Chemistry Curriculum Map and Pacing Guide Semester

Transcription

1 Chemistry Curriculum Map and Pacing Guide Semester Quarter 1 (46 Days) Chapter pic (Textbook Section) Dates # Days 2 Measurement and Analysis Lab Safety (p. 19, Flinn Safety Information) SI Units and Prefixes, Density Calculations (2.1) 8-30 Scientific Notation and Dimensional Analysis (2.2) Accuracy and Precision of Experimental Data (2.3) Calculation of Percent Error of Experimental Data (2.3) Significant Figures (2.3) Graphing Data (2.4) 3 Properties of Matter Substances (3.1) Physical and Chemical Changes in Matter (3.2) Conservation of Mass (3.2) Mixtures and Separation of Mixtures (3.3) Elements and Compounds (3.4) Law of Definite Proportions (3.4) 4.1 to to 5.2 Atomic Structure Early Theories of Matter (4.1) The Nuclear Atom (4.2) Atomic Number, Isotopes, and Mass Number (4.3) Calculating Weighted Average Atomic Mass (4.3) Light and Quantized Energy (5.1) Emission and Absorption Spectra of Atoms (5.1) Quantum Mechanical Model of the Atom and Probability (5.2) 6, 5.3 Periodicity Development of the modern periodic table (6.1) Organization of elements in periodic table (6.2) Electron Configuration, Electron Dot and Orbital Diagrams (5.3) Periodic properties and trends (6.3) Quarter 1 Benchmark Chapters Quarter 2 (39 Days) 7 Ionic Compounds Formation of Ions (7.1) Ionic Bonds (7.2) Names and Formulas for Ionic Compounds (7.3) Metallic Bonds and Properties of Metals, Alloys (7.4) 8 Covalent Bonding Covalent Bonds (8.1) Naming Molecules (8.2) Lewis Structures (8.3) VSEPR (8.4) 9 Chemical Equations Balancing Chemical Reactions (9.1) Classifying Chemical Reactions (9.2) Semester 1 Exam Review and Exam Chapters

2 Chemistry Curriculum Map and Pacing Guide Semester 2 Quarter 3 (47 Days) Chapter pic (Textbook Section) Dates # Days 10 Mole Concept Avogadro s Number and the Mole (10.1) Interconvert Mass, Moles and Number of Particles ( ) 1-24 Calculation of Molar Mass of Substances (10.3) Empirical and Molecular Formulas of Substances (10.4) Hydrates (10.5) 11 Stoichiometry Quantitative Relationships of Balanced Chemical Reaction (11.1) Stoichiometric Calculations (11.2) Limiting Reactants (11.3) Percent Yield of a Chemical Reaction (11.4) 12 States of Matter Gases and the Kinetic-Molecular Theory (KMT) (12.1) Intermolecular Forces (12.2) Liquids and Solids and KMT (12.3) Phase Changes and Phase Diagrams (12.4) 13 Gas Laws Gas Laws (13.1) Combined Gas Law and Avogadro s Principle (13.2) The Ideal Gas Law: PV=nRT (13.3) Gas Stoichiometry (13.4) Quarter 4 (46 Days) 14 Solutions Types of Mixtures (14.1) Solution Concentration (14.2) Factors Affecting Solubility (14.3) 15 Thermochemistry Energy and Specific Heat (15.1) Calorimetry (15.2) Enthalpy of Fusion and Vaporization (15.3) 16 Chemical Rates Reaction Rate and Collision Theory (16.1) Nature of Reactants (16.2) 17 Chemical Equilibrium Dynamic Equilibrium (17.1) LeChatelier s Principle (17.2) Solubility Product Constant (17.3) 18 Acids and Bases Models and Properties of Acids and Bases (18.1) Strengths of Acids and Bases (18.2) ph and poh (18.3) Titration (18.4) ACT EOCE (Chapters 2-18) Semester 2 Exam Review and Exam (Chapters 10-18)

3 Chemistry Learning Targets Semester Quarter 1 Chapter 2: Measurement and Analysis Chemists collect and analyze data to determine how matter interacts. 1. Apply safety concepts in the chemistry classroom and laboratory. (p. 19, Flinn Safety Contract) 2. Explain why mass is a quantity of matter and differentiate between mass and weight. (p. 9) 3. Use appropriate SI Units for length, time, temperature, and quantity of matter. (p ) 4. Describe the relationships among SI prefixes (e.g. centi-, milli-, kilo-, nano-) (Table 2, p. 33) 5. Explain density quantitatively as D = m/v, and solve density problems. (p ) 6. Express numbers in scientific notation when appropriate (p ) 7. Use dimensional analysis to convert between measurement units. (p ) 8. Distinguish between precision and accuracy with respect to experimental data. (p ) 9. Calculate percent error when comparing experimental data with an accepted value. (p ) 10. Use the correct number of significant figures in reporting measurements and calculations. (p ) 11. Use graphical models to express relationships inferred from sets of scientific data. (p , Density Lab) Chapter 3: Properties of Matter Everything is made of matter 1. Identify the characteristics of a substance. (p ) 2. Explain the difference between chemical and physical properties and changes. (p ) 3. Apply the law of conservation of mass to chemical reactions. (p ) 4. Classify mixtures as homogeneous or heterogeneous (p ) 5. Describe and perform common separation techniques (e.g. filtration, decanting, distillation, chromatography) (p ) 6. Distinguish between elements and compounds (p ) 7. Classify matter using Figure 19 p. 87 as a guide. 8. Explain how all compounds obey the laws of definite and multiple proportions (p ) Chapters 4.1 to 4.3, 5.1 to 5.2: Atomic Structure Atoms are the fundamental building blocks of matter. 1. Compare and contrast the atomic models of Democritus, Aristotle, and Dalton. (p ) 2. Describe the crucial contributions of scientists and the critical experiments that led to the development of the modern atomic model and distinguish between protons, electrons, and neutrons (p ) 3. Use the periodic table to determine the atomic number, atomic mass, mass number, and number of protons, electrons, and neutrons in isotopes of elements. (p ) 4. Calculate the weighted average atomic mass of an element from isotopic abundance, given the atomic mass of each contributor. (p ) 5. Describe characteristics of a wave, such as wavelength, frequency, energy, and speed. (p ) 6. Summarize the concept of Wave-Particle Duality of light and matter. (p ) 7. Describe how the absorption spectra of a material is measured using a spectrophotometer. (p ) 8. Summarize the Quantum Mechanical Model of the Atom, and the role of probability in atomic orbitals (p ) 9. Describe atomic orbitals (s, p, d, and f) and their basic shapes. (p ) Chapters 6, 5.3: Periodicity Period trends in properties of atoms allow us to predict physical & chemical properties. 1. Describe the historical development of the modern periodic table, including work by Mendeleev and Moseley. (p ) 2. Describe and explain the organization of elements into periods and groups in the periodic table. (p. 177) 3. Identify regions (e.g., groups, families, and series) of the periodic table and describe the chemical characteristics of each. (p ) 4. Apply the Aufbau process, The Pauli exclusion principle, and Hund s rule to specify orbital diagrams, electron configuration, and electron dot diagrams of elements. (p ) 5. Use the periodic table to predict and explain the valence electron configuration of the elements, to identify configuration families, and to predict the common valences of the elements. (p ) 6. Compare trends in the periodic properties of the elements (e.g., metal/nonmetal/metalloid behavior, electrical/heat conductivity, electronegativity and electron affinity, ionization energy, atomic/covalent/ionic radius) and how they relate to position in the periodic table. (p )

4 Quarter 2 Chapter 7: Ionic Compounds They are held together by chemical bonds formed by attraction of oppositely charged ions. 1. Define chemical bond and relate chemical bond formation to electron configuration. (p ) 2. Use electron dot diagrams to describe the formation of positive and negative ions. (p ) 3. Describe the formation of ionic bonds. (p ) 4. Account for the physical properties of an ionic compound. (p ) 5. Discuss the energy involved in the formation of an ionic bond. (p ) 6. Name and write formulas for ionic compounds and polyatomic ions, using IUPAC symbols for elements. (p ) 7. Interpret the information conveyed by chemical formulas for the numbers of atoms of each element represented. (ACT QC Example) 8. Describe a metallic bond. (p. 225) 9. Explain the physical properties of metals in terms of metallic bonds. (p. 226) 10. Define and describe alloys. (p ) Chapter 8: Covalent Bonding Covalent bonds form when atoms share electrons. 1. Define molecule and the covalent bond. (p ) 2. Describe the formation of single, double and triple covalent bonds and relate them to sigma and pi bonds (p ) 3. Describe the unique features of bonding in carbon compounds. (p. 243) 4. Relate the strength of covalent bonds to bond length and bond dissociation energy. (p ) 5. Identify the names of binary molecular compounds and their formulas (p ) (Note: Acids will be done in Ch. 18) 6. Draw Lewis structures for molecules and polyatomic ions, including those that must be represented by a set of resonance structures. (p ) 7. Use VSEPR theory to explain geometries of molecules and polyatomic ions. (p ) 8. Determine how electronegativity is used to determine polar, nonpolar, and polar covalent bonds (p ) Chapter 9: Chemical Equations Rxn transform reactants into products, resulting in absorption or release of energy. 1. Explain how conservation laws form the basis for balancing chemical reactions. (p. 288) 2. Write and balance chemical equations, given names of reactants and products. (p ) 3. Classify chemical reactions as being synthesis, decomposition, combustion, single replacement, or double replacement reactions. (p ) 4. Predict products of single replacement reactions using the activity series (Metal/Metal Ion Reactions Lab Simulation) 1. Inquiry in Chemistry and Science in Practice (Used throughout the course) A. Identify and clarify research questions and design experiments. B. Design experiments so that variables are controlled and the appropriate number of trials is used. C. Collect, organize, and analyze data accurately and use techniques and equipment appropriately. D. Interpret results and draw conclusions, revising hypotheses as necessary and/or formulating additional questions or explanations. E. Safely use laboratory equipment and techniques when conducting scientific investigations. F. Routinely make predictions and estimations. G. Use appropriate statistical methods to represent the results of investigations. H. Explain and apply criteria that scientists use to evaluate the validity of scientific claims and theories. I. Explain why experimental replication and peer review are essential to eliminate as much error and bias as possible in scientific claims. J. Explain the criteria that explanations must meet to be considered scientific. K. Explain why all scientific knowledge is subject to change as new evidence becomes available. L. Use a variety of appropriate sources to retrieve relevant information. M. Cite references properly. N. Identify and analyze the advantages and disadvantages of widespread use and reliance on technology. O. Compare the scientific definitions of fact, law, and theory, and give examples of each in chemistry.

5 Chemistry Learning Targets Semester Quarter 3 Chapter 10: Mole Concept The mole allows us to count extremely small things, such as atoms, by massing them. 1. Explain the meaning of mole and Avogadro s number. (p ) 2. Define the gram atomic mass of an element. (p ) 3. Interconvert between mass, moles, and number of particles. (p ; ) 4. Distinguish between formula mass, empirical mass, molecular mass, gram molecular mass, and gram formula mass (all of which is referred to as molar mass). (p ) 5. Calculate the percent composition of a substance given its formula or masses of each component element in a sample. (p ) 6. Determine the empirical formulas and molecular formulas of substances, given percent composition data or mass composition data. (p ) 7. Explain what a hydrate is and how its name and formula reflect its composition. (p ) 8. Determine percent composition experimentally and derive empirical formulas from the data for hydrates. (p ) Chapter 11: Stoichiometry The amount of each reactant present at the start determines how much product can form. 1. Describe what is represented, on a molecular and molar level, by chemical equations (p ) 2. Perform stoichiometry calculations such as mole-mole, mole-mass, and mass-mass computations (p ) 3. Identify limiting reagents when solving stoichiometric problems (p ) 4. Compute theoretical yield, actual (experimental) yield, and percent yield for chemical reactions. (p ) Chapter 12: States of Matter Kinetic-Molecular Theory explains the different properties of solids, liquids, and gases. 1. Use the Kinetic-Molecular Theory to explain the behavior of gas pressure. (p ) 2. Explain the basis for gaseous diffusion and effusion. (p ) 3. Define gas pressure and the various pressure units (e.g., torr, kilopascals, mm Hg, atmospheres). (p ) 4. Calculate the partial pressure of a gas using Dalton s law of partial pressures. (p ) 5. Compare the different types of intermolecular forces such as dispersion and van der Waals. (p ) 6. Describe hydrogen bonding and relate the special properties of water that result from hydrogen bonding. (p ) 7. Describe the phase and energy changes associated with boiling/condensation, melting/freezing, and sublimation/deposition. (p ) 8. Interpret phase diagrams relating pressure and temperature of a substance. (p ) Chapter 13: Gases Gases respond in predictable ways to changes in pressure, temperature, volume, and moles. 1. Explain the importance of absolute temperature scale and convert between Kelvin and Celsius scales. (p ) 2. Define the gas laws and solve problems based on these laws. (p ; Table 1 p. 451) 3. Describe Avogadro s principle and relate it to the concept of molar volume: 22.4 L / mol at STP (p ) 4. Apply the mathematical relationships that exist among ideal gases (PV=nRT) (p ) 5. Chose the correct value of the gas constant, R, when using the Ideal Gas Law equation. (Table 2 p. 454) 6. Apply the ideal gas law to determine the density and molar mass of a gas. (p ) 7. Solve gas stoichiometry problems at standard and nonstandard conditions. (p )

6 Quarter 4 Chapter 14: Solutions Nearly all substances that make up our world are mixtures commonly found as solutions. 1. Compare properties of suspensions, colloids, and true solutions. (p ) 2. Define solution, solute, solvent, soluble, insoluble, miscible, and immiscible. (p ) 3. Give examples of solid, liquid, or gas medium solutions. (Table 2 p. 479) 4. Determine solution concentrations by calculating percent concentration, molarity, molality, and mole fraction. (p ; ) 5. Describe the techniques and calculations needed to prepare and dilute solutions. (p ) 6. Describe how intermolecular forces affect solvation (p ) 7. Calculate the heat of solution when a solid is dissolved into a liquid. (p. 492) 8. Describe factors affecting the solubility of a solute in a given solvent and its rate of solution (p. 492) 9. Define the terms saturated, unsaturated, supersaturated, dilute, and concentrated. (p ) 10. Describe factors that affect solubility of solids and gases in liquids (p ) Chapter 15: Thermochemistry Chemical reactions usually absorb or release energy. 1. Explain the law of conservation of energy in chemical reactions. (p ) 2. Describe the concept of heat, and explain the difference between heat energy and temperature. (p. 517) 3. Interconvert energy units such as calorie, Calorie (Kcal), joules, and kilojoules. (p. 518) 4. Solve heat capacity, heat transfer, and calorimetry problems. (p ; ) 5. Explain how changes in enthalpy determine whether a reaction is endothermic or exothermic. (p ) 6. Determine heat lost or gained during changes of state (p ) Chapter 16: Chemical Rates Chem reactions proceed at a definite rate, but can be speeded up or slowed down. 1. Calculate the average rate of a chemical reaction given experimental data. (p ) 2. Use collision theory to explain how rates of chemical reactions are related to collisions between reacting particles. (p ) 3. Determine the factors that affect the rate of a chemical reaction. (p ) 4. Explain the role of catalysts and inhibitors in chemical reactions. (p ) Chapter 17: Equilibrium Chemical equilibrium is reached when both reactants and products are formed at equal rates. 1. Explain the law of mass action and write equilibrium law expressions for chemical equilibria. (p ) 2. Calculate equilibrium constants (K eq) from concentration data. (p ) 3. Apply Le Châtelier s principle to explain a variety of changes in physical and chemical equilibria. (p ) 4. Calculate the solubility of a compound from its solubility product constant (p ) Chapter 18: Acids and Bases- can be defined in terms of hydrogen and hydroxide ions or in terms of electron pairs. 1. Describe the properties of acids and bases. (p ) 2. Describe and compare the Arrhenius and Bronsted-Lowry models of acids and bases. (p ) 3. Describe the hydronium ion and the concept of a substance being amphoteric. (p ) 4. Name and write formulas for common acids and bases. (p ) 5. Describe characteristics of strong acids and bases, and identify common examples of both. (p ; ) 6. Define the water constant, K w, and the ph scale. (p. ) 7. Calculate hydrogen ion concentration, hydroxide ion concentration, ph, and poh for acidic or basic solutions. (p ) 8. Explain how neutralization reactions are used in acid-base titrations. (p ) 9. Explain how acid-base indicators work. (p. 662) 10. Define the terms titration, volumetric analysis, equivalence point, end-point, primary standard, and standardization. (p ) 11. Demonstrate ability to solve acid-base titration math problems. (See Titration Lab Notes) 12. Conduct titration experiments to standardize and determine concentration of an acid or base solution. (See Titration Lab Notes)