AP chem okamoto - unit 3 test

IMFs

Attractions that occur between molecules, affecting how much energy is required to melt, boil, and evaporate a chemical. The stronger the attractions, the more energy required to cause the chemical to melt, boil, and evaporate.

Ionic bonds

The strongest attractions between anions and cations. Solids at room temp, highest melting and boiling point

Dipole-dipole interactions

Moderate attractions which occur between POLAR COVALENT MOLECULES. Gases at room temperature

Hydrogen bonds

Strongest the of D-D interactions, attractions between a molecule with a H bonded to either O, N, or F, and a molecule containing O, N, or F. Liquid at room temp

London dispersion forces

Weak and temporary attractions between NONPOLAR COVALENT molecules. Delta ± charges are INDUCED by neighboring molecules, causing IMF attractions. lowest melting and boiling pts, gases at room temp.

Ion-dipole interactions

The attractions between anions ion and water. How solutions form. Forces of attraction increase as radius of ion increase and charge increases

PE and strength of IMFs

The PE of the electrons that are on the negative pole of the molecule decrease as they approach the positive side of another molecule. The PE decreases as molecules move closer together. Energy must be added in order to weaken or break strong IMFs, so PE increases as the molecules move away from each other.

Species with more electrons and larger electron clouds or Pi bonds are:

MORE POLARIZABLE. Weaker hold on outer electrons, can make stronger LDFs and more likely to form temporary bonds.

Stronger hydrocarbon chains are:

Longer and have more surface area to bond with, more contact points for LDF to occur, and more electrons making it more polarizable.

Properties of ionic solids

Soluble in polar solvents, conduct electricity when molten or dissolved in polar solvent, strong forces of attraction between anions and cations, high melting and boiling pt, very hard, low volatility, not malleable or ductile, brittle,

Molecular solids

Do not conduct electricity, lower melting and boiling pts than ionic solids. Held close together to maximize attractions

Heat of fusion

The amount of heat absorbed as 1 mol of a solid liquifies. Energy is required to expand IMFs as a molecule moves from the solid to liquid phase. Melting is an ENDOTHERMIC process, and heat of fusion values are always positive.

Molecular liquids

IMFs are strong (polar), but not as strong as they are in a solid, so the molecules have more freedom to move.

Heat of vaporization

The heat absorbed as 1 mol of a a liquid becomes gas. Energy is required to sever IMFs as a molecule moves from liquid to gas phase, vaporization is always endothermic, so heat of vaporization values are always positive. Ideally, no IMFs between gas particles.

Vapor pressure

When molecules leave a liquids surface to become a gas, they exert vapor pressure. Rate of evaporation, vapor pressure, and KE increase as temperature increases. Rate of evaporation will be higher in a a substance with weaker IMFs.

Boiling points

A liquid boils when its vapor pressure is equal to the atmospheric pressure. Boiling points decrease as elevation increases. Stronger IMFs require more energy to break, so they will also require a higher boiling pt.

Covalent network solids

One or two nonmetals held together by networks of covalent bonds. Carbon group elements can form covalent network solids as they can form 4 covalent bonds. The highest melting point, very hard, atoms bonded with fixed bond angles.

Graphite

Each carbon forms 3 SP2 orbitals that bond with 3 other carbon atoms. Sheets sit on top of one another, delocalized pi bonds between sheets. Weak pie bonds and LDFs allow sheets to slide (pencils). High melting point

Synthetic polymers

Non-polar, long hydrocarbon chains held together by weak LDFs, flexible/viscous liquids, increases flexibility when heated.

Metallic solids

Non-covalent bonding, bonding results from attractions between nuclei and delocalized electrons moving throughout structure (sea of electrons), conduct heat and electricity well, malleable and ductile, lack directional bonds.

When heat/KE is added to a gas:

The velocity of the gas will increase, the volume will increase

When the external pressure of a container holding a gas increases:

The volume of the container decreases, and the internal pressure increases.

Ideal gas law

No IMFs, no change in condensation, volume of gas is 0,

Miscible

Liquids mix

Immiscible (alcohol and water)

Liquids don’t mix (oil and water)

Solubility of gases

Gases are infinitely soluble in each other.

Solid-solid solutions

Formed by melting, mixing, then solidifying. Forms interstitial and substitutional alloys.

Expressing concentration with molarity (M)

Molarity = moles solute/liters solution

Expressing concentration with mole fractions

X = Moles of element/ total Moles of substance

Factors affecting solubility

“Like dissolves like”, temperature, pressure

Like dissolves like rule (not to be used as reasoning)

Substances with similar IMFs tend to be more soluble or miscible in one another.

Saturated

Only applies to liquids: a solution that holds the maximum amount of a solute it can dissolve at a given temperature

For ion dipoles, substances that are more soluble in water have:

Larger radii, more charges. They contain more electrons and have a longer bond length then the bonding nuclei, so the bond is easier to beak, and then the H2O molecules can ion-dipole with eh separated ions.

How does chromatography paper work?

The paper is composed of nonpolar carbon chains and -OH groups that can form H bonds. As the solvent carries the solutes up the paper, nonpolar solutes will travel farther up the paper as they have weak attractions for the paper and stronger IMFs for the mainly nonpolar solute. Polar solutes that can H bonds will not travel farther up the paper with the solvent as they will form H bond with the paper closer to the start point.

Fractional distillation

The separation of volatile liquids on the basis of boiling points.

Gas solubility

Gas solubility decreases as temperature increases or pressure decreases. As KE increases, all particles become less attracted and make a clear separate gas phase. The solubility of a gas is directly related to the partial pressure of the gas above the solution.

The 7 types of waves, most to least energy/frequency:

Gamma / X / UV / visible / infrared / micro / radio

Visible light spectrum, most to least energy/frequency:

Violet, Indio, blue, green, yellow, orange, red

Crest

The peak/highest point of the wavelength

Trough

The bottom or lowest point of the wavelength

Main idea of quantum theory

Energy radiated from a heated object is emitted in discrete units, or quanta

The photoelectric effect

Certain metal surfaces can shed electrons only when the frequency of the photons in the light shined on the surface reach a high enough threshold. When the threshold is reached, electrons are ejected immediately. Increasing the frequency of the light increases the velocity of the ejected electrons.

What happens when you increase the intensity of the light being shined on a metal surface in the photoelectric effect?

If the frequency of the light meets the threshold required to eject electrons, then increasing the intensity will cause the electrons to eject at a higher rate.

Photons

A stream of particles that make up a beam of light, directly related to the light’s frequency.

Continuous spectrum

Every wavelength of light is represented in the continuous spectrum, al colors are visible from violet to red.

Atomic emission spectrum

A series of colored lines at specific wavelengths against a dark background. They form when electrons fall back down to the ground state from the exited state, moving back down energy levels, and emit photons with specific energies, creating colored lines that correspond to a unique wavelength for the specific element.

Absorbing photons

When a photon is absorbed by an atom or molecule (light shined on it), an electron moves up one or more energy levels. The increase in energy is equal to the energy of the photon that was absorbed and the difference in energy between the two energy levels.

Emitting photons

When a photon (light) is emitted from an atom or molecule, an electron moves down one or more energy levels. The decrease in energy is equal to the energy of the photon that was releases and the difference in energy between the energy levels.

Spectroscopy

A method of chemical analysis which is based on the absorbance or emission of light by atoms/ions. Used to identify an element, find the concentration of a colored solution, and find the structure of a molecule.

UV/Vis Spectroscopy

Ultraviolet/visible light is shined on a sample of atoms, and the light emitted reveals the energy differences between the ground vs exited state. Helps identify the element based on its emission spectrum because the types of light emitted are unique to each element.

Visible light spectroscopy

Determining the concentration of solute in a solution using beer’s law. A colorless solute will not absorb visible light.

Infrared spectroscopy

Measures the vibrational frequencies absorbed by the molecule, and detects the presence of different types of bonds to identify bond order, carbon chains, and the compounds identity.

Microwave spectroscopy

The light absorbed changes the rotation of the bonded atoms and tells us the locations of the atoms in a molecule

Absorption spectrum

A graph with peaks that represent the wavelengths of light that correspond to the energy absorbed by an electron traveling from the ground state to the exited state. The tallest peak represents the wavelength most absorbed by electrons.

Ideal gas equation

PV=nRT