Modern Atomic Theory-Electrons crest trough photons

Transcription

1 Modern Atomic Theory-Electrons I. Waves & Energy A. Atomic model changes 1. History of the model electrons as particles a. John Dalton 1. The atom was described as a solid sphere 2. Based on evidence from combining ratios of elements and compound in reactions b. J.J. Thomson 1. Developed the Plum Pudding model 2. Electrons were distributed in a sphere of + charge c. Ernest Rutherford 1. Developed the Planetary Model 2. Negative electrons travel around a positive nucleus 2. Problems with the Rutherford model a. Opposites attract suggests that electrons should fall into the nucleus and we know that they don t. b. The model suggests that the structure is unstable but it is actually stable. B. Waves- the question was then asked- Could electrons be waves? 1. Properties of waves a. Diagram crest trough b. Terms 1. Frequency ( )- a measure of the number of waves passing a point per unit of time. It is measured in hertz, Hz, which is equal to cycles per second. 2. Wavelength ( ) - the distance between two similar points on two connecting waves. 3. Amplitude- the distance from a crest or trough to the midpoint of the wave c. Waves transmit energy from one particle to the next particle. Boats move up and down and DO NOT travel with the wave. d. As frequency increases wavelength decreases. The reverse is also true. 2. Light waves a. They transmit energy- a beam of light is a stream of photons - a discrete bundle (or quantum) of electromagnetic (or light) energy 1. The amount of energy depends on three things: -The length of time of exposure -The distance over which the light is spread -The frequency of the light -Speed of Light- because air is mostly space the value of 3.0 x 10 8 m/s is light approximate speed So, c = speed of light constant = 3.00 x 10 8 m/s = wavelength = frequency And the equation that relates these three is c = b. Researchers 1. Max Planck -Worked with light emitted by solids -Said that the energy values of atoms varied by small whole numbers 2. Albert Einstein -Worked with light energy -Determined that light was emitted as photons or quanta of energy

2 E=h ; where E= energy in Joules, h= Planck s constant x J s = frequency The energy is directly proportional to the frequency -Said that absorption of photons would explain the photoelectric effect c. Produce a continuous spectrum- Electromagnetic Spectrum -Visible light is just one portion or range of frequencies of the EMS (electromagnetic spectrum) Visible Light d. Problems with the theory that electrons are waves 1. Objects that don t burn, like iron, give off red light yellow white when heated. The wave theory would have it give off ultraviolet 2. Some materials give off electrons when light shines on them. They have a baseline for energy which they do not drop below. This is the photoelectric effect. e. Solution -Louis de Broglie said that particles can have properties of waves -Einstein said that the wave can have properties of particles -Hence we have wave-particle duality!!! II. Hydrogen Atom A. Line emission spectra-bright line spectra 1. Definition: distinct lines of color, each with its own frequency. 2. Each element has its own spectrum 3. Johann Balmer studied hydrogen and developed a mathematical relationship for the progression of the lines. R H=Rydberg Constant B. Energy levels 1. Niels Bohr- explained the spectrum 2. Diagram- In the tube, energy is absorbed as electricity and released as light 3. Photons- are little packets of light energy 4. Bohr model a. Hydrogen atoms exist in specified energy states. b. Hydrogen atoms can absorb only certain amount of energy. c. When excited, hydrogen atoms lose energy in photons. d. Different photons produce different colors in the spectrum. e. Ionization Energy- the energy required to remove one mole of electrons form one mole of atoms

3 C. Modern Model- the Quantum Model of the Atom 1. Louis de Broglie- electrons can have wave properties Diffraction- the bending of a wave as it passes by the edge of an object. Interference- occurs when waves overlap leading to an increase of energy in some places and a decrease of energy in other places. 2. Erwin Schrodinger- used the wave-particle duality idea to develop equations to describe the energy of electrons as waves 3. Their work started wave mechanics which is now called quantum mechanics in the field of physics. It incorporates the particle wave theory and probability. 4. Heisenberg-developed the Uncertainty Principle that says you cannot measure momentum and position of an electron simultaneously. D. Electron Positions 1. Orbitals- highly probable locations for electrons (region in space where an electron may be found. An orbital may hold a maximum of 2 electrons.) 2. Quantum numbers- specify the properties of atomic orbitals and electrons. a. First quantum number -Called the Principle Quantum Number -Represented by the symbol (n) -Indicates the energy level -Ranges from 1 to 7 b. Second quantum number -Called the Angular Momentum Quantum Number -Represented by the symbol (l) -Indicates the shape s, p, d, and f * s-sublevel is spherical, holds 2 electrons, 1 orbital * p- sublevel is dumbbell, hold 6 electrons, 3 orbitals * d- sublevel, holds a max of 10 electrons, 5 orbitals * f- sublevel, hold a max of 14 electrons, 7 orbitals c. Third quantum number -Called the Magnetic Quantum Number -Represented by the symbol (ml) -Gives the orientation along the x, y, & z axis d. Fourth quantum number -Called the Spin Quantum Number (ms) -Represented with +1/2 or -1/2 -Indicates the direction of the spin of the electrons. -*Pauli Exclusion Principle-no two electrons can have the same four quantum numbers. 3. Aufbau Process- When electrons enter orbitals they always fill the lowest energy level first. 4. Hund s Rule- orbitals of equal energy are each occupied by one electron before an orbital is occupied by a second electron & all electrons in singly occupied orbitals must have the same spin.

4 E. Representing Electron Configurations 1. Three Methods a. Orbital Notation -An unoccupied orbital is represented by a line -An arrow up is used to indicate an electron with positive spin -An arrow down is used to indicate an electron with negative spin -The principle quantum number and the sublevel are indicated below the line. -Examples: Hydrogen: _ 1s 2s 2p x 2p y 2p z 3s 3p x 3p y 3p z Helium: Lithium: Beryllium: Boron: b. Electron-Configuration Notation -Eliminates the lines and arrows of orbital notation -The number of electrons in a sublevel is shown by adding a superscript to the sublevel designation - Electron configuration for Krypton: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 -Examples: Hydrogen: Helium: Lithium: Beryllium: Boron: -Electron configuration exceptions In the D-block if you have a d 4 or a d 9 this is an unstable configuration. It requires an electron to move down from a higher energy level in order to make the electron configuration d 5 or d 10. This is done for stability because the electron orbitals are most stable when they are full or half full. The same rule would apply for the F block, too: f 6 or f 13 would need to become f 7 or f 14. Example: Chromium: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 Copper: Samarium: c. Noble Gas Notation -Simplifies the electron-configuration notation of third period atoms -The noble gas that precedes the atom is enclosed in square brackets [ ] -The remaining electron configuration then follows. -Examples: Electron configuration for Phosphorous: 1s 2 2s 2 2p 6 3s 2 3p 3 Noble Gas Notation for Phosphorous: [Ne] 3s 2 3p 3

5 Sodium: Magnesium: Potassium: Calcium: 2. Ground State Electrons fill the orbitals and energy levels as described above (in the normal way). This means that the electrons have their expected (normal) amount of energy. This is the most stable state for the electrons and the lowest electronic state. Example: Sodium in ground state: 1s 2 2s 2 2p 6 3s 1 3. Excited State Any state for an electron that is not at ground state. Electrons have an extra amount of energy and an electron jumps to a higher energy level than normal. Example: Sodium in an excited state: 1s 2 2s 2 2p 5 3s 2 4. Lewis Dot Diagram is a diagram that shows the number of valence electrons in an atom. Steps for drawing the diagram: - Write out the symbol for the element - Determine the number of valence electrons the element has. - Draw the valence electrons as dots around the element symbol. - You must have one dot on each side before you have two. - You can t have more than 2 dots per side. - The maximum number of dots is 8.

6 More About the Electromagnetic Spectrum Electromagnetic radiation is all around you. Your senses can only detect visible light, but there are other, invisible radiations passing through you and around you right now. They all travel at 300 million m/s the speed of light, and all of them can be reflected and refracted by suitable objects. One type is infra-red (IR), which is emitted by any hot object, including you. Electric bar fires produce IR radiation to keep you warm; grills, ovens and cookers produce it to heat food. Whenever an object absorbs any type of radiation, the object is heated, and we use this in infra-red grills. Infra-red is also used by remote controls for TVs and VCRs. Microwaves also heat substances when they are absorbed, but this only works well for certain substances such as water and fats. This is because water molecules and fat molecules absorb microwave frequencies particularly well, and the energy they absorb ends up as kinetic energy when they vibrate in other words, heat. This is why microwaves can be dangerous, because we are mainly made of water and fats. Microwaves can also cause cataracts, a clouding of the cornea in the eye. Microwaves cook things from the inside, whereas normal cooking involves heating the outside of the food. Microwaves are also used for radar, and by mobile phones. Holding a microwave transmitter to the side of your head for long periods can t be a good idea! Microwaves will induce currents in metals, which is why we must not use metal containers to heat food in microwave ovens the metal will get very hot, and may also give off sparks and wreck the oven. Having said that, some foods come in special containers with metal foil pieces especially to make use of this effect and help to cook the food. Can you think of an example? Visible light is, obviously, used for seeing things. We can refract it using lenses, and images in cameras and in our eyes. Light is not only used for seeing but can also be optical fibres, for example, in endoscopes used by doctors to see inside patients' make sent along bodies. When light travels down an optical fibre, all the light may stay inside the fibre until it emerges from the other end. This is because light travels down optical fibres by repeated total internal reflection. More information can be carried using optical fibres than by sending electrical signals through cables of the same diameter. There is also less weakening of the signal in optical fibres. Radio waves are used to transmit radio and TV programmes between different points on the Earth's surface. Longer wavelength radio waves are reflected from an electrically charged layer in the Earth's upper atmosphere, called the ionosphere. This enables them to be sent between distant points despite the curvature of the Earth's surface. Ultraviolet radiation is used in sunbeds. Special coatings that absorb ultraviolet radiation and emit the energy as light are used in fluorescent lamps and security coding. Ultra violet radiation can pass through skin to deeper tissues. The darker the skin, the more ultra violet it absorbs and the less reaches into deeper tissues; X-radiation is used to produce shadow pictures of materials which X- rays do not easily pass through including bones and metals. Gamma radiation is used to kill harmful bacteria in food, sterilise surgical instruments, and kill cancer cells. X- radiation and gamma radiation mostly pass through soft tissues, but some is absorbed by the cells. High doses of ultra violet radiation, X-radiation and gamma radiation can kill normal cells. Lower doses of these types of radiation can cause normal cells to become cancerous. 1. Name all the sources of infra-red radiation mentioned on this sheet. Write down at least two other sources of infra-red radiation 2. Infra-red radiation is absorbed by our skin and felt as heat. Use the ideas of molecules vibrating and of absorption to explain how this happens. 3. Infra-red has a higher frequency than microwave radiation, so it transfers more energy. Yet microwaves cook food faster why should this be? 4. Why should we not use metal containers in microwave ovens? 5. What are the effects of different doses of UV, X-rays and gamma rays on living cells?

7 The Electromagnetic Spectrum The electromagnetic spectrum is a continuum of all electromagnetic waves arranged according to frequency and wavelength. The sun, earth, and other bodies radiate electromagnetic energy of varying wavelengths. Electromagnetic energy passes through space at the speed of light in the form of sinusoidal waves. The wavelength is the distance from wavecrest to wavecrest (see figure below). Light is a particular type of electromagnetic radiation that can be seen and sensed by the human eye, but this energy exists at a wide range of wavelengths. The micron is the basic unit for measuring the wavelength of electomagnetic waves. The spectrum of waves is divided into sections based on wavelength. The shortest waves are gamma rays, which have wavelengths of 10e-6 microns or less. The longest waves are radio waves, which have wavelengths of many kilometers. The range of visible consists of the narrow portion of the spectrum, from 0.4 microns (blue) to 0.7 microns (red).

8 Name: Date: Pd.: Conceptual Physics Video: Light & Color Directions: As you watch the video fill-in the worksheet below: 1. Shaking a charged rubber rod produces a changing field, which in turn induces a changing field, which produces electromagnetic. 2. is electromagnetic waves generated by vibrating electrons. 3. The lowest frequency of visible light is. 4. In order form lowest to highest frequency the colors of visible light are. 5. Beyond red is red. 6. Beyond violet is violet. 7. X rays are frequency electromagnetic radiation. 8. Does light speed up or slow down when passing through a transparent material? 9. The brightest color of light from the sun is The visible spectrum can be divided into three regions, which we see as, &. 11. When red, green, and blue light overlap they produce. 12. Seawater is a greenish-blue light overlap they produce. 13. White light minus red light equals. 14. The blueness of the sky is explained by the of light from the sun by the atmosphere. 15. At sunset, the sun appears orange because frequencies of light are scattered by the atmosphere. 16. Why are clouds white? 17. A flash is sometimes seen just as the sun sets because the atmosphere acts as a prism. 18. Distant, dark objects such as mountains look while distant light objects look. 19. Next time question: Why do we see the sky s blueness, but astronauts do not? (HINT: What is the background for each situation?)

9 Name: Date: Pd.: Part A: Use the words in the box to fill in the blanks. Worksheet 1-Light & Color Before you can see an object, it must light. A material through which nearly all light passes is. A material that you cannot see through clearly is. An object cannot be seen through. light is a mixture of all visible wavelengths of the spectrum. objects absorb all colors and reflect little light. Red, blue, and green are the three colors of light. They can be mixed to produce any color. The retina contains tiny objects called cells that detect certain wavelengths of light. When the brain responds to these signals, we see. One way of producing color is by the use of a, a transparent object that some colors and allows others to pass through. The color of the filter is the same as the color of the light it. Part B: Use the words in the box to fill in the blanks. A colored material that absorbs certain colors and reflects others is a. To mix and make any color, it is necessary to have only three primary pigment colors-magenta, yellow, &. Light color is determined by the wavelength of the light from pigment particles. translucent primary opaque black transmits colors white cone absorbs transparent reflect filter cyan filter pigment additive subtractive black reflected Because primary light colors combine to produce white light, they are called colors. If all primary pigments are added equally, the result will be. Because black results from the absence of reflected light, the primary pigment colors are called colors.

10 Part C: Choose the best answer. 1. The transfer of energy by electromagnetic waves is called a. modulation b. radiation 2. Infrared radiation has a wavelength slightly longer than a. microwaves b. visible light 3. can be used for cooking. a. microwaves b. ultraviolet radiation 4. radiation has a higher frequency than visible light. a. ultraviolet b. infrared 5. Photons are tiny bundles of radiation that have no a. mass b. energy 6. Objects containing heat can emit a. X-rays b. infrared radiation 7. have the lowest photon energy. a. Radio waves b. Gamma rays 8. Radio waves are radiation with very long and very low frequencies. a. wavelengths b. photons 9. have the highest frequency of all electromagnetic waves. a. X-rays b. Gamma rays 10. Electromagnetic waves are classified according to their wavelength s on the electromagnetic. a. photon b. spectrum 11. can be used to check for broken bones. a. X-rays b. Infrared rays 12. Ozone in Earth s atmosphere blocks most of the sun s. a. ultraviolet rays b. infrared rays 13. Visible radiation is the only part of the electromagnetic spectrum you can. a. see b. feel 14. Radio waves can be changed by a process called. a. radiation b. modulation 15. are radio waves with the highest frequency and energy. a. gamma rays b. microwaves

11 Name: Date: Pd.: Worksheet 2-EMS 1. What were the contributions of each of the following people in the study of atomic structure and electrons? a. John Dalton: b. JJ Thomson: c. Rutherford: d. Niels Bohr: e. Heisenberg: f. Pauli: g. Plank: h. Einstein: 2. Define the following: a. Frequency: b. Wavelength: c. Amplitude: d. Duality: e. Photon or Quanta: f. Bright line spectrum: g. Orbital: h. Photoelectric effect:

12 i. Ground State: j. Excited state: 3. What are the four quantum numbers and what information does each provide? 4. What two things are incorporated in to quantum mechanics? 5. Between what two types of light does visible light appear on the electromagnetic spectrum? 6. What do the following statements mean? a. Energy is directly proportional to frequency. b. Frequency is inversely proportional to wavelength. 7. What are the colors of the continuous spectrum listed from lowest energy to highest energy? Highest E-,,,,,, -Lowest E 8. What three things affect the amount of energy transmitted by waves? 9. Work the following problems showing all set ups and answers with correct units. a. What is the energy of a wave with a frequency of 3.47 x Hz? b. What is the frequency of a green light wave with a wavelength of 4.34 x 10-7 meters?

13 Name: Date: Pd.: Worksheet 3- Photons & Electromagnetic Radiation 1. Explain the mathematical relationships in the following equations Use terminology such as directly proportional, inversely proportional, constant, & variable in your explanation and be sure to include proof by illustrating the relationship with examples and/or equations Wave Description of Light Particle Description of Light c = E=h c = 3.0 x 10 8 m/s E = energy of a photon = wavelength = frequency = frequency h = x J s 2. Determine the frequency of light with a wavelength of x 10-7 cm. 3. Determine the energy in joules of a photon whose frequency is 3.55 x Hz. 4. Using two equations E=h & c =, derive an equation expressing E in terms of h, c, and. 5. Given the speed of light as 3.0 x 10 8 m/s, calculate the wavelength of the electromagnetic radiation whose frequency is x Hz. 6. What is the frequency of a radio wave with an energy of 1.55 x J/photon?

14 Name: Date: Pd.: Worksheet 4- Electron Configuration 1. Sodium-23 has an atomic number of 11 and a mass number of Therefore, there are _(1)_ protons and _(2)_ neutrons in the nucleus of sodium. 3. There are electrons in a neutral sodium atom. 4. Electrons located in energy levels near the nucleus possess (more, less) energy than electrons located further away from the nucleus. 5. The probable location of an electron in an energy level is known as a. 6. How many electrons may occupy the same space orbital at any one time? 7. In order to occupy the same space orbital, two electrons must have (same,opposite) spins For each type of sublevel, furnish the desired information. SUBLEVEL NUMBER OF SPACE ORBITALS MAXIMUM NUMBER OF ELECTRONS s p d f 16. What is the shape of the s orbital? 17. What is the shape of the p orbital? 18. The arrangement of the electrons in the atom may be represented by the electron? (See ) 19. In the electron configuration, the number written first indicates the type of the electron is located in. 20. The letters s,p,d, or f indicate the type of the electron is occupying. 21. The superscript represents the of the electrons occupying that sublevel 22. The arrangement of the electrons in the atom may also be represented by which uses lines to represent the space orbitals and arrows to represent the electrons. 23. Still another method of representing an atom is the which uses the symbol of the element to represent the nucleus and all electrons except the valence electrons. 24. In this method are used to represent the electrons in the highest energy level. 25. In filling the p, d, and f orbitals, when do the electrons begin to pair up? 26. In the electron configuration for potassium, why does the 19 th electron go into the 4s orbital instead of the 3d? 27. What is the maximum number of dots in an electron dot diagram of an atom?

15 Name: Date: Pd.: Worksheet 5- More on Electron Configuration 1. What is the total number of electrons in the 2p sublevel of a chlorine atom in the ground state? a. 6 b. 2 c. 3 d Which is the electron configuration of an atom in the excited state? a. 1s 2 2s 1 b. 1s 2 2s 2 2p 1 c. 1s 2 2s 2 2p 5 d. 1s 2 2s 2 2p 2 3s 1 3. The Ca 2+ ion differs from a Ca atom in that the Ca 2+ ion has a. more protons b. fewer protons c. more electrons d. fewer electrons 4. The total number of orbital in the 4f sublevel is a. 1 b. 5 c. 3 d Which electron transition is accompanied by the emission of energy? a. 1s to 2s b. 2s to 2p c. 3p to 2s d. 3p to 4p 6. What is the total number of nucleons (protons & neutrons) in an atom of selenium? a. 34 b. 45 c. 79 d What is the total number of principal energy levels that are completely filled in Mg in it s ground state? a. 1 b. 2 c. 3 d What is the maximum number of electrons that can occupy the 4d sublevel? a. 6 b. 2 c. 10 d Which sublevels are occupied in the outermost principal energy level of an argon atom in the ground state? a. 3s & 3d b. 3s & 3p c. 2s & 3p d. 2p & 3d 10. Which element has an atom in the ground state with the most loosely bound electrons? a. He b. As c. Xe d. Cs 11. Isotopes of an element have a different a. # of electrons b. # of protons c. atomic # d. mass # 12. A neutral atom of an element has an electron configuration of What is the total number of p electrons in this atom? a. 6 b. 2 c. 10 d Which is the electron configuration of a hydrogen atom with an atomic mass of 3 and an atomic number of 1 in the ground state? a. 1s 1 b. 1s 2 c. 1s 2 2s 1 d. 1s 2 2s When an electron of an atom of hydrogen moves from the second to the first principal energy level, the result is the emission of a. a beta particle b. an alpha particle c. quantized energy d. gamma rays 15. How many occupied sublevels are in an atom of carbon in the ground state? a. 5 b. 6 c. 3 d How many orbitals in a sulfur atom in the ground state contain only one electron? a. 1 b. 2 c. 3 d. 4

16 Name: Date: Pd.: Worksheet 6- Electron Configuration Practice Write the electron configuration for the following elements. Draw their orbital notation, and draw the electron dot diagram for each. A. Sodium: Atomic # 11_ Electron configuration Electron Dot Diagram F. Sulfur: Atomic # 16_ Electron configuration Electron Dot Diagram F. Calcium: Atomic # 20_ Electron configuration Electron Dot Diagram F. Chromium: Atomic # 24_ Electron configuration Electron Dot Diagram F. Copper: Atomic # 29_ Electron configuration Electron Dot Diagram F. Bromine: Atomic # 35_ Electron configuration Electron Dot Diagram

17 G. Silver: Atomic # 47_ Electron configuration Electron Dot Diagram H. Potassium +1 ion: Atomic # Number of electrons for ion = Electron configuration I. Chlorine -1 ion: Atomic # Number of electrons in the ion = Electron configuration J. Phosphorus: Atomic # Number of electrons = Electron configuration Questions: 1. How many unpaired electrons does phosphorus have? 2. How many electrons does copper have in 4s? 3. Why does 4s fill before 3d? 4. Identify the number of valence electrons in bromine. 5. How many electrons does sulfur have in its valence level? How many of those electrons are unpaired?