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Chapter 4

Chapter 4. Arrangement of Electrons on Atoms. Development of Atomic Model. Properties of light Electromagnetic radiation Energy which travels through space as waves Speed is 3.0 x10 8 m/s Electromagnetic spectrum Gamma rays X-rays Ultra-violet rays Visible lights Infrared rays

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Chapter 4

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  1. Chapter 4 Arrangement of Electrons on Atoms

  2. Development of Atomic Model • Properties of light • Electromagnetic radiation • Energy which travels through space as waves • Speed is 3.0 x108 m/s • Electromagnetic spectrum • Gamma rays • X-rays • Ultra-violet rays • Visible lights • Infrared rays • Microwaves • Radio waves

  3. Wave length – the distance between corresponding points on adjacent waves • Frequency (V) • The number of waves that pass a point per second • Units are waves/second (often written as 1/s or s-1 ) • 1 wave /second = 1 hertz (Hz) • Speed = wavelength x frequency • C = V C = 3.0 x 108 m/s

  4. Find the frequency of radiation if the wavelength is 5.0 x 10-8 m

  5. Photoelectric Effect • Electrons are emitted when light strikes the surface of certain metals • The light must be above a minimum frequency for this to occur • Max Planck (1900) suggested that an object emits energy in small specific amounts called quanta • Quantum – the minimum energy that can be lost or gained by an atom

  6. E = energy (J), H = 6.626 x 10-34Js , V = frequency • E = h v • Albert Einstein (1906) said that light has properties of both waves and particles (photons) • Photons – a quantum of light energy • In order for the electron to be given off from the metal, the electron must be struck by a photon having enough energy to knock the electron loose

  7. Find the energy of radiation if the frequency is 6.15 x 1012 s-1

  8. Hydrogen Atom Emission Spectrum • Ground state – the lowest energy state of an atom • Excited state - a higher energy state; when an atom gains energy • When an atom returns to the ground state, it gives off energy which may include visible light • When the light emitted from hydrogen is passed through a prism a series of lines of lights of different wavelengths are seen

  9. When an atom falls from the excited state to a lower state a photon of radiation is emitted; the energy of the photon is the difference in the energy between the initial state and final state.

  10. Bohr Model of the Hydrogen Atom (1913) • The electron circles the nucleus in a definite path (orbit) • Electron can move to a higher orbit by gaining a certain amount of energy • When the electron drops down to a lower orbit, a photon is emitted; the energy of a photon is equal to the difference in energy between two orbits

  11. Quantum Model of the Atom • Electrons as waves • Louis de Broglie (1924) proposed that electrons have wave-like properties • Heisenberg Uncertainty Principal – it is impossible to determine both the position and path of the electron at the same time • Electrons are detected by their interaction with photons • Because electrons are so small, an attempt to locate an electron with a photon knocks the electron off its course

  12. Schrödinger Wave Equation • Erwin Schrodinger (1927) developed a mathematical equation that treated electrons as waves • Electrons do not travel in exact orbits; instead they exist in certain regions of space called orbitals

  13. Atomic Orbitals and Quantum Numbers • A set of four quantum numbers describes the properties of an atomic orbital. • Quantum Number Describes • Principal (n) Energy level • Angular momentum(ℓ) Shape of orbital • Magnetic (m) orientation in space • Spin spin of electron

  14. (n) (l) (n2) (2n2) Energy Level Sublevels #of obitals #of electron 1 s 1 2 2 s 1 2 p 3 6 3 s 1 2 p 3 6 d 5 10 4 s 1 2 p 3 6 d 5 10 f 7 14

  15. Electron Configuration • Electron configuration –the arrangement of the electrons in an atom • Rules for Electron configuration 1. Aufbau Principle – an electron occupies the lowest orbital possible 2. Pauli Exclusion Principal – an orbital may contain 2 electrons which have opposite spin 3. Hund’s Rule – one electron enters each orbital in a sublevel before a second electron may enter any of them

  16. Electron Configuration Notation • 1s2 means there are 2 electrons in the 1s sublevel • Use the periodic table as a guide when writing electron configurations.

  17. Examples • H – 1s1 • He – 1s2 • Li – 1s22s1 • O – 1s2 2s2 2p4 • Na – 1s22s22p63s1 • Ca – 1s22s22p63s23p64s2 • Sc –1s22s22p63s23p64s23d1 • Br – [Ar] 3d104s24p5

  18. Orbital Diagrams • 1s • H ( ) 2s • Li ( ) ( ) 2p • B ( ) ( ) ( ) ( ) ( ) • N ( ) ( ) ( ) ( ) ( )

  19. Abbreviated Electron Configurations (Noble Gas Notation) • Show the Symbol for the previous noble gas and the electron config. which follows • Br [Ar] 4s23d104p5 • Sb [Kr] 5s24d105p3 • Exceptions to the pattern for the electron arrangement • 4s 3d • Cr [Ar] ( ) ( ) ( ) ( ) ( ) ( ) • Cu [Ar] ( ) ( ) ( ) ( ) ( ) ( )

  20. Abbreviated Electron Configurations (Noble Gas Notation) • Sublevels which are filled or half filled are more stabled then other configurations.

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