Units 1 and 2: Atomic Structure and Nuclear ChemistryDay 1 8/24 What is Chemistry?

Syllabus/Safety/Measurement2.2.2 Analyze the evidence of chemical

change

Identify Lab Safety and EquipmentUse Lab Instrument for

MeasurementsObserve a chemical reaction and identify indicators of a chemical

reaction

Day 2 8/25 2.2.2 Classification of matter2.2.2 Analyze the evidence of chemical

change2.2.2 Analyze the evidence of a physical

change

Classify matterContrast and identify chemical and

physical changes

Day 3 8/26 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-2351.1.1 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

List numbers of protons, electrons, neutrons in different atoms

Use Isotope notations

Which atomic symbol represents an isotope of sulfur with 17 neutrons?a. 17

16 X b. 3316 X

c. 1732 X d.49

32X

Day 4 8/27 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, silvery-white metal that has two common isotopes, 85Rb and 87Rb. If the abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium?

Day 5 8/28 1.1.1 Identify ions using mass number and atomic number and relate to number of

protons, neutrons and electrons.

Ions and Element Math Determine the number of p, e, and n in 35 Cl 1-

Day 6 8/31 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 7 9/1 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 8 9/2 Cumulative Activity / Unit Wrap-Up / Lab Day 9 9/3 Unit Test

Unit 3: Quantum Theory/Periodic TendsDay 1 9/4 1.1.2 Analyze diagrams related to the Bohr

model of the hydrogen atom in terms of allowed, discrete 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.

Identify the parts of a wave Rank waves via energyInterpret the Bohr model to identify the type of light

An electron in an atom of hydrogen goes from energy

level 6 to energy level 2. What is the wavelength of the

electromagnetic radiation emitted?a. 410 nmb. 434 nmc. 486 nmd. 656 nm

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).

Day 2 9/8 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.1.3.2 Identify s, p, d, and f blocks on Periodic

Table.1.3.2 Identify an element based on its

electron configuration. (Students should be able to identify elements which follow the

general rules, not necessarily those which are exceptions.)

Write the electron configuration for a given element

Which is the electronic configuration of calcium?

a. 1s2 2s2 2p6 3s2 3p8b. 1s2 2s2 2p6 3s2 3p6 4s2c. 1s2 2s2 2p6 3s2 3p6 3d2

d. 1s2 2s2 2p8 3s2 3p6

Day 3 9/9 Write the orbital filing diagram for a given element

Day 4 9/10 Write the noble gas notation of a given elementDraw the Lewis dot diagram of a given element

Day 5 9/11 Chm.1.3.11.3.1 Classify the components of a periodic

table (period, group, metal, metalloid, nonmetal, transition).

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

Decode the periodic table What group is Fluorine in?

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 9/14 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 9/15Chm.1.3.2

1.3.2 Using the Periodic Table,- Define atomic radius and ionic radius.

- Know group and period general trends for atomic radius.

- Apply trends to arrange elements in order of increasing or decreasing atomic radius.

Explain the reasoning behind the trends.- Compare cation and anion radius to neutral

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?

Day 8 9/16 Predict the size, reactivity, ionization energy, or electronegativity of an element based on its location on the P.T.

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.1.3.31.3.3 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 9/17 Cumulative Activity / Unit Wrap-Up / Lab

Day 10 9/18 Unit Test

Unit 4: BondingDay 1 9/21 1.2.1 Describe metallic bonds: “metal ions Identify the type of bond between Which statement describes the

plus ‘sea’ of mobile electrons”1.2.2 Determine that a bond is

predominately 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 predominately covalent by the location of

the atoms on the Periodic Table (nonmetals combined with nonmetals) or

when EN < 1.7.

two atoms 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.

Day 2 9/22 1.2.2 Predict chemical formulas of compounds using Lewis structures.

Draw ionic Lewis and covalent Lewis structures

Draw the Lewis structure of MgCl2.

Day 3 9/24 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)

Predict the shape of a molecule based on VSEPR theory

The shape of CO2 is best described as ___________.

Day 4 9/25 Chm.1.2.31.2.3 Explain why intermolecular forces are

weaker than ionic, covalent or metallic bonds

1.2.3 Explain why hydrogen bonds are stronger than dipole-dipole forces which

are stronger than dispersion forces1.2.3 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

Define intermolecular forces. Identify the intermolecular forces

between 2 molecules

At STP, fluorine is a gas and iodine is a solid. Why?

a. Fluorine has lower average kinetic energy than iodine.

b. Fluorine has higher average kinetic energy than iodine.

c. Fluorine has weaker intermolecular forces ofattraction than iodine.

d. Fluorine has stronger intermolecular forces of

attraction

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.1.2.3 Relative strengths

(H>dipole>London/van der Waals).

than iodine.

Day 5 9/28 Chm.1.2.51.2.5 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 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.

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

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 9/29 Cumulative Activity / Unit Wrap-Up / Lab Day 7 9/30 Unit Test

Unit 5: Chemical Nomenclature / The Mole

Day 1 10/1

1.2.4 Write binary compounds of metals /nonmetals.

1.2.4 Write ternary compounds (polyatomic ions) using the polyatomic

ions on the reference table.1.2.4 Write, with charges, these polyatomic ions: nitrate, sulfate,

carbonate, acetate, and ammonium.

Write formulas for ionic compounds

What is the formula for:

(a) Magnesium Chloride?

(b) Iron (III) Oxide?

Day 2 10/2

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.1.2.4 Write, with charges, these polyatomic ions: nitrate, sulfate,

carbonate, acetate, and ammonium.

Write names for ionic compounds

What is the name for:

(a) K2S ?

(b) AuBr2 ?

Day 3 10/5

1.2.4 Write binary compounds of two nonmetals: use Greek prefixes (di-, tri-,

tetra-, …).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.

1.2.4 Write, with charges, these polyatomic ions: nitrate, sulfate,

carbonate, acetate, and ammonium.

Write names/formulas for covalent compounds

Mixed Nomenclature Review

(1) What is the name for SO2?

(2) What is the formula for NO ?

Day 4 10/6 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)

Mole Math

How many molecules are in 5.67 g of magnesium chloride?Mole Math / Nomenclature Quiz

10/7 PLAN TEST / Early Release Day

Day 5 10/8 Mole Math – Day 2

Day 6 10/9

2.2.5 Determine percentage composition by mass of a given compound.

2.2.5 Perform calculations based on percent composition.

Determine the % compositionof a substance

What is the % composition of glucose (C6H12O6)?

Day 7 10/12

2.2.5 Calculate empirical formula from mass or percent using experimental data

2.2.5 Calculate molecular formula from empirical formula using molecular weight

Determine theEmpirical & Molecular Formula of a

compound

(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?

(2) A compound consisting of 56.38% phosphorus and 43.62%oxygen has a molecular mass of 220 g/mole. What is themolecular formula of this compound?

Day 8 10/13 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 incrucible. After all of the water is driven off, 8.50 g of theanhydrous barium chloride remains in the crucible. What isthe formula of the hydrate?

10/14 PSAT TESTDay 9 10/15 Cumulative Activity / Unit Wrap-Up / Lab

Day 10 10/16 Unit TestMID-TERM REVIEW & EXAM

Day 1MID-TERM REVIEW / MID-TERM EXAMDay 2

Day 3

Unit 6: CHEMICAL REACTIONSDay 1 10/26 2.2.3 Write and balance chemical Balance chemical equations Consider this combustion

equations predicting product(s) in a reaction using the reference tables

reaction equation:

C4H10 + O2 -- > CO2 +H2O

When the equation is balanced, what will be the coefficient of O2?

a. 1 b. 7 c. 10 d. 13

Day 2 10/27

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?

C4H10 + O2 -- > CO2 +H2O

(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 mass2.2.3 Write and balance ionic equations.2.2.3 Recognize that hydrocarbons (C,H

molecule) and other molecules containing C, H, and O2 burn completely in oxygen to

produce CO2and water vapor.

Predict the products forsynthesis reactions

Predict the products forcombustion reactions

Write a balanced chemical equation demonstrating the

reaction between potassium and sulfur.

Write a balanced chemical equation representing the

combustion of pentane (C5H12)

Day 4 10/29 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 products for all types of reactions to show

Predict the products fordecomposition reaction

Write a balanced chemical equation to represent the

decomposition of magnesium hydroxide.

the conservation of mass.

Day 5 10/30

2.2.3Write 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 11/2

2.2.3Write 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 11/3

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 11/4

2.2.1 Explain collision theory – molecules must collide in order to react, and they

must collide in the correct or appropriate orientation and

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.

Collision Theory

Explain the energy contentof a chemical reaction.

Sketch a potential energy diagram, including reactants, products, activation energy,

enthalpy

Day 9 11/5 LabDay 10 11/6 Cumulative Activity / Unit Wrap-Up Day 11 11/9 Unit Test

Unit 7: Stoichiometry Day 1 11/10 2.2.4 Interpret coefficients of a balanced Identify mole ratios of products and Given the balanced chemical

equation as mole ratios2.2.4 Given moles, particles or mass of

one substance in the reactions calculate moles, particles, or mass of the desired

substance

reactants using coefficients in the balanced equation.

Calculate mole to mole, mass to mass, particle to particle using the balanced

equation

equation the reaction,P4+ 5O2 →P4O10

What mass of oxygen is needed to completely react with 7.75 g P4 ?

Day 2 11/12 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,P4+ 5O2 →P4O10

How many molecules of oxygen is needed to completely react with 7.75 g P4 ?

Day 3 11/13 Perform stoichiometry calculations to determine the limiting reactant in a

chemical reactionCompare theoretical and experimental calculations of a quantity of product to

determine the percent yield of the product formed in the lab experiment

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.

Given the balanced chemical equation the reaction,P4+ 5O2 →P4O10

What is the limiting reactant when 7.75 g P4 reacts with 7.75 g of O2 to produce P4O10?

Day 4 11/16 To prepare and determine the percent yield of sodium chloride.

2.2.4 To gain an understanding of mass relationships in chemical reactions.

Lab – React NaHCO3 with HCl to produce and measure the percent of

the NaCl product

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.

Day 5 11/17 Unit Quiz

Unit 8: Heat and States of MatterDay 1 11/18 2.1.1 Explain physical equilibrium: liquid

water-water vapor2.1.1 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

curves2.1.2 Explain phase change calculations in terms of heat absorbed or released

(endothermic vs. exothermic processes)

endothermic.d. It involves evaporation and is endothermic.

Day 2 11/19 2.1.2 Define and use the terms and/or symbols for specific heat capacity

2.1.2 Complete calculations of q = mCp∆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: qlost = (-qgained)

Q Calculations for within a phase (metals)

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?

Day 3 11/20 2.1.2 Define and use the terms and/or symbols for: heat of fusion, heat of

vaporization2.1.2 Complete calculations of q = mHf

and q = mHv using heating/cooling curve data

2.1.4 Complete calculations of: q = mCp∆T, q = mHf and q = mHv in water,

including phase changes, using laboratory data

Q Calculations for phase changes (water)

How much energy is required to turn 15.5 grams of ice at -12 °C to

steam at 124 °C?

Day 4 11/23 2.1.1 Identify pressure as well as temperature as a determining factor for

phase of matter2.1.2 Interpret phase diagrams for water

and carbon dioxide2.1.3 Draw phase diagrams of water and carbon dioxide (shows how sublimation

occurs)2.1.3 Identify regions, phases and phase

changes using a phase diagram2.1.3 Use phase diagrams to determine

Phase Change Diagram

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

What is the sublimation point of this substance at 1 atm?

Day 5 11/24 2.1.1 Appropriately use the units Celsius and Kelvin

Heat Quiz and Temperature and Pressure Conversions

2.43 atm = ________ mmHg24.6 °C = _______ K

Unit 8: Gas LawsDay 1 11/30 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 12/1 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 12/2 2.2.3 Use mole rations from the balanced equation to calculate the quantity of one substance in a reaction given the quantity

of another substance in the reaction2.1.5 One mole of any gas at STP = 22.4 L

Stoichiometry of Gases at STP 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?

Day 4 12/3 2.1.5 Dalton’s law (Pt = P1 + P2 + P3 …)2.1.5 Vapor pressure of water as a function

of temperature (conceptually)

Partial Pressure 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 12/4 2.1.5 Identify characteristics of ideal gases Stoichiometry not at STPKinetic 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 ballb. because cooling the balloon causes the gas molecules to collide more frequently

c. because cooling the balloon causes gas molecules to become smallerd. because cooling the balloon causes the average kinetic energy of the gas molecules to decrease

Day 6 12/7 Cumulative Activity / Unit Wrap-Up / Lab Day 7 12/8 Unit Test

Units 9 and 10: Solutions and Acids and BasesDay 1 12/9 3.2.4 Identify types of solutions (solid,

liquid, gaseous, aqueous)3.2.4 Define solutions as homogeneous

mixtures in a single phase3.2.4 Distinguish between elecrolytic and

nonelectrolytic solutions3.2.5 Use graph of solubility vs.

temperature to identify a substance based on solubility at a particular temperature3.2.5 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

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

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.

characteristicsDay 2 12/10 3.2.3 Compute concentration of solutions

in moles per liter3.2.3 Calculate molarity given mass of

solute and volume of solution3.2.3 Calculate mass of solute needed to create a solution of a given molarity and

volume

Molarity What’s the molartity of a 55.0 mL solution of magnesium chloride if 2.55 grams of the salt were used?

Day 3 12/11 3.2.3 Solve dilution problems3.2.4 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

Dilutions and Colligative Properties Heat is added to a solution toa. increase the solubility of a solid solute.b. increase the solubility of a gas solute.c. increase the miscibility of the solutiond. increase the degree of saturation of the solution

Day 4 12/14 3.2.1 Distinguish between acids and bases based on formula and chemical properties3.2.1 Differentiate between concentration

and strength (No calculation)3.2.1 Use pH scale to identify acids and

bases3.2.2 Distinguish properties of acids and

bases related to taste, tough, reaction with metals, electrical conductivity, and

identification with indicators such as litmus paper and phenolphthalein

Solutions Quiz and Acid/Base Properties

Given the information below, which substance is an acid?

Substance W – Does not taste sour, does not feel slippery, turns litmus paper blueSubstance X – Tastes bitter, feels slippery, turns litmus paper blueSubstance Y – Tastes bitter, feels slippery, turns litmus paper blueSubstance Z – Does not taste bitter, tastes sour, turns litmus paper red

a. Substance Wb. Substance Xc. Substance Yd. Substance Z

Day 4 12/15 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 Based on hydroxide ion concentration, which unknown substance would be an acid?a. Substance A, [OH-] = 1.0 x 10-2 Mb. Substance B, [OH-] = 1.0 x 10-4 Mc. Substance C, [OH-] = 1.0 x 10-6 Md. Substance D, [OH-] = 1.0 x 10-8 M

Day 6 12/16 pH and pOH Math

Day 7 12/17 3.2.3 Perform 1:1 titration calculations3.2.3 Determine the concentration of an

acid or base using titration. Interpret titration curve for strong acid/ strong base2.2.3 Identify acid-base neutralization as

double replacement

Titrations What volume of 0.200M HCl will neutralize 10.0mL of 0.400M KOH?a. 40.0mLb. 20.0mLc. 8.00mLd. 5.00m

Day 8 12/18 Unit Test

Unit 11: EquilibriumDay 1 1/4 3.2.1 Define chemical equilibrium for

reversible reactions3.2.1 Distinguish between equal rates and

equal concentrations3.1.3 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) + heatWhich action will increase the concentration of SO3?a. removing SO2b. increasing the temperaturec. increasing the pressured. adding a catalyst

Day 2 1/5 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

Keq Equations Write the Keq equation for the following reaction:

2SO2 (g) + O2 (g) ⇄ 2SO3 (g) + heat

proceeds to completionDay 3 1/6 – 1/8 Quiz/ Exam Review

FINAL EXAMS