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Chemistry SOL Review PacketModified from a review created by: Mrs. Joyce McAlister and Ms. Shonna Crisden, Menchville High School. http://mville.nn.k12.va.us/science/SHONNA/SOLReview/SOLTitlepage.html

Topic 1: Elements and the Periodic Table

1.1 All matter is made from about 100 different chemical elements. The Periodic Table of the Elements shows all of the known elements, arranged by increasing atomic number. Each element has a symbol. The symbol for many of the elements is one capital letter. In two-letter symbols for elements, the first letter is always an upper case letter, the second one a lower case. The smallest particle of an element is an atom. Some common elements that are gases are composed of molecules containing two atoms of the same element. Example: hydrogen H2(g) and oxygen O2(g).

2.1 Atoms are made of three types of subatomic particles: protons, neutrons and electrons. Each atom has a nucleus in the center, made of protons and neutrons packed tightly together. An electron cloud surrounds the atomic nucleus. The atomic number for an element is the same as the number of protons. All atoms of the same element have the same number of protons. A proton has a positive charge and a relative mass of one. The number of electrons is the same as the number of protons in a neutral atom. An electron has a negative charge and a relative mass of zero. A neutron has no charge and a relative mass of one.

3.1 There are only certain regions in the electron cloud where electrons are likely to be found. These regions are called energy levels. The lowest energy level is closest to the nucleus; the highest energy level is farthest away from the nucleus. Electrons will occupy the lowest available energy level(s) before they fill in higher levels. The outermost electrons in an atom are called valence electrons. The period (row) number on the periodic table corresponds to the highest energy level occupied by the valence electrons in an element. Elements in the same group (column) on the periodic table have the same number of valence electrons. All of the group 1 elements have one valence electron and group two elements have two. Group 13 elements have three valence electrons, group 14 elements have 4, group 15 have 5 and so on through group 18 elements, which have eight valence electrons. An ion is an atom that has a charge because it has gained or lost electrons. Positive ions have lost electrons; negative ions have gained electrons. The amount of charge is equal to the number of electrons lost or gained.

4.1 The principal energy levels (n) around the nucleus of an atom identify the specific regions (distances from the nucleus) where electrons are likely to be found. Principal energy levels are identified by n=1, 2, 3 . . . with n=1 closest to the nucleus. As the value of n increases, so does the distance from the nucleus. Using the periodic table, the period (row) where an element is found indicates the number of occupied energy levels for that element. The energy level of the valence electrons corresponds to the period number (row) where the element is found. Each principal energy level is divided into sublevels (s,p,d and f). In a given energy level, the s sublevel holds up to 2 electrons, and always fills before the p sublevel, which can hold up to 6 electrons. Electron configurations indicate the filling order of all of the electrons in an atom. The coefficients represent the principal energy level, the letters represent the sublevels and the superscripts represent the number of electrons in the sublevel.

5.1 Going down a group on the Periodic Table, each element has one more principal energy level filled with electrons than the element above it, so the outer electrons are farther away from the nucleus. This means the size of the atoms increases going down a group. Therefore the atomic radius increases going down a group. Going from left to right across a period of the Periodic Table the valence electrons are all in the same principal energy level, but the number of protons in the nucleus increases from one element to the next. This means that the nucleus becomes more positively charged and attracts the electrons more strongly. Therefore, the atomic radius decreases going from left to right across a period. Diatomic elements, N2, H2, O2, Cl2, I2, F2, and Br2.

6.1 Electronegativity is the ability of an atom in a bond to attract electrons. The electronegativity of an element can be judged from its position on the periodic table. The

electronegativity increases across a period of the periodic table (because the atomic radius deceases, which means that the valence electrons are held more tightly by the nucleus). The electronegativity decreases going down a group (because the valence electrons are further away from more loosely held the nucleus).

7.1 Ionization energy is the energy needed to remove a valence electron from a atom. Ionization energy increases going from left to right across a period of the periodic table because the atomic radius decreases, which means that the valence electrons are held more tightly by the nucleus. Ionization energy decreases going down a group because the valence electrons are further away from and more loosely held by the nucleus.

8.1 Metals react by losing electrons (oxidation). Group 1 metals (called the alkali metals) are the most reactive metals, because they have only one valence electron to lose. Activity of metals decreases going across each period to groups 2 and 13. Groups 3 through 12 contain the transition elements. They are all metals, but less reactive than those in group 1 and 2. Their oxidation states cannot be easily predicted, and so given in their names. Example: iron (III) means Fe+3; copper (II) means Cu+2 The nonmetals gain electrons (reduction) when they react. Group 17 nonmetals (called the halogens) are the most reactive nonmetals because they have only one electron to gain to get a stable valence shell with eight electrons. Activity decreases in groups 16 and then 15. Elements in group 18 (called the noble gases) have all their principle energy levels filled with electrons and so have little chemical reactivity.

Check Your Understanding

1. An atom with atomic number 48 and mass number 120 contains: a. 48 protons, 48 electrons, 72 neutrons b. 72 protons, 48 electrons, and 48 neutrons

c. 120 protons, 48 electrons, and 72 neutrons d. 72 protons, 72 electrons, and 48 neutrons

2. An element which has a mass number of 23 and has 12 neutrons is the element: a. Lithium b. Potassium c. Magnesium d. Sodium 3. The nucleus of the atom has a. a high density b. a low density

c. a negative charge d. no charge

4. An ion always contains a. unequal number of protons and electrons b. equal number of protons and electrons

c. unequal number of protons and neutrons d. equal number of protons and neutrons

5. The whole number that is closest to the atomic mass of an atom is the a. atomic number b. mass number

c. Avogadro's number d. number of neutrons

6. The ion with the charge of +1 and the same electron configuration as argon is a. potassium b. sodium c. neon d. magnesium 7. The tendency to lose electron_________ as we move across a period on the periodic table a. increases b. remains the same

c. decreases d. no trend exists

8. Atomic radii generally increases in size from __________ in the periodic table a. up a group and left to right across a period b. down a group and left to right across a period c. up a group and right to left across a period d. down a group and right to left across a period

9. The _____ generally have the lowest ionization energy a. noble gases b. metalloids c. nonmetals d. metals 10. Examine the following outer electron configuration and choose the correct location of the element it represents in the periodic table: 5s2 a. row 7 column 4 b. row 4 column 7

c. row 5 column 2 d. row 5 column 9

11. The element from question 10 is in a. period 5, alkali metal b. period 4, halogen

c. period 5, halogen d. period 5, transitional metal

12. Sodium and potassium have similar properties because they have the same: a. atomic radius b. number of valence electrons

c. ionization energy d. electronegativity

13. The likeliest charge an atom with 2 valence electrons would develop a. 2+ b. 6+ c. 2- d. 6- 14. The likeliest charge an atom with 6 valence electrons will develop a. 2+ b. 6+ c. 2- d. 6 15. How many electrons are in the highest occupied energy level of a sodium? a. 1 b. 2 c. 3 d. 4 16. What is the maximum number of electrons that can occupy the 5p sublevel? a. 2 b. 4 c. 6 d. 8 17. How many orbitals in the 4p sublevel? a. 1 b. 2 c. 3 d. 4 18. What is the maximum number of electrons that can occupy the 4th energy level? a. 9 b. 18 c. 20 d. 32 19. Which is the symbol of the element whose electron configuration is 1s2 2s2 2p6 3s2 3p6 4s2 a. Ca b. Li c. Rb d. Ar 20. Elements tend to gain or lose electrons in order to acquire the electron configuration of a: a. halogen b. transition metal

c. noble gas d. nonmetal

Topic 2: Compounds and Bonding

2.1 Atoms of different elements can join together by chemical bonds to form a compound. A compound has totally different properties from its elements. Chemical formulas show the ratio or number of atoms of each element in a compound. Example: 2 hydrogen atoms bonded to one oxygen atom make a water molecule (H2O).

2.2 Subscripts in a chemical formula represent the relative number of each type of atom. The subscript always follows the symbol for the element. Example: In a water molecule, H2O, there are 2 hydrogen atoms and one oxygen atom. Parentheses are used when a subscript affects a group of atoms. Example: the formula for magnesium nitrate is written Mg(NO3)2 to show that there is a ratio of one magnesium atom, 2 nitrogen atoms and 6 oxygen atoms in the compound.

2.3 Elements in groups 1, 2 and 13 (metals) will lose electrons and form positive ions. Elements in groups 15, 16 and 17 will gain electrons and form negative ions. Ionic compounds are formed by the attraction between positive and negative ions. The charges must be balanced, resulting in a compound with no net charge. The nomenclature for binary ionic compounds as directed by IUPAC, metal first keeping its name, non-metal second with an ending change of “-ide”.

Polyatomic Ions are a group of atoms bonded together that have a charge.

2.4 An ionic bond is formed when a metal (element from group 1 through 13) transfers electrons to a nonmetal (element from groups 15, 16 or 17). This is because metals form positive ions (by losing electrons) and nonmetals form negative ions (by gaining electrons), resulting in a

strong attraction between oppositely charged ions. Formulas for ionic compounds are written by balancing the ion charges. A covalent bond is formed when a nonmetal shares electrons with another nonmetal. Formulas for covalent molecules can be predicted from the dot diagrams of the combining elements. Sometimes more than one pair of electrons is shared between atoms. If two pairs are shared then there is a double covalent bond, three pairs shared is a triple covalent bond. Electron dot diagrams for elements show the number of valence electrons. Elements will transfer or share valence electrons in order to have eight valence electrons (octet rule).

2.5 The Lewis dot diagram for a covalent compound shows that each of the atoms in the molecule has a filled valence level. Each shared pair of electrons in a Lewis diagram represents a covalent bond. A covalent molecule can also be represented by a structural formula in which each covalent bond is shown as a line joining two atoms. In other words, a line in a structural formula represents two electrons. Covalently bonded compounds (molecular) use prefixes to show how many atoms of each element are present. Prefixes indicate 1 to 5 atoms are mono-, di-, tri-, tetra-, and penta-. Mono- is dropped when only one atom of the first element is present. Both ionic and covalent binary names end in “-ide”.

2.6 A covalent bond consists of electrons shared between atoms, but this sharing is not equal unless the two atoms involved are identical (like two chlorine atoms in a Cl2 molecule), because different atoms have different electronegativities. The more electronegative atom in a covalent bond will attract the electrons more strongly and this will result in it having a slight negative charge. The less electronegative atom will therefore be slightly deficient in electrons and so will have a slight positive charge. A covalent bond in which the atoms have slight electrical charges is known as a polar covalent bond. A non-polar covalent bond has equal sharing of electrons. A molecule is polar if it contains polar bonds and if its shape puts a positive and negative charge at different ends on the molecule. The shape of a molecule is determined by counting the areas of high electron density in the molecule. A linear shape results from two atoms bonded together, or from three atoms bonded together with no unshared electron pairs on the center atom (N2 and CO2). A tetrahedral shape results from four atoms bonded to a central atom and no unshared electron pairs (CH4). A pyramidal shape results three atoms bonded to a central atom and one unshared electron pair (NH3). A bent or angular shape results from two atoms bonded to a central atom and two unshared electron pairs (H2O).

2.7 Intermolecular forces are attractions resulting from forces between molecules. The strength of these forces will effect the vapor pressure (the pressure in a closed container that is due molecules that have evaporated of the liquid), boiling point (a liquid boils when its vapor pressure equals the pressure of the atmosphere) and surface tension (the stretching force among particles that produces a liquid “skin” on the surface) of the molecules. A hydrogen bond is an attraction occurring when a hydrogen atom bonded to a strongly electronegative atom is also attracted to another electronegative atom, often of a different molecule. A hydrogen bond is the strongest intermolecular force. London dispersion forces are the weakest intermolecular attraction between non-polar molecules. Dipole-dipole attraction is the attractive force resulting when polar covalent molecules line up so that the positive and negative ends are close to each other.

2.8 The Group 1 metals, the alkali metals, form ionic compounds (salts) with halogens in a ratio of 1:1. (Example: LiCl, NaCl, NaI, NaF, KCl), and with group 16 non-metals in a ratio of 2:1 (Example: Li2O, Na2O, Na2S, K2S) Group 2 metals, the alkaline earth metals, form ionic compounds with halogens in a ratio of 1:2 (Example: BeCl2, MgCl2, CaCl2, CaF2, SrI2), and with group 16 non-metals in a ratio of 1:1 (Example: CaO, MgO, CaS, MgS). Covalent bonding occurs between atoms that have relatively high electronegativities (between non-metals). Compounds can be predicted with dot diagrams. Group 18 elements (noble gases) do NOT naturally form compounds.

Check Your Understanding

1. Elements tend to gain or lose electrons in order to acquire the electron configuration of a a. halogen b. transition metal c. noble gas d. nonmetal 2. A double bond consists of a. two pairs of shared electrons b. two shared electrons

c. six shared electrons d. unshared electrons

3. When one atom is significantly more electronegative than the other one, a covalent bond between them is a. nonpolar b. hydrated c. polar d. very unstable 4. The type of chemical bonding in which electrons are transferred from one atom to another is a. nonpolar covalent b. polar covalent

c. ionic d. all of the above

5. An example of a polar covalent molecule would be a. NaCl b. HCl c. H2 d. O2 6. Unequal sharing of electrons is a characteristic of: a. covalent bonds b. ionic bonds

c. metallic bonds d. polar covalent bonds

7. Which of the following molecules does not have a linear shape? a. O2 b. H2S c. HI d. CO2 8. The structural formula for the nitrogen molecule (N2) contains a a. single bond b. double bond c. triple bond d. no bonds 9. A substance that is made up of molecules that have a partially positive and a partially negative end is a. nonpolar covalent b. polar covalent

c. ionic d. nonpolar ionic

10. The predicted geometry of NH3 is a. pyramidal b. trigonal planar c. bent d. linear 11. Choose the set of molecules which are in correct order of increasing polarity. a. C-Br , C-Cl , C-F , C-H , C-I b. C-F , C-Cl , C-H , C-Br , C-I

c. C-Cl , C-F , C-I , C-Br , C-H d. none of the above

12. Which of the following represents a tetrahedral molecule? a. CH4 b. CaCl2 c. NH3 d. Br2 13. Which of the following correctly matches the names and formulas of both compounds? a. AlCl3 . aluminum trichloride and N2O4 , nitrogen oxide b. AlCl3 , aluminum trichloride and N2O4 dinitrogen tetroxide c. AlCl3, aluminum chloride and N2O4 nitrogen oxide d. AlCl3 , aluminum chloride and N2O4 dinitrogen tetroxide 14. The name of K2SO4 is: a. dipotassium sulfate b. potassium sulfoxide

c. dipotassium sulfite d. potassium sulfate

15. The formula for Copper (I) Chloride a. CuCl2 b. Cu2Cl c. CuCl d. CuClO2 16. The formula for ammonium hydroxide is: a. AmOH b. AnOH c. NH4OH d. NH4(OH)2 17. The formula for iron (II ) oxide is : a. FeO b. Fe2O c. FeO2 d. Fe2O3 18. The formula for calcium nitrate is: a. CaNO3 b. Ca(NO3)3 c. Ca(NO3)2 d. Ca4NO2 19. The formula for ammonium phosphate is: a. NH4P b. (NH4)3PO4 c. AmPO4 d. (NH4)3(PO3)2 20. The formula for sulfur trioxide is: a. SO3 b. SO c. SO2 d. SO4

Topic 3: Kinetic Theory

3.1 Atoms and molecules are in constant motion. Particles of a gas move fastest; particles in a liquid move slower and particles in a solid move slowest. There is a direct relationship between temperature and speed of the particles. When the temperature increases, particles move faster. There is an inverse relationship between pressure and volume of gas particles. When the pressure increases, the volume decreases.

3.2 Pressure and temperature both affect the volume that a gas occupies. Pressure and volume are inversely related; if pressure increases, volume decreases.

Mathematically the relationship means that PV=k, i.e. that if all the factors remain constant then the pressure times the volume is also constant.

Absolute temperature and volume are directly related; if absolute temperature increases, volume increases when the temperature remains constant. Mathematically this relationship means that V/T=k, i.e. that if all other factors remain constant then the volume divided by the temperature is also constant.

3.3 Phase changes that require heat (like melting or boiling) are endothermic. H is positive for an endothermic change. This means heat goes in. Phase changes that give off heat (like freezing and condensing) are exothermic. H is negative. This means heat is released. Heating and cooling curves are also known as phase diagrams.

3.4 Elements form bonds to become more stable. A filled valence configuration (eight s and p electrons) s2p6 is very stable. Stability is inversely related to potential energy, therefore when atoms bond they become lower in potential energy. Potential energy is stored energy. Chemical bonds contain potential energy. Energy is required to break bonds. Breaking bonds is endothermic. Energy is released when bonds are formed. Forming bonds is exothermic.

3.5 In chemical reactions bonds are broken and new bonds are formed. The energy absorbed in breaking the bonds is never exactly equal to the energy released when the new bonds are formed. Therefore, all reactions are accompanied by a change potential energy that can be measured and is represented by the symbol H. An energy level diagram can also be used represent the energy change during a chemical reaction. Activation energy is the minimum amount of energy that must be supplied to a system to start a chemical change. Heat of reaction is the amount of energy absorbed or released during a chemical change. A catalyst is a substance added to a chemical reaction to increase the rate and that can be recovered chemically unchanged after the reaction is complete.

3.6 Dalton's Law states, the total pressure in a gas mixture is the sum of the partial pressures of the individual components. The partial pressure of any gas in a mixture can be calculated using the mole fraction of that gas in a mixture. Ptotal = P1 + P2 + P3 +..... The folowing mathematical relationship between the pressure, volume and temperature of a gas is used to describe the behavior of gases: P1 V1 = P2 V2 T1 T2

3.7 Hvap (Heat of vaporization) is the amount of energy needed for the particles of a substance to escape from the attractive forces of the other particles and escape from the surface into the gas phase. The stronger the forces of attraction between particles, the greater the heat of vaporization. Heat of fusion is the amount of energy released when the particles of a substance solidify. The stronger the attractive forces between the particles, the greater the heat of fusion. Heat of fusion (Hfus) is the amount of energy released when the particles of a substance solidify. The stronger the attractive forces between the particles, the greater the heat of fusion.

3.8 A reaction rate describes how rapidly a chemical change takes place. Reaction rates are determined experimentally by measuring a change in some physical property such as volume, temperature, color, mass or pH. There is a direct relationship between temperature and reaction rate. Reversible reactions reach equilibrium. At equilibrium, the forward and reverse reactions occur at the same rate. Le Chatelier’s Principle states, when a system at equilibrium is disturbed by applying stress, the equilibrium position shifts to relieve the stress. Stresses that can change equilibrium include changes in concentration, temperature or pressure. Check Your Understanding

1. Equal molar quantities of gases A, B, C, and D are put into an evacuated flask. The pressure of the sample is measured as 160. kPa. The pressure of gas C is a. less than 20.0 kPa b. between 20.0 and 60.0 kPa

c. between 60.0 and 100. kPa d. greater than 100.0 kPa.

2. A 30. mL sample of oxygen is at a temperature of 66oC. If the temperature is lowered to 33oC at constant pressure, the volume of the gas will become a. 15 mL b. 27 mL

c. 45 mL d. none of the above.

3. A gas at 50.0 kPa has a volume of 4.0 dm3. If the temperature is held the same and the pressure on the gas in reduced to 10.0 kPa, the volume of the gas would become a. 0.80 dm3 b. 4.0 dm3 c. 20. dm3 d. 40. dm3. 4. If 1.0 moles of carbon dioxide gas is introduced into an empty vessel at a pressure of 1.2 atm and 55oC, the volume of the gas would be a. 22.4 L b. 32.3 L c. 2270 L d. 3270 L. 5. Compute the number of moles of a gas in a 5000. L tank if the gas temperature is 44C and its pressure is 101.3 kPa. a. 0.00520 moles b. 192 moles c. 1410 moles d. 18900 moles. 6. Given the reaction at equilibrium:

H2(g) + Cl2(g) 2HCl(g) + energyThe equilibrium will shift to the right when a) chlorine gas is removed b) hydrogen gas is removed

c) hydrogen chloride gas is added d) temperature is decreased

7. In the reaction: N2(g) + 3H2(g) ---->2NH3(g) + 22.0 kcal

a) the reaction is both endothermic and exothermic b) the reaction is endothermic c) the reaction is exothermic d) the reaction is neither endothermic or exothermic 8. On the energy diagram, the energy of the activated complex is represented by:

a) a b) b c) c d) d 9. If the phases of matter are arranged in order of increasing disorder, the arrangement would be: a. solid, liguid, gas b. gas, solid, liquid c. gas, liquid, solid d. liquid, solid, gas 10. How many kilojoules of energy are needed to convert 2.5 moles of water from ice to liquid if the heat of fusion of water is 6.00 kJ/mol? a. 2.5 kJ b. 6.00 kJ c. 15 kJ d. 30kJ

11. The enthalpy change for melting a solid, such as ice, is called: a. heat of fusion b. heat of vaporization

c. heat capacity d. specific heat

12. At chemical equilibrium, the rates of the forward reaction and reverse reactions are: a. equal to 0 b. equal to each other

c. at a maximum d. at a minimum

13. The rate of a chemical reaction normally: a. increases as reactant concentration increases. b. is slowed down by a catalyst. c. decreases as temperature increases. d. decreases as surface area increases. 14. Activation energy is: a. the heat released in a reaction. b. the minimum energy colliding particles must have to react. c. the energy given off when reactants collide d. generally very high for a reaction that takes place rapidly. 15. In general, increasing the temperature causes the rate of most chemical reactions to: a. increase b. decrease

c. remain the same d. vary unpredictably

16. The principle that relates changes imposed on equilibrium systems to equilibrium position is: a. Haber's Law b. the law of chemical equilibrium

c. Le Chatelier's Principle d. Avogadro's Principle

Topic 4: The Mole and Stoichiometry

4.1 Atoms and molecules are too small to count. Mole is the unit used to tell how many particles are in a certain amount of a substance. A mole is 602,000,000,000,000,000,000,000 particles (atoms or molecules). Expressed in scientific notation, a mole is 6.02 x 1023 particles. Scientific notation is used to express very small or very large measurements in powers of ten. It expresses quantities by using a number between one and ten, which is then multiplied by ten to a power to give the quantity its proper magnitude.

4.2 The sum of the protons and neutrons in an atom is known as the mass number. The number of neutrons in an atom can be found by subtracting the atomic number from the mass number. Isotopes are atoms of the same element that have different numbers of neutrons. Some isotopes are radioactive, many are not.

4.3 The molar mass of a compound is the mass of one mole of the compound. It is found by taking the sum of the molar masses of the individual elements that make up the compound. The percent mass of an element in a compound can be determined:

% by mass of element = total mass of element in compound X 100 total mass of compound

4.4 Molar masses from the periodic table can be used to calculate the number of moles in a given mass of an element or compound. This is because the masses on the periodic table represent the number of grams in one mole. The number of moles can also be used to calculate the number of particles – atoms or molecules. The number of particles can be determined from the mass of a compound or element.

4.5 Because matter cannot be created or destroyed, the total mass of the products is equal to the total mass of the reactants in a chemical reaction. Molar masses from the periodic table

and mole ratios from the balanced equation can be used to calculate the mass of a reactant or product.

4.6 Balanced chemical reactions and numbers of moles of each substance can be used to predict the masses of reactants or products. At STP (standard temperature and pressure, 0 Celsius and 1 atmosphere OR 760 mmHg OR 101.2 kPa) the volume of 1 mole of any gas is 22.4 liters.

4.7 The number of moles of a gas (n) can be determined if the pressure (P), temperature (T) and volume (V) of the gas sample are known, using the constant R according to the following equation: PV = nRT is the ideal gas law constant and has two values depending on the pressure units. They are R = 8.314 L.Kpa/mol.K and R = 0.0821 L.atm/mol.atm. 4.8 An empirical formula shows the smallest whole number ratio of elements in a compound. Ionic solids are composed of oppositely charged ions arranged in a regular, repeating, crystal lattice structure; the empirical formula always gives the ratio of positive to negative ions. Covalent compounds are often in the form of individual molecules; the empirical formula gives the ratio of atoms in one molecule. Example: The molecular formula for glucose is C6H12O6; the empirical formula is CH2O.

Choose the best answer that either answers the question or completes the statement

1. The number 2 X 101 expressed in common numerical expression is: a) 200 b) 20 c) 2 d) 0.02 2. The number 5.10 X 102 expressed in common numerical expression is: a) 501 b) 510 c) 5100 d) .00501 3. The number 300 expressed in scientific notation is a) 3 X 10 -1 b) 3 X 10 -3 c) 3 X 10 2 d) 3 X 10 1 4. The number 0.0006 expressed in scientific notation is: a) 6 X 10 -4 b) 6 X 10 2 c) 6 X 10 -2 d) 6 X 10 4 5. How many moles are in 8.5 x 10 25 molecules of CO? a. 1.4 x 10 2 b. 7.1 x 10 -3 c. 5.1 x 10 49 d. 8.5 x 10 25 6. What is the molar mass of CO2? a. 36.0 g b. 11.0 g

c. 44.0 g d. 6.02 x 10 23 g

7. It is possible to convert moles to atoms by: a. multiplying by 6.02 x10 23 b. dividing by 6.02 x10 23

c. multiplying by the molar mass d. dividing by the molar mass

8. The percentage composition of ammonia (NH3) is: a. 78.5 % N and 21.5 % H b. 21.5 % N and 78.5 % H

c. 82.4 % N and 17.6 % H d. 17.8 % N and 82.2 % H

9. How many molecules of sulfur dioxide are present in 1.60 moles of sulfur dioxide? a. 9.63 x 10 23 b. 102.1 x10 1

c. 7.62 x 10 1 d. 3.76 x 10 23

10. Find the number of moles in 3.30 g of (NH4)2SO4 a. 132.1 b. 40.0 c. 0.279 d. 0.0250 11. What is the volume of 2.50 moles of carbon monoxide at STP? a. 0.112 L b. 3.10 L c. 56.0 L d. 8.96 L 12. The volume of 2.00 moles of any gas at STP is : a. 11.2 L b. 22.4 L c. 44.8 L d. 3.2 L 13. Which contains more atoms? a. 1.00 mole H2O2 b. 1.00 mol C2H6

c. 1.00 mol CO d. 1.00 mol H2O

14. The volume of 1.20 mol of Ar at STP is a. 170 L b. 9.9L c. 27 L d. 26.9 L 15. The percent composition of H2S is a. 5.9 % H and 94.1 % S b. 5.1 % H and 94.1 %S

c. 9.42 % H and 4.1 % S d. 50 % H and 50 % S

16. How many liters of hydrogen react with 2.00 mol of nitrogen at STP in the following reaction

N2 + 3H2 NH3?a. 3.0 L b. 22.4 L c. 67.2L d. 134 L 17. How many molecules of NO2 are produced when 2.0 x 10 20 molecules of N2O4 are decomposed according to the equation: N2O4 (g) 2NO2 (g) a. 4 b. 1.0 x 10 20 c. 2.0 x10 20 d. 4.0 x 10 20 18. Given the reaction: 2H2O 2H2 + O2, how many moles of H2O would be required to produce 2.5 moles of O2 ? a. 2.0 mol b. 2.5 mol c. 4.0 mol d. 5.0 mol 19 Given the reaction: CuO + H2 Cu + H2O, how many moles of H2O are produced when 240 grams of CuO react? a. 1.0 mol b. 3.0 mol c. 18 mol d. 54 mol 20. Gven the reaction: Zn + H2SO4 ZnSO4 + H2, how many grams of H2SO4 are required to produce 1.0 g of H2? a. 1.0 g b. 2.0 g c. 49.1 g d. 98.0 g Topic 5: Chemical Reactions 5.1 A chemical reaction is required to change one substance into another by rearranging its atoms. The only way to form a compound from elements is by a chemical reaction. Example: In a synthesis reaction (combination reaction), hydrogen gas and oxygen gas react to form water. 2H2 + O2 2H2O The only way to separate a compound into its elements is by a chemical reaction that breaks the chemical bonds, forming new substances. Example: water decomposes to form hydrogen and oxygen gas. 2H2O 2H2 + O2

5.2 A chemical equation is a record of what happens in a chemical reaction. It shows the formulas of all the reactants on the left hand side of the arrow, and the formulas for all the products on the right hand side.

5.3 Combustion reactions are exothermic reactions in which oxygen combines with other elements. One example is the reaction between methane and oxygen in a Bunsen burner: CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + energy

5.4 Because matter cannot be created or destroyed, elements must be conserved in a chemical reaction (Conservation of Mass). There must be the same number of each kind of atom on both sides of a balanced equation. The only way to balance a chemical equation is by placing coefficients in front of each substance until each side has the same number of atoms of each element.

5.5 When two or more substances combine to form a single product, the reaction is called a synthesis reaction. For example, the formation of water from hydrogen and oxygen gases is a synthesis reaction: 2H2 (g) + O2 (g) 2H2O In a decomposition reaction, a compound breaks down into two or more simpler substances. For example, in electrolysis, water is broken down into hydrogen and oxygen gases: 2H2O (l) 2H2 (g) + O2 (g)

5.6 In a single replacement reaction (review single replacement reactions here) one element takes the place of another in a compound. In a double replacement reaction the positive portions of two ionic compounds are interchanged.

5.7 In a sample of pure water a very small number of water molecules dissociate, producing equal concentrations of both hydrogen ions [H+] and hydroxide ions [OH-]. The pH of pure water is 7. H2O H+ + OH- or 2H2O H3O+ + OH-

5.8 Neutralization reactions result from the reaction of an acid with a base to form a salt and water. These reactions are usually double replacement reactions. HCl + NaOH NaCl + HOH.

Check Your Understanding

1. Which of the following types of reactions results in a single product? a. combination /synthesis b. decomposition

c. single replacement d. double replacement

2. The numbers used to balance a chemical equation are called: a. superscripts b. subscripts c. coefficients d. formula units 3. A chemical equation is balanced when ____________ . a. the equation shows an equal number of atoms for each element on both sides. b. at least one substance in each of the three physical states is present. c. the total number of moles of the reactants equals the moles of the products. d. none of the above 4. The general form for a double displacement reaction a. element + compound element + compound b. compound two or more elements or compounds c. element or compound + element or compound compound d. compound + compound compound + compound 5. In the reaction 2KClO3 2KCl + 3O2 oxygen is a _____________. a. reactant b. product c. coefficient d. subscript 6. In the equation 2Fe + 3H2O Fe2O3 + H2, iron is a ________. a. subscript b. reactant c. product d. coefficient7. The ionic compound formed in an acid-base neutralization reaction is a(n) : a. indicator b. hydroxide c. salt d. hydride 8. The products of the neutralization reaction between HNO3 and Ca(OH)2 are: a. 2CaNO3 + H20 b. Ca(NO3)2 + H2O

c. CaNO3 + H2O d. Ca(NO3)2 + 2H2O

9. In what ratio do HCl and Mg(OH)2 react through neutralization? a. 1:1 b. 2:1 c. 1:2 d. 2:2 10. The coefficients needed to balance the equation: Zn + CuSO4 ZnSO4 + Cu are a. 1,2,2,1 b. 2,1,2,1 c. 1,1,2,2 d. balanced 11. The coefficients needed to balance the equation: H2 + O2 H2O are a. 2,1,1 b. 1,1,2 c. 2,1,2 d. balanced 12. The coefficients needed to balance the equation: PCl5 PCl3 + Cl2 are a. 2,2,1 b. 1,2,1 c. 1,1,2 d. balanced 13. The coefficients needed to balance the equation: PbCl2 + Li2SO4 LiCl + PbSO4 are a. 1,1,2,1 b. 1,2,2,1 c. 2,2,1,1 d. balanced 14. The coefficients needed to balance the equation: Zn + HCl ZnCl2 + H2 are a. 1,1,2,1 b. 1,2,1,1 c. 1,1,1,2 d. balanced 15. The coefficients needed to balance the equation: AgNO3 + CaCl2 AgCl + Ca(NO3)2 are a. 2,1,2,1 b. 1,2,1,2 c. 1,1,2,2 d. balanced 16. The coefficients needed to balance the equation: Zn + FeCl2 ZnCl2 + Fe are

a. 1,3,1,3 b. 3,1,3,1 c. 1,1,3,1 d. balanced 17. The chemical reaction in which hydrogen ions from an acid react with hydroxide ions from a base to produce water is: a. ionization b. titration c. concentration d. neutralization 18. The reaction H2CO3 H2O + CO2 is an example of a reaction that is a. synthesis b. decomposition

c. single replacement d. combustion

19. The equation Cu + Hg(NO3)2 Hg + Cu(NO3)2 exemplifies a reaction that is a. synthesis b. single replacement

c. neutralization d. double replacement

20. Which of the following is a neutralization reaction? a. K + Cl2 KCl b. NaOH + NH4Cl NH4OH + NaCl

c. Cu(OH)2 CuO + H2O d. LiOH + HBr LiBr + H2O

Topic 6: Solutions

6.1 If a substance contains different types of particles, then it is called a mixture. (Example: Soapy water is a mixture of many water molecules and a few soap particles). In a heterogeneous mixture, the different parts can be easily seen (like salt and pepper mixed together). In a homogeneous mixture the particles are mixed so well that the separate parts cannot be seen (like salt dissolved in water).

6.2 A solution is a homogeneous mixture because the separate parts of the mixture cannot be seen. The solvent (usually water) is the part of the solution that is present in largest amount. The solute is the substance that is dissolved. A saturated solution has all the dissolved solute that it can hold, and can be identified by undissolved solute particles on the bottom after mixing. An unsaturated solution can still hold more solute.

6.3 Dissolving is a physical change that involves heat. Dissolving and dissociation can be represented by an equation. Example: NaCl(s) + heat Na+(aq) + Cl-(aq)

H is positive. If a solution gets cooler when a solute dissolves, it is an endothermic change and H is positive, and heat is written to the left of the arrow. If a solution gets warmer when a solute dissolves, it is an exothermic change and H is negative, and heat is written to the right of the arrow.

6.4 Solutions that contain ions are called electrolytes because they can conduct an electric current. Therefore, solutions of ionic compounds (salts) in water (aqueous solutions) are electrolytes, because ionic compounds dissociate as they dissolve. Conductivity is directly related to the number of ions in the solution. Strong electrolytes are good conductors while weak electrolytes are poor conductors. Solutions that do not contain ions are called non-electrolytes because they cannot conduct an electric current.

6.5 The concentration of a solution is the amount of solute contained in a certain volume of solution. If a solution contains a small amount of solute it is called dilute, and if it contains a large amount of solute it is called concentrated. In chemistry, concentration is given as molarity, the number of moles of the solute in one liter of solution and expressed as mol/L or just M.

6.6 The general rule for predicting solubility is "like dissolves like". Water is a polar substance, so it can dissolve ionic and polar solutes. Oil is non-polar, so oil will not dissolve in water. Oil and water don’t mix but different oils do because a non-polar solute will dissolve in a non-polar solvent.

6.7 pH (0-14) measures the hydrogen ion concentration in water. Each pH unit involves a tenfold change in hydrogen ion concentration. The pH number and the [H+] are inversely related because pH = - log [H+]

o An increase of 1 pH unit means that the hydrogen ion concentration has decreased 10 times.

o A decrease of 1 pH unit means that the hydrogen ion concentration has increased 10 times.

Acids are compounds that increase the concentration of hydrogen ions [H+] when they dissolve in water. Acid solutions have a pH below 7, taste sour and turn litmus paper red. Bases are compounds that increase the concentration of hydroxide [OH-] when they dissolve in water. Bases have a pH greater than 7, taste bitter, feel slippery and turn litmus paper blue.

6.8 Both strong acids and strong bases dissociate completely in water, therefore are strong electrolytes. In a solution of a strong acid like hydrochloric acid, almost all of the HCl molecules dissociate according to the following equation:

HCl(aq) --> H+(aq) + Cl-(aq)Weak acids and weak bases are weak electrolytes. In a solution of a weak acid like acetic acid, only a few of the CH3COOH molecules dissociate:

CH3COOH <---> H+(aq) + CH3COO-(aq) Check Your Understanding

1. Molarity is expressed as: a. moles of solvent/liter of solution b. moles of solute/liter of solution

c. moles of solute/mole of solvent d. moles of solute/kilogram of solution

2. A saturated solution: a. contains the maximum amount of solute that can be dissolved b. is concentrated c. is diluted d. none of the above 3. The molarity of a solution that contains 14 g KOH per 150 mL of solution is: a. 93 M b. 1.7M c. 0.093 M d. 11 M 4. Another name for a solution is a: a. heterogeneous mixture b. homogeneous mixture

c. compound d. element

5. The dissolved substance in a solution is called the : a. solute b. solvent c. hydrate d. tincture 6. How many moles of solute are present in 1.25 L of a 0.75M NaNO3 solution? a. 1.7 b. 0.60 c. 0.75 d. 0.94 7. If salt is dissolved in water, water serves as the: a. solute b. solvent

c. dissolved medium d. none of these

8. What is the molarity of a solution that contains 8 moles of solute in 2L of solution? a. 4M b. 8M c. 6M d. 0.25M 9. In the reaction: N2(g) + 3H2(g) <----> 3NH3(g) + heat a) the reaction is both endothermic and exothermic b) the reaction is endothermic c) the reaction is exothermic d) the reaction is neither endothermic or exothermic 10. Which of the following compounds will not dissolve in water? a. KCl b. CCl4 c. CaBr2 d. MgCl2 11. Which of the following is true? a. wood is a good conductor of electricity b. a solution that contains electrolytes conducts electricity c. a water solution of an acid will not conduct electricity d. a water solution of a base will not conduct electricity 12. A solution with a pH of 9 has an [OH-] concentration of :

a. 1.0 x 10 -14 b. 1.0 x 10 -9 c. 1.0 x 10 -5 d. 1.0 x 10 -7 13. Among the following, which is the strongest acid? a. [ H+] = 1 x 10 -5

b. pH = 3 c. [ OH -] = 1 x 10 -7

d. pH = 10

14. Which of the following is true about bases? a. have a bitter taste b. feel slippery

c. react with acids to form a salt and water d. all of these

15. A solution in which the hydroxide - ion concentration is 0.00001 is: a. acidic b. basic c. neutral d. none of these 16. In a neutral solution, the [H+] is: a. 10 -14 b. zero c. 1 x 10 -4 d. equal to [ OH-] 17. Because acids and bases are conductors of electricity, they are referred to as: a. indicators b. electrolytes c. insulators d. nonelectrolytes 18. A solution with a pH of 5.0: a. is acidic b. is basic

c. is neutral d. has a hydrogen ion concentration of 5.0 M

19. Of the following, the best conductor of electricity is: a. solid salt b. solid sugar c. aqueous salt d. distilled water 20. A solution which has a pH of 12 would be: a. acidic b. basic c. neutral d. none of these

Topic 7: Experimentation

7.1 Safety equipment is used to protect the eyes and skin from contact with laboratory chemicals and flames: goggles, aprons, gloves, safety shower, eyewash, broken glass container, fume hood and fire blanket. You must understand and follow the laboratory safety rules and procedures that are described in your Safety Contract in order to work in the chemistry laboratory.

7.2 Historically, scientists have been known to stand on the "shoulders of giants", meaning that they build on previous knowledge to make new discoveries. This is especially true as scientists began to develop theories concerning the atom and its structure. From John Dalton, to J.J Thompson to Ernest Rutherford to Neils Bohr to the modern day electron cloud (quantum mechanical) model of the atom.

7.3 Significant figures are used in making calculations with measurements made in the lab. Measure volume of a liquid in milliliters (mL) using a graduated cylinder and stating measured digits plus the estimated digit. Measure mass in grams (g) using an electronic balance and identifying the estimated digit. Significant figures are used in making calculations with measurements made in the lab. Determine the mean (average) of a set of volume or mass measurements using the rules for significant digits.

7.4 Percent error is the ratio of absolute value of the difference between the experimental value and the theoretical value to the theoretical value, multiplied by 100 |Theoretical value - Experimental value | X 100 Theoretical value

7.5 Describe and demonstrate safe techniques for lighting and using gas burners. Understand and use Material Safety Data Sheet (MSDS) warnings including: handling chemicals, lethal dose (LD) hazards, disposal and chemical spill clean-up.

7.5 Percent yield is the ratio of actual yield to theoretical yield, multiplied by 100. Actual yield X 100 Theoretical yield

7.6 Accuracy is how close a measurement is to the true value. An accurate measurement has very little error. Precision is how exact and repeatable a measurement is.

7.7 To safely dilute an acid, add acid to water. Do not add water to a acid concentrated acid. Always perform this task in a fume hood. When preparing a more dilute solution from a stock solution of known concentration, M1 X V1 = M2 X V2. The molarity of the first solution multiplied by its volume will be equal to the molarity of the new solution multiplied by its volume. Phenolphthalein is an indicator that is colorless in the presence of an acid or neutral substance but is pink in the presence of a base. Litmus is also an acid/base indicator. Base turn red litmus blue and acids turn blue litmus red.

7.8 Titration uses a buret to dispense precise amounts of solution of known concentration to determine the concentration of another solution.

Check For Understanding

1. The closeness of a measurement to it's true value is a measure of it's: a) accuracy b) repeatability c) precision d) nearness 2. According to the method of significant figures, the number of digits that are estimated in a measurement is: a) none b) one c) two d) three 3. What two components does a measurement always contain? a) a number and a decimal point b) a power of ten and a unit

c) a number and a unit d) a number only

4. Safety goggles should be worn in the chemistry laboratory a) while performing an experiment b) during clean-up c) during all parts of the lab and during clean-up d) only when using a Bunsen Burner 5. If your lab partner's clothing should catch on fire you should: a) tell him/her to run until they find a security guard. b) smother the flames with a fire blanket c) do nothing, because the flames will eventually die out d) show him/her the door and hope for the best 6. What should you do if you burn your hand? a) run hot water on it b) run cold water on it and notify the teacher immediately c) put your hand in your pocket and continue with your experiment d) don't do anything, you should have known better 7. When heating a liquid in a test tube you should: a) slant the test tube away from yourself and others b) hold the test tube in the vertical position c) point the test tube toward your lab partner d) point the test tube toward the teacher 8. How many significant figures in the measurement 2103.2g? a) 2 b) 3 c) 4 d) 5 9. How many of the zeros in the measurement 0.000040300 are significant? a) 8 b) 6 c) 5 d) 3 10. Two pieces of equipment that should be worn for every lab are: a) heavy winter coat and hat b) gym shorts and a t-shirt

c) goggles d) goggles and an apron

11. The number of significant figures in the measurement 0.070g is: a) 1 b) 2 c) 3 d) 4 12. If the mass of a dry beaker is 19.02 grams and increases to 22.40 grams when a sample is added, what is the mass of the sample? a) 22.40 g b) 41.42 g c) 3.38 g d) 1.10 g

13. If the volume of water in a cylinder is 8.0mL, but changes to 10.0mL when a solid is carefully lowered into it, the volume of the solid is: a) 2.0mL b) 10.0mL c) 18.0mL d) 80. 0mL 14. A student estimated a mass to be 250g, but upon carefully measuring it, found it to be 240g. What is the percent error of the estimated mass? a) 4.0% b) 4.2% c) -4.0% d) -4.2% 15. A beaker has a mass of 89.67g. When a solid is added, the beaker plus the solid have a mass of 92.25g. What is the mass of the solid? a) 1.92g b) 2.82g c) 2.58g d) 12.58g 16. A particular reaction is expected to produce 2.6 L of oxygen gas. In reality, the reaction only produces 1.9 L of oxygen gas. The percent yield of the reaction: a) 27 % b) 42 % c) 73 % d) 85 % 17. What type of instrument is used mainly in a titration experiment? a) buret b) bunsen burner c) beaker d) test tube 18. Phenolphthalein is an indicator for what type of solution? a) acidic solution b) basic solution

c) NaCl aqueous solution d) none of the above

Additional Resources Practice SOL Tests : http://education.jlab.org/solquiz/ Interactive Periodic Table : http://site.ifrance.com/okapi/periodic3.htm Periodic Table Trends :

http://courses.chem.psu.edu/chem12/spring/transparencies/pdfs/pm-lec12(4).pdf Bonding, naming and properties of common compounds :

http://users.senet.com.au/%7Erowanb/chem/chembond.htm Intermolecular bonds : http://www.cs.stedwards.edu/%7Ewright/text/chembond.html Lewis Structures : http://www.chem.uncc.edu/faculty/murphy/1251/slides/C18a/ Gas Law Formulas : http://www.pmel.org/HandBook/HBpage20.htm Kinetic Energy : http://plabpc.csustan.edu/general/tutorials/temperature/temperature.htm Solutions Review : http://members.aol.com/profchm/solintro.html Polarity : http://library.thinkquest.org/15567/lessons/14.html Ionization of Water :

http://www.biology.arizona.edu/biochemistry/tutorials/chemistry/page3.html Percent error : http://www.ric.edu/bgilbert/s3pcerr.htm Percent Yield : http://www.pathcom.com/~ngjdw/laws.htm Significant figures : http://dbhs.wvusd.k12.ca.us/SigFigs/SigFigRules.html Titration : http://www.dartmouth.edu/~chemlab/techniques/titration.html History of Atom : http://perso.club-internet.fr/molaire1/e_histoire.html General SOL Review : http://www.wise.k12.va.us/alted/SOL/sol.htm General SOL Review : http://www.quia.com/pages/sol12.html

Some General and SOL-specific Test Taking Hints and Strategies

1. Get a good solid night's sleep before the test. Being well rested will sharpen the mind and aid your memory. Eat a good breakfast the morning of the test.

2. Try to make the morning of the test a pleasant one. Avoid stress. Have your things ready and organized and get to school on time so that you’re not rushing around.

3. Listen carefully to the instructions from the teacher, and read the directions to each question. 4. There are 60 questions on the SOL. Only 50 count, though. The other ten are questions that are under

review. There’s no way to tell which is which, though, so do your best on all of them!5. You do NOT have to answer all the questions correctly to pass. It is not expected that you can answer

every question correctly. Currently on the Chemistry SOL, you only need to get 27 out of 50 correct. That comes out to be a 54%. Other SOL tests have other percentages.

45 out of 50 is advanced!6. Attempt to answer ALL of the questions and DO NOT leave any blanks. There is no penalty for

guessing.

7. The computer has some things that will help in taking the test - i.e., underlining important words, etc. You should always highlight words like EXCEPT, NOT and other words that limit the answers.

8. Read each question carefully and think before you answer. Be sure that you understand the question before you start to compare answer choices. This is especially important on the questions containing graphs and charts or any of these words: CHOOSE, DESCRIBE, EXPLAIN, COMPARE, IDENTIFY, SIMILAR, EXCEPT, NOT, and BUT.

9. Look at ALL of the answer choices and choose the best and most complete answer. 10. If you're not sure which answer is correct, eliminate choices that you know are incorrect. Then focus on

the remaining choices. Elimination ideas:

Eliminate the answers you know are wrong. Eliminate the answers that you know don’t have anything to do with the question. You can usually eliminate answers with ALWAYS, NEVER, EVERYONE and words

like that in them. If it’s a math question, you can usually eliminate any numbers that appear in the

question. If the questions are all numbers, usually the lowest and the highest are wrong. If you’ve already used a vocabulary word once on the test and you’re sure it’s

right, then don’t use it again.11. A word of caution about changing answers - usually your first choice is correct. Only change answers if

you’re certain the first choice was wrong. 12. If you get stuck on a question, use the Review feature and go back to it after finishing the test. You

might want to jot down a note or two about the question on your scrap paper. That way, if you come to another questions on the same subject, you can use it to help you with the first question, maybe.

13. Stay focused on the test, even if other students finish early. Don't get distracted. No one really cares if you finish first. There’s no medal or prize or extra credit and no one really thinks that you’re super-smart if you finish quickly.

14. If you need to take a break or you’re starting to feel overwhelmed, then rest a minute or two. Close your eyes, take a deep breath, and think happy thoughts. Draw a picture of your teacher on the scrap paper. Draw a large bus coming towards him. Do something to relax a bit.

15. Here’s a hint to help you double-check each and every answer. a) You get a piece of scrap paper with each test. Number one edge of the paper 1-60. b) Take the test, but instead of marking the answer on the computer screen, write it on the

scrap paper after the question number. c) Double check to make sure you marked all sixty questions. d) Flip the scrap paper over. e) Take the test again, but this time mark the correct answer on the computer screen. Do NOT

look at your paper answers!f) Once you’re finished with the test the second time, compare the answers on the screen to

the answers on your scrap paper. If you came up with the same answer twice, you’re great. If not, figure out why. If you can’t figure out why you changed your mind, stick with your first answer!

g) If you’re really, really a Type-A personality (you know who you are), you can do the whole scrap-paper thing twice if you want and then you’ll have three versions of your answer to compare.

h) Once all your answers match or you’ve figured out why you changed your mind, raise your hand to submit your answers.

The Actual Chemistry SOLs

CH.1 The student will investigate and understand that experiments in which variables are measured, analyzed, and evaluated produce observations and verifiable data. Key concepts include

designated laboratory techniques; safe use of chemicals and equipment; proper response to emergency situations; manipulation of multiple variables, using repeated trials; accurate recording, organization, and analysis of data through repeated trials; mathematical and procedural error analysis; mathematical manipulations (SI units, scientific notation, linear equations, graphing, ratio and

proportion, significant digits, dimensional analysis);

use of appropriate technology including computers, graphing calculators, and probeware, for gathering data and communicating results; and

construction and defense of a scientific viewpoint (the nature of science).CH.2 The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of

average atomic mass, mass number, and atomic number; isotopes, half lives, and radioactive decay; mass and charge characteristics of subatomic particles; families or groups; series and periods; trends including atomic radii, electronegativity, shielding effect, and ionization energy; electron configurations, valence electrons, and oxidation numbers; chemical and physical properties; and historical and quantum models.

CH.3 The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and balanced equations. Key concepts include

nomenclature; balancing chemical equations; writing chemical formulas (molecular, structural, and empirical; and Lewis diagrams); bonding types (ionic and covalent); reaction types (synthesis, decomposition, single and double replacement, oxidation-reduction,

neutralization, exothermic, and endothermic); and reaction rates and kinetics (activation energy, catalysis, and degree of randomness).

CH.4 The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. Key concepts include

Avogadro’s principle and molar volume; stoichiometric relationships; partial pressure; gas laws; solution concentrations; chemical equilibrium; and acid/base theory: strong electrolytes, weak electrolytes, and nonelectrolytes; dissociation and

ionization; pH and pOH; and the titration process.CH.5 The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction between particles. Key concepts include

pressure, temperature, and volume; vapor pressure; phase changes; molar heats of fusion and vaporization; specific heat capacity; and colligative properties.