Download honors chem 6 day review packet

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
2. ____ Na2CO3 (aq)
4. ____ helium
6. ____ H2O (l)
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.
__________ burning gas
3.
__________ decompose water into hydrogen and oxygen
4.
__________ evaporate water
Units of measurement
cm3, L or mL is __________________
g is _______________________
g/cm3, g/l or g/mL are ________________
joules , calories, kcal is ______________________
K or ºC 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)
18.63 g + 5.2 g
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?
Conversions
Given: 1 joule = 0.239 calorie. Convert 630 joules to calories.
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.
Symbol
Atomic
Number
Protons
Xe
Neutrons
Electrons
Mass
Number
77
Ba-137
27
33
27
Na+
55
22
Mn
25
17
15
18
36
O-2
Write the complete chemical symbol for the ion with 31 protons, 39 neutrons, and
28 electrons.
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 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
3.
78.3 g Pb to atoms
Final Exam Review Day 2
Name_________________
Electron Configuration (shows where all the ______________ are located)
Diagonal rule
1s2
2s2
2p6
3s2
3p6
3d10
2
6
4s
4p
4d10 4f14
5s2
5p6
5d10 5f14
2
6
6s
6p
6d10 6f14
7s2
7p6
7d10 7f14
Draw the electron configuration for the following:
Potassium
Sulfur
Cobalt
P−3
Mg+2
Valence electrons are electrons in the outer shell.
Draw the dot diagrams for the following:
Potassium
P−3
Sulfur
Mg+2
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 3-13
Group 18
___________________
___________________
___________________
Group 2
Group 17
_____________________
_____________________
The Noble gases are stable because they have ___ ______________ ____________.
Elements are placed in order of ______________ _______________ 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 __________.
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
a)
b)
c)
d)
e)
f)
g)
h)
i)
j)
k)
l)
m)
n)
o)
F
K
C
Ba
Si
N
Mn
Mg
Ge
Fe
Be
Ca
Na
Rb
Au
Se
Na
O
B
I
Cs
Cl
S
Br
Sb
F
Mg
I
Sr
Pb
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
Chemical Bonding
The three types of chemical bonds are
1.
_________________ usually between a ___________ and a ___________.
There is a __________ of electrons.
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 ______________.
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 or polyatomic ion.
Predict the shape names. Predict the intermolecular force for the molecules.
CCl4
SiO2
BaCl2
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
Name_________________
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
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 % composition by mass of each element 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 82 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)
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
F2
NO
H3PO4 (aq)
NaClO3
SO2
Recognize an acid, a base a salt and other:
SO2−2
MgSO4 _____
HCl _____
C6H12O6 _____
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 →
+
O2 (g)
→
Classify and balance the following.
________
1.
BaCl2 + (NH4)2CO3
→
KCl
________
2.
KClO3
________
3.
Na2O
+
P4O10
________
4.
C6H6
+
O2
________
5.
FeCl2 (aq)
+
+
→
→
→
BaCO3
+
NH4Cl
O2
Na3PO4
CO2
Na3PO4 (aq)
+
H2O
→
Fe3(PO4)2 (s)
+
NaCl (aq)
3 CO2(g)
+
4 H2O(g)
What do the symbols (s), (l), (g) and (aq) mean?
Which substances in question 5 are solutions?
What are the reactants in question 5?
What are the products in question 5?
What is the subscript of oxygen in question #4?
What is the coefficient of oxygen in question #4?
Stoichiometry
1.
Given the following: C3H8(g)
+
5 O2(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?
(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?
Final Exam Review Day 4
Name_________________
Gas Laws
Kinetic Molecular Theory assumes gases are made up of _________ ___________
moving in _____________ ___________, colliding into each other with
______________ collisions.
Ideal vs Real : Non-ideal behavior occurs 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 1 atm (760 mm Hg) and 273 K
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 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.
Calculate the number of grams of N2 in a 6.0 liter cylinder at 27C and 800 torr.
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
Definite shape
Definite Volume
Fluidity
Particle movement (KE)
Particle arrangement
Attractive forces
Density
Diffusion rate
Compressibility
Liquids
Gases
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°C to 30°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
Final Exam Review Day 5
Concentration
% by mass = g solute × 100
g solution
Name_________________
Molarity (M) = moles solutes
Liter solution
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 g Mg(OH)2 is
dissolved in 80 g H2O.
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?
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?
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:
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
2.
Calculate the pH of a 0.0052 M solution of HCl. pH = − log (N of acid)
3.
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.
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:
N1V1 = N2V2
How many mL of 0.15 N KOH are needed to neutralize 50.0 mL N H2S, if both
completely break into ions?
Final Exam Review Day 6
Name____________________
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?
Heat in reactions
Refer to the following diagram for the hypothetical reactions:
E
n
e
r
g
y
E
n
e
r
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)
(b)
What is the value of ∆H for the reaction?
3.
(a)
(b)
Oxidation and Reduction
Oxidation = _______ of electrons
Reduction = _______ of electrons
The oxidizing agent caused the oxidation. (It was reduced)
The reducing agent caused the reduction. (It was oxidized)
KNO3
+
CO
→
CO2
+
NO2
+
K2O
Assign oxidation numbers to the elements in each substance.
___________ was oxidized.
___________ was reduced.
___________ was the reducing agent.
___________ was the oxidizing agent.