Name: _______________________________________________ Period:_________________________
2015 Prep Chemistry Spring Semester Review
Things to know:
PHYSICAL AND CHEMICAL CHANGES AND PROPERTIES
Physical Change: changes in form (size, shape, or state), but not in chemical composition. Remains the
same substance. Examples include boiling water or chopping wood.
Key word examples of Physical changes: dissolving, melting, and boiling.
Chemical Change: new substance is made, bond between atoms are made or broken; often see an
unexplained color change, temperature change, or the formation of gas or precipitate (solid). Examples
include electrolysis of water and burning wood.
Key word examples of Chemical changes: rusting, digestion, oxidizing, reacting, and decomposition
1. Classify the following as chemical or physical changes.
A. cutting wire
P
B. ripening tomato
C
C. apple slices turning brown
C
D. compressing a gas
P
E. tearing a sheet of paper
P
Things to know:
CLASSIFICATION OF MATTER
Pure substances are elements or compounds. A pure substance has a specific melting or boiling point.
Mixtures are combinations of two or more substances. They are not chemically bonded together. They
keep their own physical properties because they do not create a new substance. You can tell if a
substance is a mixture if it can easily be separated into its components (such as through filtration or with
a magnet.)
2. Samples A, B, and C contain samples of matter.
A. Explain in terms of composition why sample A represents a pure substance.
It contains only molecules of the same arrangement and relative size.
B. Explain why sample C would represent a mixture of Fluorine and Hydrogen chloride.
Flourine is diatomic so represented by 2 symbols of the same size. HCl is then represented by
a small circle for H attached to a larger circle for Cl.
C. Contrast sample A and B in terms of pure substances and mixtures.
B and C are mixtures of 2 different molecules while A is all the same molecules.
Things to know:
PERIODIC TABLE
Periodic Table: chart of elements arranged in rows
by atomic number (number of protons)
Family Groups: columns on the periodic table;
group number is equal to the valence electrons
Chemical Symbol: element abbreviation (H, He, Cu)
Atomic Mass: protons + neutrons
Metals: conduct heat/electricity; ductile (can be
pulled into wires); malleable (bendable); shiny;
typically solids at room temperature
Nonmetals: poor conductors of heat/electricity;
brittle; dull, many are gases at room temperature
Metalloids: conduct electricity under some
conditions
Chemical Similarity: elements in the same column are most similar to one another.
Valence Electrons: outermost electrons in electron transferred during chemical reactions
Noble Gases: Group VIIIA elements, considered stable (inert) because have full outer electron shells;
noble gases have electrons, except for Helium which has 2
3. Classify the following properties as alkali metal, alkaline-earth metal, halogen, noble gas, or
transition metal.
A. contains 7 valence electrons halogens
B. represents the most reactive metals
alkali metals
C. has multiple oxidation states (multiple charges) transition metals
D. inert
noble gases
E. contains 1 valence electron alkali metals
F. elements are all gases at STP noble gases
G. contains 2 valence electrons
alkali earth metals
H. have the highest ionization energies noble gases
I. contains the atom with the largest atomic radius alkali metals
4. According to the periodic table Magnesium will most likely react with elements of which group?
A. -1
B. 3
C. 16
D. 18
Explain your answer:
Mg has two valence electrons so will relase them to an element with 6 valence electrons to achieve
A full outer shell_________________
Things to Know:
ISOTOPES
Isotopes are atoms of the same element with different masses. Isotopes have different numbers of
neutrons. Isotopes have the same number of protons and electrons.
Mass number: The total number of protons and neutrons in the nucleus of an isotope.
Atomic number: total number of protons
Isotope Notation:
Things to Know:
NUCLEAR CHEMISTRY
Unlike normal reactions, nuclear reactions affect the nucleus. They can convert one element to another
releasing particles and tremendous energy
Radioactive Decay: process in which an unstable nucleus loses energy by emitting radiation
spontaneously and changes to a more stable element
Particle
Symbol
Charge
Mass
Speed
Penetration
Alpha
Beta
Gamma



4
2𝐻𝑒
+2
Large
Slow
Low
0
−1𝑒
-1
Small
Fast
Medium
0
0𝛾
0
None
Fastest
High
5. Match the following to the correct nuclear reaction
A.
239
94𝑃𝑢
→
235
92𝑈
B.
198
79𝐴𝑢
→
198
80𝐻𝑔
C.
220
87𝐹𝑟
+ __ii_ →
D.
37
20𝐶𝑎
E.
32
15𝑃
→
37
19𝐾
+ _ii__
+ _i__
224
89𝐴𝑐
+ _iv__
+ __i___ →
i.
0
−1𝑒
ii.
4
2𝐻𝑒
iii.
0
0𝛾
iv.
0
1𝑒
32
14𝑆𝑖
6. Write an equation for the decay of polonium-218 by alpha () emission.
218/84 Pu -> 4/2 He + 214/82 Pb
7. Write an equation for the decay of carbon-14 by beta (-) emission.
14/6 C -> 0/-1 Beta + 14/7 N
8. State whether the following statements are associated with fission or fusion reactions.
A. Light mass nuclei combine to form a heavier more stable nucleus__fusion____
B. A very heavy nucleus splits into more stable nuclei of intermediate mass_fission____
C. Takes place at extremely high temperature_fusion____
D. It is used in nuclear reactors to produce electrical energy___fission___
E. Takes place in the sun and other stars fusion__
F. Is the source of energy in a Hydrogen bomb___fusion__
9. Match each description with the appropriate type of radiation – alpha, beta, positron, or gamma.
A. A negatively charged electron. beta
B. Blocked only by several feet of concrete. gamma
C. A positively charged particle stopped by lead. positron
D. Blocked by paper or clothing. alpha
E. Radiation energy with no electrical charge.
gamma
F. The atom loses the most mass when it releases this radiation alpha
G. The atomic number increases by 1 when this radiation is released beta
Things to know:
IONIC, COVALENT AND METALLIC BONDS
Ionic Bond
Contains a metal and a nonmetal, the giving and taking of electrons.
Covalent Bond
Usually all nonmetals, only. Involves the sharing of electrons.
Metallic Bonding Electrons are free to move about, like the water in the sea, “sea of electrons”. Have
characteristics of metals such as high melting and boiling points, ductility, malleability
and good conductivity.
Malleability
Can bend (a property of metals).
Ductility
To draw into thin wires (a property of metals).
10. State whether the bonds in the following compounds are ionic or covalent
A. MgO ionic
B. H2O covalent
C. LiCl ionic
D. Br2 covalent
11. Are the following properties characteristics of ionic, covalent, or metallic bonding?
A. These bonds are formed by delocalized electrons in an “electron sea.”
B. These bonds involve a transfer of electrons.
C. Substances containing these bonds are malleable and have very high melting points.
D. Substances containing these bonds do not conduct electricity and have low melting points.
E. Compounds containing these bonds have a crystal lattice structure.
F. These bonds are formed by sharing electrons.
Things to know:
CHEMICAL REACTIONS
Reactants
Elements and compounds on the left side of the yield sign (arrow).
Products
Elements and compounds on the right side of the yield sign (arrow).
Law of
Balancing equations in a chemical reaction. Number of atoms on the reactant side
Conservation
of the equation must be equal to the number of atoms on the product side of the
of Mass
equation. Reactants = Products.
Coefficient
The number in front of the element or compound in a chemical equation.
Don’t forget the diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2
Be sure to change coefficients, not subscripts.
Coefficients may only be placed in front of compounds, not between them.
Make sure that your coefficients are in the lowest terms.
Never add, delete or change subscripts.
Balance all elements first before balancing hydrogen(H) and oxygen(O).
Be sure to correctly count atoms before and after balancing.
Symbols used in Chemical Reactions
(s) means the substance is in a solid state
(g) means the substance is in a gaseous state
(l) means the substance is in a liquid state
(aq) means the substance has been dissolved in an aqueous solution (dissolved in water)
→ means “yields” and is used to indicate the result of the reaction.
∆ (over the reaction arrow) indicates the application of heat to the reactants
Write and balance the following word equation using chemical formulas, physical states, and energy.
12. Given the balanced equation: KCl(aq) + AgNO3(aq) → KNO3(aq) + X
What is the correct formula for the product represented by the letter X?
A) AgCl2(s)
B) K2Cl(aq)
C) AgCl(s)
D) KCl2(aq)
13. Given the balanced equation: H2SO4(aq) + BaCl2(aq) → 2HCl(aq) + X
What is the correct formula for the product represented by the letter X?
A) Ba(SO4)2(s)
B) H2O(l)
C) Ba2SO4(s)
D) BaSO4(s)
14. Given the balanced equation: Cl 2(g) + 2KBr(aq) → X + Br2(l)
What represents the missing product X?
A) 2K(aq)
B) 2H2O(g)
C) 2KCl(aq)
D) Cl2(g)
15. Given the balanced equation: 2K + X → 2KCl + Ca
What represents the missing reactant X?
A) CaCl2
B) 2ClC) Cl2
D) 2CaCl2
*Remember, the Law of
Conservation of Mass. The # of
atoms on the reactants side of
the equations must be equal to
the number of atoms on the
product side.
16. Given the equation: __FeCl 2 + __Na2CO3 → __FeCO3 + __NaCl
When the equation is correctly balanced using the smallest whole numbers, the
coefficient of NaCl is
A) 6
B) 2
C) 3
D) 4
17. Given the unbalanced equation: __C3H8(g) + __O2(g) → __H2O(g) + __CO2(g)
When the equation is completely balanced using smallest whole numbers, the
coefficient of O2
A) 5
B) 2
C) 3
D) 10
18. When the equation H2 + N2 → NH3 is completely balanced using the smallest whole numbers,
the sum of all the coefficients will be
A) 12
B) 7
C) 3
D) 6
Things to know:
PERCENT COMPOSITION
Percentage by mass of each element present in a compound
% composition =
mass of the element in sample
 100
total mass of sample
Empirical formula
The lowest ratio of subscripts in a chemical formula, reduced subscripts.
Molecular formula
The actual formula of a compound. Not reduced. It shows the number and kinds of
atoms in a molecule. Will need the empirical formula first.
19. Find the % composition of copper(II)chloride (CuCl2).
Cu= 47% Cl2= 53%
20. The percent composition of a compound is 40.0% C, 6.7% H, and 53.7% O. The molecular
mass of the compound is 180.0 g/mol. Find its empirical and molecular formulas.
Empirical= CH2O
Molecular= C6H12O6
Things to know:
CALCULATING ATOMS, MOLES, MOLECULES, IONS AND GRAMS
Ions = formula units
1mol = 6.02 x 1023 formula units
1mol = the atomic mass of an element on the periodic table
Avogadro’s number: 6.02 x 1023 atoms/molecules
g = grams of the element from the periodic table’s atomic mass number
Be sure to always include units! It helps to avoid mistakes.
÷ by molar mass
Mass or Grams (g)
x 6.02 x1023
Moles (mol)
x by molar mass
atoms or molecules
÷ by 6.02 x1023
21. How many magnesium sulfate molecules are in 25.0 g?
1.25 x 1023 molecules of MgSO4
22. How many atoms are there in 3 moles of calcium?
1.8 x 1024 atoms Ca
23. How many atoms are there in 12.5 g of potassium?
1.93 x 1023 atoms of K
24. How many fluoride ions are in 1.46 moles of aluminum fluoride?
2.64 x 1024 ions of Flourine (F3)
Things to know:
Always use a balanced equation.
Remember, grams is the atomic mass number found on the periodic table.
To solve stoichiometry problems:
1. Write a balanced equation.
2. Identify known & unknown.
3. Under the balanced chemical equation, write the coefficients that indicate the mole
ratios.
4. Above the balance equation write the unknown and known data from the problem in
the appropriated chemical formula.
5. When need it, multiply the moles by their respective molar masses to change them to
grams.
6. Set up a proportion to relate reactants amounts with products amounts.
7. Use cross-multiplication to solve for the unknown
1. Check answer
Using the equation below, answer the following 2 questions. Be sure to balance the equation first.
_Cr + _CuSO4 → __Cu + ___Cr2(SO4)3 2,3,3,1
25. How many grams of copper would be produced from 49.48 g of chromium?
=90.71 g Cu
26. How many grams of chromium are needed to react with in 1.25 moles of a CuSO4?
=43.33 g Cr
Things to know:
ELECTROLYTES AND NONELECTROLYTES
Electrolytes are substances that conduct electricity when dissolved in water.
Strong Electrolytes
Weak Electrolytes
Bonds
Ionic
Ionic
Completely creates ions in
Only partially creates ions
Ions
water
in water
Partially conducts
Electricity
Conducts electricity
electricity
Strong Acids: HCl, H2SO4
Weak acids: CH3COOH
Examples
Strong Bases: KOH, NaOH
Weak bases: Cu(OH)2
27. Define an electrolyte. Give examples.
Non-Electrolytes
Covalent
Creates no ions in water
Does not conduct
electricity
Sugar, Oil, Water
A substance that can conduct electricity. Ex. Ionic compounds/solutions, salts, acids, bases
_______________________________________________________________________________
28. Define a non electrolyte. Give examples.
A substance that won’t conduct electricity. Ex. Sugar, water
_______________________________________________________________________________
29. Why are strong acids also strong electrolytes?
Due to complete ionization or dissociation in solution, there are more ions in solution which allow
Electricy to flow in the solution.
________________________________________________________________________
30. Why is HCl a strong acid and HF a weak acid?
HCl ionizes or dissociates more completely than HF
____________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
31. Name two weak acids. Hydrogen peroxide, milk, acetic acid (vinegar), ascorbic acid, coffee
32. Circle the chemicals that are electrolytes:
A. NaCl
D. NaOH
B. C6H1206
E. HCl
C. CaCl2
F. HF
Things to know:
SOLUBILITY RULES
Solubility Rules: Generalizations that help you determine which compounds can dissolve in water. These
“rules” can help you predict if a compound in a reaction might form a precipitate. General solubility rules
can be found on the STAAR Chemistry reference materials sheet.
Soluble: the chemical will dissolve.
Insoluble: the chemical will not dissolve.
Precipitate: a solid product formed in a reaction.
Solid
Liquid
Gas
Temperature
Solubility Increases
No pattern
Solubility Decreases
Pressure
No effect
No effect
Solubility Increases
Surface Area
Solubility Increases
Solubility Increases
No effect
33. Use the solubility rules on the STAAR Chemistry reference materials sheet to predict whether each
of the following compounds is considered soluble or insoluble:
A. KCl
S
F. Pb(ClO3)2
S
B. NaNO3
S
G. (NH4)2S
S
C. AgCl
I
H. PbCl2
I
D. BaSO4
I
I. FeS
I
E. Ca3(PO4)2 I
J. Al2(SO4)3
S
34. Solutions of AgNO3(aq) and KCl(aq) are mixed. Will a visible reaction occur?
A. Yes, because KNO3 will precipitate out of solution.
B. Yes, because AgCl will precipitate out of solution.
C. No, because KNO3 is soluble in water.
D. No, because AgCl is soluble in water.
Things to know:
SOLUTION CALCULATIONS
Solution: homogenous mixture
Solute: the substance that is being dissolved
Solvent: the substance that is being dissolved
Molarity: unit of concentration. Formula can be found on the STAAR Chemistry Reference Materials sheet.
Dilution: Formula used to calculate concentration and volume when diluting a chemical. Formula can be
found on the STAAR Chemistry Reference Materials sheet.
35. Find the molarity of a 750 mL solution containing 346 g of potassium nitrate.
M=4.66 g/L
36. Which solution is the most concentrated?
A. 6 moles of solute dissolved in 4 liters of solution
B. 2 moles of solute dissolved in 3 liters of solution
C. 1 mole of solute dissolved in 1 liter of solution
D. 4 moles of solute dissolved in 8 liters of solution
Things to know:
TYPES OF SOLUTIONS
Saturated Solution: contains the maximum amount of dissolved solute; unable to dissolve more solute.
If more solute is added, it will fall to the bottom and not dissolve.
Unsaturated Solution: contains less than the maximum amount of solute
Supersaturated Solution: contains more than the maximum amount of solute. Can be created by
heating up a solution and dissolving as much solute as possible. Must allow it to cool undisturbed. It will
re-crystallize quickly if seed crystal is added.
37. Plot a solubility graph for AgNO3 from the following data. Place grams of solute on the vertical
axis. (Use increments of 50). Place temperature (in °C) on the horizontal axis.
Grams solute AgNO3
(per 100g H2O)
122
216
311
440
585
733
Temperature
°C
0
20
40
60
80
100
A. How does the solubility of AgNO3 vary with the temperature of the water? As solubility increase,
temperature increases
B. Estimate the solubility of AgNO3 at 35°C, 55°C, and 75°C.
~275g, ~400g, ~520g
C. At what temperature would the solubility of AgNO3 be 275g per 100g of H2O? _~35 degrees C__
D. If 100g of AgNO3 were added to 100g of H2O at 100°C, would the resulting solution be
saturated or unsaturated? _______unsaturated______
What type of solution would occur if 325g of AgNO3 were added to 100g of water at 35 °C?
____supersaturated___________
E. Using the table below, how many grams of solute will precipitate out of solution if a saturated
solution of KNO3 at 60°C is cooled down to 20°C? 74.4g
Grams KNO3
Temperature
(per 100g H2O)
(°C)
13.9
0
31.6
20
61.3
40
106
60
167
80
245
100
Things to know:
REACTIONS
Redox Reaction: A reaction in which electrons are transferred between elements. (You can identify a
redox reaction when elements change charges). Single replacement, synthesis or decomposition reactions
can be classified as Redox reaction.
Precipitate Reaction: A reaction between two aqueous solutions that forms a solid precipitate product.
Double displacement can be classified as Precipitate reaction.
Acid-Base Reaction: A reaction between an acid and a base. The products will form a salt and water.
Double displacement- Neutralization can be classified as Acid-Base reaction.
38. Label the following reactions as redox, precipitate, or acid-base reactions.
A.
AgNO3 +KI → AgI+KNO3 precip
C.
2KOH + H2SO4 → K2SO4 + H2O AB
B.
Cu + 2AgNO3 → Cu(NO3)2 + Ag Re
D.
Ba(OH)2 + 2HCl → BaCl2 +H2OAB
Things to know:
ACIDS AND BASES
pH: measures hydrogen ion [H+] concentration in solution (can use pH paper or litmus paper);
scale of 0-14; neutral substances (like water) have pH of 7; acids have pH <7; bases have pH >7. pH can
be calculated using a formula on the STAAR Chemistry Reference Materials sheet.
pOH: measure of the concentration of hydroxide [OH-] ions. The hydroxide concentration can be
calculated when you know the hydrogen ion [H+] concentration and the Kw. Formulas for calculation can
be found on the STAAR Chemistry Reference Materials sheet.
BASE
Greater than 7
Less than 7
Ends in –OH
Accepts protons
Tastes bitter, feels
slippery, electrolyte,
hazardous, chemically
corrosive
Litmus Paper
Turns paper red
Turns paper blue
Examples
HCl
NaOH
Salt: Any ionic compound that results from a reaction between an acid and a base.
pH
pOH
Arrhenius
Bronsted Lowry
Characteristics
ACID
Less than 7
Greater than 7
Begins with H
Donates proton
Tastes sour, electrolyte,
hazardous, chemically
corrosive
Neutralization: a reaction that occurs between an acid and a base. The result is a salt and water.
39.
Which compound reacts with an acid to form a salt and water?
A.
KOH
C. CH3
B.
HCl
D. KCl
40. According to the Arrhenius theory, which list of compounds includes only bases?
A.
KOH, NaOH, and LiOH
C. NH3, KOH, and ClB.
NaOH, Ca(OH)2, and CH3COOH
D. H2O, OH-, and Cl-
NO2-(aq) + H2O(l) → HNO2(aq) + OH-(aq)
41. Predict the products for the following reactions and balance the
A.
NaOH + H2SO4 → Na2SO4 + 2H2O
B.
Ca(OH)2 + HCl → CaCl2 + 2H2O
equations:
42. Calculate the pH of a solution with an OH- concentration of 1 x 10-10M.________4___________
43. Calculate the pOH of a solution with an H+ concentration of 1x 10-8 M. ________6___________
44. What is the H+ concentration of a solution with a pH of 6? _______1.0 x 10-6______
45. What is the OH- concentration of a solution with a pH of 3? ______1.0 x 10-11_______