Download Goal 1 Study Guide

Goal 1 Study Guide and Practice Problems
1. Fill in the following table:
Particle
Protons
Neutrons
Elections
Location
Nucleus
Relative Charge
Relative Mass (amu)
0
1/2000 amu
2. Isotope notations: 235
92𝑈 and U-235 represents the same isotope of an element.
a. What does “235”represent?
b. What does “92” represent?
c. What does the U represent?
3. What information does the mass number you about the nucleus of an atom?
4. What information does the atomic number tell you about the nucleus of an atom?
5. How would you determine the number of protons, electrons, and neutrons contained in a neutral
atom?
6. Practice Problems: Atomic Structure Math
Isotope
symbol
Isotope
name
Chlorine-35
Protons
Neutrons
Electrons
20
20
5
4
45
Atomic
number
Mass
number
92
35
235
7. How do isotopes of the same element differ and how are they similar?
8. How do atoms of different elements differ?
9. How are average atomic mass, actual isotopic mass, and the mass number of specific isotopes
different?
10. Why does a weighted average have to be used to calculate the average atomic mass?
11. What two factors does average atomic mass depend on?
12. A Bohr Model can be effective for describing the number of energy levels and the number of valence
electrons for only the first 18 elements, in a Bohr model how many electrons are located in each of
the first three energy levels?
13. Practice Problems: Drawing Bohr Models
a. sodium
b. sulfur
c. aluminum d. fluorine
e. calcium
f. nitrogen
14. An electron configuration can be effective for describing the number of energy levels and the
number of valence electrons for all of the elements. In the quantum mechanical model, what is the
maximum number of electrons allowed in each of the first four energy levels?
15. Fill in the blanks:
a. When an electron ____________________ an amount of energy equivalent to the energy
difference, it moves from its ground state to a higher energy level.
b. When the electron moves to a lower energy level, it ____________________ an amount of
energy equal to the energy difference in these levels as electromagnetic radiation.
c. Since the light that is released by an electron is constant (speed of light), the wavelength
and frequency of light are ____________________ related. Which means that as the
frequency of light increases, the wavelength ____________________. (c=f*λ)
d. Energy and frequency of light are ____________________ related, which means the as the
frequency of light increases, the energy ____________________. (E=h*f)
e. Niels Bohr produced a model of the hydrogen atom based on experimental observations.
This model indicated that:
i. An electron circles the nucleus only in ____________________ energy ranges called
orbits.
ii. An electron can neither gain nor lose energy inside this orbit, but could move up or
down to ____________________ orbit.
iii. The lowest energy orbit is ____________________ to the nucleus.
16. Complete the following table:
Symbol
Change in mass #
Change in atomic #
Penetrating ability
Alpha
Beta
Gamma
4
2𝐻𝑒
0
−1𝑒
γ
Decreases by 4
Paper
No change
Increases by 1
Wood
Lead or concrete
17. Practice Problems: Balancing Nuclear Equations
a. Uranium-235 (alpha decay)
b. Carbon-14 (beta decay)
18. Practice Problems: Half-life
a. In 5.49 seconds, 1.20 g of argon-35 decay to leave only 0.15 g. What is the half-life of
argon-35?
b. How many days does it take for 16 g of palladium-103 to decay to 1.0 g? The half-life of
palladium-103 is 17 days.
c. Sodium-24 has a half-life of 15 hours. How much sodium-24 will remain in an 18.0 g
sample after 60 hours?
19. How are radioactive decay, fission, and fusion different?
20. Fill in the blanks:
a. When a neutral atom ____________________ one or more electrons it becomes a
negatively charged ion, which is called a(n) ____________________.
b. When a neutral atom ____________________ one or more electrons it becomes a positively
charged ion, which is called a(n) ____________________.
21. Fill in the table:
Ionic
What happens to the
electrons?
Types of elements
ΔEN
Melting point
Boiling point
Conductivity
Other
Covalent
Metallic
Shared
Sea of electrons
Metal/nonmetal
>1.7
n/a
High
High
In a molten state or
aqueous solution
Brittle solid
22. Practice Problems: Draw Lewis structures for simple compounds.
a. SO2
b. CO3-2
c. CO2
d. ClO4-1
e. ClO2-1
23. Practice: Predicting Shapes and Bond Angles for molecular compounds
f. H2O
24.
25.
26.
27.
28.
29.
30.
31.
32.
a. SO2
b. CO3-2
c. CO2
d. ClO4-1
e. ClO2-1
f. H2O
Practice: Predict bond polarity
a. S-O
b. C-O
c. Cl-O
d. H-O
e. F-F
f. H-Cl
Rank single, double, and triple bonds in terms of strongest to weakest and longest to shortest.
Rank in order of strongest bond to weakest: Ionic bonds, metallic bonds, hydrogen bonds, dispersion
forces, dipole-dipole, and covalent.
Classify the following as intermolecular or intramolecular bonds: dispersion forces, metallic bonds,
covalent, ionic bonds, covalent bonds, and dipole-dipole.
Intermolecular forces for molecular compounds:
a. ____________________: attraction between molecules when H is bonded to N, O, or F.
b. ____________________: attractions between polar molecules.
c. ____________________: electrons of one molecule attracted to nucleus of another
molecule.
Practice: Writing formulas for compounds
a. Potassium oxide
f. silver oxide
k. barium fluoride
b. Silver nitrate
g. zinc chlorate
l. iron (III) nitride
c. Copper (II) nitride
h. iron(II) nitride
m. diphosphorus trioxide
d. Nitrogen trihydride
i. born trifluoride
n. ammonium carbonate
e. Lithium hydroxide
j. aluminum oxide
o. ammonium phosphate
Practice: Naming compounds
a. KOH
f. LiMnO4
k. Mg(NO3)2
b. Rb2SO4
g. RaCl2
l. BeS
c. N2O3
h. Cu(NO3)2
m. N2O5
d. CS2
i. (NH4)2SO4
n. F2O5
e. Zn(ClO3)2
j. Fe2(CO3)3
o. Fe(ClO3)3
____________________: vertical columns on the periodic table
Fill in the following table:
Representative Elements
Property
Number
of V.E.
Oxidation
Number
Family
Name
Gain or
lose
electrons
1A
2A
3A
4A
5A
6A
7A
8A
1
-2
Alkaline
earth
metals
n/a
n/a
n/a
Lose 3
Lose/gain
4
Gain 3
n/a
33. Reactivity ____________________ as you go down within a group for metals and
____________________ for nonmetals.
34. ____________________: horizontal rows on the periodic table.
35. Classify elements as metals, nonmetals, and metalloids based on location.
36. Classify elements as representative elements and transition elements.
37. Atomic radius ____________________ as you go down a group, and ____________________ as you
go across the period left to right.
38. The ionic radius of an anion is ____________________ than the atomic radius of its neutral atom.
The ionic radius of a cation is ____________________ than the atomic radius of its neutral atom.
39. Practice: Electron configurations
a. Write the electron configurations of the following elements:
i. Sodium
vi. Iron
ii. Bromine
vii. Barium
iii. Neptunium
viii. Cobalt
iv. Silver
ix. Tellurium
v. Radium
x. Lawrencium
b. Determine what elements are denoted by the following electron configurations:
i. 1s22s22p63s23p4
ii. 1s22s22p63s23p64s23d104p65s1
iii. [Kr]5s24d105p3
iv. [Xe]6s24f145d6
v. [Rn]7s25f11
40. Identify the s, p, d, and f blocks on the periodic table.
41. Ionization energy ____________________ as you go down a group, and ____________________ as
you go across the period left to right.
42. Electronegativity ____________________ as you go down a group, and ____________________ as
you go across the period left to right.