A1: Atomic Structure Worksheet Key (Goals 1 3, Chapter 4) 1. Democritus, who lived in Greece during the 4 th century B.C., suggested that matter is made up of tiny particles that cannot be divided. He called these particles atoms. 2. Explain why the ideas of Democritus were not useful in a scientific sense. Democritus s ideas don t explain chemical behavior & lacked experimental support. 3. The modern process of discovery about atoms began with the theories of an English school teacher named John Dalton. 4. Circle the letter of each sentence that is true about Dalton s atomic theory. a. All elements are composed of tiny, indivisible particles called atoms. b. An element is composed of several types of atoms. c. Atoms of different elements can physically mix together, or can chemically combine in simple whole number ratios to form compounds. d. Chemical reactions occur when atoms are separated, joined, or rearranged; however, atoms of one element are never changed into atoms of another element when the atoms of elements A and B combine chemically. 5. Which atomic particles carry a negative charge? Electrons carry a negative charge. 6. Thomson observed that the production of cathode rays did not depend on the kind of gas in the tube or the type of metal used in the electrodes. What conclusion did he draw from these observations? Thomson concluded that electrons must be parts of the atoms of all elements. 7. Explain Thomson s Plum Pudding model of the atom. The Plum Pudding model of the atom is a diffuse positively charged sphere with negatively charged particles dispersed throughout. 8. How many units of positive charge remain if a hydrogen atom loses an electron? one 9. The positively charged subatomic particle that remains when a hydrogen atom loses an electron is called a(n) proton. 10. Circle 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. d. The nucleus is large compared with the atom as a whole. e. Nearly all the mass of an atom is in its nucleus.
11. Describe Rutherford s gold foil experiment and explain how his results improved upon Thomson s Plum Pudding model of the atom. Rutherford shot positively charged alpha particles at a very thin sheet of gold foil. Nearly all the alpha particles went through the gold foil (the atom is mostly empty space), relatively few of the alpha particles where deflected at wide angles (positively charged particles are present in the atom), and a very small number of alpha particles bounced straight back (there is a very dense region in the center of an atom). Rutherford s gold foil experiment led to the solar system model of the atom where electrons orbit a dense positive nucleus. 12. Fill out the following table: Name Symbol Charge Mass (amu) electron e - -1 1/1837 proton p + +1 1 neutron n 0 0 1 13. Would you expect two electrons to attract or repel each other? Why? Two electrons would repel because electrons have a charge of negative 1 and like charges repel each other. 14. What is the charge, positive or negative, of the nucleus of every atom? The charge of the nucleus is positive because the nucleus is composed of protons and neutrons with a charge of positive 1 and zero, respectively. Therefore, the net charge of the nucleus is positive. 15. Why is every atom electrically neutral? Atoms are electrically neutral because they have the same number of electrons and protons. A2: Isotopes Worksheet (Goal 4 5) 1. Write the isotopic symbols for the isotopes of uranium having the following number of neutrons. 234 a. 142 neutrons 92U 235 b. 143 neutrons 92U 238 c. 146 neutrons 92U 2. Fill in the following table: Name # of protons # of neutrons # of electrons Boron-10 5 5 5 Sulfur-33 16 17 16 Iodine-127 53 74 53 Chlorine-36 17 19 17 Calcium-40 20 20 20 3. What do isotopes of the same element have in common? How do isotopes of the same element differ? Isotopes have an equal number of protons and different number of neutrons.
4. Fill in the following table: Name # of electrons # of protons # of neutrons atomic number mass number carbon-14 6 6 8 6 14 carbon-12 6 6 6 6 12 tin-119 50 50 69 50 119 tin-120 50 50 70 50 120 lithium-7 3 3 4 3 7 5. Write the isotopic symbols for argon-36, argon-38, and argon-40. 36 38 40 18Ar, 18Ar, 18Ar 6. How can there be more than 1000 different atoms when there are only about 100 different elements? Each element may have several different isotopes. 7. How is the carbon-12 atom used to define atomic mass unit? Atomic mass unit is 1/12 the mass of a carbon-12 atom. 8. Distinguish among atomic mass and mass number. atomic mass weighted average of the masses of the isotopes of an element. mass number number of protons and the number of neutrons in the nucleus of an atom 9. What data must you know about the isotopes of an element to calculate the atomic mass of the element? In order to calculate the atomic mass of an element the mass and relative abundance of each isotope is needed. 10. Element X has two naturally occurring isotopes. The isotope with mass 10.0 amu has a relative abundance of 20.0%. The isotope with a mass number of 11.0 amu has a relative abundance of 80.0%. Calculate the value of the atomic mass of element X. State the atomic number and true identity of element X. 10.0 amu x 0.200 + 11.0 amu x 0.800 = 10.8 amu X = Boron w/ 5 protons
11. The lithium found in a hearing aid battery has two naturally occurring isotopes. Lithium-6 has a mass of 6.01 amu and an abundance of 7.42%. Lithium-7 has a mass of 7.01 amu and an abundance of 92.58%. Calculate the atomic mass of lithium. 6.01 amu x 0.0742 + 7.01 amu x 0.9258 = 6.94 amu 12. Directions: complete each of the following models using what you know about atoms. The first is completed for you as an example. Atomic # 7 Atomic # _6 Atomic # 18 Atomic # _1_ Atomic Mass 14 Atomic Mass_12_ Atomic Mass_40_ Atomic Mass_1 Isotope Name: Nitrogen- 14 Symbol_ 14 N_ Atomic # _20_ Element: Carbon Symbol 12 C_ Atomic # 5_ Element: Argon Symbol_Ar_ Atomic # _13_ Element: Hydrogen Symbol _H_ Atomic # _16_ Atomic Mass_40_ Atomic Mass_11_ Atomic Mass_27_ Atomic Mass_32_ Element: Calcium Symbol_Ca Element: Boron Symbol _B_ Element: Aluminum Symbol Al_ Element:_Sulfur Symbol_S_ A3: Atomic Theory and Orbitals Worksheet (Goals 6-9, Chapter 5) 1. Describe how the quantum theory of atomic structure differs from Bohr s theory. ( Hint: Focus on electrons) Bohr electrons in orbits, Quantum electrons in orbitals 2. Differentiate between an orbit and an orbital. Orbit circular path, Orbital region of probability
3. How are electrons in the ground state different from electrons in the excited state? Ground state electrons are lower in energy than the excited state. 4. What unusual property is observed when an electron falls from excited state to a ground state? Light is given off due to the loss of energy. 5. How many orbitals are in the fourth energy level? 1 s orbital + 3 p orbitals + 5 d orbitals + 7 f orbitals = 16 orbitals total 6. How many orbitals of each type are there? s - 1 p - 3 d - 5 f - 7 7. Draw the 2s and 2px orbital. sphere, dumb-bell on the x axis. 8. What is a line spectrum? What does it represent? A line spectrum represents the energy an atom s electrons will absorb and release in the visible spectrum. 9. Fireworks give off many different colors. Using Bohr s theory and your observations from the spectroscope, explain how the firework s chemicals produce different colors. Different electron configurations of different elements will absorb and release different unique amounts of energy (colors). 10. Why do we not see a line spectrum with our eyes from the fireworks? Human eyes are not spectroscopes and can t split light up into component parts.
Atomic Orbitals and the Periodic Table Rules: 1. Aufbau Principle electrons enter orbitals of lowest energy first 2. Hund s Rule when electrons enter orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins 3. Pauli s Exclusion Principle an atomic orbital can describe at most two electrons Assigning Electrons: H 1s 1 Be 1s 2 2s 2 C 1s 2 2s 2 2p 2 Ne 1s 2 2s 2 2p 6 Na 1s 2 2s 2 2p 6 3s 1 Cl 1s 2 2s 2 2p 6 3s 2 3p 5 Fe 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 Ge 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 Kr 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 Rh 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 7 Hg 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 O 2-1s 2 2s 2 2p 6 Ca 2+ 1s 2 2s 2 2p 6 3s 2 3p 6 A4: Electron configs w/ Shorthand (Goals 6-9, Chapter 5) Rule 4: Noble Gas Shorthand Method: The last noble gas that was completed prior to arriving at your element can be written down with the symbol of that noble gas in [brackets]. Then complete the valence electrons to arrive at your element. We assume that the electrons are full in every shell up to that noble gas. Directions: Write complete Electron Configurations and Orbital Diagrams for the following. Element Long Method Noble Gas Shortcut Method Br 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 [Ar] 4s 2 3d 10 4p 5
Br -1 Mg xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx [Ar] 4s 2 3d 10 4p 6 [Ne] 3s 2 He 1s 2 xxxxxxxxxxxxxx Li 1s 2 2s 1 [He] 2s 1 B xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx [He] 2s 2 2p 1 H 1s 1 xxxxxxxxxxxxxx S 1s 2 2s 2 2p 6 3s 2 3p 4 [Ne]3s 2 3p 4 S -2 xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx [Ne]3s 2 3p 6 Ti xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx [Ar] 4s 2 3d 2 As 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 xxxxxxxxxxxxxx Na xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx [Ne]3s 1 P 1s 2 2s 2 2p 6 3s 2 3p 3 xxxxxxxxxxxxxx Cr xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx [Ar] 4s 2 3d 4 Co xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx [Ar] 4s 2 3d 7 Ag 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 9 [Kr] 5s 2 4d 9 Al 1s 2 2s 2 2p 6 3s 2 3p 1 xxxxxxxxxxxxxx Al +3 xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx [He]2s 2 2p 6 Ne 1s 2 2s 2 2p 6 xxxxxxxxxxxxxx