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Chemistry STAAR Review Ladder to Success

Rung 1

Reporting Category 1: Matter & Periodic Table

Differentiate between physical and chemical changes and properties

Physical and Chemical Properties

Physical and Chemical Changes

How can you recognize a chemical change?

Property Description Example

Change Description Example

Identify extensive and intensive properties

Types of Properties Physical and chemical properties of matter can be classified as either intensive or extensive.

Extensive Properties Intensive Properties

Compare solids, liquids, and gases in terms of compressibility, structure, shape and volume

States of Matter and Properties

Property Solid Liquid Gas

Compressible

Shape

Volume

Structure

Classify matter as pure substances or mixtures through investigation of their properties

Matter

Matter: anything that takes up space

and has mass

Pure Substance:

Element: Compound:

Mixture:

Homogeneous:

Solution:

Heterogeneous:

Chemistry STAAR Review

Ladder to Success

Rung 1: Homework

Identify the following as physical or

chemical properties. 1. The color of the house is red.

2. Oxygen is a gas.

3. A flagpole is 25 feet tall.

4. A ruby is red.

Identify the following as physical or chemical

changes:

5. Salt dissolving in water

6. Magnesium reacting with hydrochloric acid

7. Milk turning sour

8. Dry ice changing to a gas

Identify the following as element, compound or mixture.

9. An unknown, clear liquid is given to you in a beaker. You transfer the liquid from the beaker to a

clean, empty test-tube and begin to heat it. After a while, you see vapors (which on further

analysis, you discover are vapors of water) rising from the test-tube, and pretty soon, all that's left

are a few crystals of salt stuck to the edges! How would you classify this liquid?

10. You have won the world's biggest lottery, for which you are given a huge block of pure, metallic

gold. How would you classify your prize?

11. A dish is given to you, which contains a blackish-yellow powder. When you move a magnet over

it, you are amazed to see black particles (which you find out are iron) fly upwards and get stuck

to the magnet. All that's left in the dish is a yellow powder, which you discover to be sulfur. How

would you classify the initial blackish-yellow powder?

12. A substance is analyzed in a laboratory, and when viewed under an electron microscope, it is

revealed that it contains only one kind of atom. How would you classify the substance?

13. A magnesium ribbon is burnt in the air to form the grayish oxide of magnesium - magnesium

oxide (MgO). How would you classify this oxide?

Answer the following:

14. What is the density at 20°C of 12.0 milliliters of a liquid that has a mass of 4.05 grams?

15. A sample of an element has a volume of 78.0 mL and a density of 1.85 g/mL. What is the mass in

grams of the sample?


Rung 2

Reporting Category 2: Atomic Structure and Nuclear Chemistry

Understand the experimental

design and conclusions used

in the development of modern atomic

theory. Including Dalton’s

postulates, Thomson’s

discovery of the electron

properties, Rutherford’s

nuclear atom, and Bohr’s nuclear

atom.

Scientist Contribution to Modern Atomic Theory

Dalton

Thomson

Rutherford

Bohr

Modern Atomic

Theory

Understand the electromagnetic

spectrum and the mathematical relationships

between energy, frequency, and wavelength of

light.

A wave can be described by its frequency, wavelength and energy.

Frequency:

Wavelength:

Energy:

What are the mathematical relationships between wavelength and frequency? What are the mathematical relationships between energy and frequency?

Calculate the wavelength,

frequency, and energy using

Planck’s constant and the speed of

light.

Example #1: If a particular green light has a wavelength of 4.9 x 10-7 m, what is its frequency? Example #2: The human eye can see light with a frequency about as high as 7.9 x 1014 Hz, which appears violet. Calculate the energy that one photon of violet light carries. Example #3: Find the energy of violet light if the wavelength is 4 x 10-7 m.

Use isotopic composition to

calculate average atomic mass of an

element.

What is an isotope? Calculating Average Atomic Mass:

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Example: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O.

Express the arrangement of

electrons in atoms through electron

configurations and Lewis valence

electron dot structures.

Electron Configuration

Longhand configuration: Shorthand configuration:

Lewis Dot Diagrams

represent the valence e-


Rung 2: Homework

1. If the wavelength of a certain light is 6.5X10-7m what is the frequency?

2. The frequency of a wave is found to be 9.0X1014Hz. What is the wavelength?

3. The energy of one photon of light is 4.9X10-19J. What is the frequency of this light?

4. The frequency of a wave is 4.0X1014Hz. Calculate the energy.

5. Electrons travel as waves within the atom. Calculate the wavelength of a wave if the energy is 6.9X10-19J.

6. Determine the energy associated with a wave if the wavelength is 9.1X10-7m.

7. Write the electron configuration for the following atoms:

a. Bromine

b. Zirconium

c. Strontium

d. Oxygen

e. Silver

8. Draw the Lewis valence electron dot structures for the following atoms:

a. Bromine

b. Strontium

c. Hydrogen

d. Xenon

e. Calcium


Rung 3

Reporting Category 3: Bonding and Chemical Reactions

Name ionic compounds, covalent compounds, and acids using IUPAC nomenclature rules.

How are ionic compounds named using IUPAC rules? How are covalent compounds named using IUPAC rules? How are acids named using the IUPAC rules?

Write the chemical formulas of common polyatomic ions, ionic compounds, covalent compounds and acids

How do you write the chemical formulas of ionic compounds? How do you write the chemical formulas of covalent compounds? How do you write the chemical formulas of acids?

Construct electron dot formulas to illustrate ionic and covalent bonds

How to construct electron dot formulas to illustrate ionic bonds: How to construct electron dot formulas to illustrate covalent bonds:

Describe the nature of metallic bonding and apply the theory to explain metallic properties such as thermal and electrical conductivity, malleability, and ductility.

How can you describe the nature of metallic bonding? How can you apply metallic bonding theory to explain metallic properties?

Predict molecular structure for molecules with linear, trigonal planar or tetrahedral electron pair geometries using VSEPR theory.

Molecular Structures

Linear Trigonal planar Tetrahedral


Ladder to Success Rung 3: Homework

Name the following:

1. BaSO3

2. (NH4)3PO4

3. PBr5

4. MgSO4

5. CaO

6. H3PO4

7. Na2Cr2O7

8. MgO

9. SO3

10. Cu(NO3)2

Write the formula of the following:

11. hydrobromic acid

12. chromium(III) carbonate

13. magnesium sulfide

14. iodine trichloride

15. lithium hydride

16. ammonium hydroxide

17. calcium chloride

18. hydroselenic acid

19. iron(II) nitride

20. aluminum hydroxide

Draw Lewis diagrams for the following and indicate the shape.

21. CBr4

22. N2

23. AlCl3

24. SF4

25. PCl3


Rung 4

Reporting Category 4: Gases and Thermochemistry

Describe and calculate the

relations between volume,

pressure, number of moles, and

temperature for an ideal gas as

described by Boyles law, Charles law,

Avogadro’s law, Dalton’s law of

partial pressure, and the Ideal

gas law.

The following are variables in gas law calculations: P= V= n= R= T=

Gas Law Equation Description

Dalton’s Law of Partial Pressures

Boyles

Charles

Avogadro’s Law

Ideal Gas Law

Examples:

• The pressure on a balloon with a volume of 300 mL increases from 1.10 to 2.00 atm. Calculate the new volume of the balloon.

• The air in a balloon with a volume of 25 L is heated from 20 C to 60 C. If the pressure stays the same, what will be the new volume of the balloon?

• A sample of gas in a 1L flask at 1.5 atm contains 75% CO2 and 25% H2O gas. Calculate the partial pressures of each gas.

• Calculate the temperaure of 5.85 mol N2 gas in a 12 L steel bottle under 10 atm of pressure.

Perform stoichiometric calculations,

including determination

of mass and volume

relationships between

reactants and products for

reactions involving gases.

GAS STOICHIOMETRY

• Moles Liters of a Gas: – STP - use 22.4 L/mol – Non-STP - use ideal gas law

Non-STP – Given liters of gas?

• start with ideal gas law – Looking for liters of gas?

• start with stoichiometry conv. EXAMPLES:

• What is the mass of oxygen gas produced when 29.2 g of water is decomposed by electrolysis according to the balanced equation?

2H2O 2H2 + O2

• What volume in L of nitrogen dioxide gas is produced when 34 L of oxygen gas react with an excess of nitrogen monoxide? Assume conditions of STP.

2NO + O2 2NO2

Describe the postulates of

kinetic molecular

theory

• Gases are made of molecules in constant, random motion • The volume is small • Collisions are elastic • Forces (attractive and repulsive) are small • Average kinetic energy is proportional to temperature


Rung 4: Homework

1. The total pressure of a homogenous gaseous mixture is 780mmHg. If the gas mixture contains helium at a pressure of 190mmHg, oxygen at a pressure of 200mmHg, and neon what is the partial pressure of the neon gas?

2. What is the volume, in L, of 4.0 moles of carbon dioxide gas at 10.0ºC and 867mmHg?

3. A given sample of gas has a volume of 4.20 L at 60.0C and 1.00 atm. Calculate its pressure if

the volume is changed to 5.00 L and the temperature to 27C.

4. A gas sample contained in a cylinder equipped with a moveable piston occupied 300. mL at a pressure of 2.00 atm. What would be the final pressure if the volume were increased to 500. mL at constant temperature?

5. A fixed quantity of gas at 23.0C exhibits a pressure of 748 torr and occupies a volume of 10.3 L.

Calculate the volume the gas will occupy if the temperature is increased to 145C while the

pressure is held constant.

6. How many grams of AlCl3 must decompose in order to produce 3.10 dm3 of Cl2 at

50.0C and 98.4 kPa? (HINT: You must correct to STP.)

2AlCl3 2Al + 3Cl2

7. What volume of nitrogen can be produced by the decomposition of 50.0 g of NH4NO2 at

25C and 1.20 atm? (HINT: You must correct to STP.)

NH4NO2 N2 + 2H2O


Ladder to Success Rung 5

Reporting Category 5: Solutions

Describe the unique role of water in chemical and biological systems.

Factors that contribute to water’s unique properties: 1. Polarity:

2. Hydrogen bonding:

Develop and use general rules regarding solubility through investigations with aqueous solutions

Using the solubility curve:

• What is the solubility of KCl at 25C?

On the line = Below the line = Above the line= Using the solubility rules: ** Refer to your reference materials!! Example:

• Is Na2CrO4 soluble in water?

• IS BaSO4 soluble in water?

• Which product in the equation below is a precipitate? AlCl3 + K3PO4 3KCl + AlPO4

Calculate the concentration of solutions in units of molarity

How to calculate the molarity of a solution: Molarity= Example #1: What is the molarity of the solution if .25 moles of Na2SO4 is dissolved in 1.5 L solution? Example #2: If you had a 2M solution of glucose (C6H12O6) how many liters of the solution would contain 3 moles glucose? Example #3: Calculate the molarity of .51L of solution that contains 110g of NaCl.

Use molarity to calculate the dilutions of solutions

Moles of solute before dilution = Moles of solute after dilution Example #1: What volume of .5M NaOH is needed to make a .075 M in 2L solution?

Distinguish between types of solutions such as electrolytes and nonelectrolytes and unsaturated, saturated, and supersaturated solutions.

Solution Description Meaning

Electrolyte

Nonelectrolyte

Unsaturated

Saturated

Supersaturated

Investigate factors that influence solubilities and rates of dissolution such as temperature, agitation, and surface area.

Factor Effect on solubility Effect on Rate of Dissolution

Temperature

Agitation

Surface Area

How does pressure influence solubilities and rates of dissolution?


Rung 5: Homework

1. Calculate the amount in moles of hydrochloric acid needed to make 2000.mL of a 1.5M solution.

2. Find the volume of a solution if 5.00 moles of solute are present and the molarity is 2.25M.

3. If 88.0 grams of NaOH are dissolved in water and the volume of the solution is 3000.mL what is the molarity of the solution?

4. To prepare a dilute solution a student used 500.mL of a 12.0M HCl solution. The final volume of the dilute solution was 1750.mL. Calculate the molarity of the dilute solution.

5. 1.5L of a 3.0M solution was diluted to a concentration of 1.8M. What is the volume of the dilute solution?

6. Determine if the following ionic compounds are soluble or insoluble in water: a. Ba(CN)2

b. BaSO4

c. Al(OH)3

d. Sr(OH)2

e. CaCO3

f. Na2CO3


Rung 6

Reporting Category 1: Matter & Periodic Table

Explain the use of chemical and

physical properties in the

historical development of

the Periodic Table

How were chemical and physical properties used in the development of the periodic table?

Word Definition What’s so special?

Group

Period

Atomic #

Atomic mass

Valence electrons

Oxidation #

Metal

Nonmetal

Metalloid

Use the periodic table to identify

and explain periodic trends, including atomic and ionic radii,

electronegativity, and ionization

energy

Property Description Trend

Atomic radius

Ionic radius

Electronegativity

Ionization energy

Use the periodic table to identify and explain the

properties of chemical families,

including alkali metals, alkaline

earth metals, halogens, noble

gases and transition metals

Family Description Oxidation #

Alkali metals

Alkaline earth metals

Halogen

Noble gases

Transition metals


Ladder to Success Rung 6: Homework

1. Label the following groups: alkali metals, alkaline earth metals, halogens, and noble gases. Also

label the transition metals. List a few properties of each.

2. Use the periodic table to identify/label and explain periodic trends, including atomic and ionic

radii, electronegativity, and ionization energy

a. Which has the larger atomic radii? Mg or Cl

b. Which has the greater electronegativity? P or O

c. Which has the lower ionization energy? K or Br


Rung 7

Reporting Category 2: Atomic Structure and Nuclear Chemistry

Describe the characteristics of alpha, beta and

gamma radiation

Particle Alpha Beta Gamma

Symbols

Charge

Mass

Speed

Penetration

Describe radioactive decay

process in terms of balanced nuclear

equations

Alpha Example: Beta Example: Gamma: Gamma rays do not usually appear in nuclear equations since rays are emitted.

Compare fission and fusion

Fission Fusion


Rung 7: Homework

1. Which particle is more massive: alpha, beta , or a neutron?

2. What can be used to shield (or protect) someone from alpha radiation?

3. What can be used to shield (or protect) someone from beta radiation?

4. What can be used to shield (or protect) someone from gamma radiation?

5. Complete the following nuclear radioactive decay equations:

210 206

Po → Pb +

84 82 ________________

239 239

U → Np +

92 93 ________________

259 258

Md → Md +

101 101 ________________


Rung 8

Reporting Category 3: Bonding and Chemical Reactions

Define and use the

concept of a mole

How do chemist define a mole?

Representative particles=

Use the mole

concept to

calculate the

number of atoms,

ions or molecules

in a sample of

material

Example #1: A sample consists of 6.85 x 1020

atoms of carbon. How

many moles does the sample contain?

• Example #2: Another sample consists of 2.58 mol of water. What is the

number of water molecules in the sample?

• Example #3: How many atoms are in 2 moles of He?

• Example #4: Find the number of chlorine ions in 5 grams of CaCl2.

• Example #5: How many molecules are in 7.1 grams of water?

Calculate percent

composition and

empirical and

molecular

formulas

PERCENT COMPOSITION:

• Example #1: Calculate the percent composition of each element in MgCl2

EMPIRICAL FORMULAS:

• Example #1: Find the empirical formula for a compound that is 79.8% C

and 20.2% H.

MOLECULAR FORMULAS:

• Example #1: The empirical formula for a molecule is NH2. If its molecular

weight is 32 amu, what is the compound’s molecular formula?

Use the law of

conservation of

mass to write and

balance chemical

equations

LAW OF CONSERVATION OF MASS :

Total mass of reactants = Total mass of products

Balancing Equations:

H2 + N2 NH3

Perform

stoichiometric

calculations,

including

determination of

mass relationships

between reactants

and products,

calculation of

limiting reagents

and percent yield

• Example #1: Suppose 8.75 g of propane react with oxygen gas to produce

carbon dioxide and water. How many grams of water are produced?

C3H4 + 5O2 3CO2 + 4H2O

• Example #2: What is the limiting reagent when 36 g of CH4 when it

reacts with 98 g O2 to produce carbon dioxide and water?

CH4 + 2 O2 CO2 + 2H2O

• Example #3: You calculate a theoretical yield of 55 g of water but the

actual yield from your experimentation was 48.8 g of water. What is the

percent yield?


Rung 8: Homework

1. During a reaction 15.0g of magnesium reacted with excess oxygen. After the reaction students

collected 20.2g of magnesium oxide powder. Determine the percent yield for this reaction.

2Mg(s) + O2(g) → 2MgO(s)

2. When 60.0L of F2 react how many L of Cl2 are produced?

3F2(g) + 2AlCl3(s) → 3Cl2(g) + 2AlF3(s)

3. How many particles are in 10.0 moles of CaCO3?

4. How many ions are in 10.0 moles of K2SO4?

5. How many moles of particles are in 3.01X1024 CaCO3 particles?

6. How many moles of atoms are in 3.01X1024 K2SO4 particles?

7. How many moles of ions are in 3.01X1024 NaCl particles?

8. Balance the following chemical equations:

a. ___CH4 + ___O2 → ___H2O + ___CO2

b. ___LiOH → ___Li2O + ___H2O

c. ___Mg + ___Al2(CO3)2 → ___Al + ___MgCO3


Ladder to Success Rung 9

Reporting Category 4: Gases and Thermochemistry

Understand

energy and its

forms including

kinetic,

potential,

chemical and

thermal energies

What is energy?

Form Description

Kinetic

Potential

Chemical

Thermal

Understand the

law of

conservation of

energy and the

processes of heat

transfer

Law of Conservation of Energy:

Heat Transfer:

Conduction:

Convection:

Radiation:

Use

thermochemical

equations to

calculate energy

changes that

occur in

chemical

reactions and

classify

reactions as

exothermic or

endothermic

Enthalpy:

Reaction Enthalpy Change Description

Exothermic

Endothermic

Example #1: Calculate the change in energy for the following reaction at standard

conditions. Is the is reaction exothermic or endothermic?

Example #2: Calculate the change in energy for the following reaction. Is this

reaction endothermic or exothermic?

Perform

calculations

involving heat,

mass,

temperature

change, and

specific heat

Example #1: Calculate the heat absorbed by a 20 g piece of copper metal that is

heated from 25C to 125C. The specific heat of copper is .385.

Example #2: The specific heat of water is 4.18. A 1200 g water sampel at 19C loses

10,000 J of heat. What is the final temperature?

Use calorimetry

to calculate the

heat of a

chemical process

Calorimeter: tool used to measure heat of a chemical process

Energy released by the reaction = Energy absorbed by the solution


Rung 9: Homework

1. The reaction of zinc with nitric acid was carried out in a calorimeter. This reaction caused the

temperature of 72.0 grams of liquid water, within the calorimeter, to raise from 25.0C to 100.C.

If the specific heat of water is 4.18 J/(g•K) calculate the energy associated with this reaction.

2. A 4.00 gram sample of solid gold was heated from 274K to 314K. If the specific heat of gold is 0.129

J/(g•K) how much energy was involved in this change?

3. Calculate the change in enthalpy for the following reaction given that the standard enthalpy of

formation for water is -285.4kJ/mol, 0.0 kJ/mol for oxygen, and -187.8kJ/mol for hydrogen

peroxide.

2H2O(l) + O2(g) → 2H2O(l)

4. Calculate the H value for the following reaction:

CaSO4(s) CaO(s) + SO3(g) H = ? kJ/mol

5. Draw a graph to indicate exothermic and endothermic reactions.

Substance Hf (kJ/mol)

SO3(g) -395.7

CaSO4(s) -1434.5

CaO(s) -634.9


Rung 10

Reporting Category 5: Solutions

Define acids and bases and distinguish between Arhennius and Bronsted-Lowry definitions and predict the products in acid-base reactions that form water.

Theory Definition

Arrhenius Acid

Arrhenius Base

Bronsted-Lowry Acid

Bronsted- Lowry Base

Identifying Conjugate Acid: Identifying Conjugate Base:

Understand and differentiate

among acid-base reactions,

precipitation reactions, and

oxidation-reduction reactions.

Reaction Description and Example

Acid-Base

Precipitation

Oxidation-Reduction (Redox)

Define pH and use the hydrogen and

hydroxide ion concentrations to calculate the pH

of a solution

What is pH? Example 1: What is the pH of a solution with a H+ concentration of 1 x 10-5M. Example 2: What is the hydrogen concentration of a solution with a pH of 11. Example 3: What is the pH of a solution with a hydroxide concentration of 1.3 x 10-10. Example 4: What is the ph if [OH] = 8 x 10-5M

Distinguish between degrees of dissociation for strong and weak acids and bases

Strong Acid/Base: Weak Acid/Base:

chemistry staar review ladder to success rung 1 reporting ... staar review with...chemistry staar...

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