Download Atomic Theory (2

Chemistry Review Packet
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Atomic Theory (2.01, 4.01)
1.) Who were the early contributors to the atomic theory, and what was their view of the atom?
2.) Who discovered the electron?
3.) Who discovered the charge of an electron?
4.) Who discovered the nucleus?
5.) Who discovered the proton?
6.) Describe the composition of the atom and the experiments that led to that knowledge.
7.) Describe the Bohr Model- be specific.
8.) Describe the electron cloud model of the atom in terms of electron probability.
9.) What are quanta (or photons?)
10.) Describe the process by which an atom can be “excited”.
11.) What happens to the electron(s) when an atom loses energy?
12.) What is the difference between an excited state and ground state?
13.) Describe the limitations of the Bohr’s model of the atom and how quantum model provides a more
accurate picture of electron arrangement. .
Atomic Structure (2.02)
1.) What is the relative mass, relative charge, and location of each of the subatomic particles?
2.) What do the symbols A and Z mean in relation to atoms?
3.) Write the isotope notation for an isotope of Carbon in the 2 accepted isotopic notations.
4.) What element is A= 34 , Z= 16 ? How many protons does it have? Neutrons? Electrons?
5.) What mass number do most atoms of Phosphorus have?
Nomenclature (2.03)
1.) In sentence form, describe how covalent compounds are named.
2.) In sentence form, describe how binary ionic compounds are named.
3.) In sentence form, describe how polyatomic compounds are named.
4.) In sentence form, describe how binary OR polyatomic compounds are named when bonded with a
transitional metal.
5.) In sentence form, describe how binary acids are named.
6.) In sentence form, describe how polyatomic acids are named.
7.) Use the following polyatomic ions for the following questions: nitrate, sulfate, carbonate, acetate. Write
the formula and provide the name when the following metals combine with the above mentioned polyatomic
ions.
nitrate, sulfate, carbonate, acetate
a. Aluminum
b. Copper (II)
c. Silver
d. Lithium
e. Calcium
f. Ammonium
8.) Write the name for the following acids:
a. HCl
b. HNO3
c. H2SO4
d. HC2H3O2
9.) Write the formula or name for the following:
a. B2H6
b. Vanadium(III) acetate
c. Iodine trichloride
d. Zn3(PO4)2
e. Mn(ClO4)2
f. Sulfurous acid
g. Antimony pentafluoride
h. Disilicon hexachloride
i. Pb(C2H3O2)2.2H2O
j. CCl4
Physical and Chemical Properties (2.04)
1.) Identify 7 physical properties that are used to characterize a sample.
2.) Calculate mass or volume when given: (You will need to use density tables in your textbook or other
sources)
a. 123 mL of Oxygen
b. 789 g of Sulfur Dioxide
c. 467 mL of Hexane
d. 85.3 cm3 of aluminum
e. 95.1 g of Mercury
3.) A substance has a mass 55.4 grams and a volume of 77.37 mL. What is the substance?
4.) Calculate the density of Carbon monoxide at STP. Is it more or less dense than Oxygen?
Gases (2.05)
1.) What are the 5 characteristics of ideal gases?
2.) What is the volume of one mole of any gas at STP?
3.) At what temperature would 2.10 moles of N2 gas have a pressure of 1.25 atm and in a 25.0 L tank?
4.) What volume is occupied by 5.03 g of O2 at 28°C and a pressure of 0.998atm?
5.) What is the volume of 1.00 mole of a gas at standard temperature and pressure?
6.) A gas takes up a volume of 17 liters, has a pressure of 2.3 atm, and a temperature of 299 K. If I raise the
temperature to 350 K and lower the pressure to 1.5 atm, what is the new volume of the gas?
7.) Atmospheric pressure on the peak of Mt. Everest can be as low as 150 mm Hg, which is why climbers need
to bring oxygen tanks for the last part of the climb. If the climbers carry 10.0 liter tanks with an internal gas
pressure of 3.04 x 104 mm Hg, what will be the volume of the gas when it is released from the tanks?
8.) A man heats a balloon in the oven. If the balloon initially has a volume of 0.4 liters and a temperature of 20
0
C, what will the volume of the balloon be after he heats it to a temperature of 250 0C?
9.) Blast furnaces give off many unpleasant and unhealthy gases. If the total air pressure is 0.99 atm, the partial
pressure of carbon dioxide is 0.05 atm, and the partial pressure of hydrogen sulfide is 0.02 atm, what is the
partial pressure of the remaining air?
10.) 10.0 L of a gas is found to exert 97.0 kPa at 25.0°C. What would be the required temperature (in Celsius)
to change the pressure to standard pressure?
11.) A gas that has a volume of 28 liters, a temperature of 45 0C, and an unknown pressure has its volume
increased to 34 liters at STP. What is the original pressure of the gas?
Bonding (2.06, 2.07)
1.) What is an ion?
2.) How are ions formed?
3.) Why do ions form readily? (Be specific- in terms of stability)
4.) What is an anion?
5.) What is a cation?
6.) What type of elements form anions?
7.) What type of elements form cations?
8.) What is the relationship between valence electrons and ionic charges?
9.) Describe the force of attraction found in ionic bonds.
10.) Describe how to identify the type of bond by the elements present.
11.) Describe how to identify the type of bond based on electronegativity.
12.) Describe at least 4 physical properties of ionic bonds. Explain how the bond type is responsible for these
properties.
13.) Describe how a covalent bond is formed.
14.) Describe at least 4 physical properties of covalent bonds. Explain how the bond type is responsible for
these properties.
15.) Define a metallic bond.
16.) Describe at least 6 physical properties of metallic bonds. Explain how the bond type is responsible for
these properties.
17.) What are the 7 diatomic molecules? What is the saying that helps students to identify the molecules?
How are they named?
Molecular Structure Goal 2.07
1.) Of the diatomic substances, which has a double bond?
2.) Of the diatomic substances, which has a triple bond?
3.) Describe the relationship between the number of bonds (single, double, or triple) and the strength of the
bond. (In the first 2 questions, which has a stronger bond?)
4.) Draw the Lewis dot diagram for CH4.
5.) Draw the Lewis dot diagram for NH3.
6.) Draw the Lewis dot diagram for H2O. What is the shape of this molecule? Is it polar or non-polar? How
can you tell?
7.) Draw the Lewis dot diagram for CO2.
Periodic Trends (3.01)
1.) Define Group.
2.) What are the main groups on the periodic table?
3.) List 3 things that elements in the same group have in common.
4.) Describe how reactivity changes (how easily an element changes in a chemical reaction) as you go
up/down/across the periodic table for both metals and non-metals.
5.) Define Period.
6.) How many periods are found on the periodic table?
7.) What period does the D-block start in?
8.) What group does the following belong in:
a. Chlorine: ________________________
b. Neon: ________________________
c. Calcium: ________________________
d. Cesium: ________________________
e. Iron: ________________________
9.) Write the electron configurations, noble-gas configurations, AND orbital notation diagrams for the
following elements:
a. Magnesium
b. Bromine
c. Manganese
d. Phosphorus
10.) How many valence electrons do the following elements have?
a. 1s22s22p5
b. 1s22s22p63s23p6
c. 1s22s22p63s23p64s23d104p3
11.) How many electrons will be lost/gained AND what will the charge be for the following elements?
a. 1s22s22p3
b. 1s22s22p63s23p4
c. 1s22s22p63s23p64s23d104p5
d. 1s22s22p63s1
12.)
Define Ionization Energy
13.)
Explain the Period and Group Trends for Ionization.
14.)
Define atomic radius.
15.)
Explain the Period and Group Trends for Atomic Radius.
16.)
Define electronegativity
17.)
Explain the Period and Group Trends for Electronegativity.
18.)
How can be predict bond type based on electronegativity?
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Moles! (3.02)
1.) Convert the following to moles:
a. 12.04 x 1023 of C8H18
b. 6.07 x 1023 of Na2CO3
c. 1.5 x 1023 of sodium phosphate
d. 48 g of C
e. 13 g of V
f. 24.6 L of CH4
g. 75.3 L of NO2
2.) Convert the following to particles (atoms, compounds, molecules, etc):
a. 3.5 moles of Aluminum Phosphide
b. 7.9 moles of diphosphorus pentoxide
c. 1.5 moles of Carbon tetrachloride
d. 5 L of Chlorine
e. 15 L of Hydrogen gas
f. 32 g of O2
g. 62 g of P
h. 132 g of CO2
3.) Convert the following to mass
a. 32 L of N2
b. 64L of N2O5
c. 4.5 moles of CO2
d. 12.07 x 1023 atoms of BaSO4
e. 1.25 x 1023 atoms of CaCO3
4.) A compound is 43.8% N, 6.2% H, and 50.0% O. What is its molecular formula?
5.) List the steps for finding the empirical formula.
6.) A compound has the empirical formula of CH2O and a molecular mass of 180 g. What is the
molecular formula?
7.) What is the percent composition by mass of each of the atoms in methane?
8.) What is the percent composition by mass of oxygen in the formula Fe2O3?
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Stoichiometry (3.03)
1.) How many moles of chlorine gas (Cl2) would react with 5 moles of sodium (Na) according
to the following unbalanced chemical equation? Na + Cl2 --> NaCl
2) Using the equation above, determine the amount of product that can be produced from 24.7 g Na.
3) How many molecules of product would be produced from 24.7g Na?
4) In the reaction 2C8H18 + 25O2 --> 16CO2 + 18 H2O, the ratio of volumes of O2 to CO2
is _________________.
5) If 27.3g of C8H18 are burned, what mass of water will be produced?
6) If 48.3g of C8H18 are burned, how many molecules of CO2 will be produced?
7) If 87.3g of C8H18 are burned, how many Liters of CO2 will be produced?
8.) How many Liters of oxygen are produced by the burning of 45 g of C8H18?
Energy Changes Goal 4.02, Goal 2.08
1. Define endothermic
2. Define exothermic
3. For chemical reactions, complete the following table
Endothermic Reaction
Exothemic Reaction
∆ H (+ or –
joules/mole)
Product or reactant
Energy diagrams
Diagram the
energy of the
reactants and
products & the ∆H
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Heating and Cooling Curve
H
F
A
200º C
60º C
I
K
G
B
J
D
M
E
L
C
N
Energy
O
P
1.) D
e
s
c
r
i
b
e
what is happening in the following using the graph above.
a. Point A: ________________________________________________________
b. Point B: ________________________________________________________
c. Points D to E: ________________________________________________________
d. Points L to M: ________________________________________________________
e. Points F to G: ________________________________________________________
f. Points J to K: ________________________________________________________
g. Points D to E: ________________________________________________________
h. Points I to J: ________________________________________________________
i. Points C to D: ________________________________________________________
j. Points G to H: ________________________________________________________
k. Points E to F: ________________________________________________________
l. Points K to L: ________________________________________________________
m. Points M to N: ________________________________________________________
n. Area O: ________________________________________________________
o. Area P: ________________________________________________________
p. Where does melting occur? ____________________________________________________
q. Where does condensation occur? ________________________________________________
2.) Describe the above graph in terms of exothermic and endothermic.
3.) Describe the heating and cooling curve in terms of intermolecular forces and speed of molecules.
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Chemical Reactions 5.01, 5.02, 5.03
Chemical Reactions (5.01, 5.02)
1.) Identify the type of reactions:
a. ___________________ 4KMnO4  2K2O + 4MnO + 5O2
b.________________________ 2 KI + Br2  2KBr
+ I2
c. ________________________ Na2O + CO2  Na2CO3
d._________________ C12H22O11 + 12O2  12CO2 +11H2O
e. _________________ CuBr2 + ZnSO4  CuSO4 + ZnBr2
f. ___________________ What type of reaction takes place when fluorine reacts with sodium bromide?
g.________________________ What type of reaction takes place when an aqueous solution of magnesium
hydroxide is mixed with aqueous copper chloride.
h.________________________ 2 KClO3  2KCl + 3 O2
i. _______________________________ H2SO4 + NaOH  Na2SO4 + H2O
2.) Predict AND BALANCE the products using your reference table.
a.
Ba(OH)2  ________________ + ___________________
b.
Ca +
c.
N2 + 3H2  ________________________________
d.
Na3PO4 + CaCl2  ____________ + ________________
e.
C2H5OH + O2  ____________ + _______________
f.
CaO + H2O  ____________________________
g.
HBrO3  ______________ + ________________
h.
KClO3  ______________ + ____________________
HCl  ___________+ _______________
i. NaCl + Pb(NO3)2  ____________ + ______________
j. K + Cl2  ___________________________________
k. Cl2 + NaBr  _____________ + __________________
l. Cu + AgNO3  ______________ + ________________
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m. C4H10 + O2  ______________ + _________________
n. P2O5 + H2O  __________________________
o. H2O  ______________ + ________________
p. NaHCO3  ____________ + ____________ + _____________
q. LiClO3  ______________ + ________________
r. KI + Br2  ______________ + ________________
s. Li2CO3  ______________ + ________________
t.
Acid + Base  ______________ + ________________
3.) Write a complete BALANCED REACTION AND the NET IONIC REACTION for the following (don’t
forget states of matter):
a. potassium iodide added to lead(II) nitrate
b. sodium sulfite combined with acetic acid
c. a solution of sodium chloride added to dilute sulfuric acid
d. sodium hydrogen carbonate added to hydrochloric acid
4.) Using complete sentences, explain how you can use the reference table to predict the products for all types
of reactions.
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