Atomic Theory. Unit 3 Development of the Atomic Theory

Transcription

1 Atomic Theory Unit 3 Development of the Atomic Theory 1. Where is the mass of the atom concentrated? 2. What is located in the nucleus? 3. What is the negative particle that orbits the nucleus? 4. What is the sum of the protons and neutrons called? 5. With a neutral atom, what two items are equal in number? 6. What is the term for any charged particle? 7. What is the term for the positively charged ion? 8. What is the term for the negatively charged ion? 9. What type of ion is formed when electrons are gained? 10. What type of ion is formed when electrons are lost? 11. How does atomic theory today differ from Dalton s theory? 12. Which model of the atom is based on the solution to the Schrodinger equation? How is this different from the planetary model? Match each term from the experiments of J.J. Thomson with the correct description. 13. anode a. an electrode with a negative charge 14. cathode b. a glowing beam between electrodes 15. cathode ray c. an electrode with a positive charge 16. electron d. a negatively charged particle 17. The diagram shows electrons moving from left to right in a cathode-ray tube. Draw an arrow showing how the path of the electrons will be affected by the placement of the negatively and positively charged electrodes. H. Cannon, C. Clapper and T. Guillot Klein High School

2 18. Indicate the letter of each sentence that is true about atoms, matter and electric charge. a. All atoms have an electric charge. b. Electric charges are carried by particles of matter. c. Atoms always lose or gain charges in whole-number multiples of a single basic unit. d. When a given number of positively charged particles combine with an equal number of negatively charged particles, an electrically neutral particle is formed. 19. Indicate the letter next to the number of units of positive charge that remain if a hydrogen atom loses an electron. a. 0 b. 1 c. 2 d The positively charged subatomic particle that remains when a hydrogen atom loses an electron is called a(n). 21. What charge does a neutron carry? 22. Indicate the letter of each sentence that is true about the nuclear theory of atoms suggested by Rutherford s experimental results. a. An atom is mostly empty space. b. All the positive charge of an atom is concentrated in a small central region called the nucleus. c. The nucleus is composed of protons and neutrons. d. The nucleus is large compared with the atom as a whole. e. Nearly all the mass of an atom is in its nucleus. 23. According to Bohr, electrons cannot reside at in the figure below a. point A c. point c b. point b d. point d 24. According to the quantum mechanical model, point D in the above figure represents (a) th e fixed position of an electron (b) the farthest position from the nucleus that an electron can be found (c)a position where an electron probably exists (d) a position where an electron cannot exist 25. What does the atomic number of an atom represent? 3-2

3 Unit 4 Electrons in the atom Select the best possible response. 1.The type of charge on the nucleus is a. negative b. neutral c. positive 2. The number of protons in a neutral atom having 18 electrons is. 3. How many neutrons are in the isotope P-29? A neutral atom of bromine has a mass number of It has protons. 5. It has electrons. 6. It has neutrons. 7. Its nuclear composition is. One isotope of hydrogen has a mass of How many protons does it have? 9. How many neutrons does it have? An atom with a +1 charge has atomic number 19 and mass number It has protons. 11. It has electrons. The anion of oxygen has a -2 charge. 12. How many electrons does it have? 13. The number of protons in a neutral atom having 18 electrons is. An atom has a charge of +1 and has 10 neutrons with a mass number of

4 14. What is its atomic number? 15. How many electrons does it have? An atom has 10 protons, 8 neutrons, and 12 electrons. 16. What is its charge? 17. What is its mass number? 18. In forming this ion, the neutral atom gained what? An atom has 15 protons, 16 neutrons, and 16 electrons. 19. What is its mass number? 20. What is the overall charge on this ion? A neutral atom has a mass number of 24 and 11 protons. 21. How many protons does it have? 22. How many neutrons does it have? An atom with a -2 charge has atomic number 16 and mass number It has protons. 24. It has electrons. 25. It has neutrons. 4-4

5 Atomic Structure Practice Element/Isotope Atomic Mass Protons Neutrons Electrons number number 1) Zinc 2) Aluminum 3) Calcium 4) Sulfur 5) Bromine 6) Gold 7) Silver 8) Platinum 9) Uranium ) Plutonium ) Potassium 39 12) Mercury ) Titanium ) Titanium - 46 Atomic Structure Exercise I. Use the periodic table to compute the number of electrons. Neutrons and protons in the following: A. Cr B. Cl C. Mg D. Ir E. Si F. Ne II. Moseley used x rays to determine the atomic numbers of the elements identify each of the following elements by name. A. 1 Proton B. 4 Protons C. Protons D. 12 Protons E. 20 Protons F. 30 Protons III. How many protons are in the nucleus of each of the following elements A. Uranium B. Selenium C. Helium D. Magnesium IV. Give the number of neutrons in each of the following isotopes 4-5

6 A. Titanium-46 B. Nitrogen-15 C. 34 S 16 D. 85 Cu 29 The Spectra of Elements Everybody knows that a few drops of soup or milk spilled onto a gas burner will change th e blue gas flame into a mixture of colors, predominantly yellow. These colors can be used to identify the elements present in the substance dropped into the flame. We will observe the colors produced by several known substances. The color observed in the flame is the result of atoms of the element absorbing energy from the flame then reemitting it. The energy absorbed has an energy and wavelength in the visible region of the electromagnetic spectrum. What your senses detect as light is actually radiation that is part of a continuum that makes up the electromagnetic spectrum. The electromagnetic spectrum is a continuum of energies ranging f rom the high-energy gamma and x-rays to low-energy radio and microwaves. Radiation includes any energy emitted in all directions fr om a single source, not ju st nuclear decay. Light is one form of energy that can be radiated; - the sun constantly radiates energy into space as its matter is converted into energy. We see this light in a variety of colors, depending on the waveleng th of the energy when it reaches us. T he waves of light with the longest wavelength (red) are refracted (bent) to a lesser degree than the shorter waves (indigo and violet) as they pass through the atmosphere. The visible range of the spectrum is from about 400 nm to 700 nm. The wavelength () is the distance between successive peaks of a wave. The number of waves that pass a given point in space per second is the frequency (f). The unit is inverse time, sec -1 and is also called hertz (Hz). All light moves through a vacuum at the same rate of speed, about 3 X 10 8 m/s (the speed of light c). Frequency and wavelength are inversely related to each other, as the product of the two equals the speed of light: c = f. Since the electrons of an atom can only absorb certain amounts of energy, the wavelength of the energy emitted after excitement of the atoms is characteristic of the element. When viewed through a spectroscope, the spectrum of energy emitted by a sample of a particular element can be observed. 1. Using the gas tubes provided by your teacher, observe the spectrum of each gas sample through a spectrometer or diffraction grating and draw a diagram showing the colored bands produced by each different element. Each colored band corresponds to a different amount of energy (a different wavelength) of light emitted by an electron in the element. Each element has a unique spectrum based on the amount of energy emitted by its electrons. 2. In the next part of our investigation, we will use a procedure known as a flame test to become familiar with the colors produced by several common elements and use these observations to better understand the affect of energy on atoms and ions. 4-6

7 Procedure: a. Light a Bunsen burner. Make sure that the gas burns with a clear blue flame and has an easily distinguishable inner blue cone. b. Obtain samples of ions as provided by your teacher. These will be metal nitrates, dissolved in a solvent to allow easy distribution of the ion. c. Using the spray bottles, introduce a small amount of each ion (one squirt) into the flame and observe the result. d. After you have identified the color produced by each ion, mix two of the substances and see if you can detect the presence of each of the substances or if the colors mingle so that one cannot be distinguished from the other. e. Obtain an unknown sample and observe the colors produced. Try to identify the element present in the unknown sample. Color Key: Na = Yellow K = Violet Ca = Yellow-red Cu = Bluegreen Sr = Deep red Li = Crimson Ba = Green-yellow Further Analysis In 1913, Neils Bohr hypothesized that the electron in the hydrogen atom is allowed to have only certain amounts of energy. These energy levels would be orbits in which the electron in hydrogen would circle the nucleus of hydrogen. He believed that the electron would move from one energy level to another and would give off light when it jumped from a higher energy level to a lower energy level. The amount of energy would be different for jumps between different levels. 1) The diagram below shows a sketch of some of the possible orbits of the hydrogen electron and their corresponding energy values. If we think of the energy values as being on a number line, which orbit has the greatest value? 4-7

8 2) When an electron jumps from E 2 to E 1 the amount of energy in the light would be E 2 E1; or (-3.4 ev) (-13.6 ev) which would equal 10.2 ev. What would be the amount of energy in the light when an electron jumps from: a) E 3 to E 4 b) E 4 to E 2 c) E 3 to E 2 d) E 5 to E 2 3) Using Bohr s model, we would assume that the electron would only move between certain orbitals or energy levels. Every possible jump corresponds to light of a different energy. a. How many different energies of light can be emitted from hydrogen when the electron jumps down to E 2 from E 3, E 4, E 5, and E 6? b. How many bands of light did you observe when you viewed the hydrogen tube through the spectrometer? c. How do you think these two observations are related? 4) The colors corresponding to jumps to the E 1 level have higher energies or lower, than those to the E 2. a. As the electron jumped from the E 6 all the way to the E 1 level, how many different energies would be emitted? b. Did you see all of these bands? Explain why or why not. c. Bohr did not find bands corresponding to the jumps to the E3 level. Why do you think that was? Conclusions: In this activity we saw the evidence that energy is emitted in particular patterns depending on the distance that an electron travels between one orbital and another. The number of electrons in an atom and the number of energy levels those electrons occupy determine the number of bands of light and the color of the light seen. 4-8

9 What if any conclusions can you draw about the relative numbers of electrons in the atoms of our test atoms? 4-9