UNL CHEM 109 EXAM 4

Covalent Bonding

The sharing of a valence electron between atoms. They hold the atoms together by attracting the nuclei of both atoms

Ionic Bonding

The transfer of a valence electron from one atom to another. The atoms bond through electrostatic attraction between the anion and cation

lattice energy

The energy required to convert a mole of ionic solid to its constituent ions in the gas phase

Lattice Energy Formula

F ∝ charge / distance^2

F ∝ Q1 x Q2 / r^2

3 trends of Lattice energy

1) Larger ions = bigger distance = weaker attraction

2) Larger charge = stronger attraction

3) Stronger the attraction => larger the lattice energy

Is Ion Charge or Size more important in determining lattice energy

Ion charge

Born- Haber Cycle

A hypothetical series of reactions that represents the formation of an ionic compound from it's constituent elements

Born Haber Cycle Equation

∆Hf˚ = Δsub + Δdiss + IE + EA + ΔHlattꝋ

∆Hf˚ = enthalpy of formation

sub = sublimation

diss = dissolution

IE = ionization energy

EA = electron affinity

Lewis theory of bonding

a chemical bond involves atoms sharing electrons

Lewis Structure

A representation of covalent bonding in which shared electron pairs are shown as dashes between to atoms

Bond Length

The distance between the nuclei of 2 covalently bonded atoms in a molecule

Bond Length Trend

The higher the number of bonds in a molecule, the shorter the bond length

Two attractive forces in covalent bonds

1) The intramolecular bonding force that holds the atoms together in the molecule

2) The intermolecular forces between molecules

Are intermolecular bonds or intramolecular bonds stronger?

Intramolecular

3 characteristics of covalent compounds

- Usually gases, liquids or low melting solids

- Insoluble in water

- Non-conductive

Why are ionic compounds usually solid at room temperature?

Because the electrostatic attraction is usually very strong

3 characteristics of ionic compounds

- They have a high melting point

- Soluble

- Conductive

bond polarity

The extent to which electrons are equally shared between 2 atoms

polar bond

Where the valence electrons spend more time with one atom

non-polar bond

Where there is an equal sharing of valence electrons between two atoms

trend in electronegativity

Electronegativity increases to the right and up the periodic table

Electronegativity

the ability of an atom to attract electrons when the atom is in a compound

ionization energy

The energy required to remove an electron from an atom

Nonpolar bonds are

A bond between atoms whose electronegativities differ by less than 0.5

Polar bonds are

A bond between atoms who's electronegative differ by the range of0.5 to 2.0

Do ionic bonds or covalent bonds have higher bond polarity?

Ionic bonds

dipole moment

A quantitative measure of the polarity of a bond

How is dipole moment represented?

With an arrow extending from a plus sign, extending from the lower to the higher electronegativity

equation for a dipole moment

µ = Q x r

µ = Dipole moment (>0)

Q = the magnitude of partial charges

r = distance

units for a dipole moment

Debye units (D)

1 Debye unit =

3.336 x 10^-30 C.m

Formal Charge

The overall charge of a molcule

Formal Charge Formula

Formal charge = valence electrons -{bonds+dots}

Which lewis structures are preferred?

-most atoms without formal charge

-lowest magnitude of formal charges

-if there is a negative formal charge, its on the most electronegative atom

Resonance Structures

The different potential lewis structures of a molecule

Resonance Hybrid Structures

A combination of resonance structures that represent reality

Octet Rule

That most atoms want 8 valence electrons and they will react with other atoms to get them

3 exceptions to the octet rule

1) Free radicals - ions or molecules with an odd number of electrons

2) Electron deficient atoms that have less than 8 valence electrons

3) Expanded octet atoms have more than 8 valence electrons

characteristics of free radicals

They are extremely reactive and are excellent oxidizing agents because they accept electrons

What molecules can be electron deficient

Any molecule with boron as the central atom

What atoms can have an expanded octet

Any atom on the third row or below on a periodic table

Bond Enthalpy

The enthalpy change associated with breaking a particular bond in 1 mole of gaseous molecules

How do you measure the stability of a molecule

Using the average bond enthalpy

Bond Enthalpy Equation

∆H˚ = ∑(BE of reactants) - ∑(BE of products)

VSEPR

The Valence Shell Electron Pair Repulsion theory

VSEPR theory

That molecules adopt the shape tat minimizes the electron pair repulsion

How to determine the molecular shape

The relative repulsion between electron pairs around the central atom

Molecular Geometry

More specific shape and based off lone pairs. Seesaw, T shape, Bent etc.

Electron Domain Geometry

The First Shape in the Group, the broad category. Linear, Trigonal Planar, Tetrahedral, Trigonal Bipyramidal, Octahedral

difference of 0-0.5 electronegativity

That the bond is non-polar

difference of 0.5-2 electronegativity

That the bond is polar

What does a difference of >2 in electronegativity mean between two bonded atoms

That the bond is ionic

Structural Isomers

Molecules that have the same chemical formula but different arrangements of atoms

valence bond theory

States that atoms share electrons when their atomic orbitals overlap. The two shared electrons in the overlapping orbitals must have opposite spin. Formation of a bond lowers the potential energy of the system and helps the atoms in the bond achieve higher stability.

What are the 2 downsides of the valence bond theory?

1) It gives us a visualization but doesn't describe bonding

2) It does not account for the shapes of most molecules

Hybridization

Where atomic orbitals mix together to form hybrid orbitals that arrange themselves in the same way as the repulsions in the VESP theory

Can double bonded molecules freely rotate? (pi bonds)

No, they are restricted in their position

cis-trans isomers

same covalently bonded atoms but differ in spatial arrangement . cis= both on top. trans= opposite sides

pure covalent bond

neutral atoms held together by equally shared electrons

polar covalent bond

A covalent bond in which electrons are not shared equally

ionic bond

Oppositely charged ions held together by electrostatic attraction

axial positions

The positions above and below the central atom

equatorial position

one of the three positions in a trigonal bipyramidal geometry with 120° angles between them; the axial positions are located at a 90° angle