MIDTERM STUDY GUIDE. Chapter 1 Introduction to Chemistry

Transcription

1 MIDTERM STUDY GUIDE Chapter 1 Introduction to Chemistry What is chemistry? Chemical properties vs. physical properties examples of both States of matter Scientific method Chapter 2 Data Analysis SI measurement Quantity, units, symbols for measurements Density Conversion factors Accuracy vs. precision Quantitative vs. qualitative % Error Scientific Notation Chapter 3 Matter-Properties and Change Classification of matter mixtures, pure substance, heterogeneous vs. homogeneous Separation techniques Chemical properties vs. physical properties examples of both Chemical changes vs. physical changes examples of both Law of conservation of mass/matter Law of definite proportions law of multiple proportions Chapter 4 The Structure of the Atom Atomic theories Democritus, Dalton, Thomson, Rutherford Composition of the atom Protons, neutrons, electrons, nuclear forces Atomic number Mass number Average weighted atomic mass Isotopes -calculating mass number -Isotope nomenclature

2 H-Chemistry REVIEW: Mid-Year Exam Page 2 Chapter 25 Nuclear Chemistry Identify radioactive particles and radiation in term of composition and key properties Explain why certain nuclei are radioactive Write and complete balanced nuclear equations Determine if transmutation is natural or artificial Solve problems involving radioactive decay rate - half life - know how to use equation to solve problems with fractions of half life cycles Recognize, Compare and Contrast Nuclear Fission and Fusion Describe several uses of radioactivity in daily life and identify radioisotopes used in medicine and dating. Chapter 5 Electrons in Atoms Electromagnetic radiation Relationship between wavelength, frequency, speed of light, and energy of waves Energy states of atoms - Explain the formation of bright-line spectrum Bohr model of the atom Hydrogen, orbits, energy levels Quantum model of the atom Heisenberg uncertainty principle Electron configurations Aufbau rule, Hunds rule, Pauli exclusion rule Principle Energy levels 1, 2, 3, Orbitals (sublevels) s, p, d, f Orientation - shape Spin Notations Electron configuration Orbital notation Noble gas notation Chapters 6 and 7 The Periodic Table and the Periodic Law/The Elements Mendeleev, Moseley Periodic law Characteristics of the Groups Alkali metals, alkaline earth metals, transition metals, halogens, noble gases, inner transition metals - lanthanide series, actinide series Characteristics of s-block, p-block, d-block, f-block Metals vs. nonmetals Group and Period Trends in: Electronegativity Successive Ionization energy Ions and Iso-electronic species Valence electrons Atomic radii Ionic Radii

3 H-Chemistry REVIEW: Mid-Year Exam Page 3 Chapter 8 Ionic Compounds Naming/Formula writing: Ionic Compounds- Binary vs Ternary Metals & Non-metals Cations & anions & polyatomic ions Criss-Crossover rule for ionic charge Naming rules -NO prefixes -1 st name same as element -2 nd name drop ending and add ide Chemical bonds & characteristics Octet rule Ionic bonds Formula unit, Lewis-dot diagrams Metallic bonds Sea of mobile electrons Conductors, malleable, ductile Chapter 9 Covalent Bonding Chemical bonds & characteristics Octet rule Covalent bonds Molecule,, polar vs non polar bonds Lewis dot structures of covalent atoms VSEPR (molecular shapes and bond angles) Naming/Formula writing: Molecules- ONLY non-metals Always use prefixes -except when 1 st element is a single atom Memorize prefix names (1 10) Types of compounds and their properties Intermolecular forces of Attraction Exam consists of 2 parts multiple choice questions using Scantron o Free response Possible free response questions o Lewis dot structures for Molecules o VSEPR problems o Calculate average atomic mass o Electron configurations

4 H-Chemistry REVIEW: Mid-Year Exam Page 4 Electron configuration Orbital notation Noble gas notation o Wave calculations o Half-life calculations o Explain the difference between a physical and a chemical change o Explain the how atomic radius and/or ionic radius changes o Explain how elements give off light according to atomic theory (remember the pickle demo) o % error calculations o Explain accuracy vs. precision o Explain why certain groups are reactive and other groups are not. You will be provided with a Reference Table. You must bring a calculator to the exam. You must bring at least one pencil and one pen

5 H-Chemistry REVIEW: Mid-Year Exam Page 5 Honors Chemistry MID-YEAR EXAMINATION PRACTICE REVIEW Answer the following questions on a loose-leaf paper. (Copy tables). Use your neatest hand-writing. (Remember that If I cannot read it it's wrong) 1. Define chemistry: 2. Identify the quantity, unit and abbreviation represented by each of the following: Quantity Unit Abbrev. Quantity Unit Abbrev. mass liters joules mol grams/milliliter seconds K molar mass 3. Fill in the following table with either definite or indefinite in each space: State of Matter Property Solid Liquid Gas volume shape 4. Physical/Chemical Changes: A. Define physical change: B. Define chemical change: C. Identify each as a chemical or physical change, and briefly explain. i. paper burning: ii. iron rusting: iii. ice melting: iv. leaves disintegrating in the autumn: v. copper roof turning from green to orange: 5. Define matter:

6 H-Chemistry REVIEW: Mid-Year Exam Page 6 6. Define and give an example of each: A. atom: B. compound: C. mixture: D. heterogeneous: E. pure substance: F. which of the following is(are) a pure substance: element compound mixture solution 7. Define: A. quantitative data: B. qualitative data: 8. Give one example of each: A. direct relationship: B. inverse relationship: 9. Convert: A. 25 o C to K: B. 293 g/cm 3 to kg/l: 10. Express 639,000,000 in scientific notation: 11. Periodic Table: Label metals, nonmetals and metalloids. Label s-, p-, d- and f-blocks Name the group identified by each letter: A. B. or C. D. or E. F. G. H. I. J.

7 H-Chemistry REVIEW: Mid-Year Exam Page 7 Define the law of conservation of mass: 12. According to the law of conservation mass, what is the relationship between the masses of the reactant(s) and product(s) in a chemical reaction? 13. Identify the contribution of each to science: A. Democritus: B. Dalton: C. Rutherford: D. Thomson: E. Moseley: F. Mendeleev 14. Most of an atom is occupied by:. Describe the experiment that led to this conclusion 15. What are the three subatomic particles comprising an atom? 16. Which one determines the identity of the atom? 17. Define mass number: 18. Define atomic number: 19. Fill in the following spaces for the properties of the subatomic particles: Property Neutron Proton Electron symbol location relative mass charge 20. Define isotope: 21. Fill in the following spaces for isotopes: Element Isotope Atomic Number Mass Number Number Protons Number Neutrons Number Electrons Charge 16 8 O

8 H-Chemistry REVIEW: Mid-Year Exam Page Name and describe the four radioactive particles and energy. 23. What is transmutation? What is the difference between artificial and natural transmutation? Give examples. 24. Give examples for problems involving half-life calculations : Determine amount in future, amount in past, Determine time elapsed, Determine half-life. 25. What are the uses of radioactivity in real life. Give examples. 26. Name three types of electromagnetic energy: 27. What is the basic unit of electromagnetic energy? 28. What is a quantum of light called? 29. What problem was addressed by the development of the Bohr atom? 30. Draw a picture of the atom using the Bohr model. Identify the ground and excited states: 31. Describe the Bohr model of the atom: 32. Using the Bohr model of the atoms, describe how the bright line-emission spectrum of hydrogen is produced. 33. What does the Heisenberg uncertainty principle describe? 34. What is the difference between an orbit and an orbital? 35. What is the highest occupied energy level in a bromine atom? 36. Label the following orbital shapes. A. B. C. 37. Which is a higher energy level: 2s or 2p? 2s or 3s? 4s or 4d? 38. How many electrons are each of the following energy sublevels? s p d f 39. What is the orbital notation for zinc? 40. What is the orbital notation for yttrium? 41. What is the electron configuration for selenium? 42. What is the electron configuration for barium? 43. What is the noble gas configuration for iodine? 44. For what element is the noble gas configuration [Xe]4f 14 5d 10 6s 2 6p 6? 45. What are the horizontal rows on the periodic table called? 46. What are the vertical columns on the periodic table called?

9 H-Chemistry REVIEW: Mid-Year Exam Page Atomic radius: A. What is the trend for atomic radius as one goes from left to right across a period? B. Explain. C. What is the trend for atomic radius as one goes from top to bottom in a group? D. Explain. 48. Define valence electrons. 49. Define octet rule. 50. Why is it called a periodic table? 51. What is the most probable ion for each of the following? A. calcium: B. phosphorus: C. iodine D. oxygen 52. What is electronegativity? 53. What is the relationship between atomic radius and electronegativity? 54. List five properties of metals: 55. What are properties of each group? A. alkali metals: B. alkaline earth metals: C. halogens: D. noble gases: 56. What two subatomic particles are involved in a chemical bond? 57. What are the diatomic molecules? 58. What type of bond always connects two atoms in a diatomic molecule? 59. What type of bond is between each of the following pairs of atoms? A. carbon and hydrogen: B. nitrogen and phosphorus: C. lithium and fluorine D. chlorine and iodine 60. How many valence electrons are in each of the following atoms? A. carbon: B. silicon C. sodium D. bismuth

10 H-Chemistry REVIEW: Mid-Year Exam Page What is the formula unit for the compound formed by calcium and phosphorus? 62. How many electrons are in a double bond? 63. What is a sigma bond? a pi bond? 64. What determine the shape of a molecule and the bond angel? Give examples for the major six shapes you need to know. 65. What is the difference between a nonpolar and polar covalent bonds? 66. What determine if a molecule has a dipole - polar molecule? Give examples. 67. Compare the differences in physical properties (including the ability to conduct electricity when either dissolved in water or in the molten state) of metals, nonmetals, molecules, network solids, and ionic compounds. 68. Which has the most energy a single, a double, or triple bond? 69. Draw Lewis structures for following molecules/ions/ionic formula units: A. CH 4 B. NH 3 C. CO 2 D. NO 3 - E. AlBr 3 F. Fe 2 O Formula units are used for: ionic compounds or molecules? 71. What is, and give an example of, a binary substance? _ 72. What is, and give an example of, a polyatomic ion? 73. What is the correct formula for the compound formed by magnesium and phosphate? 74. What is the correct formula for the compound formed by aluminum and fluorine? 75. Name Cu 2 SO Name Ba 3 (PO 4 ) Name N 2 O How many nitrogen atoms are in NH 4 NO 3? 79. What is the formula for magnesium acetate? 80. The formula for octane is C 8 H 18. This compound is a(n) molecule, ionic compound, or metal? 81. What does a subscript in a molecular formula indicate?