Name: Period: Date: Note: Light and electrons in atoms have some properties in common so let s learn the basics about light

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1 ELECTRONS & NOTATIONS Name: Period: Date: The Atom & Unanswered Questions: 1) Recall- Rutherford s model, that atom s mass is concentrated in the & electrons move around it. a) Doesn t explain how the electrons were around the nucleus b) Doesn t explain why negatively charged electrons aren t pulled into the positively charged nucleus 2) Early 1900 s, scientists observed certain elements emitted visible when heated in a flame 3) Analysis of the emitted light revealed that an element s chemical behavior is related to the arrangement of the in its atoms. Note: Light and electrons in atoms have some properties in common so let s learn the basics about light Electromagnetic Radiation (ER): 1) = a general term for any type of energy that emanates or radiates outward in all directions 2) = radiation moving at the speed of light; a form of energy that exhibits wavelike behavior as it travels through space. 3) = all the forms of electromagnetic radiation Microwaves 1

2 4) Forms of ER included in the electromagnetic spectrum: a) Gamma rays b) c) Ultraviolet light d) Visible light e) Infrared light f) g) Radio waves 5) Properties that all ER has in common: a) Travels at the ( m/s) b) Emitted by atoms after they are or as they decay c) Acts like both a and a THE WAVE THEORY-The WAVE Nature of Light: 1) All waves can be described by several characteristics: a) ( λ ) = the distance between corresponding points on adjacent waves b) ( ν ) = the number of waves that pass a given point in a specific time, usually one second 2) = contains every wavelength between the wavelength on which the spectrum starts and the wavelength on which the spectrum ends (no gaps) 3) = all the colors of light we can see (rainbow) a) When white light passes through a prism; it is separated into a of its different components (red, orange, yellow, green blue, indigo, and violet light). 2

3 b) The frequency of light varies; increasing energy the frequency. Ex: violet light (greater frequency) has more energy than red light (lower frequency) Red Violet THE PARTICLE THEORY-The PARTICLE Nature of Light: 1) The wave model of light explain all of light s characteristics. a) Why do heated objects emit only certain frequencies of light at a given temperature or why some metals emit electrons when light of a specific frequency shines on them? b) A new model was needed to address these phenomena. 2) When objects are heated, they emit glowing. 3) By studying glowing metals, discovered that only certain wavelengths of light are emitted at each specific temperature. 4) Because temperature is a measure of energy, Planck postulated that each emitted also represents a certain energy. a) Planck concluded that energy is : that it only comes in discrete packets. b) He empirically found that at each temperature the energy was related to the frequency of light by an equation. 5) = a discrete packet of energy that can be gained or lost by an atom 3

4 6) In the, when light hits a metal surface, electrons are ejected. a) This only happens when light of a of frequency is used. b) Lower frequency light will eject electrons. c) explained the photoelectric effect by saying that light acts like which he called photons. 7) = a particle of electromagnetic radiation with no rest mass that carries a quantum of energy. 8) Mystery of the photoelectric effect- When a light s frequency was BELOW a certain minimum (depending on the metal), the photoelectric effect was observed, no matter how long the light shone on a metal. a) The wave theory of electromagnetic radiation explain this!!!!! Dual Wave-Particle Nature of Light: 1) proposed in 1905 that light (and all ER) has a nature; sometimes a beam of light exhibits wavelike and particle-like properties a) In some experiments, light acts like. Ex: Diffraction and interference of light are explained by a wave model of electromagnetic radiation. b) In some experiments, light acts like. Ex: The photoelectric effect is explained by a particle model of electromagnetic radiation. 2) a wave model and a particle model are necessary to explain all observations about electromagnetic radiation.! 4

5 Atomic Line-Emission Spectrum: 1) The wave model of light explain how is light produced in glowing tubes of neon signs 2) = lowest state of energy for an atom 3) = state in which an atom has a higher potential energy than it has in its ground state 4) Light in a neon sign is produced when electricity is passed through a tube filled with neon gas and excites the neon atoms a) The excited atoms emit light to energy. 5) When an excited atom returns to its ground state or a lower energy state, it off energy in the form of ER. 6) = the set of frequencies of the electromagnetic waves emitted by the atoms of the element. 7) If the light emitted by neon is passed through a glass, neon s atomic lineemission spectrum is produced. 8) Each element s atomic line-emission spectrum is and can be used to identify an element 9) HYDROGEN S line-emission spectrum is separated into four specific colors of the visible spectrum. 10) Why can you see just the 4 frequencies of light (lines) & not a continuous spectrum (rainbow) after being shone through a prism? 11) Niels Bohr attempted to answer this? 5

6 The Bohr Model: 1) The Bohr Model ONLY explained observed line-emission spectrum and no other element. a) The single electron in a hydrogen atom can circle the nucleus only in allowed paths ( ). b) A hydrogen atom has a fixed energy for each possible electron orbit. c) When a hydrogen atom is in its, its electron is in its energy orbit, closest to the nucleus. d) When a hydrogen atom is in an, its electron is in a energy orbit, farther away from the nucleus. 2) How does the Bohr Model explain the line-emission spectrum of hydrogen? a) It is only possible for the electron in a hydrogen atom to gain or lose quantities of energy, corresponding to the electron s allowed orbits. b) When the electron falls from a higher energy orbit to a lower energy orbit it loses energy. (Energy does not just disappear. Where did it go?) The atom emits a whose energy corresponds to the energy difference between the two energy levels.(energy lost by electron = energy of released photon) Each frequency corresponds to a certain or line on a line-emission spectrum 3) = The electrons circle the nucleus only in allowed paths (orbits). 4) Illustration of the Bohr Model of calcium: 6

7 The Quantum Mechanical Model of the Atom: 1) In 1924 proposed that have a duel wave-particle nature. Other experiments soon demonstrated wave properties of electrons. 2) In 1926 treated electrons as waves in a model called the of the atom. a) Schrödinger s equation applied equally well to elements than hydrogen. b) Schrodinger s equation: helps determine electron location in an atom 3) = a three dimensional region around the nucleus that indicates the probable location of an electron. (fuzzy electron clouds) a) The cloud has no definite boundary, it is that the electron might be found at a considerable distance from the nucleus Summary of the Models of the Atom: Quantum Mechanical Model (1926) 7

8 Quantum Numbers: Electrons are locked into fixed orbits We can only the areas where are most likely to be found Numbers are given to electrons to help with 1) = specify the properties of the atomic orbitals and the properties of electrons in those orbitals 2) The first three quantum numbers result from solutions to Schrodinger s equation and describe the orbital in which an. quantum number quantum number quantum number 3) The fourth quantum number describes an. quantum number 4) Each electron is given quantum numbers. The 4 Quantum Numbers: 1) The PRINCIPAL QUANTUM NUMBER indicates the (shell) of the orbital in which a particular electron is located. a) The quantum number is symbolized by n. b) Values of n = 1, 2, 3 (7 is the principal quantum number for any known element in its ground state). Corresponds to the 7 on the Periodic Table c) Ex: Energy Levels & the Periodic Table n=1 n=2 n=3 n=4 n=5 n=6 n=7 n=6 n=7 8

9 d) An atom s LOWEST energy level is assigned a principal quantum number of. e) The higher the principal quantum number the away from the nucleus the is f) An electron located in an orbital in a higher main energy level will have a higher. 2) The ANGULAR MOMENTUM QUANTUM number indicates the of the orbital in which a particular electron is located. a) The quantum number divides the main energy levels into smaller groups of orbitals called. b) The angular momentum quantum number is symbolized by l. c) Angular momentum quantum numbers are usually designated with letters d) The order of the sublevels can be remembered as follows: some people don t forget e) Each energy sublevel relates to orbitals of different shape. orbitals are sphere-shaped orbitals are dumbbell-shaped orbitals are double dumbbell-shaped orbitals are flower-shaped f orbitals 9

10 f) Orbitals within each main energy level occupy (subshells). Ex: An electron located in a p orbital in the 2 nd main energy level would be said to be located in the Energy levels can be thought of as rows of seats in a theater. The rows that are higher up and farther from the stage contain more seats. Similarly, energy levels related to orbitals farther from the nucleus contain more sublevels. g) The two sublevels in main energy level 2 are designated and The 2s sublevel corresponds to the 2s orbital, which is spherical like. The 2p sublevel corresponds to three dumbbell-shaped p orbital designated 2p x, 2p y, 2p z. Each of the p orbitals related to an energy sublevel has the energy. h) An electron located in different sublevels of the same energy level would in energy (for multielectron atoms). Ex: An electron located in the 4s sublevel has a lower energy than an electron located in the 4p sublevel which has a lower energy than an electron located in the 4d sublevel which has a lower energy than an electron located in the 4f sublevel. 3) The MAGNETIC QUANTUM NUMBER indicates the orientation of the in which a particular electron is located. a) The quantum number is symbolized by m. b) The orientation of an orbital is designated using a three-dimensional coordinate system with the at the center. c) Orientation of orbitals: An s orbital has possible orientation (a sphere centered on the nucleus). A p orbital has possible orientations. (p x, p y, p z ). A d orbital has possible orientations. (d xz, d yz, d xy, d x 2 - y 2, d z 2 ) An f orbital has possible orientations. (orientations not shown) d) Ex: Orbitals around the Nucleus of a Neon Atom 10

11 4) The SPIN QUANTUM NUMBER indicates the of electron on its axis a) The quantum number is symbolized by s. b) There are two possible fundamental states (spins) for an electron in an orbital. +½ and ½ are used to indicate the two possible states (spins) of an electron in an orbital. Principal Quantum Number: Main Energy Level (n) Summary of the First 4 Energy Levels Type(s) of Sublevel (orbital shapes) # of Orbitals per main energy level Maximum # of Electrons per sublevel 1 s 1 2 Number of Electrons per Main Energy Level (2n 2 ) 2 s p s p 3 6 d 5 10 s p d f

12 WHERE ARE THE ELECTRONS (e-) LOCATED? 1) = Arrangement of electrons in an atom a) Use of an element to indicate the number of. b) = an electron occupies the lowestenergy orbital that can receive it c) : Orbitals of equal energy are each occupied by electron before any orbital is occupied by a second electron All electrons in singly occupied orbitals must have the same d) : No electrons in the same atom can have the same four quantum numbers As a result, two electrons of opposite spin can occupy the same orbital 2) Ground State Electron Configurations: a) = lowest energy arrangement of electrons in an atom b) According to the Quantum Mechanical Model, electrons are said to be located in orbitals. How should this be interpreted? How is this represented? c) in atoms tend to assume arrangements that have the LOWEST possible energies. d) The Aufbau principle, Hund s rule, and the Pauli s exclusion principle can be used to determine the lowest energy, ground state electron configuration for atoms. These are your rules/guidelines! 3) Example Electron Configuration notation of Hydrogen: Hydrogen has Main Energy Level 1s 1 # Electrons in a sublevel Sublevel 12

13 4) = a tool used to help write electron configurations. o Always start with the LOWEST energy level ( ) and work your way through to the HIGHEST energy level ( ) by following the tail to head of each arrow. DIAGONAL RULE: Start here and move along the arrows one 1s 2 by one. Going from tail to head of the arrows 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6 5) Example Electron Configuration of Na (11): 6) Diagonal rule is based on trends of the Periodic Table. 13

14 Use the Diagonal Rule to find the electron configuration notation for the following: (Remember Aufbau s principle= an electron occupies the lowest-energy orbital that can receive it) Ex: Mg (12): Ex: P (15): Ex: Ni (28): Orbital Notation: 1) Shows the number of orbital(s) and the spin of electron(s). 2) Complete electron notation first 3) Orbitals are represented by circle(s): 4) Each orbital can only hold a maximum of electrons 5) Electrons are represented by slash (/) marks 6) Remember: a) Paul s Exclusion Principle: Two electrons of opposite spin can occupy the same orbital. b) Hund s Rules: Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron. All electrons in singly occupied orbitals must have the same spin. 7) Ex: Electron notation of sodium (11) = 1s 2 2s 2 2p 6 3s 1 Orbital notation of sodium = Ex: O (8): Ex: Fe (26): 14

15 Electromagnetic Radiation & Light Worksheet Directions: Answer each of the following questions in the space provided. 1) is radiation moving at the speed of light; a form of energy that exhibits wavelike behavior as it travels through space. 2) is all the forms of electromagnetic radiation (ER). 3) List 2 examples of ER. 4) What is the speed of all forms of ER? 5) Which scientist studied glowing metals and discovered that only certain wavelengths of light are emitted at each specific temperature & represented a certain energy? a) He also concluded that energy is, meaning that it only comes in discrete packets. 6) is a discrete packet of energy that can be gained or lost by an atom 7) A is a particle of electromagnetic radiation with no rest mass that carries a quantum of energy. 8) What happens to electrons in the photoelectric effect, when light of a minimum frequency hits a metal surface? 9) Who explained the photoelectric effect? a) How did he describe light behavior in order to explain the photoelectric effect? 10) Explain what is meant by the dual wave-particle nature of light. 11) Who proposed that light (and all ER) has a dual wave-particle nature? 15

16 12) is the lowest state of energy for an atom 13) is the state in which an atom has a higher potential energy than it has in its ground state 14) What happens when an excited atom returns to its ground state or a lower energy state? 15) is the set of frequencies of the electromagnetic waves emitted by the atoms of the element. 16) How is a line-emission spectrum produced? 17) How many colors of the visible spectrum are seen in hydrogen s line-emission spectrum? 18) What model was used to ONLY explain hydrogen s line-emission spectrum and no other element? 19) When an electron falls from a higher energy orbit to a lower energy orbit it loses energy. Energy does not just disappear. Where did it go? 20) Describe the Bohr Model. 16

17 Quantum Mechanical Model of the Atom Worksheet Directions: Answer each of the following questions in the space provided. 1) Who proposed that ELECTRONS have a duel wave-particle nature? 2) Who came up with the quantum mechanical model of the atom, which is the current model of the atom? 3) An is a three dimensional region around the nucleus that indicates the probable location of an electron. (fuzzy electron clouds) 4) True or False: The electron cloud has a definite boundary. 5) specify the properties of the atomic orbitals and the properties of electrons in those orbitals 6) How many quantum numbers is each electron given? 7) Fill in the information for the quantum numbers: their symbol, what each quantum number indicates. Quantum Number: Symbol: Indicates: Principal Quantum Number Angular Momentum Quantum Number Magnetic Quantum Number Spin Quantum Number 8) What is the lowest value for the principal quantum number? highest? 9) What is the maximum number of electrons in the 3 rd main energy level? 10) What principle states that no two electrons in the same atom can have the same four quantum numbers? 11) What principle states that an electron occupies the lowest-energy orbital that can receive it? 12) What rule states that 1.) orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron and 2.) all electrons in singly occupied orbitals must have the same spin? 17

18 Electron Configuration & Orbital Notations Worksheet #1 1) Write the electron-configuration notation and orbital notation for each of the following. a) B: b) F: c) Ca: d) Br: e) Zn: 18

19 f) Sn: g) U: h) N: i) Al: j) S: 19

20 1) For each of the following identify the element. a) = 1s 2 2s 2 2p 1 b) = 1s 2 2s 2 2p 6 3s 2 3p 2 c) = 1s 2 2s 2 2p 6 3s 2 3p 4 d) = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Electron Configurations Worksheet #2 e) = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 f) = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 5 g) = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 h) = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 i) = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 j) = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 8 2) Given that the electron-configuration notation of oxygen is 1s 2 2s 2 2p 4, answer the following questions: a) How many electrons are in each atom of O? b) What is the atomic number of this element? c) How many unpaired electrons does an atom of oxygen have? d) How many orbitals are only occupied by 1 electron? e) What sublevel is the unpaired electron located? f) What is the highest occupied main energy level? g) How many electron does an atom of oxygen contain in the 2 nd main energy level? 3) Given that the electron-configuration notation of phosphorous is 1s 2 2s 2 2p 6 3s 2 3p 3, answer the following questions: a) How many electrons are in each atom of P? b) What is the atomic number of this element? c) How many unpaired electrons does an atom of phosphorous have? d) What is the highest occupied main energy level? e) How many orbitals are only occupied by 1 electron? f) How many electron does an atom of phosphorus contain in the 3rd main energy level? 20

21 4) Answer each of the following questions. You may need to refer to previous pages to see the electron configurations and orbital notations. a) What is the maximum number of electrons contained in an orbital? b) What is the maximum number of electrons contained in a p sublevel? c) What is the maximum number of electrons in a s sublevel? d) What is the maximum number of electrons contained in a d sublevel? e) How many electrons are contained in a full second main energy level? f) How many main energy levels are occupied in a Sn atom? g) How many electrons are contained in the highest occupied energy level of an Al atom? h) How many unpaired electrons does an atom of Al have? i) How many orbitals are empty in an Al atom? j) How many electrons are contained in the highest occupied energy level of a S atom? k) How many electrons are contained in the 2 nd main energy level of a S atom? l) How many unpaired electrons does a S atom contain? m) Which of the following does NOT exist? 1s, 2p, 2d, 3s, 5f, 4d n) Name the element that contain a single 4s electron. o) How many orbitals contain at least one electron in a F atom? p) How many unpaired electrons does an U atom contain? q) Which of the following is the highest main energy level? 3p, 4s, 3d, 2p DIAGONAL RULE: Start here and move along the 1s 2 arrows one by one. 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6 21

22 1) List two examples of ER Electrons & Notations Review Worksheet 2) What is meant by the Dual Wave-Particle Nature of Light/Electrons? 3) In the following, 1s 2 2s 2 2p 5, how many orbitals only have one electron? 4) is radiation moving at the speed of light; a form of energy that exhibits wavelike behavior as it travels through space. 5) is the lowest state of energy for an electron 6) is a general term for any type of energy that emanates or radiates outward in all directions 7) is a discrete packet of energy that can be gained or lost by an atom 8) is a three dimensional region around the nucleus that indicates the probable location of an electron. (fuzzy electron clouds) 9) specifies the properties of the atomic orbitals and the properties of electrons in those orbitals 10) Who developed the quantum mechanical model that treated electrons in atoms as waves & helps determine probable electron location in an atom? 11) All the forms of electromagnetic radiation are called the 12) List one property that all ER has in common. 13) How many quantum numbers is each electron given? 14) What quantum number indicates the spatial orientation of the orbital about the nucleus in which an electron is located? 15) What quantum number indicates the shape of the orbital in which an electron is located? 16) What letters represent sublevels? 17) No two electrons in the same atom can have the same four quantum numbers (two electrons of opposite spin can occupy the same orbital) is known as what? 18) What is the maximum number of electrons in a s sublevel? 19) An electron occupies the lowest-energy orbital that can receive it is known as what? 22

23 20) What quantum number indicates the spin of electron on its own axis? 21) In the following, 1s 2 2s 2 2p 5, what atom is represented? 22) A is a particle of electromagnetic radiation with no rest mass that carries a quantum of energy. 23) What is the maximum number of electrons contained in an orbital? 24) All electrons in singly occupied orbitals must have the same spin is known as what? 25) How many orbitals does a f sublevel have? 26) In the following 2s 1, what does the 1 represent? 27) The state in which an electron has a higher potential energy than it has in its ground state is the 28) In the following 2s 1, what does the 2 represent? 29) Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron is known as what? 30) How many electrons are contained in a full second main energy level? 31) In the following 2s 1, what does the s represent? 32) What is the maximum number of electrons contained in a p sublevel? 33) What quantum number indicates the main energy level that the electron is found? 34) What is the maximum number of electrons contained in a d sublevel? 35) How many orbitals does a p sublevel have? 36) Explain what happens when an electron drops from a higher energy state to a lower energy state. 37) Why do we see different colors during a flame test? 38) How many orbitals does a d sublevel have? 39) In the following, 1s 2 2s 2 2p 5, how many electrons are present in the highest main energy level? 40) In the following, 1s 2 2s 2 2p 5, how many unpaired electrons are present? 41) Who explained the photoelectric effect? 42) How many colors of the visible spectrum are seen in hydrogen s line-emission spectrum? 43) What model was used to ONLY explain hydrogen s line-emission spectrum and no other element? 44) Who proposed that light (and all ER) has a dual wave-particle nature? 45) What is the lowest value for the principal quantum number? highest? 23

24 Unit Learning Map (7 days): Electrons & Notations Mrs. Hostetter Class: Academic Chemistry A - Grade 11 Unit Essential Question(s): How does electron arrangement in atoms affect chemical properties? Optional Instructional Tools: Lab Materials: Felt-tip Distribution Lab Flame Lab Guided Notes Handouts/worksheets Concept Concept Concept Concept Electromagnetic radiation What is electromagnetic radiation? Bohr Model vs. Quantum Mechanical Model What is the difference between the Bohr model & the quantum mechanical model of the atom? Electron Configurations Lesson Essential Questions: Lesson Essential Questions: Lesson Essential Questions: Lesson Essential Questions: How are electron configurations written? Vocabulary: Vocabulary: Vocabulary: Vocabulary: Radiation Electromagnetic radiation Electromagnetic spectrum Quantum Photon Line-emission spectrum Photoelectric effect Ground state Excited state Max Planck Albert Einstein Erwin Schrodinger Quantum Mechanical Model Orbital Quantum numbers Principal quantum number Angular momentum quantum number Magnetic quantum number Spin quantum number Electron configuration Ground state electron configuration Aufbau principle Hund s rule Pauli s exclusion principle 24

25 Electrons & Notations Vocabulary: 1) Radiation = a general term for any type of energy that emanates or radiates outward in all directions 2) Electromagnetic radiation (ER) = radiation moving at the speed of light; a form of energy that exhibits wavelike behavior as it travels through space. 3) Electromagnetic spectrum = all the forms of electromagnetic radiation 4) Quantum = a discrete packet of energy that can be gained or lost by an atom 5) Photon = a particle of electromagnetic radiation with no rest mass that carries a quantum of energy. 6) Photoelectric effect = when light hits a metal surface, electrons are ejected. 7) Max Planck =studied glowing metals and discovered that only certain wavelengths of light are emitted at each specific temperature & represented a certain energy. H concluded that energy is quantized; meaning that it only comes in discrete packets. 8) Bohr Model =the electrons circle the nucleus only in allowed paths (orbits). Only explained hydrogen s line-emission spectrum 9) Quantum Mechanical Model = (proposed by Erwin Schrodinger) helped determine probable electron location in an atom & applied equally well to elements other than hydrogen. 10) Albert Einstein = explained the photoelectric effect by saying that light acts like particles which he called photons & proposed that light (and all ER) has a dual wave- particle nature 11) Atomic line-emission spectrum = the set of frequencies of the electromagnetic waves emitted by the atoms of the element. 12) Ground state = lowest state of energy for an atom 13) Excited state = state in which an atom has a higher potential energy than it has in its ground state 14) orbital = a three dimensional region around the nucleus that indicates the probable location of an electron. (fuzzy electron clouds) 15) Quantum numbers = specify the properties of the atomic orbitals and the properties of electrons in those orbitals 16) Principal quantum number = Indicates the main energy level that the electron is found (#1-7) 17) Angular momentum quantum number = indicates the shape of the orbital in which an electron is located 18) Magnetic quantum number = indicates the spatial orientation of the orbital about the nucleus in which an electron is located 19) Spin quantum number = indicates the spin of electron on its own axis 20) Electron Configurations = arrangement of electrons in an atom 21) Ground state electron configuration = lowest energy arrangement of electrons in an atom 22) Aufbau Principle = an electron occupies the lowest-energy orbital that can receive it 23) Hund s Rule = a.) Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron b.) All electrons in singly occupied orbitals must have the same spin 24) Pauli s Exclusion Principle = a.) No two electrons in the same atom can have the same four quantum numbers b.) As a result, two electrons of opposite spin can occupy the same orbital 25