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Chemisty 11 Exam Review
1. Classify each of the following as homogeneous or heterogeneous mixtures:
a) blood b) chocolate chip cookies c) dissolved “Kool-Aid” d) homogenized milk e) gold jewelry.
2. Classify the following changes as physical or chemical:
a) food spoils b) an icicle melts c) a nail rusts d) oil is pumped out of a well
f) salt dissolves in water g) a window is broken
e) an egg is fried
3. Complete the following table:
Element
Symbol
Atomic #
Mass #
# of Neutrons
nitrogen-15
8
22
10Ne
beryllium-9
4
4. Give the complete symbol (include atomic number, mass number and charge) for:
(a) an atom with 17 electrons, 18 neutrons.
(b) an ion with 19 protons, 18 electrons, 20 neutrons.
(c) an ion with a -1 charge, 10 electrons, 10 neutrons.
(d) an atom with 31 protons and 39 neutrons.
5. Name the element in:
(a) group 6A, period 4
(b) the alkaline earth metals group, period 5
(c) the halogen group, period 5
(d) the oxygen group, period 3
(e) the metalloid in the halogen group
(f) the alkali metal group, period 3
(g) the noble gas group, period 6
6. State if the element is a representative element, a transition metal, an inner-transition metal, or a
metalloid:
(a) N
(b) Na
(c) Ni
(d) Nd
(g) Ga
(h) Os
(i) At
(j) Al
(e) Zr
7. Predict the monatomic ions formed by the following elements:
a) hydrogen
b) magnesium
c) sulfur
d) iodine
(f) Ge
e) aluminum
8. Define:
(a) valence electrons (b) anion
(c) cation
(d) ionization energy
(e) electronegativity
(f) octet rule
(g) ionic bond
(h) covalent bond (i) polar bond
(k) nonpolar bond
9. Give complete electron configurations for:
(a) K
(b) K+
(c) S
(d) S2-
(e) Ca
(f) Ca2+
(g) Ar
10. Describe:
(a) Dalton’s Atomic Theory (b) Thomson’s model of the atom (c) Rutherford’s model of the atom
(d) Bohr’s model of the atom (e) The Quantum Mechanical model of the atom
11. Complete the following table:
# of Valence
Dot Structure
Element
Electrons
Ion formed
Cation or
Anion
Noble gas it is
isoelectronic with
O
Na
P
Ca
12. Describe the difference between ionic and molecular compounds.
13. Classify each of the following as ionic or molecular, and give the correct chemical name for
each:
a)
b)
c)
d)
e)
f)
g)
NaI
NO
GaBr3
S4N2
ZnCl2
SnF4
P2O5
14. Give the correct chemical formula for each of the following:
a) hydrochloric acid
b) strontium nitrate
c) calcium chloride
d) ammonia
e) sulphur trioxide
f)
lithium sulfate
g) calcium hydroxide
h) ammonium sulfate
i)
lead (II) sulfite
j)
copper (I) sulfide
k) aluminum oxide
l)
magnesium bromide
m) hydrogen gas
n) nitrogen dioxide
15. Circle the element with the larger atomic size:
(a) N or O
(b) Ag or Au
(c) K or Ca
(d) Ge or Ga
16. Circle the element with the larger ionization energy.
(a) N or O
(b) Ag or Au
(c) K or Ca
(d) Ge or Ga
17. Circle the element with the larger electronegativity value.
(a) N or O
(b) Ag or Au
(c) K or Ca
(d) Ge or Ga
18. Arrange the following in order from least electronegative to most electronegative.
Ge,
Mg, Fr,
F,
C,
Te
19. Circle the bond that is the most polar:
(a) C-O,
C-N,
C-F
(c) I-Cl,
I-Br,
I-F
(b) P-N,
(d) Si-P,
As-Se,
C-N,
Br-Cl
Ge-As
20. What is the most electronegative element? What is the least electronegative element?
21. The following compounds have polar molecules, please draw a structural formula and show
which atoms are slightly positive and which are slightly negative.
a. HCl
b. H2O
22. (I) Determine if the following are ionic or covalent.
(II) If it is ionic, list the ions involved in the bonding
(III) If it is covalent, draw the Lewis structure, and state if it has polar or nonpolar bonds.
(a) MgCl2
(g) Na3N
b) H2O
(h) CaO
(c) SeBr2
(i) CO2
(d) NBr3
(j) AsCl3
(e) Al2S3
(f) CH4
23. Describe the different types of intermolecular forces.
24. Balance the following equations.
a. CF4(l) → C(s) + F2(g)
b. H2SO4(aq) + KOH(aq) → KHSO4(aq) + H2O(l)
c. ZnCl2(aq) + H2(g) → Zn(s) + HCl(aq)
d. SO2(g) + H2O(l) + O2(g) → H2SO4(aq)
e. Li(s) + H2O(l) → LiOH(aq) + H2(g)
f. H2CO3(aq) → H2O(l) + CO2(g)
g. Na2SO4(aq) + BaCl2(aq) → BaSO4(s) + NaCl(aq)
h. CH3OH(l) + O2(g) → CO2(g) + H2O(g)
25. Classify each of the above according to the 5 types of reactions (composition, decomposition,
single replacement, double replacement and combustion).
26. Write the formula for each material correctly and then balance the equation. For each reaction,
tell what type of reaction it is.
a) lead (II) nitrate and sodium iodide react to make lead iodide and sodium nitrate.
b) calcium carbonate when you heat it produces calcium oxide and carbon dioxide.
c) propane reacts with oxygen to produce carbon dioxide and water
d) copper metal and silver nitrate react to form silver metal and copper (II) nitrate.
e) sodium metal and chlorine react to make sodium chloride.
27. Calculate the molar mass of the following. Round to 2 decimal places.
A) CO2
B) Na2SO4
28. Determine the number of moles in:
(a) 25.0 g of NaCl
(b) 125 g of H2SO4
29. Determine the number of grams in:
(a) 1.70 moles of K2O
(b) 0.25 moles of KCl
30. Determine the number of moles in:
(a) 10.0 L H2(g) at STP
(b) 48.5 L O2(g)) at STP
31. Determine the number of grams in:
(a) 15.0 L Ne(g) at STP
(b) 44.0 L He(g) at STP
32. For the reaction
N2(g) + 3H2g) → 2NH3(g)
(a) How many moles of hydrogen are needed to completely react with two moles of nitrogen?
(b) How many grams of hydrogen are necessary to react completely with 50.0 g of nitrogen?
33. Write and balance the chemical equation: Sodium hydroxide and hydrochloric acid combine to
make table salt and water. 14.0 g of sodium hydroxide is added to an excess of acid. How many
moles of table salt are made? How many grams of salt is that?
34. Explain the difference between a coefficient number and a subscript number.
35. In the atmosphere, the air pollutant nitrogen dioxide reacts with water to produce nitric acid
according to the unbalanced equation:
NO2(g) + H2O(l) →
3(aq) + NO(g)
What volume of NO2 will react with 36.04 g of water at STP?