Intro to IMFs

Intramolecular Forces - forces Within the molecule (covalent bonds)

Intermolecular Forces - forces between molecules (IMFS)

IMFs are much Weaker than intramolecular Forces. If energy is applied, IMFs will break before intramolecular forces (bonds) will.

Why are IMFs important? They are responsible for molecules sticking together and help determine when things Split apart. Therefore, IMFs determine when something undergoes a phase change.

* Note - IMFs are only particularly relevant in covalent bonded compounds. They don't largely affect metallic or ionic bonded substances

Types of IMFs

• London Dispersion Forces (LDF)

• Dipole- Dipore forces

• Hydrogen bonding

London Dispersion Forces (LDFs) (weakest type of IMF) - These exist between all molecules with electrons, but are most notable in non polar covalent molecules. The strength of LDFs are directly related to the number of electrons

* Polarizability - the ability of a molecule to form instantaneous temporary dipoles. The larger the electron cloud, the higher the polarizability; thus, the Stronger the London Dispersion Forces. (as number of electrons increased, so does polarizability, and also LDFs)

Dipole Dipole - This exists between polar molecules due to a permanent, uneven distribution of charge that can be determined by looking at the Lewis Dot Structure for asymmetry or for different exterior atoms.

Symmetric Geometries include: Linear, Trigonal Planar, Tetrahedral, Trigonal Bipyramidal, Square Planar, and Octahedral.

Dipole-dipole forces tend to be much Stronger than LDFs, but Weaker than Hydrogen Bonding.

Ion-Dipole Forces - where ions are attracted to the opposItely charged dipoles of a polar molecule. Ex: NaCl in water.

Hydrogen Bonding - This is a special type of dipole-dipole attraction. However, this is Stronger due to the Electronegativity differences when hydrogen is connected to either Nitrogen, Oxygen, or Flourine (NOF).

* In order to have hydrogen bonding (IMFs) the hydrogen must be Connected to a NOF. It is not enough to just have NOF and Hydrogen present. (must be connected internally!)

To determine which molecule has stronger hydrogen bonding, count the number of hydrogen bonding regions. Hydrogen bonding regions are each set of lone pairs on NOF connected to H.

The Stronger the IMFs, the more the molecules stick together, resulting in different physical properties. (as IMFs increase, it makes the molecules harder to separate, which increases the energy needed to separate them, increasing the boiling point)

Types of Solids:

Molecular solids are formed from molecules, are held together by IMFs, have the lowest relative melting point (<100° Celsius), are never electrically conductive, and are only soluble in water when they are polar.

Covalent Network solids are formed from atoms, are held together by covalent bonds, have the highest relative melting point (1000° Celsius plus), are only electrically conductive if its graphite, and are never soluble in water. Silicone Dioxide (SiO2) is a covalent network solid.

Ionic solids are formed from cations and anions, are held together by lattice energy, have a medium relative melting point (100° Celsius plus), are only electrically conductive when liquid or aqueous, and are always soluble in water (solubility rules).

Metallic solids are formed from metals, are held together by a sea of electrons, have a medium to high relative melting point (100° Celsius plus), are always electrically conductive, and are never soluble in water.

Allotropes are different forms of an element in the same physical state. Examples: oxygen gas and ozone, graphite and diamond.

Properties of liquids

Properties - **All liquid properties depend primarily on the Strength of IMFs present.

1. viscosity -the resistance to flow (Syrup, glue, honey, rubber, silly putty).

   Related to the ease with which individual molecules can flow with respect to one another.

   Factors Affecting Viscosity

   a. Strength of IMFs between molecules - as IMFs increase, viscosity increases.

   b. Temperature - As temperature increases, IMFs break and viscosity decreases.

2. Surface Tension - a measurement of inward forces that must be overcome to break the Surface of a liquid. Occurs because of uneven, Unbalanced IMFs on the surface particles.

   As IMFs increase, surface tension increases

3. Capillary Action - the tendency of liquids to Climb narrow tubes. This is driven by adhesion and cohesion.

   Adhesion - the ability of particles to stick to other, different things.

   Cohesion - the ability of particles to stick to Themselves.

   Meniscus - the curvature that exists at the surface of a liquid.

4. Vapor Pressure - the pressure that develops in the gas phase when a liquid is placed in a closed container.

   Liquids that have high vapor pressure are said to be volatile (they vaporize easily). The amount of gas able to vaporize directly depends on IMFs.

   Vapor pressure is an equilibrium pressure and is not affected by volume.

   Volatility is the ability to turn into a gas.

   Factors Affecting Vapor Pressure

   a. IMFs - as IMFs increase, volatility decreases, and vapor pressure decreases

   b. Temperature (increased temperature breaks IMFs) - as temperature increases, volatility increases, and vapor pressure increases

   c. Molar Mass (heavier things move slower) - as molar mass increases, volatility decreases, and vapor pressure decreases

Phase diagrams

The triple point is where all three phases (gas, solid, liquid) exist at equilibrium.

Solutions

Solution - a homogeneous mixture in which one substance is dissolved in another. Example: seawater

Solute - the substance being dissolved Example: NaCl (salt)

Solvent - the substance doing the dissolving. Example: H2O (water)

Aqueous (aq) Solution - a solution in which H2O is the solvent.

Polar solvents dissolve polar molecules due to dipoles (dipoles attract one another).

Non polar solvents dissolve non polar molecules due to LDFs

Molarity - measure of concentration. (Molarity equals moles of solute divided by liters of solution. M = mol/L, also Molarity times volume equals moles)

Dilutions - adding water to a solution causes the concentration to be decreased. If you ever mix two solutions together, you must use the dilution equation.

Dilution Equation: M1V1 = M2V2

Beer's Law

Some solutions give off distinct color. The intensity of the color is based on the amount of light that is absorbed by the solution. There is a direct relationship that exists between the amount of light absorbed and the Concentration of the solution. This relationship is seen in Beer's Law: A = abc

A - Absorbance (no units)

b- path length (cm)

a - molar absorptivity (1/(M•cm))

c - concentration (M)

Note - Molar absorptivity is constant for a specific compound at a specific wavelength of light.

Identification & Separation Techniques

Different experimental techniques exist in order to identify or separate individual components of a solution.

Distillation - the process of separating out substances from one another through heating due to differences in boiling point. The differences in boiling points enables, upon heating, the lower boiling point substance to escape into the gas phase, while the higher boiling point substance remains as a liquid.

Distillate - the substance turned into a gas and later made back into a Liquid

Filtration - the process of separating out a solid from a liquid. This process is often used after a precipitate has formed. Solid is separated from a liquid by filtering the entire mixture through a porous membrane. The solid remains in the filter paper while the liquid passes through (filtrate).

Chromatography - the process of identifying the contents of a solution due to differences in relative polarity The solvent and the Stationary phase (paper) have different polarities. As the Solvent moves up the stationary phase, substances more similar to paper's polarity will stick while substances with more similarity to the Solvent travel up with the solvent.

The relative polarities can be calculated using the formula: Rf = distance traveled by the Spot (sample) / distance traveled by the water (solvent) = dS / dW

The greater the Rf value, the more similar the spot's polarity to that of the solvent.

Ideal Gas Law (IGL)

Variables

Pressure (atm) - force per area exerted by Collisions of particles with their surroundings.

Volume (Liters) - amount of space available to be taken up by particles. Determined by the container.

*Note - the presence of extra gasses does not change the volume of a rigid container.

Example: a steel box contains oxygen gas. Adding in nitrogen gas does not increase the size of the box.

Temperature (K) - average Kinetic Energy (KE) of all particles.

Amount (Mole, n) - the amount of particles present.

Ideal Gas Law

Boyle’s: P1V1 = P2V2

Charles’: V1/T1 = V2/T2

Gay-Lussac’s: P1/T1 = P2/T2

Avogadro’s: n1/V1 = n2/V2

Combined Gas Law

P1V1/(n1T1) = P2V2/(n2T2)

Ideal Gas Law

PV = nRT

Additional Gas Laws

Dalton’s Law of Partial Pressures - the sum of all the pressures of each individual substance in a container is equal to the total pressure in the container.

P total = P of A + P of B + P of C + …

Mole fraction of “A”: P of A = moles of A / total moles • P total (P_A = n_A/n• Pt)

Collecting Gas Over Water

P(gas) = P(total) - Vapor Pressure(H2O)

Graham’s Law of Diffusion and Effusion - no calculations, only comparisons (molar mass)

Diffusion - the spreading out of a gas

Effusion - gas escaping through small pores in a container (balloons)

The heavier the gas, the slower it moves (as weight increases, speed decreases)

Miscellaneous Relationships

Gases at the same temperature have the same Kinetic Energy

1. Temperature vs. Kinetic Energy

   As temperature increases, so does Kinetic Energy

2. Average Velocity

   Velocity does not equal Kinetic Energy

   As molar mass increases, velocity decreases

   As temperature increases, velocity increases (temperature makes particles move faster)

3. Maxwell-Boltzmann Diagrams - a probability distribution used for describing the speed or energy (E) of various particles.

   Higher temperature causes the distribution to move to the right because the average energy is now increasing OR lighter molecules will be represented by distributions father to the right than heavier molecules due to their differences in speed.

Kinetic Molecular Theory

The model used to predict behavior of ideal gases as seen in the ideal gas law. Real gases do not follow the behaviors predicted using the ideal gas law

Major Assumptions For Ideal Gasses

Real vs. Ideal gases

1. Gas molecules have no volume. However, they have definite mass

   • Real gasses do take up describing, though very small amounts.

2. Gas molecules move randomly and experience completely elastic collisions - there is no loss of energy. This means that individual particles have no effect on one another.

   • Real gasses have IMFs which cause them to interact with one another

Application:

Real gasses behave like ideal gasses under 3 conditions:

1. Low IMFS - reducing the amount of particle interactions reduces the amount of inelastic collisions and lost energy.

2. Raise Temperature - reduces the amount of interaction particles will have due to broken IMFs.

3. Lower Pressure - assumes a large Volume and therefore a low density of particles, reducing the amount of IMFS.