Mr. Morrow s Accelerated Chemistry Syllabus

Mr. Morrow s Accelerated Chemistry Syllabus Mr. Morrow s Accelerated Chemistry course is ten units consisting of nineteen chapters from the Glencoe Chemistry textbook. The students are expected to be in the accelerated math tract or have a strong background in math. This will allow students to apply the mathematics needed to perform mathematical computations involving chemistry concepts. Students will need a scientific calculator with a preference towards a TI-83 or TI-84. Students will be asked to have formal lab reports for twenty-one experiments performed with a partner or a small group. A student s grade will be based on an accumulation of points during each nine weeks for homework assignments, lab reports, worksheets, labs, quizzes and tests. During the year students will take a midterm at the end of each nine weeks. At the end of three of these four nine weeks there will be a lab practical that will be a component of their midterm exam grade and final exam grade. First Nine Weeks: Unit 1 Matter-Chapter 1 (Chemistry and Matter) and Chapter 3 (Matter Properties and Changes) Name elements (with correct spelling) when given their symbol or write the symbols of elements when given their names for fifty of the most common elements. (This includes Latin names for select elements.) Understand the major types of matter: substance, mixture(homogeneous and heterogeneous), element, compound Define chemistry, and list and identify examples of its branches. Compare and contrast basic research, applied research, and technological development. Describe the purpose and the major steps of the scientific method. Distinguish between a qualitative and quantitative observation. Describe the difference between hypotheses, theories (models) and law. Observe several chemical phenomena. and hypothesize explanations for each phenomenon Recognize that science is a systematic method of continuing investigation, based on observation, hypothesis testing, measurement, experimentation, and theory building, which leads to more adequate explanations of natural phenomena. Recognize that scientific knowledge and explanations are changing and almost always building on earlier knowledge. Describe problems with our vision in relationship to science. Describe the dual roles of scientific instruments and theories. Use the rowdy box to describe the law of conservation of mass. Use the Tour de France to describe the law of conservation of energy. Use the wind-up toy to describe the law of conservation of mass and energy, and to describe how this relationship is shown in Einstein s famous equation, E = mc 2. Describe how matter can be sorted. Define elements and when they do combine. They do so in whole number ratios and unique geometric patterns Distinguish between heterogeneous and homogeneous mixtures. Identify the different allotropes of carbon and their structural differences. Explain how the different structures of carbon allotropes affect their properties. Understand and follow correct lab safety procedures. Write an informal and formal lab report. Read and use a graduated cylinder to the correct number of sig figs Read and use an electronic balance to the correct number of sig figs Read and use an analytical balance to the correct number of sig figs Read and use a buret to the correct number of sig figs Determine the density of a liquid using a pipet and balance Make a heterogeneous mixture and use filtration to make a homogeneous mixture. Correctly light and use a Bunsen burner. Determine the volume and density of a metal. Determine the thickness of a piece of aluminum foil Lab 1 Observing and Inferring Lab 2 Following Directions Lab 3 Use and ID of Lab Equipment Lab 4 Specific Heat of Wood

Unit 2 SI System, Sig Figs and the Factor-Label Method-Chapter 2 Syllabus (Analyzing Data Objectives) Distinguish between a quantity, a unit, and a measurement standard. Name SI units for length, mass, time, temperature, volume, and density. Distinguish between mass and weight. Use the factor-label method to convert units and solve problems Distinguish between accuracy and precision and determine the % error and % deviation of lab results Determine the number of significant figures in measurements. Perform calculations and report answers with correct number of significant figures. Convert measurements into and out of scientific notation. Measuring and calculate the density of objects with an analytical balance and graduated cylinder. Create graphs on an Excel spreadsheet and/or TI calculator to reveal patterns for a set of data by determining the slope of a best fit line including the correlation factor (r 2 ) Determine the density of various solids by direct measurement and water displacement and liquids Lab 5 Density of a Pre and Post 1982 Pennies Lab 6 Density of Various Solids and Liquids Lab 7 Separating a Mixture (Discovery Lab Practical) Second Nine Weeks: Unit 3 Atomic Theory Development-Chapter 4 (Structure of the Atom) and Chapter 24 (Nuclear Chemistry) Explain the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. Summarize the five essential points of Dalton s atomic theory. Explain the relationship between Dalton s atomic theory and the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. Summarize the experiments of Thomson, Millikan, and Rutherford and their relationship to atomic theory (model) List the properties of electrons, protons, and neutrons Explain what a nuclide is and describe the different ways nuclides can be represented Define and relate the terms mass defect and nuclear binding energy Explain the relationship between nucleon and stability of nuclei Explain why nuclear reactions occur and how to balance a nuclear equation Students will determine if the law of conservation of mass holds during multiple chemical reactions Lab 8 Law of Conservation of Mass Unit 4 Electrons and the Periodic Table-Chapter 5 (Electrons in Atom) and Chapter 6 (The Periodic Table and Periodic Law) o Discuss the significance of the line-emission spectrum of elements to the development of the atomic model. o Explain the mathematical relationship between the speed, wavelength, and frequency of electromagnetic radiation. o Discuss the dual wave-particle nature of light. o Discuss the significance of the photoelectric effect and the line-emission spectrum of hydrogen to the development of o o the atomic model. Describe the Bohr model of the hydrogen atom Use the Schrodinger wave equation to map out any atom s electrons most probable location using electron configuration and orbital diagrams

o o State the Pauli exclusion principle, aufbau principle and Hund s rule and use these rules in writing electron configuration and orbital diagrams o Determine the number of electrons in any given energy level, sublevel and orbital o Determine any electrons four digit quantum number o Explain the roles of Mendeleev and Moseley in the development of the periodic table. o Describe the modern periodic table. o Explain how the periodic law can be used to predict the physical and chemical properties of elements. o Describe how the elements belonging to a group of the periodic table are interrelated in terms of atomic number. o Describe the relationship between electrons in sublevels and the length of each period of the periodic table. Locate and name the four blocks of the periodic table. Explain the reasons for these names. Discuss the relationship between group configurations and group numbers. Describe the locations in the periodic table and the general properties of the alkali metals, the alkaline-earth metals, the halogens, and the noble gases. Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity. Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for these variations. Define valence electrons, and state how many are present in atoms of each main-group element. Compare the atomic radii, ionization energies, and electronegativities of the d-block elements with those of the maingroup elements. Lab 9 Flame Test Lab 10 Spectroscope Unit 5 Nomenclature and Bonding-Chapter 7 (Ionic Compounds and Metals), Chapter 8 (Covalent Bonding) and Chapter 22(Organic Chemistry) Define chemical bond. Explain why most atoms form chemical bonds. Compare and contrast an ionic and a covalent bonding. Explain why most chemical bonding is neither purely ionic nor purely covalent Explain the importance and limitations of empirical, molecular and structural formulas. Classify bonding type according to electronegativity differences. Define molecule and molecular formula. Explain the relationships between potential energy, distance between approaching atoms, bond length, and bond energy. State the octet rule. Write Lewis structures for both ionic and covalent compounds. Draw Lewis structures for molecules and polyatomic ions containing single bonds, multiple bonds, or both. Relate the ability of a carbon atom to form covalent bonds to its atomic structure and hybrid orbitals. Explain how the structure and bonding of carbon lead to the diversity and number of organic compounds. Draw resonance structures for molecules and polyatomic ions Use the VSEPR theory to predict the shapes of molecules or polyatomic ions. Explain how the shapes of molecules are accounted for by hybridization theory Determine if a molecule is a dipole and indicate the δ + and δ - of a molecule Explain the significance of a chemical formula. Determine the formula of an ionic compound when given its name. Name an ionic compound given its formula. Using prefixes, name a binary molecular compound from its formula. Write the formula of a binary molecular compound given its name. Name binary molecular compounds using the IUPAC system. Define empirical formula, and explain how the term applies to ionic and molecular compounds. Lab 11 Molecular Models (VSEPR)

Third Nine Weeks: Unit 6-Descriptive Chemistry-Chapter 9 (Chemical Reactions and Chemical Equations) and Chapter 16 (Reaction Rates) List three observations that suggest that a chemical reaction has taken place. Write a word equation, skeleton equation, balanced equation and net ionic equation with correct stats of matter for a given chemical reaction. Define and give general equations for synthesis, decomposition, single-replacement, double-replacement and combustion reactions. Categorize a possible reaction, predict the products of the reactions, write a word equation, skeleton equation and balanced equation with correct stats of when given the reactants. Use an activity series and solubility chart to predict whether a given reaction will occur and what the products will be including states of matter. Predict whether a precipitate will form when solutions of soluble ionic compounds are combined, and write net ionic equations for precipitation reactions. Define oxidation and reduction. Explain what an oxidation-reduction reaction (redox reaction) is. Explain what must be conserved in redox equations. Discuss the five factors that influence reaction rate. Define catalyst, and discuss two different types. Discuss conditions under which reactions go to completion. Lab 12 Descriptive Chemistry Lab 13 Iodine Clock Reaction Unit 7- Stoichiometry-Chapter 10 (The Mole) and Chapter 11 (Stoichiometry) Explain how a mole is used to count the number of particles of matter Convert between moles, representative particles, volume and mass by using a mole map Calculate the formula mass or molar mass of any given compound. Use molar mass to convert between mass in grams and amount in moles of a chemical compound. Calculate the number of molecules, formula units, or ions in a given molar amount of a chemical compound. Calculate the percentage composition of a given chemical compound. Determine an empirical formula from either a percentage or a mass composition. Determine a molecular formula from an empirical formula when given the molecular mass. Define stoichiometry. Write a mole ratio relating two substances in a chemical equation. Calculate the amount in moles of a reactant or product from the amount in moles of a different reactant or product. Calculate the mass of a reactant or product from the amount in moles of a different reactant or product. Calculate the amount in moles of a reactant or product from the mass of a different reactant or product. Calculate the mass of a reactant or product from the mass of a different reactant or product. Determine which reactants is a limiting reactant when given the mass or moles of both reactants Calculate the amount in moles or mass in grams of a product, given the amounts in moles or masses in grams of two reactants, one of which is in excess. Distinguish between theoretical yield, actual yield, and percent yield. Calculate percent yield, given the actual yield and quantity of a reactant. Lab 14 Magnesium Oxide Empirical Formula Lab 15 Hydrate Formula Lab 16 Insoluble Salt Production (Discovery Lab Practical) Unit 8 KMT and the Gas Laws-Chapter 12 (States of Matter) and Chapter 13 (Gas Laws) State the kinetic-molecular theory of matter, and describe how it explains certain properties of matter. Describe the motion of particles in liquids and the properties of liquids according to the kinetic-molecular theory. Describe dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces. Describe the structure of a water molecule.

o Discuss the physical properties of water. Explain how they are determined by the structure of water. Discuss the process by which liquids can change into a gas. Define vaporization. Discuss the process by which liquids can change into a solid. Define freezing. Describe the motion of particles in solids and the properties of solids according to the kinetic-molecular theory. Distinguish between the types of solids. Discuss the arrangements of ions in crystals. Define lattice energy and explain its significance. Describe the electron-sea model of metallic bonding, and explain why metals are good electrical conductors. Explain why metal surfaces are shiny. Explain why metals are malleable and ductile but ionic-crystalline compounds are not. Explain the relationship between equilibrium and changes of state. Describe the processes of boiling, freezing, melting, and sublimation. Interpret phase diagrams. List the five assumptions of the kinetic-molecular theory of gases. Define the terms ideal gas and real gas Describe each of the following characteristic properties of gases: expansion, density, fluidity, compressibility, diffusion, and effusion. Describe the conditions under which a real gas deviates from ideal behavior. Define pressure, and relate it to force. Describe how pressure is measured. Convert units of pressure. State the standard conditions of temperature and pressure. Use the kinetic-molecular theory to explain the relationships between gas volume, temperature, and pressure. Use Boyle s law to calculate volume-pressure changes at constant temperature. Use Charles s law to calculate volume-temperature changes at constant pressure. Use Gay-Lussac s law to calculate pressure-temperature changes at constant volume. Use the combined gas law to calculate volume-temperature-pressure changes. Use Dalton s law of partial pressures to calculate partial pressures and total pressures. State Avogadro s law and explain its significance. Define standard molar volume of a gas, and use it to calculate gas masses and volumes. Use standard molar volume to calculate the molar mass of a gas. State the ideal gas law. Derive the ideal gas constant and discuss its units. Using the ideal gas law, calculate pressure, volume, temperature, or amount of gas when the other three quantities are known. Using the ideal gas law, calculate the molar mass or density of a gas. Explain how Gay-Lussac s law and Avogadro s law apply to the volumes of gases in chemical reactions. Use a chemical equation to specify volume ratios for gaseous reactants or products, or both. Use volume ratios and the gas laws to calculate volumes, masses, or molar amounts of gaseous reactants or products. State Graham s law of effusion. Determine the relative rates of effusion of two gases of known molar masses. State the relationship between the molecular velocities of two gases and their molar masses. Lab 17 Dry Ice Phase Diagram Lab 18 Boyle s Law Lab 19 Molar Mass of Butane Fourth Nine Weeks: Unit 9 Solutions and Colligative Properties-Chapter 14 (Mixtures and Solutions) List three different solute-solvent combinations. Compare the properties of suspensions, colloids, and solutions. Distinguish between strong electrolytes, weak electrolytes and nonelectrolytes List and explain three factors that affect the rate at which a solid solute dissolves in a liquid solvent. Explain solution equilibrium, and distinguish among saturated, unsaturated, and supersaturated solutions. Explain the meaning of like dissolves like in terms of polar and nonpolar substances. List the three interactions that contribute to the heat of solution, and explain what causes dissolution to be exothermic or endothermic.

o Compare the effects of temperature and pressure on solubility. Given the mass of solute and volume of solvent, calculate the concentration of a solution. Given the concentration of a solution, determine the amount of solute in a given amount of solution. Given the concentration of a solution, determine the amount of solution that contains a given amount of solute. Write equations for the dissolution of soluble ionic compounds in water. Compare dissociation of ionic compounds with ionization of molecular compounds. Draw the structure of the hydronium ion, and explain why it is used to represent the hydrogen ion in solution.. List four colligative properties, and explain why they are classified as colligative properties. Calculate freezing-point depression, boiling-point elevation, and solution molality of nonelectrolytic solutions. Calculate the expected changes in freezing point and boiling point of an electrolytic solution. Discuss causes of the differences between expected and experimentally observed colligative properties of electrolytic solutions. Lab 20 Electrolytes Unit 10-Acids and Bases-Chapter 17 (Chemical Equilibrium) and Chapter 18 (Acids and Bases) List three different solute-solvent combinations. Define chemical equilibrium. Explain the nature of the equilibrium constant. Write chemical equilibrium expressions and carry out calculations involving them. Predict changes in equilibrium using Le Châtelier s principle when factors that disturb equilibrium. Explain the concept of acid-ionization constants, and write acid-ionization equilibrium expressions. Define the ionization constant of water. Explain buffering. Compare cation and anion hydrolysis List five general properties of aqueous acids and bases. Name common binary acids and oxyacids, given their chemical formulas. List five acids commonly used in industry and the laboratory, and give two properties of each. Define acid and base according to Arrhenius s, Brønsted-Lowry and Lewis Explain the differences between strong and weak acids and bases. Name compounds that are acids under the Lewis definition but are not acids under the Brønsted-Lowry definition. Describe a conjugate acid, a conjugate base, and an amphoteric compound. Explain the process of neutralization. Explain how acid rain damages marble structures. o Define ph, and give the ph of a neutral solution at 25 C. Explain and use the ph scale. Given one of the following find the other three [H 3 O + ] [OH ], ph and poh. Describe how an acid-base indicator functions. Carry out an acid-base titration. Standardize a sodium hydroxide solution Calculate the molarity of a solution from titration data. Lab 21 Household Acids and Bases Lab 22 Differences in Store Bought Vinegar (Discovery Lab Practical)