Ch. 6 The Periodic Table and Periodic Law

Multiple Choice

Who received credit for the first periodic table?

• Henry Moseley
• Dmitri Mendeleev
• Niels Bohr
• John Dalton

Multiple Select

What are the 3 broad categories of elements?

• Metals

• Gases

• Nonmetals

• Metalloids

Multiple Choice

Describe the properties of metals.

• Metals are conductive, malleable, ductile, lustrous, and usually solid at room temperature.
• Metals cannot be shaped or stretched.
• Metals are always gaseous at room temperature.
• Metals are insulators and brittle.

Multiple Choice

Describe properties of nonmetals.

• Nonmetals are typically shiny and malleable.
• Nonmetals are excellent conductors of heat and electricity.
• Nonmetals are sometimes brittle, poor conductors of heat and electricity, have high ionization energies, and can exist as gases, liquids, or solids.

• Nonmetals always exist as solids at room temperature.

Multiple Choice

Describe properties of metalloids.

• Metalloids are excellent conductors of electricity.
• Metalloids have no luster and are always gases at room temperature.
• Metalloids could be semiconductors, brittle, or have metallic luster, and exhibit properties of both metals and nonmetals.

• Metalloids are always ductile and malleable.

Multiple Choice

How can an element's location on the periodic table help us to determine the electron configuration?

• The periodic table shows the atomic mass, which determines electron configuration.
• Electron configuration is solely based on the element's group number.
• The location on the periodic table has no relation to the electron configuration.
• An element's location on the periodic table indicates its electron configuration by showing the order of electron filling in shells and subshells.

Multiple Choice

Which blocks contain valence electrons?

• s and d blocks
• s and p blocks
• f and g blocks
• d and f blocks

Multiple Choice

Why do atomic radii decrease across a period?

• Atomic radii increase across a period due to decreased nuclear charge.
• Atomic radii remain constant across a period because of electron shielding.
• Atomic radii decrease across a period due to increased nuclear charge attracting electrons more strongly.
• Atomic radii decrease across a period due to increased electron repulsion.

Multiple Choice

Why do atomic radii increase down a group?

• Atomic radii increase across a period due to higher electronegativity.
• Atomic radii increase down a group due to the addition of electron shells and increased shielding effect.
• Atomic radii decrease down a group due to increased nuclear charge.
• Atomic radii remain constant down a group because of stable electron configurations.

Multiple Choice

Why do ionic radii suddenly increase then continue to decrease moving across periods?

• Ionic radii increase then decrease across periods due to the formation of anions and cations.
• Ionic radii remain constant across periods due to atomic size.
• Ionic radii decrease then increase due to electron loss.
• Ionic radii only decrease as protons are added.

Multiple Choice

When does the largest jump in ionization energy occur?

• When moving to a lower electron shell

• When moving from a lower to a higher group within the same period.
• When adding more protons to the nucleus without changing the electron shell.
• When electrons are removed from the outermost shell without changing the shell.

Multiple Choice

How are ionization energy and electronegativity related?

• Electronegativity is solely determined by an element's atomic mass.
• Ionization energy is unrelated to electronegativity and only depends on atomic size.
• Ionization energy and electronegativity are directly related with higher electronegativity requiring more ionization energy to remove electrons.

• Higher ionization energy means lower electronegativity in all cases.