DO NOT, under any circumstances, throw this away! This packet MUST be saved for the final exam.

Name: Period: Unit 10 Packet Atomic History/Structure and Quantum Mechanics Packet Contents Sheet with Objectives (This Page) Worksheet 1: Thermal energy, radiation, and color Early Discoveries About the Atom Reading Bohr s Model Reading Atomic Structure Notes Atomic Structure Practice What Ions will Main Block Elements Form? Isotopes and Average Atomic Mass Worksheet 2- Analyzing a Spectrograph Worksheet 3- Isotopes and Molar Mass Atomic History and Structure Study Guide Atom, Ion, Isotope Activity Quantum Mechanics Notes Periodic Table (to fill in with the notes) Orbital Diagramming Worksheet Study Guide: Atomic Structure/Electrons DO NOT, under any circumstances, throw this away! This packet MUST be saved for the final exam. Modeling Chemistry Unit 10 Packet Page 1

Unit 10 Structure of the Atom - Objectives Students can distinguish the line spectra of light emitted by atomic gases from the continuous spectra emitted by hot metals related to thermal energy and understand quantum mechanics (Three Main Rules) to be able to complete orbital diagrams and electron configurations for a variety of elements. For each of the scientists Thomson, Millikan, Rutherford, Moseley, Chadwick, Bohr, students can: draw the models of the atom proposed, state the problem with each of the previous models, based on experimental evidence, and describe the experiments conducted to collect the supporting evidence. Score Score 4 Score 3 Score 2 Score 1 Score 0 Scale Comment Without any major errors, students can independently: Students can analyze the line spectra of light emitted by atomic gases and the continuous spectra emitted by hot metals related to thermal energy and use quantum mechanics (Three Main Rules) to be able to complete orbital diagrams and electron configurations for a variety of elements. For each of the scientists Thomson, Millikan, Rutherford, Moseley, Chadwick, Bohr, students can: draw the models of the atom proposed, state the problem with each of the previous models, based on experimental evidence, and describe the experiments conducted to collect the supporting evidence. Without any major errors, students can independently: Students can distinguish the line spectra of light emitted by atomic gases from the continuous spectra emitted by hot metals related to thermal energy and understand quantum mechanics (Three Main Rules) to be able to complete orbital diagrams and electron configurations for a variety of elements. For each of the scientists Thomson, Millikan, Rutherford, Moseley, Chadwick, Bohr, students can: draw the models of the atom proposed, state the problem with each of the previous models, based on experimental evidence, and describe the experiments conducted to collect the supporting evidence. With one or two major errors, students can independently: Students can recognize the line spectra of light emitted by atomic gases and the continuous spectra emitted by hot metals related to thermal energy and understand that quantum mechanics (Three Main Rules) can be used to complete orbital diagrams and electron configurations for a variety of elements. For each of the scientists Thomson, Millikan, Rutherford, Moseley, Chadwick, Bohr, students can: draw the models of the atom proposed, state the problem with each of the previous models, based on experimental evidence, and describe the experiments conducted to collect the supporting evidence. With help from the teacher, students can: Students can distinguish the line spectra of light emitted by atomic gases from the continuous spectra emitted by hot metals related to thermal energy and understand quantum mechanics (Three Main Rules) to be able to complete orbital diagrams and electron configurations for a variety of elements. For each of the scientists Thomson, Millikan, Rutherford, Moseley, Chadwick, Bohr, students can: draw the models of the atom proposed, state the problem with each of the previous models, based on experimental evidence, and describe the experiments conducted to collect the supporting evidence. Even with the teachers help, students show no understanding or ability to: Students can distinguish the line spectra of light emitted by atomic gases from the continuous spectra emitted by hot metals related to thermal energy and understand quantum mechanics (Three Main Rules) to be able to complete orbital diagrams and electron configurations for a variety of elements. For each of the scientists Thomson, Millikan, Rutherford, Moseley, Chadwick, Bohr, students can: draw the models of the atom proposed, state the problem with each of the previous models, based on experimental evidence, and describe the experiments conducted to collect the supporting evidence. Modeling Chemistry Unit 10 Packet Page 2

Name Date Chemistry Unit 10 Worksheet 1 Thermal energy, radiation and color Pd 1. Based on the IR security camera images, what is the relationship between the temperature of an object and the intensity of radiation it emits? 2. Based on the demonstration of the light bulb, what other property of radiation changes with temperature, in addition to intensity? 3. Describe the relationship between the color change of the filament with temperature, and the colors in the observed spectrum of white light. 4. Order the colors of the rainbow according to the amount of energy needed to produce them, from lowest to highest. 5. Based on the spectrum of the fluorescent lamp, what is the main difference between the radiation emitted by a heated solid (filament) and the radiation emitted by an atomic gas (mercury gas in the fluorescent tube)? 6. What can you say about the energy spread of the radiation emitted by an atomic gas? Modeling Chemistry Unit 10 Packet Page 3

Early Discoveries About the Atom You have already learned about the concept of atoms and how they can combine to form compounds. You have also seen that each element has its own characteristic set of properties which help to distinguish it from all other elements. In this chapter we will study the structure of atoms and the laws governing the behavior of the particles that make up atoms. This knowledge will lead to an explanation of the properties of the elements and of their tendencies to form compounds. John Dalton regarded the atom as a particle with no internal parts. He believed an atom to be the smallest possible particle. However, certain experiments were being performed which gave definite indications that Dalton's view was not correct and that there was some sort of internal structure to the atom. It became apparent that atoms consisted of particles that had electrical charges and that these particles interacted according to the laws of electromagnetism. Charged particles carry either a positive (+) or negative (-) charge. We call two negative charges or two positive charges "like" charges. The laws of electromagnetism state that two like charges {1} each other, while unlike charges, {2} each other. Experiments performed during the late 1800's and early 1900's by chemists and physicists made it clear that atoms could, indeed, be broken into smaller parts, contrary to the ideas of {3}. In 1897, J.J. Thomson discovered that atoms could be "taken apart" when he studied the effects of electrical discharge on atoms of various gases. In his experiments he concluded that atoms were coming apart by yielding a stream of negatively charged particles with very small masses (compared to the masses of the atoms). These small negative particles became known as electrons. Thomson is credited with the discovery of electrons which were present in the atoms of all of the different gases that he examined. Another scientist, Ernest Rutherford, and his students performed experiments in England during the first decade of the 20th century in an attempt to determine the size of atoms. In 1906, Rutherford had his students direct a beam of positively-charged subatomic "alpha" particles at a very thin sheet of gold metal. It was known as the alpha scattering experiment." Since they believed that matter was mostly empty space, they expected the particles to pass through the thin sheet unhindered. To their surprise, they found that a small fraction of the particles bounced right back! This led Rutherford to believe that in the center of the atom was a small but very dense "nucleus" with which some of the alpha particles must have collided. He concluded that most of the mass of the atom was contained in the {4}. He also concluded that the nucleus was {5} charged since it repelled the positivelycharged alpha particles. He later said that "It was quite the most incredible event that has ever happened to me in my life. It was almost as if you fired a 15- inch shell into a piece of tissue paper and it came back and hit you." Rutherford realized that electrons were located at a considerable distance from the nucleus. If this were an accurate description of an atom and we could inflate the nucleus of a hydrogen atom to the size of a basketball, the electron would orbit this "basketball" nucleus at a distance of more than 15 miles away! Visualizing the atom like this enables you to realize that most of the atom is, indeed, nothing more than empty space! So much for the athlete who thinks he is "solid muscle!" With regard to the nucleus itself, it became obvious to scientists that the nucleus was composed of small particles. One of the particles in the nucleus is the proton. In 1914, Rutherford was given credit for discovering protons. A proton carries a charge equal to the charge of an electron, but opposite in character. The electron carries a charge of -1 while the proton carries a charge of {6}. In the nucleus of a neutral atom (that is, an atom with no overall electric charge), there must be an equal number of protons to balance the charges carried by the electrons. Unlike the electron, the proton is a particle with Modeling Chemistry Unit 10 Packet Page 4

a relatively large mass, in atomic terms. The proton has a mass equal to 1,836 times that of the electron! Using the common unit of the gram to measure masses, the electron has a mass equal to 9.1 X 10-28 grams, and the proton has a mass equal to 1.673 X 10-24 grams. A second component of the nucleus was discovered in 1932 by another Englishman, James Chadwick. This particle became known as the neutron. A neutron is an electrically neutral particle, meaning that it carries no electric charge. The neutron was very difficult to discover. Because it has no charge of its own, it is neither attracted to nor repelled by an electrical charge. A neutron has a mass slightly larger than that of the proton, equal to 1.675 X 10-24 grams. The presence of this particle accounted for the observed masses of atoms, which were found to be greater than that predicted if only protons were present in the nucleus. Figure 12.1 and the accompanying table present an overall summary of the locations and properties of the components of an atom. Each element differs from all others in that atoms of each element contain a specific number of electrons, protons, and {7}. Indeed, the number of protons in the nucleus determines the actual identity of an element. Determining the number of electrons, protons, and neutrons in any given element is a relatively simple process. The atomic number of an element is equal to the number of protons found in the nucleus. The element with atomic number 19, potassium, has 19 protons. For atoms to be neutral, they must have equal numbers of positive (protons) and negative (electrons) charges. This means that potassium must have 19 protons and {8} electrons. Table 12.1 Location and Properties of Subatomic Particles Particle Charge Comparative Mass Location Electron -1 1/1836 Outside nucleus Proton +1 1 Inside nucleus Neutron 0 1 Inside nucleus If an element has 10 protons in its nucleus, how many positive charges does it have?{9}. If an atom with 10 protons is neutral, how many electrons must it have?{10} The atomic number of oxygen is 8. How many protons does it have?{11} How many electrons does oxygen have?{12} Modeling Chemistry Unit 10 Packet Page 5

Bohr s Model Niels Bohr developed a theory to account for the location of electrons around the nucleus of an atom. He believed his theory would explain the bright-line spectra emitted by excited atoms. He proposed that electrons followed specific paths or orbits around the nucleus. These paths or energy levels, as they are also called, are numbered starting with the lowest one (closest to the nucleus) as 1, the next farther from the nucleus as 2, the next as 3, and so forth. Figure 12.8 illustrates 4 of the energy levels of a n atom. Bohr's model resembled a planetary system like our solar system in which he suggested that the electrons revolve around the nucleus. The first energy level, nearest the nucleus, is represented as number 1. Each level thereafter is increased by one. A total of 7 energy levels are needed to explain the structure of all of the elements. The orbits around the nucleus are called energy levels because there are different and very specific energies associated with each level. Bohr knew that energy was being added to an atom when it was being heated or when an electrical current was passed through an element. This extra energy has to go somewhere. The added energy is absorbed by the electrons that are in the outermost orbit (farthest from the nucleus). Since there are specific amounts of energy associated with each energy level, the electron that absorbs all of this extra energy can no longer stay in the orbit (energy level) in which it normally belongs which is its "ground state." Instead, it will move to another energy level. It is now in an "excited state" and is what we earlier referred to as an "excited" electron. This electron is not doomed to spend the rest of its time in this higher energy level. Indeed, the excited electron is now very unstable. Because this is an unstable state, the electron will soon return to its ground state. This is where the atomic spectra enters into the picture. The energy that the electron emits when it returns to its ground state is in the form of light and heat. Figure 12.9 gives a general picture of this process. In "A, the electron is excited and jumps to a new energy level that is further from the nucleus. As the electron falls back to its ground state in "B", energy is given off in the form of light. As electrons move to lower energy levels, the energy they emit is given off in a "burst" or quantity of energy with a well-defined wavelength. The quantities of emitted energy are called quanta. Bohr's theory of the atom gave birth to the quantum theory, a name that reflects the notion that atoms must absorb and emit energy in specific amounts. Therefore, we say that the energy is "quantized." A Modeling Chemistry Unit 10 Packet Page 6

quantum of energy can be defined as the amount of energy needed to move an electron from one energy level to the next higher one. Similarly, it can be defined as the amount of energy emitted when an electron moves from its present energy level to a lower one. In his theory, Bohr proposed that electrons were only "allowed" to exist at certain distances from the nucleus. These distances became his energy levels. He believed that electrons were not "allowed" to exist between these levels. Although the reasons for this behavior were unclear, the idea did explain the existence of bright-line spectra. The spectrum of hydrogen contains a specific set of bright lines. Of these, only 3 or 4 are clearly visible to the naked eye. (See Figure 12.2) So Bohr asked, why only this set of lines? Why does the excited hydrogen electron emit only this specific set of wavelengths? According to Bohr, the fact that there was always the same set of bright lines in the hydrogen spectrum was evidence that only certain energy changes were possible for the hydrogen electron. The electron could only make certain "jumps" and, therefore, could only emit certain wavelengths of light. The number of lines was limited because there were only a few "excited" energy levels to which the electron was "allowed" to move. He explained that the electron was - for whatever reason - not permitted" to exist between these levels. Since every element has a specific number of electrons, different wavelengths of light will be emitted by excited electrons of atoms of different elements. Even without a spectroscope you can see that the glow of a neon sign is quite different from the fluorescent lights in your classroom. You probably did not realize that what you were seeing was the result of excited electrons at work. To better understand how an electron behaves as it returns to its stable ground state, consider how you would jump down a staircase. Just as an electron can only stop at certain energy levels, you can only stop at certain levels (steps). The electron cannot exist between energy levels, and you cannot stop between the steps. To get down to the bottom of the staircase, you could jump all the way down at once; or one step at a time; or one step followed by a two-step jump; or perhaps a two-step jump followed by one step, and then by a three-step jump. A variety of different jumps is possible, each yielding a different amount of energy. Likewise, suppose an electron jumps out to level 4 in Figure 12.10. There are ten different sizes of jumps possible as it returns to its ground state. So how many different wavelengths of light could be emitted by this electron? {31} A sample containing many excited atoms of this element would emit all possible wavelengths. Modeling Chemistry Unit 10 Packet Page 7

Atomic Structure Notes Neutral Atoms Ions Isotopes Modeling Chemistry Unit 10 Packet Page 8

Atomic Structure Notes P = (+) E = (-) N = (0) Reading a Periodic Table: Li 3 Lithium 6.94 OR 7 3 Li Lithium = Element Name Li = Element Symbol 3 = Atomic Number: number of protons, defines the element * 7 = Mass Number: number of protons and neutrons combined * 6.94 = Atomic Mass: weighted average mass of all isotopes of the element *Never fractional protons, neutrons, or electrons- these must be whole #s Modeling Chemistry Unit 10 Packet Page 9

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Element or Ion Symbol/Ion Atomic # Mass # # of Protons # of Neutrons # of Electrons Nitrogen P 15 30 207 82 40 19 56 82 Oxygen Ion 16 10 17 18 5 2 Modeling Chemistry Unit 10 Packet Page 11

What Ions will the Main Block Elements Form? If an atom loses electrons it forms a ( + or - ) ion. If an atom gains electrons it forms a ( + or - ) ion. In the table below: Draw a Bohr diagram of the first element in each family of main block elements. What ion will the element most likely form to get a full outer shell? On your periodic table write the ion sign (ex: +2) above the column of the element. Main Group Element Bohr Diagram of Neutral Atom Ion Li Be B C N O F He Modeling Chemistry Unit 10 Packet Page 12

Atomic Structure Worksheet Ions Ion Type of Ion Protons Electrons Loss of e - Gain of e - H + Cation 1 0 1 X Li + Na + K + F - Cl - Br - O -2 Ca +2 Sn +4 S -2 Anion 16 18 X 2 Cu +3 N -3 Mg +2 Zn +2 Al +3 P -3 Draw a model for the following atoms and ions: Al Al +3 S S -2 Modeling Chemistry Unit 10 Packet Page 13

Name Date Chemistry Unit 10 Worksheet 2 Pd Analyzing a Spectrograph A mass spectrometer is an instrument used to separate an element's isotopes and to measure their relative abundances. Within this device, a sample of an element is vaporized, then ionized and accelerated down a tube. Near the end, the beam of ions is passed through a strong magnetic field which exerts a force on the ions. Ions of greater mass possess more inertia, or more of a tendency to continue to move in a straight line, and so deviate only slightly from their projected path. Ions of lesser mass are more greatly influenced by the field and demonstrate greater deviation. Examine the three mass spectrograph readings illustrated below and answer the questions that follow. Note that the upper scale of each spectrograph shows atomic mass (in amu). Below each spectrograph, the percents of the various isotopes present are given. Modeling Chemistry Unit 10 Packet Page 14

l. a. What is the molar mass of the isotope of the element represented by spectrum A? b. What are the name and atomic symbol of element A? 2. a. What are the symbols, including superscripts and subscripts for the isotopes in spectrum B? b. Based on the experimentally obtained values of atomic mass and percent abundance, calculate the average molar mass of this element. Show your work. c. Which isotope deviated most from its straight-line path? 3. a. Calculate the average molar mass of the element in Spectrum C. Show your work. b. What are the symbols, including superscripts and subscripts, of the isotopes of this element? c. Which isotope deviated the least from its straight-line path? Modeling Chemistry Unit 10 Packet Page 15

Name Date Pd Chemistry Unit 10 Worksheet 3 1. Element X has two natural isotopes. The isotope with a mass number of 6 has a relative abundance of 7.5%. The isotope with a mass number of 7 has a relative abundance of 92.5%. Determine the average molar mass for the element from these figures. What is the true identity and atomic number of element X? 2. The element copper is found to contain the naturally occurring isotopes 29Cu 63 and 29Cu 65. The relative abundances are 69.1% and 30.9% respectively. Calculate the average molar mass of copper. 3. Uranium has three isotopes with the following percent abundances: 92U 234 (0.0058%), 92U 235 (0.71%), 92U 238 (99.23%). What do you expect the molar mass of uranium to be in whole numbers? Why? 4. A sample of silver as it occurs in nature is 52.0% of isotope 47Ag 107 and 48.0% of isotope 47Ag 108. What is the average molar mass of silver? (Compare your result with the value given in the periodic table). Modeling Chemistry Unit 10 Packet Page 16

5. Ninety-two percent of the atoms of an element have a mass of 28.0 amu, 5.0% of the atoms have a mass of 29.0 amu, and the remaining atoms have a mass of 30.0 amu. Calculate the average molar mass and identify the element. 6. Use the following isotope data for lead to show that its molar mass is 207 amu, 82Pb 204 (1.37%) 82Pb 206 (26.26%) 82Pb 207 (20.82%) 82Pb 208 (51.55%) 7. Boron exists in the form to two stable isotopes, boron-10 and boron-11. These occur in the abundance of 19.6 percent and 80.4 percent respectively. Calculate the average molar mass of boron. 8. Precise molar masses of each isotope of magnesium are given below along with the percent abundance of each isotope: magnesium -24 23.98504 78.70% magnesium-25 24.98584 10.13% magnesium-26 25.98259 11.17% Calculate the average molar mass of magnesium. Modeling Chemistry Unit 10 Packet Page 17

Name Date Pd Chemistry Unit 10 Study Guide: Atomic History and Structure 1. Define: a. Atom b. Ion c. Isotope d. Atomic Number e. Atomic Mass f. Nucleus (Where is it? What s in it? What charge?) g. Electrons (Where are they? What charge?) 2. List the three subatomic particles that make up atoms. 3. List the charge and relative atomic mass (amu) of each of the three subatomic particles that make up atoms. 4. Atomic mass number is the number of +. 5. Which of the elements below are isotopes of each other? a. 12 C b. 1 H + c. 14 C -4 d. 2 H 6. Which of the elements above are ions? 7. A new element Cranium is discovered in Chandler. Samples show that naturally occurring Cranium is 99.2% 250 Ce, 0.6% 255 Ce, and 0.2% 260 Ce. What will the atomic mass of Ce be on the periodic table? 8. If Ce has an atomic number of 112, how many neutrons are in 250 Ce? 9. How many electrons are in a neutral atom of Cranium? Modeling Chemistry Unit 10 Packet Page 18

10. Draw the following atoms using the Bohr model. a. 31 P b. 30 P c. 30 P -3 d. 24 Mg +2 11. List the five postulates of Dalton s Atomic Theory: a. b. c. d. e. 12. Complete the following chart: Chemical Symbol Atomic number Number of protons Number of electrons Number of neutrons Mass number N 15 30 19 40 17 18 56 82 O -2 16 5 2 Charge 13. Complete the attached Atom, Ion, Isotope 4-square activity for Fluorine (F). Modeling Chemistry Unit 10 Packet Page 19

Dalton 14. For each of the scientists listed below a. Identify their major contribution to our understanding of the atom. b. Draw the atomic model of each and order them chronologically from earliest to most recent. Rutherford Democritus Thomson Bohr Chadwick 15. The gold foil experiment was used by to show that the atom had a dense central region called a. 16. The name of Thomson s experiment was and it showed the atom contained. 17. The modern day model of the atom is called the or. Draw a picture of the current day model. 18. Draw Bohr diagrams of the following neutral atoms and identify what ions they form. Element K Neutral Atom Diagram Chemical Symbol of Ion Li Cl F Modeling Chemistry Unit 10 Packet Page 20

Atom (Neutral) Definition: Complete Chemical Symbol Bohr Diagram: Subatomic Particles P = E = N = Ion Definition: Complete Chemical Symbol Bohr Diagram: Subatomic Particles P = E = N = Modeling Chemistry Unit 10 Packet Page 21

Isotope Definition: Complete Chemical Symbol Bohr Diagram: Subatomic Particles P = E = N = Modeling Chemistry Unit 10 Packet Page 22

Quantum Mechanics Notes Key Words And Questions Quantum Mechanical Model of the Atom Notes - Electrons are moving in waves, not in orderly straight orbitals Electron Movement In NOT in Basketball Court Model Structure of Energy Levels Atoms Principle Quantum # (Lowest n = 1. Highest n = 7) Sublevels: Orbitals: Total # of Electrons: s p 2 6 10 Each individual orbital holds electrons. # of sublevels in each quantum level n=1 s n=2 s,p n=3,, n=4,,, Modeling Chemistry Unit 10 Packet Page 23

Notation n Sublevel # of e - 1s 2 1 s 2 3p 6 4d 10 Electron Configuration The distribution of among the of an atom. How many electrons (we ll use this later)? H He Li Be B Mn Rules 1. Aufbau Principle (Snake Diagram): 2. Pauli Exclusion Principle: 3. Hund s Rule (Roommate Rule): Practice: Note- Use snake. Exponents must total the number of electrons in the atom H 1 electron = 1s 1 He 2 electrons = 1s 2 Li 3 electrons = 1s 2 2s 1 Be electrons = B electrons = Mn electrons = Modeling Chemistry Unit 10 Packet Page 24

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Name Date Pd Chemistry Unit 10 Orbital Diagramming Worksheet For each of the following elements draw an orbital diagram (box and arrow form) and write the correct abbreviated electron configuration. Element Symbol Orbital Diagram Electron Configuration 1. H 2. He 3. Li 4. Be 5. B 6. C 7. N 8. O 9. F 10. Ne 11. Na 12. Mg 13. Al 14. Si 15. P 16. S 17. Cl 18. Ar 19. K 20. Ca 21. Br 22. Kr 23. I Modeling Chemistry Unit 10 Packet Page 26

Name Date Pd Chemistry Unit 10 Study Guide: Atomic Structure/Electrons 1. How do isotopes of an element differ from one another? 2. How is an ion formed? 3. An atom contains 3 protons, 4 neutrons, and 3 electrons. What is its atomic number, mass number, and name? 4. What are the two names of the modern atomic model? 5. True or False: Like charges attract and unlike charges repel. Give evidence to support your answer. 6. The distance from one crest of a light wave to the next successive crest of a light wave is called what? 7. Light has a dual nature. List two ways in which light behaves. 8. What is the relationship between frequency, wavelength, and energy? 9. True or False: Waves on the high energy, short wavelength side of the electromagnetic spectrum are NOT dangerous to humans. Give evidence to support your answer. 10. Describe the difference between an orbit and an orbital. 11. What is the principal quantum number of the most stable, lowest energy level and where it is located in comparison to the nucleus? 12. True or False: Electrons must have a certain minimum amount of energy, called a quantum of energy, in order to move from one energy level to a higher energy level. Give evidence to support your answer. 13. Which principal energy level has the lowest energy? The highest? 14. Draw the general shape of the s and p orbitals. 15. How many s orbitals come in a group? p orbitals? d orbitals? f orbitals? 16. How are s orbitals represented in orbital diagrams? p? d? f? Modeling Chemistry Unit 10 Packet Page 27

17. The following electron configurations belong to which elements? 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p 5 18. Write the electron configurations for the following elements: Germanium Calcium Titanium Uranium 19. List and define the three rules that govern how electrons are placed in electron configurations. a. b. c. 20. How does the 3s orbital differ from the 2s orbital? 21. Which sublevels can be found in the fourth principal energy level of an atom? 22. Draw orbital diagrams for each of the following elements: Mg Sc Ne Pt 23. Memorize the family names and location on the periodic table. (Nothing to write here, just get it done) 24. Know the regions on the periodic table for metals, non-metals, and semi-metals. (Nothing to write here, just get it done) Modeling Chemistry Unit 10 Packet Page 28