Download 2nd Semester Exam Review

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Semester Exam Review
Chapter 9
• Reactants and Products
• NaHCO3 + HC2H3O2  CO2 + H20 + NaC2H3O2
Symbol
Purpose
Separates two or more reactants or products
Seperates reactants from products
Separates reactants from products and indicates a
reversible reactions
Solid state
Liquid state
Gas state
Water solution (aqueous state)
Balancing Chemical Equations
• Law of Conservation of Mass: mass of
reactants must = mass of products
• If I start with 12 g of Carbon I have to end with
12 g of Carbon
• Each side of the equations must contain the
same number (mass) of atoms
Mg +
O2 
Cu(SO4) +
MgO
Al 
Al2(SO4)3 + Cu
Synthesis Reactions
• 2 or more reactants join to form a single
product
•
•
H2(g) + O2(g) 
CaO(s) +
H2O(l) 
H2O(l)
Ca(OH)2(s)
Decomposition Reaction
• A single reactant breaks apart
• H2O(l) 
H2(g) +
O2(g)
• Aluminum oxide decomposes when electricity
is passed through it
Combustion
• Occurs when one of the reactants is oxygen
• CH4(g)+ O2(g)
CO2(g) + H2O(l)
•
SO3
SO2 + O2 
Single Replacement Reactions
• One element replaces the atoms of another
element in a compound
• CuCl2(aq) +
Al(s) 
AlCl3(aq)+ Cu(s)
Double Replacement Reactions
• Involves an exchange of ions between two
compounds
• Pb(NO3)2(aq) + 2KI(aq)  PbI2(s)+
2K(NO3)(aq)
• NaCl(aq) + Ag(NO3)(aq)  AgCl(s) + NaNO3
(aq)
Predict products and balance
• Li +
Cl2 
• HgO 
•
•
K+
H2 +
CaO 
KCl 
• Pb(NO3)2 +
KI 
Chapter 10
• Conversion factors:
– 1 mol =
particles
– 1 mol =
grams
– 1 mol =
Liters
• How many moles are in 12.5 g of Carbon
dioxide?
• How many molecules are in 7.25 mol of
carbon dioxide?
• How many atoms are in a 44.3 g piece of iron?
Calculating % Composition
1. Find the total molar mass of each element in
the compound.
2. Find the molar mass of the entire
compound.
3. Divide the total molar mass of each element
by the molar mass of the compound then
multiply by 100
4. Check that all your percentages add up to
100
• What is the percent composition of phosphate
in phosphoric acid?
Calculating Empirical Formulas
1. Change % to g (assume 100 g of compound
so 30% = 30 g)
2. Convert each element from g to moles
3. Divide each mole amount by the smallest
number from step 2
4. Change to a whole number = subscript in
empirical formula
• What is the empirical formula for a compound
that is 48.64% Carbon, 8.16% Hydrogen, and
43.20% Oxygen?
Calculating Molecular Formula
1. Calculate the mass of the empirical formula
2. Divide molar mass of compound (will be
given) by mass of empirical formula
3. Multiply each subscript in the empirical
formula by the answer from step 2
• A compound who’s empirical formula is NO2
has a molar mass of 92 g/mol. What is the
compound’s molecular formula?
Al + O2  Al2O3
• How many moles of Al are needed for 3.7 mol
Al2O3?
• How many moles of O2 needed to react with
14.8 mol Al?
Al + O2  Al2O3
• How many grams of Al2O3 can be produced
from 80.0 g of O2?
• How many grams of O2 will react with 79.4 g
of Al?
H 2O  H 2 + O 2
• How many molecules of oxygen from 29.2 g
water?
• How many grams of hydrogen from 4.32 x 1024
molecules of water?
Limiting and Excess Reactants
• Limiting = reactant used up first (you run out
of it)
– Determines the amount of product that can be
formed
• Excess = the other reactant (not used up)
Determining Limiting Reactant
1. Perform conversion for each reactant =
amount needed
2. Compare amount needed with amount
present
3. Whichever doesn’t have enough = limiting
Na + Cl2  NaCl
• You have 6.7 mole of sodium and 3.20 mole of
Chlorine
• Which is limiting?
• How much NaCl can be made?
Chapter 12
• Standard pressure =
– 1 atm
– 101.3 kPa
– 760 mmHg
– 760 torr
Kinetic Molecular Theory
• Describes the behavior of matter in terms of
particles in motion
– Makes several assumptions about the size,
motion, and energy of gas particles
Assumptions of the Kinetic Molecular
Theory
• Gasses consist of small particles that take up
little volume relative to the volume of empty
space around them
– Gas molecules are very far apart and therefore
don’t experience attractive or repulsive forces.
• Gas particles move in constant, random
straight lines until they collide with other
particles or with the walls of the container
– Collisions are elastic -
• The energy of gas particles is determined by
the particle’s mass and velocity
– KE =
Intermolecular forces
• Inter- means between or among
• Intermolecular forces can hold together
identical particles or two different types of
particles
• Weaker than intramolecular forces (bonds)
Dispersion Forces
• Weak forces that result from temporary shifts
in the density of electrons in electron clouds
Dipole-dipole forces
• Attraction between oppositely charged
regions of polar molecules
– Polar molecule =
• Neighboring polar molecules orient
themselves so that oppositely charged regions
align
Hydrogen Bonds
• Dipole-dipole attraction that occurs between
molecules containing a hydrogen atom
bonded to a flourine, oxygen, or nitrogen
atom
• Viscosity - measure of the resistance of a
liquid to flow
– Attractive forces – stronger intermolecular forces
= higher viscosity
– Particle size – larger molecules = higher viscosity
– Temperature – lower temperature = higher
viscosity
• Surface tension – the energy required to
increase the surface area of a liquid by a given
amount
– Caused by intermolecular forces pulling down on
the particles on the surface of a liquid which
stretches it tight like a drum
Phase changes that require energy
• Melting
– Heat flows from an object at a higher temperature
to an object at a lower temperature
– Ice absorbs heat which does not raise
temperature but is used to break hydrogen bonds
– When hydrogen bonds are broken molecules can
move further apart into the liquid phase
• Vaporization – process by which liquid
changes to vapor
– Vapor – gaseous state of a substance that is
normally liquid at room temperature
– Evaporation – when vaporization occurs only at
the surface of a liquid
– Vapor pressure – the pressure exerted by a vapor
over a liquid
• Sublimation – changing from solid to gas
without becoming a liquid
– Dry ice
– Moth balls
– Solid air fresheners
Phase changes that release energy
• Freezing
– Heat flows out of warmer object into cooler object
– Molecules slow down & become less likely to flow
past one another
– Intermolecular forces cause the molecules to
become fixed into set positions
– Freezing point – temperature in which a liquid
becomes a solid
• Condensation – process by which a gas or
vapor becomes a liquid
• Deposition – substance changes from gas or
vapor to solid without first becoming a liquid
– frost
• Triple point = point at which all three phases
exist at the same time
Chapter 13
•
•
•
•
•
Boyle’s Law =
Charle’s Law =
Gay-Lussac’s Law =
Combined gas Law =
Ideal gas law =
– Which value of R to use
• For all laws temperature must be in K!
• STP
• Molar volume of a gas is the volume that 1
mol occupies at STP
– STP = 0 oC & 1 atm or 101.3 kPa
– 1 mol = 22.4 L of gas
• How many moles of carbon dioxide gas are in
a 1.0 L balloon?
• What volume will 9.22 g of H2 gas occupy at
STP?
• The volume of a gas at 99.0 kPa is 300.0 mL. if
the pressure is increased to 188 kPa, what will
be the new volume?
• A helium balloon in a closed car occupies a
volume of 2.32 L at 40.0 oC. If the car is parked
on a hot day and the temperature inside rises
to 75.0 oC what is the new volume of the
balloon?
• The pressure of the oxygen gas inside a
canister is 5.00 atm at 25.0 oC. The canister is
located at a camp high on Mount Everest. If
the temperature there falls to -10.0 oC, what is
the new pressure inside the canister?
• A gas at 110 kPa and 30.0oC fills a flexible
container with an initial volume of 2.00 L. If
the temperature is raised to 80.0 oC and the
pressure increases to 440 kPa, what is the new
volume?
• Calculate the number of moles of ammonia
gas contained in a 2.0L vessel at 300.0 K with a
pressure of 1.50 atm
Chapter 14
Heterogeneous mixtures
• Suspensions –
– Particle size:
• Colloids –
– Particle size:
– Brownian motion:
– Tyndall effect:
Homogeneous Mixtures
• Solutions
– Solute
– Solvent
– Can be solid, liquid, or gas
• Soluble • Insoluble • Miscible • Immiscible -
• Molarity (M)
• What is the molarity of a solution containing 3
moles of solute in 1.5 L of solution?
• What is the molarity of 155 mL of solution
containing 1.55 g dissolved KBr?
• Diluting molar solutions:
• What volume of 2.00M CaCl2 stock solution
would you use to make 0.50L of 0.300M CaCl2
solution?
• Preparing molar solutions
• How many grams of CaCl2 would be dissolved
in 1.0 L of a 0.10 M solution of CaCl2?
Ways to increase solvation
• Increase collisions between solute and solvent
particles
• Agitation
• Increase surface area
• Increase temperature
Solubility
• Unsaturated solutions
• Saturated solutions
• Supersaturated solutions
Colligative Properties of Solutions
• Colligative properties depend on number of
solute particles in a solution
• Vapor pressure lowering –
– Vapor pressure:
– Adding a nonvolatile solute to a solution lowers
the solvent’s vapor pressure
• Boiling point elevation –
– Boiling point
– Because vapor pressure is lowered, boiling point
increases
• Freezing point depression – solute particles
interfere with attractive forces holding solvent
particles together
Chapter 15 Review Problem 1
• A 225 g sample of iron at 98.5 oC is placed into
72.4 g of water at 22.0 oC. The final
temperature of the mixture is 41.2 oC. The
specific heat of water is 4.18 J/g oC, what is
the specific heat of iron?
Review problem 2
• What is the
Hrxn for the reaction:
• P4(s) + 6H2O(l)  4H3PO4(l)
Review problem 3
• What amount of heat is released when 252 g
of tin at 112 oC cools to 37.5 oC. The specific
heat of tin is 0.226 J/g oC
• How much heat is released when 48.0 g of
glucose is burned? ΔHcomb = -2808 kJ
• What is the energy change for the following?
H2O2(l)  2H2O(l) + O2(g)
a. 2H2(g) + O2(g)  2H2O(l) ΔH = -572kJ
b. H2(g) + O2(g)  H2O2(l) ΔH = -188kJ
Law of conservation of energy
• 1st law of thermodynamics – energy cannot be
created or destroyed, only transferred
Heat
• Heat = energy
• Symbol = q
• Always flows from warmer object to cooler
object
Measuring Heat
• calorie – the amount of energy needed to
raise the temperature of one gram of pure
water by one degree Celsius
• Joule – the SI unit of energy
• ΔHrxn = Hproducts – Hreactants
• Exothermic reaction – negative ΔH
• Endothermic reaction – positive ΔH
Standard heat of formation
• The change in enthalpy that accompanies the
formation of one mole of the compound in its
standard state
• ΔHf
• ΔHf of an element = 0
Entropy
• Entropy(S) – a measure of the number of
possible was that the energy of a system can
be distributed
– Determined by the freedom of the systems
particles to move and the number of ways they
can be arranged
– Disorder or randomness of a system
• Second law of thermodynamics – spontaneous
processes always proceed in such a way that
the entropy of the universe increases
Changes resulting in decreased
entropy
• Phase changes that decrease molecule
movement
– (g)  (l)
– (l)  (s)
• Number of particles decreases in a reaction
• Dissolving of gas in a solvent
• Decrease in temperature
• Predict the sign of ΔS for the following:
– ClF(g) + F2(g)  ClF3(g)
– NH3(g)  NH3(aq)
– CH3OH(l) CH3OH(aq)
– C10H8(l)  C10H8(s)