UWorld Interactions of Chemical Substances

Covalent Bonds

Bonds between atoms that are made from sharing electrons, which is made possible from the overlap of the atom’s orbitals from each atom.

Sigma Bonds

Covalent bonds made by the end to end overlap of atomic orbitals

Pi Bonds

Covalent bonds made by the side to side overlap of p orbitals

Double Bond

One pi and one sigma bond

Triple Bond

One sigma bond and two pi bonds

Overall Bond Dissociation Energy

The total energy required to break both the sigma and the pi bonds. Increases with each additional pi bond.

Sigma Bonds vs Pi Bonds

Sigma bonds are lower in energy, more stable, and have a greater dissociation energy than pi bonds

Coordinate Covalent Bonds

Formed between two atoms when both shared electrons are donated by the same atom. Such coordinate bonds are often formed between electron poor metal ions and molecules called ligands that contain one or more electron rich atoms with available lone pair electrons.

Complex

The coordinately bonded metal and its ligands

Coordination Number

The number of coordinate bonds formed between the central metal ion and its neighboring atoms.

Hybridization

Atomic orbitals combine to form new hybrid orbitals that assume specific orientations to minimize electric repulsions between orbitals.

Counting bonds in hybridization

A double bond counts as one, so a double bond and two single bonds = sp²

Coordination Complex

A central metal ion surrounded by one or more molecules called ligands. The ligands act as a lewis base (electron pain donor) and the metal center acts as a lewis acid (electron pair acceptor)

IR-Active

If a net charge in a dipole is produced by a particular vibrational mode and is not cancelled out by another dipole in the molecule. the vibration is IR-Active and will absorb IR light at particular frequencies proportional to the frequency of the bond vibration.

VSEPR Theory (Electron Geometry)

Electron pairs around a central atom in a molecule will adopt an electron geometry/spatial arrangement that maximizes the separation between the electron domains and minimizes the repulsions between electrons.

VSEPR Theory (Molecular Geometry)

The molecular geometry/shape of a molecule can be predicted from its lewis structure by evaluating the orientations that the bonded atom will adopt due to the number of bonding and nonbonding electron domains present around the central atom.

[(CO) OH]+2 Oxidation Number

Everything in [] should equal oxidation number +2

Kw

[H3O+][OH-] (Hydronium and hydroxide) As H3O+ decreases OH- increases and vice versa

Reaction Mechanism

A sequence of elementary reaction that yield an overall reaction. Must be experimentally tested first.

Experimental methods to test proposed reaction mechanisms

1) Verifying the reaction’s rate law. If a reaction follows a proposed mechanism, the rate of the reaction will change according to the mechanism’s rate law as the initial concentrations of the reactants are changed

2) Detecting an Intermediate. If a reaction follows a mechanism that forms an intermediate, adding a substance that reacts only with the intermediate will form a side product that is detectable by instrumental methods (spectroscopy). If no side product is detected, the intermediate does not form and the mechanism is wrong.

3) Interfering with a mechanistic pathway. If there is a unique feature in one mechanism that is not present in another mechanism, an experiment can be modified in a way that would interfere with one mechanism but would not impact the other mechanism.

Intermolecular Forces from weakest to strongest

1. London Dispersion Forces (Between all molecules)

2. Dipole Dipole Interactions (Occur between the negative part of a bond and the positive part of the other)

3. Hydrogen Bonding (Bond with a hydrogen atom)

4. Ion Dipole Interactions

Percent Yield

Actual Yield / Theoretical Yield x 100 (Use ratio from balanced chemical equation to simplify theoretical yield)

Theoretical Yield

The maximum amount of the product that can form if 100% of the limiting reactant is converted into products.

Dipole

Consists of opposite electric charges separated across a distance (chemical bond)

Permanent Dipoles

Form across some covalent bonds due to unequal sharing of electrons between two atoms caused from large differences in electronegativity. The region with greater electron density gains a negative charge and the region with less electron density gains a positive charge.

Dipole Dipole Interaction

Occurs when the opposite partial charges of the permanent dipoles of the polar bonds in neighboring molecules form a mutual attraction.

Ion Dipole Interaction

Full charge of an ion is attracted to the opposite polar charge of the partial charge of a polar bond dipole.

Dipole-Induced Dipole Interaction

The permanent dipole from the polar bond distorts the electron cloud from the neighboring nonpolar bond and creates a weak temporary dipole.

London Forces

Two nonpolar bonds make dipoles with each other.

Electrolytes

Solutes that enable the conduction of electricity within a solvent

Ionic Compounds

Consists of oppositely charged metals and nonmetal ions bound by strong electrostatic attractions (ionic compounds) that results in solids with very high melting points. Only conduct electricity as liquids with free electrolytes and not solid.

Metallic Compounds

Consists of metal nuclei surrounded by delocalized electrons that move freely. Strong metallic bonds give metals a high melting point and make them electrical conductors in all states.

Molecular Compounds

Consists of covalently bonded nonmetals that have weak, nonpolar, intermolecular interactions between neighboring molecules that result in relatively low melting points. Because the bonding electrons are localized (confined within the covalent bonds), they are nonconductive as solids and liquids.

Lewis Acid

Electron pair acceptor

Lewis Base

Electron pair donor

Bronsted Lowry Acid

A molecule that donates a proton ion H+

Bronsted Lowry Base

A molecule that accepts a proton ion H+

Bond Dissociation Energy

Energy required to break a bond. Shorter bonds are stronger and require more energy to break. Short atomic radii form stronger bonds than big atomic radii with weak bonds,

Charge Neutralization in Ionic Bonds

Ionic bonds form when positive charges neutralize negative charges. One ion can carry multiple charges and must be neutralized by the same number of opposite charges.

Net Dipole Moment (Separation of Charge)

Occur when the individual dipoles within the molecule do not cancel each other. Symmetrical molecules do not have net dipole moments while asymmetrical do.

Steps to solving theoretical yield

1. Balance chemical equation

2. Use moles from balanced equation to find grams of each reactant

3. Reactant with less moles is the limiting reactant

4. Grams formed from limiting reactant through molar mass is the theoretical yield

Hydronium Concentration

-log(H+)

pH Equation

pH = 14 - log pOH

How can pH be increased and decreased

pH can be increased by adding hydroxide ions or by removing hydronium ions.

Precipitate

An insoluble solid

Chelate Formation Reactions

A metal cation and a ligand react to form one or more rings via a coordinate bonding arrangement.

Combination Reaction

Two reactants form a single ionic product

Hydrophilic Molecules

Molecules with many polar bonds that promote dipolar interactions with water (More hydrogen bonds = more hydrophillic and polar)

Hydrophobic Molecules

Molecules with many nonpolar bonds that lack attractive dipolar interactions (More nonpolar bonds and less hydrogen bonds)

Magnitude of a Dipole Moment across a Polar Covalent Bond

Equal to the magnitude of the partial charge multiplied by the bond length separating the charge. The difference in electronegativity between two covalently bonded atoms is indicative of the magnitude of the partial charge between atoms. Lower electronegativity = smaller dipole moment and bigger electronegativity = bigger dipole moment

Single Replacement Reaction

Reaction involves the replacement of one atom with another in a reaction between a compound and a neutral element to form a new compound and neutral element.

Decomposition Reaction Example

CaCO3 (s) → CaO (s) + CO2 (g)

Boiling Points and Intermolecular Forces

Structures with weaker intermolecular forces have lower boiling points because less energy is required for a phase change from liquid to gas. Larger molecules with strong IMFs have higher boiling points.

How to find mole consumption

1. Take mass reacting from the question

2. Use molar mass to find moles of the element asked

3. Use mole ratio from balanced chemical equation to find moles consumed asked

What is the mass of F(2) ? (MM = 19 g)

38 g

Trigonal Pyramidal

3 bonds and 0 lone pairs

Greater difference in electronegativity equals:

A larger dipole moment (higher polarity)

T Shaped

3 bonds and 2 lone pairs

Tetrahedral

4 bonds and 0 lone pairs

Trigonal Planar

3 bonds and 1 lone pairs

Polar Molecules

Have a significant net dipole that depends on the number of polar bonds, the strength of the dipoles, and the shape of the molecule.

Complex Ions/Coordination Complexes

Consists of a central metal ion surrounded by one or more ions or molecules called ligands that are bound to the metal by coordinate bonds.

Metals vs nonmetals (cations and anions)

Metals tend to lose electrons to form cations whereas nonmetals tend to gain electrons to form anions. The charge of an atom increases one unit for each electron lost and decreases one unit for each electron gained. Metal cations with a small ionic radius and high positive charge are stronger lewis aids than those with large ionic radius and low charge.

Hydrogen Bonds

Covalently bonded with, nitrogen, oxygen, and fluorine atoms from polar bonds that yield dipoles.

Square Planar

4 bonds and 2 lone pairs

See-saw

4 bonds and 1 lone pair

Identical Particles

Must have the same elemental composition, the same number of electrons, and the same orbital configuration.

Ionic Character

The difference in electronegativity between two bonded atoms is directly proportional to the degree of ionic character of the bond between atoms.

Formal Charge Equation

FC = Valence Electrons - Nonbonding electrons - (Bonding electrons / 2)

Sigma Bonds vs Pi Bonds

Sigma bonds are stronger and more stable than pi bonds, and therefore have a higher dissociation energy than pi bonds. Sigma bonds form first between two atoms, and every bond thereafter is a pi bond. Although individual pi bonds are weaker than sigma bonds, a double bond is composed of a sigma and pi bond, so it is stronger than a single bond.

Resonance Structures

1. Atoms never move, only electrons

2. All resonance structures must have the same total number of valence electrons

3. The octet rules must be obeyed for first and second row elements

4. Only electrons in pi bonds (double or triple bonds) or nonbonding lone pairs can move, not electrons in sigma bonds (single bonds)

5. Electron movement should only be to adjacent atoms when going from one resonance structure to the other.

6. The overall charge of the molecule most not change

Hybrid Orbitals

Formed by combining atomic orbitals for a given atom. To identify hybrid orbitals, add the number of lone pairs and bonds to find the hybrid orbital

Trigonal Pyramidal

3 bonds and 1 lone pairs

Trigonal Planar

3 bonds and 0 lone pairs

Degree of potentiation / Change in system’s response to stimulus is:

Best measured as a ratio. The ratio shows the percent change in signal’s intensity in the presence of potentiator.

As pi bonds decrease:

Bond dissociation energy and strength decreases, bond length increases, and decreases the rigidity

Stronger Lewis Base

1. Charged oxygen atom

2. Additional lone pair electrons

3. More readily donates electrons

Weaker Lewis Base

1. Uncharged oxygen atom

2. Fewer lone pair electrons

Lewis base and acids in coordination complexes

A ligand acts as a lewis base and the metal center is the lewis acid

Coordinate Covalent Bond

Forms by a lewis acid base interaction in which a lone pair of electrons from an electron rich atom in a ligand is shared with an electron deficient metal cation via the overlap atomic orbitals without forming additional ions.

Resonance can influence:

The variables of the dipole moment. The magnitude of the dipole moment of a polar bond is the product of the magnitude of the partial charge and its separation distance

Van Der Waals Forces

Noncovalent interactions between the dipoles of two neutral molecules. Van Der Waals interactions include:

1. Dipole Dipole interactions (interactions between permanent dipoles)

2. Dipole Induced Interactions (Attractions between a permanent dipole and an induced dipole)

3. London Dispersion Forces (Attractions between two induced dipoles)

London Dispersion Forces (Size)

London dispersion forces tend to be more pronounced in large molecules with larger, more polarizable electron clouds

Hydrophobic Molecules

Insoluble in water and nonpolar

Hydrophilic Molecules