KAP Chemistry Fall Final Review 2024 PDF

KAP Chemistry Final Exam Review: Fall Atomic Structure and Nuclear Chemistry 1. Describe the structure of an atom. What is it composed of? An atom is composed of a dense but small nucleus surrounded by an electron (negatively charged) cloud. The nucleus contains the protons (posit...

KAP Chemistry Final Exam Review: Fall Atomic Structure and Nuclear Chemistry 1. Describe the structure of an atom. What is it composed of? An atom is composed of a dense but small nucleus surrounded by an electron (negatively charged) cloud. The nucleus contains the protons (positive charge) and neutrons (neutral). 2. Compare and contrast the subatomic parts of an atom. Protons; positive charge, about 2000 x as massive as an electron, approx. equal to mass of a neutron: Neutrons, neutral, similar mass to protons: Electrons; negative charge, much smaller and lighter than protons and neutrons. 3. What is the smallest part of an element that retains the properties of that element? Atom 4. What accounts for the vast majority of the mass of an atom? And the volume? Mass – nucleus, volume – electron cloud 5. Define isotope. What remains the same in two isotopes of the same element? Atoms of the same element that have different numbers of neutrons (and therefore different mass numbers). Number of protons and electrons stays the same, as do chemical properties. 6. What are two ways to express the isotope of an atom with 15 protons and 16 neutrons? Phosphorus-31 7. What are two ways to express the isotope of an atom with a mass number of 80 and 35 electrons? Bromine-80 8. What atom has an average atomic mass in nature of approximately 24 amu? A magnesium atom 9. Describe the gold foil experiment. Who performed it? What part of the atom was discovered because of it? The gold foil experiment was performed by Rutherford, He shot a stream of alpha particles at a thin sheet of gold foil, expecting them to pass through undeflected. However, several particles were significantly deflected and some bounced straight back. This led to the conclusion that the positive charge in an atom was not spread evenly throughout but was concentrated in the nucleus (this is the only way there would be something massive enough in the atom to cause the relatively large alpha particle to be deflected). Thus, the nucleus was discovered. 10. List the 5 components of Dalton’s Atomic Theory. Which parts have now been corrected based on information we have acquired since the time of Dalton? All matter is composed of extremely small particles called atoms Atoms of the same element are identical and have the same properties; atoms of different elemtns have different properties. (X) Atoms are indivisible and indestructible. They cannot be created, destroyed or changed into another element. (X) Atoms of different elements combine in whole number ratios to form compounds. This is known as the “Law of Definite Proportions”. When atoms react, they can sometimes combine in more than one whole number ratio to form different compounds. This is known as the “Law of Multiple Proportions”. Since Dalton's time, we have learned more about the atom, providing more information about points 2 and 3 (above): Atoms can be divided into smaller subatomic particles known as protons, neutrons and electrons. 1  Atoms of the same element are NOT all identical. They can have different masses because they differ in the number f neutrons. These “varieties” of one element are known as isotopes. During nuclear reactions (which are fundamentally different to chemical reactions), it is possible for one atom to be transformed into a different type of atom. 11. Describe the Bohr model of the atom and state the flaw in this model based on what we now know to be true. Electrons are restricted to specific energies and follow paths called orbits a fixed distance from the nucleus, similar to the way that planets orbit the sun. However, the Quantum Mechanical Model later showed that electrons DO NOT have neat orbits like the planets. 12. Describe the current model of the atom, The Quantum Mechanical Model. This is a mathematical model, developed in the 1920s by the work of many scientists including De Broglie, Schrödinger, Pauli and Heisenberg. In the quantum model, electrons are described as standing waves of energy. Electrons exist in regions of space around the nucleus called orbitals. The paths of the electrons within the orbitals are random and therefore cannot be predicted. We can only talk about the probability of an electron being in a certain region. Electrons have a high probability of being located near the nucleus and are in constant motion. 13. Carbon exists as carbon-12 and carbon-14. Which isotope occurs in greater abundance? Explain. The average atomic mass of carbon is 12.011 amu which is weighted much more towards 12 than 14, indicating that there is more of 12 than 14 in the average. 14. Determine the average atomic mass of neon, composed of neon-20 (mass 19.99 amu, 90.15%), neon-21 (mass 20.99, 0.27%) and neon-22 (mass 21.00 amu, 9.22%) 20.01 amu 15. An atom of I has a mass number of 127. How many neutrons does it contain? 74 16. What is the atomic number of potassium-39? What is its mass number and how many neutrons does it have? Atomic number, 19; Mass number, 39, Number of neutrons, 20 17. An atom has atomic number P and mass number M. How many neutrons does it have? Neutrons = M - P 18. Define radioactive decay. A process in which certain isotopes of some elements spontaneously break down to produce a smaller element along with particles (alpha, beta) and energy (gamma radiation). 19. Define half-life. The time during which a given atom of a radioactive isotope has a 50:50 chance of decaying OR the time it takes for half of a sample of a radioactive isotope to decay. 20. Is the half-life the same for all atoms of a specific radioactive isotope? For all isotopes of different elements? Yes; for example, all atoms of carbon-14 have the same half-life. No; each radioactive isotope has a unique half-life different from others. For example, carbon-14 has a different half-life than cobalt-60. 21. Determine the original and final amounts for the following: Time elapsed Half-life Original Amount Final amount 35 s 70 s 100 g 25 g 20 min 2 hours 38.4 g 0.6 g 2  1.6 days 8 days 6000 mg 187.5 mg 4 years 16 years 120 g 7.5 g 22. Define alpha (α) particles and their characteristics and know the symbol. Helium nucleus, , more massive, most damage over short range, least penetrating (stopped by a sheet of paper) 23. Define beta (β) particles and their characteristics and know the symbol. High speed electron, , less massive, less damage over short range, moderately penetrating (stopped by a thin sheet of aluminum foil) 24. Define gamma (γ) radiation and its characteristics. No mass, all energy, cause least damage over a short range, most penetrating (stopped by a thick sheet of lead) 25. Write the complete reaction for the following radioactive decay processes: 26. Contrast nuclear fission with nuclear fusion. Write equations for both and state the uses and/or where these reactions are observed. Fission: Fusion: Developed during WWII for use in nuclear Occurs on the stars, including our sun. They weapons (uranium or plutonium fuel). Currently liberate 3-10 times more energy than fission used as the energy source in nuclear power plants, reactions and do not generate nuclear waste but man has not found a way to carry out fusion reactions in an efficient way. Electrons 27. What is the maximum number of electrons that can occupy the 4th energy level? 32 (cool hack: 2n2 where n = the energy level. Be aware that this is totally theoretical, and the larger energy levels may never hold their maximum capacity because even our largest atoms are not large enough to supply that number of electrons). 28. Name and explain the three rules that describe how electrons enter an atom’s orbitals. Rule 1: Aufbau Principle. Lowest energy orbitals fill first, therefore 1s, 2s, 2p, 3s, 3p, 4s, 3d, etc. Since the orbitals within a subshell are degenerate (of equal energy), the entire subshell of a particular orbital type is filled before moving to the next subshell of higher energy. Rule 2: Pauli Exclusion Principle. Only two electrons are permitted per orbital and they must be of opposite spin. If one electron within an orbital possesses a clockwise spin, then the second electron within that orbital must possess a counter-clockwise spin. 3  Rule 3: Hund’s Rule. The most stable arrangement of electrons in a subshell occurs when the maximum number of unpaired electrons exist, all possessing the same spin direction. This occurs due to the degeneracy of the orbitals, all orbitals within a subshell are of equal energy. Electrons are repulsive to one another and only pair after all the orbitals have been singly filled. 29. What is the electron configuration for the atom with atomic number 12? 1s2 2s2 2p6 3s2 30. What is the electron configuration for the atom with atomic number 24? 1s2 2s2 2p6 3s2 3p6 4s1 3d5 31. How many valence electrons do the following elements have a. S 6 b. Cl 7 c. K 1 d. Kr 8 32. What is the atomic number of the element with electron configuration 1s2 2s2 2p6 3s1? 11 33. Draw an orbital diagram for phosphorus. 34. Draw an electron dot diagram for phosphorus. What do the dots represent? , dots represent valence electrons 35. What is the maximum number of d electrons in the 4th energy level? 10 36. What is the maximum number of p orbitals in the 3rd energy level? 3 37. Which energy level is the first to have an occupied f level? 4 38. An atom has electrons in the 3p, 2p, 3s, 2s, and 1s orbitals. Which electrons are at the lowest energy? Those in the 1s orbital 39. Calculate the wavelength (in meters) of radiation with a frequency of 5.00 x 1014 Hz. In what part of the electromagnetic spectrum will this radiation be found? 6.00 x 10-7 m or 600 nm. Within visible light (yellow). (See EM Spectrum chart in notes). 40. A certain source of radiation emits a wavelength of 500.0 nm. What is the energy in J? 3.98 x 10-19J 41. What is a photon? Quantum packet of light, a “particle light 42. Arrange the types of electromagnetic radiation by increasing wavelength. Now arrange by increasing frequency. Finally arrange in order of increasing energy. Wavelength: gamma, x-ray, UV, Visible (violet-red), Infrared, microwave, radio Frequency (opposite): radio, microwave, infrared, visible (red-violet), UV, x-ray, gamma Energy: same as frequency Periodic Table 43. In general metals have few/many valence electrons and they lose/gain electrons to become stable. Non- metals have few/many valence electrons and they lose/gain electrons to become stable. (Circle the correct answer). 44. Define Ionization Energy and state the trend down a group and across a period. Explain why. 4  The amount of energy required to remove an electron from a gaseous atom, producing a cation (positively-charged ion). Down a group the IE decreases because the electrons are further from the positive pull of the nucleus and the shielding effect of the other electrons in the atom. This means that it is easier to remove an electron. Across a period, IE increases because increasing nuclear charge results in more tightly held electrons, requiring more energy to remove. 45. Define Electronegativity and state the trend down a group and across a period. Explain why. Electronegativity is the tendency of an atom in a compound to attract the bonding electrons to itself. EN decreases down a group because the shielding effect makes it more difficult for the nucleus to attract bonding electrons. EN increases across a period because increasing nuclear charge results in stronger attraction for the bonding electrons. 46. State the trend for atomic radius across a period and explain why the trend is observed. Atomic radius decreases across a period because increasing nuclear charge pulls the electrons closer to the nucleus. 47. State the trend for atomic radius down a group and explain why the trend is observed. Atomic radius increases down a group because added energy levels increase the size of the electron cloud. 48. Who is credited with developing the first Periodic Table? What procedure did he use? Dmitri Mendeleev collected information on the elements and wrote them on cards. He arranged the cards in order of increasing atomic mass and grouped elements with similar properties in the same columns. By observing gaps in his card system, he was able to predict the existence of elements that had not yet been discovered. 49. Henry Mosely realized that Medeleev’s table was actually arranged in increasing atomic number. 50. State the periodic law. The physical and chemical properties of the elements are periodic functions of their atomic numbers. 51. Elements in the same group in the periodic table can be expected to have similar properties. 52. Describe the halogens in terms of electron configuration pattern and properties. Electron configs end in p7. Halogens are the most reactive non-metals. They have a strong odor and will burn flesh. 53. Describe the alkali metals in terms of electron configuration pattern and properties. Electron configs end in s1. They are the most reactive metals, are shiny and soft, are not found uncombined in nature and are typically stored in oil. 54. Describe the noble gases in terms of electron configuration pattern and properties. Electron configs end in s2p6. They are largely unreactive (resist forming compounds) beause they have an octet of valence electrons. 55. What is the most electronegative element on the periodic table? And the least (excluding noble gases)? Most - Fluorine Least - Francium Bonding 56. Which of the following contain(s) covalent bonds? CO2 Yes K2O No NaCl No SO42- Yes How can you tell? Covalent bonds are formed between nonmetals 57. Correctly represent a formula unit of sodium chloride, potassium phosphate and magnesium sulfate. NaCl, K3PO4, MgSO4 58. Which of the following do not have the same electron configuration: Mg 2+ and Mg No Mg 2+ and Ne Yes Cl- and Ar Yes Br- and Xe No 5  59. How many protons and electrons are in the following ions: F- 9p, 10e Li+ 3p, 2e S2- 16p, 18e Al3+ 13p, 10e 60. In what state(s) can an ionic compound conduct electricity? When molten or aqueous (dissolved in water). (Cannot conduct in solid state). 61. Describe the electron sea model of metals. What properties of metals can be accounted for by this model? Metals have few electrons in their highest energy levels meaning there are vacant orbitals. These vacant orbitals are used by the valence electrons to move freely throughout the metal creating an “sea” of delocalized electrons. This accounts for metals’ high electrical and thermal conductivity. 62. Draw a Lewis Structure for the phosphate ion. How many unshared (lone) pairs of electrons are there? 12 lone pairs (24 electrons) 63. How many valence electrons (A, available) are in a molecule of ammonia? 8 64. Draw all possible resonance forms for the sulfur trioxide molecule. 65. What is the molecular geometry (shape) of carbon tetrachloride? Tetrahedral 66. What is the molecular geometry (shape) of carbon dioxide? Linear 67. What is the molecular geometry (shape) of phosphorus trichloride? Trigonal pyramidal 68. What is the molecular geometry (shape) of the nitrate ion? Trigonal planar 69. Which of the following molecules is/are polar? CO2, H2O, NH3 H2O and NH3 are polar. CO2 contains polar bonds but is a non-polar molecule 70. Which of the following molecules is/are nonpolar? CCl4, C3H6, OCl2 CCl4 and C3H6 are non-polar. OCl2 is polar due to bent geometry Nomenclature: Chemical Names and Formulas 71. What are the individual parts of an ionic compound called? Formula units (made up of ions) 72. What are the individual parts of a molecular compound called? Molecules (made up of atoms) 73. Define a cation and an anion 6  Cation – positively charged ion. Anion – negatively charged ion 74. State the number of electrons lost or gained when these elements form ions. K 1 lost Br 1 gained Al 3 lost S 2 gained Ca 2 lost O 2 gained N 3 gained Na 1 lost 75. What is the formula for silver sulfide? Ag2S 76. What is the formula for lead (IV) phosphate? Pb3PO4 77. What is the formula for tetraphosphorus decoxide? P4O10 78. What is the formula for barium sulfate? BaSO4 79. What is the name for HNO2? Nitrous acid 80. What is the name for H2CO3? Carbonic acid 81. What is the name for Cu2O? Copper (I) oxide 82. What is the name for N2O5? Dinitrogen pentoxide Significant Figures 83. Round 0.000362548 to three significant figures. 0.000363 or 3.63 x 10-4 84. The number of significant figures in 0.02810 is 4 85. The number of significant figures in 21 000 is 2 86. Work out the following calculation and express to the correct number of significant figures. 25.36 + 21.0 - 36.2256 = 10.1 87. Work out the following calculation and express to the correct number of significant figures. (3.56 x 0.07100)/220 = 0.0011 or 1.1 x 10-3 Dimensional Analysis 88. Convert 0.68 mg to g 0.00068 g or 6.8 x 10-4 g 89. Convert 1036 kg to cg 103,600,00 cg or 1.036 x 108 cg 90. Convert 9.6 m to dm 96 dm 91. Convert 160.8 cm3 to L 0.1608 L 92. How many seconds are in 3 days? 259,200 s (3 x 105 s with SF) 93. Convert the speed of light (3.0 x 108 m/s) to miles/hour. (1mile = 1.609 km) 6.7 x 108 mi/hr 94. A primitive society uses a barter system to purchase goods. A farmer wants to exchange his corn for a dairy cow. How many bushels of corn will he need? (1 cow = 95 baskets of pears; 4 baskets of pears = 42 tubs of butter; 2 bushels of corn = 15 tubs of butter) 133 bushels (SF do not apply here; exact counted objects) 95. A metal has a density of 3.8 g/cm3. How many L of water will a 38g sample of the metal displace? 10. L 7 Chemical Quantities: Moles 96. What is the definition of molar mass? What is the correct unit? The mass in grams of 1 mole of a substance. Unit = g/mol 97. What is the molar mass of aluminum sulfate? 342.17 g/mol 98. How many moles of potassium hypochlorite are present in 45.0g of potassium hypochlorite? 0.497 mol 99. How many nitrate ions are present in 3.00 mol of copper (II) nitrate? 3.61 x 1024 NO- ions 100. How many hydroxide ions are present in 1 formaula unit of aluminum hydroxide? 3 OH- ions 101. How many fluorine atoms are present in 1.50 mol of carbon tetrafluoride? 3.61 x 1024 F atoms 102. How many molecules of sulfur dioxide are present in 4.25g of sulfur dioxide? 3.99 x 1022 SO2 molecules 103. What is the percentage composition of glucose, C6H12O6? 40.0% C, 6.7% H, 53.3% O 104. Determine the mass of oxygen in 34.5 grams of potassium chlorate? 13.5 g oxygen 105. What is the empirical formula for a compound that is 32.4% sodium, 22.5% sulfur and 45.1% oxygen by mass? Na2SO4 106. What is the empirical formula for a compound that is 54.5% carbon, 9.1% hydrogen and 36.4% oxygen by mass? C2H4O 107. The molecular formula for butene is C5H10. What is its empirical formula? CH2 108. The empirical formula of a compound is NO2. Its molar mass is 92.02 g/mol. What is the molecular formula of the compound? N2O4 109. The empirical formula of a compound is C4H10O. Its molar mass is 222.42 g/mol. What is the molecular formula of the compound? C12H30O3 110. How many atoms of Nitrogen are in 7.50L of nitrogen gas at STP? 4.03 x 1023 atoms of N (Remember that the representative particles of nitrogen gas are diatomic molecules). 8