CK-12 Chemistry Concepts - Intermediate
Answer Key
Chapter 6: The Periodic Table
6.1 Early History of the Periodic Table
Review
Questions
1. List some elements known since ancient times.
2. What properties were the basis of the triad system?
3. Why did Dobereiner believe that lithium, sodium, and potassium belonged in a
triad?
4. What was a shortcoming of the triad system?
5. How did Newlands arrange the element?
6. What was a problem with the “Law of Octaves”?
Answers
1. Copper and lead.
2. Physical and chemical properties.
3. The average of the atomic weights for lithium and potassium was approximately
equal to the atomic weight for sodium.
4. By atomic mass
5. Some elements were missing and the system did not work for elements above
calcium.
6.2 Mendeleev’s Periodic Table
Review
Questions
1.
2.
3.
4.
5.
When did Mendeleev publish his periodic table?
Who else came out with a periodic table at about the same time?
Why was Mendeleev’s table considered to be superior?
What element did Mendeleev predict to exist?
What element was named after Mendeleev?
Answers
1. 1869.
2. Lothar Meyer.
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3. Mendeleev’s table allowed for predictions.
4. Gallium.
5. Mendelevium.
6.3 Periodic Law
Review
Questions
1. Did Mendeleev know about the nucleus of an atom?
2. Who discovered the relationship between wavelength of X-rays and atomic
number?
3. What did Moseley conclude from his research?
4. What is the “periodic law”?
5. What do the vertical columns (groups) in the periodic table represent?
Answers
1. No
2. Henry Moseley.
3. Elements of the periodic table should be arranged in order of atomic number
instead of atomic weight.
4. When elements are arranged in order of increasing atomic number, there is a
periodic repetition of their chemical and physical properties.
5. Elements with similar chemical properties.
6.4 Modern Periodic Table
Review
Questions
1. How is today’s periodic table different from the one that Mendeleev published?
2. Are all the elements in today’s periodic table naturally occurring? Explain your
answer.
3. What is a “period”? What does it represent?
4. What is a “group”? What does it represent?
5. Why are there two different numbering systems for groups?
Answers
1. New elements are included and elements rearranged based on common
properties.
2. No, some elements have been artificially synthesized.
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3. A horizontal row on the table. Organized by energy levels of electrons.
4. A vertical row on the table. Organized by outer shell electron configuration.
5. U.S. and Europe used two different systems.
6.5 Metals
Review
Questions
1.
2.
3.
4.
5.
What properties of an element are affected by electron distribution?
Define malleability.
Define ductility.
State one reason gold is used in jewelry.
Why is mercury no longer used in many devices?
Answers
1. Chemical properties.
2. Malleable means that they can be hammered into very thin sheets without
breaking.
3. Ductile means that they can be drawn into wires.
4. Gold is soft and easily shaped into a variety of items.
5. It was commonly used until it was discovered that it is highly toxic.
6.6 Nonmetals
Review
Questions
1. What are the properties of nonmetals?
2. List the states of matter in which nonmetals can exist and give one example of
each state.
3. What are the physical properties and uses of sulfur?
4. What are the physical properties and uses of bromine?
5. What are the physical properties and uses of helium?
Answers
1. Low boiling point, poor conductors of heat, dull and brittle.
2. Solid (carbon and sulfur), liquid (bromine), and gas (nitrogen and oxygen)
3. Sulfur is non-conductive, not reflective, and is used in insecticide and
manufacturing materials such as rubber for tires.
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4. Bromine is versatile and used for water treatment, flame retardant materials, and
in medicines.
5. Helium is a gas that is used in balloons, lasers, and in magnetic imaging
technology.
6.7 Metalloids
Review
Questions
1. Define “metalloid”.
2. Why would it be difficult to decide whether or not an element is a metalloid based
on its properties?
3. Why is silicon used in electronics?
4. What are borosilicates used for?
5. Why is the use of arsenic as an insecticide decreasing?
6. Why is antimony often mixed with lead?
Answers
1. Metalloids are elements with properties that fall between metals and nonmetals.
2. It would be difficult to identify if an element is metalloid based on its properties
because it shares many properties with both metals and nonmetals.
3. It’s electrical conductivity is between that of a metal and a nonmetal.
4. To make glass resistant to rapid changes in temperature known as thermal
shock.
5. Arsenic is now known to be very toxic so alternative insecticides are being
developed for safety reasons.
6. To increase its hardness and strength.
6.8 Blocks of the Periodic Table
Review
Questions
1.
2.
3.
4.
What are the horizontal rows of the periodic table called?
Which sublevel is being filled in period 1?
Which sublevels are being filled in period 7?
How does the electron configuration of an element gives information about the
period it is in?
5. What block of elements has the d sublevels being filled?
Answers
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1.
2.
3.
4.
5.
Periods.
1s
7s, 5f, and 6d
By showing what sublevel is being filled.
Transition metals.
6.9 Hydrogen and Alkali Metals
Review
Questions
1.
2.
3.
4.
What group are the alkali metals and hydrogen in?
What is the outer shell configuration in this group?
How reactive are alkali metals with oxygen?
How reactive are these metals with water?
Answers
1. Group 1.
2. n​s1​ ​.
3. Alkali metals are very reactive. They quickly react rapidly with oxygen to produce
metal oxides.
4. Alkali metals are all very reactive with water.
6.10 Alkaline Earth Metals
Review
Questions
1. Why are these elements known as the “alkaline earth” elements?
2. How many electrons are in the outer shell of alkaline earth elements?
3. Are the alkaline earth elements more or less reactive than the alkali metals?
Explain your answer.
4. Is radium usually considered as part of the alkaline earth category in terms of
chemistry? Explain your answer.
Answers
1. These elements form basic solutions (pH > 7) when dissolved in water.
2. Two.
3. Less reactive. Alkaline earth metals contain two electrons in their valence shell
and require more energy input to remove both.
4. Because Radium is radioactive these properties overshadow its characteristics
as an alkaline earth metal. It does share some common properties with barium.
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6.11 Noble Gases
Review
Questions
1.
2.
3.
4.
5.
6.
What elements comprise the noble elements?
What state are they in at room temperature?
Why is helium non-reactive?
Why were the other noble gases believed to be non-reactive?
When was the first compound formed from xenon?
What happens when an electric current is passed through tubes containing a
noble gas?
Answers
1. Noble gases are extremely unreactive elements in Group 18 including neon,
argon, and helium.
2. Noble gases are monatomic gases at room temperature.
3. Helium has a full shell of two s electrons. Since there are no electrons shielding
this shell from the nucleus, these two electrons will be very difficult to remove,
making helium unreactive.
4. The elements in the group have full outer shells of eight electrons.
5. In 1962, XePtF​6​ was made with xenon.
6. The different gases glow when an electric current is passed through them.
6.12 Halogens
Review
Questions
1. Pick two elements that are halogens. For each, write the name, chemical symbol,
and atomic weight.
2. What does the term “halogen” mean?
3. What is the outer shell electron configuration for halogens?
4. What is the physical state of each halogen at room temperature?
5. Where are the halogens found?
Answers
1. Fluorine, F, 18.998 amu; Chlorine, Cl, 35.453 amu; Bromine, Br, 79.904 amu;
Iodine, I, 126.904 amu; Astatine, At, 210 amu (Radioactive, no average mass).
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2. It means “salt-former” in ancient greek; halogens frequently form salts with alkali
metals and alkaline earth metals.
3. ns​2​np7​
4. Fluorine and chlorine are gaseous. Bromine is a liquid. Iodine and astatine are
solids.
5. Most halogens are found in minerals or dissolved in the oceans.
6.13 Transition Metals
Review
Questions
1. List five different transition elements, giving their name, chemical symbol, and
atomic number.
2. What is unique about the transition elements in terms of electron configurations?
3. Why are these elements often referred to as “d-block” elements?
4. Why do many transition element compounds have bright colors?
Answers
1.
2.
3.
4.
Answers will vary.
The inner-shell d orbitals are not all filled, but the outer s orbitals are.
Because the d orbitals are unfilled.
Electron transitions of inner-shell electrons release energy that can be seen as
light.
6.14 Lanthanides and Actinides
Review
Questions
1.
2.
3.
4.
5.
6.
What electron sublevel is being filled in the lanthanides?
What electron sublevel is being filled in the actinides?
What sublevel is filled just prior to the filling of this sublevel?
Which actinides are found naturally on earth?
List some uses for lanthanides.
List some uses for actinides.
Answers
1. 4​f
2. 5​f
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3.
4.
5.
6.
6​s
Th, Pa, U, Np
Alloys, oil refining, optical devices.
energy sources
6.15 Periodic Trends: Atomic Radius
Review
Questions
1.
2.
3.
4.
5.
Define “atomic radius.”
What are the units of measurement for atomic radius?
How does the atomic radius of different elements change across a period?
How does atomic radius change from top to bottom within a group?
Explain why the atomic radius of hydrogen is so much smaller than the atomic
radius for potassium.
Answers
1.
2.
3.
4.
5.
One-half the distance between two nuclei in a diatomic atom.
picometers.
It decreases from left to right.
It increases from top to bottom.
There are more electrons in the potassium atom. The outer electrons in
potassium are at higher energy levels.
6.16 Ion
Review
Questions
1.
2.
3.
4.
5.
6.
Why are atoms neutral in electric charge?
Define ion.
Compare and contrast cations and anions, and give an example of each.
Describe how ions form.
List properties of ions.
The model in the illustration below represents an atom of lithium (Li). If the lithium
atom becomes an ion, which type of ion will it be, a cation or an anion? What will
be the electric charge of this ion? What will the ion be named? What symbol will
be used to represent it?
Answers
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1. Atoms have equal numbers of positive protons and negative electrons, so they
are neutral in electric charge.
2. Ions are atoms that have a positive or negative charge because they have
unequal numbers of protons and electrons.
3. If atoms lose electrons, they become positive ions, or cations. Example: iron (Fe)
may lose two electrons to form Fe2+. If atoms gain electrons, they become
negative ions, or anions. Example: the fluoride ion F4. Most ions form when atoms transfer electrons to or from other atoms or
molecules. Ionization may also occur when atoms are exposed to high levels of
radiation that gives their outer electrons enough energy to escape from the
attraction of the positive nucleus.
5. Ions are reactive, attracted or repulsed by other charged particles, and deflected
by a magnetic field.
6. Cation, charge 1+, lithium ion, Li​+
6.17 Periodic Trends: Ionization Energy
Review
Questions
1. Define “ionization energy.”
2. Do valence electrons have larger or smaller ionization energies that the
inner-shell kernel electrons?
3. What is electron shielding?
4. Describe the trends in ionization energy from left to right across the periodic
table.
5. Describe the trends in ionization energy from top to bottom of a group in the
periodic table.
6. Why is the second ionization energy for lithium so much larger than the first
ionization energy?
Answers
1.
2.
3.
4.
5.
6.
The energy required to remove an electron from a specific atom.
Smaller.
Outer electrons are shielded from the nucleus by inner electrons.
Ionization energy increases from left to right.
Ionization energy decreases from top to bottom.
After the first electron is lost, there is one more positive charge than there is
negative charge, so there is a stronger attraction by the nucleus.
6.18 Electron Shielding
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Review
Questions
1. Define “electron shielding.”
2. Why do group 13 elements have lower ionization energies than group 2
elements?
3. What influence does a larger shielding effect have on ionization energy?
4. How do ​s​ orbit electrons affect the ionization energy of a ​p​ electron in the same
shell?
5. Why do group 16 atoms have lower ionization energies than the corresponding
group 15 atoms?
Answers
1. The inner electrons partially shield the outer electrons from the attractive force of
the nucleus.
2. Group 13 elements have more inner electrons for shielding.
3. A larger shielding effect leads to easier ionization.
4. The two s electrons provide additional shielding.
5. In group 16, there is one set of paired p electrons, which repel each other and
make it easier to remove one of the electrons.
6.19 Periodic Trends: Electron Affinity
Review
Questions
1. Define “electron affinity.”
2. Does the addition of an electron to a neutral atom require energy to be absorbed
or released?
3. Describe the general trend for electron affinity values moving from left to right on
the periodic table.
4. Describe the general trend for electron affinity values moving from top to bottom
in a group on the periodic table.
5. Why is more energy released in the formation of a halide ion than with other
elements?
Answers
1.
2.
3.
4.
The energy change that occurs when a neutral atom acquires an electron.
In most cases, energy is released.
In general, the values become more negative.
In general, the values become less negative.
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5. Addition of more electrons gives the ion the configuration of an inert gas.
6.20 Periodic Trends: Ionic Radii
Review
Questions
1. How are ionic radii measured?
2. Explain why the radius of the rubidium ion is smaller than the radius of the
rubidium atom.
3. Explain why the radius of the tellurium ion is larger than the radius of the
tellurium atom.
4. Why is the oxygen anion larger than the fluoride anion?
5. Why is the sodium cation larger than the magnesium cation?
Answers
1.
2.
3.
4.
5.
Using solid crystals.
Rubidium loses an electron to form the ion.
Tellurium gains two electrons to form the ion.
Oxygen gains two electrons while fluorine only gains one.
Sodium only loses one electron, but magnesium loses two and the remaining
electrons are pulled closer to the nucleus.
6.21 Periodic Trends: Electronegativity
Review
Questions
1.
2.
3.
4.
5.
Define “electronegativity.”
How does electronegativity differ from electron affinity?
Why are the electronegativity values of metals generally low?
Describe the trend in electronegativities across the periodic table.
Describe the trends in electronegativities in a group of the periodic table.
Answers
1. The ability of an atom to attract electrons to itself.
2. Electron affinity involves actual energy release, but electronegativity is a relative
scale.
3. They increase stability by losing electrons.
4. Increase from left to right across a period.
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5. Decrease from top to bottom of a group.
6.21 Periodic Trends: Metallic and Nonmetallic Character
Review
Questions
1.
2.
3.
4.
5.
Define “metallic character.”
Define “non-metallic character.”
Describe the trend in metallic character going down a group.
Describe the trend in non-metallic character going across the periodic table.
Why does the metallic character increase as you go down a group?
Answers
1.
2.
3.
4.
5.
Level of reactivity of a metal.
Tendency to accept electrons during chemical reactions.
Increases.
Increases from left to right.
Easier to lose electrons as you go down a group.
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