Atomic Emission Spectra. and. Flame Tests. Burlingame High School Chemistry

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1 Atomic Structure

2 Atomic Emission Spectra and Flame Tests

3 Flame Tests Sodium potassium lithium When electrons are excited they bump up to a higher energy level. As they bounce back down they release energy in the form of light.

4 Bohr Model In a Bohr atom, the electron is a particle that travels in specific, fixed orbits, but never in the space between orbits. This arrangement expresses energy quantization, and accounts for atomic emission spectra.

5 Energy absorption and light emission in a Bohr atom When an atom absorbs energy, an electron is excited to a higher energy orbit. The electron then transitions back to a lower energy orbit and emits a photon of light.

6 Flame tests reveal the number of levels an electron travels and allows us to predict the amount of energy released. Each color represents a different amount of energy

7 Electromagnetic Spectrum

8 Electromagnetic Spectrum Electromagnetic radiation has a dual behavior. It has properties of a particle called a photon and as a wave traveling at the speed of light. Characterized by a wavelength and frequency.

9 Rainbows from white light When a narrow beam of ordinary white light is passed through a glass prism, different wavelengths travel through the glass at different rates and appear as different colors A similar effect occurs when light passes through water droplets in the air, forming a rainbow.

10 Electromagnetic Radiation Wavelength-λ The distance between two corresponding points on a wave. λ

11 Electromagnetic Radiation Frequency-ν The number of wave crests passing a given point per unit time.

12 Electromagnetic Radiation c = the speed of light = λν c = 3.00 x 10 8 m/s

13 Frequency and Wavelength long wavelength = small frequency Short wavelength = high frequency Qui ckti me and a Graphi cs decompressor are needed to see thi s pi cture. See Screen 7.4 increasing frequency increasing wavelength

14 Bohr Model and Atomic Spectra Niels Bohr ( ) Bohr s s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the SHARP LINE SPECTRA of excited atoms.

15 Bohr Model of Atom An atom has a number of definite and discrete energy levels in which an electron can exist. Electrons can move from one energy level to another. Electron moves in circular orbit.

16 The Bohr Model and Quantized Energy Orbiting Electrons e - e - + Energies are quantized in other words the Energies are limited to discrete values.

17 The Bohr Model and Quantized Engery Input Photon with Energy=hv + e - e -

18 The Bohr Model and Quantized Engery Excited State Electron + e - High Energy Orbit Low Energy Orbit e - Ground State Energy

19 The Bohr Model and Quantized Engery Output Photon with Energy=hv + e - e -

20 Line Spectra of Excited Atoms Excited atoms emit light of only certain wavelengths The wavelengths of emitted light depend on the element Each element gives a different flame color and has a different spectral fingerprint

21 SPECTRAL LINES Each line corresponds to one exact amount of energy being emitted. The emission spectrum of each element is unique to that element

22 Bright Line Spectra High E Short λ High ν Low E Long λ Low ν Visible lines in Hydrogen atom spectrum

23 Atomic Spectra Continuous spectra from sun contain all wavelengths. Line spectra have discrete lines from atoms.

24 Absorption Atomic Spectra The process where an electron moves from a lower to higher energy state, resulting with the gain of energy. From the ground state to an excited state. Emission The process where an electron drops from a higher to lower energy state, resulting with the loss of energy (visible as light) From an excited state back down to the ground state.

25 36Kr = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3d 10 4p p Electron Configuration 10Ne = 1s 2 2s 2 2p 6

26 Electron Configurations The way in which electrons are distributed among the various orbitals is called the electron configuration. Orbitals are filled in order of increasing energy, with no more than two electrons per orbital

27 Aufbau Principle The electron enters the lowest-energy atomic orbital available

28 Pauli Exclusion Principle Only two electrons per orbital, each with opposite spins

29 Hund s Rule When several orbitals of equal energy are available, electrons enter singly with parallel spins. Only after all orbitals in a sublevel are halffilled do electrons pair up.

30

31

32 Paired vs. Unpaired Electrons Electrons having opposite spins are said to be "paired" electrons, as with the electrons occupying the Li 1s orbital. Likewise, the single electron in the 2s orbital (for Li) is said to be "unpaired"

33 Shapes of the s orbitals Representations of (a) 1s, (b) 2s, and (c) 3s orbitals. Cutaway views of these spherical orbitals are shown on the top, with the probability of finding an electron represented by the density of the shading.

34 Shapes of the p-orbitals Each p orbital has two lobes of high electron probability separated by a nodal plane passing through the nucleus. The different shadings of the lobes reflect different algebraic signs analogous to the different phases of a wave.

35 Shapes of the d orbitals Four of the orbitals are shaped like a cloverleaf and the fifth is shaped like an elongated dumbbell inside a donut.

36 Energy Levels The order of filling is determined by the energy of each orbital.

37 Shells are organized into subshells

38 Energy levels for Electrons

39 Energy Level Diagram Notice how the energy levels overlap when put in order of increasing energy. Electrons must fill the lowest energy level first.

40

41

42 p-orbital { Periodic Table d-orbital IA VIIIA 1 2 H He IIA IIIA IVA VA VIA VIIA Li Be B C N O F Ne Na Mg Al Si P S Cl Ar IIIB IVB VB VIB VIIB VIIIB IB IIB K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe {f-orbital Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn (209) (210) (222) Fr Ra Ac (223) Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu (145) Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr (244) (243) (247) (247) (251) (252) Burlingame (257) (258) High (259) School (260) Chemistry { {s-orbital

43

44

45 Order of Filling 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 8s

46 Orbital Shapes s orbital p orbital d orbital

47 Orbital Box Diagram

48 Lithium 3s 2s 3p 2p Atomic number = 3 1s 2 2s 1 3 total electrons 1s

49 Beryllium 3s 2s 3p 2p Atomic number = 4 1s 2 2s 2 4 total electrons 1s

50 Boron 3s 3p Atomic number = 5 1s 2 2s 2 2p 1 5 total electrons 2s 2p 1s

51 Carbon 3s 2s 1s 3p 2p Atomic number = 6 1s 2 2s 2 2p 2 6 total electrons Here we see for the first time HUND S S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible.

52 Nitrogen 3s 2s 3p 2p Atomic number = 7 1s 2 2s 2 2p 3 7 total electrons 1s

53 Oxygen 3s 2s 3p 2p Atomic number = 8 1s 2 2s 2 2p 4 8 total electrons 1s

54 Fluorine 3s 3p Atomic number = 9 9 total electrons 1s 2 2s 2 2p 5 2s 2p 1s

55 Neon 3s 2s 1s 3p 2p Atomic number = 10 1s 2 2s 2 2p 6 10 total electrons Note that we have reached the end of the 2nd period, and the 2nd shell is full!

56 Sodium Atomic number = electrons 1s 2 2s 2 2p 6 3s 1 3s 3p 2s 2p 1s

57

58 Electronic Configuration Exceptions [Ar]4s 2 3d 4 [Ar]4s 3d 5

59 Electronic Configuration of The noble gases are chemically stable as individual atoms and have a full complement of outer groups s and p electrons. 2He = 1s 2 10Ne = 1s 2 2s 2 2p 6 Nobel Gases 18Ar = 1s 2 2s 2 2p 6 3s 2 3p 6 36Kr = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 54Xe = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6

60 Nobel Gases Complete energy levels Neither gain nor lose electrons easily Do not form compounds readily Also called inert gases

61 Noble Gas Notation Chlorine has 17 electrons Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 Cl: [Ne] 3s 2 3p 5 Titanium has 22 electrons Ti: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 Ti: [Ar] 4s 2 3d 2

62 Aluminum Atomic number = electrons 1s 2 2s 2 2p 6 3s 2 3p 1 Or [Ne]] 3s 2 3p 1 3s 2s 3p 2p 1s

63 Phosphorus Atomic number = electrons 1s 2 2s 2 2p 6 3s 2 3p 3 or [Ne]] 3s 2 3p 3 3s 2s 3p 2p 1s

64 Calcium Atomic number = electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 [Ar]] 4s 2

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