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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?
Chemistry STAAR Review Ladder to Success
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:
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-
Chemistry STAAR Review Ladder to Success
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
Chemistry STAAR Review Ladder to Success
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
Chemistry STAAR Review
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
Chemistry STAAR Review Ladder to Success
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
Chemistry STAAR Review Ladder to Success
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
Chemistry STAAR Review
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?
Chemistry STAAR Review Ladder to Success
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
Chemistry STAAR Review Ladder to Success
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
Chemistry STAAR Review
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
Chemistry STAAR Review Ladder to Success
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
Chemistry STAAR Review Ladder to Success
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 ________________
Chemistry STAAR Review Ladder to Success
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?
Chemistry STAAR Review Ladder to Success
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
Chemistry STAAR Review
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
Chemistry STAAR Review Ladder to Success
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
Chemistry STAAR Review Ladder to Success
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: