1. Modern Model of the Atom The emission of light is fundamentally related to the behavior of electrons.

2. Hydrogen-Atom Line-Emission Spectrum Hydrogen-Atom Line-Emission Spectrum Objective: Discuss the significance of the line-emission spectrum of hydrogen to the development of the atomic model and how electrons exist in specific (quantized) energy levels. Ground State  The lowest energy state of an atom Excited state  State in which an atom has higher potential energy than in the ground state. Line-Emission Spectrum  Light is given off by excited atoms as they return to lower energy states.  Light is given off in very definite wavelengths.

3. Classical verses Quantum Theory Classical theory predicted that the hydrogen atoms would be excited by whatever amount of energy was added and thus observe the emission of a continuous range of frequencies of EM radiation, or a continuous spectrum would appear. However, lines of particular colors appeared giving a line-emission spectrum.  Definite frequency and definite wavelength. E2E2 E1E1 E photon = E 2 – E 1 = h ν When an excited atom with energy E 2 falls back to energy E 1, it releases a photon that has a quantized amount of energy therefore a definite frequency.

4. The Bohr Model of the Atom The Bohr Model of the Atom Objective: Compare and contrast the Bohr model and the quantum model of the atom. Electron Orbits, or Energy Levels  Electrons can only be in certain discrete orbits, and that they absorb or emit energy in discrete amounts as they move from one orbit to another.  Each orbit thus corresponds to a definite energy level for the electron.

5. Pair and Share Distinguish between ground state and an excited state of an atom. Distinguish between continuous and line- emission spectrum. Explain the meaning of quantized energy levels in an atom and show how these levels relate to the discrete lines in the spectrum of that atom.

6. Quantum Model of the Atom Quantum Model of the Atom Objective: Discuss Louis de Broglie’s role in the development of the quantum model of the atom. Electrons as Waves Electrons as Waves  Louis de Broglie suggested that matter in motion has properties that are normally associated with waves.  The wave properties are especially applicable to very small particles, such as electrons.  Each particle’s wavelength is related to its mass, its velocity and Planck’s constant.  Smaller the mass, and greater the velocity, the more wavelike the characteristics.

7. Heisenberg Uncertainty Principle Heisenberg Uncertainty Principle Objective: Explain the significance of the Heisenberg Uncertainty Principle on the development of the modern atomic model. States that it is impossible to determine simultaneously both the position and velocity of an electron.  Electrons are detected by their interaction with photons.  Because photons have about the same energy as electrons, any attempt to locate a specific electron with a photon knocks the electron off its course.

8. Schrödinger Wave Equation Schrödinger Wave Equation Objective: Discuss the impact of the Schrödinger Wave Equation on the development of the modern atomic model and the location of electrons around the atom. Predicts a three dimensional region around the nucleus called an atomic orbital that describes the electron’s probable location. The boundary of the atomic orbital is defined as the volume that encloses a 90% probability of containing its electrons.

9. Atomic Orbitals An atomic orbital is a region of space in which the probability of finding an electron is high. Electrons have a designated arrangement in all atoms. The atomic orbitals have specific energy levels and shapes in which the electrons are distributed. These are referred to as the following:  Principal energy levels  Energy Sublevels  Orbitals  Number of electrons in an orbital (electron spin).

10. Principal Energy Level Principal Energy Level Objective: Identify the principle energy levels in an atom and state the energy trend among them. The principal energy level, symbolized by n, indicates the main energy level occupied by the electron.  Values of n are positive integers only: n = 1,2,3,…..  Increases in energy and size as one moves away from nucleus.  Also related to the periods of the periodic table.  The number of orbitals possible per energy level is equal to n 2 n=1 n=2 n=3 n=4 n=5 EnergyEnergy

11. Energy Sublevels Energy Sublevels Objective: For each principle energy level, state the number of sublevels, identify them, and state the energy trend among them. Indicates the shape of the orbital n Number of orbital shapes allowed in an energy level equals n.  Shapes of the first four orbitals are designated s, p, d, f

12. Electron Orbitals Electron Orbitals Objective: Sketch the general shapes of the s and p orbitals. State the number of orbitals in each sublevel and the distribution of electrons in s, p, d, and f sublevels. Each sublevel has a different number of orbitals.  s – one possible orbital in space  p – three possible orbitals in space  d – five possible orbitals in space  f – seven possible orbitals in space

13. Electron Spin A single orbital can contain only two electrons, which must have opposite spins.

14. Shapes of Orbitals

15. Shapes Cont.

16. Modern Atom Model Chart Organizer Modern Atom Model Chart Organizer Objective: Relate the number of sublevels corresponding to each of an atom’s main energy levels, the number of orbitals per sublevel, the number of electrons per sublevel and main energy level. Principal n Energy Level (n) Sublevels in main energy level n (n sublevels) Number of orbitals per sublevel Number of electrons per sublevel n Number of electrons per principal energy level (2n 2 ) 1s122 2spsp 1313 2626 8 3spdspd 135135 2 6 10 18 4spdfspdf 13571357 2 6 10 14 32