1
Student Name: ______________________ The following booklet is designed to provide a review of important concepts that have been covered throughout your science education, with particular emphasis on concepts from Chemistry 11. The first few days of Adv./Chemistry 12 will be dedicated to the completion of this booklet. Students are expected to review any relevant material using their textbooks and to complete the work independently. Teacher support will be provided, however material will not be re-taught.
Topics included are:
Significant digits Metric system Scientific notation Working with mathematical formulas – specific heat capacity Identifying & naming compounds & calculating molar mass Writing chemical formulas Balancing chemical equations Identifying reaction types Predicting products and writing chemical reactions Mole-mass conversions and stoichiometry Introduction to acids and bases Molecular structure and polarity Solubility Hydrocarbon derivatives
2 Significant Digits A measurement can only be as accurate and precise as the instrument that produced it. A scientist must be able to express the accuracy of a number, not just its numerical value. We can determine the accuracy of a number by the number of significant digits it contains. Rules for identifying the number of Significant digits a measurement contains.
1) All digits 1-9 inclusive are significant. Example: 129 has 3 significant digits. 2) Zeros between significant digits are always significant. Example: 5,007 has 4 significant digits. 3) Trailing zeros in a number are significant if the number contains a decimal point. Trailing zeros in a number may be significant if the number does not contain a decimal point. Example: 100.0 has 4 significant digits. 100 has 1, 2 or 3 significant digits (best to convert to scientific notation so rule 5 applies). 4) Zeros at the beginning of a number whose only function is to serve as a place holder are NOT significant. Example: 0.0025 has 2 significant digits. 5) Zeros following significant decimal places (i.e. trailing zeros) are significant. Example: 0.000470 has 3 significant digits. 0.47000 has 5 significant digits. Let’s Practice: Determine the number of significant digits in the following numbers. 1) 0.015
_________
6) 7 004 003
_________
2) 81040
_________
7) 0.00043
_________
3) 0.11020 _________
8) 150
_________
4) 50000.
9) 0.03250
_________
10) 6199.5
_________
_________
5) 7.00 x 10-3 _________
3 Calculations Using Significant Digits: There are two rules for calculations involving measured values. The key concept to remember is that an answer cannot be more precise than your least precise measurement. 1) When multiplying and dividing, round the final answer to the least number of significant digits in any of the measurements. Example:
23.0cm x 432cm x 19cm = 188 784cm3
The answer is expressed as 190 000cm3 since 19cm has only two significant digits. (You could also express the answer using scientific notation = 1.9 x 105cm3)
2) When adding and subtracting, round your answer to contain the least number of significant decimal places in any of the measurements. Example:
123.25ml + 46.0ml + 86.257ml = 255.507ml
The answer is expressed as 255.5ml since 46.0ml has only one significant decimal place.
Let’s Practice: Perform the following operations expressing the answer with the correct number of significant digits. Don’t forget to include units! 1) 1.353m x 2.4670m
_______________
2) 1047m2 57m
_______________
3) 12.01ml + 35.2ml + 6ml
_______________
4) 55.46mm - 28.9mm
_______________
5) 0.21cm x 3.2cm x 100.1cm
_______________
6) 0.15cm + 1.15cm + 2.051cm
_______________
7) 150 L3 4 L
_______________
8) 1.1252mm x 0.115mm x 0.012mm
_______________
9) 1.278 x 103m2 1.4267 x 102m
_______________
10) 605.1kg - 450.25kg
_______________
4 Metric System Converting from one unit to another mega M 1 000 000 106
kilo
hecto deca basic unit K h da g, L, m 1000 100 10 1 3 2 1 10 10 10 100
deci centi milli
micro
d
c
m
.1 10-1
.01 10-2
.001 .000001 10-3 10-6
The factor label method can be used to solve virtually any problem involving a change in units. It is especially useful in making complex conversions dealing with concentrations and derived units. Example:
100dag 0.005 kg x = 0.5 dag 1kg
Remember: When converting measurements to new units the number of significant digits should not change.
Let’s Practice: Complete the following chart: Amount Conversion 35 mL =
cL
950 g =
kg
275 mm =
cm
1000 L =
kL
1000 L =
mL Scientific Notation
Scientists have developed a short hand method to express very large or very small numbers. This method is called scientific notation. Scientific notation is based on powers of the base number 10. For Example: The number 123,000,000,000 in scientific notation is written as:
1.23 x 1011
5 Let’s Practice: Convert the following measurements to scientific notation. Be careful to consider significant digits. 1) 0.0015m
_________________
2) 5050m
_________________
3) 0.2150m
_________________
4) 0.000083m _________________ 5) 743000m
_________________
Convert the following measurements to standard notation. Be careful to consider significant digits. 1) 1.15 x 103cm
_________________
2) 3.75 x 10-4cm
_________________
3) 1 x 104cm
_________________
4) 1.2 x 10-4cm
_________________
-5
5) 2.20 x 10 cm
_________________ Working with Mathematical Formulas
Using the formula, q=mct, answer each of the following questions. Where “q” is heat in Joules (J), “m” is mass in grams (g), “c” is specific heat capacity in Joules per gram degree Celsius (J/goC), and “t” is temperature change in degrees Celsius (oC). Example: How many joules of heat are given off when 5.2g of water cools from 70.oC to 25oC? (Specific heat capacity of water = 4.184J/goC) q = mct q = (5.2g)(4.184 J/goC)(25-70oC) q = -980J or -9.80 x 102J 1.
How many joules of heat are given off when 45g of water cools from 90.0 oC to 30.0oC?
2.
What is the temperature change that takes place if 824 joules of energy is released when 10.0g of water cools?
3.
What mass of water gives off 1200 joules of energy when there is a temperature change of 35 oC?
6 Identifying & Naming Compounds and Calculating Molar Mass Complete the following table: Formula
Type of Compound
Name of Compound
Molar Mass
Calcium oxide
40.08 + 16.00 = 56.08 g/mol
(ionic vs. molecular)
Example: CaO KMnO4 Al(SCN)3
N3O6
Ca(NO3)2
PCl3
Zn(C2H3O2)22H2O
Na3PO4
NH4OH
SF6
CaSO4 PbO2
Ionic
7 Writing Chemical Formulas Write the correct formulas for each of the following compounds. Name
Formula
Sodium hydroxide lithium chloride
NaOH
Example:
phosphoric acid iron(III) chloride magnesium phosphate carbon tetrachloride barium carbonate ammonium chloride dinitrogen pentoxide hydrofluoric acid lead(IV) sulfate
Balancing Chemical Equations Balance the following chemical equations: 1) ___ KClO3 → ___ KCl + ___ O2 2) ___ AgNO3 + ___ MgCl2 → ___ AgCl + ___ Mg(NO3)2 3) ___ AlBr3 + ___ K2SO4 → ___ KBr + ___ Al2(SO4)3 4) ___ C8H18 + ___ O2 → ___ CO2 + ___ H2O 5) ___ H2S + ___ KOH → ___ H2O + ___ K2S 6) ___ H3PO4 + ___ NH4OH → ___ H2O + ___ (NH4)3PO4 7) ___ NaHCO3 → ___ Na2CO3 + ___ CO2 + ___ H2O 8) ___ NaHSO4 + ___ NaClO + ___ NaCl → ___ Cl2 + ___ H2O + ___ Na2SO4 9) ___ HCl + ___ CaCO3 → ___ CaCl2 + ___ H2O + ___ CO2 10) ___ Ba(OH)2●8H2O + ___NH4SCN → ___NH3 + ___Ba 2+ + ___SCN- + ___ H2O
8 Identifying Reaction Types, Predicting Products and Writing Chemical Reactions Write names of products, balanced chemical equations and the reaction type for the following reactions. a. Carbon tetrachloride decomposes. ____________________________________ Balanced equation: _______________________________________________ Reaction type: _________________________ b. Propane (C3H8) burns. ______________________________________________ Balanced equation: _______________________________________________ Reaction Type: _________________________ c. Aluminum reacts with gold (III) phosphate. ______________________________ Balanced equation: _______________________________________________ Reaction Type: _________________________ d. Oxygen and aluminum combine. ______________________________________ Balanced equation: _______________________________________________ Reaction Type: _________________________ e. Calcium chloride decomposes. _______________________________________ Balanced equation: _______________________________________________ Reaction Type: _________________________ f. Sodium phosphide and ammonium nitrate react. _________________________ Balanced equation: _______________________________________________ Reaction Type: _________________________ g. Calcium oxide and water combine, forming one product. ___________________ Balanced equation: _______________________________________________ Reaction Type: _________________________ h. Sodium chloride and lead react. ______________________________________ Balanced equation: _______________________________________________ Reaction Type: _________________________
9 Mole-Mass Conversions 1. Convert each of the following masses into moles. a. 8.40 g of NaOH __________________________ b. 4.2 kg of H2O __________________________ 2. Convert each of the following moles into masses (in grams). a. 0.456 mol of Al2(SO4)3 ___________________________ b. 0.518 mol of CuSO4 • 5H2O ___________________________ Stoichiometry & The Mole Mole-Mole Problems: 1)
2KClO3 → 2KCl + 3O2 How many moles of oxygen are produced by decomposing six moles of potassium chlorate?
2)
C3H8 + 5O2 → 3CO2 + 4H2O How many moles of oxygen are needed to completely react with four moles of propane?
3)
K3PO4 + Al(NO3)3 → 3KNO3 + AlPO4 How many moles of potassium nitrate are produced when two moles of potassium phosphate react with two moles of aluminum nitrate?
Stoichiometry: Mixed Problems: 1)
N2 + 3H2 → 2NH3 What volume of NH3 at STP is produced if 25.0g of N2 is reacted with an excess of H2?
2)
H2SO4 + 2NaOH → 2H2O + Na2SO4 How many molecules of water are produced if 2.0g of sodium sulfate are produced in the above reaction?
3)
2KClO3 → 2KCl + 3O2 How many potassium ions are produced if 5.0grams of KClO3 decompose?
10 Introduction to Acids & Bases The pH scale is a scale from 0 to 14. 0 represents a strong acid and 14 represents a strong base. 7, or the middle, is neutral. Let’s look at this in diagram format and see some examples of compounds that have different pH values. Orange juice, Soda
Acid (HCl) of Ammonia Baking stomach lining Black “Pure” Oven Solution Soda coffee Water Cleaner Battery Lemon juice, Tomato Sea Milk of Drain Soapy Juice acid vinegar Saliva Water Magnesia Cleaner Water
0
1
2
3
5
4
6
7
8
9
10
11
12
13
14
Well what does pH really mean? pH is a measure of acidity. Acidity comes from the Latin term acidus meaning sharp and if you think about it and look at the acidic compounds in the pH chart above (lemon juice, vinegar), these solutions all taste sharp when you taste them. The pH scale is great for determining how much an acid has dissociated or broken apart. If an acid is strong, it means that the acid completely dissociates or breaks apart into ions. For example, HCl is a strong acid. It is found in the lining of your stomach. HCl → H+ + ClH+
Cl-
→
H+
+
Cl-
The same is true for strong bases. NaOH → Na+ + OHNa+
OH-
→
Na+
+
OH-
But for weak acids and bases, there are lesser dissociations. What this means is that when the acid or base dissociates, only a portion of the ions can be found in the solution. Example: Vinegar (Acetic Acid) HC2H3O2 H+ + C2H3O2H+
C2H3O2+-
→
H+ +
C2H3O2+-
11 What this means is that only a small portion of the vinegar actually dissociates or breaks apart. More of the HC2H3O2 molecules stay together than HCl because vinegar is a weaker acid (pH = 2) than HCl. If we looked at soda, that contains phosphoric acid (H3PO4), with a pH = 3, or coffee, with a pH = 5, there would be even fewer H+ ions broken off and found in solution. The greater the proportion of ions, the stronger the acid or base. Another way to look at it is that if there are a lot of H+ ions in the solution you have a strong acid, if there are a lot of OHions in solution, we have a strong base. It is like a continuum. Let’s Practice: 1. For each of the following, indicate whether each of the following is an acid or a base. Name the substances. a) KOH
______________________
e) H2CO3
______________________
b) HNO3
____________________ __
f) NH4OH
______________________
c) NaOH
______________________
g) H2C2H3O2 ______________________
d) Ca(OH)2 ______________________
h) H2O
______________________
Molecular Structure and Polarity There is a relationship between the shape of a molecule and its polarity. 1) If the shape is linear, and the surrounding atoms are the same, the molecule will be non-polar. If the shape is linear and the surrounding atoms are different, the molecule will be polar. 2) Rule #1 also applies to tetrahedral and trigonal planar shapes. 3) Bent and pyramidal shapes are always polar. ex. :N ≡ N: is a linear molecule which is non-polar. Note: A general rule is that if a molecule is symmetrical, it is non-polar. Let’s Practice: Draw the correct 3-D (VSEPR) structure and determine whether it is polar or non-polar. 1) HF
4) CF2S
2) PH3
5) CCl4
3) C2H4
6) Br2O
12 Solubility 1. For each of the following, indicate whether the ionic compound will be aqueous or solid when in water at 25oC. a) KOH
______________________
e) PbCl2
______________________
b) Mg(NO3) ______________________
f) NH4OH
______________________
c) NaF
g) CaCO3
______________________
______________________
d) Ca(OH)2 ______________________
h) Ba3(PO4)2 ______________________
Hydrocarbons and Hydrocarbon Derivatives 1. Draw a structural diagram for each substance shown below: a) Butane
b) Cyclobutene
c) Butanal
d) Butene
e) Cyclobutyne
f) Butanoic acid
g) Butyne
h) Butanol
i) Butyl Butanoate
j) Cyclohexane
k) 2-Hexanone
l) Hexanamine
m) Hexanamide
n) Dibutyl ether
o) 2-bromohexane
