Chem chapter 3 study guide

3.2 The Periodic Law and the Periodic Table

Introduction to the Periodic Table

The periodic table is a fundamental tool in chemistry, organizing elements based on atomic number and recurring chemical properties. Its structure reflects periodic trends and relationships among elements.

• Mendeleev's Organization: Dmitri Mendeleev developed the modern form of the periodic table, grouping elements with similar characteristics.

• Periodic Law: The periodic law states that the properties of elements recur periodically when arranged by increasing atomic number.

• Periodic Table Summary: Arranging elements by atomic number reveals strong correlations with elemental properties.

3.3 Electron Configurations: How Electrons Occupy Orbitals

Quantum Numbers and Electron Arrangement

Electrons occupy atomic orbitals according to specific rules and quantum numbers, which determine their arrangement and chemical behavior.

• Spin Quantum Number (ms): Can have values of +½ or –½, representing the two possible spin states of an electron.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers. Each orbital can hold a maximum of two electrons with opposite spins.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing, maximizing total spin.

• Subshells and Electron Filling: Subshells within a principal energy level are filled in order of increasing energy, following the concept of electron shielding.

• Writing Electron Configurations: Use the periodic table to determine the order in which orbitals are filled and write configurations for any element.

Example:

• Carbon (Z = 6): $1s^2 2s^2 2p^2$

3.4 Electron Configurations, Valence Electrons, and the Periodic Table

Core and Valence Electrons

Understanding the distinction between core and valence electrons is essential for predicting chemical reactivity and bonding.

• Core Electrons: Electrons in inner shells, not involved in bonding.

• Valence Electrons: Electrons in the outermost shell, responsible for chemical properties.

• Blocks of the Periodic Table: The table is divided into s, p, d, and f blocks based on the type of orbital being filled.

• Practice: Refer to problems 49-54 for additional exercises.

3.5 Electron Configurations and Elemental Properties

Valence Electrons and Chemical Properties

The number and arrangement of valence electrons determine the chemical properties and reactivity of elements.

• Periodic Table Analysis: The periodic table helps distinguish between metals, nonmetals, and semimetals.

• Main-Group vs. Transition Elements: Main-group elements have predictable valence electron configurations, while transition elements may have variable configurations.

• Bond Formation: Elements with certain valence electron arrangements are energetically favorable for forming stable bonds.

3.6 Periodic Trends in Atomic Size and Effective Nuclear Charge

Atomic Radius and Shielding

Periodic trends in atomic size and effective nuclear charge ($Z_{eff}$) are crucial for understanding element properties.

• Atomic Radius: The distance from the nucleus to the outermost electron shell. Increases down a group and decreases across a period.

• Effective Nuclear Charge ($Z_{eff}$): The net positive charge experienced by valence electrons, accounting for shielding by inner electrons.

• Shielding: Inner electrons reduce the attraction between the nucleus and outer electrons.

Formula:

• $Z_{eff} = Z - S$ Where $Z$ is the atomic number and $S$ is the number of shielding electrons.

3.7 Ions: Electron Configurations, Magnetic Properties, Radii, and Ionization Energy

Ion Formation and Properties

Ions are formed by the gain or loss of electrons, affecting their electron configurations, magnetic properties, and size.

• Electron Configurations for Ions: For anions, add electrons; for cations, remove electrons from the highest principal energy level.

• Paramagnetic vs. Diamagnetic: Paramagnetic ions have unpaired electrons and are attracted to magnetic fields; diamagnetic ions have all electrons paired and are repelled.

• Ion Size: Cations are smaller than their neutral atoms; anions are larger.

• Ionization Energy: The energy required to remove an electron from an atom. First ionization energy decreases down a group and increases across a period.

• Successive Ionization Energies: Each subsequent electron removed requires more energy, especially after removing all valence electrons.

• Practice: Refer to problems 81-94 for more exercises.

Formula:

• $\text{Ionization Energy} = E_{\text{final}} - E_{\text{initial}}$

3.8 Electron Affinities and Metallic Character

Electron Affinity and Metallic Trends

Electron affinity measures the energy change when an atom gains an electron, while metallic character describes how readily an element loses electrons.

• Electron Affinity: Generally becomes less exothermic down a group and more exothermic across a period.

• Metallic Character: Increases down a group and decreases across a period.

• Practice: Refer to problems 99-102 for additional exercises.

Example Table: Predicted Trends in Atomic Properties

Property	Down a Group	Across a Period
Atomic Radius	Increases	Decreases
Ionization Energy	Decreases	Increases
Electron Affinity	Less exothermic	More exothermic
Metallic Character	Increases	Decreases