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Summer Work
Grade 7
American Program
Table of Contents
Part I – Structure of the Atom ---------------------------------------------------------------- Pages 1–3
Part II –Chemical Reactions ------------------------------------------------------------------- Pages 4–7
Part III –Chemical Bonding ------------------------------------------------------------------ Pages 8–10
Part IV – Organic Chemistry --------------------------------------------------------------- Pages 11–13
Part V – Electrochemistry ------------------------------------------------------------------ Pages 14–16
Part I – Structure of the Atom
First Exercise: Read the following information on structure of the atom. Fill in the blanks where
necessary.
1. The 3 particles ofthe atom are:
a. ______________________
b. ______________________
c. ______________________
2. Their respective charges are:
a. ______________________
b. ______________________
c. _____________________
3. The number of protons in one atom of an element determines the atom’s
__________________ , and the number of electrons determines ___________________
of an element.
4. The atomic number tells you the number of ______________________ in one atom of an
element. It also tells you the number of ______________________ in a neutral atom of
that element. The atomic number gives the “identity “ of an element as well as its
location on the Periodic Table.
5. No two different elements will have the ______________________ atomic number.
6. The ______________________ of an element is the average mass of an element’s
naturally occurring atom, or isotopes, taking into account the ______________________
of each isotope.
7. The ______________________ of an element is the total number of protons and neutrons
in the______________________ of the atom.
8. The mass number is used to calculate the number of ______________________ in one
atom of an element. In order to calculate the number of neutrons you must subtract the
______________________ from the ______________________.
Second Exercise: Give the symbol and number of protons in one atom of:
Lithium __________________
Bromine __________________
Iron __________________
Copper __________________
Oxygen __________________
Mercury __________________
Krypton __________________
Helium __________________
Third Exercise: Give the symbol and number of electrons in a neutral atom of:
Uranium __________________
Chlorine __________________
Boron __________________
Iodine __________________
Antimony __________________
Xenon __________________
Fourth Exercise: Give the symbol and number of neutrons in one atom of: Show your
calculations.
Barium __________________
Bismuth __________________
Carbon __________________
Hydrogen __________________
Fluorine __________________
Magnesium __________________
Europium __________________
Mercury __________________
Fifth Exercise: Name the element which has the following numbers of particles:
1.
2.
3.
4.
5.
6.
26 electrons, 29 neutrons, 26 protons _____________________
53 protons, 74 neutrons _____________________
2 electrons (neutral atoms) _____________________
20 protons _____________________
86 electrons, 125 neutrons, 82 protons (charged atom) _____________________
0 neutrons _____________________
Sixth Exercise: Fill in the following table:
Name
Symbol
Z
A
#p
#e
#e
Atomic
representation
Na
17
Potassium
P
Iron
53
Silver
36
Seventh Exercise: Approximately 75% of the chlorine atoms found in nature have a mass of 35.
The other 25% have a mass of 37.Whatshould we report as the average atomic weight for
chlorine?
Eighth Exercise: Suppose that a new element (E) were discovered that existed as three natural
isotopes. 25% ofthe atoms had a mass of 278, 38% had a mass of 281, and the remainder had a
mass of 285. What would be listed as the atomic weight of this element?
Ninth Exercise: Consider 200 grams of H2SO4. Calculate the number of moles present in this
mass, as well as the number of molecules. M(H)=1g/mol M(O)=16g/mol M(S)=32g/mol
Part II –Chemical Reactions:
First Exercise: Balance the following reactions:
1. ____ NaBr + ____ Ca(OH)2  ___ CaBr2 + ____ NaOH
2. ____ NH3+ ____ H2SO4  ____ (NH4)2SO4
3. ____ C5H9O + ____ O2  ____ CO2 + ____ H2O
4. ____ Pb + ____ H3PO4  ____ H2 + ____ Pb3(PO4)2
5. ____ Li3N + ____ NH4NO3  ___ LiNO3 + ___ (NH4)3N
6. ____ HBr + ___ Al(OH)3  ___ H2O + ___ AlBr3
Second Exercise: Indicate which type of chemical reaction (synthesis, decomposition, singledisplacement, double-displacement or combustion) is being represented in the reactions below:
1. Na3PO4 + 3 KOH  3 NaOH + K3PO4
Reaction Type _______________________
2. MgCl2 + Li2CO3  MgCO3 + 2 LiCl
Reaction Type _______________________
3. C6H12 + 9 O2  6 CO2 + 6 H2O
Reaction Type _______________________
4. Pb + FeSO4  PbSO4 + Fe
Reaction Type _______________________
5. CaCO3  CaO + CO2
Reaction Type _______________________
6. P4 + 3 O2  2 P2O3
Reaction Type _______________________
7. 2 RbNO3 + BeF2  Be(NO3)2 + 2 RbF
Reaction Type _______________________
8. 2 AgNO3 + Cu  Cu(NO3)2 + 2 Ag
Reaction Type _______________________
9. C3H6O + 4 O2  3 CO2 + 3 H2O
Reaction Type _______________________
10. 2 C5H5 + Fe  Fe(C5H5)2
Reaction Type _______________________
11. SeCl6 + O2  SeO2 + 3Cl2
Reaction Type _______________________
12. 2 MgI2 + Mn(SO3)2  2 MgSO3 + MnI4
Reaction Type _______________________
13. O3  O. + O2
Reaction Type _______________________
14. 2 NO2  2 O2 + N2
Reaction Type _______________________
Third Exercise: For the reaction below, which change would cause the equilibrium to shift to
the right?
CH4(g) + 2H2S(g) ↔CS2(g) + 4H2(g)
(a) Decrease the concentration of dihydrogen sulfide.
(b) Increase the pressure on the system.
(c) Increase the temperature of the system.
(d) Increase the concentration of carbon disulfide.
(e) Decrease the concentration of methane.
Fourth Exercise: Predict the effect of decreasing the volume of the container for each
equilibrium.
(a) 2H2O(g) + N2(g) ↔ 2H2(g) + 2NO(g)
(b) SiO2(s) + 4HF(g) ↔ SiF 4(g) + 2H2O(g)
(c) CO(g) + H2(g) ↔C(s) + H2O(g)
Fifth Exercise: Complete the following chart on page 7 by writing left, right or none for
equilibrium shift, and decreases, increases or remains the same for the concentrations of
reactants and products, and for the value of K.
N2(g) + 3H2(g) ↔ 2NH3(g) + 22.0 kcal
Stress
Add N2
Add H2
Add NH3
Remove N2
Remove H2
Remove NH3
Increase
Temperature
Decrease
Temperature
Increase Pressure
Decrease Pressure
Equilibrium
Shift
Concentration
of N2
Concentration
of H2
Concentration
of NH3
Part III – Chemical Bonding:
First Exercise: Bonding of NaCl
1. Fill in the Lewis dot symbols for Na (11) and Cl (17), below, and complete the shorthand
electron configuration for each:
Na
__s
Cl
__s __ p
2. Now allow each atom to complete its octet, by losing or gaining electrons. Then write the
shorthand electron configuration for each ion and its charge.
Na
Cl
3. Show the formation of the bond. Indicate its type.
Second Exercise: Show the transfer of electrons between the following atoms to form cations
and anions, and show the bond, in addition to indicating its type.
a) K (19) and S (16)
b) O (8) and Ba (56)
Third Exercise: Label "i" (ionic) or "c" (covalent) the bonds in the following compounds.
1. CO2
_____
2. Na2O
_____
3. SF2
_____
4. N2O
_____
5. CaO
_____
6. Na2CO3 _____
Fourth Exercise: Answer the following multiple choice questions by choosing the correct letter.
1) How many electrons are shared in a single covalent bond?
A) 2
B) 3
C) 8
D) 1
1. ________
E) 4
2) How many electrons are shared in a double covalent bond?
A) 3
B) 6
C) 8
D) 2
E) 4
3) How many unshared pairs of electrons does the nitrogen atom in
ammonia possess?
A) 1
B) 2
C) 4
D) 5
2. ________
3. ________
E) 3
4) How do atoms achieve noble-gas electron configurations in single
covalent bonds?
4. ________
A) Two atoms share two electrons.
B) Two atoms share two pairs of electrons.
C) Two atoms share one electron.
D) One atom completely loses two electrons to the other atom in the bond.
5) Why do atoms share electrons in covalent bonds?
A) to attain a noble-gas electron configuration
B) to become ions and attract each other
C) to become more polar
D) to increase their atomic numbers
5. ________
6) Which of the following is the name given to the pairs of valence
electrons that do not participate in bonding in diatomic oxygen molecules?
6. ________
A) outer pair
B) unvalenced pair
C) bound pair
D) unshared pair
E) inner pair
7) Which elements can form diatomic molecules joined by a single covalent bond? 7.________
A) hydrogen only
B) halogens and members of the oxygen group only
C) halogens only
D) hydrogen, halogens, and members of the oxygen group
E) hydrogen and the halogens only
8) A molecule with a single covalent bond is _____.
B)
A) CO2
CO
C) Cl2
8. ________
D) N2
9) A diatomic molecule with a triple covalent bond is _____.
A) O2
B) F2
C) N2
9. ________
D) H2
10) Which of the following diatomic molecules is joined by a double
covalent bond?
A) Cl2
B) O2
C) N2
10. _______
D) H2
Fifth Exercise: Draw Lewis dot structures for the following compounds:
CO2
C2H2
HCN
Part IV – Organic Chemistry:
First Exercise: Draw the complete structural formula and condensed molecular formula for each
compound.
IUPAC Name
butane
2-methylhexane
3-ethyl-2-methylnonane
propene
ethyne
2-octyne
cyclopropane
3-pentene
Complete Structural Formula
Condensed Formula
Second Exercise: Use IUPAC naming rules to name the following hydrocarbon compounds:
CH2 ─ CH3
|
a) CH3─ CH ─CH2─CH ─ CH ─ CH3
|
|
CH3
CH2
|
CH3
b)
H
CH3
\
/
C=C
/
\
H
CH2 ─CH3
c)
CH ≡ C─ CH3
d) CH3 -CH2-CH2-CH=CH-CH3
e)
f)
g)
h)
Third Exercise: Consider the following table of alkanes:
Alkane
Chemical
Formula
Melting
Point (oC)
Boiling Point
(oC)
Methane
Ethane
Propane
Butane
Pentane
Hexane
Heptane
CH4
C2H6
C3H8
C4H10
C5H12
C6H14
C7H16
-182
-183
-190
-138
-130
-95
-91
-161
-89
-42
-1
36
69
98
1. Draw the condensed formulas of all the above alkanes.
2. What is the general formula of alkanes?
______________________________________________________________________________
3. How does the boiling point change with the increase in number of carbons?
______________________________________________________________________________
4. How does the melting point change with the increase in number of carbons?
______________________________________________________________________________
Part V – Electrochemistry:
First Exercise: Assign oxidation numbers to each of the atoms in the following compounds:
CO2:
O=
C=
CH4:
H=
C=
HClO4
O=
H=
MnO2
O=
Mn =
SO32-
O=
S=
Cl =
Second Exercise: Consider the following compounds:
1. Assign an oxidation state to each of the following nitrogen containing compounds:
Formula
Oxidation Number of N
NH3
NO2
NO3NH2OH
N2H4
2. Indicate the type of the following reaction according to N. Indicate the role of Nitrogen.
N2(g) + 3 H2(g) → 2 NH3(g)
Third Exercise: In each of the following reactions, assign oxidation numbers to all of the
elements, identify the type of the reaction done by each element, and identify the oxidizing and
reducing agents.
a) 4 Fe + 3 O2 → 2 Fe2O3
b) Cr2O72-
+
2OH- →
c) NH4NO2 → N2 +
d) P4
+
2 CrO42- +
H2O
2 H2O
10 Cl2 → 4 PCl5
e) 2 Cr3+ +
H2O
+
6 ClO3- →
Cr2O72- +
6ClO2
+
2 H+
Fourth Exercise: Consider the following galvanic cell. Answer the question on page 16.
C.
D.
Anode
Cathode
B.
F.
1. Complete the labeling the above diagram:
A. _________________
B. _________________
C. _________________
D. _________________
E. _________________
F. __________________
G. _________________
2. Write the half-reaction occurring at each electrode and indicate its type.
Anode: ___________________________________
type: ________________
Cathode: ___________________________________ type: ________________
3. What is the overall equation?
______________________________________________
4. What is the charge of the anode? ______________
5. What is the charge of the cathode? ______________
6. The electrons flow from:
a) Zn to Cu
b) Cu to Zn
7. Give an example of a salt bridge, and indicate to which direction its ions will go?
______________________________________________________________________________