UNIT 5 NOTES: MODELS OF ATOMIC THEORY

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1 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 1 NAME PERIOD UNIT 5 NOTES: MODELS OF ATOMIC THEORY STUDENT OBJECTIVES: Your fascinating teachers would like you amazing learners to be able to 1. Perform calculations involving the speed of light (c), the wavelength (λ), and the frequency (ν) of electromagnetic radiation (EMR). 2. Perform calculations involving the energy, the wavelength (λ), and the frequency (ν) of electromagnetic radiation. 3. Place forms of electromagnetic radiation in correct energy order on an electromagnetic spectrum. 4. Compare the energy, frequency, and wavelength of EMR. 5. Label an energy level diagram, indicating both excitation and relaxation. 6. Describe the four quantum numbers used to define the region of space that has a 90% probability of finding an electron with a given energy. 7. Explain the information provided by each of the quantum numbers. 8. Distinguish between an energy level, sublevel, and an orbital. 9. List the allowable quantum numbers for each energy level. 10. Draw an orbital diagram using the Aufbau Principle, Hund s Rule, and the Pauli Exclusion Principle. 11. Write the electron configuration for and atom or an ion. 12. Be able to identify the highest-energy electrons, the valence electrons, and the number of valence electrons. 13. Using the periodic table, describe the periodic trends in atomic radii, ionic radii, ionization energy, reactivity, and electronegativity.

2 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 2 LEGGETT Atom EMR (6:56) I. LIGHT AND ELECTRON CONFIGURATION: Much of what we know about the atom has been learned through experiments with light; thus, we need to know some fundamental concepts of light in order to understand the structure of the atom, especially the placement of the electrons. 1. Characteristics of Light: Has "Dual" nature (or split personality) - there are times in which it behaves like a continuous WAVE and other times when it behaves like a PACKET of light (a blip of light!!!) a. Light as a Wave i. Speed of light in a vacuum is the most accurately known constant, c, in the universe. c = 3.0 x 10 8 meters/sec OR 3.0 x nm/sec (assuming a vacuum) ii. Wavelength is represented by the Greek letter, ( ): It is the length between corresponding parts of adjacent crests and can be expressed in ANY units of length you choose (feet, inches, meters, kilometers, etc) iii. Frequency is represented by the Greek letter, ( ): It is the number of wave crests which pass a given point in 1 second. Its units are, most commonly, cycles / second, simply sec -1. The unit, Hertz (Hz) is also used as a unit label. b. Electromagnetic Spectrum i. Any visible light will have a wavelength of between nm and nm. The longest wavelength in visible is red (about 700 nm) and the shortest visible wavelength is violet (about 400 nm) The only thing that makes one color of light different from another is its and - the velocity is always the same (the speed of light!). The colors of visible light in order from longest wavelength to shortest spell out ROY G BIV. ii. NOT ALL LIGHT IS VISIBLE! The entire range of light energy is considered the Electromagnetic Spectrum. Example 5-1. Label the following electromagnetic spectrum from lowest energy on the left hand side to highest energy on the right hand side. 700 nm R O Y G B I V 400 nm RED MARTIANS INVADE ROYGBIV USING X-RAY GIZMOS

3 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 3 Leggett PAP Chem Atomic Models 5-2 (14) Example 5-2. Do you think a longer or shorter wavelength would have the HIGHER frequency? Example 5-3. What mathematical relationship can you draw about wavelength and frequency? II. CALCULATIONS WITH LIGHT FORMULAS ON THE PERIODIC TABLE: hc c = E photon = h E photon λ CONSTANTS ON THE PERIODIC TABLE: c = 3.0 x 10 8 m/s (speed of light in meters) h= 6.63 x J sec (Planck s constant) Note: BE CAREFUL OF CHOOSING SPEED OF LIGHT FOR YOUR FORMULA!!! You must make the units of wavelength match up with your units for speed of light. Remember that c can also be 3.0 x nm/s. Units of frequency can be reported as 1 s (per second), second 1, or Hertz. They are interchangeable with each other! Example 5-4. If the wavelength is known to be 550 nm, what is its frequency? Example 5-5. If the frequency of light is known to be 9.45 x s -1, is the light visible? If not, is it UV or IR? HINT: Calculate the wavelength in nm. Example 5-6. If the wavelength of light is 7.23 x 10-5 nm, what is the frequency of light? Example 5-7. What is the energy of a photon of light whose frequency is 7.85 x Hz? Is the light visible? How do you know? Example 5-8. If a light has a wavelength of 550 nm, what is the energy of one photon of this light? Example 5-9. If one photon of light is known to have energy of 3.33 x J, is it visible?

4 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 4 LEGGETT Atom Bohr & AES (13:49) III. BOHR ATOM: Don t Forget! Bohr said that electrons are in energy levels which is true! And his research with light & electron energy levels is very important and lead to a lot of what we know about the atom! But he also proposed the planetary model which is not true. a. Electron cloud made up of which are like the of a ladder and are BUT they are not. b. Amount of energy possessed by an electron determines its. c. Levels farther from nucleus represent energy. d. Electrons like to hang out where they are most, which is their. This is the exact amount of energy they like to possess it s the electron s home. All electrons have their own ground state on Level 1, Level 2 or Level 8! e. When electrons gain a quanta (amount) of energy (through sources such as extreme heat or electricity), the electron will up energy levels in a process called. This excited state is very!!! f. Electrons can t stay excited at the higher energy level forever! When they fall back to their ground state, they will lose the extra energy in the form of a packet of light emitted in a process called. The larger the fall, the larger energy of light released! When the electrons lose the energy they gained, they release LIGHT (EMR) n= n= n= n= n= e - energy Process e - EMR (energy) Process Ground State:

5 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 5 g. Each possible energy fall corresponds to a amount which means that it corresponds to a different of. Some of the falls can be detected by the human eye (as they fall in the visible range), but other falls that are not visible must be detected by instruments (as they are infrared or ultraviolet). h. Different elements have different energies associated with their energy levels. This means that similar falls in will produce. i. When we observe a color being produced from an element undergoing relaxation, we are observing all of the VISIBLE falls simultaneously occurring. We can use things like a prism or a spectroscope to break apart that light into its individual colors resulting in a! j. We can use the resulting bright-line spectra to identify an element! Example Based on the Bright-Line Spectra shown above, identify the elements found in the mixture.

6 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 6 Leggett PAP Chem Atomic Models 5-4 (12) IV. SCHRODINGER MODEL: Quantum Mechanical Model (1926) a. Schrodinger with Bohr about quantized energy levels. b. However, Schrodinger in the fact that he does not define an path called an orbit. Schrodinger s model describes a region in space with a high of finding an electron. Unfortunately, we still call this an. ANALOGY: It is probable that you will be in the region of space defined by my classroom during this specified period BUT we don t know where you are in the room and how you are moving around the room. We re just 90% sure we will find you (maybe you are absent, in the office, etc.). My classroom is the orbital! c. ELECTRON LOCATION: Electrons are distributed around the nucleus in specific ways and each electron has four coordinates to describe its most probable location or distribution: 4 quantum numbers, n, l, m, and s. Think of these as being the address for an electron s probable location! (State, city, street, house number each gets more and more specific.) n = (Principle Quantum Number) Indicates how far from the nucleus the electron is located and its energy. It is designated by the Principle Quantum Number, n. The higher the value of n, the away from the nucleus and the the energy level The number of electrons in any energy level: (n=energy level) Example Fill in the chart with how many electrons each energy level can hold. First Energy Level Second Energy Level Third Energy Level Fourth Energy Level Leggett PAP Chem Atomic Models 5-5 (8.5) l = (Second Quantum Number) ** Each energy level is divided into which are designated by the Second Quantum Number, l. Actually not a number at all, but rather a letter! Here is a trick for remembering the order: Stupid People Die First. ** Indicates the of the region. 1. Sublevels are labeled by the letters s, p, d, and f. 2. The first energy level will have 1 sublevel and its label is "s". 3. The 2 nd energy level will have 2 sublevels; "s" and "p". 4. The 3 rd energy level will have 3 sublevels; "s", "p", and "d" 5. The 4 th energy level will have 4 sublevels; "s", "p", "d", "f"

7 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 7 s sublevels have a shape. p sublevels have a shape. d and f sublevels have really weird shapes!!! Leggett PAP Chem Atomic Models 5-6 (10) m l = (Third/Magnetic Quantum Number) Each sublevel contains which are regions of high for finding electrons the m l indicates the of those orbitals! The number of orbitals is related to the sublevels (l) Each orbital can hold only electrons The surface of the orbital is drawn to represent area where any particular electron can be found of the time. Each sublevel has a set number of orbitals: s sublevel - orbital p sublevel - orbitals d sublevel - orbitals f sublevel - orbitals

8 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 8 Leggett PAP Chem Atomic Models 5-7 (7) Example Complete the table: Energy Level Sublevel(s) Orbitals (use lines to depict) # e per sublevel 1 # e per energy level m s = (Fourth Quantum Number) ** The electrons within an orbital must have spins to overcome their repulsion, which we represent with opposite arrows. ** Spin is either + ½ ( ) or - ½ ( ) We are going to use these FOUR Quantum Numbers to describe the location of electrons, like your ticket describing your seat for a sporting event! Each get more and more specific. Think of it like this: level n = stadium level (lower, platinum, nosebleed) = energy l = section number = sublevel m l = row = orbital s = seat number = spin

9 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 9 LEGGETT Atom Models (12:46) We will look very soon at how to use all of these numbers to describe an electron s probable location, just like level, section, row, and seat number would help us find you at the game (assuming of course you are not out getting nachos)!!! V. ELECTRON CONFIGURATIONS AND ORBITAL DIAGRAMS Knowing the rules for electron arrangement allows us to write what is called the electron configuration, which gives a description of where in that atom the electrons are actually located. We can also draw an orbital diagram which gives the same information with more detail. provide a lot of information (since they show ALL FOUR quantum numbers), but they become annoying to draw. are a shorter way of showing arrangements of electrons around the nucleus. They don t provide all of the information that an orbital diagram would, but they do show us quite a bit! Each element has its own electron configuration. Electron configuration gives specifically the energy level, sublevel, and number of electrons in that sublevel. WE HAVE SOME BACKGROUND INFO WE NEED TO GO OVER BEFORE WE CAN WRITE THESE ELECTRON CONFIGURATIONS & ORBITAL DIAGRAMS Remember that n provides information on both energy and distance. The relationship between n and DISTANCE stays true! Electrons in higher n energy levels are more likely to be found further from the nucleus. In a perfect world we would have the following energy order: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f etc. UNFORTUNATELY, things get a little mixed up in terms of energy meaning it doesn t fill up in the order we quite expect. We will learn the order for filling sublevels in just a minute. FOR NOW: trust us! We are going to ultimately show you that the periodic table provides us with the perfect arrangement to break the code of filling order! The electrons are the electrons that are in the highest energy level, n. The electrons are in the last sublevel that was filling. ALSO: There are some rules we must follow for arranging electrons Aufbau Principle: Electrons are placed in orbitals of energy first. Pauli Exclusion Principle: An orbital may hold only electrons. Think of it like an exclusive club with only 2 members! Why can t we have an orbital with more electrons? Well, NO TWO ELECTRONS CAN HAVE THE SAME 4 QUANTUM NUMBERS. The combination of 4 quantum numbers is for that electron only! If they have the same energy (n), are in the same sublevel (l), and are in the same orbital (m l), then they must have opposite spins meaning 4 different quantum numbers. REMEMBER: We designate those spins with arrows! for +1/2 for 1/2 Hund s Rule: When filling a sublevel (which contains orbitals of equal energy), one electron enters each orbital until each, then they!

10 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 10 ABOUT IONS: Most metals will electrons to become with the nearest noble gas. When they lose electrons, the ion will be smaller than the original atom. Non-metals will electrons when they are with metals to become with the nearest noble gas. When they gain electrons, the ion will be larger than the original atom. In Noble gases, electrons have maximized attraction to the nucleus and minimized repulsion with each other as much as possible, due to their electron configuration. They are STABLE and do not react aka INERT. When substances become ions, they want to copy-cat the ideal electron configurations of the noble gases!!! Example For each of the following, predict the charge of the most likely ion and indicate which sublevel gains or loses electrons. Group Electron Configuration of Valence Electrons? Number of Valence Electrons? Lose or Gain Electrons to look like Noble Gas? Common Charge of Ions? Would resulting ion be larger or smaller than the original atom? nonmetals 16 nonmetals 17 nonmetals

11 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 11 PERIODIC TABLE and ELECTRON CONFIGURATION/ORBITAL DIAGRAMS 1 2 Group 1 (IA) Group 18 (VIIIA) 1 2 1s 1 1s 2 Group 2 (IIA) Group 13 (IIIA) Group 14 (IVA) Group 15 (VA) Group 16 (VIA) Group 17 (VIIA) ATOMIC NUMBER s 1 2p 1 2p 3 2p 4 2p s 2 Period 4 Group 3 (IIIB) Group 4 (IVB) Group 5 (VB) Group 6 (VIB) Group 7 (VIIB) Group 8 (VIIIB) Group 9 (VIIIB) Group 10 (VIIIB) Group 11 (IB) Group 12 (IIB) d 2 4p s 1 4d 3 4d 5 4d 8 5p s 2 5d 10 6p 3 6p d 4 7p 1 7p f 1 4f f 13 Example Example Fill in the missing outer configurations in the periodic table above. Answer the following question about the periodic table above. a) How many electrons does an s- sublevel hold? b) How many elements are in a row in the s block? Hmmm c) How many elements are in a row in the p block? d) Remember that each orbital can hold TWO electrons How many orbitals would be needed for the p block? Hmmm e) How many elements are in a row in the d block? f) Remember that each orbital can hold TWO electrons How many orbitals would be needed for the d block? Hmmm are you noticing a recognizable pattern? g) How many elements are in a row in the f block? h) Remember that each orbital can hold TWO electrons How many orbitals would be needed for the f block? Hmmm i) What is the correlation between the period number and the energy level, n, for the s and p blocks? j) What is the correlation between the period number and the energy level, n, for the d block? k) What is the correlation between the period number and the energy level, n, for the f block?

12 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 12 5s 4s 3s 5p 4p 3p 4d 3d 2s 2p 1s Example (a) Write the electron configuration for Cobalt. Then, fill out the orbital diagram above for Cobalt. (b) What are the highest-energy electrons in Cobalt? (c) What are the outer-most (valence) electrons in Cobalt? (d) How many valence electrons does Cobalt have?

13 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 13 Example Work the following regarding Sulfur. (a) Write the electron configuration. (b) Write the orbital diagram across from left to right (rather than upwards). (c) What are the highest-energy electrons? (d) What are the outer-most (valence) electrons? (e) How many valence electrons are there? (f) What is the most likely ion for Sulfur? Did we lose or gain electrons to form the ion? (g) What would be the new electron configuration for the Sulfur ion? (h) Would the resulting ion be larger or smaller than the original atom? (i) With which noble gas is the Sulfur ion isoelectronic? Example Work the following regarding Strontium. (a) Write the electron configuration. (b) Write the orbital diagram across from left to right (rather than upwards). (c) What are the highest-energy electrons? (d) What are the outer-most (valence) electrons? (e) How many valence electrons are there? (f) What is the most likely ion for Strontium? Did we lose or gain electrons to form the ion? (g) What would be the new electron configuration for the Strontium ion? (h) Would the resulting ion be larger or smaller than the original atom? (i) With which noble gas is the Strontium ion isoelectronic?

14 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 14 Isotope: ALUMINUM 27 (NO CHARGE) Nuclear Symbol: Atom or Ion? # of Protons: # of Neutrons: # of Electrons: Circle the Highest-Energy Electrons. Box the Valence Electrons. NUCLEUS Number of Valence Electrons: Electron Configuration: Orbital Diagram: Example Answer the following regarding the Aluminum isotope above. (a) What is the most likely ion for Aluminum? Did we lose or gain electrons to form the ion? (b) What would be the new electron configuration for the Aluminum ion? (c) Would the resulting ion be larger or smaller than the original atom? (d) With which noble gas is the Aluminum ion isoelectronic?

15 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 15 Isotope: BROMINE - 71 (NO CHARGE) Nuclear Symbol: Atom or Ion? # of Protons: # of Neutrons: # of Electrons: Circle the Highest-Energy Electrons. Box the Valence Electrons. NUCLEUS Number of Valence Electrons: Electron Configuration: Orbital Diagram: Example Answer the following regarding the Bromine isotope above. (a) What is the most likely ion for Bromine? Did we lose or gain electrons to form the ion? (b) What would be the new electron configuration for the Bromine ion? (c) Would the resulting ion be larger or smaller than the original atom? (d) With which noble gas is the Bromine ion isoelectronic?

16 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 16 Isotope: ZINC - 64 (NO CHARGE) Nuclear Symbol: Atom or Ion? # of Protons: # of Neutrons: # of Electrons: Circle the Highest-Energy Electrons. Box the Valence Electrons. NUCLEUS Number of Valence Electrons: Electron Configuration: Orbital Diagram: Example Answer the following regarding the Zinc isotope above. (a) What is the most likely ion for Zinc? Did we lose or gain electrons to form the ion? (b) What would be the new electron configuration for the Zinc ion? (c) Would the resulting ion be larger or smaller than the original atom? (d) Is there any noble gas that the Zinc ion is isoelectronic with?

17 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 17 Example Let s see how to include the F Block!!! Write the electron configuration for the following. a) Tungsten ( #74) b) Gold (#79) ABBREVIATED ELECTRON CONFIGURATIONS Writing these long, repetitive configurations can be a pain in the neck and takes up a lot of space. Space is at a premium on a periodic table, so scientists developed an abbreviated form. To write abbreviated electron configurations, look to the noble gas in the prior period. Brackets around the symbol indicate that the element has the same electron configuration of that noble gas PLUS whatever follows. Note: You cannot use these if a question says to write the complete electron configuration! For iron (Fe): 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 BECOMES: [Ar] 4s 2 3d 6 Example Write the abbreviated electron configurations for the following elements. a) Antimony (#51) b) Palladium (#46) c) Mercury (#80) IDENTIFYING ELEMENTS USING ELECTRON CONFIGURATIONS Using either the total count of electrons (by totaling up the exponents on a complete electron configuration), or by looking at the last component of the electron configuration, we can determine what element we have by using the electron configuration. Example Identify the neutral elements that have the following electron configurations. 1s 2 2s 2 2p 6 3s 2 3p 3 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 7 [Xe] 6s 2 4f 14 5d 10 6p 3 [Rn]7s 2 5f 14 6d 5

18 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 18 VI. PERIODIC TRENDS Use these notes to help with your Inquiry! GROUPS (aka Families) run VERTICALLY on the periodic table. Trends (such as reactivity and some physical properties) tend to be consistent within a group. PERIODS run HORIZONTALLY on the periodic table. As you work through the periods on the periodic table, trends tend to repeat themselves from period to period (such as number of valence electrons). However, there are some specific trends that can be summarized by looking at the periodic table as a whole. The following contains a description of those trends. Realize with all of these trends, we are giving general descriptions of how these trends work as with anything in chemistry; there are exceptions to the rules! ATOMIC SIZE (RADIUS) The radius is ½ of the distance measured from the center of one atom to the center of an adjacent identical atom. The bulk of this distance is the electron playground! Atomic size decreases across a period, due to the increased number of protons having a greater pull on the valence electrons more as you move across. Atomic size increases down a group, as you are adding more energy levels as you go down. This would mean that Fr (Francium) would theoretically have the largest atomic size. He (Helium) would have the smallest atomic size. EX: Which would have a larger atomic size Ca, Br, or Ba? IONIZATION ENERGY This refers to the amount of energy required to remove one outer-most (or VALENCE) electron. Since 8 valence electrons indicate a full outer shell (s & p sublevels), the closer the number of valence electrons is to 8, the more difficult it will be to pull off an electron - meaning it will take more ionization energy. Ionization energy increases across a period, because the atom is more tightly held together as the valence sublevels become full meaning it will take more energy to pull an electron off. Ionization energy decreases down a group, because if you have more energy levels, the outer-most electrons are further away, meaning they aren t held on as tight meaning it takes less energy to pull them off. This would mean that He (Helium) would theoretically have the highest ionization energy. (Yes, he only has two valence electrons, but that s all he can handle!) Fr (Francium) would have the lowest ionization energy. EX: Which would have a higher ionization energy Ca, Br, or Ba?

19 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 Page 19 ELECTRONEGATIVITY This refers to the attraction an atom has to gain a nearby electron. The closer an atom is to 8 valence electrons (in the s & p sublevels), the higher the electronegatvitiy (because the nucleus has a stronger + charge!) However, would an atom that already has 8 valence electrons want to gain another? Electronegativity increases across a period, because as you move across, the atom has a greater pull to gain another electron (to achieve the 8 valence electrons). Electronegativity decreases down a group, because as the valence electrons get further from the nucleus, it has less of a pull to gain more electrons. Noble (Inert) gases have no electronegativity, because they already have 8 valence electrons (filled s & p sublevels), so they have no pull to gain any more! This is also part of the reason they are considered INERT meaning unreactive. This would mean that F (Fluorine) would theoretically have the highest electronegativity. Fr (Francium) would have the lowest electronegativity. EX: Which would have a higher electronegativity Ca, Br, or Ba? USE THE PERIODIC TABLE BELOW TO MAKE SUMMARIES ABOUT PERIODIC TRENDS!!!