Chemistry Video 7 Flashcards
Ionic bonds
Cations and anions interact via ionic bonds. Strong electrostatic interactions that allow for the formation of ionic compounds.
Occurs when the difference in electronegativity is greater than ~1.8 and electron transfer will occur
Between metal and non-metal
Lattice energy
Ionic bonds; Anions and cations arranged in a. grid. Anions are arranged in a geometry that will allow for the maximization of electrostatic interactions between oppositely charged particles.
Stronger interactions = greater lattice energy = greater energy needed to disrupt the lattice and separate all the ions
2 factors that influence lattice energy
- Magnitude of charge on the ions. Greater magnitude of the charge = greater lattice energy.
- Ion size (ionic radius). Smaller radius = Closer nuclei = stronger lattice attraction
Covalent bonds
Share electrons between 2 atoms. Can be between atoms of the same element or atoms of different elements that do not have large difference in electronegativity.
Covalent compounds are neutral in charge.
Each electron feels the pull of its own nucleus and the pull of the other nucleus
Between 2 non-metals
Ideal internuclear distance
The ideal covalent bond length between 2 atoms. The distance between the 2 protons of 2 atoms. There is a dip down in energy because if each electron is interacting with both protons, then it is a stabilizing and energetically favourable situation.
If the atoms are pushed too close together, it becomes energetically unfavourable because of proton-proton interaction
i.e. 1s orbitals of 2 atoms overlap until ideal bond length is formed.
Nonpolar covalent bond
2 electrons shared precisely evenly between 2 atoms due to an electronegativity difference of less than 0.4
Polar covalent bond
2 electrons shared unevenly between 2 atoms due to an electronegativity difference that is NOT 0
Partial positive and partial negative charges exist. These partial charges are not formal charges
EN difference of 0.4 to 1.8
Dipole arrow
Points in the direction of excess electron density. Larger arrow indicates more polarity
Metallic bond
Pure metals. Make bonds amongst itself.
The metal is positively charged metal cations, with the valence electrons delocalized. The atoms are regarded as being of neutral charge. Allows metals to conduct electricity
Lewis dot structures
For individual atoms. Chemical symbol in centre, surrounded by valence electrons represented by dots
Add or delete electron dots to show ions and add the - or + symbol to the top right corner
Can be used for ionic or covalent compounds
Covalent bonds are represented as a line, instead of 2 electron dots. Lone pairs remain as dots.
Octet rule
Certain elements (C, N, O) prefer to be surrounded by 8 electrons.
However, some elements with access to d orbitals can expand their octets and have more bonds.
Violations of octet rule
Boron and aluminum have 3 valence electrons and can only form 3 covalent bonds.
Hypervalent elements are phosphorus that make 5 bonds and sulfur that make 6 bonds. This occurs because these elements can access d orbitals.
Xenon can make covalent bonds with select elements, such as with 4 Fluorine and 2 sets of lone pairs
VSEPR Theory
Valence shell electron pair repulsion. Electron domains will repel each other.
Linear electron domain geometry
2 electron domains 180 degrees apart. The central atom is sp hybridized
Trigonal planar electron domain geometry
3 electron domains 120 degrees apart. The central atom is sp^2 hybridized, which means 1 s orbital and 2 p orbitals
Tetrahedral electron domain geometry
4 electron domains 109.5 degrees apart. The central atom is sp^3 hybridized, which means 1 s orbital and 3 p orbitals
Trigonal bipyramidal electron domain geometry
5 electron domains has both 90 and 120 degrees apart for bond angles. The central atom is sp^3d hybridized, which means 1 s orbital, 3 p orbitals and 1 d orbital
Axial and equatorial positions are not the same. Lone pairs prefer equatorial positions (120 degrees apart)
Octahedral electron domain geometry
6 electron domains 90 degrees apart for bond angles. The central atom is sp^3d^2 hybridized, which means 1 s orbital, 3 p orbitals and 2 d orbitals
Electron domain geometry
Describes atoms and lone pairs
Molecular shape
Describes atoms only
If molecule has 0 lone pairs
Electron domain geometry = molecular shape
If molecule has some lone pairs
Electron domain geometry IS NOT EQUAL to molecular shape
H2O angle
Angle is 104.5 degrees.
Methane CH4 angle
Angle is 109.5 degrees
Ammonia NH3 angle
Angle is 107 degrees
Net dipole moment
Add up all the bond dipole moment vectors within a molecule.
If there is net dipole, the molecule is polar. If there is no net dipole, the molecule is nonpolar overall, but may have polar bonds
Geometry is a factor.
Polarity influences…
Melting point, boiling point, solubility
Atomic vs hybrid orbitals
Atomic orbitals are like 2s, 2p.
Hybrid orbitals are like sp, sp^2, sp^3 etc. which are equal in energy to each other and lower energy than p orbitals.
sp (2 electron domains) combines s and px. sp^2 (3 electron domains) combines s, px and py. sp^3 (4 electron domains) combines s, px, py and pz.
sp^3d is 5 electron domains and combines s, px, py, pz and 1 d.
sp^3d^2 is 6 electron domains and combines s, px, py, pz and 2 d.
Sigma bond
Direct overlap of covalent orbitals on the internuclear axis.
Pi bond
Lateral overlap of covalent orbitals perpendicular to the internuclear axis. Weaker than sigma bond. But adding pi bonds increases the overall bond strength.
Node
The internuclear axis, no probability of electron density. In pi bonds
Single, double, triple bond
Single = 1 sigma Double = 1 sigma, 1 pi Triple = 1 sigma, 2 pi
Bond strength
Triple bond stronger than double bond stronger than single bond
Resonance structures
Multiple ways of distributing pi electrons (unhybridized electrons in a pi orbital (AKA pi bonds) or lone pairs) within a molecule. The pi electron density is delocalized. The molecule does not flip between the different resonance structures
i.e. partial double bond character
Formal charge
valence electrons normally - # valence electrons around the atom in the current molecule
Composite resonance structure
[ ] Box brackets around entire molecule with dotted lines to show the partial (i.e. double) bond characters.
Partial charge (+ or -) is shown on the top right outside of the box brackets surrounding the molecule
Comparing resonance structure stability
- Fewer formal charges = more stable
- Formal charges are better on specific atoms. More electropositive = better for positive charge. (C+ > N+ > O+). More electronegative = better for negative charge. (O- > N- > C-).
- More full octets = more favourable. Avoid lack of octet on C/N/O.
Molecular orbital theory
Combines atomic orbitals to make molecular orbitals that are delocalized around the entire molecule
Valence bond theory vs molecular orbital theory
Valence bond theory: bonds are localized between 2 atoms, bonds are overlapping atomic/hybrid orbitals, generates sigma/pi bonds, predicts molecular shape
Molecular orbital theory: electrons are delocalized throughout the molecule, combines atomic orbitals to form molecular orbitals, bonding/antibonding interactions, predicts arrangement of electrons
Constructive interference
Increase amplitude of waves.
Occurs when lobes of the same phase overlap.
Electrons are 3D waves that combine with in-phase waves to make a region with high probability of electron density.
Lower energy, sigma bonding orbital. Or pi bonding orbital with the node on the internuclear axis.
Electrons occupy the lowest energy orbital, thus bonding orbitals are filled first.
Destructive interference
Decrease amplitude of waves
Occurs when lobes of opposite phase overlap.
Electrons are are 3D waves that combine with out-of-phase waves to make a region with no probability of electron density. Which is called nodes.
Higher energy, sigma star () antibonding orbital that has a node between the nuclei. OR pi star () antibonding orbital that has a node between the nuclei, which means that the electrons cannot be shared between them.
Molecular orbital diagrams
Atomic orbitals of each atom are shown on either side of the molecular orbital diagram. Fill according to Aufbau Principle and Hund’s Rule.
Calculating bond order
Bonding orbitals increase bond strength. Antibonding orbitals decrease bond strength.
Bond order dictates the number of pairs of bonding electrons between 2 atoms.
Bond order of 1 means single bond. Bond order of 2 means double bond. Bond order of 3 means triple bond.
Bond order = (# of bonding electrons - # of antibonding electrons) / 2
Higher bond order means more bonds and a stronger bond
If the value of the bond order is between 2 integers (AKA a decimal value), it indicates the presence of partial pi bonds. i.e. resonance.
If bond order is 0, it means that a stable bond between the 2 atoms will not form
