Curriculum Mapping 2005 Content: Chemistry

District: Harrisville Central School Grade: 11

C. Ritchie

Months	Learning Outcome	Content	Standards	Skills	Assessment
September Atom Nucleus Electron Proton Neutron Isotopes Mass Number Atomic Mass Protium Deuterium Tritium AMU Gram Atomic Mass Avogadro’s Number Quanta Principal Energy Level Sublevels Wave-Mechanical Model Ground state Excited state Spectral Lines Radiant energy Spectroscope Valence electrons Kernel Electron dot symbols	3.1a The modern model of the atom has evolved over a long period of time through the work of many scientists. 3.1b Each atom has a nucleus, with an overall positive charge, surrounded by negatively charged electrons 3.1c Subatomic particles contained in the nucleus include protons and neutrons. 3.1d The proton is positively charged and the neutron has no charge. The electron is negatively charged. 3.1e Protons and electrons have equal but opposite charges. 3.1f The mass of each proton and each neutron is approximately equal to one atomic mass unit. An electron is much less massive than a proton or a neutron. 3.1h In the wave-mechanical (electron cloud) model, the electrons are in orbitals, which are regions of most probable electron location (ground state). 3.1i Each electron in an atom has its own distinct amount of energy	Atomic Concepts	Matter is made of particles whose properties determine the observable characteristics of matter and its reactivity Explain the properties of matter in terms of the arrangement and properties of the atoms that compose them	Safety rules, math skills, identifying properties of matter (melting point/ freezing point), measurement, Lab write up, CEI , significant figures, scientific notation, order of magnitude The student should be able to: relate experimental evidence (given in the introduction of Key Idea 3) to models of the atom (3.1ii) use models to describe the structure of an atom (3.1i) determine the number of protons or electrons in an atom or ion when given one of these values (3.1iii) calculate the mass of an atom, the number of neutrons or the number of protons, given the other two values (3.1iv)	Bell Ringer Unit test based on Regents exam Atom project Quizzes Labs Lab Safety Lab Techniques Spectra of Elements Flame Tests

Months	Learning Outcome	Content	Standards	Skills	Assessment
Sept. Cont	3.1j After an electron in an atom gains a specific amount of energy, the electron is at a higher energy level (excited state). 3.1k When an electron returns from a higher energy state to a lower energy state, energy is emitted. This emitted energy corresponds to a specific wavelength in the electromagnetic spectrum. 3.1l Wavelengths can be used to identify a substance. Each kind of atom or molecule can gain or lose energy in discrete amounts, and thus can absorb or emit energy only at wavelengths corresponding to these amounts. 3.1m In general, the outermost electrons in an atom are called the valence electrons. The number of valence electrons determines the chemical properties of an element. The chemical reactivity of an atom is dependent on its size. 3.1n Atoms of an element that contain the same number of protons but a different number of neutrons are called isotopes of that element. 3.1o The average atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.			distinguish between ground state and excited state electron configurations, e.g., 2-8-2 vs. 2- 7-3 (3.1v) identify an element by comparing its bright-line spectrum to given spectra (3.1vi) draw a Lewis electron-dot structure of an atom (3.1viii) distinguish between valence and non-valence electrons, given an electron configuration, e.g., 2-8-2 (3.1vii) given an atomic mass, determine the most abundant isotope (3.1xi) calculate the atomic mass of an element, given the masses and ratios of naturally occurring isotopes (3.1xii)

Months	Learning Outcome	Content	Standards	Skills	Assessment
Sept –Oct Periodic Table Isotopes Metals Nonmetals Metalloids Alloys Chemical Properties Physical Properties Periods Rows Groups Families Allotropes Atomic Radius Kernel Electronegativity Ionization Energy Shielding effect Alkali Metals Alkaline Earth Metals Halogens Noble Gases Monoatomic gases Diatomic	3.1y The placement or location of elements on the periodic table gives an indication of physical and chemical properties of that element. The elements on the periodic table are arranged in order of increasing atomic number. 3.1g The number of protons in an atom (atomic number) identifies the element. The sum of the protons and neutrons in an atom (mass number) identifies an isotope. Examples of common notations that represent isotopes include: 6^14C, ^14C, carbon-14, C-14. 3.1v Elements can be classified as metals, nonmetals, metalloids, and noble gases.) 3.1w Elements can be differentiated by their physical properties. Physical properties of substances, such as density, conductivity, malleability, solubility, and hardness, differ between elements. 3.1x Elements can be differentiated by chemical properties. Chemical properties describe how an element behaves during a chemical reaction. 5.2f Some elements exist as allotropes. Allotropes are two or more forms of the same element that differ in their molecular or crystalline structure, and hence in their properties.	Periodic Table	Matter is made of particles whose properties determine the observable characteristics of matter and its reactivity Explain the properties of materials in terms of the arrangement and properties of the atoms that compose them.	The student should be able to: explain the placement of an unknown element in the Periodic Table based on its properties (3.1xvi) interpret and write isotopic notation (3.1x) classify elements as metals, nonmetals, metalloids, or noble gases by their properties (3.1xiii) describe the states of the elements at STP (3.1xviii) determine the group of an element, given the chemical formula of a compound, e.g., XCl or XCl2 (3.1xv) compare and contrast properties of elements within a group or a period for Groups 1, 2, 13- 18 on the Periodic Table (3.1xiv)	Quick quiz Bell Ringer Homework checks Wipe board work Lab report Teacher observation Student responses to teacher questions Rubrics for evaluating projects Labs: Periodic Law Lab

Months	Learning Outcome	Content	Standards	Skills	Assessment
Cont.	3.1z For Groups 1,2, and 13-18, elements within the same group have the same number of valence electrons (helium is an exception) and therefore similar reactivity. 3.1aa The succession of elements within the same group demonstrates characteristic trends, e.g., differences in atomic radius, ionic radius, electronegativity, first ionization energy. These trends can be explained in terms of atomic structure (size). 3.1bb The succession of elements across the same period demonstrates characteristic trends, e.g., differences in atomic radius, ionic radius, electronegativity, first ionization energy. These trends can be explained in terms of atomic structure (nuclear charge

Months	Learning Outcome	Content	Standards	Skills	Assessment
October – November Compound Chemical Formula Empirical Formula Molecular Formula Structural Formula Polyatomic Ion Binary Acid Ternary Acid Binary Salt Coefficients Aqueous Solution Gram Atomic Mass Gram Formula Mass Mole Stoichiometry Percent Composition Synthesis Composition Decomposition Analysis Single Replacement Double Replacement Precipitation Reaction	3.1cc compound is a substance composed of two or more different elements that are chemically combines in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific formula and assigned a name based on the IUPAC system. 3.1dd A chemical formula can be represented as an empirical formula, a structural formula, or a molecular formula. 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different than the molecular formula, which is the actual ratio of atoms in a molecule of that compound. 3.3e The molecular formula is a whole-number multiple of the empirical formula. 3.3a In all reactions there is a conservation of mass, energy, and charge. 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction. 3.3f The formula of a substance is the sum of the masses of its atom. The gram-formula mass of a substance equals one mole of that substance.	Moles and Stoichiometry	Matter is made of particles whose properties determine the observable characteristics of matter and its reactivity Apply the principle of conservation of mass to chemical reactions. Use atomic and molecular models to explain common chemical reactions	The student should be able to: determine the molecular formula, given the empirical formula and molecular mass (3.3vii) determine the empirical formula from a molecular formula (3.3v) interpret balanced chemical equations in terms of conservation of matter and energy (3.3ii) balance equations, given the formulas for reactants and products (3.3i) interpret balanced chemical equations in terms of conservation of matter and energy (3.3ii) create and use models of particles to demonstrate balanced equations (3.3iii) calculate simple mole-mole stoichiometry problems, given a balanced equation (3.3iv) calculate the formula mass and the gram-formula mass (3.3viii) determine the number of moles of a substance, given its mass (3.3ix) determine the mass of a given number of moles of a substance (3.3vi)	Quick quiz Bell Ringer Homework checks Wipe board work Check problem Homework pairs Lab report Mini-test Student responses to teacher questions Rubrics for evaluating projects Labs: Formulas & Oxidation Numbers Quantitative Determination of an Empirical Formula Mole Relationship in a Chemical Reaction Hydrated Crystals

Months	Learning Outcome	Content	Standards	Skills	Assessment
Cont.	3.3g The percent composition by mass of each element in a compound can be calculated mathematically. 3.2b There are many types of chemical reactions, e.g., synthesis, decomposition, single replacement, and double replacement.			identify types of chemical reactions (3.2ii)

Months	Learning Outcome	Content	Standards	Skills	Assessment
November Chemical Bond Inert Ionic Bond Ionic Solid Covalent Bond Nonpolar Covalent Polar Covalent Dipole Coordinate Covalent Molecular Substances Molecular Solid Network Solid Metallic Bond Ionic Radius Dipoles Hydrogen Bonding Van Der Waals Forces Weak Intermolecular Attractions Molecule – Ion Attractions Polyatomic Ions	5.2g Compounds differ in composition as well as chemical and physical properties. Two major categories of compounds are ionic compounds and molecular (covalent) compounds. 5.2a Chemical bonds are formed when valence electrons are: - Transferred from one atom to another (ionic) - Shared between atoms (covalent) - Mobile within a metal (metallic) 5.2e In a multiple covalent bond, more than one pair of electrons is shared between two atoms. Unsaturated organic compounds contain at least one double or triple band. 5.2l Molecular polarity can be determined by the shape and the distribution of charge. Examples of symmetrical (nonpolar) molecules include CO2, CH4 and diatomic elements. Examples of asymmetric (polar) molecules include HC1, NH3, and H2O. 5.2c When an atom gains one or more electrons, it becomes a negative ion and its radius increases. When an atom loses one or more electrons, it becomes a positive ion and its radius decreases	Chemical Bonding	Energy and matter interact through forces that result in changes in motion. Explain chemical bonding in terms of the behavior of electrons Explain the properties of materials in terms of the arrangement and properties of the atoms that compose them.	The student should be able to: distinguish among ionic, molecular, and metallic substances, given their properties (3.1xix) demonstrate bonding concepts using Lewis dot structures representing valence electrons: transferred (ionic bonding); shared (covalent bonding); in a stable octet (5.2i) determine the noble gas configuration an atom will achieve when bonding (5.2iv) demonstrate bonding concepts, using Lewis dot structures representing valence electrons: transferred (ionic bonding); shared (covalent bonding); in a stable octet (5.2i distinguish between nonpolar covalent bonds (two of the same nonmetals) and polar covalent bonds (5.2v)	Quick quiz Bell Ringer Homework checks Wipe board work Check problem Homework pairs Mini-lab Lab report Unit tests Test revisions Chapter tests Teacher observations Student responses to teacher questions Rubrics for evaluating projects Labs: Conductivity and Chemical Bonding Its all in the Shape Christmas Tree Lab Boiling Point of Sucrose

Months	Learning Outcome	Content	Standards	Skills	Assessment
Continued	5.2h When a bond is broken, energy is absorbed. When a bond is formed, energy is released. 5.2b Atoms tend to bond so that a stable valence electron configuration, like that of a noble gas, is achieved. 5.2d Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements and compounds. 5.2f Electronegativity indicates how strongly an atom of an element attracts electrons in a chemical bond. 5.2i Electronegativity values are assigned according to arbitrary scales. 5.2k The electronegativity difference between two bonded atoms is used to assess the degree of polarity in the bond.

Months	Learning Outcome	Content	Standards	Skills	Assessment
Dec-Jan Extensive Properties Intensive Properties Solid Phase Super-cooled liquid Crystals Strength Malleability Amorphous Vapor Pressure Evaporation Boiling Point Exothermic Reaction Endothermic Reaction Melting Point Heat of Fusion Sublimation Heat of Vaporization Substance Compound Mixture Homogeneous Uniform Solution Heterogeneous Filtration Distillation Density Particle Size Magnetic Separation	3.1r Matter is classified as a substance or a mixture of substances. 3.1jj The three phases of matter, i.e., solids, liquids, and gases, have different properties 3.1s A substance (element or compound) has a constant composition and constant properties throughout a given sample, and from sample to sample 3.1u Elements are substances that are composed of atoms that have the same atomic number. Elements cannot be broken down by chemical change. 3.1t A mixture is not a substance because it is made up of two or more different elements and/or compounds. The proportions of components in a mixture can be varied. Each component in a mixture retains its original properties. 3.1mm When different substances are mixed together, a homogeneous or heterogeneous mixture is formed. Mixtures are composed of two or more different substances that can be separated by physical means. 3.1nn Differences in physical properties such as mass, particle size, molecular polarity, boiling point and freezing point, and solubility permit physical separation of the components of the mixture	PHYSICAL BEHAVIOR OF MATTER	Observe and describe transmission of various forms of energy, and explain heat in terms of kinetic molecular theory. Explain the properties of materials in terms of the arrangement and properties of the atoms that compose them. Use kinetic molecular theory to explain rates of reactions and the relationship among temperature, pressure and volume. Explain heat in terms of kinetic molecular theory. Explain heat in terms of kinetic molecular theory. Explain chemical bonding in terms of the behavior of electrons.	The student should be able to: use a simple particle model to differentiate among properties of a solid, a liquid, and a gas (3.1xxii) use particle models/diagrams to differentiate among elements, compounds, and mixtures (3.1xxxvi) describe the process and use of filtration, distillation, and chromatography in the separation of a mixture (3.1xxiv) interpret and construct solubility curves (3.1xxv) use solubility curves to distinguish among saturated, supersaturated and unsaturated solutions (3.1xxviii) apply the adage "like dissolves like" to real-world situations (3.1xxvi) describe the preparation of a solution, given the molarity (3.1xxx) interpret solution concentration data (3.1xxx) calculate solution concentrations in molarity (M), percent mass, and parts per million (ppm) (3.1xxix) distinguish between heat energy and temperature in terms of molecular motion and amount of matter (4.2i)	Quick quiz Bell Ringer Homework checks Wipe board work Check problem Homework pairs Mini labs Lab report Unit checks Test revisions Chapter test Teacher observations Student responses to teacher questions Rubrics for evaluating projects Labs: Charle’s Law Heating Curve of Water

Months	Learning Outcome	Content	Standards	Skills	Assessment
Continued Boiling Point Freezing Point Soluability Solutions Solvent Solute Solvation Hydration Dissolution Molarity Miscible Immiscible Saturated Solution Unsaturated Solution Parts per million Colligative Properties Molality Boiling Point Elevation Freezing Point depression Electrolytes Potential Energy Kinetic Energy Heat Temperature Thermometer Celsius Kelvin Absolute zero Specific Heat	3.1oo A solution is a homogeneous mixture of a solute dissolved in a solvent. The solubility of a solute in a given amount is dependent on the temperature, pressure, and the chemical natures of the solute and solvent. The concentration of a solution is expressed in molarity (M), percent by volume, percent by mass, or parts per million (ppm). 3.1pp The addition of a nonvolatile solute to a solvent causes the boiling point of the solvent to increase and the freezing point of the solvent to decrease. The greater the concentration of particles, the greater the effect. 4.1a Energy can exist in different forms, e.g., chemical, light, heat, nuclear. 4.2a Heat is a transfer of energy (usually thermal energy) from a body of higher temperature to a body of lower temperature. Thermal energy is the energy associated with the random motion of atoms and molecules. 4.2b Temperature is a measurement of the average kinetic energy of the particles in a sample of material. Temperature is not a form of energy. 3.4b Kinetic molecular theory (KMT) for an ideal gas states: - All particles are in random, constant, straight-line motion. - Gas molecules are separated by great distances relative to their size; the volume of the gas molecules is considered negligible.		Explain the properties of materials in terms of the arrangement and properties of the atoms that compose them. Use atomic and molecular models to explain common chemical reactions. Observe and describe transmission of various forms of energy.	qualitatively interpret heating and cooling curves in terms of changes in kinetic and potential energy, heat of vaporization, heat of fusion, and phase changes (4.2iii) distinguish between heat energy and temperature in terms of molecular motion and amount of matter (4.2i) explain phase changes in terms of the changes in energy and intermolecular distance (4.2ii)

Months	Learning Outcome	Content	Standards	Skills	Assessment
Continued Ideal gas Gases Boyle’s Law Charle’s Law STP Combined Gas Law Avagadro’s Hypothesis	- The molecules have no attractive forces between them. - Collisions between gas particles may result in the transfer of energy between gas particles 3.4d Particles are in constant motion except at absolute zero (zero kelvin). 3.4c Kinetic molecular theory describes the relationship of pressure, volume, temperature, velocity, and frequency and force of collisions. 4.2c The concepts of kinetic and potential energy can be used to explain physical processes that include: fusion (melting), solidification (freezing), vaporization (boiling, evaporation), condensation, sublimation, and deposition. 3.2a A physical change results in the rearrangement of existing particles in a substance. A chemical change results in the formation of different particles with changed properties.) 4.1b Chemical and physical changes can be exothermic or endothermic. 3.1ii The structure and arrangement of particles and their interactions determine the physical state of a substance at a given temperature and pressure. 5.2m Intermolecular forces created by the unequal distribution of electrons result in varying degrees of attraction between molecules. Hydrogen bonding is an example of a strong intermolecular force.

Months	Learning Outcome	Content	Standards	Skills	Assessment
Feburary Kinetics Collision Theory Effective Collision Activation Energy Activated Complex Catalyst Heat of Reaction Delta H Enthalpy Surface Area Equilibrium Phase Equilibrium Dynamic Equilibrium Solution Equilibrium Chemical Equilibrium Le Chetlier’s Principal Spontaneous Reactions Energy Changes Entropy Changes Common Ion Effect	4.1e The stability of a compound is dependent on the amount of energy absorbed or released during the formation of the compound from its elements. 3.4f Collision theory states that a reaction is most likely to occur if reactant particles collide with the proper energy and 3.4f The rate of a chemical reaction is the change in concentration of a reactant or product per unit time. 3.4g The rate of chemical reaction depends on several factors: temperature, concentration, and nature of reactants, surface area, and the presence of a catalyst. 3.4i Some chemical and physical changes can reach equilibrium. 3.4j At equilibrium the rate of the forward reaction equal the rate of the reverse reaction while the measurable quantities of reactants and products remain constant. 3.4k LeChatelier’s principle can be used to predict the effect of stress (change in pressure, volume, concentration, and temperature) on a system at equilibrium. 4.1c Energy released or absorbed by a chemical reaction can be represented by a potential energy diagram. 4.1d Energy released or absorbed during a chemical reaction is equal to the difference between the potential energy of the products and the potential energy of the reactants.	Kinetics & Equilibrium	Matter is made of particles whose properties determine the observable characteristics of matter and its reactivity Observe and describe transmission of various forms of energy Explain the properties of materials in terms of the arrangement and properties of the atoms that compose them. Use kinetic molecular theory to explain rates of reactions and the relationships among temperature, pressure and volume of a substance.	The student should be able to: use collision theory to explain how various factors, such as temperature, surface area, and concentration, influence the rate of reaction (3.4vi) identify examples of physical equilibria as solution equilibrium and phase equilibrium, including the concept that a saturated solution is at equilibrium (3.4 vii) describe the concentration of particles and rates of opposing reactions in an equilibrium system (3.4iv) qualitatively describe the effect of stress on equilibrium, using LeChatelier's principle (3.4v) read and interpret potential energy diagrams: PE of reactants and products, activation energy (with or without a catalyst), heat of reaction (4.1ii) compare the entropy of phases of matter (3.1xxiii)	Quick quiz Bell Ringer Homework checks Wipe board work Mini- lab Lab report Mini-test Teacher observations Student responses to teacher questions Rubrics for evaluating projects Lab: Chemical Equilibrium A Study in Chemical Equilbrium

Months	Learning Outcome	Content	Standards	Skills	Assessment
Continued	3.4h A catalyst provides an alternate reaction mechanism, which has a lower activation energy than an uncatalyzed reaction. 3.1kk Entropy is a measure of the randomness or disorder of a system. A system with greater disorder has greater entropy. 3.1ll Systems in nature tend to undergo changes in energy and entropy.

Months	Learning Outcome	Content	Standards	Skills	Assessment
March Organic Chemistry Organic Inorganic compound Conductivity Melting Point Organic Reaction Tetrahedron Isomer Saturated Unsaturated Single Bond Double Bond Triple Bond Hydrocarbon Homologous series Alkanes Alkenes Alkynes Benzene Toluine Alkyl Group Functional Group Alcohols Organic Acids Aldehydes Ketones Ethers Esters Amines Amino Acids Substitution Addition Fermentation Combustion	3.1ee Organic compounds contain carbon atoms which bond to one another in chains, rings, and networks to form a variety of structures (polymers, oils, and other large molecules). (3.1ee) 3.1gg Functional groups impart distinctive physical and chemical properties to organic compounds. 3.1ff Hydrocarbons, organic acids, alcohols, esters, amines, amides, and amino acids are categories of organic molecules that differ in their structural formulae as a result of different functional groups. 3.1hh Hydrocarbons, organic acids, alcohols, and esters are named using the IUPAC system. The IUPAC system. The IUPAC system provides a method of distinguishing among isomers of organic compounds. 3.1ii Unsaturated organic compounds contain at least one 3.2c Addition, hydrogenation, substitution, polymerization, esterification, fermentation, saponification, oxidation, and combustion are examples of organic reactions.	Organic Chemistry	Matter is made of particles whose properties determine the observable characteristics of matter and its reactivity Explain chemical bonding in terms of the behavior of electrons Explain the properties of materials in terms of the arrangement and properties of the atoms that compose them. Use atomic and molecular models to explain common chemical reactions.	The student should be able to: classify an organic compound based on its structural or condensed structural formula (3.1xvii) draw structural formulas for alkanes, alkenes, and alkynes containing a maximum of ten carbon atoms (3.1xxi) classify an organic compound based on its structural or condensed structural formula (3.1xvii) draw a structural formula with the functional group(s) on a straight chain hydrocarbon backbone, when given the correct IUPAC name for the compound (3.1xx) identify types of organic reactions (3.2iv) determine a missing reactant or product in a balanced equation (3.2iii)	Quick quiz Bell Ringer Homework checks Wipe board work Homework pairs Mini- lab Lab report Teacher observations Student responses to teacher questions Rubrics for evaluating projects Lab: Preparation of Organic Chemicals

Months	Learning Outcome	Content	Standards	Skills	Assessment
Continued Esterification Soponification Oxidation Polymerization Polymer Condensation Addition

Months	Learning Outcome	Content	Standards	Skills	Assessment
April Reduction Oxidation Oxidizing Agent Reducing Agent REDOX Composition Decompostion Single Replacement Corrosion Electrochemical Cells Electrodes Salt Bridge Electrolytic Cell Electrolysis Battery Electroplating	3.2d An oxidation-reduction (redox) reaction involves the transfer of electrons (e^-). 3.2e Reduction is the gain of electrons. 3.2f A half-reaction can be written to represent reduction. 3.2g Oxidation is the loss of electrons. Historically, oxidation was explained as the chemical combination with oxygen. 3.2h A half-reaction can be written to represent oxidation. 3.3b In a redox reaction there is a conservation of charge and mass. The number of electrons lost is equal to the number of electrons gained. 3.2i Oxidation numbers can be assigned to atoms and ions. Changes in oxidation numbers indicate oxidation and reduction. 3.2j Corrosion of metals, combustion of fuels, and spoilage of foods are examples of redox reactions. 3.2k In an electrochemical cell, oxidation occurs at the anode and reduction at the cathode. 3.2l A voltaic cell can operate without an outside energy source. 3.2m Batteries are practical applications of voltaic cells. 3.2n An electrolytic cell requires an outside energy source. 3.2a Refining of metals and electroplating are practical applications of voltaic cells	Oxidation-Reduction	Matter is made of particles whose properties determine the observable characteristics of matter and its reactivity Apply the principle of conservation of mass to chemical reactions. Use atomic and molecular models to explain common chemical reactions	The student should be able to: determine a missing reactant or product in a balanced equation 3.2iii) write and balance half-reactions for oxidation and reduction of free elements and their monatomic ions (3.2vi) compare and contrast voltaic and electrolytic cells (3.2ix) identify and label the parts of a voltaic cell (cathode, anode, salt bridge) and direction of electron flow, given the reaction equation (3.2vii) use an activity series to determine whether a redox reaction is spontaneous (3.2x) identify and label the parts of an electrolytic cell (anode, cathode) and direction of electron flow, given the reaction equation (3.2viii)	Quick quiz Bell Ringer Homework checks Wipe board work Check problem Homework pairs Lab report Unit checks Mini tests Test revisions Chapter test Teacher observation Student responses to teacher questions Rubrics for evaluating projects Lab: Redox Reactions

Months	Learning Outcome	Content	Standards	Skills	Assessment
May Electrolyte Acid Base Arrhenius Theory Bronstead-Lowery Theory Indicators Amphoteric Substance Neutralization Salts PH Titration	3.1uu Behavior of many acids and bases can be explained by the Arrhenius theory. Arrhenius acids and bases are electrolytes. 3.1qq An electrolyte is a substance which, when dissolved in water, forms a solution capable of conducting an electric current. The ability of a solution to conduct an electric current depends on the number of ions. 3.1vv Arrhenius acids yield H3O^+(hydronium ion) as the only positive ion in an aqueous solution 3.1ww Arrhenius acids react with active metals to produce hydrogen gas. 3.1xx Arrhenius bases yield OH^-(hydroxide ion) as the only negative ion in an aqueous solution. 3.1yy In the process of neutralization, an Arrhenius acid and an Arrhenius base react to form a salt and water. 3.1zz Titration is a laboratory process in which a volume of solution of known concentration is used to determine the concentration of another solution. 3.1rr An aqueous solution of a salt conducts electricity. The solution can have a pH greater than, equal to, or less than 7. 3.1ss The acidity or alkalinity of a solution can be measured by its pH value. The relative level o f acidity or alkalinity of a solution can be shown by using indicators.	Acids, Bases, & Salts	Matter is made of particles whose properties determine the observable characteristics of matter and its reactivity Explain the properties of materials in terms of the arrangement and properties of the atoms that compose them.	The student should be able to: given properties, identify substances as Arrhenius acids or Arrhenius bases (3.1xxxi) write simple neutralization reactions when given the reactants (3.1xxxiv) calculate the concentration or volume of a solution, using titration data (3.1xxxv) interpret changes in acid-base indicator color (3.1xxxiii) identify solutions as acid, base, or neutral based upon the pH (3.1xxxii)	Quick quiz Bell Ringer Homework checks Wipe board work Check problem Homework pairs Mini lab Lab report Teacher observation Student responses to teacher questions Rubrics for evaluating projects Lab: Acid – Base Titrations

Months	Learning Outcome	Content	Standards	Skills	Assessment
Continued	3.1tt The mathematical definition of pH is: pH=-log [H3O^+]. On the pH scale, each decrease of one unit of pH represents a ten-fold increase in hydronium ion concentration.

Months	Learning Outcome	Content	Standards	Skills	Assessment
May – June Half Life Radioactivity Alpha Decay Beta Decay Gamma Decay Emanation Accelerator Mass Defect Binding Energy Fission Fusion Deuterium Tritium Radioisotope Tracer	3.1p Isotopes of atoms can be stable or unstable. Stability of isotopes is based on the number of protons and neutrons in its nucleus. Some nuclei are unstable and spontaneously decay, emitting radiation. 4.4a Each radioactive isotope has a specific mode and rate of decay (half-life). 3.1q Spontaneous decay can involve the release of alpha particles, beta particles, or gamma radiation from the nucleus of an unstable isotope. These emissions differ in mass, charge, and penetrating and ionizing power. 5.3a Any change in the nucleus of an atom that converts it from one element to another is called transmutation. This can occur naturally or can be induced by the bombardment of the nucleus by high-energy particles. 4.4b Nuclear reactions can be represented by equations that include symbols which represent atomic nuclei (with the mass number and atomic number), subatomic particles (with mass number and charge), and/or emissions such as gamma radiation. 5.3b Energy released in a nuclear reaction (fission/fusion) comes from the fractional amount of mass converted into energy. Nuclear changes convert matter into energy according to E=mc^2.	Nuclear Chemistry	Matter is made of particles whose properties determine the observable characteristics of matter and its reactivity Explain the properties of materials in terms of the arrangement and properties of the atoms that compose them. Explain the benefits and risks of radioactivity Compare energy relationships within an atom’s nucleus to those outside the nucleus	The student should be able to: calculate the initial amount, the fraction remaining, or the halflife of a radioactive isotope, given two of the three variables (4.4i) determine decay mode and write nuclear equations showing alpha and beta decay (3.1ix) compare and contrast fission and fusion reactions (4.4ii) complete nuclear equations; predict missing particles from nuclear equations (4.4iii) identify specific uses of some common radioisotopes, such as: I-131 in diagnosing and treating thyroid disorders; C-14 to C-12 ratio in dating living organisms; U-238 to Pb-206 ratio in dating geological formations; Co-60 in treating cancer (4.4iv)	Quick quiz Bell ringer Homework checks Wipe board work Check problem Homework pairs Mini lab Lab report Teacher observation Student responses to teacher questions Rubrics for evaluating projects Lab: Half Life Energy of Fusion

Months	Learning Outcome	Content	Standards	Skills	Assessment
Continued	5.3c Energy released during nuclear reactions is much greater then the energy released during chemical reactions. 4.4d There are inherent risks associated with radioactivity and its uses, such as long-term storage and disposal of radioactive isotopes, and nuclear accidents. 4.4c Radioactive isotopes have many beneficial uses. Radioactive isotopes are used in medicine and industrial chemistry, e.g., radioactive dating, tracing chemical and biological processes, industrial measurement, detection and treatment of diseases.