Chemistry CP Midterm Exam Review Guide
Atomic Structure
1. What is an atom’s atomic number?
2. What is an atom’s mass number?
3. Define the term isotope of an element. How do we name an isotope?
4. Draw the nucleus of a nitrogen-15 atom. Make a key for the protons and neutrons.
5. Complete the following table.
Isotope Symbol
40
19
K
18
9
F
Isotope Name
Atomic Number
16
Mass Number
23
Number of Protons
12
Number of Neutrons
15
Number of Electrons
16
10
6. Chlorine has two major isotopes: Chlorine-35 and Chlorine-37.
a. Using the atomic mass on the Periodic Table compare the natural abundance of the two
isotopes of chlorine.
b. Write the symbol notation for the heavier isotope of chlorine.
(Make sure you include the atomic number, mass number, and symbol.)
Chemistry CP Midterm Exam Review Guide
7. Using the picture below answer the following questions
What is the mass number of the atom shown?
Name this isotope:
Write the isotope notation for this atom:
Light, Energy, and Electron Configuration
Bohr’s Atomic Structure:
1. How is light produced?
2. What is the meaning of ground state and excited state? Look at the pictures below and label it
accordingly.
3. What occurs when an electron absorbs energy?
Chemistry CP Midterm Exam Review Guide
4. Explain how light is produced in a neon sign. Your answer should include a discussion of energy
levels in the atom, the transition of electrons between energy levels, and the absorption and
release of energy.
Use the pictures below to support your answer, complete it and label it. How is it possible for one
atom to give off more than one type of radiation?
5. Energy Emission (Light):
When the hydrogen atom is in its ground state, the single electron is in the n = 1 orbit (no energy
radiated). When energy is absorbed, the electron moves to a higher-energy orbit (raises atom to an
excited state).
When the electron drops from the ________________-energy orbit to a ________________-
energy orbit the atom emits a ________________ (electromagnetic radiation).
6. Define Valence Electrons:
7. Fill in the table below:
Element
Number of
valence electrons
Na
Al
Pb
P
Chemistry CP Midterm Exam Review Guide
8. Fill in the following chart. Draw the Bohr diagram for the isotope. Use rings to represent each
energy level with appropriate number of electrons in each energy level. Include the proper number
of protons and neutrons in the nucleus.
sulfur-32
(Bohr model)
valence
electrons
Lewis dot
Structure
lose or gain
electrons to
form ion
isotope
notation
𝑨
𝐙
X
9. Electron configuration:
Write the electron configuration notation and the Noble gas electron configuration notation for
each of the elements listed below.
Element
Electron configuration
Noble gas configuration
beryllium (Be)
oxygen (O)
sodium (Na)
sulfur (S)
chlorine (Cl)
Chemistry CP Midterm Exam Review Guide
Periodic Table and Trends
1. The elements in the modern periodic table are arranged according to increasing ______________
number as a result of the work of Henry Moseley. This arrangement is based on the number of
__________________ in the nucleus of an atom of the element. The modern form of the periodic
table results in the periodic law, which states that when elements are arranged according to
increasing atomic number, there is a periodic repetition of their chemical and physical
________________________.
2. Which has the largest atomic radius: carbon (C), fluorine (F), beryllium (Be), or lithium (Li)?
3. Explain your answer in terms of trends in atomic radii.
4. Describe ionization energy trends on the periodic table by completing the statements below.
a. Ionization energy generally ______________ as you move left-to-right across a
______________. Increased nuclear charge leads to increased __________________
attraction on valence electrons.
b. Ionization energy generally _______________ when you move down a ____________. Less
energy is required to remove _______________ electrons because they are
_______________ from the nucleus.
c. The Octet Rule states that atoms tend to gain, lose, or share _______________ in order to
acquire a full set of _______________________. First period elements are the
_______________ to this rule.
Ionization energy is the amount of energy required to remove an electron from an atom.
Knowing this, look at the following successive ionization energies (in
KJ
/
mole
) for an element.
Ionization Energy (kJ/mol)
1
st
IE
2
nd
IE
3
rd
IE
4
th
IE
577
1820
2740
11600
5. Based off of the values of successive ionization energy
a. Explain how many valence electrons this element has
b. Justify how you determined your answer.
Chemistry CP Midterm Exam Review Guide
Reactivity:
6. The most reactive metals are located in group ________________.
7. The most reactive nonmetals are located in group ______________.
8. Why are the noble gases nonreactive?
9. Which element is more reactive, Magnesium or Barium? Support your answer with information
about ions formed, and the periodic trends
10. The first and second ionization energies of magnesium are both relatively low, but the third
ionization energy jumps to five times the previous level. Explain the reason.
11. What does the electronegativity of an element indicate?
12. Identify two reasons why the relative size of an ion becomes smaller due to the loss of electrons:
Chemical Bonding and VSEPR
The electronegativity difference (EN) determines the bond character between atoms
Electronegativity difference
(EN)
< 0.4
0.4 β€” 1.7
> 1.7
Bond character
Non-polar covalent
Polar covalent
ionic
1. Distinguish between ionic, polar covalent and nonpolar covalent bonds in terms of what occurs
with the electrons when two atoms bond.
Chemistry CP Midterm Exam Review Guide
2. Describe what occurs in terms of electrons when a sodium atom reacts with a chlorine atom to
form sodium chloride. Be specific.
3. Define Electronegativity
4. In terms of difference in electronegativity (EN), how is a covalent bond distinguished from an
ionic bond?
5. When two atoms form an ionic bond, which atom, the one with higher or the one with lower
electronegativity, becomes the negative ion?
6. What is a cation? What is an anion? Give examples.
7. In terms of difference in electronegativity (EN), how is a nonpolar covalent bond distinguished
from a polar covalent bond?
8. When a bond is determined to be polar covalent, one atom is said to be partially positive (+),
while the other is partially negative (-). How do you determine which is which?
9. For each of the following bonds, determine EN and whether a bond between the two atoms
results in a nonpolar covalent, polar covalent, or an ionic bond.
Electronegativity Values for Selected Elements
Na = 0.93
Cl = 3.16
P = 2.19
Ca = 1.00
O = 3.44
C = 2.55
H = 2.20
Na with Cl in NaCl
EN =
Bond Type:
P with Cl in PCl
3
EN =
Bond Type:
Ca with O in CaO
EN =
Bond Type:
C with O in CO
2
EN =
Bond Type:
H with O in H
2
O
EN =
Bond Type:
Chemistry CP Midterm Exam Review Guide
10. Complete the following:
Elements
carbon dioxide
dihydrogen monosulfide
Ξ”EN
Bond
type
Lewis
Structure
Elements
aluminum (Al) and chlorine(Cl)
barium (Ba) and fluorine (F)
Ξ”EN
Bond
type
Lewis
Structure
Chemistry CP Midterm Exam Review Guide
11. Draw the Lewis Structure for each of the following covalent molecules and determine the VSEPR
shape of each:
Molecule
Total #
valence
e-
Lewis
structure
#
Electron
domains
Molecular
shape
Molecular
polarity
PH
3
NH
3
CF
4
CO
2
Properties of ionic compounds:
12. When dissolved in water, ionic compounds _________________________.
13. Ionic compounds can conduct electricity as they contain _________________________.
14. Ionic crystals have high melting points and high boiling points because their bonds are relatively
_________________________.
15. A three-dimensional arrangement of particles in an ionic solid is called a(n) _________________.
16. In electron transfer involving a metallic atom and a nonmetallic atom during ion formation, which
of the following is correct?
A. The metallic atom gains electrons from the nonmetallic atom.
B. The nonmetallic atom gains electrons from the metallic atom.
C. Both atoms gain electrons.
D. Neither atom gains electrons.
Chemistry CP Midterm Exam Review Guide
Intermolecular Forces
οƒ˜ London dispersion: force which results from the temporary shift of electron cloud in a molecule.
οƒ˜ Dipole – dipole: polar molecules contain permanent dipoles (+ and -)
οƒ˜ Hydrogen bond: dipole interaction that specifically occurs between molecules that contain
hydrogen atoms bonded to either fluorine, nitrogen or oxygen (hydrogen just wants to have FON)
οƒ˜ Ion – dipole: occurs between ionic substances dissolving in a polar solution
1. What are intramolecular forces and intermolecular forces?
2. Number the intermolecular forces below from 1 to 4 with 1 being the strongest and 4 being the
weakest:
Dipole-Dipole ______
Hydrogen bonding _____
Dispersion Forces _____
Ion-Dipole _____
3. In a polar molecule, which atom will have the greatest partial negative charge (-)?
4. Illustrate the interaction between HI molecules. What is the force of interaction between them?
5. As shown in the figure below, which intermolecular force depends on the formation of temporary
dipoles?
6. Draw the interaction between 4 molecules of water. What is intermolecular force between them
called?
7. Write the dissociation equation for magnesium chloride (MgCl
2
) dissolving in water.
_____________ (s) β†’ ________ (aq) + ________ (aq)
8. What is the intermolecular force involved in dissolving of the substance?
Chemistry CP Midterm Exam Review Guide
9. Illustrate the interaction. You MUST include the Lewis structure for at least six (6) water
molecules in your picture indicating the intermolecular force(s) of attraction between the water
and MgCl
2
.
10. If you spilled a few drops of octane (C
8
H
18
) on a lab bench, would you expect it to form beads of
liquid on the surface, or spread out evenly? Why? What about water?
11. Put the following compounds in order from the lowest melting point to the highest melting point:
H
2
O, C
2
H
6
, He, CaO, K
2
O
12. Investigate to find out what factors influences the state of matter of a substance
Substance
Melting
point (Β°C)
Boiling
point (Β°C)
Polarity
Intermolecular
force
State of matter
at room
temperature
(22 Β°C)
Fluorine- F
2
-219.62
-188.14
Chlorine- Cl
2
-100.98
-34.6
Bromine- Br
2
-7.2
58.8
Iodine- I
2
114
183
Methane- CH
4
-162
-182
Methanol- CH
3
OH
-98
64.7
a. Based on the information that you have, what factors can be used to form a general rule to
determine if a substance is a solid, liquid, or a gas at room temperature?
b. Which substance has the lowest boiling point and melting point? Why?