Unit 13 Flashcards
Intermolecular interactions
-an attractive force that arises between the positive components (or protons) of one molecule and the negative components (or electrons) of another molecule
-for two or more molecules to interact, the molecules must come into contact with one another
Ideal gas law
-equation of state of a hypothetical ideal gas, approximation of the behavior of many gases under many conditions
-PV = nRT
Variables in ideal gas law
-PV = nRT
-P= pressure of the gas (atm)
-V= volume of the gas (L)
-n= number of moles of gas
-T= temperature of the gas (K)
-R= gas constant (0.08206 L atm/mol K)
What can pressure be increased by?
-increase in number of collisions (leads to increase in n)
-increase in the momentum of colliding gas particles (causes increase in KE of gas and increase in temperature, (KE of gas correlates to temperature))
Pressure for mixture of ideal gases
P(total)= (nA + nB + nC)/V
or
P(A) + P(B) + P(C)
mole fraction of a gas in an ideal gas mixture
- (n(A))/(n(total))
Low pressure
-P(actual)= P(ideal)
-ideal gas behavior
-gas molecules occupy a negligible volume compared to the volume of their container
-gas molecules have elastic collisions
-atoms/molecules, on average, are very large distances apart
-interactions are negligible
elastic collision
a collision in which there is no net loss in kinetic energy in the system as a result of the collision
medium pressure
-P(actual)< P(ideal)
-non-ideal behavior
-gas molecules are mutually attractive and cluster together
-average distance between atoms/molecules is small enough that attractive interactions dominate
-when clusters of molecules form, number of collisions decrease
-ATTRACTIONS DOMINATE
high pressure
-P(actual)> P(ideal)
-non-ideal behavior
-the volume occupied by the gas is no longer negligible and the molecules repel each other
-average distance between atoms/molecules is very small
-REPULSIONS DOMINATE
circumstances in which we compare low to medium to high pressure
-V and T are fixed
-n is varied (increases) to see effect on P
Ideal behavior
-occurs when molecules collide and behave as billiard balls
-collisions are described as “elastic”
-collide and transfer momentum (total momentum is conserved)
-have negligible volume
-no intermolecular forces between them (no attraction/repulsion) and no interactions after collision
-only collide elastically with each other and the container walls, meaning they move randomly with no energy loss during collisions
-all gases display ideal behavior over some range of temperature and pressure (usually high T, low P)
Real behavior
-deviates from ideal gas behavior
-particles have a finite volume
-particles experience intermolecular attractive/repulsive forces
-occurs usually at low T, high P
Dalton’s Law of Partial Pressure
-the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the gases in the mixture
-each individual gas exerts a partial pressure that contributes to the total pressure
-can be used to quantify the pressure of the entire mixture (Ideal Gas Law can quantify partial pressure of each component)
Vanderwaals equation
-extends the ideal gas law to include the non-zero size of gas molecules and the interactions between them (real behavior)
-P= ((nRT)/(V-nb)) - ((an^2)/(v^2))
-a: positive constat, related to attractive interactions (large a, lower P)
-b: positive constant, related to repulsive interactions (large b, higher P)
how medium/high pressure correlates to vanderwaals equation
-high pressure: more repulsive forces=higher b, lower P
-medium pressure: more attractive force=higher a, lower P
-low pressure: equal a and b, medium P
molecular dipole
-the separation of charge within a molecule
-one side of the molecule has a partial positive charge and the other side has a partial negative charge
-arises from unequal sharing of electrons between atoms due to differences in electronegativity
-measure of molecule’s overall polarity
-sum of all bond dipoles
bond dipole
-the separation of partial positive and negative charges within a covalent bond
-occurs when electrons are shared unevenly between two atoms due to differences in their electronegativity
-M>0, molecule is polar
-A bond dipole always points from the less electronegative atom towards the more electronegative atom
Bond dipole vs molecule dipole
-A “bond dipole” refers to the polarity of a single chemical bond between two atoms within a molecule, determined by the difference in their electronegativity
-a “molecular dipole” represents the overall polarity of the entire molecule, calculated by summing up the individual bond dipoles considering the molecule’s geometry
-a molecular dipole considers the combined effect of all the bond dipoles in a molecule, taking into account their spatial arrangement
How to determine if a bond is polar or nonpolar
- draw all bond dipoles in correct geometry
- sum dipole vectors head-to-tail
-if bond dipoles point same direction, permanent molecular dipole exists
-if bond dipoles point different directions, no permanent molecular dipole exists (often occurs in symmetrical molecules where bond dipoles cancel one another)
permanent vs temporary molecular dipole
-A permanent molecular dipole refers to a molecule that has a consistent separation of charge due to unequal sharing of electrons between atoms with different electronegativities meaning it always has a positive and negative end
-a temporary/induced dipole is a fleeting separation of charge that occurs due to random fluctuations in electron distribution within a molecule, only present for a brief moment in time
-a non-polar molecule can temporarily become polar due to these fluctuations
Polarizability
-the ability of electron density in the atom/molecule to be distorted when interacting with a nearby atom/molecule/charge
-for any type of intermolecular interaction, the interaction is made stronger when the molecules involved are more polarizable
Factors that affect polarizability
- molecular size (larger molecules are more polarizable (more points of interaction))
- heavier atoms (Cl2 vs I2, I2 is more polarizable due to Z)
- σ vs π electrons (molecules with π electrons are more polarizable)
- surface area (extended molecules are more polarizable (have larger surface area) then compact molecules)
Hydrocarbons
-neutral molecules that contain only C and H
-assume they are nonpolar unless otherwise told
-more extended molecular chains are more polarizable and have stronger id-id interactions
Boiling points (Tbp)
-temperature at which the vapor pressure of a liquid equals the external pressure
-boiling points give us info about the relative strengths of intermolecular interactions
-the higher the boiling point, the more energy is required to break intermolecular interactions (not covalent bonds)
-higher the boiling point, the more energy required to break intermolecular interactions and therefore stronger the intermolecular itneractions
-at temperatures above boiling point, kinetic energies of molecules are high enough to overcome dissociation energy and separate from other molecules
Dissociation Energy (Do)
-the amount of energy needed to break a specific chemical bond in a molecule
-A higher dissociation energy generally corresponds to a higher boiling point because a higher dissociation energy indicates stronger bonds within a molecule, meaning more energy is needed to separate the molecules and cause them to vaporize
Types of intermolecular interactions
-dipole-dipole interactions (d-d)
-dipole induced itnereactions (d-id)
-instantaneous dipole-induced dipole (id-id)
-hydrogen bonding (H-bond)
Dipole-Dipole Interactions
-approximate strength= 2 kJ/mol
-occurs between 2 polar molecules
-strongest interactions typically occur with head-to-tail alignment of molecular dipoles
-ex. HCl-HCl
Dipole-Induced Interactions
-approximate strength= <2 kJ/mol
-occurs between a polar and nonpolar molecule
-occurs when 2 molecules closely approach one another
-permanent dipole of polar molecule influences electron density of nonpolar molecule, causing nonpolar’s electron density to adopt induced dipole in direction of permanent dipole
-ex. HCl-I₂
Instantaneous Dipole-Induced Dipole
-approximate strength= «2 kJ/mol
-sometimes called ‘London Dispersion forces’ or ‘Vanderwaals forces’
-occurs between 2 nonpolar species (atoms/molecules)
-random fluctuation in electron density results in an instantaneous dipole
ex. Xe-Xe
Hydrogen bonding
-approximate strength= 20 kJ/mol
-special class of d-d interactions
-interaction between an acceptor bond dipole and a donor dipole (can assume donor dipole is a lone pair dipole)
-occurs when element is O, N, or F (due to high electronegativity, small size, and lone pairs)
-H-bonds are usually very strong because H is really small, allowing for very close approach of interacting acceptor and donor dipole
-strongest interaction aligns lone pair dipole and bond dipole head to tail with 180 angle
-high electronegativity of N, O, and F cause H electrons to migrate towards these elements giving H a slightly positive charge. This causes H to be attracted to the negativity of lone pairs on different molecules.
Intermolecular Interactions ranked by strength
- Hydrogen bonding (20 kJ/mol)
- Dipole-Dipole (2 kJ/mol)
- Dipole-Induced Dipole (<2 kJ/mol)
- Instantaneous dipole-Induced dipole (dispersion) («2 kJ/mol)
How to tell what molecule has higher boiling point
-decide intermolecular interaction strength, polarity, and size/surface area
-polar, H-bonding, or large surface area have higher boiling point
-compounds with larger molar mass have higher boiling points (increased polarizability of heavier compounds)
Binding Energy
-amount of energy needed to separate two molecules/atoms
Low temperatures
-intermolecular forces dominate
-low kinetic energies permit molecules to be bound together
