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Transcript
Semester I CP Chemistry
Review
2009/2010
1. What has to happen for a hypothesis to
become a theory?
It has to be supported by many, many
experiments.
2. Compare the size and mass of the
nucleus in an atom. (where is most of
the mass of an atom found?)
The atom is mostly empty space. Although
the nucleus occupies only a small part of
the atom’s volume it contains most of its
mass.
3. All atoms of the same element have the
same number of what subatomic
particle?
Protons
4. What subatomic particles are in the
nucleus and outside the nucleus?
In nucleus = protons + neutrons
Orbiting around nucleus = electrons
5a. 131I undergoes beta (β )emission
13153I  13154Xe + 0 -1β
(a loss of a beta particle = a gain of 1 in
atomic # and a no change in mass)
5b. 230Th undergoes alpha(α)
emission
23090Th  22688Ra + 4 2α
4 2α aka 4 2He
A loss of 4 atomic mass and a loss of
2 in atomic #
230
5c.
90Th undergoes gamma emission
23090Th  23090Th + gamma
No change just release of energy
6. Put the three types of decay (alpha,
beta, gamma) in order from biggest to
smallest. Then circle the type of
decay that is the easiest to block.
Alpha particles are the biggest and
gamma the smallest
Alpha easiest to block, gamma harder to
block
Alpha low energy and gamma high energy
7. What are alpha particles? Beta
particles? Gamma particles?
Alpha are helium nuclei
Beta are electrons
Gamma are photons
8. What occurs in a fusion reaction? What
occurs in a fission reaction? Give an
example of each.
Fusion – two smaller nuclei combine to
make one larger nucleus
Fission – One large nucleus breaks down
into 2 or + smaller nuclei
9. How does the energy produced in a
fission or fusion reactions compare to
the energy produced in a chemical
reaction?
Both fission and fusion release A LOT
more energy than a chemical reaction.
18. What is the force that holds the
nucleus together called?
Strong nuclear forces
19. What is an isotope?
Atoms with the same # of protons but
different # of neutrons.
20. How do you know if an isotope of an
element is radioactive or stable?
(remember there is ratio, one for
elements #20 and under and one for
those larger than #21)
For elements #20 > the ratio of protons to
neutrons is 1:1, then stable
For elements #21 < the ratio of protons to
neutrons is 1:1.5, then stable
Or if in band of stability on graph (dots
part)
20.b Describe how an atom of Li would be
changed for each of the following
situations:
a. Li atom loses or gains a neutron – loses
or gains atomic mass
 b. Li atom loses or gains an electron –
loses becomes a + ion, gains becomes a –
ion.
 c. Li atom loses or gains a proton –
becomes a different element
22. Identify the element from the given
electron configuration
 1s2 2s2 2p6 3s2 3p5 – add p’s = 17 so Cl
 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 – p’s
= 38, Sr
 1s2 2s2 2p4 – p’s = 8, O
23. Draw the electron dot (Lewis dot)
diagram that indicates the number of
valance electrons for #22 a, b and c.
Cl
Sr
O
24. Bromine (Br) commonly has a -1 charge while
Potassium (K )has a +1 charge. Explain why
this is and how it relates to the noble gas
configuration. Why do ions commonly lose or
gain ions to have a noble gas configuration?
 Br has 7 valence electrons and would like 8 to
have a full outer orbital just like the noble gases,
so it steals 1 electron away from another atom.
 K has 1 valence electron and would like to lose it
so then it would be left with the next lower
energy level which is full, just like a noble gas.
So K likes to give its 1 valence electron away.
25. Indicate which element has the
greater atomic radius and give a reason
for your choice.
 a. Na or Li - Na because atomic radius
increases as you go down the group, plus Na
has 3 energy levels and Li only has 2.
 b. Sr or Mg – Sr. Sr has 5 energy levels and
Mg only has 3.
 c. C or Ge – Ge has 4 energy levels, C has 2
 d. Se or O – Se has 4 energy levels, O has 2
32. What is the difference between a
group and a period on the periodic table?
A group is a column and all in a group
have the same # of valence electrons and
behave similarly in a chemical reaction.
A period is a row and tells you how many
energy levels the element has.
33. What is electronegativity and what is
its trend down a group and across a row
in the periodic table?
The ability of an atom to attract electrons
to itself.
F has highest, Fr lowest
Increases from left to right. Decreases as
you go down a group.
34. For each of the following pairs which
atom/ion will have the HIGHER
ionization energy? Explain.
a. Ca or Ba? - Ca is higher because it has
only 4 energy levels so its valence
electrons are closer to its nucleus
b. Al or Cl? – Cl, they both have 3 energy
levels but Cl has more protons attracting
its electrons
35. Where are the metals, nonmetals and
metalloids on the periodic table?
Metals to the left, below staircase
Nonmetals to the right, above staircase
Metalloids along the staircase
36. How many valence electrons does the
atom have?




a. Na – 1 (it’s in group IA)
b. N – 5 (it’s in group VA)
c. Ga - 3 (it’s in group IIIA)
d. Mg – 12 (it’s in group IIA)
37. What type of elements form ionic
bonds ? Describe what happens to
the electrons in an ionic bond? Why
do the ions in an ionic compound stay
together?
Ionic – metal with nonmetal
Electrons are transferred from metal to
nonmetal
Metal becomes + and nonmetal becomes –.
Opposites attract. Electrostatic forces.
36. What type of elements form
covalent bonds? What are the
electrons doing in a covalent bond?
2 or more nonmetals
Electrons are being shared between the
atoms
37. Describe what the electrons are doing
in a metallic bond. What types of atoms
form metallic bonds?
Sea of electrons. Valence electrons are
floating from one atom to another.
Metallic bonds form only between metal
atoms!!
38. What type of ions to metals form?
How about non-metals?
Metals – cations (+ ions)
Nonmetal – anions (- ions)
39. What is a coefficient?
The number before a formula in a
chemical equation. It tells you how many
atoms or molecules of that substance is
involved in the reaction.
40. Balance the following equations
 2 Al +
2 H3PO4 
 __P4 + __S8 
2AlPO4 +
3 H2
__ P4S8
 2 C2H6 + 7 O2 
4 CO2 +
 2 AgNO3 + ___K2S 
6 H2O
___Ag2S + 2 KNO3
 ___H2CO3  ____H2O + ____CO2
40b. Why must chemical equations be
balanced?
Because matter cannot be created nor
destroyed in a chemical reaction.
So the numbers of elements in the
reactants must equal the numbers of
elements in the products
41. What is the mass in grams of each of
the following:
 a. 1 mole of Al = 27 g

b. 0.5 moles of S = 16 g

c. 1 mole of NaOH = 40 g (Na = 23 +
o = 16 + H = 1)

d. 3 moles of O2 = 32 g (O = 16 x 2)
You must use the atomic masses from the
periodic table
42. How many moles are in the following:
 b. 22 grams of N2O
22g N2O 1 mole N2O = 0.5 mole
44g N2O
 b. 342 grams of Al2(SO4)3
 342 g Al2(SO4)3
1 mole Al2(SO4)3 = 1 mole
342 g Al2(SO4)3
c. 10 grams of Fe2O3
10g Fe2O3
1 mole Fe2O3 = 0.06 mole
160 g Fe2O3
43. How many atoms of molecules are in
the following?
 a. 2 moles of sugar – 1.2 x 10 24
 b. 0.5 moles of H2O – 3.0 x 10 23
 c. 42 grams of C3H6
42g C3H6 1 mole C3H6 6.0 x 10 24molecules
42 g C3H6
1 mole C3H6
= 6.02 x 10 23
43b. How many atoms in
 a. 32 g of O
32g O 1mole O 6.02 x 10 23 atoms
16 g O
1 mole O
= 1.2 x 10 24 atoms
 b. 23 g of Na
23g Na
1mole Na 6.02 x 10 23 atoms
23g Na
1 moleNa
= 6.02 x 10 23 atoms
c. 36 g of C
36g C
1mole C
12 g C
= 1.8 x 1024 atoms
6.02 x 10 23 atoms
1 mole C
43c. How many atoms in 1 mole of Br?
= 6.02 x 1023 atoms
How many molecules in 1 mole of Fe2O3?
= 6.02 x 1023 molecules
44. Calculate the volume of these gases
at STP
a. 1.5 moles of Neon
1.5 mole Ne 22. 4 L
= 33.6 L
1 mole Ne
b. 32 grams of oxygen
32g O
1 mole O
22.4 L
= 44.8 L
16g O
1 mole O
48. Given the following balanced
equation: 2NH3 + 3O2 + 2CH4 
2HCN + 6H2O
 A. How many moles of HCN are
produced from 9 moles of O2?
9 mole O2 2 mole HCN = 6 moles HCN
3 mole O2
 B. How many moles of NH3 are needed
to produce 24 moles of H2O?
24 moles H2O 2 moles NH3 = 8 moles H2O
4 moles of H2O
49. Given the following balanced
equation: SiO2 + 4HF  SiF4 + 2H2O
a. If 5 grams of HF reacts with SiO2,
what mass of H2O forms?
5g HF
1 mole HF
20 g HF
= 2.25 g H2O
2mole H2O 18g H2O
4 mole HF
1 mole H2O
b. If 120 grams of SiO2 reacts with excess
HF, how many grams of water will form?
120g SiO2
= 72 g H2O
1mole SiO2
60g SiO2
2mole H2O 18g H2O
1 mole SiO2 1mole H2O