Chemistry Chapter on Noble Gases & Electron Config

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Questions and Answers

What is the common valence electron configuration for alkali metals?

ns1 (correct)

ns2

ns2 np5

ns2 np6

Which group of elements has incompletely filled d subshells?

Alkaline earth metals

Transition metals (correct)

Noble gases

Lanthanides

What distinguishes Group 2B elements from transition metals?

They have completely filled p subshells.

They form unipositive cations.

They possess a valence electron configuration of ns1.

They are not transition metals or representative elements. (correct)

Which electron configuration represents noble gases?

ns2 np6 Signup and view all the answers

What type of chemical reactivity do noble gases generally exhibit?

Low reactivity due to filled subshells Signup and view all the answers

How many valence electrons do alkaline earth metals have?

2 Signup and view all the answers

What does the configuration ns2 np5 represent?

Halogens Signup and view all the answers

Which element type is characterized by incompletely filled f subshells?

Actinides and Lanthanides Signup and view all the answers

What is the common outer electron configuration for elements in Group 14?

ns2np2 Signup and view all the answers

Why are noble gases considered chemically inert?

They have fully filled ns and np subshells. Signup and view all the answers

Which of the following statements is true about metals?

Metals can be malleable and ductile. Signup and view all the answers

Which element is not classified as a nonmetal?

Silicon Signup and view all the answers

What category do elements that can conduct electricity under certain conditions belong to?

Metalloids Signup and view all the answers

Which of the following is true regarding nonmetals at room temperature?

Many nonmetals are gases. Signup and view all the answers

How are elements organized in the periodic table according to IUPAC?

Columns are numbered sequentially from 1 to 18. Signup and view all the answers

Which statement best describes metalloids?

They have intermediate properties between metals and nonmetals. Signup and view all the answers

What does the effective nuclear charge (Z eff) represent?

The nuclear charge felt by an electron accounting for shielding Signup and view all the answers

How does the atomic radius change across a period?

It decreases due to increasing effective nuclear charge Signup and view all the answers

What is the primary reason for the increase in atomic radius down a group?

Addition of electrons in larger orbitals Signup and view all the answers

Which statement about effective nuclear charge is true as one moves down a group?

Z eff remains almost constant despite increased shielding Signup and view all the answers

Which of the following best describes the role of core electrons?

They shield valence electrons from the nucleus more effectively than other valence electrons Signup and view all the answers

Why does effective nuclear charge increase from left to right within a period?

The valence electrons are added without increased shielding Signup and view all the answers

Which factor primarily determines how much an electron is attracted to the nucleus?

The amount of shielding from other electrons Signup and view all the answers

If an element has a larger atomic radius, what can be inferred about its position in the periodic table?

It is located at the bottom of a group Signup and view all the answers

Which ion is larger, S2- or Cl-?

S2- Signup and view all the answers

Between Mg2+ and Ca2+, which ion has a larger radius?

Ca2+ Signup and view all the answers

In comparing Fe2+ and Fe3+, which one is larger?

Fe2+ Signup and view all the answers

What defines ionization energy?

Energy required to remove an electron from a gaseous atom Signup and view all the answers

How does the nuclear charge affect the ionic radius of isoelectronic ions?

Greater nuclear charge results in a smaller radius Signup and view all the answers

What is the first ionization energy represented by?

A(g) + 1é → A+(g) Signup and view all the answers

What happens to ionization energy as atomic number increases in the periodic table?

Ionization energy generally increases Signup and view all the answers

Which statement is true regarding the ionization energy trends in the periodic table?

It increases from top to bottom in a group Signup and view all the answers

Which statement accurately describes the trend of ionization energy across a period?

Ionization energy increases due to increasing effective nuclear charge. Signup and view all the answers

What is the relationship between successive ionization energies?

They increase as more electrons are removed. Signup and view all the answers

How does ionization energy change as you move down a group in the periodic table?

It decreases due to increased distance between nucleus and valence electrons. Signup and view all the answers

What effect does a larger effective nuclear charge have on ionization energy?

It increases the attraction between the nucleus and valence electrons. Signup and view all the answers

Which group of elements has the lowest first ionization energies?

Alkali metals Signup and view all the answers

What factors contribute to the decrease in ionization energy within a group?

Increased distance between the nucleus and valence electrons. Signup and view all the answers

Why is it more difficult to remove an electron from a positively charged ion compared to a neutral atom?

The nuclear charge remains constant while electrons are removed. Signup and view all the answers

What describes the trend of ionization energy based on atomic number in a period?

Ionization energy increases steadily. Signup and view all the answers

Study Notes

Noble Gases and Electron Configuration

Noble gases have completely filled p subshells, with helium having a fully filled 1s subshell.
Transition metals are located in Groups 1B and 3B through 8B and have incompletely filled d subshells or can form cations with such characteristics.
Group 2B includes zinc (Zn), cadmium (Cd), and mercury (Hg), which are not classified as representative elements or transition metals.
Lanthanides and actinides are characterized by incompletely filled f subshells.

Electron Configuration and Chemical Reactivity

Valence electrons are the outermost electrons that determine chemical reactivity, while core electrons are nonvalence electrons.
Elements in the same group (column) of the periodic table share the same number and type of valence electrons, suggesting similar chemical behaviors.
Alkali metals (Group 1) have a valence electron configuration of ns1 and typically lose one electron to form +1 cations.
Alkaline earth metals (Group 2) possess a configuration of ns2 and tend to lose two electrons, forming +2 cations.
Transition metals have a valence configuration of (n-1)dx ny, where the values x and y vary.
Halogens (Group 17) have the configuration ns2 np5, while noble gases (Group 18) have ns2 np6, except helium (1s2), signifying stability and low reactivity.

Physical Properties and Classification

Periodic Table elements can be categorized as metals, nonmetals, or metalloids.
Metals are typically good heat and electricity conductors, malleable, and ductile, with most being solid except mercury, which is liquid.
Nonmetals are usually poor conductors, with gases, liquids (like bromine), and solids (like carbon) represented among them.
Metalloids possess intermediate properties, behaving like nonmetals but capable of conducting electricity under specific conditions.

Periodicity in Properties

Effective nuclear charge (Z_eff) represents the nuclear charge experienced by an electron, factoring both the actual charge (Z) and the shielding effects from other electrons.
Within a period, Z_eff increases, leading to a greater pull on valence electrons.
In groups, the effective nuclear charge remains constant, but the increase in principal quantum number decreases the attraction of the nucleus to valence electrons.

Atomic and Ionic Radius

The atomic radius decreases from left to right across a period due to increasing Z_eff.
Conversely, the atomic radius increases down a group because of an increase in the principal quantum number (n).
Ionic radii increase with added negative charges and decrease with increased positive charges; ionic radii comparisons can be made using isoelectronic series and charge considerations.

Ionization Energy

Ionization energy is the minimum energy required to remove an electron from a gaseous atom; energy levels are defined for successive ionization processes.
Ionization energy generally increases across a period due to greater Z_eff and decreases down a group due to increased distance between the nucleus and valence electrons.
Alkali metals exhibit the lowest first ionization energies among groups due to their single valence electron being effectively shielded.

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Description

This quiz explores the properties of noble gases and their electron configurations, including the distinction between valence and core electrons. It also covers the reactivity of different groups in the periodic table and the characteristics of transition metals and lanthanides. Test your knowledge on these essential chemical concepts!