Untitled

• Molecular polarity refers to how charge gets distributed in a molecule.

Polarity of Molecules

• All compounds can be classified by the type of chemical bond formed when the atoms combined.
• Classification determines if a molecule is polar or nonpolar.

Covalent and Ionic Compounds

• Covalent compounds are molecules formed by covalent bonds where atoms share one or more valence electrons.
• Ionic compounds consist of ions, not molecules.
• Ionic compound atoms transfer electrons from one atom to another, forming an ionic bond that relies on electrostatic attraction.

Two Types of Covalent Bonds

• Polar Covalent Bond:

  • Exists when two bonded atoms aren't equally distributed.
  • One atom is more electronegative than the other.
  • One atom is partially positive, the other partially negative.

• Non-Polar Covalent Bond:

  • Created when two atoms have similar electronegativity and will share electrons equally.

• Two factors determine the polarity of molecules:

  • Polarity of bonds between atoms studied based on electronegativity.
  • The geometrical shape of the molecule predicted via the Valence Shell Electron Pair Repulsion (VSEPR) theory.

Electronegativity Defined

• Electronegativity is a measure of an atom's tendency to attract electrons when chemically combined with another atom.
• The higher the electronegativity, the more it attracts electrons.

Polar Covalent Bonds

• Polar covalent bonds occur when electron pairs are unequally shared, with a significant difference in electronegativity between atoms.

• Examples of compounds exhibiting polar covalent bonds include:

• HCl: EN of H = 2.1, EN of Cl = 3.0, ΔEN = 0.9

• HF: EN of H = 2.1, EN of F = 4.0, ΔEN = 1.9

• Separation of charges makes the bond polar.

• An electric dipole is created, referring to "two poles," with a positive and negative pole within the molecule.

• Elements with higher EN values become the partial negative pole.

• Elements with lower EN values become the partial positive pole, this turns the molecule into a polar molecule.

• Electronegativity difference to determine bond polarity:

  • Greater than or equal to 1.7 = Ionic
  • 0.5 - 1.6 = Polar
  • Less than 0.5 (Covalent) = Non-Polar

Bond Types and Polarity

• Ionic Bond:
  • Electrons are transferred.
  • Creates oppositely charged ions.
  • Occurs between metals and nonmetals.
  • Examples: Sodium Chloride (NaCl), Magnesium Oxide (MgO).
• Covalent Bond:
  • Electrons are shared.
  • Strong bonds form between nonmetals.
  • Example: Water (H2O), Carbon Dioxide (CO2)
• Nonpolar Covalent Bond:
  • Electrons are shared equally.
  • No partial charges on atoms.
  • Occurs between atoms with similar electronegativity.
  • Examples: Methane (CH4), Oxygen (O2).

Molecular Geometry and VSEPR Theory

• Molecular geometry and the valence shell electron pair repulsion (VSEPR) theory help determine the spatial arrangement of atoms in a polyatomic molecule.

Shapes Under VSEPR Theory

• Linear
• Trigonal Planar
• Bent or Angular
• Tetrahedral
• Trigonal Pyramidal
• Trigonal Bipyramidal
• Octahedral

Shapes, Descriptions, Bond Angles, and Examples

• Linear shape:
  • Atoms are arranged in a straight line
  • 180° bond angle
  • Example: Carbon Dioxide (CO2)
• Trigonal Planar shape:
  • Three atoms are around a central atom in a plane.
  • 120° bond angle.
  • Example: Boron Trifluoride (BF3)
• Tetrahedral shape:
• Four atoms around a central atom create a tetrahedron
• 109.5° bond angle
• Creates Methane (CH4)
• Trigonal Bipyramidal shape:
• Five atoms around a central atom, three in a plane, two above and below.
• 90°, 120° bond angles.
• Example: Phosphorus Pentachloride (PCl5)
• Octahedral shape:
  • Six atoms around a central atom forming an octahedron.
  • 90° bond angle.
• Example: Sulfur Hexafluoride (SF6)
• Bent/Angular shape:
  • Similar to trigonal planar but with lone pairs
  • < 120° bond angle.
  • Example: Water (H2O).

Lewis Dot Structure

• Lewis dot structures (aka electron dot structures) are diagrams describing chemical bonding between atoms in a molecule.
• They display the number of lone pairs in each of the atoms that make up the molecule.

How to Draw a Lewis Structure

• Find total valence electrons.
• Draw the skeleton structure sing single bonds.
• Assign remaining valence electrons.
• Make multiple bonds if all octets aren't filled.
• Add bracket and charge last.

Valence Electrons

• Valence electrons are defined as the electron shell can hold a number of elections.

Polarity, Solubility, and Miscibility

• Polarity = a separation of electric charge leading to a molecule or its chemical groups. It has an electric dipole moment, with a negatively charged end with a positively charged end.
• Polar molecules = one or more polar bonds due to electronegativity differences.
• Solubility = a solid substance's ability to be dissolved in a solvent.
• Miscibility = ability of two liquids to combine/mix in all proportions, to create a homogenous mixture.

Solubility and Miscibility Rule

• "Like dissolves like" or "like mixes with like" means:
  • Polar substances dissolve/mix with polar substances.
  • Nonpolar substances are soluble/miscible with nonpolar substances.

Intermolecular Forces

• Intermolecular forces of attraction = forces that hold molecules together in a substance.
  • This force holds the molecules, and can either be attractive or repulsive.

London Dispersion Forces (LDF)

• Weakest force, found in all molecules.
• Caused by temporary shifts in electron movement.
• Example: Oxygen gas (O2), noble gases such as helium (He).

Dipole-Dipole Forces

• Happens occur with polar molecules (molecules with positive and negative sides).
• Stronger than LDF, but weaker than hydrogen bonding.
• Example: Water (H2O) and acetone.

Hydrogen Bonding

• Special strong dipole-dipole force, when hydrogen bonds with Fluorine, Oxygen, or Nitrogen.
• Makes water molecules stick together. This gives water its' unique properties.
• Example: Water (H2O), DNA structure, and alcohol.

Ion-Dipole Forces

• Strongest intermolecular force - occurs with ions and polar molecules.
• Helps dissolve ionic compounds in water.
• Example: Salt (NaCl) dissolving in water.

Bond Strength and Physical Properties

• Intermolecular forces of attraction (IMFA) exist between molecules.
• There are different types of IMFA, which are arranged from strongest to weakest: Ion-dipole, Hydrogen-bonding, Dipole-dipole, Dipole-induced dipole, London forces of attraction.
• The strength of IMFA affects the the physical properties of substances.

General Properties of Polar vs Nonpolar Molecules

• Polar molecules:
• H-bonding and dipole-dipole IMFA
• Exist as solids/liquids at room temperature
• High boiling point
• High melting point
• High surface tension
• Low vapor pressure
• Low volatility
• Soluble in water
• Nonpolar molecules:
  • London dispersion IMFA
  • exist as gases at room temperature
  • Low boiling point
  • Low melting point
  • Low surface tension
  • High vapor pressure
  • High volatility
  • Insoluble in water

Physical Property Definitions;

• Boiling Point = temperature at which the vapor pressure and atmospheric pressure of a liquid substance are equal.

• Melting point = temperature at which solid becomes liquid.

  • Solid and liquid phases exist in equilibrium.

• Surface tension = energy needed to increase the surface area by a unit amount.

• Viscosity = the resistance of a liquid to flow.

• Vapor Pressure = pressure exerted by a substance in its gaseous state.

• Volatility = measures the rate at which a substance vaporizes (changes from liquid to gas).

• Boiling point, melting point, viscosity, and surface tension increase as the strength of intermolecular forces increases.

• Vapor pressure and volatility decrease with increasing strength of IMFA.

• London dispersion forces increase as the molecular mass of a substance increases.

• In H-bonding, the molar mass increases but the boiling point, melting point, viscosity and surface tension decrease.