Atomic Structure. Part 3: Wave-Mechanical Model of the Atom. Key Question: How does the wave mechanical model explain the location of electrons?

Transcription

1 Name Chemistry Atomic Structure Essential Question: How was the structure of the atom determined? Vocabulary: bright-line spectrum electron configuration excited state ground state orbital wave-mechanical model Part 3: Wave-Mechanical Model of the Atom Key Question: How does the wave mechanical model explain the location of electrons? 1. Dalton s Atomic Theory (1803): Modifications to Dalton s Theory: JJ Thomson (1897); model based on his discovery of the electron Ernest Rutherford (1911): 1. Atom is mostly 2. At center of atom is a dense positively charged 3. Niels Bohr and other physicists studied light given off by elements with a Why were they studying light? They could apply principles about the nature of light to the nature of subatomic particles. What is light? 1. Newton (1600 s): light travels through space as a 2. Maxwell (1864): light travels through space as a beam of 4. Electromagnetic Spectrum: (wavelengths in meters) to 7 x Gamma X rays UV Visible light IR Micro Radio waves Wavelength increases FM short AM Frequency decreases Energy decreases Speed is constant = 3.00 X 10 8 m/sec

2 5. Properties of Electromagnetic Waves: Wavelength ( ): distance between two in a wave Frequency: of waves that pass a given point per second Speed: measured in per second (3.00 x 10 8 m/s) 6. Bohr: From his observations about light given off by elements Knew that the color of light emitted by a star or element indicated how much was being released. Thought that as an atom gained energy from heat, electricity, etc., that it was the that gained the energy. Electrons normally in the, close to the nucleus Electrons become when energy is added and they move away from the nucleus Excited state is : electrons return to the ground state releasing energy as,, Bright-line shows energy released from the electrons as they return to the ground state 7. Bohr Experiment with the Bright-line Spectrum of Hydrogen

3 Line Spectra and the Energy Level Diagram 1. In this activity we will investigate the spectrum of hydrogen just as Niels Bohr did in Remember that energy may be added to electrons in the form of heat, light, or electricity. As this energy is added, the electrons are promoted to higher, or excited, energy states. As the electrons lose this energy, light is emitted. The wavelengths of the light emitted enable us to determine the energy levels that the electrons occupied. 2. This emitted light may be measured using a spectroscope or diffraction film. From the wavelength, the ENERGY associated with each color may be determined using the following formula: E KJ = 119,557 This equation comes from work done by Max Planck and Louis de Broglie who were investigating the wave-like properties of matter. Energy is measured in kilojoules. 3. Your teacher will place a tube filled with hydrogen gas in a high voltage apparatus. As electricity flows through the hydrogen, observe the spectrum seen using the diffraction film in the glasses. For each color observed, record both the color and the wavelength in the table below. The wavelength Using the equation above, calculate the energy values and record in the table. Approximate Color Wavelength Energy (kilojoules) 4. The energy transitions you have seen occurred when the electrons fell back to the second energy level. Transitions back to the first level were not seen since they are in the ultraviolet region. 5. Use these values to construct the energy level diagram for hydrogen on the back of this page. For this diagram you will need to know at what energy the highest energy level is located. This can be found by using Table S to look up the first ionization energy for hydrogen. This energy is the amount of energy needed to strip an electron from the atom of hydrogen and therefore corresponds to the highest possible energy level. First Ionization energy for hydrogen: 6. After completing the plotting of the energies you observed, draw arrows to correspond to the electrons returning to the second energy level.

4 E N E R G Y (kj/ mol) Questions to answer: a. Which color of light contains the greatest amount of energy? b. The least amount of energy? c. To which energy level did all the excited electrons return to? d. An electron absorbs a specific amount of energy, called a, when it jumps to a higher energy level. e. Each color represents a specific amount of energy released by an electron and is called a. f. Each line on the diagram represents a different level in the hydrogen atom.

5 Bohr Model of the Atom 1. Parts of Bohr s model: 1.) Electrons orbit the nucleus in orbits at fixed distance called 2.) Electrons farther from the nucleus contain (more, less) energy 3.) When electrons gain energy, they move from the to higher energy levels 4.) Electrons return to the ground state by releasing energy in specific amounts called 2. Problems with the Bohr Model 1.) It did not account for all the in the spectra of the other elements. 2.) With better instruments, the model did not even work for. 3. Heisenberg Uncertainty Principle (1927): it is not possible to know the exact and of a subatomic particle. 4. Wave Mechanical or Quantum Model: based on the of finding an electron in a given region of space around the nucleus. Proposed by Erwin Schrodinger and Louis DeBroglie and uses Bohr s as distances from the nucleus. 1.) Electrons do not move in paths around the nucleus. 2.) Electrons farther from the nucleus have energy. 3.) An is a region in space with a high probability of finding an electron. 5. Demonstration- Penny Toss!

6 Wave Mechanical, or Quantum, Model of the Atom 1. Principal Energy Level (Bohr) = Principal Quantum Number, n I N C R n = 4 E A n = 3 S I N G n = 2 E N E R G Y n = 1 2. Energy Sublevels: found within the Principal Energy Level (account for the extra lines in the spectra of other elements) 1.) Number of Sublevels = Principal Quantum Number, n n = 1, sublevel n = 2, sublevels n = 3, sublevels 2.) Energy of each sublevel is slightly. 3.) Each electron in a given sublevel has amount of energy. 4.) Sublevel symbols are s, p, d, f Draw the sublevels in the diagram above.

7 3. Orbitals: regions with a sublevel where electrons may be found 1.) Each orbital can hold or electrons. 2.) Number of orbitals in each sublevel s sublevel = orbital p sublevel = orbitals d sublevel = orbitals f sublevel = orbitals 3.) Number of orbitals in a Principal Energy Level = n 2. Use on the diagram on the front to draw the orbitals in an energy level. 4. Electron Spin: electron rotates on its axis like the 1.) Opposite spins: and 2.) Pauli Exclusion Principle: only electrons can occupy the same orbital, each with opposite spin 5. Shapes of Orbitals: 1.) s orbital: shape; higher Principal Energy Levels has diameter. See p. 131 in textbook. Diagram below. 2.) p orbital: shape around the nucleus and each axis. Diagram below. 6. Hund Rule: must place 1 electron in each p orbital before a second electron can go in. Think of the Birthday Rule!!

8 7. Fill in orbital structures for the atoms below. Use these three rules: Aufbau principle: electrons occupy the orbitals of energy first. Pauli exclusion principle: an atomic orbital may describe at most electrons. Hund s rule: one electron enters each orbital until all orbitals contain one electron with spin Element Atomic 1s 2s 2p 3s 3p 4s Number H Li B C O F Ne Na Al S Ar K Ca

9 8. Summary of Wave Mechanical Model Principal Quantum Number n Number of Orbitals n 2 Sublevel s p d f Number of Electrons 2n 2 9. Electron Configurations: shorthand method for showing electron location in an atom. Write the electron configurations for the atoms listed in #7 on the previous page. H Li B C O F Ne Na Al S Ar K Ca 10. Write the electron configuration for an atom of copper. Cu

10 Look at the Periodic Table. What does it give for copper s electron configuration? Orbital Notation Practice and Questions 1. Complete the orbital notation for the elements below. Atomic Number 1s 2s 2p 3s 3p 3d 4s 4p H Li Ne Si Fe Br 1 st 2 nd 3 rd 4 th Principal Energy Levels 2. Complete the following table (Note PQL = Principal Quantum Level, n) Atomic Number H= Li= Ne= Si= Fe= Br= Number of Electrons Number Occupied Orbitals Number Filled Orbitals Number of Unpaired Electrons Number of Electron Pairs Number of Occupied Sublevels Number of Filled Sublevels Number of Occupied PQL s Number of Filled PQL s PQL of Highest Occupied PQL Number of Valence Electrons (outermost energy level) Number of Kernel Electrons (non-valence electrons) Number of Unpaired Valence Electrons

11 1. Periodic Table Relationships Significance of Wave Mechanical Model 1.) Period (row) corresponds to the Level being filled with electrons example: For Period 3, which Principal Energy Level is filling? 2.) Electron configurations given in block for each element in corner. You need to know the filling order!! 3.) Blocks of elements correspond to the sublevel being filled s block: Groups and filling p block: Groups,,,,, filling d block: Groups through filling (includes the elements) f block: and Series filling 2. Belated filling: when an sublevel fills before a lower p sublevel. Example: 4s fills before the 3. Valence Electrons: electrons in the Principal Energy Level for an element Valence Shell: corresponds to the ; determines the properties of an atom Filled Valence Shell: Group elements called the or Gases 8 electrons in valence shell indicates chemical. Group: elements in the same vertical column with the number of valence electrons 4. Kernel: nucleus and all electrons. 5. Excited State Configurations: look for electrons not in lowest possible energy levels Atomic Ground State Excited State Number oxygen 1s 2 2s 2 2p 4 1s 2 2s 2 2p 3 3s Ions: atoms with or electrons Atomic Ground State Ion Number oxygen 1s 2 2s 2 2p 4 1s 2 2s 2 2p 6 2-6