pacing guide 2018 - 2019 units 1 and 2: atomic structure...

PACING GUIDE 2018 - 2019 Units 1 and 2: Atomic Structure and Nuclear Chemistry

Day 1 8/27 ● What is Chemistry? ● Syllabus/Safety/Measurement

2.2.2 Analyze the evidence of chemical change

Identify Lab Safety and Equipment Use Lab Instrument for

Measurements Observe a chemical reaction and identify indicators of a chemical

reaction

Day 2 8/28 Significant digits, scientific notation, metric conversions, and measurements

Sig figs and measurement lab

Day 3 8/29 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 8/30 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

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 5 8/31 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-

9/3 LABOR DAY HOLIDAY

Day 6 9/4 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 QUIZ #1 - Matter, Sig. Figs.

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 7 9/5 1.1.4 Use the symbols for and distinguish between alpha ( 42He), 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 9/6 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 9/7 CANDIUM Lab / QUIZ #2 - Atom Math UNITS 1 & 2 TEST DATE: 9/11/18

Day 10 9/10 Units 1 & 2 Test Day Wrap-up Study Guide (Required) / Tutoring

Day 11 9/11 Units 1 & 2 Test

Unit 3: Quantum Theory/Periodic Trends Day 1 9/12

EARLY

RELEASE

DAY

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. 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).

Identify the parts of a wave Rank waves via energy Interpret 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 nm b. 434 nm c. 486 nm d. 656 nm

Day 2 9/13 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 3p8 b. 1s2 2s2 2p6 3s2 3p6 4s2 c. 1s2 2s2 2p6 3s2 3p6 3d2

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

Day 3 9/14 Write the orbital filling diagram for a given element

Day 4 9/17 Write the noble gas notation of a given element Draw the Lewis dot diagram of a given element

Day 5 9/18 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.

Draw the Lewis dot diagram of a given element Predict an ions charge based on its location on the periodic table

What would be the charge of a sodium ion?

UNIT 3 QUIZ #1

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.

9/19 TEACHER WORKDAY

Day 6 9/20 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 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.

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.

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?

- 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 7 9/21 1.3.1 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

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.

UNIT 3 QUIZ #2

Day 8 9/24 Cumulative Activity / Unit Wrap-Up / Lab

Day 9 9/25 Unit 3 Test

Unit 4: Bonding

Day 1

9/26 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.

Identify the type of bond between two atoms

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.

Day 2 9/28 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 10/1 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 10/2 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

Polarity and Bond Properties 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

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)

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 5 10/3 1.2.3 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 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. 1.2.3 Relative strength(H>dipole>London)

Define intermolecular forces. Identify the intermolecular forces

between 2 molecules

Unit 4 QUIZ

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 of attraction than iodine.

d. Fluorine has stronger intermolecular forces of

attraction than iodine.

Day6 10/4 Cumulative Activity / Unit Wrap-Up / Lab


Unit 5: Chemical Nomenclature / The Mole

Day 1 10/8 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 10/9

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,


Write names for ionic compounds

What is the name for:

(a) K2S ?

(b) AuBr2 ?

Day 3

PSAT

TEST

DATE

10/10

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,


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 10/11 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?

Day 5 10/12

Mole Math – Day 2

UNIT 5 QUIZ - NOMENCLATURE

Day 6

10/15

2.2.5 Determine percentage composition by mass of a given compound.

2.2.5 Perform calculations based on percent composition.

Determine the % composition

of a substance

What is the % composition of

glucose (C6H12O6)?

Day 7 10/16

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 the Empirical & 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 the molecular formula of this compound?

Day 8

10/17 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?

10/18

Day 9 10/19 MOLE AIRLINE/QUIZ

Day 10 10/22 UNIT 5 CUMULATIVE REVIEW

Day 11 10/23 UNIT 5 TEST

10/24 Parts of Rxn/Balancing

10/25 Types of Rxns/Midterm Review

10/26 Midterms/Midterm Review

10/29 Midterms/Types of Rxn

10/30 Midterm Make-Up/Predicting Products

10/31 Teacher Workday

Unit 6: CHEMICAL REACTIONS

Day 1 10/24 2.2.3 Write and balance chemical

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

Balance chemical equations

Consider this combustion 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 10/25

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 10/30

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 O2 burn completely in oxygen to

produce CO2 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 (C5H12)

10/31 TEACHER WORKDAY

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 products for all types of reactions to show

the conservation of mass.

Predict the products for decomposition reaction

Write a balanced chemical equation to represent the

decomposition of magnesium hydroxide.

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 QUIZ - BAL EQUATIONS

Predict the products for and balance the following chemical

reaction:

Ag + KNO3→

Day 6 11/5

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 11/6

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 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 content of a chemical reaction.

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

enthalpy

Day 9 11/8 Chemical Reactions Lab (Lime Water Demo)

Day 10 11/9 Cumulative Activity / Unit Wrap-Up

11/12 Veteran’s Day

Day 11 11/13 UNIT 6 TEST

Unit 7: Stoichiometry Day 1 11/14 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

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, P4+ 5O2 →P4O10 What mass of oxygen is needed to completely react with 7.75 g P4 ?

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 2 11/15 Perform stoichiometry calculations to determine the limiting reactant in a

chemical reaction

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 3 11/16 Compare theoretical and experimental calculations of a quantity of product to

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

% Yield % Yield Lab Welch - Lab

Wilson - Review

Day 4 11/19 UNIT 7 REVIEW Welch - Review Wilson - Lab

Day 5 11/20 UNIT 7 QUEST 11/21 -

11/23 THANKSGIVING HOLIDAYS

Unit 8: Heat and States of Matter

Day 1 11/26 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

average kinetic energy 2.1.2 Interpret heating and cooling curves 2.1.2 Explain phase change calculations in

terms of heat absorbed or released (endothermic vs. exothermic processes)

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 endothermic. d. It involves evaporation and is endothermic.

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 2 11/27 2.1.2 Define and use the terms and/or symbols for: heat of fusion, heat of

vaporization 2.1.2 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 11/28 2.1.1 Identify pressure as well as temperature as a determining factor for

phase of matter 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

Phase Change Diagram

What is the sublimation point of this

substance at 1 atm?

Day 5 11/29 2.1.1 Appropriately use the units Celsius and Kelvin

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)

Temperature and Pressure Conversions

Combined Gas Law

2.43 atm = ________ mmHg 24.6 °C = _______ K

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?

12/3 Unit 8 Heat Quest Review Start Gas Laws

Unit 8 Quest - HEAT

Unit 9: Gas Laws

Day 2 12/5 Early

Release

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 12/6 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

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

12/7 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 12/10 Dry Ice Lab Cumulative Activity / Unit Wrap-Up


Units 10: Solutions

Day 1 12/12 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

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

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.

Day 2 12/13 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

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

Day 3 12/14 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

Dilutions and Colligative Properties 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 12/17 UNIT 10 Solutions REVIEW

Day 5 12/18 UNIT 10 Solutions QUEST 12/19 -

1/2 WINTER BREAK

Units 11: Acids and Bases

Day 1 1/3/19 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, tough, reaction

with metals, electrical conductivity, and identification with indicators such as

litmus paper and phenolphthalein

Acid/Base Properties Acid/Base Nomenclature

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 2 1/4/19 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 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 3 1/7/19 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

Unit 12: Equilibrium

Day 1 1/8/18 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 1/9/18 Early

Release

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

Keq Equations Write the Keq equation for the following reaction:

2SO2 (g) + O2 (g) ⇄ 2SO3 (g) + heat

1/10-1/11 Exam Review

1/12-1/19 FINAL EXAMS

pacing guide 2018 - 2019 units 1 and 2: atomic structure...

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