Download PAP Chemistry - Fall Final Review

PAP Chemistry - Fall Final Review
Chapter 1 (Matter) & 2 (Measurements)
1. Know the rules for significant figures (what’s significant and what’s not)
2. Be able to determine the number of significant figures present in a given number
a. 0.00203 3
b. 123
3
c. 100
1
d. 100.
3
e. 100.050 6
Chapter 3 (Atoms)
3. Understand and apply the Law of Conservation of Mass
4. Understand the Law of Multiple Proportions
5. Define atom, nucleus, electron, neutron, proton (including relative size and charge of
subatomic particles)
6. What did Rutherford discover from the Gold Foil Experiment – p.72
The nucleus and that the atom was mostly empty space
7. When is a bright-line spectrum produced by an atom? IE – How does an atom give off
color (especially when burned)?
The resting state or the ground state is when the electron is closest to the nucleus.
It will jump to outer levels gaining energy and this is called the excited state. When
the electron falls or drops back to the ground state, it will lose energy in the form of
light. When burned, the heat is the energy lost or given off.
8. Use the mass number and atomic number to determine the element and its number of
protons, electrons, and neutrons
9. Be able to determine the atomic number and mass number of an element when the
number of protons, neutrons, and electrons is specified
10. How does mass number relate to number of protons when talking about isotopes?
The number of protons does not change. An isotope is the different mass numbers
for the same element.
11. What is Avogadro’s Number? 6.02 X 1023
12. How many atoms are in 1 mole of each element? 6.02 X 1023 Does this number change?
No, since the equation is 1 mole = 6.02 X 1023 atoms
13. Be able to do mol-mass conversions
14. Determine the number of protons, electrons, and neutrons in the following isotopes:
An isotope only deals with the different mass of the same element. The equation is
mass = protons + neutrons. The number of protons will always equal the atomic
number. If it is a neutral atom or an isotope, the number of protons and electrons
are equal. The number of electrons changes if we are talking about an ion.
a. sodium-23
protons = 11, neutrons = 12, electrons = 11
b. calcium-40
protons = 20, neutrons = 20, electrons = 20
c. Cu
protons = 29, neutrons = 35, electrons = 29
d. Ag
protons = 47, neutrons = 61, electrons = 47
15. Determine the mass in grams of the following:
a. 2.00 mol N
2.00 mol N2
| 28.014 g N2
= 56.0 g N2
|
1 mol N2
b. 3.01 x 1023 atoms Cl
3.01 x 1023 atoms Cl2 |
1 mol Cl2
| 70.905 g Cl2 = 3.55 x 101 g Cl2 or 35.5 g
23
| 6.02 x 10 atoms Cl2 |
16. Determine the amount of moles of the following:
a. 12.15 g Mg
12.15 g Mg
|
1 mol Mg = .4999 mol Mg
|
24.305 g Mg
b. 1.50 x 1023 atoms F
1.50 x 1023 atoms F2 |
1 mol F2
| 37.996 g F2
= 9.47 g F2
| 6.02 x 1023 atoms F2 | 1 mol F2
17. Determine the number of atoms present in a 251.2 g sample of manganese.
251.2 g Mn
| 6.02 x 1023 atoms Mn = 2.74 x 10 24 atoms Mn
|
54.938 g Mn
Chapters 4 (Electrons in Atoms) & 5 (Periodic Law)
19. How many electrons are held is each quantum sublevel (s = 2, p = 6, d = 10, f = 14)
20. Write electron configurations and orbital filling diagrams (including energy level,
sublevels, and orbitals for any element)
Example: chlorine
Electron configuration: 1s2 2s2 2p6 3s2 3p5
Orbital notation:
↑↓
1s
↑↓
2s
↑↓
↑↓
↑↓
↑↓
3s
↑↓
↑↓
↑ .
2p
3p
21. Be able to locate the alkaline earth, alkali metals, noble gases, and halogens
22. Periodic Trends – know atomic radius and electronegatitivity and how they compare
IE = increase across, decrease down ( ,  )
EA = increase across, decrease down ( ,  )
EN = increase across, decrease down ( ,  )
AR = decrease across, increase down (  ,  )
IR = decrease across, increase down (  ,  )
SHIELDING = remains constant across, increases down ( constant  ,  )
23. What is the octet rule? What orbitals must be filled for it to be satisified?
There must be 8 electrons in the outermost shell, orbital. The “s” and “p” orbitals,
24. How is the periodic table organized (atomic number, atomic mass)
By increasing atomic number.
25. Define periodic law and why is it important?
26. What is the general electron configuration for a noble gas? (s, p, d, f)
Example: noble gas configuration for strontium
[ Kr ] 5s2
27. Write the electron configuration for the following
a. carbon
1s2 2s2 2p2
b. neon
1s2 2s2 2p6
c. sulfur
1s2 2s2 2p6 3s2 3p4
d. potassium
1s2 2s2 2p6 3s2 3p6 4s1
e. vanadium
1s2 2s2 2p6 3s2 3p6 4s2 3d3
Chapter 7 (Formulas and Compounds)
39. Be able to name compounds
a. Pb3(PO4)2
lead II phosphate
b. NaF
sodium fluoride
c. C2O
dicarbon monoxide
d. H2SO4
sulfuric acid
e. HF
hydrofluoric acid
40. Be able to convert between gramsmolesatoms.
a. How many grams of Al2S3 are in 2.00 moles of Al2S3?
2.00 mol Al2S3
|
150.162 g Al2S3
= 300. g Al2S3
|
1 mol Al2S3
b. How many atoms are found in 1.00 moles of Na?
1.00 mol Na
|
6.02 x1023 atoms Na
|
1 mol Na
= 6.02 x 1023 atoms Na
c. How many Na+ ions are found in 1.00 moles of NaF? OMIT
41. In a compound, the sum of all the oxidation number of the atoms must equal? zero
42. In an ion, the sum of all of the oxidation numbers of the atoms must equal? The charge
on the ion.
43. What is the oxidation number of C in CO2?
CO2
oxygen has a -2 charge so multiple it by its subscript
.
-4 = 0
What number does it take to add to -4 to get zero? +4
+4 -4 = 0
44. What is the oxidation number of P in PO43PO4-3
the -3 is the charge of the polyatomic ion so the whole formula
must equal a -3
PO4
oxygen has a -2 charge so multiple it by its subscript
.
- 8 = -3
phosphorus need a +5 in order for the equation to
+5 - 8 = -3
produce a -3
45. What is % composition?
% composition =
part
x 100
whole
46. Determine the % composition of NaCl.
% composition =
Na wt
x 100
Na wt + Cl wt
=
22.990
x 100
22.990 + 35.453
47. Determine the % composition of magnesium hydroxide.
% composition =
Mg wt
Mg wt + (O wt x 2) + (H wt x 2)
=
=
% composition =
=
24.305
24.305 + 31.998 + 2.016
x 100
x 100
41.676% Mg
(O wt x 2)
Mg wt + (O wt x 2) + (H wt x 2)
31.998
24.305 + 31.998 + 2.016
x 100
x 100
=
% composition =
=
54.867% O
(H wt x 2)
Mg wt + (O wt x 2) + (H wt x 2)
2.016
24.305 + 31.998 + 2.016
x 100
x 100
=
3.457% H
48. What is empirical and molecular formula?
Empirical – a formula with the subscripts in the lowest terms
Molecular – a formula that is a multiple of the empirical formula
49. Be able to determine the empirical formula.
a. What is the empirical formula for a compound that is found to contain 63.52% iron
and 36.48% sulfur? (If given the number as a percent, you divide the % number by
the weight of the element. If the given number is in grams, you convert to moles)
Step 1 : divide to get a number
.6352 Fe
= .01137
55.845
.3648 S = .01137
32.066
Step 2: divide each number above by the smallest number
.01137 = 1
.01137
.01137 = 1
.01137
Step 3: these numbers are the subscripts for the elements. So, the empirical
formula is: FeS
50. Given the empirical formula and a molar mass, be able to determine the molecular
formula.
a. The empirical formula is CH5. Its formula mass is 85.0 g/mol. What is the molecular
formula? (Since you are already given the empirical formula, you just have to
do the step to calculate the molecular formula. The equation you use is the “me”
equation:
molecular wt = a # that is the multiplier
empirical wt
85.0
Wt of C + (5 x wt of H)
85.0
17.051
=
= 4.98 = round this to 5
So, the take the 5 and multiply each subscript by it. The molecular formula is C5H25
Chapter 8 (Rxns)
51. What is a precipitate? A solid in a liquid solution
52. When is a chemical equation balanced?
When it follows the Law of Conservation: the same number of each atom on both
sides of the arrow.
53. What is a subscript? # of atoms What is a coefficient? # of mol Where are the products
and reactants located in a chemical reaction?
54. Be able to balance chemical equations.
a. Al4C3 + H2O  CH4 + Al(OH)3
Al4C3 + 12H2O  3CH4 + 4Al(OH)3
55. Be able to recognize a synthesis reaction, a single replacement reaction, a double
replacement reaction, and a decomposition reaction.
56. When a double replacement reaction produces a precipitate, what happened?
It is insoluble.
57. What products are formed when Lithium metal reacts with water?
Li(s) + H2O(L) → H2(g) + LiOH(aq)
58. What is the activity series of metals? What does it indicate about the loss or gain of
electrons?
It tells which elements can or cannot replace which in a single replacement reaction.
Those elements that can lose few electrons are higher on the activity series so they
Are stronger and can replace other easier.
59. Given solubility rules, be able to determine what products will form precipitates (see the
lab that took forever)