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Chemistry

Chemistry 1120 - Test 3 Review Sheet

Chapter 10: Chemical Bonding and Molecular Structure

Octet Rule

  • The octet rule states that atoms tend to prefer to have eight electrons in their valence shell to achieve stability.

Electron Dot Diagrams (Lewis Dot Diagrams)

  • Steps Taken to Create Lewis Dot Diagrams:

    1. Calculate Total Electrons Available:

    • Include the total number of valence electrons from all atoms involved, taking into account the charge of any polyatomic ion.

    1. Determine the Central Atom:

    • Typically the least electronegative atom, which usually is placed at the center of the structure.

    1. Arrange Remaining Atoms:

    • Position atoms around the central atom as needed.

    1. Connect With Single Bonds:

    • Form single bonds between the central atom and surrounding atoms.

    1. Fill Remaining Electron Pairs:

    • Distribute remaining electrons to fill octets for surrounding atoms as much as possible.

    1. Move Pairs of Electrons:

    • If octets are not fulfilled for the central atom, shift pairs of electrons from outer atoms to form double or triple bonds.

    1. Resonance Structures:

    • If double bonds are formed, check for possible resonance structures.

    1. Include Brackets and Charges:

    • If applicable, denote overall molecular charge.

Resonance Structure

  • A resonance structure is one of two or more valid Lewis structures for a molecule that cannot be represented accurately by one structure alone. These structures illustrate the delocalization of electrons.

Electronegativity

  • Electronegativity is the measure of an atom's ability to attract and hold electrons when it is in a compound.

    • Periodic Trends in Electronegativity:

      • Increases left to right across a period.

      • Decreases as one moves down a group.

Polar Bond

  • A polar bond occurs when there is an uneven sharing of electrons between two atoms with different electronegativities, resulting in partial charges.

    • Use electronegativity values to identify polar bonds.

Molecular Geometry and Shapes

  • VSEPR Theory (Valence Shell Electron Pair Repulsion Theory):

    • Used to predict the shape of individual molecules based on the extent of electron-pair repulsion.

  • Basic Molecular Shapes:

    • Tetrahedral:

    • Four bonded atoms and no lone pairs around the central atom.

    • Trigonal Pyramidal:

    • Three bonded atoms and one lone pair on the central atom.

    • Bent:

    • Two bonded atoms and two lone pairs on the central atom.

    • Trigonal Planar:

    • Three bonded atoms and no lone pairs on the central atom.

    • Linear:

    • Two bonded atoms or a straight chain of atoms.

  • Effect of Lone Pairs on Angles in Tetrahedral Shape:

    • The angles become less than $109.5^{ ext{o}}$ as lone pairs are added since they occupy more space than bonding pairs.

  • Assessing Molecular Polarity:

    • A molecule is non-polar if it is symmetrical (the same all around).

    • If the central atom has lone pairs, the molecule tends to be polar.

Chapter 12: Intermolecular Forces

Definition of Intermolecular Forces

  • Intermolecular forces refer to the forces that act between molecules, affecting their physical properties.

Phase Changes

  • Evaporation/Vaporization:

    • The process by which liquid converts into gas due to sufficient thermal energy to overcome intermolecular forces holding the liquid molecules together.

  • Condensation:

    • The process where gas molecules lose enough energy to form a liquid through intermolecular forces that cause molecules to ``stick'' together.

  • Boiling Point/Condensation Point:

    • The temperature at which a liquid turns to gas (boiling point) and the reverse process, where gas turns to liquid (condensation point).

  • Melting Point/Freezing Point:

    • The temperature at which a solid transitions to liquid (melting point) and where a liquid transitions to solid (freezing point).

Types of Intermolecular Forces

  • Dispersion Forces (London Forces):

    • Very weak forces due to temporary alignments of electrons; strength increases with increasing molar mass; affects all molecules and is the only force impacting non-polar molecules.

  • Dipole-Dipole Interactions:

    • Occur between polar molecules due to interactions between their partial charges; stronger than dispersion forces since the charges are permanent.

  • Hydrogen Bonding:

    • A special case of dipole-dipole interaction, occurring when hydrogen is bonded to an electronegative atom (such as oxygen, nitrogen, or fluorine) resulting in strong intermolecular attractions.

Non-Molecular Solids

  • Ionic Solids:

    • Composed of a lattice arrangement of positive and negative ions. Proportional to the charge and size of ions, they exhibit significantly stronger interactions than molecular solids.

  • Network Solids (Network Covalent Solids):

    • Solids where every atom is covalently bonded to its neighbors, leading to exceptionally strong materials.

  • Metals (Metallic Bonding):

    • Characterized by a "sea" of delocalized electrons that are shared among all metal atoms, contributing to their conductivity and malleability.

Chapter 13: Solutions

Definition of a Solution

  • A solution is a homogeneous mixture comprising two or more substances.

Possible Combinations of Solutions

  • Solid in Liquid:

  • Solid in Solid: (only when melted first)

  • Liquid in Liquid:

  • Gas in Liquid:

  • Gas in Gas: (example - air you are currently breathing)

Components of a Solution

  • Solvent:

    • The substance present in the larger quantity that performs the dissolving action.

  • Solute:

    • The substance being dissolved, typically present in a smaller quantity compared to the solvent.