Download Final Exam Review Day 1

Acl Chem Final Exam Review - Day 1
Name __________________
Classify the following matter as an element, a compound or a mixture.
1. ____HCl
3. ____ Cl2 (g)
5. ____ barium
2 ____ HCl (aq)
4. ____ diet coke
6. ____ calcium nitrate
Classify the following matter as a pure substance or a mixture
1. ____ Na2CO3 (s)
3. ____ air
5. ____ dirt
7. _____ MgCl2
2. ____ Na2CO3 (aq)
4. ____ helium
6. ____ H2O (l)
8. _____ tea
Classify the following matter as homogeneous or heterogeneous.
1. ____ tossed salad
3. ____ stainless steel
5. ____ Ca(OH)2 (aq)
2. ____ Cl2
4. ____ saline solution
6. ____ MgSO4 (s)
Classify the following as a physical change or a chemical change.
1.
__________ dissolving sugar into water
2.
__________ decompose water into hydrogen and oxygen
3.
__________ burning gas
4.
__________ evaporate water
Classify the following as matter or energy
1. ____ air
2. ____ light
3. ____ heat
7. _____ Ne
Classify the following as a metal, nonmetal or metalloid.
1. ____ Na
2. ____ Kr
3. ____ As
7. _____ H
6. ____ brittle
7. ____ ductile
8. ____ insulator
9. _____ has luster
Units of measurement
cm3, L or mL is _____________________
g is _______________________
g/cm3, g/l or g/mL are ________________
K or ºC is __________________
joules , calories, kcal is _______________
Temperature conversions
ºC = K – 273
20º C = __________K
Metric Conversions
________ mm = 1 m
___________ cm = 1m
__________ m = 1 km
Significant Figures
1.
Reading Instruments Properly – give all figures known and estimate the last one
2.
Determine the number of significant figures in the following measurements.
3.
a)
87 000 000 000
__________
b)
0.000 607 0
__________
c)
320.00
__________
Calculate each of the following. Report answers using significant figures.
a)
(4.15 × 105) m × (6.024 × 1023) m
b)
__________3.56_g__________
3.6 cm × 2.5 cm × 5.2215 cm
c)
18.60 g + 5.20 cg
Density
D = M /V
Calculate the density of an object measuring 21 cm by 6.0 cm by 2.12 cm and
having a mass of 522.2 grams
A piece of granite has a mass of 55 g. When I place it into a graduated cylinder
that has 50.0 mL of water in it the level changes to 70.0 mL. What is the density
of this piece of granite?
% Error
% Error = (Experimental – Actual) ÷ Actual × 100
The density of a sample of water was experimentally determined to be 1.08 g/ml.
What is the % error?
2
Conversions
Given: 1 joule = 0.239 calorie. Convert 630 joules to kilocalories.
Given: 2.24 lb = 1 kg. Convert 18 lbs to kg.
Subatomic Particles and their Jobs
# protons = atomic number (identifies the element)
neutrons define the mass of the atom
# protons + neutrons = mass number
# electrons determines the charge
Fill in the following table:
Subatomic particle
Relative Mass
(amu)
Charge
Location in atom
0
0
1
Fill in the following table:
Symbol
Atomic
Number
Protons
Xe
Neutrons
Electrons
Mass
Number
77
Ba-137
27
33
27
Na+
55
Mn
15
22
25
O-2
Write the complete chemical symbol for the ion with 31 protons, 39 neutrons, and 28
electrons.
How many neutrons are in chromium−52?
3
Iron consists of four natural isotopes:
Isotope Mass (amu) Percent Abundance
54
Fe
53.9696
5.82
56
Fe
55.9349
91.66
57
Fe
56.9354
2.19
58
Fe
57.9333
0.33
Calculate the average atomic mass of iron to five significant digits.
Convert the following
1.
45.6 g Mg to atoms
2.
45.9 moles gold to g
Final Exam Review Day 2
Electron Configuration (shows where all the ______________ are located)
Diagonal rule
1s2
2s2
2p6
2
3s
3p6
3d10
4s2
4p6
4d10 4f14
2
6
5s
5p
5d10 5f14
6s2
6p6
6d10 6f14
2
6
7s
7p
7d10 7f14
Draw the electron configuration for the following:
Potassium
Sulfur
Cobalt
P−3
Mg+2
Valence electrons are electrons in the outer shell.
Determine the number of valence electrons and draw the dot diagrams for the
following:
Potassium
Sulfur
P−3
Mg+2
4
Draw the shortcut electron configuration for each of the following
Potassium
Sulfur
Cobalt
P−3
Mg+2
Identify the following
[Xe]4f 46s2
contains 3 electrons in its sixth and outer main energy level
the element that has 2 electrons in the p sublevel in its second main energy level
4s24p5
Be able to locate s, p, d, and f blocks on the periodic table
The ___________ ____________ _____________ is the same as the period number.
There are ________ main energy levels. The ____________ ___________ _________
is caused when electrons emit energy as they fall back to a lower energy level.
Periodic Table
Know the name and location of the following groups
Group 1
___________________
Group 2
_____________________
Group 3-12
___________________
Group 17
_____________________
Group 18
___________________
The Noble gases are stable because they have ___ ______________ ____________.
Elements are placed in order of increasing ______________ _______________ and
placed in groups according to their ___________________ __________________.
Rows = ______________ =___________ __________ _______
Columns =______________ = _______________
Metals are located _________________________________
Nonmetals are located______________________________
Metalloids are located __________________________________
Metals _____________ electrons. They form __________________.
Nonmetals ______________________ electrons. They form __________.
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Trends
Atomic radius (size)
ionization energy (energy to remove an electron)
Electronegativity (ability to gain an electron)
Reactivity for metal (metallic character)
Reactivity for nonmetal
Ionic Radius
a)
b)
c)
d)
e)
f)
g)
h)
i)
j)
k)
l)
m)
n)
o)
p)
F
K
C
Ba
Si
N
Mn
Mg
Ge
Fe
Be
Ca
Na
Rb
Au
Ca+2
Se
Na
O
B
I
Cs
Cl
S
Br
Sb
F
Mg
I
Sr
Pb
K+1
is more reactive.
is the smaller atom.
has the higher ionization energy.
has only s electrons in its outer shell
has 4 valence electrons
is a metal
last electron is d5
goes to a charge of −2
has higher electronegativity
is a metalloid
is a halogen
is more reactive
is more likely to combine with oxygen
loses electrons more readily
is a transition metal
has large ionic radius
Chemical Bonding
The three types of chemical bonds are
1.
_________________ usually between a ___________ and a ___________.
There is a __________ of electrons. The bond is due to attraction between
_______________ and _______________.
2.
_________________ usually between a ___________ and a ___________. The
electrons are _____________.
a) _____________ _______________ between 2 same nonmetals. Electrons
are ____________ shared.
b) _____________ _______________ between 2 different nonmetals. Electrons
are ____________ shared.
3.
_________________ usually between a ___________ and a ___________.
Many electrons are shared. They are ______________.
Classify the following chemical bonds as ionic, covalent or metallic:
_____ CuO _____AlBr3 _____ H2O _____ PCl5 _____ MgSO4
6
The type of bond is determined according to the difference in ___________________
between the two atoms.
|----|---------------|-------------------|
0 .3
1.7
4
Characteristics of ionic vs covalent. Ionic bonds are strong. Ionic substances have
orderly pattern and a high ____________ ___________. They form ______________.
Lewis Structures
Predict the type of bonding. Draw Lewis structures for molecules. Predict the shape of
the molecule.
CCl4
SiO2
O2
H2O
CO3−2
NH3
What is an intermolecular force? What are the three types of intermolecular forces?
What is the strongest type?
Final Exam Review Day 3
Calculate the oxidation numbers of the atoms in each of the following compounds or
polyatomic ions:
H3AsO4
H2Cr2O7
TiO2
SO3−2
Calculate the molar mass of the following
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CaCl2
(NH4)2CO3
Mn3(PO4)2
Convert each of the following.
a.
5.6 × 1026 molecules NaOH to grams NaOH
b.
45g KClO3 to moles KClO3
c.
7.4 moles Sn(SO4)2 to molecules
d.
3.25 moles (NH4)2CO3 to grams
Find the % of phosphorus in Fe3(PO4)2.
A compound contains 68.04 g N and 155.52 g O. Find its empirical formula.
A compound contains 68.04 g N and 155.52 g O. Its molar mass is 92 g. Find its
molecular formula.
Naming Compounds and writing formulas
Binary Ionic
MgCl2
CuCl2
Tertiary Ionic
MgSO4
Covalent
CO2
Binary acids
HI (aq)
Oxyacids
HNO3 (aq)
HNO2 (aq)
8
Write the formula for each of the following substances:
Strontium hypobromite
Phosphorus trichloride
Sodium chromite
aluminum sulfite
Magnesium nitrite
chloride
Calcium sulfide
iron (II) hydroxide
Sulfuric acid
hydrosulfuric acid
Phosphorus trioxide
hydrogen
Name the following substances
HCl (aq)
CaO
SO2
F2
NO
SO2-2
H3PO4 (aq)
NaClO3
Recognize an acid, a base a salt and other:
MgSO4 _____
C6H12O6 _____
HCl _____
NaOH _____
Chemical Reactions:
Synthesis (element + element → compound)
Mg (s)
Decomposition (compound → element + element)
HgO (s) →
Single replacement (element + cpd→element + cpd)
K + Ni(NO3)2 →
Double replacement (cpd + cpd → cpd + cpd)
Pb(NO3)2 + NaBr →
Combustion (CxHy + O2 → CO2 + H2O)
C3H8 + O2 →
Neutralization (acid + base → salt + water)
LiOH + H2SO4 →
Classify and balance the following.
________
1.
BaCl2 + (NH4)2CO3
→
KCl
________
2.
KClO3
________
3.
Na2O
+
P4O10
________
4.
C6H6
+
O2
________
5.
FeCl2 (aq)
+
→
+
→
→
+
O2 (g)
→
NH4Cl
O2
Na3PO4
CO2
Na3PO4 (aq)
BaCO3
+
+
→
H2O
Fe3(PO4)2 (s)
+
NaCl (aq)
9
Identify the precipitate in question #5.
What are the products in question 5?
What do the symbols (s), (l), (g) and (aq)
mean?
What is the subscript of oxygen in
question #4?
Which substances in question 5 are
solutions?
What is the coefficient of oxygen in
question #4?
What are the reactants in question 5?
Predict whether the following single replacement reactions will occur.
1.
________
Ca
2.
________
NaCl
3.
________
Li
+
AgNO3
I2
→
K2O
→
+
+
→
Predict if the following double replacement reactions will occur.
1.
________
Na2CO3(aq) + Hg(NO3)2(aq) → 2 NaNO3(aq) + HgCO3(s)
2.
________
Ba(NO3)2(aq) + MgCl2(aq) → BaCl2(aq) + Mg(NO3)2(aq)
3.
________
Ca(OH)2(aq)
4.
________
FeS(s) + 2 HCl(aq) → H2S(g) + FeCl2(aq)
+
H2SO4(aq)
→
Stoichiometry
1.
Given the following: C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g)
a)
How many moles of C3H8 are needed to produce 9 moles of CO2?
b)
How many grams of CO2 are produced when 2.0 moles of C3H8 react with
excess oxygen?
c)
How many grams of oxygen are needed to react with 22.0 g C3H8?
d)
How many grams of oxygen are needed to react with 3.48 liters of C3H8
at STP?
2.
3.
Given the following: 2 Na
+
Cl2
→
2 NaCl
(a)
If 23.0 g Na and 71.0 g Cl2 react, what is the theoretical mass of NaCl
produced? Why? What is the limiting reactant? What is the excess
reactant?
b)
An excess of sodium is reacted with 142 g Cl2. If 199 g of NaCl are
actually produced, what is the percent yield?
Aluminum reacts with hydrochloric acid according to the following unbalanced
equation:
2 Al + 6 HCl → 2 AlCl3 + 3 H2
(a)
If 18 g Al are combined with 75 g HCl, which reactant is the limiting
reactant? Which is the excess reactant?
(b)
What mass of AlCl3 is theoretically formed?
(c)
An excess of HCl is reacted with 52 g of Al. If 5.0 g H2 are actually
produced, what is the percent yield?
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Final Exam Review Day 4
Gas Laws
Kinetic Molecular Theory assumes gases are made up of _________ ___________
moving in _____________ ___________, colliding into each other with
______________ collisions. As temperature increases, the particle movement also
_____________________.
Gases do not behave ideally when gases stop moving (or move slowly)
That is at _________ temperatures and __________ pressures.
P1V1
T1
=
P2V2
T2
PV = nRT
Ptot = P1 + P2 + P3 …
STP means ______________ (760 mm Hg) and _______________
1.
A sample of gas occupies 400.0 mL at a pressure of 1 atm. What will the volume
be if the pressure is changed to 2.00 atm while the temperature remains constant?
2.
Calculate the volume occupied by 64.0 grams of O2 at a pressure of 850 torr and a
temperature of 25C.
3.
A gas occupies 250 mL at STP. It expands to 375 mL as it is heated to 75ºC.
What is the new pressure?
4.
A gas sample is collected over water when at a pressure of 760 mm Hg.
What is the pressure of the dry gas if the partial pressure of the water vapor is
11.5 mm Hg?
5.
A 600. mL sample of gas at STP expands to 900. mL at constant pressure. What
is the new temperature?
6.
Calculate the number of grams of N2 in a 6.0 liter cylinder at 27C and 800. torr.
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Gas Stoichiometry
1.
Given:
2 H2(g)
+
O2(g)
→
2 H2O(g)
If this reaction occurs at 200.C and 1500 torr, how many liters of O2 will react
with 20.0 L of H2?
2.
Given:
C3H8(g)
+
5 O2(g)
→
3 CO2(g)
+
4 H2O(g)
How many liters of CO2 at 500.C and 850 torr are produced from 6.60 g C3H8?
Phases and Phase Diagrams
Solids
Liquids
Gases
Definite shape
Definite Volume
Fluidity
Particle movement (KE)
Particle arrangement
Attractive forces
Density
Diffusion rate
Compressibility
13
Heat Calculations
Q = m∙Cp∙∆T
Q = m∙∆H
1.
What quantity of heat, in joules, is required to raise the temperature of
22.5 g of lead from 18.0°C to 30.0°C? The specific heat of lead is 0.129 J/g·°C.
2.
Calculate the quantity of heat energy (to the nearest calorie) required to convert
50.0 g of ice at 0°C to steam at 110°C. Be sure to draw a diagram showing
the temperature and phase changes and the heat energy increases.
Specific heats: water = 1.00 cal/g·°C
steam = 0.480 cal/g·°C
∆Hvap = 540 cal/g; ∆Hfus = 80 cal/g
14
Final Exam Review Day 5
Concentration
% by mass = g solute × 100
g solution
Molarity (M) = moles solute
Liter solution
Molality(m) = moles solute
Kg solvent
1.
Calculate the molarity of 95.0 grams of Al(OH)3 dissolved in 500. mL of solution.
2.
Calculate the % of a solution containing 35 g NaCl in 110 g H2O.
3.
Calculate the % concentration by mass of a solution in which 20.0 g Mg(OH)2 is
dissolved in 80.0 g H2O.
4.
What is the molality of a solution made by dissolving 2.5 g Ca(OH)2 in 250 g of
water?
5.
How many g of NaCl are needed to make 2000. mL of a 2.50 M solution?
6.
Find the freezing point depression and the boiling point elevation of each of the
following ELECTROLYTES in solution.
a.
18.0 g MgCl2 in 250. g water
b.
80.0 NH4 Br in 100. g water
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Solubility
1.
What is the solubility of KCLO3 at 30 ºC?
2.
50 g of KCl are added to 100 g of water at 40°C. Is the solution saturated
or unsaturated?
If saturated, how many more grams of KCl remain undissolved?
If unsaturated, how many more grams of KCl could be dissolved?
3.
How does temperature affect the solubility of a solid in a liquid?
4.
How does temperature affect the solubility of a gas in a liquid?
5.
How does pressure affect the solubility of a solid in a liquid?
6.
How does pressure affect the solubility of a gas in a liquid?
Dilutions: C1V1 = C2V2
How many mL of a 12.0 M HCl solution are needed to make 100 mL of a 1.0 M HCl
solution?
16
Ions in solution
1.
Write an equation showing the dissociation of AlCl3 in water.
2.
For the following write (a) the ionic equation; and (b) the net ionic equation.
formula equation: 2 AlCl3(aq) + 3 Na2CO3(aq) → Al2(CO3)3(s) + 6 NaCl (aq)
ionic equation:
net ionic equation:
What is the precipitate in the equation above?
3.
Write an equation showing the ionization of H3PO4 in water.
4.
Which of the following are electrolytes?
CaBr2
CH4
Ca(OH)2
Na3PO4
HF
Acids and Bases
Be able to identify a neutralization reaction.
1.
Calculate the Normality of: Normality (N) = M × Subscript of H+ or OH−
a.
0.45 M H2SO4
b.
0.82 M Ca(OH)2
17
2.
pH = − log (N of acid)
Calculate the pH of :
a. 0.0052 M HCl
3.
b.
0.001 N HNO3
In an aqueous solution if H3O+ = 1  10 −5, then OH− = ________,
pH=
______, and pOH =______.
Know pH scale
pH + pOH = 14
|-----------------|-----------------|
acid
←7 →
base
Acids donate a proton. Bases accept a proton.
Show a single step of the ionization of the following acids in water. Label the
conjugate acid-base pairs.
(a) HCrO2-1
+
H2O
→
(b) HBr
+
H2O
→
Indicate each of the following as an acid, base, or salt:
H2SO4
KBr
HNO3
Ba(OH)2
Na2SO4
Properties of acid
1.
Properties of base
1.
2.
2.
3.
3.
4.
4.
5.
5.
Titration:
NAVA = NBVB
How many mL of 0.15 N KOH are needed to neutralize 50.0 mL 0.25 N H2S, if both
completely break into ions?
18
Final Exam Review Day 6
Equilibrium
Keq= Products
reactants
1.
A(g)
Given:
+
2 B(g)
→
AB2(g)
+
heat
Write the equation for the equilibrium constant, Keq.
If at equilibrium the concentrations are: [A] = 1.2 × 10−2 M; [B] = 2.0 × 10−2 M;
[AB2] = 1.5 × 10−5 M ; calculate the value of the equilibrium constant.
Is the reaction endothermic or exothermic?
2.
3.
For the same equation given above, predict the effect of the following (get more
products, get more reactants, or no effect on the amount of products or reactants)
Adding more B
Removing some A
Heating
decreasing the pressure
Adding a catalyst
removing AB2
Ksp = products
The higher the Ksp the greater the solubility.
For the reaction:
Al(OH)3 (s) → Al+3 (aq) + 3 OH−1(aq)
Write the expression for the solubility product, Ksp.
4.
Name 4 ways to increase reaction rate.
5.
What must you have in order for a reaction to occur?
19
Heat in reactions
Refer to the following diagram for the hypothetical reactions:
70
E
n
e (kJ) 40
r
20
g
y
70
E
60
n
e (kJ)
r
20
g
y
Time
1.
Time
Is the reaction endothermic or exothermic?
(a)
2.
(b)
What is the value for the activation energy for the forward reaction?
(a)
3.
(b)
What is the value of ∆H for the reaction?
(a)
(b)
Classify the following processes as endothermic or exothermic.
1.
__________ clothes dry
2.
__________ condense gas
3.
__________ a bomb explodes
4.
__________ H2O(g)
+
____________∆H = -314.43 kJ
5.
C(s)
+
131.3 kJ
→
CO(g)
+
H2(g)
Oxidation Numbers
Assign oxidation numbers to the elements in each substance
H2SO3
HNO3
P2O5
MnO4−2
20