Chapter 4 - Structure and Properties of Substances
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
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