Chemistry:  Methods and Measurement

• State the definition of chemistry and discuss its interrelationship with other fields of science and medicine.
• Describe the approach to science, the scientific method.
• Distinguish among the terms hypothesis, theory and scientific law.
• State both the differences and relationships between science and technology.
• Distinguish between data and results.
• Learn the major units of measure in the English and metric systems, and be able to convert from one system to another.
• Compare and contrast the terms error, accuracy, precision, and uncertainty
• Report data and results using scientific notation and the proper number of significant figures.
• Use appropriate experimental quantities in problem solving.
• Calculate the density of an object from mass and volume data and calculate the specific gravity of an object from its density.

The Composition and Structure of the Atom

• Describe the properties of the solid, liquid, and gaseous state.
• Classify properties as chemical or physical.
• Classify observed changes in matter as chemical or physical.
• Provide specific examples of physical and chemical properties.
• Distinguish between intensive and extensive properties.
• Classify matter as element, compound, or mixture.
• Recognize the interrelationship of the structure of matter and its physical and chemical properties.
• Describe the important properties of protons, neutrons, and electrons.
• Calculate the number of protons, neutrons, and electrons in any atom.
• Distinguish among atoms, ions, and isotopes.
• Trace the history of development of atomic theory, beginning with Dalton.
• Summarize the experimental basis for the discovery of the charged particles and the nucleus.
• Explain the critical role of spectroscopy in the development of atomic theory and in our everyday lives.
• State the basic postulates of Bohr’s theory.
• Compare and contrast Bohr’s theory and the more sophisticated “wave-mechanical” approach.

Elements, Atoms, Ions, and the Periodic Table

• Recognize the important subdivisions of the periodic table:  periods, groups (families), metals, and nonmetals
• Use the periodic table to obtain information about an element.
• Describe the relationship between the electronic structure of an element and its position in the periodic table.
• Write electron configurations for atoms of the most commonly occurring elements.
• Know the meaning of the octet rule and its predictive usefulness.
• Use the octet rule to predict the charge of common cations and anions.
• Utilize the periodic table and its predictive power to estimate the relative sizes of atoms and ions, as well as relative magnitudes of ionization energy and electron affinity.
• Use values of ionization energies and electron affinities to predict ion formation.

Structure and Properties of Ionic and Covalent Compounds

• Classify compounds as having ionic, covalent, or polar covalent bonds.
• Name common inorganic compounds using standard conventions and recognize the common names of frequently used substances.
• Write the formulas of compounds when provided with the name of the compound.
• Predict the differences in physical state, melting and boiling points, solid-state structure, and solution chemistry that result from differences in bonding.
• Draw Lewis structures for covalent compounds and complex inorganic ions.
• Describe the relationship between stability and bond energy.
• Predict the geometry of molecules and ions using the octet rule and Lewis structure.
• Understand the role that molecular geometry plays in determining the solubility and melting and boiling points of compounds.
• Use the principles of VSEPR theory and molecular geometry to predict relative melting points, boiling points, and solubilities of compounds.

Calculations and the Chemical Equation

• Know the relationship between the mole and Avagadro’s number, and the usefulness of these quantities.
• Perform calculations using Avagadro’s number and the mole.
• Write chemical formulas for common inorganic substances.
• Calculate the formula weight and molar mass of a compound.
• Know the major function served by the chemical equation, the basis for chemical calculations.
• Balance chemical equations given the identity of products and reactants.
• Calculate the number of moles of product resulting from a given number of moles of reactants or the number of moles of reactant needed to produce a certain number of moles of product.
• Perform calculations involving limiting reactants to determine the theoretical yield of a reaction.
• Calculate theoretical and percent yield.

Reactions and Solutions

• Classify chemical reactions by type:  combination (synthesis), decomposition, or replacement (single or double).
• Recognize the various classes of chemical reactions:  precipitation, reactions with oxygen, acid-base, and oxidation-reduction
• Distinguish among the terms solution, solute, and solvent.
• Describe various kinds of solutions, and give examples of each.
• Describe the relationship between solubility and equilibrium.
• Calculate solution concentration in weight/volume percent and weight/weight percent.
• Calculate solution concentration using molarity.
• Perform dilution calculations.
• Interconvert molar concentration of ions and milliequivalents/liter.
• Describe and explain concentration-dependent solution properties.
• Describe why the chemical and physical properties of water make it a truly unique solvent.
• Explain the role of electrolytes in blood and their relationship to the process of dialysis.

States of Matter:  Gases, Liquids, and Solids

• Describe the behavior of gases expressed by the gas laws:  Boyle’s law, Charles’s law, combined gas law, Avagadro’s law, the ideal gas law, and Dalton’s law.
• Use gas law equations to calculate conditions and changes in conditions of gases.
• Describe the major points of the kinetic molecular theory of gases.
• Explain the relationship between the kinetic molecular theory and the physical properties of macroscopic quantities of gases.
• Describe properties of the liquid state in terms of the properties of the individual molecules that comprise the liquid.
• Describe the process of melting, boiling, evaporation, and condensation.
• Describe the dipolar attractions known collectively as van der Waals forces.
• Describe hydrogen bonding and its relationship to boiling and melting temperatures.
• Relate the properties of the various classes of solids (ionic, covalent, molecular, and metallic) to the structure of these solids.

Chemical and Physical Change:  Energy, Rate, and Equilibrium

• Correlate the terms endothermic and exothermic with heat flow between a system and its surroundings.
• State the meaning of the terms enthalpy, entropy, and free energy and know their implications.
• Describe experiments that yield thermochemical information and calculate fuel value based on experimental data.
• Describe the concept of reaction rate and the role of kinetics in chemical and physical change.
• Describe the importance of activation energy and the activated complex in determining reaction rate.
• Predict the way reactant structure, concentration, temperature, and catalysis affect the rate of a chemical reaction.
• Write rate equations for elementary processes.
• Recognize and describe equilibrium situations.
• Write equilibrium-constant expressions and use these expressions to calculate equilibrium constants.
• Use LeChatelier’s principle to predict changes in equilibrium position.

Charge Transfer Reactions:  Acids and Bases and Oxidation-Reduction

• Identify acids and bases and acid-base reactions.
• Write equations describing acid-base dissociation and label the conjugate acid-base pairs.
• Describe the role of the solvent in acid-base reactions, and explain the meaning of the term pH.
• Calculate pH from concentration data.
• Calculate hydronium and/or hydroxide ion concentration from pH data.
• Provide examples of the importance of pH in chemical and biochemical systems.
• Describe the meaning and utility of neutralization reactions.
• State the meaning of the term buffer and describe the applications of buffers to chemical and biochemical systems, particularly blood chemistry.
• Describe oxidation and reduction, and describe some practical examples of redox processes.
• Diagram a voltaic cell and describe its function.
• Compare and contrast voltaic and electrolytic cell.

The Nucleus, Radioactivity, and Nuclear Medicine

• Enumerate the characteristics of alpha, beta, and gamma radiation.
• Write balanced equations for common nuclear processes.
• Calculate the amount of radioactive substance remaining after a specified number of half-lives.
• Describe the various ways in which nuclear energy many be used to generate electricity:  fission, fusion, and the breeder reactor.
• Explain the process of radiocarbon dating.
• Cite several examples of the use of radioactive isotopes in medicine.
• Describe the use of ionizing radiation in cancer therapy.
• Discuss the preparation of radioisotopes for use in diagnostic imaging studies.
• Explain the difference between natural and artificial radioactivity.
• Describe the characteristics of radioactive materials that relate to radiation exposure and safety.
• Be familiar with common techniques for the detection of radioactivity.
• Know the common units in which radiation intensity is represented:  the curie, roentgen, rad, and rem.

An Introduction to Organic Chemistry:  The Saturated Hydrocarbons

• Compare and contrast organic and inorganic compounds.
• Draw structures that represent each of the families of organic compounds.
• Write the names and draw the structures of the common functional groups.
• Write condensed and structural formulas for saturated hydrocarbons.
• Describe the relationship between the structure and physical properties of saturated hydrocarbons.
• Use the basic rules of the IUPAC Nomenclature Sysytem to name alkanes and substituted alkanes.
• Draw constitutional isomers of simple organic compounds.
• Write the names and draw the structures of simple cycloalkanes.
• Draw cis and trans isomers of cycloalkanes.
• Describe conformations of alkanes.
• Draw the chair and boat conformations of cyclohexane.
• Write equations for the combustion reactions of alkanes.
• Write equations for halogenation reactions of alkanes.

The Unsaturated Hydrocarbons:  Alkenes, Alkynes, and Aromatics

• Describe the physical properties of alkenes and alkynes.
• Draw the structures and write the IUPAC names for simple alkenes and alkynes.
• Write the names and draw the structures of simple geometric isomers of alkenes.
• Write equations predicting the products of addition reactions of alkenes:  hydrogenation, halogenation, hydration, and hydrohalogenation.
• Apply Markovnikov’s rule to predict the major and minor products of the hydration and hydrohalogenation reactions of unsymmetrical alkenes.
• Write equations representing the oxidation of simple alkenes.
• Write equations representing the formation of addition polymers of alkenes.
• Draw the structures and write the names of common aromatic hydrocarbons.
• Write equations for substitution reactions involving benzene.
• Describe heterocyclic aromatic compounds and list several biological molecules in which they are found.

Learning Goals...with more specifics

Introduction:  Matter and Measurement

Matter:  Elements, Compounds, and Mixtures

• Distinguish between physical and chemical properties and also between simple and physical and chemical changes.
• Differentiate between the three states of matter.
• Distinguish between elements, compounds, and mixtures.
• Give the symbols for the elements discussed in class.

Physical Quantities and Units 

• You should be able to list the basic SI and metric units and the commonly used prefixes in scientific measurement.

Uncertainty in Measurements:  Significant Figures

• Determine the number of significant figures in a measured quantity.
• Express the result of a calculation with the proper number of significant figures.

Temperature and Density:  Intensive Properties

• Convert temperatures among the Fahrenheit, Celsius, and Kelvin scales.
• Perform calculations involving density.

Dimensional Analysis

• You should be able to convert between units by using dimensional analysis.

Atoms, Molecules, and Ions

Atoms

• Describe the composition of an atom in terms of protons, neutrons, and electrons.
• Give the approximate size, relative mass, and charge of an atom, proton, neutron, and electron.
• Write the chemical symbol for an element, having been given its mass number and atomic number and perform the reverse operation.
• Describe the properties of the electron as seen in cathode rays.  Describe the means by which J.J. Thomson determined the ratio e/m for the electron.
• Describe Millikan's oil-drop experiment and indicate what property of the electron he was able to measure.
• Cite the evidence from studies of radioactivity for the existence of subatomic particles.
• Describe the experimental evidence for the nuclear nature of the atom.

Molecules and Ions:  Relationships in the Periodic Table

• Write the symbol and charge for an atom or ion, having been given the number of protons, neutrons, and electrons, and perform the reverse operation.
• Use the periodic table to predict the charges of monatomic ions.
• Use the periodic table to predict whether an element is a metal or a nonmetal.
• Distinguish between empirical formulas, molecular formulas and structural formulas.

Nomenclature

• Write the simplest formula for a compound, having been given the charges of the ions from which it is made.
• Write the name of a simple inorganic compound, having been given its chemical formula, and perform the reverse operation.

Stoichiometry:  Calculations with Chemical Formulas and Equations

Chemical Equations:  Balancing and Predicting Products of Reactions

• Balance chemical equations.
• Predict the products of a chemical reaction, having seen a suitable analogy.
• Predict the products of the combustion reactions of hydrocarbons and simple compounds containing C, H, and O atoms.

Atomic Weight, Molecular Weight, and the Mole

• Calculate the atomic weight of an element given the abundances and masses of its isotopes.
• Calculate the molecular weight and molar mass of a substance from its chemical formula.
• Interconvert number of moles and mass of a substance.  Use Avogadro's number and molar mass to calculate the number of particles making up a substance, and vice versa.

Determination of Empirical and Molecular Formulas

• Calculate the empirical formula of a compound, having been given appropriate analytical data such as elemental percentages or the quantity of CO₂ and H₂O produced by combustion.
• Calculate the molecular formula, having been given the empirical formula and molecular weight.

Chemical Equations:  Mass and Mole Relationships

• Calculate the mass of a particular substance produced or used in a chemical reaction (mass-mass problem).
• Determine the limiting reagent in a reaction.
• Calculate the theoretical and actual yields of chemical reactions given the appropriate data.

Aqueous Reactions and Solution Stoichiometry

Aqueous Solutions:  Electrolytes and Acids and Bases

• Predict whether a substance is a nonelectrolyte, strong electrolyte, or weak electrolyte from its chemical behavior.
• Predict the ions formed by electrolytes when they dissociate of ionize.
• Identify substances as acids, bases, or salts.

Precipitation Reactions:  Ionic Equations

• Use solubility rules to predict whether a precipitate forms when electrolyte solutions are mixed.
• Predict the products of metathesis reactions (including both neutralization and precipitation reactions) and write balanced chemical equations for them.
• Identify the spectator ions and write the net ionic equations for solution reactions, starting with their molecular equations.

Oxidation and Reduction:  Oxidation Numbers and Activity Series

• Determine whether a chemical reaction involves oxidation and reduction.
• Assign oxidation numbers to atoms in molecules and ions.
• Use the activity series to predict whether a reaction will occur when a metal is added to an aqueous solution of either a metal salt of an acid; and write the balanced molecular and net ionic equation for  the reaction.

Concentration of Solutions

• Calculate molatiry; solution volume, or number of moles of solute given any two of these quantities.
• Calculate the volume of a more concentrated solution that must be diluted to obtain a given quantity of a more dilute solution.

Solution Stoichiometry

• Calculate the volume of a solution required to react with a volume of a different solution using molarity and the stoichiometry of the reaction.
• Calculate the amount of a substance required to react with a given volume of a solution using molarity and the stoichiometry of the reaction.
• Calculate the concentration or mass of solute in a sample from titration data.

Thermochemistry

Thermodynamics:  The First Law and Internal Energy Changes

• Give examples of different forms of energy.
• List the important units in which energy is expressed and convert from one to another.
• Define the first law of thermodynamics both verbally and by means of an equation.
• Describe how the change in internal energy of a system is related to the exchange of heat and work between the system and its surroundings.
• Define the term state function and describe its importance in thermochemistry.

Enthalpy:  Heats of Reactions

• Define enthalpy, and relate the enthalpy change in a process occurring at constant pressure to the heat added to or lost by the system during the process.
• Sketch an energy diagram, given the enthalpy changes in the processes involved, and associate the sign of DH with whether the process is exothermic or endothermic.
• Calculate the quantity of heat involved in a reaction at constant pressure given the quantity of reactants and the enthalpy change for the reaction on a mole basis.

Calorimetry:  Fuel Values

• Define the terms heat capacity and specific heat.
• Calculate any one of the following quantities given the other three:  heat, quantity of material, temperature change, and specific heat.
• Calculate the heat capacity of a calorimeter, given the temperature change and quantity of heat involved; also calculate the heat evolved or absorbed in a process from a knowledge of the heat capacity of the system and its temperature change.
• Define the term fuel value; calculate the fuel value of a substance given its heat of combustion or estimate the fuel value of a material given its composition.
• List the major sources of energy on which humankind must depend, and discuss the likely availability of these for the foreseeable future.

Hess's Law

• State Hess's law, and apply it to calculate the enthalpy change in a process, given the enthalpy changes in other processes that could be combined to yield the reaction of interest.

Heats of Formation:  Calculating Heats of Reactions

• Define and illustrate what is meant by the term standard state, and identify the standard states for the elements carbon, hydrogen, and oxygen.
• Define the term standard heat of formation, and identify the type of chemical reaction with which it is associated.
• Calculate the enthalpy change in a reaction occurring at constant pressure, given the standard enthalpies of formation of each reactant and product.

Electronic Structures of Atoms

Electromagnetic Radiation

• Describe the wave properties and characteristic speed of propagation of radiant energy (electromagnetic radiation).
• Use the relationship ln = c, which relates the wavelength (ll) and the frequency (n) of radiant energy to its speed (c).

Quantization of Energy

• Explain the essential feature of Planck's quantum theory, namely, the smallest increment, or quantum, of radiant energy of frequency, n, that can be emitted or absorbed is hn, where h is Planck's constant.
• Explain how Einstein accounted for the photoelectric effect by considering the radiant energy to be a stream of particle-like photons striking a metal surface.  In other words, you should be able to explain all the observations about the photoelectric effect using Einstein's model.

Line Spectra and the Bohr Model

• Explain the origin of the expression line spectra.
• List the assumptions made by Bohr in his model of the hydrogen atom.
• Explain the concept of an allowed energy state and how this concept is related to the quantum theory. 
• Calculate the energy differences between any two allowed energy states of the electron in hydrogen.
• Explain the concept of ionization energy.

Principles of Modern Quantum Theory

• Calculate the characteristic wavelength of a particle from a knowledge of its mass and velocity.
• Describe the uncertainty principle and explain the limitation it places on our ability to define simultaneously the location and momentum of a subatomic particle, particularly an electron.
• Explain the concepts of orbital, electron density, and probability as used in the quantum-mechanical model of the atom.   Explain the physical significance of Y².
• Describe the quantum numbers, n, l, m, used to define an orbital in an atom and list the limitations placed on the values each may have.
• Describe the shapes of the s, p, and d orbitals.

Energies of Orbitals in Many-Electron Atoms

• Explain  why electrons with the same value of principal quantum number (n) but different values of the azimuthal quantum number (l) possess different energies.

Electronic Structure of Many-Electron Atoms

• Explain the concepts of electron spin and the electron spin quantum number.
• State the Pauli exclusion principle and Hund's rule, and illustrate how they are used in writing the electronic structures of the elements.
• Write the electron configuration for any element.
• Write the orbital diagram representation for electron configurations of atoms.

The Periodic Table:  Periodic Arrangement of Electron Configurations and Valence Electrons Electromagnetic Radiation

• Describe what we mean by the s, p, d, and f blocks of elements.
• Write the electron configuration and valence electron configuration for any element once you know its place in the periodic table.

Periodic Properties of the Elements

Atomic Properties:  Atomic Size, Ionization Energy and Electron Affinity

• Explain the effect of increasing nuclear charge on the radial density function in many-electron atoms.
• Explain the variations in bonding atomic radii among the elements, and predict the relative sizes of atoms based on their positions in the periodic table.
• Explain the observed changes in values of the successive ionization energies for a given atom.
• Explain the general variations in first ionization energies among the elements, and relate these variations to variations in atomic radii.
• Explain the variations in electron affinities among the elements.

Overview:  Metals and Nonmetals

• Describe the periodic trends in metallic and nonmetallic behavior.
• Describe the general differences in chemical reactivity between metals and nonmetals.

Group Trends Exemplified:  The Active Metals

• Describe the general physical and chemical  behavior of the alkali metals and alkaline earth metals, and explain how their chemistry relates to their position in the periodic table.
• Write balanced equations for the reaction of hydrogen with metals to form metal hydrides. 
• Write balanced equations for simple reactions between the active metals (groups 1A and 2A) and the nonmetals in groups 6A and 7A.

Group Trends Exemplified:  Selected Nonmetals

• Write balanced equations for the reaction of hydrogen with non-metals such as oxygen and chlorine.
• Describe the allotropy of oxygen.
• Explain the dominant chemical reactions of oxygen and relate this behavior to it position in the periodic table.
• Describe the physical states and colors of the halogens, and explain the trends in reactivity with increasing atomic number in the family.
• Explain the very low chemical reactivity of the noble gas elements.

Basic Concepts of Chemical Bonding

Lewis Symbols:  Octet Rule

• Determine the number of valence electrons for any atom, and write its Lewis symbol.
• Recognize when the octet rule applies to the arrangement of electrons in the valence shell for an atom.

Ionic Bonding:  Energy, Ions, and Ionic Size

• Describe the origin of the energy terms that lead to stabilization of ionic lattices.
• Predict on the basis of the periodic table the probable formulas of ionic substances formed between common metals and nonmetals.
• Describe how the radii of ions relates to those of atoms.
• Explain the concept of an isoelectronic series and the origin of changes in ionic radius within a series.

The Lewis Model for Covalent Bonding

• Describe the basis of the Lewis theory, and predict the valence of common nonmetallic elements from their positions in the periodic table.
• Be able to describe a covalent bond in terms of sharing of electron density between bonded atoms.
• Describe the formation of a covalent compound using Lewis symbols.
• Be able to look at a Lewis structure and determine if it properly fits the Lewis model.
• Describe a single, double, and triple covalent bond.

Bond Polarity, Electronegativity, and Nomenclature

• Explain the significance of electronegativity and in a general way relate the electronegativity of an element to its position in the periodic table.
• Predict the relative polarities of bonds using either the periodic table or electronegativity values.
• Name a binary compound given its chemical formula or write the chemical formula given its chemical name.

Drawing Lewis Structures

• Write the Lewis structures for molecules and ions containing covalent bonds, using the periodic table.
• Write resonance forms for molecules or polyatomic ions that are not adequately described by a single Lewis structure.

Exceptions to the Octet Rule

• You should be able to write the Lewis structures for molecules and ions containing covalent bonds that have an odd number of electrons, a deficiency of electrons, or an expanded octet.

Strengths of Covalent Bonds

• You should be able to relate bond enthalpies to bond strengths and use bond enthalpies to estimate DH for reactions.

Molecular Geometry and Bonding Theories

VSEPR Model:  A Tool for Predicting Molecular Structure and Dipole Moments

• Relate the number of electron domains in the valence shell of an atom in a molecule to the geometrical arrangement around the atom.
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Curriculum Guide for NCHS Chemistry By Objective Goals Content Honors/Advanced COMPETENCY GOAL 1: The learner will build an understanding of the structure and properties of matter. (30% of curriculum) 1.01 Summarize the development of current atomic theory. (2%) John Dalton's atomic theory. J. J. Thomson - discovery of the electron E. Rutherford gold foil experiment, nucleus R. A. Millikan - charge on the electron N. Bohr - Hydrogen spectrum and electron arrangement (No names will be tested)   1.02 Examine the nature of atomic structure: 1.021 Protons. 1.022 Neutrons. 1.023 Electrons. 1.023 Atomic mass. 1.024 Atomic number. 1.025 Electron configuration. 1.026 Energy levels. 1.027 Isotopes. (4%) Properties of sub-atomic particles: relative mass, charge, and location in atom Symbols: A and Z, Principle quantum numbers; s, and p sublevels; Electron configuraton Orbitals notation using up/down arrows for opposite spin Valence electrons Lewis electron dot diagrams for atoms Isotope notation:E (ie.U or U-238) Identify isotopes by mass and atomic number Quantum numbers: azimuthal, magnetic and spin Computation of energies and wavelengths in H spectrum. 1.03 Apply the language and symbols of chemistry (4%) Binary nomenclature: Stock system for metal-nonmetal compounds and Greek prefix system for nonmetal-nonmetal compounds. Stock system for compounds with polyatomic ions. State symbols: (s), (l), (g) Name the 6 strong acids and acetic. Arrows indicating reactions and equilibria. Organic nomenclature, functional groups and named reactions. 1.04 Identify substances using their physical properties: 1.041 Melting points. 1.042 Boiling points. 1.043 Density. 1.044 Color. 1.045 Solubility. (4%) Identify substances using their physical properties. Students should be able to read and apply information from the reference tables.   1.05 Analyze and explain the nature and behavior of the atomic nucleus including radioactive isotopes and their practical application. (4%) Characteristics of alpha, beta, gamma radiation: Relative masses, charges, symbols, penetrating ability; Shielding: air (alpha), metal (beta), and distance (qualitative use of inverse square law). Concepts of half-life, fission, and fusion. Uses: dating, cancer therapy, smoke detectors, imaging. Decay equations. Inverse square law calculations. 1.06 Analyze the basic assumptions of kinetic molecular theory and its applications: 1.061 Ideal Gas Equation. 1.062 Combined Gas Law. 1.063 Graham’s Law. 1.064 Dalton’s Law of Partial Pressures. (4%) Five assumptions of KMT: Avogadro's Law, PV=nRT, Boyle's Law, Charles' Law, P1V1/T1=P2V2/T2, 1 mole of any gas at STP = 22.4 L Differentiate between real and ideal gases (factors nor calculations) Graham's Law Pt=P1+P2+ ... ; collecting a gas over water and vapor pressure of water. Calculations of KE or speeds of molecules, Maxwell's distribution Calculate molecular weight from effusion of gases 1.07 Assess the structure of compounds relating bonding and molecular geometry to chemical and physical properties; 1.071 Ionic bonds. 1.072 Covalent bonds. 1.073 Metallic bonds. (3%) Electronegativity general trend - predict nature of bond. Ion formation and stable arrangements (i.e. inert gas structure) prediction of physical properties based on bonding (melting point etc.) Lewis structures including single, double, triple bonds VSEPR Theory: Geometry: linear, bent, trigonal planar and tetrahedral, trigonal pyramidal. Polar / nonpolar bonds, polar / nonpolar molecules and solubility in polar or nonpolar solvents. ("like dissolves like"). Include intermolecular forces to explain polarity. Resonance. Formal charge calculations. Geometries: trigonal bipyramidal, octahedral. Hybrid orbital theory. Molecular orbital theory. COMPETENCY GOAL 2: The learner will build an understanding of regularities in chemistry. ( 36% of curriculum) 2.01 Analyze periodic nature of trends in chemical properties and examine the use of the Periodic Table to predict properties of elements; 2.011 Symbols. 2.012 Groups(families). 2.013 Periods. 2.014 Transition elements. 2.015 Ionization energy. 2.016 Atomic and ionic radii. 2.017 Electronegativity (5%) Define family (group) and period. Location on PT of alkali metals, alkaline earth metals, transition metals, rare earth metals, metalloids, halogens, inert gases. Also s, p, and d block elements. General trends of electronegativity, and ionization energy. Use PT to predict chemical and physical properties as well as charge of ions. General trends in atomic and ionic radii, Relate periodicity to electron configurations. Students will always have PT to use.   2.02 Analyze the mole concept and Avogadro's number and use them to calculate: 2.021 Mole to molecule. 2.022 Mass to moles. 2.023 Volume of a gas to moles. 2.024 Solution concentrations. (5%)   Conversion factors using moles, mole-mole, mole-mass, mass-mass. 1 mole of any gas at STP = 22.4 L Molarity. Limiting factors, theoretical and actual yields. Gravimetric and volumetric analysis. Determine empirical and molecular formulas. Normality. % concentration. Molality. 2.03 Identify various types of chemical equations and balance those equations: 2.031 Single replacement. 2.032 Double replacement. 2.033 Decomposition. 2.034 Synthesis. 2.035 Combustion. (7%) Use references table on reaction types to identify reaction types and predict products. Use activity series for single replacement. Use solubility table and/or solubility rules for double replacement. Write ionic and net ionic equations. Arrhenius acid/base neutralization reactions.   2.04 Calculate quantitative relationships in chemical reactions. (stoichiometry) (7.5%) stoichiometry mole-mole problems mass-mass problems mass-volume problems volume-volume problems gas laws and PV=nRT molarity   2.05 Identify the indicators of chemical change: 2.051 Formation of a precipitate. 2.052 Evolution of a gas. 2.053 Color change. 2.054 Absorption or release of heat. (4%) Recognize occurrence of reaction based on indicators of change such as formation of precipitate, evolution of a gas, color change and/or energy changes. Use the solubility rules and activity series in reference materials to predict the outcome of reactions.   2.06 Track the transfer of electrons in oxidation/reduction reactions and assign oxidation numbers: 2.061 Identify the oxidizing and reducing agents 2.062 Assess practical applications of oxidation and reduction reactions. (4%) Using PT and ion chart, assign oxidation state for each element in a compound. Show transfer of electrons by writing simple half reactions. Only simple metal/metallic ions will be tested. Determine Voltage calculations. Identify the element or ion oxidized, element or ion reduced, oxidizing agent, and reducing agent. Know that redox reactions occur in batteries, combustion, corrosion, and electroplating. Redox equation balancing. Primary vs secondary cells. COMPETENCY GOAL 3: The learner will build an understanding of energy changes in chemistry. (18% of curriculum) 3.01 Observe and interpret changes (emission/absorption) in electron energies in the hydrogen atom including the quantized levels and their relationship to atomic spectra: 3.011 Electromagnetic radiation. 3.012 Light. 3.013 Photons. (3%) Hydrogen spectrum and Bohr model, electron transfer between "orbits" and relation to energy given off as light. Use reference table. Use electromagnetic spectrum to relate wavelength and energy. Use equations only as illustration of relationship between energy and wavelength. c= fl, E=hf. Particle and wave nature of light. No calculation of wavelength between two Bohr orbits. 3.02 Analyze the law of conservation of energy, energy transformation, and various forms of energy involved in chemical reactions. (5%) Connect to 3.04 - calorimetry, calculations of heat based on temperature change of a quantity of water, q=mcDT, definitions of enthalpy, exothermic, endothermic, heats of reaction and stoichiometry. Energy vs pathway diagram showing energy of reactants, energy of products, enthalpy change, activation energy for exothermic and endothermic reactions. Heating and cooling curves. Phases Diagrams Hess's law, enthalpy calculations, heats of formation. 3.03 Compare and contrast the nature of heat and temperature. (4%) Temperature as measure of average kinetic energy of molecules. Heat as q=mcDT, energy transferred from hot to cold. Specific heat Specific heat.   3.04 Analyze calorimetric measurement in simple systems and the energy involved in changes in state. (5%) Calorimetry applications. Heating curve for ice and or water showing plateaus at phase changes, include heat of fusion, vaporization for water.   3.05 Analyze the relationship between energy transfer and disorder in the universe: 3.051 Nuclear. 3.052 Fossil fuels, Solar, Alternative sources. (2%) General knowledge of how a nuclear reactor works. Describe energy sources and the pros/cons of each energy source. Definition of entropy and its implications. Calculations of entropy, enthalpy or Gibbs free energy. COMPETENCY GOAL 4: The learner will build an understanding of equilibrium and kinetics. (16% of curriculum) 4.01 Explain the dynamics of physical and chemical equilibria: 4.011 Phase changes. 4.012 Forward and reversible reactions. (3%) Understand ice/water and water/vapor equilibrium. Understand that some reactions don't go to completion, and an equilibrium is established. Write equilibrium express but no calculations. Phase diagrams. Equilibrium expressions or Calculations. Triple point. 4.02 Explain the factors that alter the equilibrium in a chemical reaction. (4%) Le Chatelier's Principle - concentration, pressure, temperature. Use the terms "shift to the right, shift to the left" or make more produce, make more reactant to describe changes. Equilibrium expression as it relates to weak and strong acids but no calculations. Equilibrium expressions or calculations. 4.03 Assess reaction rates and factors that affect reaction rates. (4%) Rate as change in concentration (or pressure) as function of time. Factors affecting rate: concentration, pressure, temperature, catalyst (lower activation energy). Reaction order, time/concentration equations. Reaction mechanisms or rate determining steps. 4.04 Compare and contrast the nature, behavior, concentration, and strength of acids and bases: 4.041 Acid-base neutralization. 4.042 Degree of dissociation or ionization. 4.043 Electrical conductivity. 4.044 pH. (5%) Properties of acids and bases; Strength vs concentration; strength of weak acids and bases - partial dissociation. Arrhenius and Bronsted.Lowery theories Acid/base titration and stoichiometry. (nMV = nMV) Weak Vs strong acids pH scale and calculations with pH=-log[H+], pOH=-log[OH-], pH + pOH = 14, [OH]= 10-pOH[H+]=10-pH Buffer systems. (qualitative discussion only) Acid-base equilibria equations. Lewis theory. Ka and Kb calculations

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