energies would be emitted? b. Did you see all of these bands? Explain why or why not. c. Bohr did not find bands corresponding to the jumps to the E 3 level. Why do you think that was? Conclusions: In this activity we saw the evidence that energy is emitted in particular patterns depending on the distance that an electron travels between one orbital and another. The number of electrons in an atom and the number of energy levels those electrons occupy determine the number of bands of light and the color of the light seen. What if any conclusions can you draw about the relative numbers of electrons in the atoms of our test atoms? 3-8
Arrangement of Electrons in Atoms 1. What is the difference between the earlier models of the atom and the modern quantum mechanical model? 2. What is a quanta and who developed the concept of quantum energy? 3. How many quantum numbers are used to describe the energy state of an electron in atom? a) 1 (b) 2 (c) 3 d) 4 4. What is the Heisenberg uncertainty principle? 5. The energy level of an electron is the region around the nucleus where 6. In general, the higher the electron is on the energy ladder, the it is to/from the nucleus? 7. A quantum of energy is the amount of energy required to a. move and electron from its present energy level to the next lower one b. maintain an electron in its present energy level c. move an electron from its present energy level to the next higher one 8. True or False: the electrons in an atom can exist between energy levels. 9. True or False: The quantum mechanical model of the atom estimates the probability of finding an electron in a certain position. 10. Which name describes the major energy levels of electrons? a) atomic orbitals b) quantum mechanical numbers c) quanta d) principal quantum number 11. What formula represents the maximum number of electrons that can occupy a principal energy level (n = principal quantum number)? a) 2n 2 b)n 2 c) 2n d) n 12. A spherical electron cloud surrounding an atomic nucleus would best represent _. a. an s orbital b. a p orbital c. a combination of two different p orbitals d. a combination of an s and a p orbitals 13. An energy level of n = 4 can hold electrons. (a) 32 (b) 24 (c) 8 (d) 6 14. An energy level of n = 2 can hold electrons. (a) 32 (b) 24 (c) 8 (d) 6 15. An electron for which n = 4 has more than an electron for which n = 2. 3-9
(a) spin (b) stability (c) energy (d) wave nature Behavior of electrons in the atom 1. How did de Broglie conclude that electrons have a wave nature? 2. Identify each of the four quantum numbers and the properties to which they refer. 3. How did the Heisenberg uncertainty principle contribute to the idea that electrons occupy "clouds," or "orbitals"? 4. Complete the following table. Principal quantum number, n 1 2 3 4 Number of Sublevels Types of Orbitals 5. The way in which electrons are arranged around the nuclei of atoms is called. Match the name of the rule used to find the electron configurations of atoms with the rule itself. 6. aufbau principle a. When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins 7. Pauli exclusion principle b. Electrons enter orbitals of lowest energy first 8. Hund s rule c. an atomic orbital may describe at most two electrons moving in opposite directions 9. In the shorthand method for writing an electron configuration, what does the sum of the superscripts equal? Write the electron configuration and orbital notation for each of the following atoms. 10. Nitrogen 11. Aluminum 12. Argon 3-10
13. Which guideline, Hund's rule or the Pauli exclusion principle, is violated in the following orbital diagrams? 14. What is the relationship between the principal quantum number and the electron configuration? 15. How does the figure above illustrate Hund's rule? 16. How does the figure above illustrate the Pauli exclusion principle? 17. True of False: The aufbau principle works for every element in the periodic table. 18. Filled energy sublevels are more than partially filled sublevels. 19. Half-filled sublevels are not as stable as levels but are more stable than other configurations. 20. Write the electron configuration of the following atoms: a. carbon b. potassium c. gallium d. copper 21. What is an electron dot structure? 3-11
22. Draw the electron dot structure of each of the following atoms. a. argon b. calcium c. iodine 23. Write the electron configurations for these metals and circle the electrons lost when each metal forms a cation. a. Mg b. A1 c. K Match the noble gas with its electron configuration. 24. argon a. 1s 2 25. helium b. 1s 2 2s 2 2p6 26. neon c. 1s 2 2s 2 2p 6 3s 2 3p 6 27. krypton d. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 13. What is the electron configuration called that has 18 electrons in the outer energy level and all of the orbitals filled? 14. Write the electron configuration for zinc. 15. Fill in the electron configuration diagram for the copper(i) ion. Electron Review 3-12
Electron Configuration 1. Write the electronic configurations, orbital notation, Lewis dot structure for the following: a. S b. C c. P d. Ca e. Zn f. Fe g. A1 2. How many dots would appear in the Lewis electron dot diagram for an atom whose electron configuration ended 4S 2 3d 10 4p 3? 3. How many unpaired electrons does the Lewis dot structure of N have? 4. How many pairs of electrons does the Lewis dot structure of O have? 5. Why does copper have 1 valence electron? 3-13
Electron Configuration Charting Oxidation Number Complete the following chart. You may wish to use the periodic table in your text. 3-14