Chapter 4_ Arrangement of Electrons in Atoms

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Chapter 4: Arrangement of Electrons in Atoms Properties of Light: Light acts as a wave when propagating through space and as a particle when interacting with matter. The Wave Nature of Light: (fig. 2 pg. 98) 1. Properties of a wave a. wavelength ( ) – distance between two identical points on a wave (or distance between crests or troughs). This is measure in units of _________ (m, nm, etc). b. frequency ( ) – how many wave cycles (the distance of one wavelength) pass a point in a certain amount of time. This is measured in cycles per second called hertz ( or 1/s). c. Amplitude ( )-the displacement of the wave from the equilibrium position (how far from zero). 2. The wavelength and the frequency of a wave are inversely related (meaning as the wavelength gets longer the frequency is _____________). The product of the two is equal to the wave speed. Wave speed (m/s) = wavelength (m) . frequency (Hz or 1/s) 3. Light is a member of a family of waves called the electromagnetic spectrum for all these waves are partially electrical and partially magnetic (perpendicular to an electrical field is a magnetic field). These waves need no _____________ in which to travel (sound requires a medium to travel and is called a mechanical wave). See fig. 1 pg. 98. (You must memorize the relative positions of each of the bands of the electromagnetic spectrum and know that visible light lies between 400 and 700 nm). Draw the electromagnetic spectrum below.

4. Visible, ultraviolet light, and x-rays can be produced from electrons jumping from higher to lower _________ levels. Elements give off a unique color when __________ . The heat causes electrons to jump up to higher energy levels. Later when the electrons jump _________ to lower energy levels, they give off light. Only certain colors are given off for a given element (Bohr studied hydrogen) and because of this Bohr stated that energy levels are _____________ (meaning can be one value or another, nothing in between). From this, the Bohr planetary model of the atom was created which will be discussed in more detail later. 5. Electromagnetic waves all travel the speed of light (c = 3.00.108 m/s). c = λν λ = wavelength (m) ν = frequency (hz) Practice problem: a. Determine the wavelength of the light emitted by a sodium vapor lamp if the frequency of the radiation is 5.10.1014 hz. In what region of the electromagnetic spectrum does this light exist?

b. Determine the frequency of light whose wavelength is 0.500 nm. In what region of the electromagnetic spectrum is this radiation located?

Homework: pg. 124: 1; 4; 6; 10

1.What is the wavelength of light with a frequency of 6.65.10 Hz? 2.What is the frequency of light if its wavelength is 695 nm? bonus: pg. 124: 13 (c = d / t)

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Planck’s Hypothesis: 1. German scientist Max Planck stated that energy (light) instead of being radiated (given off) continuously, was given off in packets or bundles of energy called ___________ . 2. Light particles contain specific amounts of energy (light particles are called photons). 3. Planck came up with the idea of energy being quantized by the analysis of a heated piece of _____. As the iron is heated it changes in _______ from black to red, to yellow, to white to blue. He stated that energy changed in specific amounts and was therefore _________ (a wild idea at the time for it clashed completely with classical physics). (Bohr used Planck’s ideas then to create his planetary model of the atom where electrons were quantized in orbits of fixed energy. That led to the modern quantum mechanical model and the idea of orbitals). 4. Planck found a direct relationship between energy and frequency of a photon (as frequency of a radiant particle increases, the energy of the photon _______________). E=hν E = energy (J) Since ν = c / λ h = Planck’s constant (6.63.10-34 Js) E = hc/λ (this is a derived equation and will not be given on the test)

5. Since x-rays have ________ frequencies, they have ________ energies. Since radio waves have _________ frequencies, they have _________ energies. Since gamma rays have _________ wavelength, they have _____________ energies (so energy and wavelength have an _____________ relationship with each other). Practice Problems: a. Calculate the energy of a photon whose frequency is 5.00.1015 hz.

b. Calculate the energy of a photon whose wavelength is 400. nm.

Homework: pg. 124: 5 (only speak of 1 experiment); 11

1. How much energy is in a photon of light with wavelength 580 nm? 14 2. How much energy is in a photon of light with a frequency of 5.47.10 Hz? 3. What is the frequency of light with energy of 4.05 x 10-19J? 4. What is the frequency of light with wavelength of 423 nm?
5. Draw a diagram of 2 waves of different frequency and label the amplitude and wavelength on each. Label one as high frequency and one as low frequency. 6. At what speed do all electromagnetic waves travel? 7. What scientist determined the relationship of frequency to energy of electromagnetic radiation? 8. Infrared radiation, uv radiation from the sun, a green traffic light, the signal from a radio tower, dental x-rays, microwaves. a. arrange the radiation shown above in order of increasing wavelength. b. Arrange the radiation shown above in order of increasing frequency. c. Which order (a or b) will be correct for increasing energy? Bonus: (must be solved as a factor-label conversion problem for credit) Determine the amount of time in minutes that will be required for light to travel from Earth to Mars? (the distance from Earth to Mars is 1.29.105 miles)

Photoelectric Effect and Atomic Spectra of Elements The Photoelectric Effect: (see pg. 99-100) 1. The particle nature of light was proposed by ___________ in the 1600’s, while evidence for its wave nature was shown by ____________. The debate as to whether light behaved as a particle at all raged on until the 1900’s when Albert Einstein showed that although light behaves as a _________ when traveling through space, it behaves as a ______________ when interacting with matter. 2. Einstein showed that light behaved as a particle when interacting with matter through his explanation of the photoelectric effect. The photoelectric effect is the ejecting of ___ from an active metal’s surface by shining light on the surface of the metal. It was observed that very ________ red light would not eject electrons from the metal surface but very ____ yellow light would eject electrons from an active metal’s surface. This phenomenon could not be explained by wave theory which states the brighter the light is, the greater the wave energy would be. 3. Einstein, using Planck’s hypothesis, reasoned that each photon of the _________ frequency red light had a smaller energy than each photon of energy from the yellow light (like an ant can’t kick a football off of a tee (red light) but a football kicker can (yellow light)). 4. The frequency of light that would just suffice to give enough energy to _______ the electron is called the threshold frequency. a. The brighter the light shown on the metal surface (above the threshold frequency), the _______ electrons that will be ejected (like lots of football kickers each with their own football on a tee). b. The greater the frequency of the light (or shorter wavelength), above the threshold frequency, the ____________ (more kinetic energy) the electron will be ejected from the metal surface (like an elementary football kicker (low frequency above the threshold) compared to an NFL kicker (higher frequency above threshold frequency)). c. Applications of Photoelectric effect: (electric eyes and photocells) Pictures of Photoelectric Effect bright red light dim yellow light

bright yellow light

dim blue light

(no e- ejected)

(e- ejected)

(more e- ejected)

(faster e- ejected)

Atomic Spectra of Elements (absorption and emission spectroscopy): Absorption Spectroscopy (see fig. 8a pg. 102) 1. When light is shown on a piece of material (like a shirt), some of the wavelengths may be __________. The absorbed wavelengths provide an energy ________ to a difference in energy between energy levels of the atom absorbing the light (E = hν, c = λν). What light you see emanating from the shirt is all the light that was not absorbed (the absorbed light warms up the shirt). 2. When an electron absorbs a light photon it ________ into a higher energy level (here the electrons goes from a ground state to an excited state). 3. Since elements have different arrangements of electrons, they will have different energy values between energy levels and thus different wavelengths will be ____________, so different elements are different colors. Draw a picture of what is happening on an atomic scale in absorption spectroscopy.

Emission Spectroscopy: (fig. 5; 7 pg. 101) 1. In emission spectroscopy, gaseous atoms are placed into an __________ state by being heated (electricity works as well). Nanoseconds after being excited, the electron _________ back into the ground state where it will emit light of an energy equal to the difference in energy between the energy levels (this difference in energy will determine the wavelength of the light (E = hν, c = λν) 2. The light can be separated through the use of a _________ or diffraction grating. Draw a picture of what is happening on the atomic scale for emission spectroscopy.

3. Robert Bunsen observed that placing different salts in a flame produced different colors, now referred to the _________ test. 4. An element’s absorption or emission spectrum is like a _________________ that can be used to identify the element. Absorption spectroscopy applications: Colors you observe from materials, determination of concentrations of solutions, identification of elements. Emission spectroscopy applications: Lights from incandescent bulbs, light from stars, determination of the elemental composition of stars (or heated gaseous elements) from their spectral patterns, fireworks, LASERS, neon signs. Homework: 1. If the energy required to eject an electron from a photoactive metal is 3.10.10-19 J, determine the threshold frequency (E = h ν) for this metal. 2 a. Calculate if 600. nm light will eject an electron from the metal from problem #1? b. Will very dim 600. nm light eject an electron, from this metal? c. Will very bright light with a frequency of 4.50.1014 hz will eject the metal’s electron? d. What would be different of an ejected electron if 100 nm light were shown on the metal? 3. What scientist presented experimental evidence of light’s wave nature? 4. What scientist presented experimental evidence of light’s particle nature? 5. Give 2 applications of the photoelectric effect. 6. Give 2 applications of absorption and emission spectroscopy. 7 a. When looking at a yellow shirt through a spectroscope (diffraction grating), only the colors green and red are observed, what has become of the blue light? b. Draw a picture on the atomic scale to show what has become of the blue light (label the ground state and excited state atoms). a. When viewing a heated tube of an elemental gas through a spectroscope several bright lines are observed. What is the source of the bright lines? b. Draw a picture on the atomic scale to show the source of one of the lines (label the ground state and excited state atoms).

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1. 2. 3. 4. 5. 6. 7.

8. 9. 10.

11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23.

24. 25. 26. 27.

Spectrum of Kodachromium http://jersey.uoregon.edu/vlab/elements/Elements.html Kodachromium is an element whose first 7 energy levels contain the following energy values. n = 1 (0.000.10-19 J) n = 2 (2.880 . 10-19 J) n = 3 (3.616 . 10-19 J) n = 4 (3.968 . 10-19 J) n = 5 (4.368 . 10-19 J) n = 6 (4.960 . 10-19 J) n = 7 (5.088 . 10-19 J) Part 1 1. Show all possible transitions from higher energy levels to the n=1 energy level. 2. Calculate the wavelength for each of the transitions. On the picture show the waves being emitted as colors (if they are in the visible spectrum). If they are not in the visible spectrum label what portion of the Electromagnetic spectrum they exist. You can use black for portions outside the visible spectrum. 3. Draw an emission spectrum for the element (see pg. 97). Draw the spectrum between the wavelengths of 700-400 nm. 4. Draw an absorption spectrum for the element (see pg. 95) using the same guidelines as the emission spectrum.

Wavelengths of Colors: Violet 400-440 nm Blue 440-480 nm Green 480-530 nm Ultraviolet light – below 400 nm

yellow 530-590 nm orange 590-630 nm red 630-700 nm Infrared – above 700 nm

Calculations: (work down not across!!) Transition: n=7 to n=1

Transition: n=6 to n=1

Transition: n=5 to n=1

Transition: n=4 to n=1

Transition: n=3 to n=1

Transition: n=2 to n=1

Emission Spectrum

700nm

650nm

600nm

550nm

500nm

450nm

400nm

Absorption Spectrum

700nm

650nm

600nm

550nm

500nm

450nm

400nm

Flame Test Lab
Metal Ion Sodium Na+ Potassium K+ Strontium Sr+2 Lithium Li+ Copper Cu+2 Barium Ba+2 Unknown A Unknown B Unknown C Flame Color Orange Orange, red Orange, Red Magenta Green Yellow-Green

Name ___________________________
Color of spectral lines 590 nm 404 nm, 691 nm, 694 nm 422 nm, 496 nm, 641 nm, 650. nm 610 nm, 671 nm 406 nm, 465 nm, 515 nm, 522 nm 500. nm; 505 nm; 535 nm; 630. nm

Number of spectral lines 1 3 4 2 4 4

Unknown A ________

Unknown B _________

Unknown C ________

Results and Conclusions: 1. . Which element releases light with the longest wavelength? ___________ a. Calculate the frequency. (show calculation below) b. Calculate the energy. (show calculation below)

2. Which element releases light with the highest frequency? a. Calculate the frequency. (show calculation below)

b. Calculate the energy. (show calculation below)

3. What kind of relationship exists between: (direct or inverse) i. wavelength and frequency ii. frequency and energy iii. energy and wavelength ______________________ ______________________ ______________________

2. In terms of the atom, what causes the production of several spectral lines for some elements?

Outline of Chapter 4 Part I Test The Wave Nature of Light: 1. Properties of a wave a. wavelength (λ) – distance between two identical points on a wave (or distance between crests or troughs). This is measure in units of distance (m, nm, angstrom etc). b. frequency (ν) – how many wave cycles (the distance of one wavelength) pass a point in a certain amount of time. This is measured in cycles per second called hertz (s-1). c. Amplitude (A)-the displacement of the wave from the equilibrium position (how far from zero). The wavelength and the frequency of a wave are inversely related. c=λν 2. Light is a member of a family of waves called the electromagnetic spectrum for all these waves are partially electrical and partially magnetic (perpendicular to an electrical field is a magnetic field). These waves need no medium in which to travel (sound requires a medium to travel and is called a mechanical wave) and all travel the same speed (c = 3.00.108 m/s). (You must memorize the relative positions of each of the bands of the electromagnetic spectrum and know that visible light lies between 400 and 700 nm).
104 wavelength (m) 102 1m 10-2 10-4 10-6 10-8 10-10 10-12 10-14

infrared

uv

x-ray

gamma

visible
frequency (Hz) 104 106 108 1010 1012 1014 1016 1018 1020 1022

Light: A historical perspective: 1. In the late 1600’s Isaac Newton reasoned that light consisted of particles, while one of his contemporaries (Christian Huygens) reasoned that light behaved as waves, for it was proven by him to have wave properties. Later, Einstein showed that light had a particle nature (now accepted that light acts as a particle and a wave) through his explanation of the photoelectric effect. So it was shown that light can act as a wave when traversing through space or a particle when interacting with matter. 2. In the photoelectric effect, light at or above a threshold frequency will provide the work function (Energy) required to eject an electron from the surface of a photoactive metal. A frequency of light above the threshold frequency will eject electrons of a certain kinetic energy. The higher the frequency the light (above the threshold frequency), the faster the electron will be ejected. The brighter the light (above the threshold frequency) shining on the photoactive metal, the more electrons that will be ejected from the photoactive metal. Applications: Electric Eyes and photocells. Planck’s Hypothesis: 1. German scientist Max Planck stated that energy (light) instead of being radiated (given off) in a continuous fashion, instead energy is given off in packets or bundles of energy called quanta. 2. Quanta of radiant energy are referred to as photons (or light particles). 3. Planck came up with the idea of energy being quantized by the analysis of a heated piece of iron. As the iron is heated it changes in color from black to red, to yellow, to white to blue. He stated that energy changed in discrete units and was therefore quantized (a wild idea at the time for it clashed completely with classical physics). (Bohr used Planck’s ideas then to create his planetary model of the atom where electrons were quantized in orbits of fixed energy. That led to the modern quantum mechanical model and the idea of orbitals). Planck found a direct relationship between energy and frequency of a photon. E = h ν since c = λν E = hc/ λ (Planck’s constant (h = 6.63.10-34 J s)). 4. Visible, ultraviolet light, and x-rays can be produced from electrons jumping from higher to lower energy levels. Elements give off a unique color when heated (as observed for heating salts in a flame). The heat causes electrons to jump up to higher energy levels. Later when the electrons jump down to lower energy levels, they give off light whose energy is equal to the difference in energy between the 2 energy levels. Only certain colors are given off for a given element (Bohr studied hydrogen) and because of this Bohr stated that energy levels are quantized (and thus the Bohr planetary model of the atom was created which will be discussed in more detail later). Bohr Model- atom is mostly empty space but electrons surround nucleus in orbits of fixed energy (like planets around the sun). When light is emitted an electron jumps from an excited state to a ground state. When light is absorbed an electron jumps from a ground state to an excited state. An emission spectrum is a series of bright colored lines on a dark background. The colored lines are the result of energy being added to the atom to make electrons jump to high energy levels, soon after they relax to lower energy levels. The color observed has energy equal to the difference in energy between the 2 energy levels.

An absorption spectrum is a series of black lines on a continuous (ROYGBIV) color spectrum. The source of the black lines is that line of all colors (ROYGBIV) being exposed to the atom. Any colors whose energy matches the difference in energy between 2 energy levels is absorbed (removed) to cause electrons to jump to higher energy levels. The remaining light passes through the atom. -

e

e

e
Absorbed λ

-

e-

Ground State Excited State This is a picture of light being absorbed

Unabsorbed light passes through atom

Ground State Excited State (heat or electricity excites the electron) This is a picture of light being emitted

Absorption spectroscopy applications: Colors you observe from materials, determination of concentrations of solutions, identification of elements. Emission spectroscopy applications: Lights from incandescent bulbs, light from stars, determination of the elemental composition of stars (or heated gaseous elements) from their spectral patterns, fireworks, LASERS, neon signs.

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