Electrons! Chapter 5

Electrons! Chapter 5

I.Light & Quantized Energy A.Background 1. Rutherford s nuclear model: nucleus surrounded by fast-moving electrons; no info on how electrons move, how they re arranged, or differences in chemical behavior a. Example: Cl, Ar, K are #s 17, 18, 19 -- consecutive atomic numbers, but very different properties -- why? 2. Early 1900s: scientists noticed that some elements produced light when placed in a flame a. Analyzing the specific light emitted showed that chemical behavior was related to electron arrangement

B. Light: Wave Nature 1. Wavelength (λ): distance between two neighboring crests or troughs a. unit: usually nanometers (1 x 10-9 m = 1 nm) 2. Frequency (ν): number of waves (measured in # of crests) which pass a given point each second a. unit: hertz (1 Hz = 1 wave/second), kilohertz, megahertz b. in mathematical calculations, use s -1 (1 s -1 = 1 Hz) 3. Amplitude: height of wave from origin to crest or origin to trough 4. Speed of light (c) = 3.00 x 10 8 m/s c = λν ^see p. 119

i. Example: A certain red light has a wavelength of 650 nm. What is its frequency? (BRINGIN BACK G.U.E.S.S) G: wavelength = 650 nm, c = 3.00 x 10 8 m/s, 1 Hz = 1 s -1 U: frequency (in Hz) E: c = λν can change to ν = c/λ S: 3.00 x 10 8 m 1 nm 1 1 Hz 1 s 1 x 10-9 m 650 nm 1 s -1 S: 4.62 x 10 14 Hz

5. Electromagnetic (EM) radiation: form of energy released by EM processes that acts like a wave as it travels through space; consists of synchronized electric & magnetic fields a. Electromagnetic (EM) spectrum: spectrum of all EM radiation resulting from fluctuations or vibrations of charged particles (p. 120) i. Includes radio waves, microwaves, radar waves, infrared light, visible light (aka continuous spectrum), ultraviolet, X-rays, and gamma rays ii. Each part of the continuous spectrum corresponds to a specific wavelength & frequency, which are refracted at different angles -- produces colors (ROY G BIV)

Practice Problems: p. 121, #1-4

C. Light: Particle Nature 1. Wave model of light does not describe all of its interactions with matter, like iron emitting different wavelengths of light (colors) at different temperatures -- so how can these interactions be explained? 2. Max Planck (German physicist, 1890s) wanted to figure this out! a. Studied light emitted from heated objects b. Found that frequencies & energy changed in tiny, specific amounts, and matter can only gain or lose energy in these small amounts (quanta), NOT in any random amount c. Quantum: minimum amount of energy that can be gained or lost by an atom; plural quanta; also called photons d. All energy, including light, comes in quanta -- this is quantum theory e. Each color of light has a specific sized quanta; red is small, violet is large

3. Planck proposed a formula that expresses the energy of a single quantum (photon) of radiation of any given frequency E = energy (in joules) E = hν h = Planck s constant, 6.626 x 10-34 J s ν = frequency in s -1 a. Example: how much energy is contained in one quantum of a gamma ray? Frequency of gamma ray is 1 x 10 24 Hz G: h = 6.626 x 10-34 J s, ν = 1 x 10 24 Hz, 1 Hz = 1 s -1 U: energy (in joules) E: E = hν S: 6.626 x 10-34 J s 1 x 10 24 Hz 1 s -1 1 1 1 Hz S: 6.626 x 10-10 J b. Example 2: how much energy is contained in one quantum of a radio wave? Frequency of radio wave is 1 x 10 4 Hz

4. Albert Einstein (German physicist, 1905) studied the photoelectric effect: light of a certain frequency (color) shining upon a material & causing electrons to be discharged from the material a. Theoretically, a very lowfrequency light directed at a material for a long period of time will eventually accumulate & provide enough energy for the material to eject electrons BUT, that didn t happen -- why?

b. Einstein found that when the intensity of light of a particular wavelength was reduced, fewer electrons were ejected, & vice versa i. Based on wave theory, less intense light should cause electrons to be ejected with lower kinetic energy, & vice versa ii. However, electrons emitted at lower intensities still had the same kinetic energy as when the intensity was greater iii.conclusions: Light has both wave & particle properties Light consists of photons: particles of EM radiation carrying a quantum of energy (plain English: light is a stream of tiny bundles of energy) Photons have their own specific energy for a given wavelength; at minimum, one photon possesses the energy necessary to eject an electron from a material

Practice Problems: p. 124, #5-6 Homework: p. 126, #7-12

Homework Answers p. 873 3) 3.61 x 10-19 J 4) 8.48 x 10-16 J p. 124 5a) 4.19 x 10-13 J 6a) gamma ray or X ray 5b) 6.29 x 10-20 J 6b) infrared 5c) 6.96 x 10-18 J 6c) ultraviolet p. 126 7. speed, wavelength, frequency, amplitude; EM waves travel at c 8. Wave model treats light as an EM wave; particle model treats it as comprised of photons. Wave model couldn t explain photoelectric effect, color of hot objects, & emission spectra. 9. A quantum is the minimum amount of energy that can be lost or gained by an atom; matter loses or gains energy in multiples of the quantum 10. Continuous spectrum contains all colors, emission spectra contains only specific colors 11. Einstein proposed that EM radiation has a wave-particle nature, that the energy of a photon depends on the frequency of the radiation, and that a photon s energy is given by E photon = hv 12. a: 3 (radio waves) b: 1 (gamma rays) c: 1 (gamma rays)

D. Light Spectra 1. Light is not always absorbed or emitted by a material continuously, as the wave model would suggest -- instead, it is often emitted in bursts by excited electrons 2. Atomic emission spectrum: frequencies of light emitted by a given material a. When substances are passed through a flame, they glow a particular color (example: copper - green, sodium - yellow) b. When this light is passed through a prism & divided, a series of colored lines appears called spectral lines c. Each element has a unique set of lines it produces, so atomic emission spectra (the sets of lines) can be used to identify elements d. Example: hydrogen s emissions spectrum, p. 126

II. Quantum Theory & the Atom A.Bohr s Model 1. Proposed a quantum atomic model that explained why atomic emission spectra are only made up of certain frequencies of light, not all continuously 2. Key points: a. There are certain definite orbits in which electrons can travel b. An electron in a given orbit has a certain definite amount of energy c. The greater the distance from the nucleus, the greater the energy of the electron d. Possible electron orbits are called energy levels e. An electron can gain energy from various sources, causing it to jump to a higher level f. When an electron drops to a lower energy level, it loses energy in the form of a photon of radiation g. Energy of the photon corresponds to the difference in energy levels h. An electron cannot drop into an energy level which already holds its maximum number of electrons

3. Also mapped out energy levels a. Lowest energy level is called the ground state b. When an atom gains energy, it enters an excited state Bohr described 7 possible excited states/energy levels of hydrogen & assumed it was the same for other atoms c. Atom gains energy electrons jump to higher energy levels d. Atom loses energy in the form of a photon of EM radiation electrons fall back to a lower energy level e. Amount of energy photon carries is determined by the difference in energy levels when an electron falls from an upper level to a lower level ΔE = E higher-energy orbit E lower energy orbit = E photon = hf

B. Quantum-Mechanical Model 1. Bohr s model explained hydrogen well, but not really other elements; primarily his ideas of quantized energy & energy levels were correct & expanded upon by several scientists 2. Louis de Broglie suggested that all moving particles have wave properties a. de Broglie equation: λ = h/mv wavelength of any particle is equal to Planck s constant divided by mass times velocity b. He was correct! 3. Heisenberg s uncertainty principle: there is always some uncertainty as to the location & velocity of an electron it is impossible to know both simultaneously 4. Schrödinger treated electron as a wave & came up with an equation to describe its wave-like behavior confirmed electron has both wave & particle properties 5. Wave model predicts that electrons will be found in an atomic orbital: a 3D region around the nucleus

C. Hydrogen s Atomic Orbitals 1. Principal quantum number (n): indicates sizes & energies of atomic orbitals a. Corresponds to the energy levels suggested by Bohr (1,2,3,4,5,6,7) called principal energy levels b. Electrons may be found in each energy level of an atom (7 shown on the periodic table) c. Largest possible number of electrons possible in any one level is 2n 2 2. Energy sublevels: contained within principal energy levels a. Number of sublevels contained within each principal energy level = n b. All electrons in each energy level do not have the same energy as earlier assumed c. Determines the shape of the electron cloud d. Four types of sublevels are possible, each with a defined shape & number of electrons: s sharp : lowest sublevel p principal : second sublevel d diffuse : third sublevel f fundamental : fourth sublevel 2e - (1 pair) 6e - (3 pair) 10e - (5 pair) 14e - (7 pair)

Copy in notes: Table 5-1 Figure 5-10 a & b