honors chemistry unit 4 study guides

Key Concepts

• Chemical bonds form when atoms share or transfer electrons to achieve a stable electronic configuration
• Valence electrons in the outermost shell participate in bonding
• Octet rule states atoms tend to gain, lose, or share electrons to attain a full valence shell with 8 electrons (except hydrogen which requires 2)
• Ionic bonds involve the complete transfer of electrons from one atom to another, forming ions with opposite charges that attract each other
• Covalent bonds involve the sharing of electrons between atoms, creating a stable bond through shared electron pairs
• Electronegativity difference between bonded atoms determines the polarity of the bond
• VSEPR theory predicts the geometry of molecules based on the number of electron domains around the central atom

Types of Chemical Bonds

• Ionic bonds form between metals and nonmetals through the complete transfer of electrons (sodium chloride, NaCl)
• Covalent bonds occur between nonmetals through the sharing of electrons
  • Single covalent bonds share one pair of electrons (hydrogen gas, H2)
  • Double covalent bonds share two pairs of electrons (carbon dioxide, CO2)
  • Triple covalent bonds share three pairs of electrons (nitrogen gas, N2)
• Coordinate covalent bonds involve one atom providing both electrons in the shared pair (ammonium ion, NH4+)
• Metallic bonds occur between metal atoms, characterized by a sea of delocalized electrons surrounding positively charged metal ions
• Polar covalent bonds have unequal sharing of electrons due to electronegativity differences between atoms (water, H2O)
• Nonpolar covalent bonds have equal sharing of electrons between atoms with similar electronegativities (methane, CH4)

Electronegativity and Polarity

• Electronegativity is the ability of an atom to attract electrons in a chemical bond
• Pauling scale assigns relative electronegativity values to elements, with fluorine being the most electronegative at 4.0
• Electronegativity difference between bonded atoms determines bond polarity
  • Nonpolar covalent bonds have an electronegativity difference of less than 0.4
  • Polar covalent bonds have an electronegativity difference between 0.4 and 1.7
  • Ionic bonds have an electronegativity difference greater than 1.7
• Dipole moment is a measure of the polarity of a molecule, represented by an arrow pointing from the positive to the negative end
• Polar molecules have an uneven distribution of charge and a net dipole moment (hydrogen fluoride, HF)
• Nonpolar molecules have a symmetrical distribution of charge and a net dipole moment of zero (carbon tetrachloride, CCl4)

Molecular Geometry

• VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry based on the number of electron domains (bonding and lone pairs) around the central atom
• Electron domains arrange themselves to minimize repulsion and maximize stability
• Two electron domains result in a linear geometry (carbon dioxide, CO2)
• Three electron domains result in a trigonal planar geometry (boron trifluoride, BF3)
• Four electron domains can result in tetrahedral (methane, CH4), trigonal pyramidal (ammonia, NH3), or bent (water, H2O) geometries, depending on the number of lone pairs
• Five electron domains result in a trigonal bipyramidal geometry (phosphorus pentachloride, PCl5)
• Six electron domains result in an octahedral geometry (sulfur hexafluoride, SF6)
• Molecular geometry influences polarity, as the arrangement of polar bonds can cancel out or reinforce each other

Intermolecular Forces

• Intermolecular forces are attractive forces between molecules, weaker than intramolecular bonds (covalent, ionic)
• Dipole-dipole forces occur between polar molecules, as the positive end of one molecule attracts the negative end of another
• London dispersion forces (induced dipole forces) occur between nonpolar molecules due to temporary fluctuations in electron distribution
  • Strength of London dispersion forces increases with molecular size and surface area
• Hydrogen bonding is a strong dipole-dipole interaction involving hydrogen atoms bonded to highly electronegative elements (fluorine, oxygen, nitrogen)
  • Hydrogen bonds are responsible for the unique properties of water, such as high boiling point and surface tension
• Ion-dipole forces occur between ions and polar molecules (sodium ion and water)
• Intermolecular forces influence physical properties such as boiling point, melting point, and solubility

Bond Energy and Strength

• Bond energy is the amount of energy required to break a chemical bond
• Bond strength is directly proportional to bond energy; stronger bonds have higher bond energies
• Covalent bond strength depends on the number of shared electron pairs (single, double, triple bonds)
  • Triple bonds are the strongest, followed by double bonds, then single bonds
• Ionic bond strength depends on the charge and size of the ions
  • Ions with higher charges and smaller radii form stronger ionic bonds
• Bond length is inversely proportional to bond strength; shorter bonds are generally stronger than longer bonds
• Resonance structures contribute to bond stability by delocalizing electrons over multiple atoms (benzene, C6H6)
• Hybridization of atomic orbitals (sp, sp2, sp3) influences bond strength and geometry

Applications in Real-World Chemistry

• Ionic compounds (salts) have high melting and boiling points, conduct electricity when molten or in solution, and are often soluble in water (sodium chloride, NaCl)
• Covalent compounds have lower melting and boiling points, do not conduct electricity, and are often insoluble in water (glucose, C6H12O6)
• Hydrogen bonding in water is responsible for its unique properties, such as high specific heat capacity, which helps regulate Earth's climate
• Intermolecular forces influence drug design, as the solubility and bioavailability of a drug depend on its ability to form interactions with water and other molecules
• Polymers are large molecules composed of repeating units (monomers) held together by covalent bonds (polyethylene, nylon)
  • The strength and flexibility of polymers can be tuned by adjusting the type and number of bonds between monomers
• Enzymes rely on precise molecular geometry and intermolecular forces to bind substrates and catalyze reactions in biological systems
• Materials science exploits the properties of chemical bonds to create new materials with desired characteristics (carbon fiber, graphene)

Practice Problems and Review

• Determine the type of bond (ionic, covalent, polar covalent) formed between the following pairs of atoms: a) potassium and chlorine (ionic) b) carbon and oxygen (polar covalent) c) nitrogen and nitrogen (covalent)
• Calculate the electronegativity difference between sulfur (2.58) and oxygen (3.44), and determine the bond polarity (polar covalent, electronegativity difference = 0.86)
• Draw the Lewis structure and predict the molecular geometry of sulfur dioxide, SO2 (bent, two bonding domains and one lone pair)
• Rank the following compounds in order of increasing boiling point based on intermolecular forces: methane (CH4), water (H2O), and ethanol (C2H5OH) (methane < ethanol < water)
• Compare the bond strengths of a carbon-carbon single bond, double bond, and triple bond (single < double < triple)
• Explain why the boiling point of water (100°C) is much higher than that of methane (-161°C), despite both being covalent compounds (water exhibits strong hydrogen bonding, while methane has only weak London dispersion forces)
• Discuss the role of chemical bonds in the function of a real-world application, such as the structure of DNA or the properties of a specific polymer (Kevlar, a high-strength polymer used in bulletproof vests, derives its strength from extensive hydrogen bonding between polymer chains)