Day 1 8/29 What is Chemistry? Syllabus/Safety/Measurement 2.2.2 Analyze the evidence of chemical change Units 1 : Math & Measurement Identify Lab Safety and Equipment Observe a chemical reaction and identify indicators of a chemical reaction Day 2 8/30 Significant digits and measurements Sig figs and measurement lab Day 3 8/31 2.2.2 Classification of matter 2.2.2 Analyze the evidence of chemical change 2.2.2 Analyze the evidence of a physical change Classify matter Contrast and identify chemical and physical changes Day 4 9/1 1.1.1 Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). 1.1.1 Use symbols: A= mass number, Z=atomic number 1.1.1 Use notation for writing isotope symbols: 235 92U or U-235 1.1.1 Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. 1.1.1 Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes Unit 2: Atomic Structure and Nuclear Chemistry List numbers of protons, electrons, neutrons in different atoms Use Isotope notations Demo and Pretest/Post-test Highlight key words to use in a concept map Post definitions (word wall) Carbon Snake Demo Which atomic symbol represents an isotope of sulfur with 17 neutrons? a. 17 16 X b. 33 16 X c. 17 32 X d. 49 32X Day 5 9/2 1.1.1 Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. 1.1.1 Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes Use isotopes to calculate average atomic mass of an element Rubidium is a soft, silverywhite metal that has two common isotopes, 85 Rb and 87 Rb. If the abundance of 85 Rb is 72.2% and the abundance of 87 Rb is 27.8%, what is the average atomic mass of rubidium?
Day 6 9/6 1.1.1 Identify ions using mass number and atomic number and relate to number of protons, neutrons and electrons. Labor Day Ions and Element Math Determine the number of p, e, and n in 35 Cl 1- Day 7 9/7 1.1.4 Use the symbols for and distinguish between alpha ( 4 2He), and beta ( -1 0e) nuclear particles, and gamma (γ) radiation. 1.1.4 Compare the penetrating ability of alpha, beta, and gamma radiation. 1.1.4 Use shorthand notation of particles involved in nuclear equations to balance and solve for unknowns. 1.1.4 Using symbols to represent simple balanced decay equations Nuclear Decay Write the balanced decay equation for the alpha particle decay of U-238. Day 8 9/8 1.1.4 Compare radioactive decay with fission and fusion. 1.1.4 Half-life (including simple calculations) Half-life The half-life of a radioactive isotope is 20 minutes. What is the total amount of 1.00 g of sample of this isotope remaining after 1 hour? Day 9 9/9 Half-Life Lab Day 10 9/12 Cumulative Activity / Unit Wrap-Up Day 11 9/13 UNIT TEST Unit 3: Quantum Theory/Periodic Trends Day 1 9/14 1.1.2 Analyze diagrams related to the Bohr model of the hydrogen atom in terms of allowed, discrete Identify the parts of a wave An electron in an atom of hydrogen goes from
energy levels in the emission spectrum. 1.1.3 Understand that energy exists in discrete units called quanta. 1.1.3 Describe the concepts of excited and ground state of electrons in the atom: 1.1.3 Gaining energy results in the electron moving from its ground state to a higher energy level. 1.1.3 When the electron moves to a lower energy level, it releases the energy difference in the two levels as electromagnetic radiation (emissions spectrum). Rank waves via energy Interpret the Bohr model to identify the type of light energy level 6 to energy level 2. What is the wavelength of the electromagnetic radiation emitted? a. 410 nm b. 434 nm c. 486 nm d. 656 nm Day 2 9/15 1.3.2 Write electron configurations, including noble gas abbreviations (no exceptions to the general rules). Included here are extended arrangements showing electrons in orbitals. Write the electron configuration for a given element Day 3 9/16 Write the orbital filling diagram for a given element 1.3.2 Identify s, p, d, and f blocks on Periodic Table. Day 4 9/19 Write the noble gas notation of a given 1.3.2 Identify an element based on its electron element configuration. (Students should be able to identify elements which follow the general rules, not necessarily those which are exceptions.) Draw the Lewis dot diagram of a given element Which is the electronic configuration of calcium? a. 1s2 2s2 2p6 3s2 3p8 b. 1s2 2s2 2p6 3s2 3p6 4s2 c. 1s2 2s2 2p6 3s2 3p6 3d2 d. 1s2 2s2 2p8 3s2 3p6 Day 5 9/20 Chm.1.3.1 1.3.1 Classify the components of a periodic table (period, group, metal, metalloid, nonmetal, transition). Decode the periodic table In which group is Fluorine located?
1.3.1 Using the Periodic Table: - Identify groups as vertical columns on the periodic table. - Know that main group elements in the same group have similar properties, the same number of valence electrons, and the same oxidation number. - Summarize that reactivity increases as you go down within a group for metals and decreases for nonmetals. - Identify periods as horizontal rows on the periodic table. - Metals/Nonmetals/Metalloids - Identify regions of the periodic table where metals, nonmetals, and metalloids are located. - Classify elements as metals/nonmetals/metalloids based on location. -Representative elements (main group) and transition elements -Identify representative (main group) elements as A groups or as groups 1, 2, 13-18. - Identify alkali metals, alkaline earth metals, halogens, and noble gases based on location on periodic table. Identify transition elements as B groups or as groups 3-12.
6 Day 9/21 1.3.2 Infer the physical properties (atomic radius, metallic and nonmetallic characteristics) of an element based on its position on the Periodic Table. 1.3.3 Infer the atomic size, reactivity, electronegativity, and ionization energy of an element from its position on the Periodic Table. 1.3.3 Using the Periodic Table, predict the number of electrons lost or gained and the oxidation number based on the electron configuration of an atom. Predict an ions charge based on its location on the periodic table What would the charge of sodium be? Day 7 9/22 Chm.1.3.2 1.3.2 Using the Periodic Table, - Define atomic radius and ionic radius. Day 8 9/23 Predict the size, reactivity, ionization - Know group and period general trends for atomic energy, or electronegativity of an element radius. based on its location on the P.T. - Apply trends to arrange elements in order of increasing or decreasing atomic radius. Explain the reasoning behind the trends. Predict the size, reactivity, ionization energy, or electronegativity of an element based on its location on the P.T. Which is more likely to gain electrons, B or K? - Compare cation and anion radius to neutral atom. - Determine the number of valence electrons from electron configurations. - Compare the metallic character of elements. - Use electron configuration and ion formation to justify metallic character. (Metals tend to lose electrons in order to achieve the stability of a filled
octet.) - Relate metallic character to ionization energy and electronegativity. Chm.1.3.3 1.3.3 Using the Periodic Table, - Predict the number of electrons lost or gained and the oxidation number based on the electron configuration of an atom. - Explain how the general size of an atom contributes to its reactivity sharing, gaining or losing electrons. - Compare reactivity of elements within groups and periods of the periodic table. - Define ionization energy and know group and period general trends for ionization energy. Explain the reasoning behind the trend. - Apply trends to arrange elements in order of increasing or decreasing ionization energy. - Define electronegativity and know group and period general trends for electronegativity. Explain the reasoning behind the trend. - Apply trends to arrange elements in order of increasing or decreasing electronegativity. Day 9 9/26 Cumulative Activity / Unit Wrap-Up / Lab Day 10 9/27 Unit Test
Unit 4: Bonding Day 1 9/28 1.2.1 Describe metallic bonds: metal ions plus sea of mobile electrons 1.2.2 Determine that a bond is predominantly ionic by the location of the atoms on the Periodic Table (metals combined with nonmetals) or when EN > 1.7. 1.2.2 Determine that a bond is predominantly covalent by the location of the atoms on the Periodic Table (nonmetals combined with nonmetals) or when EN < 1.7. Day 2 9/29 1.2.2 Predict chemical formulas of compounds using Lewis structures. Day 3 9/30 1.2.5 Apply Valence Shell Electron Pair Repulsion Theory (VSEPR) for these electron pair geometries and molecular geometries, and bond angles - Electron pair - Molecular (bond angle); Linear framework linear; Trigonal planar framework trigonal planar, bent; Tetrahedral framework tetrahedral, trigonal pyramidal, bent; Bond angles (include distorting effect of lone pair electrons no specific angles, conceptually only) Identify the type of bond between two atoms Draw ionic Lewis and covalent Lewis structures Predict the shape of a molecule based on VSEPR theory Which statement describes the compound formed between sodium and oxygen? a. It is NaO2, which is ionic. b. It is NaO2, which is covalent. c. It is Na2O, which is ionic. d. It is Na2O, which is covalent. Draw the Lewis structure of MgCl 2. The shape of CO 2 is best described as. 10/3 CMS ANNUAL LEAVE DAY - NO SCHOOL Day 4 10/4 Chm.1.2.3 1.2.3 Explain why intermolecular forces are weaker than ionic, covalent or metallic bonds Define intermolecular forces. Identify the intermolecular forces between 2 molecules At STP, fluorine is a gas and iodine is a solid. Why?
1.2.3 Explain why hydrogen bonds are stronger than dipole-dipole forces which are stronger than dispersion forces 1.2.3 Apply the relationship between bond energy and length of single, double, and triple bonds (conceptual, no numbers). 1.2.3 Describe intermolecular forces for molecular compounds. 1.2.3 H-bond as attraction between molecules when H is bonded to O, N, or F. Dipole-dipole attractions between polar molecules. 1.2.3 London dispersion forces (electrons of one molecule attracted to nucleus of another molecule) i.e. liquefied inert gases. a. Fluorine has lower average kinetic energy than iodine. b. Fluorine has higher average kinetic energy than iodine. c. Fluorine has weaker intermolecular forces of attraction than iodine. d. Fluorine has stronger intermolecular forces of attraction than iodine. 1.2.3 Relative strengths (H>dipole>London/van der Waals). Day 5 10/5 Chm.1.2.5 1.2.5 Explain how ionic bonding in compounds determines their characteristics: high MP, high BP, brittle, and high electrical conductivity either in molten state or in aqueous solution. 1.2.5 Explain how covalent bonding in compounds determines their characteristics: low MP, low BP, poor electrical conductivity, polar nature, etc. 1.2.5 Explain how metallic bonding determines the characteristics of metals: high MP, high BP, high Predict the properties of a substance. An unknown substance is tested in the laboratory. The physical test results are listed below. - Nonconductor of electricity - Insoluble in water - Soluble in oil - Low melting point
conductivity, malleability, ductility, and luster. 1.2.5 Describe bond polarity. Polar/nonpolar molecules (relate to symmetry) ; relate polarity to solubility like dissolves like 1.2.5 Describe macromolecules and network solids: water (ice), graphite/diamond, polymers (PVC, nylon), proteins (hair, DNA) intermolecular structure as a class of molecules with unique properties. Based on these results, what is the unknown substance? a. ionic and polar. b. ionic and nonpolar. c. covalent and polar. d. covalent and nonpolar. Day 6 10/6 Cumulative Activity / Unit Wrap-Up / Lab Day 7 10/7 Unit Test Unit 5: Chemical Nomenclature / The Mole 1.2.4 Write binary compounds of metals /nonmetals. What is the formula for: Day 1 10/10 1.2.4 Write ternary compounds (polyatomic ions) using the polyatomic ions on the reference table. Write formulas for ionic compounds (a) Magnesium Chloride? 1.2.4 Write, with charges, these polyatomic ions: nitrate, sulfate, carbonate, acetate, and ammonium. (b) Iron (III) Oxide? Day 2 10/11 1.2.4 Write binary compounds of metal/nonmetal. Write names for ionic compounds What is the name for: 1.2.4 Write ternary compounds (polyatomic ions) using the polyatomic (a) K 2 S?
ions on the reference table. 1.2.4 Write, with charges, these polyatomic ions: nitrate, sulfate, carbonate, acetate, and ammonium. (b) AuBr 2? 1.2.4 Write binary compounds of two nonmetals: use Greek prefixes (di-, tri-, tetra-, ). Day 3 10/12 - EARLY RELEASE DAY 1.2.4 Write binary compounds of metal /nonmetal*. 1.2.4 Write ternary compounds (polyatomic ions)* using the polyatomic ions on the reference table. Write names/formulas for covalent compounds Mixed Nomenclature Review (1) What is the name for SO 2? (2) What is the formula for NO? 1.2.4 Write, with charges, these polyatomic ions: nitrate, sulfate, carbonate, acetate, and ammonium. Day 4 10/13 2.2.4 Use mole ratios from the balanced equation to calculate the quantity of one substance in a reaction given the quantity of another substance in the reaction. (given moles, particles, mass, or volume and ending with moles, particles, mass, or volume of the desired substance) Day 5 10/14 Mole Math Day 2 Mole Math Mole Math / Nomenclature Quiz How many molecules are in 5.67 g of magnesium chloride? Day 6 10/17 2.2.5 Determine percentage composition Determine the % composition What is the % composition of
by mass of a given compound. of a substance glucose (C 6 H 12 O 6 )? 2.2.5 Perform calculations based on percent composition. 2.2.5 Calculate empirical formula from mass or percent using experimental data (1) Elemental analysis of a compound showed that it consisted of 81.82% carbon and 18.18% hydrogen by mass. What is the empirical formula of the compound? Day 7 10/18 2.2.5 Calculate molecular formula from empirical formula using molecular weight Determine the Empirical & Molecular Formula of a compound (2) A compound consisting of 56.38% phosphorus and 43.62% oxygen has a molecular mass of 220 g/mole. What is the molecular formula of this compound? Day 8 10/19 2.2.5 Determine the composition of hydrates using experimental data. Determine the composition of a hydrate A 10.10 g sample of barium chloride hydrate is heated in crucible. After all of the water is driven off, 8.50 g of the anhydrous barium chloride remains in the crucible. What is the formula of the hydrate? Day 10/20 Cumulative Activity / Unit Wrap-Up / Lab Day 10 10/21 MID-TERM REVIEW & EXAM
Day 1 10/24 Day 2 10/25 MIDTERM REVIEW / MIDTERM EXAM Day 3 10/26 Unit 6: CHEMICAL REACTIONS Consider this combustion reaction equation: C 4 H 10 + O 2 -- > CO 2 +H 2 O Day 1 10/27 2.2.3 Write and balance chemical equations predicting product(s) in a reaction using the reference tables Balance chemical equations When the equation is balanced, what will be the coefficient of O 2? a. 1 b. 7 c. 10 d. 13 Day 2 10/26 2.2.3 Write and balance chemical equations predicting product(s) in a reaction using the reference tables Identify the type of chemical reaction Write balanced chemical equations from a word problem (1)What type of chemical reaction is represented by the equation below? C 4 H 10 + O 2 -- > CO 2 +H 2 O
(2)Write a balanced equation to represent the following chemical reaction: Zinc and lead (II) nitrate react to form zinc nitrate and lead. Day 3 10/28 2.2.3 Use reference table rules to predict products for all types of reactions to show the conservation of mass 2.2.3 Write and balance ionic equations. 2.2.3 Recognize that hydrocarbons (C,H molecule) and other molecules containing C, H, and O 2 burn completely in oxygen to produce CO 2 and water vapor. Predict the products for synthesis reactions Predict the products for combustion reactions Write a balanced chemical equation demonstrating the reaction between potassium and sulfur. Write a balanced chemical equation representing the combustion of pentane (C 5 H 12 ) 10/31 TEACHER WORKDAY - All Hallow s Eve Day 4 11/1 2.2.3 Write and balance chemical equations predicting product(s) in a reaction using the reference tables 2.2.3 Use reference table rules to predict Predict the products for decomposition reaction Write a balanced chemical equation to represent the decomposition of magnesium hydroxide.
products for all types of reactions to show the conservation of mass. Day 5 11/2 2.2.3 Write and balance chemical equations predicting product(s) in a reaction using the reference tables. 2.2.3 Use activity series to predict whether a single replacement reaction will take place. Apply the activity series of metals to predict products for single replacement reactions Predict the products for and balance the following chemical reaction: Ag + KNO3 Day 6 11/3 2.2.3 Write and balance chemical equations predicting product(s) in a reaction using the reference tables. 2.2.3 Use the solubility rules to determine the precipitate in a double replacement reaction if a reaction occurs Apply solubility rules to predict products for double replacement reactions (1) Is NaOH soluble or insoluble? (2) Predict the product(s) of the following reaction: CaS (aq) + NaCl (aq) Day 7 11/4 2.2.3 Use the solubility rules to determine the precipitate in a double replacement reaction if a reaction occurs. 2.2.3 Write and balance net ionic equations for double replacement reactions. Write balanced net ionic equations for double replacement reactions. Write the net ionic equation for the reaction between lead (II) nitrate and sodium chloride. Day 8 11/7 2.2.1 Explain collision theory molecules must collide in order to react, and they must collide in the correct or appropriate Collision Theory Explain the energy content of a chemical reaction. Sketch a potential energy diagram, including reactants, products, activation energy,
orientation and enthalpy Day 9 11/9 with sufficient energy to equal or exceed the activation energy. 2.2.1 Interpret potential energy diagrams for endothermic and exothermic reactions including reactants, products, and activated complex with and without the presence of a catalyst. 11/8 WORKDAY (ELECTION DAY) Day 10 11/10 LAB 11/11 HOLIDAY Day 11 11/14 Cumulative Activity / Unit Wrap-Up Day 12 11/15 Unit Test Day 1 11/16 2.2.4 Interpret coefficients of a balanced equation as mole ratios 2.2.4 Given moles, particles or mass of one substance in the reactions calculate moles, particles, or mass of the desired substance Unit 7: Stoichiometry Identify mole ratios of products and reactants using coefficients in the balanced equation. Calculate mole to mole, mass to mass, particle to particle using the balanced equation Given the balanced chemical equation the reaction, P 4 + 5O 2 P 4 O 10 What mass of oxygen is needed to completely react with 7.75 g P 4? Day 2 11/17 2.2.4 Given moles, particles or mass of one substance in the reaction, calculate moles, particles, or mass of the desired substance Calculate between moles, mass and particles in a chemical reaction Given the balanced chemical equation the reaction, P 4 + 5O 2 P 4 O 10 How many molecules of oxygen
Day 3 11/18 Perform stoichiometry calculations to determine the limiting reactant in a chemical reaction Compare theoretical and experimental calculations of a quantity of product to determine the percent yield of the product formed in the lab experiment Day 4 11/21 To prepare and determine the percent yield of sodium chloride. 2.2.4 To gain an understanding of mass relationships in chemical reactions. Day 5 11/22 Unit Quiz Define limiting reactant and recognize its effect on production of products in a chemical reaction. Use quantities of reactants and mole ratios to determine the limiting reactant in a chemical reaction. Lab React NaHCO 3 with HCl to produce and measure the percent of the NaCl product is needed to completely react with 7.75 g P 4? Given the balanced chemical equation the reaction, P 4 + 5O 2 P 4 O 10 What is the limiting reactant when 7.75 g P 4 reacts with 7.75 g of O 2 to produce P 4 O 10? 40.0 grams of hydrogen reacts with an excess of Oxygen to produce water. At the end of the lab, the water collected weighed 340. grams. What is the theoretical yield of water? Calculate the percent of water. Unit 8: Heat and States of Matter Day 1 11/28 2.1.1 Explain physical equilibrium: liquid water-water vapor 2.1.1 Explain how the energy (kinetic and potential) of the particles of a substance changes when heated, cooled, or changing phase 2.1.1 Contrast heat and temperature, including temperature as a measure of States of Matter Properties and Heating and Cooling Curve What causes the process of perspiration to be cooling for human skin? a. It involves condensation and is exothermic. b. It involves evaporation and is exothermic. c. It involves condensation and is
average kinetic energy 2.1.2 Interpret heating and cooling curves endothermic. d. It involves evaporation and is endothermic. 2.1.2 Explain phase change calculations in terms of heat absorbed or released (endothermic vs. exothermic processes) Day 2 11/29 2.1.2 Define and use the terms and/or symbols for specific heat capacity 2.1.2 Complete calculations of q = mc p T using heating/cooling curve data 2.1.4 Recognize that, for a closed system, energy is neither lost nor gained only transferred between components of the system 2.1.4 Complete calculations of: q lost = (- q gained ) Day 3 11/30 2.1.2 Define and use the terms and/or symbols for: heat of fusion, heat of vaporization Q Calculations for within a phase (metals) Q Calculations for phase changes (water) An 8.80 g sample of metal is heated to 92.0 C and then added to 15.0 g of water at 20.0 C in an insulated calorimeter. At thermal equilibrium the temperature of the system was measured as 25.0 C. What is the identity of the metal? How much energy is required to turn 15.5 grams of ice at -12 C to steam at 124 C? 2.1.2 Complete calculations of q = mh f and q = mh v using heating/cooling curve data 2.1.4 Complete calculations of: q = mc p T, q = mh f and q = mh v in water, including phase changes, using laboratory data
Day 4 12/1 2.1.1 Identify pressure as well as temperature as a determining factor for phase of matter Phase Change Diagram 2.1.2 Interpret phase diagrams for water and carbon dioxide 2.1.3 Draw phase diagrams of water and carbon dioxide (shows how sublimation occurs) 2.1.3 Identify regions, phases and phase changes using a phase diagram 2.1.3 Use phase diagrams to determine information such as (1) phase at a given temperature and pressure, (2) boiling point or melting point at a given pressure, (3) triple point of a material Day 5 12/2 2.1.1 Appropriately use the units Celsius and Kelvin Heat Quiz and Temperature and Pressure Conversions What is the sublimation point of this substance at 1 atm? 2.43 atm = mmhg 24.6 C = K Unit 8: Gas Laws Day 1 12/5 2.1.5 Combined gas law and applications holding one variable constant (students should be able to derive and use these gas laws, but are not necessarily expected to memorize their names) Combined Gas Law When volume increases and temperature is held constant, what happens to pressure? When pressure is held constant and temperature is increased, what happens to volume?
Day 2 12/6 2.1.5 Ideal gas equation Ideal Gas Law How many grams of oxygen gas will take up 1.75 liters at 2.30 atm and 315 K? Day 3 12/7 2.2.3 Use mole ratios from the balanced equation to calculate the quantity of one substance in a reaction given the quantity of another substance in the reaction 2.1.5 One mole of any gas at STP = 22.4 L Day 4 12/8 2.1.5 Dalton s law (P t = P 1 + P 2 + P 3 ) 2.1.5 Vapor pressure of water as a function of temperature (conceptually) Stoichiometry of Gases at STP Partial Pressure 12.0 grams of oxygen are added to hydrogen and water is produced, at STP. How many liters of hydrogen should be added to completely use up the oxygen? If a 12 L container of oxygen gas, at 1.3 atm, and a 12 L container of nitrogen gas, at 755 mmhg, are combined in a 12 L container, what is the total pressure? Day 5 12/9 2.1.5 Identify characteristics of ideal gases Stoichiometry not at STP Kinetic Molecular Theory (again) What causes an inflated balloon to shrink when it is cooled? a. because cooling the balloon causes gas to escape from the ball b. because cooling the balloon causes the gas molecules to collide more frequently c. because cooling the balloon causes gas molecules to become smaller d. because cooling the balloon
causes the average kinetic energy of the gas molecules to decrease Day 6 12/12 Cumulative Activity / Unit Wrap-Up / Lab Day 7 12/13 Unit Test Units 9 and 10: Solutions and Acids and Bases Day 1 12/14 3.2.4 Identify types of solutions (solid, liquid, gaseous, aqueous) 3.2.4 Define solutions as homogeneous mixtures in a single phase 3.2.4 Distinguish between electrolytic and non electrolytic solutions 3.2.5 Use graph of solubility vs. temperature to identify a substance based on solubility at a particular temperature 3.2.5 Use graph to relate the degree of saturation of solutions to temperature 3.2.6 Develop a conceptual model for the solution process with a cause and effect relationship involving forces of attraction between solute and solvent particles. A material is insoluble due to a lack of attraction between particles Solutions Intro/Definitions and Solubility Curves When considering the energetics of the solution process, which process is always exothermic? a. Solute particles separate from one another. b. Solvent particles separate from one another. c. Solute and solvent particles form attractions for one another. d. Solution formation as a whole is always endothermic.
3.2.6 Describe the energetic of the solution process as it occurs and the overall process as exothermic or endothermic 3.2.6 Explain solubility in terms of the nature of solute-solvent attraction, temperature and pressure (for gases) 2.1.5 Apply general gas solubility characteristics Day 2 12/15 3.2.3 Compute concentration of solutions in moles per liter 3.2.3 Calculate molarity given mass of solute and volume of solution 3.2.3 Calculate mass of solute needed to create a solution of a given molarity and volume Day 3 12/16 3.2.3 Solve dilution problems 3.2.4 Summarize colligative properties (vapor pressure reduction, boiling point elevation, freezing point depression, and osmotic pressure) 2.1.1 Explain physical equilibrium: liquid water-water vapor. Vapor pressure depends on temperature and concentration of particles in solution Molarity Dilutions and Colligative Properties What s the molarity of a 55.0 ml solution of magnesium chloride if 2.55 grams of the salt were used? Heat is added to a solution to a. increase the solubility of a solid solute. b. increase the solubility of a gas solute. c. increase the miscibility of the solution d. increase the degree of saturation of the solution
Day 4 12/19 UNIT 9 - SOLUTIONS - QUEST Day 5 12/20 3.2.1 Distinguish between acids and bases based on formula and chemical properties 3.2.1 Differentiate between concentration and strength (No calculation) 3.2.1 Use ph scale to identify acids and bases 3.2.2 Distinguish properties of acids and bases related to taste, touch, reaction with metals, electrical conductivity, and identification with indicators such as litmus paper and phenolphthalein Given the information below, which substance is an acid? Substance W Does not taste sour, does not feel slippery, turns litmus paper blue Substance X Tastes bitter, feels slippery, turns litmus paper blue Substance Y Tastes bitter, feels slippery, turns litmus paper blue Substance Z Does not taste bitter, tastes sour, turns litmus paper red a. Substance W b. Substance X c. Substance Y d. Substance Z Day 5 12/21 GOLD PENNY LAB
12/22-1/3 WINTER BREAK Day 6 1/4 3.2.1 Relate the color of indicator to ph using ph ranges provided in a table 3.2.1 Compute ph, poh, [H+], and [OH-] 3.2.1 Interpret ph scale in terms of the exponential nature of ph values in terms of concentrations ph and poh Math ph and poh Math Based on hydroxide ion concentration, which unknown substance would be an acid? a. Substance A, [OH-] = 1.0 x 10-2 M b. Substance B, [OH-] = 1.0 x 10-4 M c. Substance C, [OH-] = 1.0 x 10-6 M d. Substance D, [OH-] = 1.0 x 10-8 M Day 7 1/5 3.2.3 Perform 1:1 titration calculations 3.2.3 Determine the concentration of an acid or base using titration. Interpret titration curve for strong acid/ strong base 2.2.3 Identify acid-base neutralization as double replacement Titrations What volume of 0.200M HCl will neutralize 10.0mL of 0.400M KOH? a. 40.0mL b. 20.0mL c. 8.00mL d. 5.00m Day 8 1/6 CUMULATIVE ACTIVITY
Day 9 1/9 ACID/BASE QUEST Unit 11: Equilibrium Day 1 1/10 3.2.1 Define chemical equilibrium for reversible reactions 3.2.1 Distinguish between equal rates and equal concentrations 3.1.3 Determine the effects of stresses on systems at equilibrium (adding/removing a reactant or product; adding/removing heat/ increasing/decreasing pressure) 3.1.3 Relate the shift that occurs in terms of the order/disorder of the system Le Chatelier s Principle For the reaction 2SO2 (g) + O2 (g) 2SO3 (g) + heat Which action will increase the concentration of SO3? a. removing SO2 b. increasing the temperature c. increasing the pressure d. adding a catalyst Day 2 1/11 3.2.1 Explain equilibrium expressions for a given reaction 3.2.1 Evaluate equilibrium constants as a measure of the extent that the reaction proceeds to completion K eq Equations Write the K eq equation for the following reaction: 2SO2 (g) + O2 (g) 2SO3 (g) + heat 1/12 Early Release Day 1/16 HOLIDAY Exam Review
1/17-1/23 FINAL EXAMS (estimated dates)