Module 2-1

• The Atomic Structure

  • The nucleus contains positively charged protons and uncharged neutrons. Most of the mass of the atom is contained in the nucleus.
  • The electron cloud is composed of negatively charged electrons. The electron cloud comprises most of the volume of the atom.
  • A cation is positively charged and has fewer electrons than protons.
  • An anion is negatively charged and has more electrons than protons.
• Quantum mechanics - electrons have properties of both a particle and wave.
• Orbital = space where an electron spends ~95% of its time
• Atomic orbitals (AOs)

  • electrons are arranged in “shells” at various distances from the nucleus
  • shells are arranged in increasing energy
  • within each shell, electrons are paired into orbitals with each orbital holding two electrons
  • paired electrons within an orbital have opposite spins
• 1 s-orbital - spherical (2 e- )

  3 p-orbitals - dumbbell shaped (6 e- )

  5 d-orbitals - four clover leaf & one dumbbell w/ a ring (10 e- )
• Aufbau principle: Orbitals are filled so that those of lowest energy are filled first.
• Pauli exclusion principle: A maximum of two electrons may be placed in each orbital but only when the spins of the electrons are paired.
• Hund’s rule: If two or more degenerate orbitals are available, one electron is placed in each orbitals with their spins parallel until all are half-full.
• The periodic Table of elements

  • Rows represent the period.
  • Columns represent groups.
  • Elements in the same group have similar properties.
  • Chemical Abstracts Services labeled the groups in roman numerals with A/B, with A for main groups & B for transition metals.
• Valence Electrons

  • electrons in the outermost shell; electrons involved in bonding
• Octet Rule

  • atoms will give up, accept, or share electrons in order to achieve a noble gas configuration
• Formal charge

  • The charge assigned to individual atoms in a Lewis structures.
• The best Lewis structure…

  • is the one with the fewest charges.
  • puts a negative charge on the most electronegative atom.
• Bond Length

  • Distance between the nuclei of two atoms joined by a covalent bond
  • Decreases across a row of the periodic table as the size of the atom decreases.
  • Increases down a column of the periodic table as the size of an atom increases
• Bond Angle

  • Determines the shape around any atom bonded to two other atoms.
  • Two group= linear (180°), three group = trigonal planar (120°), four groups = tetrahedral (109.5°)
• VSEPR Theory

  • Valence-shell Electron-pair Repulsion (VSEPR) Theory – the most stable arrangement keeps the groups as far away as possible from the others to minimize repulsions.
• VSEPR is based on the fact that electron pairs repel each other, so we need to consider the lone pairs.
• Bonding is the joining of two atoms in a stable arrangement
• Through bonding, atoms attain a complete outer shell of valence electron
• Two types of bonding

  • ionic bonds - result from the transfer of electrons from one element to another.
  • Covalent bonds - result from the sharing of electrons between two nuclei.
• Covalent bonds are much more common in organic chemistry than ionic bonds.
• A covalent bond consists of the simultaneous attraction of two nuclei for one or more pairs of electrons.
• Types of Covalent Bond

  • A sigma (σ) bond results when the bonding orbitals point along the axis connecting the two bonding nuclei.
  • A pi (π) bond results when the bonding orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei.
  • The interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore σ bonds are stronger than π bonds.
• Hybridization

  • Orbitals are used to form bonds.
  • All single bonds are σ bonds
  • A σ bond concentrates electron density on the axis that joins two nuclei.
• Hybridization: sp3 hybrids

  • These hybrid orbitals are called sp3 hybrids because they are formed from one s orbital and three p orbitals.
• Hybridization: sp2 hybrids

  • These orbitals are formed when 1 s orbital combines with 2 p orbitals, hence the name sp2 hybrid orbitals. *Note that 1 p orbital is not used to form the hybrid orbital.
• Hybridization: sp2 hybrids

  • A σ bond is formed by end-on overlap of two sp2 hybrid orbitals.
  • A π bond is formed by side-by-side overlap of two 2p orbitals.
• Hybridization: sp hybrids

  • These orbitals are formed when 1 s orbital combines with 1 p orbital, hence the name sp hybrid orbitals. *Note that 2 p orbitals are not used to form the hybrid orbital.
• Importance of Lewis structures

  • A properly drawn Lewis structure shows the number of bonds and lone pairs present around each atom in a molecule.
  • This is the first step needed to determine many properties of a molecule.
• Resonance structures

  • Some molecules can’t be adequately represented by a single Lewis structure.
• Resonance structures

  • Two Lewis structures having the same placement of atoms but a different arrangements of electrons, represented by a double-headed arrow.
• Drawing Resonance structures

  • Curved arrow notation – shows the movement of an electron pair.
  • Tail – shows where the electrons are coming from; always begins at an electron pair (a bond or lone pair)
  • Head – shows where the electrons are going.
• • Resonance structures must be valid Lewis structures.
  • Never exceed an octet for second-row elements.
  • Never break a single bond!
• • positively charged carbon atoms are called carbocations.
  • Carbocations are unstable intermediates because they contain a carbon atom that is lacking an octet of electrons
• Drawing Resonance structures

  1. A lone pair next to a pi bond.
  2. A lone pair next to a positive charge.
  3. A pi bond next to a positive charge.
  4. A pi bond between two atoms, where one of those atoms is electronegative.
  5. Pi bonds going all the way around a ring.
• The resonance hybrid is more stable than any resonance structure because it delocalizes electron density over a larger volume.
• • Delocalized refers to electrons that participate in resonance.
  • Localized refers to electrons that does not participate in resonance
• Isomers

  • Different compounds that have the same molecular formula.
• Constitutional isomers differ only in the way atoms are connected
• Constitutional isomers

  • Different physical properties
  • Different names