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Goal 1 Study Guide and Practice Problems 1. Fill in the following table: Particle Protons Neutrons Electrons Location Nucleus Nucleus Electron cloud Relative Charge +1 0 -1 Relative Mass (amu) 1 amu 1 amu 1/2000 amu 2. Isotope notations: 235 92π and U-235 represents the same isotope of an element. a. What does β235βrepresent? Mass number b. What does β92β represent?Atomic number c. What does the U represent? Chemical symbol for uranium 3. What information does the mass number you about the nucleus of an atom? The mass number refers to the total number of protons and neutrons in the nucleus of an atom. 4. What information does the atomic number tell you about the nucleus of an atom? The atomic number refers to the number of protons in the nucleus of an atom. 5. How would you determine the number of protons, electrons, and neutrons contained in a neutral atom? a. Protons = atomic number b. Electrons = protons (in a neutral atom) c. Electrons = protons β electrons (for an ion) d. Neutrons = mass number β protons (or mass number β atomic number) 6. Practice Problems: Atomic Structure Math Isotope symbol Isotope name Protons Neutrons Electrons Atomic number Mass number ππ πππͺπ Chlorine-35 17 18 17 17 17 ππ πππͺπ Calcium-40 20 20 20 20 40 π ππ©π Beryllium-9 4 5 4 4 9 πππ πππΌ Uranium-235 92 143 92 92 235 ππ πππ©π Bromine-80 35 45 35 35 80 7. How do isotopes of the same element differ and how are they similar? Isotopes of the same element have a different number of neutrons, but have the same number of protons. 8. How do atoms of different elements differ? Atoms of different elements have a different number of protons. 9. How are average atomic mass, actual isotopic mass, and the mass number of specific isotopes different? The average atomic mass is the weighted average mass of the naturally occurring isotopes of an element. The actual isotopic mass is the mass of one of the isotopes of an element. The mass number is the number of protons and neutrons in the nucleus of an atom. 10. Why does a weighted average have to be used to calculate the average atomic mass? Some isotopes are not as abundant as the other isotopes of an element. 11. What two factors does average atomic mass depend on?Actual isotopic mass and relative abundance. 12. A Bohr Model can be effective for describing the number of energy levels and the number of valence electrons for only the first 18 elements, in a Bohr model how many electrons are located in each of the first three energy levels? 1st- 2electrons, 2nd- 8 electrons and 3rd-8 electrons. 13. Practice Problems: Drawing Bohr Models a. sodium b. sulfur c. aluminum d. fluorine e. calcium f. nitrogen b. 2-8-1 b. 2-8-6 c. 2-8-3 d. 2-7 e. 2-8-8-2 f. 2-8-5 14. An electron configuration can be effective for describing the number of energy levels and the number of valence electrons for all of the elements. In the quantum mechanical model, what is the maximum number of electrons allowed in each of the first four energy levels? a. 1st energy level = 2 b. 2nd energy level = 8 c. 3rd energy level = 18 d. 4th energy level = 32 15. Fill in the blanks: a. When an electron gains/absorbs an amount of energy equivalent to the energy difference, it moves from its ground state to a higher energy level. b. When the electron moves to a lower energy level, it loses/releases an amount of energy equal to the energy difference in these levels as electromagnetic radiation. c. Since the light that is released by an electron is constant (speed of light), the wavelength and frequency of light are inversely related. Which means that as the frequency of light increases, the wavelength decreases. (c=f*Ξ») d. Energy and frequency of light are directly related, which means the as the frequency of light increases, the energy increases. (E=h*f) e. Niels Bohr produced a model of the hydrogen atom based on experimental observations. This model indicated that: i. An electron circles the nucleus only in fixed energy ranges called orbits. ii. An electron can neither gain nor lose energy inside this orbit, but could move up or down to another orbit. iii. The lowest energy orbit is closest to the nucleus. 16. Complete the following table: Symbol Change in mass # Change in atomic # Penetrating ability Alpha Beta Gamma 4 2π»π 0 β1π Ξ³ Decreases by 4 Decreases by 2 Paper No change Increases by 1 Wood No change No change Lead or concrete 17. Practice Problems: Balancing Nuclear Equations πππ π a. Uranium-235 (alpha decay) πππ πππΌ β ππ―π + πππ»π b. Carbon-14 (beta decay) ππππͺ β βπππ + ππππ΅ 18. Practice Problems: Half-life a. In 5.49 seconds, 1.20 g of argon-35 decay to leave only 0.15 g. What is the half-life of argon-35? T1/2 = 1.83 s b. How many days does it take for 16 g of palladium-103 to decay to 1.0 g? The half-life of palladium-103 is 17 days. T = 68 days c. Sodium-24 has a half-life of 15 hours. How much sodium-24 will remain in an 18.0 g sample after 60 hours? M = 1.125 g 19. How are radioactive decay, fission, and fusion different? a. Radioactive decay- when an unstable nucleus breaks apart into smaller nuclei, or change in some other way to make it more stable. (Spontaneous and random) b. Nuclear fission- is a process by which a large atomic nucleus breaks up to form smaller ones. c. Nuclear fusion- is a process by which small nuclei combine to form larger ones. This process is accompanied by an even greater production of energy than nuclear fission. 20. Fill in the blanks: a. When a neutral atom gains one or more electrons it becomes a negatively charged ion, which is called a(n) anion. b. When a neutral atom loses one or more electrons it becomes a positively charged ion, which is called a(n) cation. 21. Fill in the table: What happens to the electrons? Types of elements ΞEN Melting point Boiling point Conductivity Other Ionic Covalent Metallic Transferred Shared Sea of electrons Metal/nonmetal >1.7 High High In a molten state or aqueous solution Brittle solid two nonmetals <1.7 Low Low Two metals n/a High High Poor conductors Good conductors Polar nature Malleable and ductile 22. Practice Problems: Draw Lewis structures for simple compounds. Cannot post these diagrams. Ask in class if you would like to see these. a. SO2 b. CO3-2 c. CO2 d. ClO4-1 e. ClO2-1 f. H2O 23. Practice: Predicting Shapes and Bond Angles for molecular compounds a. SO2 bent; 120 b. CO3-2 trigonal planar; 120 c. CO2 linear; 180 d. ClO4-1 tetrahedral; 109 e. ClO2-1 bent; 109 f. H2O bent; 109 24. Practice: Predict bond polarity a. S-O 3.44-2.58=0.86; polar b. C-O 3.44-2.55=0.87; polar c. Cl-O 3.44-3.16=0.28; nonpolar d. H-O 3.44-2.20=1.24; polar e. F-F 3.98-3.98=0; nonpolar f. H-Cl 3.16-2.20=0.96; polar 25. Rank single, double, and triple bonds in terms of strongest to weakest and longest to shortest. a. Strongest to weakest: triple, double, single b. Longest to shortest: single, double, triple 26. Rank in order of strongest bond to weakest: Ionic bonds, metallic bonds, hydrogen bonds, dispersion forces, dipole-dipole, and covalent. Covalent, ionic, metallic, hydrogen, dipole-dipole, and dispersion. 27. Classify the following as intermolecular or intramolecular bonds: dispersion forces, metallic bonds, covalent, ionic bonds, covalent bonds, and dipole-dipole. a. Intermolecular forces: hydrogen bonding, dipole-dipole, and dispersion. b. Intramolecular forces: ionic, covalent, and metallic bonds 28. Intermolecular forces for molecular compounds: a. Hydrogen bonding: attraction between molecules when H is bonded to N, O, or F. b. Dipole-dipole: attractions between polar molecules. c. Dispersion forces: electrons of one molecule attracted to nucleus of another molecule. 29. Practice: Writing formulas for compounds a. K2O f. Ag2O k. BaF2 b. AgNO3 g. Zn(ClO4)2 l. Fe(NO3)3 c. Cu3N2 h. Fe3N2 m. P2O3 d. NH3 i. BF3 n. (NH4)2CO3 e. LiOH j. Al2O3 o. (NH4)3PO4 30. Practice: Naming compounds a. KOH potassium hydroxide b. LiMnO4 lithium permanganate c. Mg(NO3)2 magnesium nitrate d. Rb2SO4 rubidium sulfate e. RaCl2 radium chloride f. BeS beryllium sulfide g. N2O3 dinitrogen trioxide h. Cu(NO3)2 copper(II) nitrate i. N2O5 dinitrogen pentoxide j. CS2 carbon disulfide k. (NH4)2SO4 ammonium sulfate l. F2O5 difluorine pentoxide m. Zn(ClO3)2 zinc chlorate n. Fe2(CO3)3 iron(III) carbonate o. Fe(ClO3)3 iron(III) chlorate 31. Groups or families: vertical columns on the periodic table 32. Fill in the following table: Representative Elements Property Number of V.E. Oxidation Number Family Name Gain or lose electrons 1A 2A 3A 4A 5A 6A 7A 8A 1 2 3 4 5 6 7 8 +1 +2 +3 +/-4 -3 -2 -1 0 Alkali metals Alkaline earth metals n/a n/a n/a n/a Halogens Noble Gases Lose 1 Lose 2 Lose 3 Lose/gain 4 Gain 3 Gain 2 Gain 1 Already stable 33. Reactivity increases as you go down within a group for metals and decreases for nonmetals. 34. Periods: horizontal rows on the periodic table. 35. Classify elements as metals, nonmetals, and metalloids based on location. Metals are located left of the stair-step line, nonmetals are to the right of the stair-step line, and metalloids are located along the stair-step line. 36. Classify elements as representative elements and transition elements. Representative elements are the A-group elements. Transition elements are the B-group elements. 37. Atomic radius increases as you go down a group, and decreases as you go across the period left to right. 38. The ionic radius of an anion is larger than the atomic radius of its neutral atom. The ionic radius of a cation is smaller than the atomic radius of its neutral atom. 39. Practice: Electron configurations a. Write the electron configurations of the following elements: i. Sodium 1s22s22p63s1 ii. Bromine 1s22s22p63s23p64s23d104p5 iii. Neptunium [Rn]7s25d5 iv. Silver [Kr]5s24d9 v. Radium [Rn]7s2 vi. Iron 1s22s22p63s23p64s23d6 vii. Barium [Xe]6s2 viii. Cobalt 1s22s22p63s23p64s23d7 ix. Tellurium [Kr]5s24d105p4 x. Lawrencium [Rn]7s25f146d1 b. Determine what elements are denoted by the following electron configurations: i. 1s22s22p63s23p4 sulfur 2 2 6 2 6 2 10 6 1 ii. 1s 2s 2p 3s 3p 4s 3d 4p 5s rubidium iii. [Kr]5s24d105p3 antimony 2 14 6 iv. [Xe]6s 4f 5d osmium v. [Rn]7s25f11 einsteinium 40. Identify the s, p, d, and f blocks on the periodic table. 41. Ionization energy decreases as you go down a group, and increases as you go across the period left to right. 42. Electronegativity decreases as you go down a group, and increases as you go across the period left to right.
