CRHS Academic Chemistry Unit 4 Electrons. Notes. Key Dates

Transcription

1 Name Period CRHS Academic Chemistry Unit 4 Electrons Notes Key Dates Quiz Date Exam Date Lab Dates Notes, Homework, Exam Reviews and Their KEYS located on CRHS Academic Chemistry Website:

2 Page 2 of 20 Unit 6 Notes 4.1 WAVE NATURE OF LIGHT Before 1900, scientist thought that light behaved solely as a. This belief changed when it was later discovered that light also has particle-like characteristics leading to the modern Quantum Atomic Theory. Still, many of light s properties can be described in terms of waves. What follows is a quick review of the wavelike properties of light. Electromagnetic Radiation that exhibits behavior as it travels through space. Examples: *Light is forms of electromagnetic radiation (below). Visible light is the very narrow part of the electromagnetic spectrum we can see. All of these types of light travel in a form. Wave Part LOCATION Highest point on a wave Lowest point on a wave Midline of a wave

3 Unit 6 Notes Page 3 of 20 Wavelength on a wave, the distance between two or. Symbol is (lambda) Units distance ( or ) λ Frequency - # of to pass a given point in. Symbol is Units - ( ) f Amplitude wave s from to origin or to trough Units - displacement ( or ) Increasing frequency

4 Page 4 of 20 Unit 6 Notes All forms of light travel at the same, 3.00 x 10 8 m/s (denoted by c). This can be calculated by the below equation, which is true for forms of light. Wavelength (λ) and frequency (f)are related As wavelength () frequency (f) As wavelength () frequency (f) Example: What is the frequency of green light, which has a wavelength of 4.90 x 10 7 m? Practice: 1. Your lunch table engages in a 15 minute argument about the virtues of popular music. As a group, you conclude that the greatest top 40 pop radio station in all the land broadcasts their music at a frequency of x 10 6 Hz. What is the wavelength of the radio station broadcasting this music? 2. Our atmosphere allows in enough ultraviolet radiation from our sun to suitably burn the skin of most humans, given enough time and exposure. Of particular danger to us is ultraviolet B, even though it makes up less than 5% of all UV radiation reaching the earth. Ultraviolet B has a wavelength of 3.24 x 10 7 m. What is the frequency of UVB?

5 Unit 6 Notes Page 5 of PARTICLE NATURE OF LIGHT In the early 1900 s scientist conducted two experiments involving interactions of light and matter that could not be explained by the wave theory of light. One experiment involved a phenomena known as the. This refers to the emission of electrons from a metal when light shines on the metal. At that time, the mystery of the photoelectric effect was why no electrons are emitted if the light s frequency is below a certain frequency regardless of the intensity. Wave theory of light predicted that light of any frequency could supply enough energy to eject an electron. The Concept of the Quantum The explanation of the photoelectric effect dates back to 1900, when German physicist Max Planck was studying the emission of light by hot objects. As a result of that work, Plank proposed that a objects can or energy only in small distinct units called (singular is quantum) the energy of the wave (photon) is calculated by multiplying the frequency by Planck s constant (below) energy is proportional to frequency of light Max Planck So... as frequency (f), energy (E). Given by the Planck-Einstein formula: E = (units: Joules) h = Planck s constant: 6.63 x J-s f = of radiation (units: Hertz) Examples: 1. Calculate the energy of a wave with a frequency of 2.11 x Hz.

6 Page 6 of 20 Unit 6 Notes 2. Find the energy of a wave with a wavelength of 7.34 x 10-6 m. Practice: 1. What is the energy of a wave with a frequency of 3.45 x Hz? 2. In your position as physicist and lead astronomer at the McDonald Observatory in the Davis Mountains of west Texas, you spend a lot of time evaluating emanations from space. One day you discover a millisecond event in which an unknown energy source emits energy at 2.28 x 10 9 J. What is the frequency of this light? Light as a Particle In 1905, Albert Einstein, referencing the photoelectric effect, proposed the radical idea that light also behaves as a (which he called a ). Albert Einstein Each photon is a particle of electromagnetic energy having zero mass and carrying a of energy. Einstein won a Nobel prize for explaining the photoelectric effect by proposing that electromagnetic radiation is absorbed by matter only in of photons. For an electron to be ejected from a metal surface, the electron must be struck by a single possessing the minimum energy required. The energy of a photon depends on the frequency of the radiation or E photon = h x f

7 Unit 6 Notes Page 7 of 20 SO, light behaves both as a wave and a particle Properties of a the entire electromagnetic spectrum, light can be bent and redirected or refracted Properties of a quantum theory of the atom, the photoelectric effect All atoms emit light when excited by electricity or heat electrons at their lowest possible energy level are in the state electrons that are at one or more quanta above their ground state are in the state When an electron loses energy to return to the ground state, a of light is emitted. The energy of the photon that is emitted is determined by the size of the energy gap between the excited state energy level and the ground state. Practice: Draw your own e jumping to a higher energy level, losing energy, and falling back to its ground state. Photon of visible colored light

8 Page 8 of 20 Unit 6 Notes What is an Atomic Emission Spectrum? Passing electric current through the gas or vapor of a pure substance causes the substance to. In this process, electrons in atoms of the substance are raised to a number of possible higher states, but quickly return to the single state. If we pass this light through a, we get a distinct color pattern to that substance. All elements emit light in unique, distinct patterns of visible light called. The spectrum corresponds to in the energy of the different excited states and energy value of the single ground state. Spectra can be used to an element.

9 Unit 6 Notes Page 9 of ELECTRON CLOUD: QUANTUM MODEL The explanation for atomic emission spectra lead to an entirely new atomic theory called. This breakthrough was significant because it predicted the location and properties of electrons in atoms with more than one electron. Bohr s planetary model could only predict the atomic spectrum of hydrogen. Key attributes of the Atomic Quantum theory 1. Describes the - properties of electrons and other small particles. 2. Heisenberg determined that it is to determine simultaneously determine both the position and the velocity of an electron (or any other particle). 3. As a result of the, the solutions to the Schrodinger equations called wave functions only give a of finding an electron at a given place around the nucleous. 4. These regions are called. 5. It takes four to completely describe the properties of atomic orbitals and the properties of the electrons in those orbitals. The Four Quantum Numbers: 1. Principal Quantum Number (n) a. Identifies the of the orbital where the electron is located. b. There are 7 possible primary energy levels to an atom. c. Primary energy levels roughly correspond to on the periodic table. d. As n increases the electron is located from the nucleus.

10 Page 10 of 20 Unit 6 Notes 2. Angular Momentum Quantum Number (l) a. The angular momentum sublevel, symbolized by l, indicates the shape of the orbital. b. Orbitals of different shape within the same primary energy levels are called. c. Orbital sublevels are assigned a letter l Letter 0 s 1 p 2 d 3 f d. Orbital sublevels correspond to a on the periodic table S D F P

11 e. The number of sublevels that can be occupied are shown below. Unit 6 Notes Page 11 of 20 Energy Level Sublevels Occupied 1 spdf 2 spdf 3 spdf 4 spdf 5 spdf 6 spdf 7 spdf Note: The only reason that 6f, 7d, and 7f sublevels are not occupied is lack of large elements to fill those orbitals. 3. Magnetic Quantum Number (m) a. The magnetic quantum number signifies the orientation of an atomic orbital around the nucleous b. One orbital holds electrons c. The four sublevels have different shaped orbitals and different numbers of orbitals per sublevel. SUBLEVEL Shape # of orbitals Max #of e s 1 p 3 d 5 f 7 Shape of orbitals

12 Page 12 of 20 Unit 6 Notes Summary Drawing of Atomic Orbitals Magnetic Quantum Number (m) Sublevels (l) m = 0 m = -1, 0, 1 m = -2, -1, 0, 1, 2 m = -3, -2, -1, 0, 1, 2, 3 4. Spin Quantum Number (+1/2, -1/2) The of an electron. Electrons in same orbital must have spins clockwise () counterclockwise ()

13 Unit 6 Notes Page 13 of ELECTRON CONFIGURATION Q: What is Electron Configuration? A: A way to express complete electron around an atom. These configurations are the - configurations for each element. Example: Oxygen s Electron Configuration: 1s 2 2s 2 2p 4 How is it written? Hydrogen = 1s 1 large number is the the letter is (section of periodic table) the superscript is of electrons in the sublevel (each square on periodic table = electron) Process to write the electron configuration of any atom: Step 1. Determine total number of electrons in atom Step 2. Identify the energy levels of the orbitals using the Periodic Table as a guide and reading it like a book. ALL sublevels are written in the order filled: left to right, top to bottom on periodic table (

14 Page 14 of 20 Unit 6 Notes Step 3. Add electron to the orbitals, one by one, according to the Aufbau Principal. Aufbau Principle an electron occupies the energy level possible. Each orbital holds up to electrons, so sublevels can hold: sublevel # orbitals Max electrons s 1 p 3 d 5 f 7 Note: See graphic on page 19 that illustrates the relative energies of the sublevels. Practice: List the sublevel from lowest energy to highest energy and include the maximum number of electrons in the superscript. Practice: 1. Name the elements whose electron configurations are: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 9 1s 2 2s 2 2p 6 3s 2 3p 6 2. Write electron configurations for these elements: Potassium - Copper - Bromine -

15 Unit 6 Notes Page 15 of 20 How do you check your answer? Option 1 Total of superscripts equals which equals the atomic number in an atom. Sodium = 1s 2 2s 2 2p 6 3s1 ( = 11) Atomic # = 11 Option 2 The last segment of an electron configuration identifies location of element on the periodic table. 3s 1 energy level sublevel electron sodium! What is Noble Gas Configuration? Answer: A shorter way to write electron configuration in which a noble gas symbol substitutes for a portion of the configuration Write the electron configuration for Neon. Write the electron configuration for Sodium. Above, notice what is the same and what is different by putting brackets around what is the same. To write noble gas configuration 1. Write the noble gas that the element in question 2. Write the electron configuration that the noble gas to reach the element in question. Example Magnesium Practice Bromine

16 Page 16 of 20 Unit 6 Notes 4.5 ORBITAL NOTATION Q: What is Orbital Notation? A: A representation of electron configuration that illustrates all four quantum numbers for each electron. Example: Spin (-1/2, +1/2) Carbon - 1s 2s 2p Primary Energy Level (n) SubLevel (l) Magnetic Quantum Number (m) In addition to the Aufbau Principal covered previously, two other rules are applied when writing for any atom in the ground state. Hund s Rule Orbitals of are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin. Remember Pauli Exclusion Principal No two electrons in an atom can have the same set of four. This means that electrons in the same orbital must have opposite spin states. An electron occupies the lowest energy level possible. See graphic on page 19 that illustrates orbital energies Orbitals of same energy levels are each occupied by e before any one orbital is occupied by a e. Electrons in singly-occupied orbitals must have the spin. Electrons in the same orbital must have opposite spin. Process to write orbital notation 1. Use to represent orbitals. 2. Use to represent electron spin. 3. Group sublevel orbitals together.

17 Unit 6 Notes Page 17 of 20 Practice: Fill in missing arrows Element # of e Orbital Notation Electron Configuration Li 3 1s 2s Be 4 1s 2s B 5 1s 2s 2p N 7 1s 2s 2p Ne 10 1s 2s 2p Na 11 1s 2s 2p 3s 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 2p 1 1s 2 2s 2 2p 3 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 1 PRACTICE: Write orbital notation for these elements: Potassium Copper Selenium

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19 Unit 6 Notes Page 19 of 20 The following is a graphic representation of the energy levels of the electron orbitals. Increasing Energy! Remember, lowest energy orbitals filled first The following schematic is a graphic tool guides you on the order to fill electron orbitals. It is consistent with the energy diagram above.

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