Chemistry I Honors Semester 1 Exam Study Guide PDF

# Chemistry I Honors Semester I Exam Study Guide ## Exam Details: - 100 Multiple Choice Questions (Scantron) - 1 hour and 45 minutes - Periodic Table and VSEPR Table will be provided - You may use a calculator and a 3x5 notecard, front and back, hand generated. The notecard will be turned in upon...

# Chemistry I Honors Semester I Exam Study Guide ## Exam Details: - 100 Multiple Choice Questions (Scantron) - 1 hour and 45 minutes - Periodic Table and VSEPR Table will be provided - You may use a calculator and a 3x5 notecard, front and back, hand generated. The notecard will be turned in upon completion of the exam. ## Study Guide Resources: - Physical/Chemical Notepacket - Atomic Structure Notepacket - Periodicity Notepacket - Bonding Notepacket - Phase/IMF Notepacket - Nomenclature Notepacket - Supplementary Packets ## Exam Content: 1. **Endothermic and Exothermic Reactions:** - **Endothermic:** Energy is absorbed. - **Exothermic:** Energy is released. 2. **Elements:** Substances which cannot be further decomposed by ordinary chemical means. 3. **Physical Changes of Matter:** Examples include state of matter change, cutting, crushing, dissolving. 4. **Chemical Changes of Matter:** Examples include decomposition, rusting, corrosion, combustion, and neutralization. 5. **Atomic Composition:** All atoms of the same element have the same number of protons. 6. **Isotope Composition:** The Chlorine-35 isotope has 17 protons, 18 neutrons, and 17 electrons. 7. **Electron Configuration:** The 5d¹⁰ sublevel has 10 electrons with energy level 5, sublevel d. 8. **Atomic Structure:** The major portion of an atom’s volume is empty space, often referred to as the electron cloud. 9. **Atomic Number:** An ion with 12 protons, 13 neutrons, and a charge of +2 has an atomic number of 12. 10. **Electron Energy Levels:** When electrons change to a higher level they absorb energy, and when they change to a lower level they emit energy. 11. **Electron Capacity:** The total number of electrons in a completely filled fifth principal energy level is 50. 12. **Ion Charge and Mass:** When an atom loses an electron, it becomes an ion with a positive charge and little change in its mass. 13. **Electron Spin:** When occupying the same orbital, two electrons must have opposite spins. 14. **Quantum Theory:** - Electrons form clouds based on probability of location. - Electrons occupy the lowest energy levels before moving into higher energy levels. - Electronic clouds form characteristic shapes due to repelling negative charges. 15. **Indications of Chemical Changes:** Color change, heat transfer, gas production, and precipitate formation are all indicators of chemical changes. 16. **Magnesium Ion:** An Mg⁺² magnesium ion has 10 electrons. 17. **Weighted Atomic Mass:** The weighted atomic mass of element X with isotopes X-79 (50.69% abundance, 78.918 amu) and X-81 (49.31% abundance, 80.917 amu) is 79.9 amu. The element is Bromine. 18. **Hund’s Rule:** The orbital diagram shown for the element Phosphorus violates Hund's Rule, which states that orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron. 19. **Orbital Occupancy:** A single atomic orbital can contain 0, 1, or 2 electrons. 20. **Nuclear Properties:** The nucleus of an atom has a high density, contains the mass of the atom, is positively charged, and has large amounts of energy that holds the nucleus together. 21. **Periodic Trends:** Elements in the same group have the same properties and valence electrons. 22. **Most Active Nonmetal:** Fluorine (F) is the most active nonmetal. 23. **Group 15 Trends:** The elements in Group 15 progress from nonmetals to metalloids to metals. 24. **Group Similarity:** Elements with the most similar chemical properties belong to the same group (column) on the periodic table. 25. **Electronegativity Trend:** Electronegativity increases across the periodic table, with the exception of the noble gases. 26. **Least Tendency to Lose Electrons:** The element in Period 4 with the least tendency to lose an electron is a nonmetal. 27. **Nonmetal Properties:** Nonmetals are solid, brittle, and have a low melting point. An example is Bromine (Br). 28. **Chemically Inactive Elements:** Elements that have high ionization energy and tend to be chemically inactive are classified as noble gases. 29. **Valence Electrons:** The number of valence electrons in an atom with an electron configuration of 1s²2s²2p⁶3s²3p⁵ is 7. 30. **Periodic Table Organization:** The present periodic table is organized based on increasing atomic number and the number of protons. 31. **Electron Configuration Period:** The element with the electron configuration 1s²2s²2p⁶3s²3p⁴4s²3d¹⁰4p⁶5s²4d¹⁰5p² is in period 5. 32. **Element in Period 4, Group 6:** The element found in Period 4, Group 6 is Chromium (Cr). 33. **Reactive Elements:** Groups 1, 2, and 17 on the periodic table are too reactive to be found as monoatomic atoms in nature. 34. **Nonmetal Properties:** Nonmetals have high ionization energy and low electrical conductivity. 35. **Group Labels:** A: Alkali metals, B: Alkaline Earth Metals, C: Halogens, D: Noble gases. 36. **Valence Electrons:** A: 1 valence electron, B: 2 valence electrons, C: 7 valence elections, D: 8 valence electrons 37. **Electron Configurations:** A and B: The final electron configuration ends with s². C and D: the final electron configuration ends with p⁶. 38. **Metal-Nonmetal Reactions:** When a metal reacts with a nonmetal, the metal will lose electrons and form a positive ion. 39. **Least Ionic Character:** Nonmetals have the least ionic character. 40. **Carbon Bonding:** Carbon can form 4 covalent bonds. 41. **Electronegativity:** Se is more electronegative than Te, meaning it attracts bonding electrons more strongly. 42. **Molecular Shape:** The major factors that determine a molecule’s shape are repulsive forces between the outer levels of electron clouds. ## Ionic Compounds and Molecules 43. **Ionic Compounds and Molecules:** - Co: molecule, K2S: ionic compound, CuSO4: Ionic compound, FeBr3: ionic compound. 44. **Electronegativity Difference:** When the electronegativity difference between bonded atoms is greater than 1.70, the substance is classified as ionic. 45. **Shared Electrons:** There are 4 shared electrons between the carbon atoms in H2-C=C-H2. 46. **Stability:** Most atoms are stable with 8 outer electrons. 47. **Valence Electrons:** The total number of valence electrons in CCl4 is 32. 48. **Atom Bonding:** Atoms form bonds to become more stable and to lower their overall energy. 49. **Ionic Bond Properties:** Ionic bonds are electrostatic attraction between oppositely charged ions. 50. **Metallic Bonding:** Metallic bonding occurs between metals (any metal). 51. **Predominant Bond:** The predominant type of bond formed between atoms with ground state electron configurations 1s²2s² and 1s²2s²2p⁶3s²3p⁶ is an ionic bond. 52. **Polar Molecules:** Uneven sharing of electrons or nonsymmetrical molecular geometry are factors that cause polar molecules. 53. **SO₂ Structure:** The VSEPR electron arrangement and molecular geometry for SO₂ is bent, with two bonding pairs and one lone pair. 54. **CH₂O Structure:** There is one double bond in the Lewis dot structure for CH₂O. 55. **Molecular Geometry:** - H₂O: bent - SO₃: trigonal planar - O₃: bent - PCl₃: trigonal pyramidal 56. **Multiple Bonds:** Multiple bonds are required when all electrons have been distributed, but not all atoms have an octet, or 8 electrons. 57. **Phase Diagram:** - **A:** Triple point, where all phases occur at once. - **B:** Critical point, above which gas cannot be condensed down to a liquid. - **C → D:** Sublimation - **D → C:** Deposition 58. **Evaporation:** The escape of gas molecules from the surface of an uncontained liquid is called evaporation. 59. **Hydrogen Bonding:** Hydrogen bonding occurs between polar molecules, such as H-F, H-N, and H-O. 60. **AgNO₃:** The attractive force that holds AgNO₃ together is ionic, or electrostatic forces. 61. **Strongest Van der Waals Force:** The strongest of the Van der Waals forces is hydrogen bonding. 62. **Intermolecular Forces:** - **Hydrogen bonding:** Occurs between H and N,O, or F. - **Dipole-dipole interactions:** Occurs between polar molecules. - **London dispersion forces (LDF):** Occurs between nonpolar molecules 63. **Intermolecular and Intramolecular Forces:** Intermolecular forces are overcome when ice melts, and intramolecular forces are overcome to allow silver to tarnish in a chemical reaction. 64. **Heating Curve:** - **1:** Solid - **2:** Melting - **3:** Liquid - **4:** Boiling - **5:** Gas 65. **Kinetic Energy Changes:** Segments 1, 3, and 5 on the heating curve represent changes in kinetic energy. 66. **Potential Energy Changes:** Segments 2 and 4 on the heating curve represent changes in potential energy. 67. **Boiling Point:** The normal boiling point of the substance in the phase diagram is approximately 100 °C. 68. **Pressure Change:** If the pressure of the substance is increased from 0.250 atm to 0.75 atm at 60 °C, condensation will occur. 69. **Critical Point:** The critical point on the phase diagram is where "e" is located. 70. **Common Element:** Zinc perchlorate [Zn(ClO₄)₂] and aluminum nitrate [Al(NO₃)₃] both contain oxygen. 71. **Multiple Oxidation States:** The elements with multiple positive oxidation states include Iron (Fe), Copper (Cu), Cobalt (Co), Tin (Sn), Lead (Pb), Mercury (Hg), Manganese (Mn), and Chromium (Cr). 72. **Mercury Oxidation State:** The oxidation state of mercury in HgS is +1. 73. **Chromium Ion Charge:** The charge on each chromium ion in Cr₂O₃ is +3. 74. **Ionic Compounds:** When combined with iodine, metals would most likely form ionic compounds. 75. **Halogen Oxidation Trend:** Halogens in Group 17 of the periodic table tend to gain 1 electron and form a -1 charge. 76. **Binary Molecule Naming:** The number of atoms of the second element in a binary molecule is indicated using prefixes. 77. **Stock System:** The Stock System uses Roman numerals in parentheses to indicate the charge of a transition metal ion. 78. **-ite and -ate Endings:** The -ite and -ate endings in compound names indicate the presence of a polyatomic ion, which is an ion with two or more atoms. 79. **Anions:** Anions are negatively charged ions. 80. **Binary Molecular Compounds:** Binary molecular compounds are compounds that contain two elements. Examples include: CO₂, FeO, NaCl, and CH₃OH. 81. **Neutral Compounds:** The total oxidation of a neutral compound formed from ions must equal zero. 82. **Polyatomic Ions:** Polyatomic ions of chlorate and chlorite differ in the number of oxygen atoms. 83. **Naming Compounds:** Ionic compounds and molecules often use the suffix "ide" when the nonmetal is the second listed element, i.e. oxide (O^-2). 84. **Compound Formula:** If a compound has the formula X₃Y and there 15 X atoms, there will be 45 Y atoms. 85. **Compound Names:** - NH₃: Ammonia - HCIO₃: Chloric acid - ICI: Iodine monochloride - Sn₃(PO₄)₂: Tin (II) phosphate 86. **Chemical Formulas:** - Aluminum cyanide: Al(CN)₃ - Plumbic chromate: Pb(CrO₄)₂ - Hydrofluoric acid: HF - Sulfuric acid: H₂SO₄ - Diphosphorus tetroxide: P₂O₄ 87. **Testing Fumes:** When testing the odor of fumes, gently wave the fumes toward your nose and softly inhale. 88. **Chemical Spills:** If you accidentally obtain too much of a specific chemical, dispose of it rather than contaminating the original stock. 89. **Safety Symbols:** Review the safety symbols to identify the type of hazard for each symbol. 90. **Lab Equipment:** Review lab equipment to identify the name of pictures of lab equipment.