1 Unit 1 Measurement, Matter and Change Part one: Scientific Measurement 1. Apply the rules of safety in the Chemistry laboratory 2. Distinguish between quantitative and qualitative measurements 3. Distinguish between precision and accuracy. 4. Distinguish between an observation and an inference 5. Write measurements in scientific notation 6. Identify the number of significant figures in a measurement. 7. Apply mathematical rules for significant figures to report measurements correctly. 8. List and define SI units of measurement. Also, know the prefixes used in Chemistry (Giga through nano) 9. Convert one measurement of unit to another using the factor label method. 10. Calculate and compare the densities of solid s, liquids and gases. 11. Graph experimental data (mass, volume) to determine the density and identify a pure substance. 12. Distinguish between an independent and dependent variable in an experiment and on an XY graph 13. Calculate the percent error of an experimentally determined measurement Part Two: States of Matter 1. Characterize the three states of matter by their particle arrangement. 2. List the physical properties of the states of matter; identify physical changes of matter. 3. List the chemical properties of matter and be able to identify chemical changes in matter. 4. Compare and contrast physical and chemical changes in matter. 5. Determine the temperature of phase changes from warming/cooling curves 6. Classify matter as a pure substance (element or compound), or a mixture (homogenous and heterogeneous) 7. Distinguish between an element and a compound. 8. Distinguish between a homogenous and heterogeneous mixture. 9. Compare and contrast the methods of separating a mixture (filtration, distillation, - simple and fractional, chromatography) and decomposing a compound. 10. Distinguish the symbols of common elements, and match the names of common elements to their symbols. 11. State the Law of Conservation of Mass and Energy and explain their roles in chemical and physical changes. 12. Distinguish between exothermic and endothermic chemical and physical changes
2 Unit 2- The Mole and Chemical Equations 1. Identify reactants and products in a chemical reaction. 2. Balance chemical equations when given the formulas for all reactants and products in a chemical reaction. 3. Identify substances that are commonly measured by: count, mass, and volume. 4. Describe how Avogadro s number is related to a mole of any substance. 5. Convert between number of particles and moles using the factor-label method. 6. Identify the diatomic molecules: Br, I, N Cl H O F (Brinklehoff) 7. Determine the atomic mass, molar mass, and formula mass of atoms and compounds. 8. Convert between moles and mass of a substance using the factor label method. 9. Convert between moles and volume of a gas using the volume of one mole of any STP (22.4 L). 10. Convert among measurements of mass, volume, and number of particles using the mole. 11. Determine the density of a STP. Note: Should we save the percent comp, empirical and molecular formulas for stoich?
3 Unit 3- Phases of Matter Part ONE: Defining the Physical States of Matter 1. Define kinetic and potential energy. 2. Describe the arrangement of particles in solids, liquids and gases in terms of particle motion (i.e. their kinetic energy) 3. Define intermolecular force. (IM Force) 4. Describe how particle organization (their kinetic energy and IM forces) distinguishes solids from liquids and gases. 5. Define temperature as a measure of the kinetic of particles. 6. Explain the high surface tension, the high specific heat, the high heat of vaporization, the high boiling point and the low vapor pressure using the concept of hydrogen bonding (the strongest IM force.) Part TWO: Measuring Energy Changes When Matter Changes 1. Describe in words and with diagrams the changes that occur in melting, freezing, boiling and condensing. 2. Explain how heat is a form of energy and how heat changes accompany physical and chemical changes. 3. Distinguish between specific heat capacity and heat capacity and their SI units. 4. Describe heat changes in terms of a system and its surroundings. 5. Calculate the heat changes that accompany physical and chemical changes. 6. Measure and calculate energy changes using a calorimeter. 7. Construct equations that show the heat changes for physical and chemical processes. 8. Identify endothermic and exothermic changes by the heat energy term in a chemical equation.
4 Part THREE: Characteristics of Solids, Liquids, and Gases 1. Interpret a phase diagram of a substance at any given temperature and pressure. 2. Explain the vaporization of liquids using kinetic theory. 3. Describe what happens on a particle level at the boiling point of a liquid. 4. Determine the boiling point using a vapor pressure curve. 5. Compare and contrast the effect of intermolecular forces on the boiling point of substances using the data on a vapor pressure curve. 6. Define sublimation. 9.1 ENERGY CONVERSIONS 1. Describe the effects of adding energy to matter in terms of the motion of atoms and molecules, and the resulting phase changes. 2. Explain how energy is transferred by conduction, convection and radiation. 3. Describe energy transformations among heat, light, electricity and motion. Points of emphasis: solar energy 9.7 ENERGY RESERVOIRS 1. Explain how solar energy causes water to cycle through the major earth reservoirs. 2. Explain how internal energy of the Earth causes matter to cycle through the magma and the solid earth. Points of emphasis: Earth s internal & external sources of energy
5 Unit 4- Atomic Structure and Nuclear Chemistry 1. Describe the five models in the historical development of modern atomic theory (Dalton, Thomson, Rutherford, Bohr, and Quantum Mechanical Model) 2. Distinguish among protons, neutrons, and electrons in terms of their relative masses, charges, and location with respect to the nucleus. 3. Infer the number of protons, neutrons, and electrons using the atomic number and mass number of an element from the periodic table and symbol notation. 9.4 ATOMS TO MOLECULES Describe the general structure of the atom, and explain how the properties of the first 20 elements in the Periodic Table are related to their atomic structures. 4. Explain how isotopes of an element differ. 5. Explain why the atomic masses of elements are not whole numbers. 6. Calculate the average atomic mass of an element from isotope data. 7. Differentiate between an atom and an ion. 8. Determine ion charge given proton, neutron, and electron data. 9. Distinguish between isotopes and radioisotopes in terms of stability and radioactive decay. 10. Contrast the characteristics of alpha, beta, and gamma radiation during radioactive decay. 11. Use symbol notation for subatomic particles and particle radiation. 12. Write balanced nuclear equations for alpha and beta decay processes. 13. Compute the amount of radioisotope remaining at a given time using the halflife method. 14. Write equations to show how transuranium elements are synthesized by transmutation. 15. Distinguish between artificial and natural transmutation. 16. Write balanced nuclear equations for artificial transmutation. 17. Compare and contrast nuclear fusion and nuclear fission. 18. Calculate the energy released during a nuclear reaction using E = mc Contrast the simple nuclear fission chain reaction to those used in the fission bombs of World War II: what led to the development of the fission bombs; who was involved? 20. Compare simple nuclear fission with the controlled fission reactions that used to generate energy in a nuclear reactor. 21. Is nuclear energy a viable alternative to fossil fuel energy? What are the real and imagined dangers of a nuclear reactor? We ll determine the human elements of those reactor accidents that have occurred, how they could have been avoided, and appreciate their effects on society. 22. What is the future of nuclear power as an alternative fuel source? 9.3 GLOBAL ENERGY SOURCES 1. Explain how heat is used to generate electricity. 2. Describe the availability, current uses and environmental issues related to the use of fossil and nuclear fuels to produce electricity. 3. Describe the availability, current uses and environmental issues related to the use of hydrogen fuel cells, wind and solar energy to produce electricity. Points of emphasis: fossil fuels and electricity, nuclear fission
6 Unit 5: Electrons and the Periodic Table 1. Explain the atomic emission spectrum of an atom using Bohr s model of the atom. 2. Use flame tests to identify transition metal elements. Explain flame test emissions in terms of electron behavior. 3. Calculate the frequency and wavelength of electromagnetic radiation. (EMR) 4. Calculate the energy of a photon associated with a given wavelength or frequency of EMR. Describe the properties of the different types of EMR. 5. Calculate the frequency and wavelength of light emitted for a specific energy level transition in the hydrogen atom. Identify which electron transitions produce emission spectra lines in the Balmer, Lyman, and Paschen series. 6. Describe the Quantum Mechanical Model of the atom. Apply the Aufbau Principle, the Pauli Exclusion Principle, and Hund s Rule to write electron configurations and orbital diagrams of elements. 7. Identify the information provided by the four quantum numbers and determine the four unique quantum numbers for a given electron. 8. Describe the origins of the modern periodic table. Describe the organization of the periodic table (periods, groups, periodic law), and categorize the elements as noble gases, transition metals, inner transition metals, and representative elements (halogens, alkali metals, alkaline earth metals). 9. Contrast the physical and chemical properties of metals, non-metals and metalloids. Also, one should be able to locate them on the periodic table. 10. Explain the relationship between the electron configuration of an element, its position on the periodic table, and its chemical properties. 11. State the trends of properties of elements within periods and groups of the Periodic Table. Interpret the trend shown by atomic radii, ionic radius, electronegativities, electron affinity, and 1 st, 2 nd, and 3 rd ionization energies within the Periodic Table, including exceptions. 12. Determine the number of number of valence electrons and/or predict the stable ion formation by a representative element, using the periodic table. Use the periodic table to predict the charge of the ion formed by a representative element and write Lewis Dot structures for atoms and ions.
7 Unit 5B- Bonding 1. Describe the formation of a cation from an atom of a metallic element, using the octet rule and the importance of noble-gas electron configurations. 2. Describe the formation of an anion from an atom of a nonmetallic element. 3. List the characteristics of an ionic bond and explain the electrical conductivity of melted and of aqueous solutions of ionic compounds. 9.2 ENERGY and ELECTRICITY 1. Explain the relationship among voltage, current and resistance in a simple series circuit. 2. Explain how electricity is used to produce heat and light in incandescent bulbs and heating elements. 3. Describe the relationship between current and magnetism. Points of emphasis: magnets and electrical currents 4. Explain the physical properties of metals using the theory of metallic bonding. 5. Describe the formation of a covalent bond between two nonmetallic elements. 9.4 ATOMS TO MOLECULES Describe how atoms combine to form new substances by transferring electrons (ionic bonding) or sharing electrons (covalent bonding). 6. Describe double and triple covalent bonds and draw Lewis structures to represent covalent bond structures containing single, double, and triple bonds. Also, be able to draw Lewis Dot structures for exceptions to the octet rule. 7. Explain the formation of a coordinate covalent bond. 8. Explain the modern interpretation of resonance bonding. 9. Describe the molecular orbital theory of covalent bonding, including sigma and pi bonding. 10. Predict the shape, polarity and bond angles of molecules (VSEPR theory). 11. Describe the shape of simple molecules using orbital hybridization. 12. Identify bonds as ionic, polar covalent, non-polar covalent using Electronegativity values. 13. Identify examples, structure and properties of polar and non-polar molecules 14. Define bond dissociation energy.
8 15. Identify properties of molecular substances. 16. Identify an organic compound s functional group. (Alcohols, esters, ethers, aldehydes, ketones, carboxylic acids). Be able to draw their structures as well. 17. Explain physical property differences in a class of organic compounds based on their molecular structure. 9.5 CARBON MOLECULES 1. Explain how the structure of the carbon atom affects the type of bonds it forms in organic and inorganic molecules. 2. Explain the general formation and structure of carbon-based polymers, including synthetic polymers, such as polyethylene, and biopolymers, such as carbohydrate. 18. Describe the characteristics of the formation of a polymer. 9.6 MOLECULES, HEALTH, AND TECHNOLOGY 1. Explain how simple chemical monomers can be combined to create linear, branched, and/or cross-linked polymers. 2. Explain how the chemical structure of polymers affects their physical properties. Points of emphasis: petroleum and its products, plastics 19. Classify hydrocarbons as alkanes, alkenes, alkynes or aromatic and give examples of each. Be able to draw these structures as well.
9 Unit 6- Nomenclature 1. Distinguish among atoms, molecules and formula units. 2. Distinguish between ionic and molecular compounds. 3. Write chemical formulas and names of binary molecular compounds using Greek prefixes. 4. Explain how a compound obeys the Law of Definite Proportions. 5. Explain how two different compounds composed of the same elements obey the Law of Multiple Proportions. 6. Distinguish between an ion and a polyatomic ion. 7. Memorize the names, formulas, and charges of the common polyatomic ions. 8. Memorize the charges of common monoatomic ions. 9. Write chemical formulas for binary ionic compounds. 10. Name binary ionic compounds when given the chemical formula. 11. Identify by name and write the chemical formulas for ternary ionic compounds (with polyatomic ions). 12. Use the stock naming system (roman numerals) for naming ionic compounds when appropriate. 13. Identify by name and write formulas for common acids. 15. Name hydrocarbons using the IUPAC system and write the structural formula given its name. 16. Name simple functional group compounds using the IUPAC system.
10 Unit 7- Chemical Reaction Types 1. Write word equations from chemical equations and vice versa. 2. Distinguish between these five types of reactions: combination, decomposition, single replacement, double replacement, combustion of hydrocarbons, and acid/base reactions (with basic ph). 9.4 ATOMS TO MOLECULES Explain the chemical composition of acids and bases, and explain the change of ph in neutralization reactions. 9.5 CARBON MOLECULES Describe combustion reactions of hydrocarbons and their resulting byproducts. 9.8 HUMAN RESOURCES Explain how the release of sulfur dioxide (SO 2 ) into the atmosphere can form acid rain, and how acid rain affects water sources, organisms and human-made structures. Points of emphasis: Air pollution from CO 2, SO 2, NO x s 3. Predict the products and balance simple combination and decomposition reactions, including reactions of carbonates, chlorates, hydrogen peroxide and water. 4. Predict the products and balance single replacement reactions using activity series of metals. 5. Predict the products, identify the precipitate, and balance double replacement reactions using memorized solubility rules and gaseous decomposition products. 6. Predict the products and balance combustion reactions for hydrocarbons. 7. Identify the oxidizing and reducing agent in a redox reaction and give the characteristics of a redox reaction. 8. Compute the oxidation number of an atom of any element in a pure substance. 9. Define oxidation and reduction in terms of a change in oxidation number and identify atoms being oxidized or reduced in redox reactions. 10. Distinguish between redox and nonredox reactions. 11. Describe the characteristics of a substitution reaction. 12. Describe esterifiaction. 13. Describe reactions of esters and define saponification. 14. List the structures that are products of an addition reaction.
11 Unit 8- Acids and Bases Fundamentals 1. List the general properties of aqueous acids and bases. 2. Explain the difference between strong acid and bases and weak acids and bases and give examples. 3. Define and give examples of Arrhenius acids and bases. 4. Classify substances as acids or bases, and identify conjugate acid-base pairs in acid-base reactions using the Bronsted-Lowry Theory. 5. Explain why proton-transfer reactions favor the production of the weaker acid and the weaker base. 6. Define an amphoteric substance and give examples. 7. Classify a solution as neutral, acidic, or basic, given the hydronium-ion or hydroxide-ion concentration. 8. Calculate the ph or poh of a solution given the hydronium-ion or hydroxideion concentration. 9. Calculate the hydronium-ion or hydroxide-ion concentration given the ph or poh. 10. Define acid anhydride and basic anhydride and give examples. 11. Write equations for the reactions of acid anhydrides with water. 12. Write equations for the reaction of base anhydrides with water. 13. Classify substances as Lewis acids or bases. 14. Explain how the formation of Acid Rain occurs. Also, explain the environmental implications of this phenomenon and possible solutions to this issue.
12 Unit 9- Stoichiometry 1. Calculate the percent composition by mass of a substance from its chemical formula or experimental data. 2. Derive the empirical formula of a compound from experimental data or percent composition data. 3. Derive the molecular formula of a compound from experimental data. 4. Construct mole ratios from balanced chemical equations. 5. Calculate stoichiometric quantities, using the factor-label method, from balanced chemical equations using units of moles, mass, number of particles, and volumes of STP. 6. Identify the limiting reagent in a chemical reaction and use it to calculate stoichiometric quantities and the amount of excess reagents. 7. Calculate the theoretical yield and percent yield for a given chemical reaction using experimental data.
13 Unit 10- Gas Laws 1. Describe the motion of particles of a gas according to the kinetic theory. 2. Explain gas pressure in terms of the kinetic theory. 3. Describe the design and function of a thermometer. 4. Describe the design and function of a barometer. 5. Read an open-ended and close-ended manometer to determine the pressure of a gas sample. 6. Define the temperature of a substance as a measure of the kinetic energy of the particles in the substance. 7. Boyle s Law: A) Calculate the pressure or volume from the pressure-volume relationship of a contained gas at constant temperature. 8. The Gay-Lusaac Law: A) Calculate the temperature or pressure from the temperature-pressure relationship of a contained gas at constant volume. 9. The Combined Gas Law: A) Calculate pressure, volume, or temperature from the pressure-volumetemperature relationship of a contained gas. B) Calculate the amount of gas at any specified conditions of pressure, volume, and temperature. 10. Describe the Ideal Gas Law: PV = nrt. Calculate pressure, temperature, volume, moles, grams, or molecules using the Ideal Gas Law/ 11. Using the Ideal Gas Law: A. Calculate the molar mass or density of a gas. B. Calculate the amount of gas at any specified conditions of pressure, volume, and temperature. 12. Calculate the total pressure of a mixture of gases or the partial pressure of a gas in a mixture of gases 13. Explain Avogadro s hypothesis using the kinetic theory. 14. Explain, using kinetic theory, why molecular of small mass diffuse more rapidly than molecules of large mass.
14 15. Use a chemical equation to specify volume ratios for gaseous reactants and/or products. 16. Use volume ratios, standard molar volume, and the gas laws where appropriate to calculate volumes, masses or molar amounts of reactants involving gases. 17. Distinguish between real and ideal gases. 18. Explain why no gas behaves as an ideal gas at all temperatures and pressures. 19. Describe and analyze the set up for the preparation of different gases as well as test for H 2, O 2, and CO Be able to calculate rates of diffusion and effusion of gases using Graham s Law.
15 Unit 11- Solutions and Thermodynamics Part ONE: Solutions 1. Define the terms solution, aqueous solution, solute and solvent and give an example of each. 2. Describe the role of solvation in the dissolving process and use the rule like dissolves like to predict solubility of one substance in another. 3. Distinguish colloids and suspensions from solutions by discussing their properties. 4. List three factors that determine how fast a soluble substance dissolves. 5. Explain the difference among saturated, unsaturated, and supersaturated solutions. 6. Apply information provided by a solubility curve. 7. Calculate the molarity of a solution (Review from Unit 2). 8. Determine the number of moles or grams of solute given a molar solution. 9. Prepare dilute solutions of given concentrations from concentrated solutions of known molarity using appropriate calculations. 10. Calculate percent by mass, and percent by volume for solutions. 11. What are colligative properties? Explain on a particle basis why a solution has a lower vapor pressure, an elevated boiling point, and a depressed freezing point than the pure solvent of that solution. 12. Be able to calculate molality and mole fraction of a solution. 13. Calculate the freezing point depression and boiling point elevation of aqueous solutions. 14. Calculate molecular mass of an unknown from experimental freezing point depression or boiling point elevation measurements.
16 Part TWO: Thermodynamics 1. Calculate the enthalpy change for a reaction for varying amounts of a reactant. 2. Write a balanced standard formation reaction for a compound. 3. Review calculations for heats of physical change of substances 4. Define Heat of Solution. 5. Apply Hess s Law of Summation to find heat changes for chemical and physical processes. 6. Apply Standard Heats of Formation ( H f ) to find enthalpy changes for chemical and physical processes. 7. Apply average bond energies to find enthalpy changes for chemical and physical processes. 8. Define free energy. 9. Contrast spontaneous and non-spontaneous reactions. 10.Be able to draw potential energy graphs to signify exothermic and endothermic processes. 11. Show how changes in entropy relate to a change of state, a change in temperature, and a change in the number of product particles compare with reactant particles. 12. Apply Standard Entropy values ( S o ) to calculate the entropy change of chemical and physical processes. 13. Apply Gibbs Free Energy Equation to explain how changes in enthalpy and entropy influence the spontaneity of a reaction.
17 Unit 12- Kinetics, Equilibrium and Advanced Acid/Base Chemistry 1. Interpret and express the meaning of the rate of a chemical reaction. 2. Explain how the rate of a chemical reaction is influenced by the temperature, concentration, particle size of the reactants, and catalysts using collision theory. 3. Define a reaction mechanism. 4. Explain how the rate-determining step of the reaction mechanism affects the overall rate of the reaction. 5. Write the rate law fro a reaction given experimental data. 6. Define chemical equilibrium in terms of a reversible reaction. 7. Explain the process of reaching equilibrium. 8. Identify chemical reactions that go to equilibrium and those that go to completion. 9.Explain the nature of the equilibrium constant. 10. Write an equilibrium constant expression for a reaction and complete its value from experimental data. 11. Determine initial and final equilibrium concentrations for any reaction at equilibrium. 12. Predict changes in the equilibrium position due to changes in concentration, temperature, and pressure using Le Chatelier s principle. 13. Calculate ion concentrations of slightly soluble salts using the solubility product constant. 14. Explain the common-ion effect using Le Chatelier s principle. 15. Calculate an acid dissociation constant K a, and a base dissociation constant, K b, from concentration and ph measurements of weak acids and bases. 16. Chose the best indicator for a given acid-base titration. 17. Describe the process of salt hydrolysis and calculate the ph of a salt solution. 18. Define a buffer and, using equations, show how a buffer system works. 19. Use the Henderson-Hasselbach Equation to calculate ph of a buffer system.
18 Unit 13- Electrochemistry 1. Apply the oxidation-number change method to balance redox reactions 2. Apply the half-reaction method to balance redox equations. 3. Describe the production of electric current in an electrochemical (galvanic) cell. 4. Explain a voltaic cell using a sketch and labeling he anode, cathode, and direction of electron and ion flow. 5. Identify the chemical reactions in an electrochemical (galvanic) cell. 6. Define cell potential and describe how it is determined. 7. Define the standard electrode potential of an electrode. 8. Compute the standard emf of a cell using standard electrode potentials. 9. Distinguish between a voltaic and an electrolytic cell. List some examples of each. 10. Describe the chemical reactions in an electrolytic cell. 11. Explain an electrolytic cell using a sketch and labeling the anode, cathode, electric charges on electrodes, and direction of electron and ion flow. 12. Identify the chemical reactions in an electrolytic cell. 13. Explain the process of corrosion. 14. Describe how commercial cells produce an electric current. 15. Calculate the amount of product in an electrolytic cell based on the current (amperes) or charge (coulomb, Faraday) involved.
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1. The average kinetic energy of water molecules increases when 1) H 2 O(s) changes to H 2 O( ) at 0ºC 3) H 2 O( ) at 10ºC changes to H 2 O( ) at 20ºC 2) H 2 O( ) changes to H 2 O(s) at 0ºC 4) H 2 O( )
CHEM 120 Online Chapter 7 Date: 1. Which of the following statements is not a part of kinetic molecular theory? A) Matter is composed of particles that are in constant motion. B) Particle velocity increases
Practice Multiple Choice Questions: 1) Which of the following is NOT a laboratory safety rule? a) You should never mix acids with bases b) You should tie back your long hair c) You should never add water
Content Outline for Physical Sciences Section of the MCAT GENERAL CHEMISTRY ELECTRONIC STRUCTURE AND PERIODIC TABLE A. Electronic Structure 1. Orbital structure of hydrogen atom, principal quantum number
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Chapter 2 The Chemical Context of Life Multiple-Choice Questions 1) About 25 of the 92 natural elements are known to be essential to life. Which four of these 25 elements make up approximately 96% of living
P.S./CEMISTRY The University of the State of New York REGENTS IG SCOOL EXAMINATION PYSICAL SETTING CEMISTRY Wednesday, January 28, 2015 1:15 to 4:15 p.m., only The possession or use of any communications
EXPERIMENT 1: Survival Organic Chemistry: Molecular Models Introduction: The goal in this laboratory experience is for you to easily and quickly move between empirical formulas, molecular formulas, condensed
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Acids and : A Brief Review Acids: taste sour and cause dyes to change color. : taste bitter and feel soapy. Arrhenius: acids increase [H ] bases increase [OH ] in solution. Arrhenius: acid base salt water.
Acids and Bases Know the definition of Arrhenius, Bronsted-Lowry, and Lewis acid and base. Autoionization of Water Since we will be dealing with aqueous acid and base solution, first we must examine the
DP Chemistry Review Topic 1: Quantitative chemistry 1.1 The mole concept and Avogadro s constant Assessment statement Apply the mole concept to substances. Determine the number of particles and the amount
Name: 1) Which molecule is nonpolar and has a symmetrical shape? A) NH3 B) H2O C) HCl D) CH4 7222-1 - Page 1 2) When ammonium chloride crystals are dissolved in water, the temperature of the water decreases.
P.S./EMISTRY The University of the State of New York REGENTS IG SOOL EXAMINATION PYSIAL SETTING EMISTRY Wednesday, January 29, 2014 1:15 to 4:15 p.m., only The possession or use of any communications device
Chemical Principles As Applied to Soils I. Chemical units a. Moles and Avogadro s number The numbers of atoms, ions or molecules are important in chemical reactions because the number, rather than mass
5s Solubility & Conductivity OBJECTIVES To explore the relationship between the structures of common household substances and the kinds of solvents in which they dissolve. To demonstrate the ionic nature
PS-3.1 Distinguish chemical properties of matter (including reactivity) from physical properties of matter (including boiling point, freezing/melting point, density [with density calculations], solubility,
The Advanced Placement Examination in Chemistry Part I Multiple Choice Questions Part II Free Response Questions Selected Questions from1970 to 2010 Atomic Theory and Periodicity Part I 1984 1. Which of
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Freezing Point Depression: Why Don t Oceans Freeze? Teacher Advanced Version Freezing point depression describes the process where the temperature at which a liquid freezes is lowered by adding another
HOMEWORK 5A Barometer; Boyle s Law 1. The pressure of the first two gases below is determined with a manometer that is filled with mercury (density = 13.6 g/ml). The pressure of the last two gases below
CHAPTER 10 REVIEW States of Matter SECTION 1 SHORT ANSWER Answer the following questions in the space provided. 1. Identify whether the descriptions below describe an ideal gas or a real gas. ideal gas
Chemistry 151 Final Exam Name: SSN: Exam Rules & Guidelines Show your work. No credit will be given for an answer unless your work is shown. Indicate your answer with a box or a circle. All paperwork must
Basic Chemistry Why do we study chemistry in a biology course? All living organisms are composed of chemicals. To understand life, we must understand the structure, function, and properties of the chemicals
Chapter 13 - LIQUIDS AND SOLIDS Problems to try at end of chapter: Answers in Appendix I: 1,3,5,7b,9b,15,17,23,25,29,31,33,45,49,51,53,61 13.1 Properties of Liquids 1. Liquids take the shape of their container,
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