Differences between Indiana Chemistry I Standards from 2000 and 2010

The new standards are more than just an organizational change as can be seen by the below comparison. The Indiana Department of Education came out with its initial proposal for Indiana's Science Standards in April of 2010. They are still tweaking it and will come out with a final version in March of 2011. The new changes will take affect in the 2011-2012 school year. 

2000 Indiana Standards	2010 Indiana Standards
C.1.1         Differentiate between pure substances and mixtures based on physical properties such as density, melting point, boiling point, and solubility.	C.1.1    Based on physical properties, differentiate between pure substances and mixtures.
C.1.2         Determine the properties and quantities of matter such as mass, volume, temperature, density, melting point, boiling point, conductivity, solubility, color, numbers of moles, and pH (calculate pH from the hydrogen-ion concentration), and designate these properties as either extensive or intensive.	C.1.2   Observe and describe chemical and physical properties of different types of matter and designate them as either extensive or intensive. C.8.4   Given the hydronium (H3O+) ion concentration in a solution, calculate the pH and vice versa.  Explain the meaning of these values.
C.1.3         Recognize indicators of chemical changes such as temperature change, the production of a gas, the production of a precipitate, or a color change.	C.1.3    Recognize observable indicators of chemical changes.
C.1.4         Describe solutions in terms of their degree of saturation.	Not expected
C.1.5         Describe solutions in appropriate concentration units (be able to calculate these units), such as molarity, percent by mass or volume, parts per million (ppm), or parts per billion (ppb).	C.7.3   Describe the concentration of solutes in a solution in terms of molarity.  Perform calculations using molarity, mass and volume.
C.1.6         Predict formulas of stable ionic compounds based on charge balance of stable ions.	C.3.4    Write chemical formulas for ionic compounds given their names and vice versa.
C.1.8         Use formulas and laboratory investigations to classify substances as metal or nonmetal, ionic or molecular, acid or base, and organic or inorganic.	C.3.5   Compare and contrast ionic, covalent network, metallic and polar and non-polar molecular crystals with respect to constituent particles, strength of bonds, melting and boiling points and conductivity; provide examples of each type. C.8.1   Use Arrhenius and Brønsted-Lowry definitions to classify substances as acids or bases.
C.1.9         Describe chemical reactions with balanced chemical equations.	C.4.2   Balance chemical equations using the law of conservation of mass and use them to describe chemical reactions.
C.1.10       Recognize and classify reactions of various types such as oxidation-reduction.	C.4.5   Describe, classify and give examples of various kids of reactions-synthesis (i.e., combination), decomposition, single displacement, double displacement and combustion.
C.1.11       Predict products of simple reaction types including acid/base, electron transfer, and precipitation.	C.4.1   Predict products of simple reactions such as synthesis, decomposition, single replacement and double replacement.
C.1.12       Demonstrate the principle of             conservation of mass through laboratory investigations.	C.1.6   Explain and apply the law of conservation of mass as it applies to chemical processes.
C.1.13       Use the principle of conservation of mass to make calculations related to chemical reactions. Calculate the masses of reactants and products in a chemical reaction from the mass of one of the reactants or products and the relevant atomic masses.	C.4.3   Given mass of the sample, use the mole concept to determine the number of moles and number of atoms or molecules in samples of elements and compounds.
C.1.14       Use Avogadro’s law to make mass-volume calculations for simple chemical reactions.	C.5.3   Given the equation for a chemical reaction involving one or more gases as reactants, products or both, calculate the volumes of gas when assuming the reaction goes to completion and the ideal gas law holds.
C.1.15       Given a chemical equation, calculate the mass, gas volume, and/or number of moles needed to produce a given gas volume, mass, and/or number of moles of product.	C.4.3   Given mass of the sample, use the mole concept to determine the number of moles and number of atoms or molecules in samples of elements and compounds. C.4.4   Using a balanced chemical equation, calculate the quantities of reactants needed and products made in a chemical reaction that goes to completion. C.5.3   Given the equation for a chemical reaction involving one or more gases as reactants, products or both, calculate the volumes of gas when assuming the reaction goes to completion and the ideal gas law holds.
C.1.16       Calculate the percent composition by mass of a compound or mixture when given the formula	C.4.7   Perform calculations to determine the composition of a compound or mixture when given the formula.
C.1.17       Perform calculations that demonstrate an understanding of the relationship between molarity, volume, and number of moles of a solute in a solution.	C.7.3   Describe the concentration of solutes in a solution in terms of molarity.  Perform calculations using molarity, mass and volume.
C.1.18       Prepare a specified volume of a solution of given molarity.	C.7.4   Prepare a specific volume of a solution of a given molarity when provided with a known solute.
C.1.19       Use titration data to calculate the concentration of an unknown solution.	C.8.5    From acid-base titration data, calculate the concentration of an unknown solution.
C.1.20       Predict how a reaction rate will be quantitatively affected by changes of concentration.	C.7.5   Explain how the rate of a reaction is qualitatively affected by changes in concentration, temperature, surface area and the use of a catalyst.
C.1.21       Predict how changes in temperature, surface area, and the use of catalysts will qualitatively affect the rate of a reaction.	C.7.5   Explain how the rate of a reaction is qualitatively affected by changes in concentration, temperature, surface area and the use of a catalyst.
C.1.22       Use oxidation states to recognize electron transfer reactions and identify the substance(s) losing and gaining electrons in an electron transfer reaction.	C.4.6   Determine oxidation states and identify the substances gaining and losing electrons in redox reactions.
C.1.23       Write a rate law for a chemical reaction using experimental data.	Not expected
C.1.24       Recognize and describe nuclear changes.	C.2.7   Compare and contrast nuclear reactions with chemical reactions. C.2.8. Describe how fusion and fission processes transform elements present before the reaction into elements present after the reaction.
C.1.25       Recognize the importance of chemical processes in industrial and laboratory settings, e.g., electroplating, electrolysis, the operation of voltaic cells, and such important applications as the refining of aluminum.	Not expected
C.1.26       Describe physical changes and properties of matter through sketches and descriptions of the involved materials.	C.1.4    Describe physical and chemical changes at the molecular level. C.1.5   Describe the characteristics of solids, liquids and gases and changes in state at the molecular level.
C.1.27       Describe chemical changes and reactions using sketches and descriptions of the reactants and products.	C.1.4   Describe physical and chemical changes at the molecular level.
C.1.28       Explain that chemical bonds between atoms in molecules, such as H2, CH4, NH3, C2H4, N2, Cl2, and many large biological molecules are covalent.	C.3.5   Compare and contrast ionic, covalent network, metallic and polar and non-polar molecular crystals with respect to constituent particles, strength of bonds, melting and boiling points and conductivity; provide examples of each type. C.9.2    Illustrate the variety of molecular types formed by the covalent bonding of carbon atoms and describe the typical properties of these molecular types.
C.1.29       Describe dynamic equilibrium.	C.7.6    Write equilibrium expressions for reversible reactions.
C.1.30       Perform calculations that demonstrate an understanding of the gas laws. Apply the gas laws to relations between pressure, temperature, and volume of any amount of an ideal gas or any mixture of ideal gases.	C.5.2    Using the ideal gas equation of state PV = nRT, calculate the change in one variable when another variable is changed and the others are held constant.
C.1.31       Use kinetic molecular theory to explain changes in gas volumes, pressure, and temperature (Solve problems using pV=nRT).	C.5.1   Use kinetic molecular theory to explain changes in gas volumes, pressure, moles and temperature.
C.1.32       Describe the possible subatomic particles within an atom or ion.	C.2.2   Describe how the subatomic particles (i.e., protons, neutrons and electrons) contribute to the structure of an atom and recognize that the particles within the nucleus are held together against the electrical repulsion of the protons.
C.1.33       Use an element’s location in the Periodic Table to determine its number of valence electrons, and predict what stable ion or ions an element is likely to form in reacting with other specified elements.	C.2.6   Use the periodic table and electron configuration to determine an element's number of valence electrons and its chemical and physical properties. C.4.6   Determine oxidation states and identify the substances gaining and losing electrons in redox reactions.
C.1.34       Use the Periodic Table to compare attractions that atoms have for their electrons and explain periodic properties, such as atomic size, based on these attractions.	C.2.2   Describe how the subatomic particles (i.e., protons, neutrons and electrons) contribute to the structure of an atom and recognize that the particles within the nucleus are held together against the electrical repulsion of the protons.
C.1.35       Infer and explain physical properties of substances, such as melting points, boiling points, and solubility, based on the strength of molecular attractions.	C.3.5   Compare and contrast ionic, covalent network, metallic and polar and non-polar molecular crystals with respect to constituent particles, strength of bonds, melting and boiling points and conductivity; provide examples of each type.
C.1.36       Describe the nature of ionic, covalent, and hydrogen bonds and give examples of how they contribute to the formation of various types of compounds.	C.3.1   Describe, compare and contrast the characteristics of the interactions between atoms in ionic and covalent compounds.
C.1.37       Describe that spectral lines are the result of transitions of electrons between energy levels and that these lines correspond to photons with a frequency related to the energy spacing between levels by using Planck’s relationship (E=hv).	Not expected
C.1.38       Distinguish between the concepts of temperature and heat.	C.6.2   Distinguish between the concepts of temperature and heat flow in macroscopic and microscopic terms.
C.1.39       Solve problems involving heat flow and temperature changes, using known values of specific heat and latent heat of phase change.	C.6.4   Solve problems involving heat flow and temperature changes by using known values of specific heat, phase change constants (i.e., latent heat values) or both.
C.1.40       Classify chemical reactions and/or phase changes as exothermic or endothermic.	C.6.3   Classify chemical reactions and phase changes as exothermic or endothermic.
C.1.41       Describe the role of light, heat, and electrical energies in physical, chemical, and nuclear changes.	Not expected
C.1.42       Describe that the energy release per gram of material is much larger in nuclear fusion or fission reactions than in chemical reactions. The change in mass (calculated by E=mc2) is small but significant in nuclear reactions.	C.2.7   Compare and contrast nuclear reactions with chemical reactions.
C.1.43       Calculate the amount of radioactive substance remaining after an integral number of half-lives have passed.	C.2.9    Understand that the radioactive decay process is random for any given atom but that this property leads to a predictable and measurable exponential decay of a sample of radioactive material. Know how to calculate the initial amount, the fraction remaining or the half-life of a radioactive isotope when given two of the other three variables.
C.1.44       Convert between formulas and names of common organic compounds.	C.9.1   Use structural formulas to illustrate carbon atoms’ ability to bond covalently to one another to form many different substances.
C.1.45       Recognize common functional groups and polymers when given chemical formulas and names.	C.9.2    Illustrate the variety of molecular types formed by the covalent bonding of carbon atoms and describe the typical properties of these molecular types.
C.2.3         Explain that John Dalton’s modernization of the ancient Greek ideas of element, atom, compound, and molecule strengthened the new chemistry by providing physical explanations for reactions that could be expressed in quantitative terms.	C.2.1   Describe how models of atomic structure changed over time based on available experimental evidence and understand the current model of atomic structure.
C.2.1         Explain that Antoine Lavoisier invented a whole new field of science based on a theory of materials, physical laws, and quantitative methods, with the conservation of matter at its core. Recognize that he persuaded a generation of scientists that his approach accounted for the experimental results better than other chemical systems. C.2.2         Describe how Lavoisier’s system for naming substances and describing their reactions contributed to the rapid growth of chemistry by enabling scientists everywhere to share their findings about chemical reactions with one another without ambiguity. C.2.4         Explain how Frederich Wohler’s synthesis of the simple organic compound urea from inorganic substances made it clear that living organisms carry out chemical processes not fundamentally different from inorganic chemical processes. Describe how this discovery led to the development of the huge field of organic chemistry, the industries based on it, and eventually to the field of biochemistry. C.2.5         Explain how Arrhenius’ discovery of the nature of ionic solutions contributed to the understanding of a broad class of chemical reactions. C.2.6         Explain that the application of the laws of quantum mechanics to chemistry by Linus Pauling and others made possible an understanding of chemical reactions on the atomic level. C.2.7         Describe how the discovery of the structure of DNA by James D. Watson and Francis Crick made it possible to interpret the genetic code on the basis of a sequence of “letters.”	Not expected
	New Standards: C.2.3   Determine the number of protons, neutrons, and electrons in isotopes and in those isotopes that comprise a specific element.  Relate these numbers to atomic number and mass number. C.2.4   Calculate the average atomic mass of an element from isotopic abundance data. C.2.5   Write the electron configuration of an element and relate this to its position on the periodic table. C.3.2   Compare and contrast how ionic and covalent compounds form. C.3.3   Draw structural formulas for and name simple molecules.  C.6.1   Explain that atoms and molecules are in constant motion and that this motion increases as thermal energy increases. C.7.1   Describe the composition and properties of types of solutions. C.7.2   Explain how temperature, pressure and polarity of the solvent affect the solubility of a solute. C.8.2   Describe the characteristic properties of acids and bases. C.8.3   Compare and contrast the dissociation and strength of acids and bases in solutions. Process Standards