Final exam study guide notes chm1046 PDF

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chem notes for the final exam and cumulative information from the whole course. Really helpful when I took it...

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Lecture 1: Part 1: Intermolecular Forces 

Describe the types of intermolecular forces possible between atoms or molecules in condensed phases (dispersion forces, dipole-dipole attractions, and hydrogen bonding) o

Dispersion forces 

Dispersion Forces: The force of attraction between an instantaneous dipole and an induced dipole.

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Polarizability: Measure of ease with which electron charge density is distorted by an external electrical field: reflects the facility with which a dipole can be induced. 

The greater the polarizability of molecules, the stronger the intermolecular forces between them

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Larger molecules are more polar because the outer electrons are more loosely bound in larger molecules than smaller molecules so they have greater polarizability

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The higher the atom and molecular masses, the higher the polarizability and intermolecular forces

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Elongated Molecules vs Compact Molecules 

Elongated have greater intermolecular forces because they cover a greater distance

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Dipole-Dipole attractions 

Dipole-Dipole Attraction: When opposite charges, the positive dipole and the negative dipole, on neighboring molecules to attract each other

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Polar molecules have a: permanent separation of charge within their bonds

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Permanent Dipole: When one end of the molecule is electron deficient with a partial positive charge (δ+), while the other end is electron rich with a partial negative charge (δ−)

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The strength of the intermolecular force depends on what? 

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The distance and the relative orientation of the dipoles

What is a property of polar compounds when compared to non-polar compounds of similar molecular masses? 

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Higher melting and boiling points

Dipole-Dipole forces are stronger than this other force: London Dispersion Forces

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o

What types of compounds does Dipole-Dipole attraction occur between?  Any polar compounds Hydrogen Bonding 

Hydrogen Bond Definition: attraction between a highly electronegative atom in one polar bond and a slightly positive hydrogen atom in another polar bond

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Qualities that water has due to the hydrogen bonds: 

High heat capacity

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High heat of evaporation

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High cohesion surface tension

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Excellent solvent for other polar molecules

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Hydrophilic interactions: h- bonds can form between any 2 polar molecules

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Hydrophobic interactions: non polar molecules (cannot dissolve in water)

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Why do strong ionic attractions dissolve in water?

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In the presence of water, attractions between ions can be "replaced" by attractions to the + or - ends of waters polar bonds, so salts dissolve easily in water

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Identify the types of intermolecular forces experienced by specific molecules based on their structures o

The structure of a molecule plays a very important role in determining its intermolecular forces and its chemical properties.

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The method we will be using to predict the shape, or geometry, of a molecule is called the valence shell electron pair repulsion (VSEPR) theory.

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The VSEPR theory is about geometry of compounds and electron location. Electron pairs in a molecule repel one another; therefore, electrons will distribute themselves in positions around the central atom that are as far away from each other as possible.

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The number of electron pairs, non-shared electron pairs, and shared electron pairs, in the valence energy level of the central atom determines the shape of a molecule. The VSEPR shape of a molecule can be determined after drawing a Lewis dot structure.

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o

London dispersion forces: 

Experienced by all molecules and particles.

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Weakest intermolecular force.

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Examples: Ne, O2, CO2.

Dipole-dipole forces: 

Experienced by all polar molecules.

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The strength depends on the difference in electronegativity between the two atoms.

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Examples: H2S, SO2.

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Hydrogen bonding: 

Experienced by polar molecules in which hydrogen is covalently bonded to a highly electronegative element.

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A particularly strong dipole-dipole force where a H atom in one molecule is electrostatically attracted to the N, O, or F atom in another molecule.

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Examples: HF, H2O, NH3, alcohols, and macromolecules.

Ion-dipole forces: 

Experienced by ions interacting with polar molecules in solution.

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They become the stronger as either the charge on the ion increases or as the magnitude of the dipole of the polar molecule increases.

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Examples: NaCl + H2O.

Explain the relation between the intermolecular forces present within a substance and the temperatures associated with changes in its physical state o

Lower temperatures: stronger intermolecular forces

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Higher temperatures: weaker intermolecular forces

Part 2: Properties of Liquids 

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Distinguish between adhesive and cohesive forces o

Adhesive forces are between unlike substances

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Cohesive forces are between like substances.

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(glue to paper//water to water)

Define viscosity, surface tension, and capillary rise o

Viscosity: the state of being thick, sticky, and semifluid in consistency, due to internal friction

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Surface tension: the tension of the surface film of a liquid caused by the attraction of the particles in the surface layer by the bulk of the liquid, which tends to minimize surface area

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Capillary rise: Capillary action is the ability of a liquid to flow in narrow spaces without the assistance of, or even in opposition to, external forces like gravity

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Describe the roles of intermolecular attractive forces in each of these properties/phenomena o

Viscosity: the stronger the IMF, the higher the viscosity.

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Surface tension: the stronger the IMF, the stronger the surface tension

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Capillary rise: occurs when the intermolecular forces in the liquid are weaker than the adhesive forces with the outside substance.

Lecture 2: Part 1: Phase Transitions 

Define phase transitions and phase transition temperatures o

Phase transitions: describe transitions between solid, liquid, and gaseous states of matter 

During a phase transition of a given medium, certain properties of the medium change, often discontinuously, as a result of the change of external conditions, such as temperature, pressure, or others.

o

Phase transition temperatures: the temperature required to induce a phase transition in a certain substance.

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Explain the relation between phase transition temperatures and intermolecular attractive forces o

The stronger the IMF, the higher the phase transition temperatures will be.

Part 2: Phase Diagrams •

Explain the construction and use of a typical phase diagram

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Phase diagrams consist of the following 

Triple point: the point on a phase diagram at which the three states of matter: gas, liquid, and solid coexist

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Critical point: the point on a phase diagram at which the substance is indistinguishable between liquid and gaseous states

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Fusion(melting) (or freezing) curve: the curve on a phase diagram which represents the transition between liquid and solid states

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Vaporization (or condensation) curve: the curve on a phase diagram which represents the transition between gaseous and liquid states

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Sublimation (or deposition) curve: the curve on a phase diagram which represents the transition between gaseous and solid states

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The use of a phase diagram is to be able to distinguish which phase a certain substance will be at a certain temperature and pressure. This allows for the further determination of the properties of the substance in the phase determined.

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Use phase diagrams to identify stable phases at given temperatures and pressures, and to describe phase transitions resulting from changes in these properties o

Use the phase diagram to identify the physical state of a sample of water under specified conditions of pressure and temperature 

For example, using the diagram below, a pressure of 50 kPa and a temperature of −10 °C correspond to the region of the diagram labeled “ice.” Under these conditions, water exists only as a solid (ice). A pressure of 50 kPa and a temperature of 50 °C correspond to the “water” region—here, water exists only as a liquid.

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Describe the supercritical fluid phase of matter o

Supercritical fluid phase: any substance at a temperature and pressure above its critical point, where distinct liquid and gas phases do not exist.

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Behavior of supercritical fluid phase: A supercritical fluid is any substance at a temperature and pressure above its critical point, where distinct liquid and gas phases do not exist. It can effuse through solids like a gas, and dissolve materials like a liquid. Small changes in pressure or temperature result in large changes in density, allowing many properties of a supercritical fluid to be dependent on these changes.

Lecture 3: Part 1: Solid State of Matter •

Define and describe the bonding and properties of ionic, molecular, metallic, and covalent network crystalline solids o

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Ionic 

Bonding: ionic, bonds are electrostatic

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Properties: high melting point, brittle, hard

Molecular 

Bonding: hydrogen bonding, dipole-dipole, London Dispersion forces

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Properties: low melting point, nonconducting

Metallic 

Bonding: metallic bonding

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Properties: Variable hardness and melting point (depending upon strength of metallic bonding), conducting

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Covalent network 

Bonding: covalent bonding

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Properties: high melting point, hard, nonconducting

Describe the main types of crystalline solids: ionic solids, metallic solids, covalent network solids, and molecular solids 

Ionic solids: 

Consists of positively and negatively charged ions held together by electrostatic forces.

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The strength of the attractive forces depends on the charge and size of the ions that compose the lattice and determines many of the physical properties of the crystal.

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Lattice energy: the energy required to separate 1 mole of a crystalline ionic solid into its component ions in the gas phase

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Metallic solids 

The valence electrons are no longer exclusively associated with a single atom, instead these electrons exist in molecular orbitals that are delocalized over many atoms, producing an electronic band structure.

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Essentially, the metallic crystal consists of a set of metal cations in a sea of electrons.

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Covalent network solids 

Formed by networks or chains of atoms or molecules held together by covalent bonds

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A crystal of a covalent solid is one single large molecule

Molecular solids 

Consist of atoms or molecules held to each other by dipole–dipole interactions, London dispersion forces, or hydrogen bonds, or any combination of these

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The intermolecular interactions in a molecular solid are relatively weak compared with ionic and covalent bonds, so molecular solids tend to be soft, low melting, and easily vaporized

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Explain the ways in which crystal defects can occur in a solid o

Crystal defects are imperfections in the geometrical shape of atoms within a crystalline solid

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These defects occur through deformation of the solid which can be caused from the rapid cooling from a high temperature or high energy radiation striking the solid

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These defects can have significant effects on the mechanical, electrical, or optical behavior

Part 2: Lattice Structures •

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Describe the arrangement of atoms and ions in crystalline structures o

Ions and atoms in a crystal lattice structure are bound by electrostatic attraction

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The exact arrangement of ions in a crystal depends on the size of the ions in the crystal

Compute ionic radii using unit cell dimensions

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The relationships between edge length and radius are as follows (for simple cubic cell, body center cubic cell, and face center cubic cell):

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Use given cell dimensions to find edge length (a), and then use the above formulas in order to determine the radius (r)

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Explain the use of X-ray diffraction measurements in determining crystalline structures o

X-ray diffraction is used for phase identification of a crystalline material and can provide information on unit cell dimensions

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The instrument used accelerates electrons towards the sample, and when the electrons are strong enough to dislodge inner shell electrons of the sample, an x-ray spectrum is produced, allowing for the visualization of the structure of the sample

Lecture 4: Part 1: The Dissolution Process •

Describe the basic properties of solutions and how they form o

Solutions form by solvation: the breaking apart of a solute in a solvent to form a homogenous solution

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Solvent particles surround the solute particles

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The solvation process is thermodynamically favorable through the increase in entropy and the exothermic decrease in energy

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Properties 

Rate of solubility: greater solubility = faster solvation

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Concentration: larger ratio of solute to solvent = higher concentration

Predict whether a given mixture will yield a solution based on molecular properties of its components o

Solubility: the maximum amount of a solute that dissolved in a solvent at a given temperature

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o

A higher temperature will allow higher solubility

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Stirring solution will increase solubility

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Smaller particle size will increase solubility (higher surface area of solute)

Explain why some solutions either produce or absorb heat when they form o

Due to the endothermic or exothermic nature of the chemical reaction (heat of solution) which is due to the intermolecular forces

o

Heat of solution is the sum of the enthalpy changes in the separation of solute molecules, separation of solvent molecules, and the formation of solute-solvent interactions

Part 2: Electrolytes •

Define and give examples of electrolytes o

Electrolytes: substances that are capable of carrying electric current when placed in water

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Strong electrolytes: completely dissociate into ions in solution (NaCl)

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o

Weak electrolytes: produce a few ions and do not dissociate fully (HF)

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Nonelectrolytes: do not produce ions, but dissolve in water due to polarity (sugar)

Distinguish between the physical and chemical changes that accompany dissolution of ionic and covalent electrolytes o

Ionic: completely dissociated into its ions

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Covalent: dipole-dipole interactions between the molecule with water rather than the ions with water

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Relate electrolyte strength to solute-solvent attractive forces o

The stronger the electrolyte, the stronger the solute-solvent attractive forces

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This is due to the ability to overcome the intermolecular forces within the molecule and to completely dissociate the molecule into its ions

Part 3: Solubility •

Describe the effects of temperature and pressure on solubility o

Temperature 

Higher temperature increases solubility for solids and liquids (in most cases)

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Lower temperature increase solubility for gases and some exothermic reactions 

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Pressure 

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Keep gas in the solution instead of evaporating away

Higher pressure increases solubility for gases

Why? 

Increases the amount of solute/solvent collisions and increases speed/energy of particles (for solids and liquids

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State Henry’s law and use it in calculations involving the solubility of a gas in a liquid

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Henry’s Law: the solubility of a gas in a liquid is proportional to the pressure of the gas over the solution

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c=kP 

c: concentration in molarity

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k: constant that depends on the temperature (different value for different temperature), units: mol/L*atm

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P: pressure outside of the solution

Explain the degrees of solubility possible for liquid-liquid solutions o

Immiscible: two liquids that do not have any possible intermolecular forces with one another; will not mix or dissolve at all 

o

Nonpolar liquids are immiscible in water

Miscible: two liquids which share intermolecular forces and are able to dissolve or mix with one another 

Polar liquids or substances that form hydrogen bonds are miscible in water

Lecture 5: Part 1: Colligative Properties •

Express concentrations of solution components using mole fraction and molality o

Mole fraction 

The number of moles of a single solute divided by the total moles of all solutes/solvents

o

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Mole fraction Xi= moles of i/moles of total solution

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All mole fractions of individual solutes should total to 1

Molality 

Calculated by taking the moles of solute and dividing by the kilograms of solvent

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Designated by a lower case m

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m= moles of solute/ kg of solvent

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Useful when temperature changes are considered and reports true concentration independent on temperature 

Temperature therefore changes the volume of solutions where masses are not affected by temperature

Part 2: Units of Concentration (Mole Fraction, Molarity, and Molality) 

Mole fraction o

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