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Unit 4 Notes: Chapter 5 Electrons in Atoms Page |1 Section 5.1- Models of the atom Goals for this section: - Identify the inadequacies in Rutherford’s model - Identify the new proposal in the Bohr model of the atom - Describe the energies and positions of electrons according to the quantum mechanical model of the atom - Describe the shapes of the orbitals Unit 4 Notes: Chapter 5 I. II. III. Electrons in Atoms Page |2 Identify the number of orbitals and electrons in each sublevel Identify the sublevels in each principle energy level Identify the last sublevel and principle energy level ofelectrons in an element using the periodic table Identify the number of electrons in the last sublevel and principle energy level in an element using the periodic table Know the following terms: Energy level, quantum, quantum mechanical model, atomic orbital, Rutherford’s model A. Rutherford’s model did not explain of elements B. His model also did not explain why objects when heated Bohr’s Model A. Bohr proposed that an electron is found only in , around the nucleus. B. Each possible electron orbit in Bohr’s model has a . C. The fixed energies an electron can have are called . D. A quantum of energy is the amount of energy required to move an electron from E. The higher the energy level occupied by an electron, the it takes to move from that energy level to the next higher energy level. Current atomic model- Quantum Mechanical Model A. Still has the nucleus containing in the center. B. Still has electrons outside the nucleus in a low density area C. The quantum mechanical model determines the an electron can have and how likely it is to find the electron in various locations around the nucleus. D. This model is based on equations developed by Erwin Schrodinger E. The probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud. The cloud is where the probability of finding the electron is high. 1) The fuzzy cloud is called an electron cloud F. Principle Energy Levels of Electrons 1) Energy levels and # of electrons in each: a) Level 1 Holds b) Level 2 Holds c) Level 3 Holds d) Level 4 Holds 2) Principle energy levels are broken down into sublevels called: 3) An atomic orbital is often thought of as a region of space in which there is a of finding an electron. 4) Each energy sublevel corresponds to an orbital of a , which describes where the electron is likely to be found. 5) Each orbital contains up to 6) Orbitals include: 7) Summary of Energy Levels and Orbitals Unit 4 Notes: Chapter 5 Sublevel Electrons in Atoms # of Orbitals Page |3 Maximum # of Electrons in sublevel S P D F 8) Summary of Energy levels and orbitals Principle # of Sublevels Identity of Energy Level Sublevels 1 2 3 4 Shape # of orbitals Maximum # of electrons The n represents the principle energy level The values for “n” are the same as the period numbers. Notice the s block has 2 columns because s can hold 2 electrons. The d block has 10 columns because d can hold up to 10 electrons The p block has 6 columns because p can hold up to 6 electrons You can use this diagram to help you identify the last electrons in the series of electrons in an element. For example: magnesium is in period 3 column 2 in the s block. This means the last electrons in the series are in the principle energy level of 3 and the sublevel of s. Since it is in column 2, there are 2 electrons in the s sublevel. Sulfur is in period 3 in the 4th column of the p block. This means the last electrons are in the principle energy level of 3 in the sublevel of p. There are 4 electrons in the p sublevel. F block is the part of the periodic table that is written below the main portion. It’s principle energy levels are indicated by n-2. Unit 4 Notes: Chapter 5 Electrons in Atoms Page |4 A summary of terms: - Principle Energy level- How far an electron is from the o The value for the outermost s and p electrons is found by looking at the period number of the element o The value for the outermost d electrons is found by looking at the period number and subtracting 1. o The value for the outermost f electrons is found by looking at the period number and subtracting 2. - Sublevel- Indicates the in which the electrons are found- includes s, p, d, and f o The value for the outermost electrons if found by looking at the “block” of the periodic table - Number of electrons in the outermost sublevel o Found by looking at the really small number above the column in which the element is found. - Put the principle energy level, sublevel, and number of electrons together like this: Section 5.2- Electron Configurations Goals for this section: - Write electron configurations for various elements using the orbital diagram method, the aufbau chart method and the periodic table method - Explain why electrons fill orbitals the way they do - List and explain the 3 rules for writing electron configurations - Explain why some elements have electron configurations different from those predicted by the aufbau principle - Identify the number of valence electrons in elements using the electron configurations - Know these terms: electron configuration, aufbau principle, Pauli exclusion principle, Hund’s rule I. Stable Atoms A. Change generally proceeds toward the B. High energy systems are unstable and to become more stable 1) (This is what you see when the crushed wintergreen mints emit light or when the samples of gases glowed- electrons are falling from a high energy level to a lower energy level) C. Electrons are arranged with possible level This arrangement is called the electron configuration II. Electron Configurations A. The ways in which electrons are arranged in various around the nuclei of atoms are called electron configurations. B. Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell you how to find the electron configurations of atoms. C. Aufbau Principle 1) According to the aufbau principle, electrons occupy the orbitals of first. D. According to the Pauli exclusion principle, an atomic orbital may hold at most Unit 4 Notes: Chapter 5 Electrons in Atoms Page |5 1) To occupy the same orbital, two electrons must have 2) The opposite spins are indicated by pointing in 3) Spinning electrons produce , which then allow the electrons to to each other E. Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible. 1) In other words, electrons fill orbitals and have spins. 2) After all orbitals in a sublevel have one electron, added electrons double up in orbitals and have opposite spins to electrons already there. 3) Example: F. How to write electron configurations without making an orbital diagram: a. Look at the periodic table to find the correct number of i. (look at the atomic number to find protons and for neutral atoms, the number of electrons matches the number of protons) b. Use the Aufbau chart to figure out where the electrons go. Start with and follow the arrows. c. The numbers at the top of the chart indicate the of electrons that can be placed in that sublevel. d. Fill up each sublevel until you have reached the correct number of electrons, which you found earlier on the periodic table Unit 4 Notes: Chapter 5 Electrons in Atoms Page |6 Aufbau Diagram Check your work: Use the periodic table to help you check your work. The last electrons written in your electron configuration should be the same as the principle energy level, sublevel, and number of electrons in the outermost energy level as indicated on your periodic table! G. Valence Electrons electrons that determine the of elements 1) Since they are in the outer shell, you can find out how many are present by counting the electrons in the 2) Oxygen example: H. Shorthand electron configurations 1) find the element you want on the 2) look up and find the in that row 3) Write the noble gas chemical symbol in brackets 4) Look back at the periodic table and find the information above the element you’re working with 5) Fill in the values for n and n-1 by looking at the row the element is in. 6) Make sure you pay attention to the presence of possible d and f orbitals if your element is in the p block. 7) You’ll need to look at the aufbau chart for this 8) Example of Shorthand notation: a) I. Remember to use your periodic table to help you find the electron configuration for the electrons after the noble gas. Exceptions to the Aufbau principle Unit 4 Notes: Chapter 5 Electrons in Atoms Page |7 1) According to the aufbau chart, copper should have this electron configuration: 2) Instead, it has this electron configuration: 3) Exceptions occur in groups 6B and 1B because: a. Sublevels are most stable when they are b. Sublevels are fairly stable when they are c. Sublevels lack stability when they are d. The d sublevel becomes more stable in groups 6B and 1B by stealing an electron from the Section 5.3- Physics and the Quantum Mechanical Model Goals for this section: - Describe the relationship between wavelength and frequency of light - Identify the source of atomic emission spectra - Explain how the frequencies of light are related to changes in electron energies - Distinguish between quantum mechanics and classical physics - Know these terms: amplitude, wavelength, frequency, hertz, electromagnetic radiation, spectrum, ground state, photons, Heisenberg uncertainty principle I. Physics and the Quantum Mechanical Model A. Electrons and light a. Electrons and light both exhibit at the same time! b. The “particle” of light is known as a B. What is a wave? a. Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays. b. All electromagnetic waves travel in a vacuum at a speed of c. The wavelength and frequency of light are to each other C. Terms associated with a wave: 1) The amplitude of a wave is the wave’s from zero to the crest. 2) The wavelength, represented by (the Greek letter lambda), is the distance between the . 3) The frequency, represented by (the Greek letter nu), is the number of to pass a given point per unit of time. a) The SI unit of cycles per second is called a hertz (Hz). A hertz = 1/second 4) Equation and value of constant: 5) Example problem: D. Separating Light a. A prism separates light into the colors it contains. b. When white light passes through a prism, it produces a rainbow of colors. c. The rainbow is called a Unit 4 Notes: Chapter 5 Electrons in Atoms Page |8 E. Atomic emission spectrum : the formed when atoms absorb energy, forcing electrons into energy levels. These electrons the lose energy by when they return to energy levels. a) The light is made up of only a few specific , depending on the element b) Each frequency is a different c) The light is emitted as electrons fall from one to another, like from n=4 to n=1 d) They are like atomic fingerprints- every element is unique 1) An Explanation of Atomic Spectra a) When the electron has its , the atom is in its ground state. b) Excitation of the electron by raises the atom from the ground state to an excited state. (not the orbital predicted by aufbau chart) c) A quantum of energy in the form of light is emitted when the electron drops back to a lower energy level. a. This energy is directly proportional to frequency, which determines the light’s color. d) Light emitted by an electron moving from a higher to lower energy level has a frequency to the energy change of the electron 1. Equation describing energy change of the electron 2. E = h x ν (h = 6.626 x 10-34J*s) 3. So different energy level drops result in different frequencies (and colors) of light F. Why does it matter that an electron behaves as both particle and wave? 1) The fact that electrons behave as waves leads to some odd observations, like: 2) Heisenberg’s uncertainty principle- it is impossible to know exactly both the of a particle at the same time. a. This limitation is critical in dealing with small particles such as electrons. i. Just the act of observing an electron changes what that electron does! b. This limitation does not matter for ordinary-sized object such as cars or airplanes. 3) When photons hit metals, they can result in the - This is the photoelectric effect a. We use this property to create electron microscopes, which allow for clearer images since electrons have than light Unit 4 Notes: Chapter 5 Electrons in Atoms Page |9 Extra Into: Quantum Numbers: • Principal Quantum Numbers Represented by: n Main E level of electron # E sublevels = # of main level N cannot be zero and must be a positive integer • l Quantum number indicating orbital or sublevel – angular quantum number l is any integer between 0 and n-1 So for n = 2, l can be 0, 1 l = 0 means s l = 1 means p l = 2 means d l = 3 means f Magnetic quantum number m can be any integer from - l to + l So if l = 0, m must be 0 (indicates s with 1 orbital) l = 1, m must be -1, 0, or 1 (indicates p with 3 orbitals) l = 2, m must be -2, -1, 0, 1, 2 (indicates d with 5 orbitals) l = 3, m must be -3, -2, -1, 0, 1, 2, 3 (indicates f with 7 orbitals) Spin quantum number- s- indicates angluar momentum of electrons Is either +1/2 or -1/2 Written as (n, l, m, s) How do orbitals relate to quantum numbers: http://en.wikipedia.org/wiki/Atomic_orbital Principle Possible Orbitals Magnetic Spins Quantum quantum # (n) angular present quantum numbers quantum numbers written numbers (l) (n, l, m, s) 1 2 3 4 Unit 4 Notes: Chapter 5 Electrons in Atoms P a g e | 10