Periodic Table and the Atom Answers
Complete the table. There is enough information given for each element to determine all missing numbers.
Symbol
23
Atomic
Number
Mass
Number
Number
of
Protons
Number
of
Electrons
Number
of
Neutrons
11
23
11
11
12
Na
19
K
19
38
40
90
9
19
Sr
21
19
38
9
38
9
F
10
21
20
Ca2+
20
Sn
50
53
131
12
26
131
26
41
22
52
18
50
53
50
53
72
78
12
12
14
I
Mg
Review:
1. An element's or isotope's atomic number tells you ___ how many protons are in its atoms.______.
2. An element's or isotope's mass number tells you __ how many protons and neutrons in its atoms.__.
3. The heaviest part of an atom is the __nucleus________, which contains both ____protons____ and
___neutrons________. The _____electrons________ are found in a cloud surrounding the nucleus.
4. If an atom was a penny inside a football stadium, the penny would represent the
______nucleus_______ and the football stadium would represent __the atom or electron cloud____.
5. Which has a higher atomic number?
Helium or Hydrogen ____Helium_____________
Magnesium or Manganese _____Manganese________
6. Which has a lower atomic mass?
Carbon or Calcium _____Calcium____________
Xenon or Radon ________Radon_______________.
7. Generally speaking, how does atomic mass change throughout the periodic table? ___atomic mass______
___increases as you go across and down on the periodic table._____________
The Groups and Periodic Trends Practice Problem Answers
Identify the following elements as metal, nonmetals, or metalloid.
1) Boron
2) Carbon
3) Gold
4) Lead
5) Hydrogen
Metalloid
Nonmetal
Metal
Metal
Nonmetal
Identify the following elements by which group they belong to on the periodic table.
6) Flourine
7) Argon
8) Calcium
9) Potassium
10) Carbon
halogen
noble gas
alkaline earth metal
alkali metal
nonmetal
11) Which of these elements has the largest atomic radius?
a) aluminum
b) calcium
c) fluorine
d) potassium
e) sulfur
12) Which of these elements has the smallest atomic radius?
a) potassium
b) iron
c) arsenic
d) bromine
e) krypton
13) Which of these elements has the highest first ionization energy?
a) oxygen
b) oxygen
c) fluorine
d) carbon
e) boron
14) Which of these elements has the highest electronegativity?
a) lithium
b) nitrogen
c) potassium
d) arsenic
e) beryllium
Lewis Dot Structure Practice Problem Answers
1)
PBr3
2)
N2H2
3)
CH3OH
4)
NO2-1
5)
C2H4
6)
BSF
7)
HBr
8)
C2H5OH (ethanol)
9)
N2F4
10)
SF6
Chemical Bonding Practice Problem Answers
1) How are ionic bonds and covalent bonds different?
Ionic bonds result from the transfer of electrons from one atom to another; Covalent bonds result from two
atoms sharing electrons.
2) Describe the relationship between the length of a bond and the strength of that bond.
Strength of a bond increases as the bond gets shorter (inverse relationship)
3) Identify the type(s) of bond(s) found in the following molecules:
a. CCl4
___covalent________________________
b. Li2O
___ionic________________________
c. NF3
___covalent________________________
d. CaSO4
___ionic and covalent________________________
e. SO2
___covalent________________________
f.
Mg(OH)2 ___ionic and covalent________________________
4) Determine if the bond between atoms in each example below is nonpolar covalent, polar covalent, or
ionic.
a. H2
___npc____________
b. PCl
___pc_____________
c. F2
___npc____________
d. NaBr
___ionic__________
e. NF
__pc______________
f. MgO
__ionic____________
g. CH
__npc_____________
h. HCl
__pc______________
5) Proteins are large biological molecules. What type of bonds do they form? covalent
6) Carbohydrates are large biological molecules. What type of bonds do they form? covalent
7) Lipids are large biological molecules. What type of bonds do they form? covalent
8) Sugars are large biological molecules. What type of bonds do they form? covalent
Types of Intermolecular Forces Practice Problem Solutions
What is the strongest intermolecular force present for each of the following compounds?
1)
water
hydrogen bonding
2)
carbon tetrachloride
3)
ammonia
hydrogen bonding
4)
carbon dioxide
London dispersion forces
5)
phosphorus trichloride
dipole-dipole forces
6)
nitrogen
London dispersion forces
7)
ethane (C2H6)
8)
acetone (CH2O)
dipole-dipole forces
9)
methanol (CH3OH)
hydrogen bonding
10)
borane (BH3)
London dispersion forces
London dispersion forces
dipole-dipole forces
For each of the following compounds, determine the main intermolecular force. You may find it useful to
draw Lewis structures for some of these molecules:
11)
nitrogen – Van der Waals forces
12)
carbon tetrachloride – Van der Waals forces
13)
H2S – dipole-dipole forces
14)
sulfur monoxide – dipole-dipole forces
15)
N2H2 – hydrogen bonding
16)
boron trihydride – Van der Waals forces
17)
CH4O – hydrogen bonding
18)
SiH2O – dipole-dipole forces
Balancing Chemical Equations – Answer Key
Balance the equations below:
1)
1 N2 + 3 H2  2 NH3
2)
2 KClO3  2 KCl + 3 O2
3)
2 NaCl + 1 F2  2 NaF + 1 Cl2
4)
2 H2 + 1 O2  2 H2O
5)
1 Pb(OH)2 + 2 HCl  2 H2O + 1 PbCl2
6)
2 AlBr3 + 3 K2SO4  6 KBr + 1 Al2(SO4)3
7)
1 CH4 + 2 O2  1 CO2 + 2 H2O
8)
1 C3H8 + 5 O2  3 CO2 + 4 H2O
9)
2 C8H18 + 25 O2  16 CO2 + 18 H2O
10)
1 FeCl3 + 3 NaOH  1 Fe(OH)3 + 3 NaCl
11)
4 P + 5 O2  2 P2O5
12)
2 Na + 2 H2O  2 NaOH + 1 H2
13)
2 Ag2O  4 Ag + 1 O2
14)
1 S8 + 12 O2  8 SO3
15)
6 CO2 + 6 H2O  1 C6H12O6 + 6 O2
16)
1 K + 1 MgBr  1 KBr + 1 Mg
17)
2 HCl + 1 CaCO3  1 CaCl2 + 1 H2O + 1 CO2
18)
1 HNO3 + 1 NaHCO3  1 NaNO3 + 1 H2O + 1 CO2
19)
2 H2O + 1 O2  2 H2O2
20)
2 NaBr + 1 CaF2  2 NaF + 1 CaBr2
21)
1 H2SO4 + 2 NaNO2  2 HNO2 + 1 Na2SO4
Mole Problems Answers
Question 1
How many moles of copper are in 6,000,000 atoms of copper? 9.96 x 1019 moles of copper
Question 2
How many atoms are in 5 moles of silver? 3.01 x 1024 atoms of silver
Question 3
How many atoms of gold are in 1 gram of gold? 3.06 x 1021 atoms of gold
Question 4
How many moles of sulfur are in 53.7 grams of sulfur? 1.67 moles of sulfur
Question 5
How many grams is a sample containing 2.71 x 1024 atoms of iron? 251.33 grams of iron.
Question 6
How many moles of lithium (Li) are in 1 mole of lithium hydride (LiH)? 1 mole of lithium
Question 7
How many moles of oxygen (O) are in 1 mole of calcium carbonate (CaCO3)? 3 moles of oxygen
Question 8
How many atoms of hydrogen are in 1 mole of water (H20)? 1.20 x 1024 atoms of hydrogen
Question 9
How many atoms of oxygen are in 2 moles of O2? 2.41 x 1024 atoms of oxygen
Question 10
How many moles of oxygen are in 2.71 x 1025 molecules of carbon dioxide (CO2) 90 moles
Question 11
Predict the mass of a mole of magnesium atoms. 24.3 grams
Question 12
Calculate the molecular weights of carbon dioxide (CO2) and sugar (C12H22O11) and the mass of a mole of
each compound. 44 g CO2, 324 g C12H22O11
Question 13
Describe the difference between the mass of a mole of oxygen atoms and a mole of O2 molecules. The mass
of a mole of oxygen (O) is only 16 grams, while the mass of the diatomic oxygen (O2) is twice that, or 32
grams because each molecule contains two oxygen atoms.
Question 14 Calculate the mass in grams of a single 12-C atom. 12 grams
Moles, Molecules, and Grams Worksheet – Answer Key
1)
How many molecules are there in 24 grams of FeF3? 1.28 x 1023 molecules
2)
How many molecules are there in 450 grams of Na2SO4? 1.91 x 1024 molecules
3)
How many grams are there in 2.3 x 1024 atoms of silver? 421 grams
4)
How many grams are there in 7.4 x 1023 molecules of AgNO3? 209 grams
5)
How many grams are there in 7.5 x 1023 molecules of H2SO4? 122 grams
6)
How many molecules are there in 122 grams of Cu(NO3)2? 3.92 x 1023 molecules
7)
How many grams are there in 9.4 x 1025 molecules of H2? 312 grams
8)
How many molecules are there in 230 grams of CoCl2? 1.07 x 1024 molecules
9)
How many molecules are there in 2.3 grams of NH4SO2? 1.69 x 1022 molecules
10)
How many grams are there in 3.3 x 1023 molecules of N2I6? 430 grams
11)
How many molecules are there in 200 grams of CCl4? 7.82 x 1023 molecules
12)
How many grams are there in 1 x 1024 molecules of BCl3? 195 grams
13)
How many grams are there in 4.5 x 1022 molecules of Ba(NO2)2? 17.1 grams
14)
How many moles are in 15 grams of lithium? 0.46 moles
15)
How many grams are in 2.4 moles of sulfur? 77.0 grams
16)
How many moles are in 22 grams of argon? 0.55 moles
17)
How many grams are in 88.1 moles of magnesium? 2141 grams
18)
How many moles are in 2.3 grams of phosphorus? 0.074 moles
19)
How many grams are in 11.9 moles of chromium? 618.8 grams
20)
How many moles are in 9.8 grams of calcium? 0.24 moles
21)
How many grams are in 238 moles of arsenic? 17,826 grams
Solutions for the Stoichiometry Practice Worksheet:
For both of the problems on this worksheet, the method for solving them can be found elsewhere in the “Mr.
Guch’s Helpdesk” section of my website (http://www.chemfiesta.com). If you’re having problems with
stoichiometry problems, I would highly suggest consulting this section of the site before answering these
questions.
When doing stoichiometry problems, people are frequently worried by statements such as “if you have an
excess of (compound X)”. This statement shouldn’t worry you… what it really means is that this isn’t a
limiting reagent problem, so you can totally ignore whatever reagent you have an excess of. Don’t even give
it a second thought, because if you do, you’ll run into trouble.
1)
355.3 grams of Na2SO4
2)
313.6 grams of LiNO3
3) Write
the balanced equation for the reaction of acetic acid with aluminum hydroxide to form water
and aluminum acetate:
3 C2H3O2H + Al(OH)3  Al(C2H3O2)3 + 3 H2O
4)
Using the equation from problem #1, determine the mass of aluminum acetate that can be made if I
do this reaction with 125 grams of acetic acid and 275 grams of aluminum hydroxide.
Two calculations are required. One determines the quantity of aluminum acetate that can be
made with 125 grams of acetic acid and the other determines the quantity of aluminum acetate
that can be made using 275 grams of aluminum hydroxide. The smaller of these two answers is
correct, and the reagent that leads to this answer is the limiting reagent. Both calculations are
shown below – the correct answer is circled.
5)
What is the limiting reagent in problem #2?
Acetic acid.
6)
How much of the excess reagent will be left over after the reaction is complete?
Pressure and Diffusion Practice Problem Answers
1) Simple diffusion is defined as
a) molecules from areas of higher concentration to areas of lower concentration.
b) molecules from areas of lower concentration to areas of higher concentration.
c) water molecules across a membrane.
d) gas molecules across a membrane.
e) water or gas molecules across a membrane.
2) When sugar is mixed with water, equilibrium is reached when
a) molecules of sugar stop moving.
b) water and sugar molecules are moving at the same speed.
c) the dissolved sugar molecules are evenly distributed throughout the solution.
d) there are the same number of water molecules as dissolved sugar molecules.
e) two tablespoons of coffee are added.
3) The rate of diffusion is affected by which of the following?
a) temperature
b) size of molecules
c) steepness of the concentration gradient
d) A and B
e) A, B, and C
4) The molecules in a solid lump of sugar do not move.
a) True
b) False
5) Which of the following is the best explanation of why a decrease in volume causes an increase in
pressure?
a) At a smaller volume the atoms will move faster and hit the sides more often.
b) At a smaller volume the atoms will slow down and so they will have more contact with the walls
of the container.
c) At a smaller volume the atoms will have less room to move around, so they will collide with
the sides more often.
d) The initial statement is false. Gas pressures do not increase when the volume is decreased.
6) What are the five assumptions we make about an ideal gas?
1.
All gas particles are in constant, random motion.
2. All collisions between gas particles are perfectly elastic (meaning that the kinetic energy of the
system is conserved).
3. The volume of the gas molecules in a gas is negligible.
4. Gases have no intermolecular attractive or repulsive forces.
5. The average kinetic energy of the gas is directly proportional to its Kelvin temperature and is
the same for all gases at a specified temperature.
Gas Laws Practice Problem Answers
Use Boyles’ Law to answer the following questions:
1)
1.00 L of a gas at standard temperature and pressure is compressed to 473 mL. What is the new
pressure of the gas? 2.11 atm
2)
In a thermonuclear device, the pressure of 0.050 liters of gas within the bomb casing reaches 4.0 x
106 atm. When the bomb casing is destroyed by the explosion, the gas is released into the
atmosphere where it reaches a pressure of 1.00 atm. What is the volume of the gas after the
explosion?
2.0 x 105 L
3)
The temperature inside my refrigerator is about 40 Celsius. If I place a balloon in my fridge that
initially has a temperature of 220 C and a volume of 0.5 liters, what will be the volume of the balloon
when it is fully cooled by my refrigerator? 0.47 L
4)
A man heats a balloon in the oven. If the balloon initially has a volume of 0.4 liters and a
temperature of 20 0C, what will the volume of the balloon be after he heats it to a temperature of 250
0
C? 0.71 L
5)
If I have 4 moles of a gas at a pressure of 5.6 atm and a volume of 12 liters, what is the temperature?
205 K
6)
If I have an unknown quantity of gas at a pressure of 1.2 atm, a volume of 31 liters, and a temperature
of 87 0C, how many moles of gas do I have?
1.26 moles
Convert the following temperatures into the unit required. 7) 100 K into C
8) 323 K into C
50 C
9) 100 C into K
373 K
10) 25 C into K
298 K
-173 C
pH Scale Practice Problem Answers
1) Five solutions A, B, C, D, E when tested with universal indicator showed a pH of 4, 1, 11, 7, and 9
respectively.
a) Which solution is (i) neutral, (ii) strongly alkaline, (iii) strongly acidic, (iv) weakly acidic, and (v) weakly
alkaline?
b) Arrange the pH in increasing order of hydrogen ion comcentration.
2) Define the term "pH"; what does" pH" stand for?
Answer: The term "pH" is defined as the negative logarithm of H+ ion concentration of a given solution;
the concentration being expressed as moles per litre. Mathematically pH = - log [H+] 'pH' stands for: Power
of hydrogen ion concentration, 'p' for power and 'H' for H+ ion concentration.
3) What is 'pH' scale? Explain briefly.
Answer: The strength of an acid or a base is expressed in terms of hydronium ion concentration. This is
expressed on a scale known as 'pH' scale. It is a 14 point scale; i.e., it has values ranging from 0 to 14,
indicating the value of negative logs of H+ ion concentration of the solution. Some important benchmark
values in the pH scale are: pH = 7 indicates neutral solutions e.g., aqueous solutions. pH > 7 to 14 indicates
alkaline solutions and pH < 7 to 0 indicate acidic solutions
4) What is the 'pH' of pure water and that of rain water? Explain the difference.
Answer: The pH of pure water is seven. Rain water is slightly acidic because as rain drop fall, the carbon
dioxide in the air dissolves with drops to form very weak carbonic acid. Accordingly, rain water has a pH
that is slightly below 7.
5) What is the pH of solution 'A' which liberates CO2 gas with a carbonate salt? Give the reason?
Answer: The pH of solution 'A' is lesser than 7. Carbonates salts react with acids (A) to liberate CO2 gas.
6) What is the pH of solution 'B' which liberates NH3 gas with an ammonium salt? Give reason?
Answer: The pH of solution 'B' is lesser than 7 because 'B' is an alkali as it liberates NH3 gas.
7) How do you increase or decrease the pH of pure water?
Answer: By adding a few drops of alkali to pure water, it's pH increases; and by adding a few drops of an
acid decreases the pH of pure water.
8) What are indicators?
Answer: Indicators are chemicals that show whether the given solution is acidic or basic, by the sudden
change of color.
Types of Acids and Bases Practice Problem Answers
1) What is the pH of the solution with a hydronium concentration [H3O+] 1.47 x 10-4?
What is the pOH of this solution?
pH = 3.83
pOH = 10.17
2) What is the pOH of the solution with a hydroxyl concentration [OH-] 2.98 x 10-2?
What is the hydronium concentration [H3O+] of this solution?
pOH = 1.53 hydronium concentration [H3O+] = 3.37 x 10-13 OR pH = 12.47
hydronium concentration [H3O+] = 3.38 x 10-13
3) What is the hydronium concentration [H3O+] of a solution with a pH of 7.84?
What is the hydroxyl concentration [OH-] of this solution?
hydronium concentration [H3O+] = 1.45 x 10-8
hydroxyl concentration [OH-] = 6.90 x 10- OR pOH = 6.16 hydroxyl concentration [OH-] = 6.9 x 10-7
4) What is the hydroxyl concentration [OH-] of a solution with a pH of 3.76?
hydroxyl concentration [OH-] = 5.75 x 10-11
5) What is the hydronium concentration [H3O+] of a solution with a pOH of 2.47?
hydronium concentration [H3O+] = 2.95 x 10-12
Identify the following compounds as a strong acid, weak acid, strong base, or a weak base.
NaOH
strong acid
HCl
strong acid
H2CO3
weak acid
KOH
strong base
H2SO4
strong acid
LiOH
strong base
NH3
weak base
Solute, Solvent/ Dissolving Process Practice Problem
Answers
Distinguish in the following examples between the solute and the solvent.
nail polish (solute) acetone (solvent)
glue (solute) acetone (solvent)
eggshells (solute) vinegar (solvent)
iodine (solute) hexane (solvent)
chromium (solute) hydrochloric acid (solvent)
Kool Aid (solute) water (solvent)
Describe the difference between unsaturated, saturated, and supersaturated solutions.
At the maximum solubility the solution is saturated and in dynamic equilibrium with the insoluble part of
solute. Such a solution is called saturated. Solution with less concentration is call unsaturated.
NaCl(s) <==> Na+(aq) + Cl (aq)
If there is more solute dissolved than saturation allows, the solution is said to be supersaturated.
Draw a picture showing a salt crystal.
Now draw how that salt crustal dissolves in a water solution.
Dissolving and Solution Concentration Answers
1)
How many grams of beryllium chloride are needed to make 125 mL of a 0.050 M solution?
0.50 grams
2)
How many grams of beryllium chloride would you need to add to 125 mL of water to make a
0.050 molal solution?
0.50 grams
3)
The density of ethanol is 0.789 g/mL. How many grams of ethanol should be mixed with
225 mL of water to make a 4.5% (v/v) mixture?
7.99 grams
4)
Explain how to make at least one liter of a 1.25 molal ammonium hydroxide solution.
Dissolve 43.8 grams (1.25 moles) of ammonium hydroxide in 1 L H2O.
5)
What is the molarity of a solution in which 0.45 grams of sodium nitrate are dissolved in 265
mL of solution.
0.020 M
6)
What is the mole fraction of sulfuric acid in a solution made by adding 3.4 grams of sulfuric
acid to 3,500 mL of water?
1.8 x 10-4
7)
What will the volume of a 0.50 M solution be if it contains 25 grams of calcium hydroxide?
680 mL
8)
How many grams of ammonia are present in 5.0 L of a 0.050 M solution?
4.3 grams
Endothermic and Exothermic Practice Problem Answers
1) What type of chemical reaction absorbs energy and requires energy for the reaction to occur?
a. endothermic
b. exothermic
c. synthesis
d. both A and B
2) What type of reaction releases energy and does not require initial energy to occur?
a. endothermic
b. exothermic
c. decomposition
d. both A and B
3) Which of the following are examples of an exothermic chemical reaction? Check all that apply.
a. photosynthesis
b. burning a piece of wood
c. freezing ice into water
d. none of the above
4) Which type of reactions cannot occur spontaneously?
a. endothermic
b. exothermic
c. neither
5) Any type of reaction that involves burning (combustion) can be classified as which of the following types
of reactions?
a. synthesis
b. endothermic
c. exothermic
d. all of the above
6) Burning sugar is an exothermic process.
a. True
b. False
7) What is enthalpy?
a. heat content
b. absolute amount of energy is a chemical system
c. the reactants of the chemical reaction
d. none of the above
8) A 100 g sample of water at 25 oC and 1 atm. of pressure _?_ than 100 g of water that has recently been
heated to 100 oC from 0 oC and then cooled to 25 oC at 1 atm. of pressure.
a. has more internal energy than
b. has less internal energy than
c. has the same internal energy as
9) Water has a specific heat of 4.184 J/g deg while glass (Pyrex) has a specific heat of 0.780 J/g deg. If 10.0 J
of heat is added to 1.00 g of each of these, which will experience the larger increase of temperature?
a. glass
b. water
c. They both will experience the same change in temperature since only the amount of a substance
relates to the increase in temperature.
Specific Heat Practice Problem Answers
1) cp = 0.0042 J/gC
2) final temperature of metal and water is 76C
3) final temperature of metal and water is 97C
4) 8.8 x 105 J
5) 3.2 x 104 J
6 a) 5.52 ºC
b) 10.1 ºC
Changing the Rate of a Reaction Problem Answers
1.
A
The activation energy is the energy that must be overcome for the reaction to proceed. Also remember that
for a reaction to occur, the collisions between molecules must be sufficiently energetic and of the proper
geometric orientation.
2.
C
The energy change for the overall reaction is simply the difference between the energies of the products and
reactants, and this is indicated by the letter C on the diagram.
3.
T, T
(Fill in CE.) The energy change indicated by A on the diagram represents the activation energy of the
reaction—the energy investment required to form the activated complex Y, also known as the energy that
must be put into the system to make the reaction go. B on the diagram represents the energy released when
the unstable transition state molecule Y goes to a lower energy state as the products Z. The reaction is
exothermic when the energy payoff exceeds the energy investment, and since the second statement is the
reason for the first statement, you would fill in the CE oval.
4.
T, F
(Do not fill in CE oval.) The first statement is true—when a chemical reaction is at equilibrium, the rate of
the forward reaction is equal to the rate of the reverse reaction, in which the reactants are formed. However,
statement II is incorrect and is a common misconception. The amount of reactant and product remain
constant at equilibrium but usually do not equal each other. Since the second statement is false, you would
not fill in the CE oval.
5.
E
Two conditions must be met in order for a chemical reaction to occur. First of all, the molecules must collide
with sufficient energy, and second, the molecules must collide with such an orientation that the product
bonds can be formed.
6.
C
The addition of a catalyst lowers the activation energy, thus speeding up a chemical reaction.
Le Chatelier’s Practice Problem Answers
1.
E
This question combines two concepts. The reaction is exothermic, so think of heat as a product. Increasing
the temperature has the same effect as increasing a product’s concentration, so it causes a shift to the left,
meaning statement I cannot be in the answer choice. Decreasing the temperature (removing heat) would have
the same effect as removing a product (since the reaction is exothermic), so this would cause a shift to the
right, and II must be in the correct answer choice. Finally, since all reactants and products are in the gas
phase, and there is a total of four moles of gas on the left and a total of two moles of gas on the right,
increasing the pressure will push the reaction toward the side with the fewest moles of gas. In this case, the
side with the fewer moles of gas is the products side, so this also causes a shift to the right. III is also correct,
and answer choice E is correct.
2) Suggest four ways to increase the concentration of SO3 in the following equilibrium reaction.
Be specific.
2 SO2(g) + O2(g)
2 SO3(g) + 192.3 kJ
1. _______increase the pressure_____________________________________________________________
2. _______add more SO2__________________________________________________________________
3. _______add more O2 gas _______________________________________________________________
4. ______decrease the temperature __________________________________________________________
3) Use Le Chatelier's Principle to predict how the changes listed will affect the following equilibrium
reaction:
2 HI (g) + 9.4 kJ
H2 (g) + I2 (g)
a) Will the concentration of HI increase, decrease, or remain the same if more H2 is added?
increase
b) What is the effect on the concentration of HI if the pressure of the system is increased?
no effect
c) What is the effect on the concentration of HI if the temperature of the system is increased?
decrease
d) What is the effect on the concentration of HI if a catalyst is added to the system?
decrease
e) Write the equilibrium constant expression for this reaction.
Keq = [H2] [I2]
[HI]2
Organic Chemistry Practice Problem Answers
1. Functional Groups: Methyl & ester. Properties: Neutral, water soluble.
2. Functional Groups: Methyl, amino, hydroxyl & carboxyl. Properties: Has both acidic and
basic functional groups, carboxyl is a stronger acid than amino is a base so overall this
molecule should act as an acid, water soluble.
3. Functional Groups: Methyl, hydroxyl & ketone. Properties: Neutral, hydrophobic &
therefore insoluble.
4. Functional Groups: Hydroxyl & aldehyde. Properties: Neutral, very water soluble.
5. Functional Groups: Amino, hydroxyl & phosphate. Properties: Acidic (phosphates are the
most acidic functional group), water soluble.
6. Functional Groups: Hydroxyl & ketone. Properties: Neutral, very water soluble.
Nuclear Practice Problem Answers
1) Select the correct equation when: Americium-244, Am, undergoes decay to form Curium-244, Cm..
.
.
.
YES
2) Which of the following statements is true?
a. The atomic number is always greater than the atomic mass.
b. The mass number is the same for all atoms of the same element.
c. The atomic number is the sum of the number of particles in the nucleus
d. The difference between the atomic mass and the atomic number is the number of neutrons
e. All the above are correct
3) Neon-21 undergoes beta particle emission. Indicate the correct equation
YES
Directions: Identify the following as alpha, beta, gamma, or neutron.
1.
1
0
n neutron
2.
0
1
e
beta
3.
4
2
He
beta
4.
239
Pu
94
→
7. Least penetrating nuclear decay
alpha
8. Most damaging nuclear decay to the human body
gamma
9. Nuclear decay that can be stopped by skin or paper.
alpha
10. Nuclear decay that can be stopped by aluminum.
beta
Complete the following nuclear equations.
11. 42
K → -10 e +
__________
19
12.
13.
14.
9
4
Be
→
9
4
Be + __________
235
92 U
0
0
γ gamma___
4
He
2
+ __________
→ _________ +
231
90
Th