Periodic Table and Periodic Trends: Key Concepts and Study Guide

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Periodic Table and Periodic Trends

Introduction

The periodic table is a fundamental tool in chemistry, organizing elements based on their properties and atomic structure. Understanding its layout and the trends within it is essential for predicting chemical behavior and properties.

5.1 The Original Periodic Law (Mendeleev)

Periodic Law (Mendeleev): Dmitri Mendeleev proposed that the properties of elements are a periodic function of their atomic masses.
Significance: This arrangement allowed for the prediction of undiscovered elements and their properties.
Example: Mendeleev left gaps for elements like germanium and predicted their properties before discovery.

5.2 Modern Periodic Law (Moseley)

Modern Periodic Law: Henry Moseley refined the periodic law, stating that properties of elements are a periodic function of their atomic numbers.
Atomic Number: The number of protons in the nucleus of an atom.
Example: The periodic table is now arranged by increasing atomic number, not atomic mass.

5.3 Organization of Elements in the Periodic Table

Periods: Horizontal rows in the periodic table. Elements in the same period have the same number of electron shells.
Groups/Families: Vertical columns. Elements in the same group have similar chemical properties due to similar valence electron configurations.
Example: Alkali metals (Group 1) are highly reactive and have one valence electron.

5.4 Designating Groups: Old and IUPAC Conventions

Old Convention: Groups labeled as IA-VIIIA and IB-VIIIB.
IUPAC Convention: Groups numbered 1-18 from left to right.
Example: Group 17 (IUPAC) is the same as Group VIIA (old convention), known as the halogens.

5.5 Periodic Trends: Atomic Size

Atomic Radius: The distance from the nucleus to the outermost electron shell.
Trend Across a Period: Atomic size decreases from left to right due to increasing nuclear charge.
Trend Down a Group: Atomic size increases due to addition of electron shells.
Example: Sodium (Na) is larger than chlorine (Cl) in the same period.

5.6 Predicting Chemical Formulas and Comparing Elements

Chemical Formula Prediction: Use valence electrons and group properties to predict formulas for compounds.
Example: Elements in Group 1 (alkali metals) form compounds with Group 17 (halogens) in a 1:1 ratio, e.g., NaCl.

5.7 Electron Sublevels and Energy Order

Electron Sublevels: s, p, d, f sublevels fill in a specific order based on energy.
Order of Filling: Sublevels fill according to increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
Example: The electron configuration of calcium (Ca):

5.8 Electron Configuration Prediction

Electron Configuration: The arrangement of electrons in an atom's orbitals.
Method: Use the periodic table and the order of sublevel filling to write configurations.
Example: Oxygen (O):

5.9 Valence Electrons

Valence Electrons: Electrons in the outermost shell, responsible for chemical bonding.
Prediction: The group number (for main group elements) indicates the number of valence electrons.
Example: Carbon (Group 14) has 4 valence electrons.

5.10 Ion Formation and Ionic Charge Prediction

Ionic Charge: Determined by the loss or gain of electrons to achieve a stable electron configuration.
Metals: Lose electrons to form positive ions (cations).
Nonmetals: Gain electrons to form negative ions (anions).
Example: Sodium (Na) loses one electron to form ; chlorine (Cl) gains one electron to form .

5.11 Ionization Energy

Ionization Energy: The energy required to remove an electron from an atom in the gaseous state.
Trend Across a Period: Increases from left to right due to increasing nuclear charge.
Trend Down a Group: Decreases due to increased distance from the nucleus.
Example: Helium has the highest ionization energy; cesium has one of the lowest.

5.12 Electron Pair Ionization Energy

Electron Pair: A pair of electrons in the same orbital.
Prediction: The pair with the highest ionization energy is typically found in elements with high nuclear charge and small atomic radius.
Example: Fluorine has a high ionization energy due to its small size and high effective nuclear charge.

5.13 Ionic Charge Prediction for Representative Elements

Representative Elements: Elements in groups 1, 2, and 13-18.
Prediction: Group 1 forms ions, Group 2 forms ions, Group 17 forms ions, etc.
Example: Magnesium (Group 2) forms ; oxygen (Group 16) forms .

5.14 Electron Configuration Diagrams

Diagramming: Use the periodic table and sublevel order to diagram electron configurations for selected ions and elements.
Example: : (same as neon)

Summary Table: Periodic Trends

Trend	Across a Period (Left to Right)	Down a Group
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electronegativity	Increases	Decreases
Metallic Character	Decreases	Increases

Additional info: Some explanations and examples have been expanded for clarity and completeness.