Science Exam Reviewer: Electronic Structure & Chemical Bonding PDF

**SCIENCE EXAM REVIEWER** **[Electronic Structure of Matter]** ### **I. Overview of the Electronic Structure of Matter** - **Electronic Structure** refers to how electrons are arranged in atoms, ions, or molecules. - It determines chemical properties, reactivity, bonding behavior, and...

**SCIENCE EXAM REVIEWER** **[Electronic Structure of Matter]** ### **I. Overview of the Electronic Structure of Matter** - **Electronic Structure** refers to how electrons are arranged in atoms, ions, or molecules. - It determines chemical properties, reactivity, bonding behavior, and more. ### **II. Basic Atomic Structure** - **Subatomic Particles**: - **Protons**: Positive charge (+1), found in the nucleus. - **Neutrons**: No charge, found in the nucleus. - **Electrons**: Negative charge (-1), orbit the nucleus in energy levels. ### **III. Energy Levels and Sublevels** 1. **Energy Levels (Shells)** - Denoted by **principal quantum number (n)**. - Example: n = 1 (first shell), n = 2 (second shell), etc. 2. **Sublevels within Energy Levels**\ Each energy level contains one or more sublevels (s, p, d, f). - **s-orbital**: 1 orbital, holds 2 electrons. - **p-orbital**: 3 orbitals, holds 6 electrons. - **d-orbital**: 5 orbitals, holds 10 electrons. - **f-orbital**: 7 orbitals, holds 14 electrons. ### **IV. Quantum Numbers and Their Meaning** 1. **Principal Quantum Number (n)**: - Describes the main energy level (1, 2, 3...). 2. **Azimuthal Quantum Number (l)**: - Describes the shape of the orbital (0 = s, 1 = p, 2 = d, 3 = f). 3. **Magnetic Quantum Number (mₗ)**: - Describes the orientation of an orbital (--l to +l). 4. **Spin Quantum Number (ms)**: - Describes the spin of the electron (+½ or -½). ### **V. Electron Configuration and Rules for Filling Orbitals** 1. **Electron Configuration** - Describes the arrangement of electrons in an atom\'s orbitals.\ Example: **Carbon (C)**: [1*s*^2 ^2*s*^2^ 2*p*^2^]{.math.inline} 2. **Principles Governing Electron Configuration** - **Aufbau Principle**: Fill lower-energy orbitals first (e.g., 1s before 2s). - **Pauli Exclusion Principle**: No two electrons can have the same set of quantum numbers. - **Hund's Rule**: Electrons fill orbitals of equal energy one by one before pairing up. **The Noble Gas Configuration** is the shorthand way of writing the electron configuration. **Energy Level Order:** 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s ### **VI. Periodic Trends Related to Electronic Structure** 1. **Atomic Radius** - Increases down a group and decreases across a period. 2. **Ionization Energy** - Energy required to remove an electron. - Increases across a period and decreases down a group. 3. **Electronegativity** - Tendency of an atom to attract electrons. - Increases across a period and decreases down a group. 4. **Electron Affinity** - Energy change when an atom gains an electron. **[Chemical Bonding ]** ### **I. Overview of Chemical Bonding** Chemical bonding refers to the interaction that holds atoms or ions together to form compounds and molecules. The goal of bonding is to achieve a more stable configuration, often fulfilling the **octet rule** (8 valence electrons). ### **II. Types of Chemical Bonds** 1. **Ionic Bond** - **Definition**: A bond formed through the **transfer of electrons** from a metal to a non-metal. - **Examples**: NaCl (Sodium Chloride), MgO (Magnesium Oxide). - **Properties**: - High melting and boiling points. - Soluble in water. - Conduct electricity when dissolved or molten. 2. **Covalent Bond** - **Definition**: A bond formed by **sharing electron pairs** between two non-metals. - **Examples**: H₂O (Water), CO₂ (Carbon Dioxide). - **Types of Covalent Bonds**: - **Non-polar covalent**: Equal sharing of electrons (e.g., O₂). - **Polar covalent**: Unequal sharing of electrons (e.g., H₂O). - **Properties**: - Low melting and boiling points. - Poor electrical conductors. - Can be polar or non-polar. 3. **Metallic Bond** - **Definition**: A bond where metal atoms share a \"sea of delocalized electrons\" that move freely between atoms. - **Examples**: Fe (Iron), Cu (Copper). - **Properties**: - Good conductors of heat and electricity. - Malleable and ductile. - High melting points. 4. **Hydrogen Bond** (Special type of attraction) - **Definition**: A weak attraction between a hydrogen atom (bonded to F, O, or N) and another electronegative atom. - **Examples**: H₂O (between water molecules), NH₃ (Ammonia). - **Importance**: Affects the properties of water and DNA stability. ### ### **III. Key Concepts in Chemical Bonding** 1. **Octet Rule** - Atoms tend to gain, lose, or share electrons to have **8 electrons in their outer shell** (except for hydrogen and helium). 2. **Electronegativity** - **Electronegativity Difference** determines bond type: - 0--0.4 → Non-polar covalent - 0.5--1.7 → Polar covalent - 1.7 → Ionic 3. **Lewis Structures** - Diagrams showing the arrangement of electrons in a molecule. - **Lone pairs**: Non-bonding electrons. - **Bond pairs**: Electrons shared between atoms. ### **Summary of Key Points** - **Quantum numbers** define the position and behavior of electrons. - **Electron configurations** explain how electrons are arranged in atoms. - **Periodic trends** like atomic radius and ionization energy are influenced by electronic structure. - **Bonding behavior** (ionic, covalent, metallic) is determined by the arrangement of valence electrons. - **Ionic bonds** involve electron transfer between metals and non-metals. - **Covalent bonds** involve electron sharing between non-metals, which can be polar or non-polar. - **Metallic bonds** feature delocalized electrons among metal atoms, allowing conductivity. - **Hydrogen bonding** is a weak interaction that plays a crucial role in biological systems. Goodluck. ;D