• Questions at the end of standards-based textbook readings and/or activities cited in the right-hand column after each standard/benchmark can be considered as potential standards based assessment questions for quarter "mid-terms" or semester "end-of-term" finals.·
• YELLOW is used to draw attention to core instructional vocabulary.
• BLUE is used to draw attention to instructional "experiences" that students should have.
• GREEN is used to draw attention to expected student opportunities (some requiring the application student initiated metacognitive skills) based on state framework suggestions.
• RED is used to draw attention to issues that might affect the scope and sequence of how the standard based material is presented.
• PINK is used to draw attentions to items that can be used for "cross-curricular" integration of "Language Arts" standard-based items.
• GRAY is used to draw attentions to items that can be used for "cross-curricular" integration of "Mathematics" standard-based items.
• Notations like "Q 2" beside a standard or benchmark mean that the standard or benchmark in question will be covered during the 2nd Quarter. L means late in the quarter and E means early in the quarter.
• The "Content Standard Summary " and annotations after each standard and benchmark are from the California Science Framework (scroll through this document to the Chemistry Standards).
• A * benchmark means it is not considered an "Essential Standard" (Standards considered "essential" are those that are included in the state CST blueprint for a given subject area test).
• "(13% - 8 items) " means that 13% or 8 questions on the HS Chemistry CSS CRT have been written using framework descriptions for this standard and its benchmarks (Chem. CST Blueprint, 2002).
• NE - Considered a Non-Essential Standard (Standards considered "essential" are those that are included in the state CST blueprint for a given subject area test)
• PE means Pupil's Edition and TE means Teacher’s Edition.
• Resources used:

• Course: Introduction to Chemistry
• Text: Modern Chemistry, Holt Rinehart, Winston, 1999
• Criteria: California Science Content Standards for High School Chemistry

1st Content Standard

Starting with atomic and molecular structure in the study of high school chemistry is important because this topic is a foundation of the discipline. However, because structural concepts are highly theoretical and deal with the quantum realm, they can be hard to relate to real-world experience. Ideally, from grades three through eight, students have been gradually introduced to the atomic theory; and by the end of the eighth grade, they should have covered the major concepts in the structure of atoms and molecules. By the time students reach high school, they should be familiar with basic aspects of this theory.

The study of structure can begin with the simplest element, hydrogen. Students can progress from a simple model of the atom (see standards 1.a, 1.d, 1.e, 1.h*, and 11.a in this section) to the historic Bohr model (see Standard 1.i* in this section) and, finally, to a quantum mechanical model (see standards 1.g* and 1.j* in this section), which is the picture of the atom that students should ultimately develop. Students should learn that the quantum mechanical model takes into account the particle and wave properties of the electron and uses mathematical equations to solve for electron energies and regions of electron density.

Students should understand that the energy carried by electrons either within an atom or as electricity can be transformed into light energy. Those who have completed high school physics will be familiar with the properties of electromagnetic waves. In chemistry students learn to apply the equations E = hv and c = λv. Students without significant training in physics will need to understand electromagnetic radiation as energy, frequency, and wavelength. The necessary mathematical background for the study of chemistry includes algebraic isolation of variables, use of conversion factors, and manipulation of exponents, all of which are covered in the mathematics standards for the middle grades.

Chapters 1, 2, 4, 5

Atomic and Molecular Structure (10% of CST: 6 items)- Q1

1. The periodic table displays the elements in increasing atomic number and shows how periodicity of the physical and chemical properties of the elements relates to atomic structure. As a basis for understanding this concept:

- Q1

a. Students know how to relate the position of an element in the periodic table to its atomic number and atomic mass

PE - 1 PE -23,126, 140-142, 156-159

TE - 124 , 140, 141, 142

An atom consists of a nucleus made of protons and neutrons that is orbited by electrons. The number of protons, not electrons or neutrons, determines the unique properties of an element. This number of protons is called the element’s atomic number. Elements are arranged on the periodic table in order of increasing atomic number. Historically, elements were ordered by atomic mass, but now scientists know that this order would lead to misplaced elements (e.g., tellurium and iodine) because differences in the number of neutrons for isotopes of the same element affect the atomic mass but do not change the identity of the element.

- Q1

b. Students know how to use the periodic table to identify metals, semimetals, nonmetals, and halogens.

PE - 21, 125-126 , 129-134 , 136-138, 156-159

TE - 21, 124, 128, 129, 130, 136, 138, 139

Most periodic tables have a heavy stepped line running from boron to astatine. Elements to the immediate right and left of this line, excluding the metal aluminum, are semimetals and have properties that are intermediate between metals and nonmetals. Elements further to the left are metals. Those further to the right are nonmetals. Halogens, which are a well-known family of nonmetals, are found in Group 17 (formerly referred to as Group VIIA). A group, also sometimes called a “family,” is found in a vertical column in the periodic table.

- Q1

c. Students know how to use the periodic table to identify alkali metals, alkaline earth metals and transition metals, trends in ionization energy, electronegativity, and the relative sizes of ions and atoms.

PE - 128-132, 134-136 , 140-146,151-154, 728, 734, 740

TE - 128, 129, 142, 146, 147, 152, 153

A few other groups are given family names. These include the alkali metals (Group 1), such as sodium and potassium, which are soft and white and extremely reactive chemically. Alkaline earth metals (Group 2), such as magnesium and calcium, are found in the second column of the periodic table. The transition metals (Groups 3 through 12) are represented by some of the most common metals, such as iron, copper, gold, mercury, silver, and zinc. All these elements have electrons in their outer d orbitals.

Electronegativity is a measure of the ability of an atom of an element to attract electrons toward itself in a chemical bond. The values of electronegativity calculated for various elements range from one or less for the alkali metals to three and one half for oxygen to about four for fluorine. Ionization energy is the energy it takes to remove an electron from an atom. An element often has multiple ionization energies, which correspond to the energy needed to remove first, second, third, and so forth electrons from the atom. Generally in the periodic table, ionization energy and electronegativity increase from left to right because of increasing numbers of protons and decrease from top to bottom owing to an increasing distance between electrons and the nucleus. Atomic and ionic sizes generally decrease from left to right and increase from top to bottom for the same reasons. Exceptions to these general trends in properties occur because of filled and half-filled subshells of electrons.

- Q1

d. Students know how to use the periodic table to determine the number of electrons available for bonding.

PE - 150-151, 169-170, 172, 174

TE - 150, 151

Only electrons in the outermost energy levels of the atom are available for bonding; this outermost bundle of energy levels is often referred to as the valence shell or valence shell of orbitals. All the elements in a group have the same number of electrons in their outermost energy level. Therefore, alkali metals (Group 1) have one electron available for bonding, alkaline earth metals (Group 2) have two, and elements in Group 13 (once called Group III) have three. Unfilled energy levels are also available for bonding. For example, Group 16, the chalcogens, has room for two more electrons; and Group 17, the halogens, has room for one more electron to fill its outermost energy level.

To find the number of electrons available for bonding or the number of unfilled electron positions for a given element, students can examine the combining ratios of the element’s compounds. For instance, one atom of an element from Group 2 will most often combine with two atoms of an element from Group 17 (e.g., MgCl2) because Group 2 elements have two electrons available for bonding, and Group 17 elements have only one electron position open in the outermost energy level. (Note that some periodic tables indicate an element’s electron configuration or preferred oxidation states. This information is useful in determining how many electrons are involved in bonding.)

- Q 1

e. Students know the nucleus of the atom is much smaller than the atom yet contains most of its mass.

PE - 70, 72-74

TE - 72

The volume of the hydrogen nucleus is about one trillion times less than the volume of the hydrogen atom, yet the nucleus contains almost all the mass in the form of one proton. The diameter of an atom of any one of the elements is about 10,000 to 100,000 times greater than the diameter of the nucleus. The mass of the atom is densely packed in the nucleus.

The electrons occupy a large region of space centered around a tiny nucleus, and so it is this region that defines the volume of the atom. If the nucleus (proton) of a hydrogen atom were as large as the width of a human thumb, the electron would be on the average about one kilometer away in a great expanse of empty space. The electron is almost 2,000 times lighter than the proton; therefore, the large region of space occupied by the electron contains less than 0.1 percent of the mass of the atom.

NE- Q

f.* Students know how to use the periodic table to identify the lanthanide, actinide, and transactinide elements and know that the transuranium elements were synthesized and identified in laboratory experiments through the use of nuclear accelerators.

PE - 21, 126, 130-131, 138-139, 712, 720-721

TE - 21, 130

The lanthanide series, or rare earths, and the actinide series, all of which are radioactive, are separated for reasons of practical display from the main body of the periodic table. If these two series were inserted into the main body, the table would be wider by 14 elements and less manageable for viewing. Within each of these series, properties are similar because the configurations of outer electrons are similar. As a general rule elements in both series appear to have three electrons available for bonding. They combine with halogens to form compounds with the general formula MX3, such as LaFz3. The transactinide elements begin with rutherfordium, element 104.

All the elements with atomic numbers greater than 92 were first synthesized and identified in experiments. These experiments required the invention and use of accelerators, which are electromagnetic devices designed to create new elements by accelerating and colliding the positively charged nuclei of atoms. Ernest O. Lawrence, at the University of California, Berkeley, invented one of the most useful nuclear accelerators, the cyclotron. Many transuranium elements, such as 97-berkelium, 98-californium, 103-lawrencium, and 106-seaborgium, were first created and identified at the adjacent Lawrence Berkeley National Laboratory. Today a few transuranium elements are produced in nuclear reactors. For example, hundreds of metric tons of plutonium have been produced in commercial nuclear reactors.

NE- Q

g.* Students know how to relate the position of an element in the periodic table to its quantum electron configuration and to its reactivity with other elements in the table.

PE - 101-107, 110-116, 121, 128-134, 136-139

TE - 101, 103, 105, 106, 107, 113,114, 116, 128, 129

Each element has a unique electron configuration (also known as quantum electron configuration) that determines the properties of the element. Quantum mechanical calculations predict these electron energy states, which provide the theoretical justification for the organization of the periodic table, previously organized on the basis of chemical properties.

Students can learn the principal quantum numbers which are 1, 2, 3, 4, 5, 6, and 7.and the corresponding periods, or horizontal rows, on the periodic table. They can learn the angular momentum quantum numbers that give rise to s, p, d, and f subshells of orbitals and the rules for the sequence of orbital filling. The electrons in the highest energy orbitals with the same principal quantum number are the valence electrons. For example, for all elements in Group 1, the valence electron configuration is ns1, where n is the principal quantum number. Analogously, all elements in Group 16 have valence electron configurations of ns2np4. Particular configurations of valence electrons are associated with regular patterns in chemical reactivity. Generally, those elements with one electron in excess or one electron short of a full octet in the highest occupied energy level, the alkali metals and halogens, respectively, are the most reactive.

NE- Q 1

h.* Students know the experimental basis for Thomson’s discovery of the electron, Rutherford’s nuclear atom, Millikan’s oil drop experiment, and Einstein’s explanation of the photoelectric effect.

PE - 70-73, 93-94

TE - 72, 73

In 1887 J. J. Thomson performed experiments from which he concluded that cathode rays are streams of negative, identical particles, which he named electrons. In 1913 E. R. Rutherford headed a group that shot a beam of helium nuclei, or alpha particles, through a very thin piece of gold foil; the unexpectedly large deflections of some helium nuclei led to the hypothesis that the charged mass of each gold atom is concentrated in a small central nucleus. Robert Millikan confirmed Thomson’s conclusion that electrons are identical particles when he balanced tiny, electrically charged oil droplets between electric and gravitational fields and so discovered that the droplets always contained charge equal to an integral multiple of a single unit.

Albert Einstein found he could explain the photoelectric effect, in which light ejects electrons from metal surfaces, by proposing that light consists of discrete bundles of energy, or photons, each photon with an energy directly proportional to the frequency of the light, and by proposing that each photon could eject one and only one electron. Photons with sufficient energy will eject an electron whose kinetic energy is equal to the photon energy minus the energy required to free the electron from the metal. If the frequency of the light and therefore the energy of each proton is too low to free an electron, then merely increasing the light’s intensity (that is, merely producing more photons) does not cause electrons to eject.

NE- Q 1

i.* Students know the experimental basis for the development of the quantum theory of atomic structure and the historical importance of the Bohr model of the atom

PE - 94-104

TE - 96, 98, 100, 101

Niels Bohr combined the concepts of Rutherford’s nuclear atom and Einstein’s photons with several other ideas to develop a model that successfully explains the observed spectrum, or wavelengths, of electromagnetic radiation that is emitted when a hydrogen atom falls from a high energy state to a low energy state. In classical physics all accelerating charges emit energy. If electrons in an atom behaved in this way, light of ever-decreasing frequencies would be emitted from atoms. Bohr explained why this phenomenon does not occur when he suggested that electrons in atoms gain or lose energy only by making transitions from discrete energy levels. This idea was a key to the development of nonclassical descriptions of atoms. Louis de Broglie advanced the understanding of the nature of matter by proposing that particles have wave properties. On the basis of these ideas, Erwin Schrodinger and others developed quantum mechanics, a theory that describes and predicts atomic and nuclear phenomena.

NE - Q

j.* Students know 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).

PE - 92-95

TE - 92, 94, 95

The Bohr model gives a simple explanation of the spectrum of the hydrogen atom. An electron that loses energy in going from a higher energy level to a lower one emits a photon of light, with energy equal to the difference between the two energy levels. Transitions of electrons from higher energy states to lower energy states yield emission, or bright-line, spectra. Absorption spectra occur when electrons jump to higher energy levels as a result of absorbing photons of light. When atoms or molecules absorb or emit light, the absolute value of the energy change is equal to hc/λ, where h is a number called Planck’s constant, c is the speed of light, and λ is the wavelength of light emitted, yielding Planck’s relationship E =hv.

The transition from the Bohr model to the quantum mechanical model of the atom requires students to be aware of the probabilistic nature of the distribution of electrons around the atom. A “dart board” made of concentric rings can serve as a two-dimensional model of the three-dimensional atom. A graph as a function of radius of the number of random hits by a dart can be compared with a similar graph as a function of radius of the probability density of the electron in a hydrogen atom.

Chemical Bonds (11.7% of CST: 7 items)- Q2

2. Biological, chemical, and physical properties of matter result from the ability of atoms to form bonds from electrostatic forces between electrons and protons and between atoms and molecules. As a basis for understanding this concept:

- Q2

a. Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds.

PE - 161-165, 167-182, 196-199

TE - 161, 162, 164, 165, 167, 170, 176, 178, 181

In the localized electron model, a covalent bond appears as a shared pair of electrons contained in a region of overlap between two atomic orbitals. Atoms (usually nonmetals) of similar electronegativities can form covalent bonds to become molecules. In a covalent bond, therefore, bonding electron pairs are localized in the region between the bonded atoms. In metals valence electrons are not localized to individual atoms but are free to move to temporarily occupy vacant orbitals on adjacent metal atoms. For this reason metals conduct electricity well.

When an electron from an atom with low electronegativity (e.g., a metal) is removed by another atom with high electronegativity (e.g., a nonmetal), the two atoms become oppositely charged ions that attract each other, resulting in an ionic bond. Chemical bonds between atoms can be almost entirely covalent, almost entirely ionic, or in between these two extremes. The triple bond in nitrogen molecules (N2) is nearly 100 percent covalent. A salt such as sodium chloride (NaCl) has bonds that are nearly completely ionic. However, the electrons in gaseous hydrogen chloride are shared somewhat unevenly between the two atoms. This kind of bond is called polar covalent. (Note that elements in groups 1, 2, 16, and 17 in the periodic table usually gain or lose electrons through the formation of either ionic or covalent bonds, resulting in eight outer shell electrons. This behavior is sometimes described as “the octet rule.”)

- Q2

b. Students know chemical bonds between atoms in molecules such as H 2 , CH 4 , NH 3 , H 2 CCH 2 , N 2 , Cl 2 , and many large biological molecules are covalent

PE - 164-165, 167-168, 172, 761-766

TE - 164, 167

Organic and biological molecules consist primarily of carbon, oxygen, hydrogen, and nitrogen. These elements share valence electrons to form bonds so that the outer electron energy levels of each atom are filled and have electron configurations like those of the nearest noble gas element. (Noble gases, or inert gases, are in the last column on the right of the periodic table.) For example, nitrogen has one lone pair and three unpaired electrons and therefore can form covalent bonds with three hydrogen atoms to make four electron pairs around the nitrogen. Carbon has four unpaired electrons and combines with hydrogen, nitrogen, and oxygen to form covalent bonds sharing electron pairs.

The great variety of combinations of carbon, nitrogen, oxygen, and hydrogen make it possible, through covalent bond formation, to have many compounds from just these few elements. Teachers can use ball and stick or gumdrop and toothpick models to explore possible bonding combinations.

- Q2

c. Students know salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by electrostatic attraction.

PE - 176-178

TE - 176, 177, 178

The energy that holds ionic compounds together, called lattice energy, is caused by the electrostatic attraction of cations, which are positive ions, with anions, which are negative ions. To minimize their energy state, the ions form repeating patterns that reduce the distance between positive and negative ions and maximize the distance between ions of like charges.

- Q2

d. Students know the atoms and molecules in liquids move in a random pattern relative to one another because the intermolecular forces are too weak to hold the atoms or molecules in a solid form.

PE - 189-193

TE - 190, 192

In any substance at any temperature, the forces holding the material together are opposed by the internal energy of particle motion, which tends to break the substance apart. In a solid, internal agitation is insufficient to overcome intermolecular or interatomic forces. When enough energy is added to the solid, the kinetic energy of the atoms and molecules increases sufficiently to overcome the attractive forces between the particles, and they break free of their fixed lattice positions. This change, called melting, forms a liquid, which is disordered and non-rigid. The particles in the liquid are free to move about randomly although they remain in contact with each other.

- Q2

e. Students know how to draw Lewis dot structures.

PE - 170-175, 180, 184-185, 187, 196-197

TE - 170, 171, 174

A Lewis dot structure shows how valence electrons and covalent bonds are arranged between atoms in a molecule. Teachers should follow the rules for drawing Lewis dot diagrams provided in a chemistry textbook. Students should be able to use the periodic table to determine the number of valence electrons for each element in Groups 1 through 3 and 13 through 18. Carbon, for example, would have four valence electrons. Lewis dot diagrams represent each electron as a dot or an x placed around the symbol for carbon, which is C. A covalent bond is shown as a pair of dots, or x’s, representing a pair of electrons. For example, a Lewis dot diagram for methane, which is CH4, would appear as shown in Figure 3. Lewis dot diagrams provide a method for predicting correct combining ratios between atoms and for determining aspects of chemical bonds, such as whether they are covalent or consist of single, double, or triple bonds.

NE- Q2

f.* Students know how to predict the shape of simple molecules and their polarity from Lewis dot structures.

PE - 183-189

TE - 184, 185, 187

Using information obtained from Lewis dot structures of covalently bonded molecules, students can predict the overall geometry of those molecules. This model assumes that valence electron pairs repel each other and that atoms covalently bonded around a central atom position themselves as far apart as possible while maintaining the covalent bond. The model also assumes that double or triple bonds define a single electronic region.

To predict shapes, students start with a correct Lewis dot structure. From the number of electron pairs or regions, both bonded and nonbonded, students can determine the molecular geometry of the molecule because Lewis dot structures, although drawn in two dimensions, represent the three-dimensional symmetry of the molecule. A symmetrical distribution of the electron clouds around a central atom leads to a nonpolar molecule in which charge is evenly distributed. Students should be able to predict that a molecule such as methane, with one carbon and four hydrogen atoms, forms a tetrahedral shape and because of its symmetry is a nonpolar molecule.

- Q2

g.* Students know how electronegativity and ionization energy relate to bond formation.

PE - 145-146, 151-154, 157, 161-163

TE - 145, 146, 147, 151, 152, 153

During bond formation atoms with large electronegativity values, such as fluorine and oxygen, attract electrons away from lower electronegativity atoms, such as the alkali metals. The difference in electronegativity between the two bonding atoms gives information on how evenly an electron pair is shared. A large difference in electronegativity leads to an ionic bond, with essentially no sharing of electrons.

This phenomenon usually occurs between metal and nonmetal atoms. A small difference in electronegativity leads to a covalent bond with more equal sharing of electrons. This result typically occurs between two nonmetal atoms. Electronegativity is related to ionization energy, the energy needed to remove an electron from an isolated gaseous atom, leaving a positively charged ion. High ionization energies usually correlate with large electronegativities.

- Q2

h.* Students know how to identify solids and liquids held together by Van der Waals forces or hydrogen bonding and relate these forces to volatility and boiling/melting point temperatures.

PE - 165, 167-169, 182, 192

TE - 165, 167, 168, 192

Liquids and solids that are held together not by covalent or ionic bonds but only by weaker intermolecular forces tend to have low to moderate melting and boiling points and to be from somewhat to very volatile. The volatility of a substance means how readily it evaporates at ordinary temperatures and pressures.

Two kinds of intermolecular forces are hydrogen bonding and van der Waals attractions (often referred to as London dispersion forces). Hydrogen bonding is essential to life and gives water many of its unusual properties. A hydrogen bond occurs when a hydrogen atom on one molecule, bonded directly to a highly electronegative atom (fluorine, oxygen, or nitrogen), is weakly attracted to the electronegative atom on a neighboring molecule. In the important case of water, this attraction exists between the hydrogen on one water molecule and the oxygen on a neighboring water molecule. This attraction happens because of water’s polar nature and bent shape

Van der Waals forces exist between all molecules, polar or nonpolar. Even when the molecule is nonpolar, the electrons move around and may sometimes find themselves temporarily closer to one nucleus than to the other. The atom with the greater share of the electron density becomes, for an instant, slightly negatively charged, and the other atom becomes a little bit positive. If the same thing happens in a nearby molecule, the positive and negative centers on the two molecules temporarily attract each other. Bigger molecules have more electrons, and their van der Waals forces are greater. This phenomenon leads to molecules with lower volatility and higher melting and boiling temperatures

3rd Content Standard

Standard Set 3 demands more facility with mathematics than do the previous two chemistry standard sets. For this reason students need prerequisite mathematical skills in two broad categories. First, they must be able to manipulate very large and very small numbers by using exponents as expressed in scientific notation and they should learn the rules of significant digits when reporting measurements and the results of calculations. Second, students must be able to manipulate simple equations in symbolic form, such as the isolation of variables, and to write equations numerically with correct units. Handling units successfully is necessary for problems involving mole-to-mass and mass-to-mole conversions. The ability to make other types of unit conversion and to square and cube linear measurements will also be required. An understanding of ratios will help students to see the logic behind problems related to Standard Set 3.

Simple metric conversions are relatively easy to grasp, and they contain the basic elements of stoichiometric calculations. For example, to convert 510 nanometers to meters, students write the unknown unit, which in this case is meters; set it equal to the given unit, which in this case is nanometers; and then multiply by the correct conversion factor, making sure that the desired unit is in the numerator and the unit to be cancelled is in the denominator. The following equation demonstrates this procedure:

510 nm = 510 nm × (10−"9 m/1 nm) = 510 × 10−"9 m = 5.1 × 10−"7 m

This technique converts any units however complicated. For example, centimeters can be converted to nanometers by multiplying by two factors: one that converts from centimeters to meters and another that converts from meters to nanometers. Converting to an area requires the square of a conversion factor; converting to a volume requires the cube. Once students have tackled simple problems, they and move on to more complex stoichiometric relations. They will also need to learn to balance equations easily.

Chapters 1, 2, 3, 7, 8, 9,11, 12

Conservation of Matter and Stoichiometry (16.7% of CST: 10 items)- Q1, 2, 3

3. The conservation of atoms in chemical reactions leads to the principle of conservation of matter and the ability to calculate the mass of products and reactants. As a basis for understanding this concept:

- Q2

a. Students know how to describe chemical reactions by writing balanced equations.

PE - 243-254, 264, 270-273

TE - 243, 245, 247, 248, 251, 252, 253, 254

Reactions are described by balanced equations because all the atoms of the reactants must be accounted for in the reaction products. An equation with all correct chemical formulas can be balanced by a number of methods, the simplest being by inspection. Given an unbalanced equation, students can do an inventory to determine how many of each atom are on each side of the equation. If the result is not equal for all atoms, coefficients (not subscripts) are changed until balance is achieved. Sometimes, reactions refer to substances with written names rather than to chemical symbols. Students should learn the rules of chemical nomenclature. This knowledge can be acquired in stages as new categories of functional groups are introduced.

- Q1

b. Students know the quantity one mole is set by defining one mole of carbon 12 atoms to have a mass of exactly 12 grams.

PE - 81

TE - 81

The mole concept is often difficult for students to understand at first, but they can be taught that the concept is convenient in chemistry just as a dozen is a convenient concept, or measurement unit, in the grocery store. The mole is a number. Specifically, a mole is defined as the number of atoms in 12 grams of carbon-12. When atomic masses were assigned to elements, the mass of 12 grams of carbon-12 was selected as a standard reference to which the masses of all other elements are compared. The number of atoms in 12 grams of carbon-12 is defined as one mole, or conversely, if one mole of 12C atoms were weighed, it would weigh exactly 12 grams. (Note that carbon, as found in nature, is a mixture of isotopes, including atoms of carbon-12, carbon-13, and trace amounts of carbon-14.) The definition of the mole refers to pure carbon-12.

The atomic mass of an element is the weighted average of the mass of one mole of its atoms based on the abundance of all its naturally occurring isotopes. The atomic mass of carbon is 12.011 grams. If naturally occurring carbon is combined with oxygen to form carbon dioxide, the mass of one mole of naturally occurring oxygen can be determined from the combining mass ratios of the two elements. For example, the weight, or atomic mass, of one mole of oxygen containing mostly oxygen-16 and a small amount of oxygen-18 is 15.999 grams.

- Q 1

c. Students know one mole equals 6.02 x 10 ²³ particles (atoms or molecules).

PE - 80-81, 89

TE - 81

A mole is a very large number. Standard 3.b describes the mole as the number of atoms in 12 grams of 12C. The number of atoms in a mole has been found experimentally to be about 6.02 x 10 ²³. This number, called Avogadro’s number, is known to a high degree of accuracy.

- Q2

d. Students know how to determine the molar mass of a molecule from its chemical formula and a table of atomic masses and how to convert the mass of a molecular substance to moles, number of particles, or volume of gas at standard temperature and pressure.

PE - 82-85, 88, 222-225, 238, 277, 281-286, 296, 816-818

TE - 82, 83, 84, 222, 223, 224, 225, 281, 283, 284, 285

The molar mass of a compound, which is also called either the molecular mass or molecular weight, is the sum of the atomic masses of the constituent atoms of each element in the molecule. Molar mass is expressed in units of grams per mole. The periodic table is a useful reference for finding the atomic masses of each element. For example, one mole of carbon dioxide molecules contains one mole of carbon atoms weighing 12.011 grams and two moles of oxygen atoms weighing 2 × 15.999 grams for a total molecular mass of 44.009 grams per mole of carbon dioxide molecules.

The mass of a sample of a compound can be converted to moles by dividing its mass by the molar mass of the compound. This process is similar to the unit conversion discussed in the introduction to Standard Set 3. The number of particles in the sample is determined by multiplying the number of moles by Avogadro’s number. The volume of an ideal or a nearly ideal gas at a fixed temperature and pressure is proportional to the number of moles. Students should be able to calculate the number of moles of a gas from its volume by using the relationship that at standard temperature and pressure (0°C and 1 atmosphere), one mole of gas occupies a volume of 22.4 liters.

- Q2, 3

e. Students know how to 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.

PE - 275, 286-287, 296

TE - 275, 286, 287

Atoms are neither created nor destroyed in a chemical reaction. When the chemical reaction is written as a balanced expression, it is possible to calculate the mass of any one of the products or of any one of the reactants if the mass of just one reactant or product is known. Students can be taught how to use balanced chemical equations to predict the mass of any product or reactant. Teachers should emphasize that the coefficients in the balanced chemical equation are mole quantities, not masses. Here is an example: How many grams of water will be obtained by combining 5.0 grams of hydrogen gas with an excess of oxygen gas, according to the following balanced equation?

This calculation is often set up algebraically, for example, as

and can be easily completed by direct calculation and unit cancellation (dimensional analysis). Students should learn to recognize that the coefficients in the balanced equations refer to moles rather than to mass.

NE- Q3

f.* Students know how to calculate percent yield in a chemical reaction.

PE - 288-291, 293-294, 297-299, 819-821

TE - 288, 289, 290, 293, 294, 820

Students can use a balanced equation for a chemical reaction to calculate from the given masses of reactants the masses of the resulting products. The masses so calculated represent the theoretical 100 percent conversion of reactants to products.

When one or more of the products are weighed, the masses are often less than the theoretical 100 percent yield. One explanation is that the reaction may not go to completion and therefore will not convert all the reactants to products. A second possible explanation is that product material is lost in the separation and purification process. A third reason is that alternative reactions are taking place, leading to products different from those predicted. Percent yield is a standard way to compare actual and theoretical yields. It is defined as:

NE- Q4

g.* Students know how to identify reactions that involve oxidation and reduction and how to balance oxidation-reduction reactions.

PE - 591-617, 618-621, 878-883

TE - 594, 595, 596, 597, 598, 600, 601, 607, 610

Oxidation of an element is defined as an increase in oxidation number, or a loss of electrons. Reduction of an element is defined as a decrease in oxidation number, or a gain in electrons. The assignment of oxidation numbers, or oxidation states, is a bookkeeping device. This process may be defined as the charge assigned to an atom, as though all the electrons in each bond were located on the more electronegative atom in the bond. Students should be taught how to assign oxidation numbers to atoms in free elements and in compounds. Lists of rules for this procedure are commonly found in chemistry textbooks.

In many important chemical reactions, elements change their oxidation states. These changes are called redox, or oxidation-reduction reactions. Respiration and photosynthesis are common examples with which students are familiar. Any chemical reaction in which electrons are transferred from one substance to another is an oxidation-reduction reaction. Transfer can be determined by checking the oxidation states of atoms in reactants and products. A single displacement reaction, such as zinc metal in a copper sulfate solution, is a typical redox reaction because the zinc atoms lose two electrons and the copper ions gain two electrons.

Redox reactions may be balanced by ensuring that the number of electrons lost in one part of the reaction equals the number of electrons gained in another part of the reaction. This principle and the conservation of atoms will hold true in any balanced oxidation/reduction reaction. Students should be taught how to use the half-reaction method for balancing redox equations. This method divides the reaction into an oxidation portion and a reduction portion, both of which are balanced separately and then combined into an overall equation with no net change in number of electrons.

Content Standards

Standard Set 4 requires a knowledge of applicable physical concepts and sufficient skills in mathematical problem solving to describe (1) gases at the molecular level; (2) the behavior of gases; and (3) the measurable properties of gases. As background, students should know the physical states of matter and their properties and know how mixtures, especially homogeneous mixtures, differ from pure substances. Students should know that temperature measures how hot a system is, regardless of the system’s size, and that heat flows from a region of higher temperature to one of lower temperature. Familiarity with the Celsius and Kelvin temperature scales is necessary. Knowledge of the motion of particles (and objects) and of kinetic energy is also required. Students should have the mathematical background to recognize and use directly and inversely proportional relationships. The ability to solve algebraic equations with several given quantities and one unknown is essential.

The main assumptions of the kinetic molecular theory should be covered in presenting this material, and a connection should be made to the behavior of ideal gases. Students are often more comfortable with studying the volume and temperature of a gas, but the concept of pressure also needs to be dealt with thoroughly. With knowledge of the kinetic molecular theory, students should see how motions and collisions of particles produce measurable properties, such as pressure.

Chapters 10, 11, 14

Gases and Their Properties (10% of CST: 6 items) - Q2, 3, 4

4. The kinetic molecular theory describes the motion of atoms and molecules and explains the properties of gases. As a basis for understanding this concept:

- Q2, 3

a. Students know the random motion of molecules and their collisions with a surface create the observable pressure on that surface.

PE - 303-304

TE - 304

Fluids consist of molecules that freely move past each other in random directions. Intermolecular forces hold the atoms or molecules in liquids close to each other. Gases consist of tiny particles, either atoms or molecules, spaced far apart from each other and reasonably free to move at high speeds, near the speed of sound. In the study of chemistry, gases and liquids are considered fluids. Pressure is defined as force per unit area. The force in fluids comes from collisions of atoms or molecules with the walls of the container. Air pressure is created by the weight of the gas in the atmosphere striking surfaces. Gravity pulls air molecules toward Earth, the surface that they strike. Water pressure can be understood in the same fashion, but the pressures are much greater because of the greater density of water. Pressure in water increases with depth, and pressure in air decreases with altitude. However, pressure is felt equally in all directions in fluids because of the random motion of the molecules.

- Q2, 3, 4

b. Students know the random motion of molecules explains the diffusion of gases.

PE - 305-306, 351, 353

TE - 305, 351, 353

Another result of the kinetic molecular theory is that gases diffuse into each other to form homogeneous mixtures. An excellent demonstration of diffusion is the white ammonium chloride ring formed by simultaneous diffusion of ammonia vapor and hydrogen chloride gas toward the middle of a glass tube. The white ring forms nearer the region where hydrogen chloride was introduced, illustrating both diffusion and the principle that heavier gases have a slower rate of diffusion.

- Q2, 3

c. Students know how to apply the gas laws to relations between the pressure, temperature, and volume of any amount of an ideal gas or any mixture of ideal gases.

PE - 313-323,328-331

TE - 313, 314, 315, 316, 317, 319, 320, 321, 322, 323

A fixed number of moles n of gas can have different values for pressure P, volume V, and temperature T. Relationships among these properties are defined for an ideal gas and can be used to predict the effects of changing one or more of these properties and solving for unknown quantities. Students should know and be able to use the three gas law relationships summarized in Table 1, “Gas Law Relationships.”

The first expression of the gas law shown in Table 1 is sometimes taught as Boyle’s law and the second as Charles’s law, according to the historical order of their discovery. They are both simpler cases of the more general ideal gas law introduced in Standard 4.h in this section. For a fixed number of moles of gas, a combined gas law has the form PV/T = constant, or P1V1/T1 = P2V2/T2. This law is useful in calculations where P, V, and T are changing. By placing a balloon over the mouth of an Erlenmeyer flask, the teacher can demonstrate that volume divided by temperature equals a constant. When the flask is heated, the balloon inflates; when the flask is cooled, the balloon deflates.

- Q2, 3

d. Students know the values and meanings of standard temperature and pressure (STP).

PE - 312

TE - 312

Standard temperature is 0°C, and standard pressure (STP) is 1 atmosphere (760 mm Hg). These standards are an agreed-on set of conditions for gases against which to consider other temperatures and pressures. When volumes of gases are being compared, the temperature and pressure must be specified. For a fixed mass of gas at a specified temperature and pressure, the volume is also fixed.

- Q2, 3

e. Students know how to convert between the Celsius and Kelvin temperature scales.

PE - 317

TE - 316

Some chemical calculations require an absolute temperature scale, called the Kelvin scale (K), for which the coldest possible temperature is equal to zero. There are no negative temperatures on the Kelvin scale. In theory if a sample of any material is cooled as much as possible, the lowest temperature that can be reached is 0 K, experimentally determined as equivalent to -273.15°C. The Kelvin scale starts with absolute zero (0 K) because of this theoretical lowest temperature limit. A Kelvin temperature is always 273.15 degrees greater than an equivalent Celsius temperature, but a Kelvin temperature is specified without the degree symbol. The magnitude of one unit of change in the K scale is equal to the magnitude of one unit of change on the degree °C scale.

- Q2, 3

f. Students know there is no temperature lower than 0 Kelvin.

PE - 317

TE - 316

The kinetic molecular theory is the basis for understanding heat and temperature. The greater the atomic and molecular motion, the greater the observed temperature of a substance. If all atomic and molecular motion stopped, the temperature

of the material would reach an absolute minimum. This minimum is absolute zero, or -273.15°C. The third law of thermodynamics states that this temperature can never be reached. Experimental efforts to create very low temperatures have resulted in lowering the temperature of objects to within a fraction of a degree of absolute zero.

- Q2, 3

g.* Students know the kinetic theory of gases relates the absolute temperature of a gas to the average kinetic energy of its molecules or atoms.

PE - 303-306

TE - 303, 304, 306

The value of the average kinetic energy for an ideal gas is directly proportional to its Kelvin temperature. Average kinetic energy can be related to changes in pressure and volume as a function of temperature. At 0 K all motion in an ideal monatomic gas ceases, meaning that the average kinetic energy equals zero.

NE- Q2, 3

h.* Students know how to solve problems by using the ideal gas law in the form PV =nRT.

PE - 340-350, 357-361

TE - 340, 343, 344, 345, 346, 347, 348, 349, 350

The relationships among pressure, volume, and temperature for a fixed mass of gas can be expressed as the ideal gas law, PV = nRT, where n represents moles and R represents the universal gas constant, which is 0.0821 liter-atmosphere per mole- Kelvin (this unit can be abbreviated as (L atm)/(mol K) or as L atm mol^-1K^-1).

NE- Q2, 3

i.* Students know how to apply Dalton’s law of partial pressures to describe the composition of gases and Graham’s law to predict diffusion of gases.

PE - 322-325, 329, 334

TE - 323, 324, 325

It is important to distinguish clearly between diffusion and effusion. Diffusion is the process by which separate atoms or molecules intermingle as a result of random motion. Effusion is the process by which gas molecules pass from one container to another at lower pressure through a very small opening.

Graham’s law states that the rates of effusion of two gases at the same temperature and pressure are inversely proportional to the square roots of their molar masses. Graham’s law also approximately applies to rates of diffusion for gases, although diffusion is a more complicated process to describe than effusion. Dalton’s law of partial pressures states that total pressure in a gas-filled container is equal to the sum of the partial pressures of the component gases. This law can be introduced by showing that ideal gases have properties based solely on the number of moles present in the sample, without regard to the chemical identities of the gas particles involved acids.

5th Content Standard

Students who learn the concepts in this standard set will be able to understand and explain aqueous acid–base reactions, properties of acids and bases, and pH as a measure of acidity and basicity. Careful thought should be given to how this standard set best fits with the other chemistry standards. It may be desirable to cover Standard Set 6, “Solutions,” in this section first so that students will know about the aqueous dissolving process and about concentration calculations and units. Familiarity with Standard Set 9, “Chemical Equilibrium,” in this section may help students to conceptualize the strengths of acids and bases.

Students should be able to represent and balance chemical reactions. They should also be able to interpret periodic trends in electronegativity for the upper two rows of the periodic table (see Standard Set 1, “Atomic and Molecular Structure, in this section). Hydrogen with its low electronegativity easily forms the positive hydrogen ion, H+. Students need to know the charge and formula of the hydroxide ion, OH−". With knowledge of polar covalent bonding, students should be able to distinguish between two important types of neutral molecular compounds: those that dissolve in an aqueous solution and remain almost completely as neutral molecules and those that dissolve in an aqueous solution and partially or almost completely produce charged ions (see Standard Sets 1 and 2 in this section).

Students should be able to compare the three descriptions of acids and bases—the Arrhenius, Brønsted-Lowry, and Lewis acid–base definitions—and recognize electron lone pairs on Lewis dot structures of molecules (see Standard Set 2, “Chemical Bonds,” in this section). To calculate pH, students should understand and be able to use base-10 logarithms and antilogarithms and know how to obtain logarithms by using a calculator. Students should become proficient at converting between pH, pOH, [H+], and [OH−"].

Chapter15, 16,