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2.2 Chemistry Notes: The Periodic Table and Atomic Structure
A Predictive Model
Electron Configuration and Properties:
An element's chemical and physical properties are determined by its electron configuration, which includes the arrangement of valence electrons around the nucleus.
The periodic table structure helps predict these properties, especially within chemical families like alkali metals, alkaline earth metals, halogens, noble gases, and transition metals, by recognizing patterns in valence electrons.
Periodic Table Blocks (s, p, d, f):
The periodic table is divided into four blocks (s, p, d, f) which help visualize and predict an element’s reactivity and chemical properties.
s-block: Includes Groups 1 and 2. Elements have 1 or 2 electrons in their outer s subshell.
p-block: Includes Groups 13–18. Elements have p subshell electrons.
d-block: Includes transition metals (Groups 3–12), with filled or partially filled d subshells.
f-block: Includes lanthanides and actinides, where f subshells are being filled.
Group-Specific Reactivity and Properties
Alkali Metals (Group 1):
Have one electron in the outer s subshell.
Highly reactive due to having only one valence electron.
Alkaline Earth Metals (Group 2):
Have two electrons in the outer s subshell.
Less reactive than alkali metals due to a filled s orbital.
Halogens (Group 17):
One electron short of a full octet.
Very reactive, especially in forming compounds.
Noble Gases (Group 18):
Have a full valence shell.
Relatively nonreactive due to stable electron configuration.
Transition Metals (Groups 3–12):
Contain filled or partially filled d subshells.
Generally less reactive than Groups 1 and 2 elements.
The Periodic Table as a Predictive Model (Continued)
Predicting Electron Configurations:
The periodic table can predict electron configurations based on an element’s position, reflecting systematic changes across periods.
Core and Valence Electrons:
Example: Sodium (Na)
Atomic number: 11
Configuration: [Ne] 3s¹ (10 core electrons and 1 valence electron)
Core electrons are those in the inner energy levels, while valence electrons are in the outermost energy level and participate in chemical reactions.
Developing Models:
Chlorine (Cl) Example:
Use chlorine’s position to model its electron configuration, including protons in the nucleus, core electrons, and valence electrons.
Coulomb’s Law
Interactions of Charged Particles:
Coulomb’s Law describes the force (F) between two charged particles, depending on the charge (q) of each particle and the distance (d) between them.
Formula: F=keq1q2d2F = k_e \frac{q_1 q_2}{d^2}F=ked2q1q2
The force is directly proportional to the product of the charges and inversely proportional to the square of the distance.
Implications of Coulomb’s Law:
Larger charges result in stronger attractions or repulsions.
Greater distances between charges weaken the force.
The Shielding Effect and Effective Nuclear Charge
Shielding Effect:
Core electrons shield valence electrons from the full attractive force of the nucleus.
The repulsive forces between electrons reduce the effective nuclear charge experienced by each electron.
Effective Nuclear Charge (Zeff):
Formula: Zeff=Z−SZ_{\text{eff}} = Z - SZeff=Z−S
ZZZ: Nuclear charge (number of protons).
SSS: Shielding constant (approximated by the number of core electrons).
Trends in Zeff:
Increases from left to right across a period because the nuclear charge increases while the number of core electrons remains constant.
Example of increasing Zeff:
Sodium (Na): Zeff=11−10=+1Z_{\text{eff}} = 11 - 10 = +1Zeff=11−10=+1
Magnesium (Mg): Zeff=12−10=+2Z_{\text{eff}} = 12 - 10 = +2Zeff=12−10=+2
Aluminum (Al): Zeff=13−10=+3Z_{\text{eff}} = 13 - 10 = +3Zeff=13−10=+3
Periodic Trends:
Zeff increases across a period and decreases down a group.
Comparative Models:
Comparing elements like silicon and germanium, which are in the same group, can show how Zeff changes due to differences in shielding and nuclear charge.
2.2 Chemistry Notes: The Periodic Table and Atomic Structure
A Predictive Model
Electron Configuration and Properties:
An element's chemical and physical properties are determined by its electron configuration, which includes the arrangement of valence electrons around the nucleus.
The periodic table structure helps predict these properties, especially within chemical families like alkali metals, alkaline earth metals, halogens, noble gases, and transition metals, by recognizing patterns in valence electrons.
Periodic Table Blocks (s, p, d, f):
The periodic table is divided into four blocks (s, p, d, f) which help visualize and predict an element’s reactivity and chemical properties.
s-block: Includes Groups 1 and 2. Elements have 1 or 2 electrons in their outer s subshell.
p-block: Includes Groups 13–18. Elements have p subshell electrons.
d-block: Includes transition metals (Groups 3–12), with filled or partially filled d subshells.
f-block: Includes lanthanides and actinides, where f subshells are being filled.
Group-Specific Reactivity and Properties
Alkali Metals (Group 1):
Have one electron in the outer s subshell.
Highly reactive due to having only one valence electron.
Alkaline Earth Metals (Group 2):
Have two electrons in the outer s subshell.
Less reactive than alkali metals due to a filled s orbital.
Halogens (Group 17):
One electron short of a full octet.
Very reactive, especially in forming compounds.
Noble Gases (Group 18):
Have a full valence shell.
Relatively nonreactive due to stable electron configuration.
Transition Metals (Groups 3–12):
Contain filled or partially filled d subshells.
Generally less reactive than Groups 1 and 2 elements.
The Periodic Table as a Predictive Model (Continued)
Predicting Electron Configurations:
The periodic table can predict electron configurations based on an element’s position, reflecting systematic changes across periods.
Core and Valence Electrons:
Example: Sodium (Na)
Atomic number: 11
Configuration: [Ne] 3s¹ (10 core electrons and 1 valence electron)
Core electrons are those in the inner energy levels, while valence electrons are in the outermost energy level and participate in chemical reactions.
Developing Models:
Chlorine (Cl) Example:
Use chlorine’s position to model its electron configuration, including protons in the nucleus, core electrons, and valence electrons.
Coulomb’s Law
Interactions of Charged Particles:
Coulomb’s Law describes the force (F) between two charged particles, depending on the charge (q) of each particle and the distance (d) between them.
Formula: F=keq1q2d2F = k_e \frac{q_1 q_2}{d^2}F=ked2q1q2
The force is directly proportional to the product of the charges and inversely proportional to the square of the distance.
Implications of Coulomb’s Law:
Larger charges result in stronger attractions or repulsions.
Greater distances between charges weaken the force.
The Shielding Effect and Effective Nuclear Charge
Shielding Effect:
Core electrons shield valence electrons from the full attractive force of the nucleus.
The repulsive forces between electrons reduce the effective nuclear charge experienced by each electron.
Effective Nuclear Charge (Zeff):
Formula: Zeff=Z−SZ_{\text{eff}} = Z - SZeff=Z−S
ZZZ: Nuclear charge (number of protons).
SSS: Shielding constant (approximated by the number of core electrons).
Trends in Zeff:
Increases from left to right across a period because the nuclear charge increases while the number of core electrons remains constant.
Example of increasing Zeff:
Sodium (Na): Zeff=11−10=+1Z_{\text{eff}} = 11 - 10 = +1Zeff=11−10=+1
Magnesium (Mg): Zeff=12−10=+2Z_{\text{eff}} = 12 - 10 = +2Zeff=12−10=+2
Aluminum (Al): Zeff=13−10=+3Z_{\text{eff}} = 13 - 10 = +3Zeff=13−10=+3
Periodic Trends:
Zeff increases across a period and decreases down a group.
Comparative Models:
Comparing elements like silicon and germanium, which are in the same group, can show how Zeff changes due to differences in shielding and nuclear charge.
