IMF and Solutions

TEST 3•29•23

Solute

Substance being dissolved

Solvent

Substance doing the dissolving

* present in greater amounts
* ^^Aqueous^^: solvent is water

Types of Solutions

Solutions created with solvents in different stages of matter

* Gaseous
* gas + gas (air)
* liquid + gas (humid air)
* Liquid
* gas + liquid (carbonated drinks)
* liquid + liquid (vinegar)
* solid + liquid (salt water)
* Solid
* liquid + solid (dental amalgam)
* solid + solid (sterling silver)

Dissociation

Strong Ionic compounds dissociate when dissolved in water

* Dissociate: Ions are separated and dispersed through solution

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Forces of attraction between ions are overome by the attraction between the ions and the polar water molecules.

“Like dissolves like”

Polar solutes dissolve in polar solvents

Nonpolar solutes dissolve in nonpolar solvents

* Explains why oil (nonpolar) and water (polar) will not mix

Spectial Heterogeneous Mixtures

Suspension

* Particles are fairly evenly dispersed
* Usually can be separated by filtration or settling

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Colloid

* Usually not considers a true solution, particles stay dispersed
* Tends to scatter light

Concentrated

Large amount of solute

Dilute

Small amount of solute

Concentration Measurements

A. Mass %

* (Mass solute) / (Mass solution) x 100

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B. Molarity (M)

* (Moles solute) / (Liters solution)

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C. Molality (m)

* (Moles solute) / (kg solvent)

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D. Mole Fraction

* (Moles) / (Total moles)

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Molarity “M”

M = mols of solute / L of solution

* Convert mass to moles by dividing by molar mass
* Convert mL to Liters by dividing by 1,000

Saturation

A saturated solution has the maximum amount of solute dissolved in teh solvent at a specific temperature

* Solution is satured if there are some undissolved crystals

Solubility

Solubility of a substance is the amount that gives an exactly saturated solution

* Expressed in grams solute / 100g solvent at a specific temperature
* Solubitily directly proportionaly to temperature

Supersaturation

Supersaturated solution temporarily holds more solute than is possible at that temperature

* Solute is dissolved at a higher temperature and then solution is lowered back down to original temperature (normally room temperature: 20℃)

Solubility of Gases

* Higher temperature decreases the solubility of gases in a liquid (inversly proportional)
* Higher pressure increases solubility of gases in liquids (directly proportional)
* (S 1) x (P 1) = (S 2) x (P 2)

Dilution

The process of adding more solvent to a solution to achieve a new desired lower concentration

* Concentration (M) will decrease as the Volume (V; in Liters) increases (Indirectly proportional)
* (M 1)(V 1) = (M 2)(V 2)

Solution Concentration

Concentration can be determined by finding mass %, molarity, density, or molaity

* Boiling off extra solvent in a NaCl solution could help us find the mass of solute in the solution
* Measuring the density of sugar water as more sugar is added

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If given 1 or more of these issues:

* Quantity is very small
* There are other solutes in the mixture
* Solvent can’t be boiled off

The you use Spectrophotometry

Spectrophotometry

* Absorption: The cauntity of light that a solution is able to absorb (Differes depending on the wavelength of the light)
* Transmittance: The amount of light able to pass through a solution
* Reflection: The amount of ligth that refracts off of a solution

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Spectrophotometry in the measurement of how much a solution absorbs or transmits

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A Spectrophotometer is the instrument used for spectrophotometry. It measure the amount of light that is transmitted by a sample solution at a particular wavelength

* Different solutions absorb the most at different wavelengths

Beer’s Law

For a given substance, the amount of light absrobed depends on:

* The concentration
* The cell or path length
* The wavelength
* The solvent

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*A* = *εbC*

* *A*: Absorbance
* *ε*: Molar Absorptivity
* *b*: Length of Light Path
* *C*: Concentration (Molarity)

If *ε*, *b*, and *C* are kept constant, the absorbance is proportional to the concentration/molarity of the sample

Determining Concentration

Given a solution with a known identity and unknown concentration, you can determine the concetration via:

* Evaporation
* Measure and weigh a solution sample
* Boil off the solvent
* From the solute mass remaining, calcuate the original concentration
* Conductivity
* An ionic solution will conduct currect
* A conductivity meter can measure the quantity of current
* A curve is developed with known concentrations
* Measure the conductivity of an unknown and extrapolate its concentration from a graph
* Absorbance

Electrolytes

Solutions that carry and electric current

Molality (m)

m = (moles of solute) / (kg solvent)

* At room temperature, 1 L water has a mass of 1 kg
* At room temperature g = mL and kg = L

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Molality is used whenever temperature may change becase the molarity will change with temperature but molality will note

* Molality is used when determining a change in Boiling Point and Freezing Point (colligative properties)

Colligative Properties of Solutions

Properties that depend on the number of solute particles in a solution rather than their identity

* Boiling Point Elevation
* Freezing Point Depression
* Osmotic Pressure
* Vapor Pressure

Boiling Point Elevation

A solute lowers the vapor pressure of the solvent which in turn elevates the boiling point

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ΔT = i • Kb • m-solute

* ΔT = Change in temperature above normal tempertature (boiling = Change from 100℃)
* i = Integer representing Van’t Hoff Factor
* Kb = Molal boiling point elevation constant
* Water: 0.51 ℃/m
* m = Molality

Van’t Hoff Factor “i”

i = number of ions per formula unit

* Nonelectrolytes (molecular): i = 1
* Electrolytes (ions) = Find number of ions in solution due to dissociation
* Count amount of atoms in the compound, but remember that polyatomic ions will count as one single atom

Freezing Point Depression

Solutes lower the freezing point of a solvent

* Salt on icy roads

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ΔT = i • Kf • m-solute

* ΔT = Change from original freezing point (normally 0 ℃)
* i = Van’t Hoff Factor
* Kf = Molal freezing point depression constant
* Water: 1.86 ℃/m
* m = Molality

Osmotic Pressure

Semipermeable Membrance

* Holes of a size where the solvent can pass through by not the solute

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Osmosis:

* Flow of solvent into solution through a membrance

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Osmotic Pressure (π) = minimun pressure that would stop the osmosis

* π = i • c • R • T
* i = Van’t Hoff index
* c = Molar concentration (molarity)
* R = Ideal Gass constant
* T = Temperature in Kelvin

Vapor Pressure of Solutions

The presence of a nonvolatile solute lowers the vapor pressure of the solvent

* Raoult’s Law: P-solv = X-solv • P0-solve
* X = Mole Fraction (mols solvent / total mols)
* P0 = Vapor Pressure of the pure solvent (standard Vapor Pressure)
* “i still plays a factor: If you have an ionic compoud mulitply total moles present in solution by it’s “i” index

Surface Tenstion

When liquid molecules are strongly attracted to eachother (strong intermolecular forces) they minimize the surface area

* When you push water in a pool, it comes in large collections rather than droplets

Capillary Action

The force of adhesion to the wall of a vessel is stonger than the molecular forces of the liquid

* Reason why meniscus is formed in graduated cylinder

Polarity

You can determine a molecue’s polarity based off of its lone pairs and molecular geometry.

* Hydrocarbons have no lone pairs and are normally nonpolar
* Oxygen and Nitrogen have lone electron pairs and normally make organic molecules polar

Intermolecular Forces

Molecules interact with each other or with other molecules

* Explain physical properties (Boiling Point, Freezing Point, Density, etc)
* Subsance with stronger attrctions will have higher:
* Melting/Boiling Points
* Vapor Pressure
* Surface Tension
* Viscosity

Intra- vs. Inter-

Intra:

* Occur within the molecule
* Hold atoms of a molecule together
* About electron and electronegativity (polarity)

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Inter:

* Occur among other molecules
* Weaker than intramolecular forces

Types of Intra/Inter Molecular Forces

Dipole-Dipole

* Bewteen polar molecules (Based on electronegativity; 0.5 - 1.9)

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London Dispersion Forces (LDF)

* Between nonpolar molecules (Based on electronegativity; 0 - 0.49)
* Weakest Force
* Noble Gases and nonpolar

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Hydrogen Bonding

* H bond is a special dipole-dipole force
* Present in molecules where Hydrogen (H) is bonded to Oxygen (O), Nitrogen (N) or Fluorine (F)
* VERY Polar
* Hydrogen a small atom; bonds/atoms are very close
* Water made up of Hydrogen Bonds

Evaportation

Evaporation (Endothermic) requires energy to overcome IF strength in water.

* In an open container evaportation will continue until all of the liquid is vaporized.

Vapor Pressure

Vapor Pressure - The pressure exerted by a vapor in equilibrium wih a liquid.

* In closed container, the rate of evaporation and condensation will eventually equal
* ^^Equilibrium^^
* At a higher temperature Vapor Pressure will be higher (Dependent on temperature)
* A liquid boild when its vapor pressure equals the ^^atmospheric pressure^^
* at a particular temperature, the higher the vapor pressure of a liquid and the weaer the intermolecular forces, meaning a lower boiling point.
* If substance has a high vapor pressure it is called ^^volatile^^

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Volatile Substances

* High Vapor Pressure
* Low Intermolecular Forces
* Evaporates easily

Viscocity

Resistance of a liquid to flow

* Strong intermolecular forces

Chromatography

A technique for separating mixtures into their components in order to analyze, identify and purify the mixture/components

* Uses intermolecular forces to separate compounds
* Travel at different rates based on attraction for mobile phase or stationary phase

Types of Solids

* Amorphus: “Without from”
* Crystal: Pure Substances; Regular/Ordered Arrangement/Repeating Pattern

Molecular Crystals

Strong covalent bonds hold molecules together

* Sugar, ice, dry ice

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Weak intermolecular forces

* Soft, low melting points, easy to dissolve

Netwrok Covalent Solids

Single giant repeating molecule with covalent bonds

* Carbon: Graphite, Diamond, Fullerences

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Hard with very high melting points

Allotropes

A different compoun with the same atoms arranged in diferent formations

* Graphite vs. Diamond vs. Fullerenes/Buckyball

Graphene

Single, 2D sheet of #D Graphite

* Semi-Conductor electrical connectors and batteries

Ion Crystals

Oppositely charged ions packed together in “lattice”

* Strong electrostatic forces between ions
* High melting/boild points
* Dense, retain shape and difficult to compress
* Break, don’t bend

Metallic Solids

Do not form bonds with each other; valence electron decentralize, creating an “electron sea”

* Malleable and ductile, packed into plane structures that can slide

Alloys

Mixture of elements which has improved metallic properties

* Substitutional
* Some atoms relaced by other of similar size (Sterling Silver)
* Interstitial
* Small holes filled by smaller atoms (Steel)