WEST ESSEX REGIONAL SCHOOL DISTRICT
Science Department
I. COURSE NAME: Honors Chemistry
II. COURSE PREREQUISITES: Biology Honors, Algebra I Honors
III. GRADE LEVEL: 10, 11
IV. COURSE DESCRIPTION: This course has been designed with respect to and in compliance with the expectations set forth in the New Jersey Core Curriculum Content Standards.
V. COURSE OBJECTIVES:
This course will provide students with the ability to:
· learn how careful observations and other activities of science are useful in developing a model of the structure of matter and its interactions.
· develop an appreciation for the importance of uncertainty in measurements and how they apply to laboratory data.
· become familiar with calculations based upon stoichiometric relationships in chemical reactions and chemical solutions using unit analysis.
· develop an understanding of atomic theory.
· demonstrate an understanding of electromagnetic radiation, electron configuration, molecular architecture, and their relationships to chemical bonding.
· study the various phases of matter and draw relationships to chemical bonding.
· demonstrate an understanding of the relationships and trends among the atoms in a chemical family, the difference between chemical families and the periodic law.
· develop an understanding of the interactions between matter and energy in chemical and nuclear changes.
· develop an appreciation for the factors governing the rates of chemical reactions.
· develop an understanding of the laws of chemical equilibrium and its applications in chemical reactions, solubility, and acid base chemistry.
· gain an appreciation for the career opportunities available in the field of chemistry.
Unit 1: Introduction to Matter and Scientific Measurement
Upon completion of this unit, students should be able to:
· Identify branches of chemistry
· Describe practical applications of chemistry
· Determine the classifications of matter
· Describe basic separation techniques
· Use symbols to convey chemical information
· Use the SI system for scientific measurement
· Determine the uncertainty in a measurement
· Use measurements and calculations with the correct number of significant figures
· Use dimensional analysis to convert between units
· Conduct research with academic integrity
· Develop basic laboratory techniques and skills
· Demonstrate laboratory findings in written and oral formats
· Develop collaborative skills in a laboratory and research context
Unit 2: Calorimetry/Heating Curves/States of Matter
· Demonstrate proficiency in the use of calorimeters and related equipment.
· Define heat and state its units.
· Perform specific heat calculations.
· Describe the motion of particles in liquids and the properties of liquids according to the kinetic-molecular theory.
· 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 two types of solids.
· Describe the different types of crystal symmetry. Define crystal and unit cell.
· Explain the relationship between equilibrium and changes of state.
· Interpret and draw phase diagrams.
· Explain what is meant by equilibrium vapor pressure.
· Describe the processes of boiling, freezing, melting, and sublimation.
· Describe the structure of the water molecule.
· Discuss the physical properties of water. Explain how they are determined by the structure of water.
· Calculate the amount of energy absorbed or released when a quantity of water changes state.
Unit 3 Physical Characteristics and Molecular Composition of Gases
· List the five assumptions of the kinetic molecular theory of gases.
· Describe the conditions under which a real gas deviates from an “ideal” gas.
· Define pressure and relate it to force.
· Convert units of pressure.
· State the standard conditions of temperature and pressure.
· Given experimental data, determine the molar volume of a gas.
· Collect gas by water displacement and make a standard pressure and temperature comparison to the accepted value
· Use Boyle’s Law, Charles’ Law, Gay-Lussac’s Law, and the combined gas law to calculate pressure-volume-temperature changes.
· Use Dalton’s Law of Partial Pressures to calculate partial pressures and total pressures.
· State the ideal gas law.
· Derive the ideal gas constant and discuss its units.
· Use the ideal gas law to calculate the molar mass or density of a gas.
· Use volume ratios and the gas laws to calculate volumes, masses, or molar amounts of gaseous reactants and products.
Unit 4 Mole Concept
· Explore, by analogy, how chemists count atoms and molecules.
· Evaluate the accuracy of a “counting by weighing” method
· Write Lewis structures for a polyatomic ion given the identity of the atoms combined and other appropriate information
· Define mole in terms of Avogadro’s number and define molar mass.
· Solve problems involving mass in grams, and number of atoms of an element.
· Calculate the percentage composition of a given chemical compound
· Define mole in terms of Avogadro’s number and define molar mass.
· Solve problems involving mass in grams, and number of atoms of an element.
· Calculate the percentage composition of a given chemical compound.
· Determine empirical formula, and explain how the term applies to ionic and molecular compounds.
· Determine an empirical formula from either a percentage or a mass composition.
· Explain the relationship between empirical formula and the molecular formula of a given compound.
· Determine a molecular formula from an empirical formula.
· Given the mass of solute and volume of solvent, calculate the concentration of a solution.
· Given the concentration of 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.
· Given experimental data, determine the moles of water in a given hydrate.
Unit 5: Atomic Theory, Electrons, Periodic Table and Trends
· 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 observed properties of cathode rays that led to the discovery of the electron.
· Summarize the experiment carried out by Rutherford and his co-workers that led to the discovery of the nucleus.
· List the properties of protons, neutrons, and electrons.
· Define atom.
· Explain what isotopes are.
· Define atomic number and mass number, and describe how they apply to isotopes.
· Given the identity of a nuclide, determine its number of protons, neutrons, and electrons.
· Define mole in terms of Avogadro’s number, and define molar mass.
· Solve problems involving mass in grams, amount in moles, and number of atoms of an element.
· Explain the mathematical relationship between the speed, wavelength, and frequency of electromagnetic radiation.
· Discuss the dual wave-particle nature of light.
· Discuss the significance of the photoelectric effect and the line-emission spectrum of hydrogen to the development of the atomic model.
· Describe the Bohr model of the hydrogen atom.
· Discuss Louis de Broglie’s role in the development of the quantum model of the atom.
· Compare and contrast the Bohr model and the quantum model of the atom.
· Explain how the Heisenberg uncertainty principle and the Schrödinger wave equation led to the idea of atomic orbitals.
· List the four quantum numbers, and describe their significance.
· Relate the number of sublevels corresponding to each of an atom’s main energy levels, the number of orbitals per sublevel, and the number of orbitals per main energy level.
· List the total number of electrons needed to fully occupy each main energy level.
· State the Aufbau principle, the Pauli exclusion principle, and Hund’s rule.
· Describe the electron configurations for the atoms of any element using orbital notation, electron-configuration notation, and, when appropriate, noble-gas notation.
· Explain the roles of Mendeleev and Moseley in the development of the periodic table.
· Describe the modern periodic table.
· Explain how the periodic law can be used to predict the physical and chemical properties of elements.
· Describe how the elements belonging to a group of the periodic table are interrelated in terms of atomic number.
· 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 main-group elements.
Unit 6 Chemical Bonding and Nomenclature
· Define chemical bond
· Explain why most atoms from chemical bonds.
· Describe ionic and covalent bonding.
· Explain why most bonding is neither purely ionic nor purely covalent.
· Classify bonding type according to electronegativity differences.
· Explain the significance of a chemical formula.
· Determine the formula of an ionic compound formed between two given ions.
· Name an ionic compound given its formula.
· Using prefixes, name a binary molecular compounds from its formula.
· Write the formula of a binary molecular compound given its name.
· List the rules for assigning oxidation numbers.
· Give the oxidation number for each element in the formula of a chemical compound.
· Name binary molecular compounds using oxidation numbers and the Stock system.
· Explain the relationships between potential energy, distance between approaching atoms, bond length, and bond energy.
· Draw Lewis structures for molecules containing single bonds, multiple bonds, or both.
· Explain why resonance structures are used to represent some molecules.
· Draw Lewis structures for molecules containing single bonds, multiple bonds, or both.
· Explain why resonance structures are used to represent some molecules.
· Explain VSEPR theory.
· Predict the shapes of molecules and polyatomic ions using VSEPR theory.
· Explain how shapes of molecules are accounted for by hybridization theory.
· Describe dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces.
· Explain what determines molecular polarity.
Unit 7 Chemical Equations and Reactions
· List three observations that suggest a chemical reaction has taken place.
· List three requirements for a correctly written chemical equation.
· Write a word equation and a formula equation for a given chemical reaction.
· Balance a formula equation by inspection.
· Define and give general equations for synthesis, decomposition, single-replacement, and double replacement reactions.
· Classify a reaction as synthesis, decomposition, single-replacement, double replacement reaction, or combustion.
· List three types of synthesis reactions and six types of decomposition reactions.
· Predict the products of simple equations given the reactants.
· Classify a reaction as single-replacement.
· Use an activity series to predict whether a given reaction will occur and predict the products.
· Write equations for the dissolution of soluble ionic compounds in water.
· Predict whether a precipitate will form when solutions of soluble ionic compounds are combined.
· Observe a variety of chemical reactions and identify the patterns in the conversion of reactants into products.
· Analyze reactions to classify the chemical reactions into different groups.
· Use an analysis scheme for identifying both cations and anions in solution in order to determine the ions present in known and unknown solutions.
Unit 8 Stoichiometry
· Define stoichiometry.
· Describe the importance of the mole relationship in stoichiometric calculations.
· Calculate the amount in moles of a reactant or product from the amount in moles of a different reactant or product.
· Calculate the amount in mass of a reactant or product from the amount in mass of a different reactant or product.
· Calculate the amount on moles or mass in grams of a product, given the amounts of 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.
Unit 9 Solution Chemistry and Colligative Properties
· Give 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
· Determine the volume of stock solution and the amount of solvent needed to dilute a solution to a given concentration.
· Solve problems using the dilution formula.
· Prepare a series of solutions using the molarity and dilution equations.
· Investigate the relationship between concentration of a solution and absorbance.
· Determine the accuracy of the solution preparation and dilution procedures.
· Prepare a series of solutions using the molarity and dilution equations.
· Investigate the relationship between concentration of a solution and absorbance.
· Determine the accuracy of the solution preparation and dilution procedures.
· List four colligative properties, and explain why they are classified as colligative properties.
· Calculate freezing-point depression, boiling point elevation, and solution molality of nonelectrolyte 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.