Ch. 6 The Periodic Table and Periodic Law

Presentation

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Chemistry

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9th - 12th Grade

•

Practice Problem

•

Easy

•

NGSS

HS-PS1-1, HS-PS1-2

Standards-aligned

Mindy Creel

Used 12+ times

FREE Resource

60 Slides • 12 Questions

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.

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