Chapter 4.1: Chemical Bonding

• Chemical bonds: electrostatic forces that hold atoms together in compounds lower-energy systems tend to have greater stability than higher-energy systems Bonded atoms are more stable

• Lewis structures: atomic symbol in the middle with its outer valence electrons around

  • Lines: bonding pairs between atoms

  • Dots: lone pairs

• Ionic bonding occurs with atoms of elements of large differences in electronegativity

  • s block metals paired with p block non-metals

    • Can be drawn with an orbital diagram or Lewis structures

  • Bonding allows atoms to obtain valence electron configuration of the closest Noble gas → octet rule is observed

• Ionic solids have:

  • Crystalline with smooth, shiny surfaces

  • Hard and brittle

  • Non-conductors of electricity and heat

  • High melting points

• Lattice energy: energy is given off when ionic crystal forms from the gaseous ions

  • The energy of the same amount is needed to break turn back to the gaseous ions

  • Adding energy decreases stability and processes that readily give off energy become stable

• Covalent bonding: electron-sharing to acquire noble gas configuration

• Bond energy: energy required to break the force of attraction between two atoms and to separate it

  • An increase of bonds means more bond energy

• Covalent compounds:

  • Can be soft solids, liquids, or gas at room temperature

  • Low melting points and boiling points

  • Poor conductors of electricity

  • Not usually soluble in water (acetone is an exception)

• Can use electronegativity difference to determine bonding type

• Ionic bonds: described as “electron transfer”, “losing electrons”, “gaining electrons”

• Polar covalent bond: covalent bond with unequal shared pair of electrons between atoms: ∆EN (change between electronegativities) between 0.4 and 1.7

• Bonded atoms with ∆EN between 1.7 and 3.3 are mostly ionic

• Metallic Bonding

  • Free-electron model: pictures metals as being composed of densely packed core of metallic cations with free roaming electrons

  • Metabolic bonding: force of attraction between positively charged cations and the pool of valence electrons that moves among them

  • Characteristics of metallic bonds:

    • Conductivity

    • Malleability and ductility

    • High melting and boiling points

Chapter 4.2: Molecular Shape and Polarity

• How to draw a Lewis Structure:

  • Least electronegative center (hydrogen and fluoride located at the edge)

  • Determine the total number of valence electrons of each atom (pay attention to charges)

    • Determine the total number of valence electrons to achieve the noble gas configuration

    • If a negative 2 change then add the total number of valence electrons to the calculated valence electrons

    • Add when negative charge, subtract when a positive charge

  • Subtract the number of valence electrons from a number of electrons needed to satisfy the octet rule

    • Divide it by two to give the number of bonds

    • Double bonds are two bonds and triple bonds are three bonds

  • Subtract the number of shared electrons from a number of valence electrons to get lone pairs to draw around the central atom

• Coordinate covalent bond: when a filled atomic orbital overlaps with an empty atomic orbital

• Resonance structure: models that give the same relative position of atoms but show different places for their bonding and lone pairs

  • Do not exit in reality, instead exists as a hybrid

• Expanded valence energy level: when a central atom has more than eight electrons in the valence energy level of the central atom

• Lewis models do not communicate about the molecule’s shape, you need 3D models for that

• Valence-Shell Electron-Pair Repulsion: bonding pairs and lone, non-bonding pairs of electrons in the valence level of an atom repel one another

  • LP-LP > LP-BP> BP-BP

    • LP (lone pair)

    • BP (bond pair)

    • Repulsion greatest in lone pair repulsions

• When all the electron groups are bonding pairs, a molecule will have one of the five geometrical arrangements. If one or more of the electron groups include a lone pair of electrons, variations in one of the five geometric arrangements result.

  • Linear: (2 BP)

  • Trigonal planar: (3BP) or (2 BP, 1 LP)

  • Tetrahedral: (4 BP) or (3 BP, 1LP) or (2 BP, 2 LP)

  • Trigonal bipyramidal: (5 BP) or (4 BP, 1LP) or (3 BP, 2 LP) or (2 BP, 3 LP)

  • Octahedral: (6 BP) or (5 BP, 1LP) or (4 BP, 2LP)

• How to find the molecular shape

  • Determine Lewis structure

  • Find the number of electrons around the central atom

  • Determine which five geometrical arrangement

  • Determine using positions of bonding pairs and lone pairs

• Dipole: charge separation for the entire molecule

Chapter 4.3: Intermolecular Forces in Liquids and Solids

• Intramolecular forces: forces exerted within a molecule or polyatomic ion

• Intermolecular forces: forces that influence physical properties of substances; forces of attraction and repulsion that act between molecules

• Dipole-dipole forces: attractions between opposite charges

  • Ex. Positive and negative

• Ion-dipole forces: the force of attraction between ion and polar molecule

  • Ex. Positive ion attracts with the negative end of the water

• Induced Intermolecular forces:

  • Ion-induced dipole forces: ion in close proximity to a nonpolar molecule that distorts electron density of non-polar molecule making it polarized

    • Ex. Nonpolar O2 with Fe2+

  • Dipole-induced dipole force: similar to an ion-induced dipole force, a polar molecule distorts the electron density of a nonpolar molecule making it polarized

• Dispersion (London) force: the force of attraction between nonpolar molecules larger number of electrons cause an uneven distribution of charge large shape of the molecule are stronger and results in increased boiling point

• Hydrogen Bonding: a strong form of dipole-dipole attraction that exists between a hydrogen atom in a polar-bonded molecule that contains bonds such as O, N, F. There are also unshared pairs of electrons from these electronegative atoms

  • Water pack in a specific way where the solid-state is less dense than the liquid state

  • Soluble for other polar covalent compounds such as alcohols

• Crystalline solids: organized particle arrangements of solids

  • Ex. Amethyst, garnet

    • Bonding and properties of Crystalline Solids:

      • Atomic solids: made up of individual atoms held by dispersion forces

      • Molecular solids: molecules held by dispersion forces, dipole-dipole forces, or hydrogen bonding

      • Network solids: continuous two or three-dimensional arrays

      • Carbon-based network solids:

        • Allotropes: different crystalline or molecular forms of the same element that differ in physical and chemical properties

• Amorphous solids: indistinct shapes because particle arrangements lack order

  • Ex. glass rubber

• Ionic Crystals: array of ions, arranged at regular positions in a crystal lattice

• The crystal lattice is made up of identical repeating unit cells which gives the crystal its unique shape