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mcat gen. chem


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What is ionization energy and what is its periodic trend?
IE = energy required to completely remove an electron from an atom or ion First IE is ALWAYS > Second IE IE increases from left to right and from bottom to top
What is the periodic trend for atomic radius?
Radius decreases from left to right and increases down a group - Effective nuclear charge increases across a period, so radius decreases
What is electron affinity and what is its periodic trend?
EA = energy change that occurs when an electron is added to a gaseous atom - Positive EA represents energy release when an electron is added to an atom - EA increases as effective nuclear charge increases Generalizations about EA: - Alkaline earths have low EAs (full s subshell) - Halogens have high EAs (can get to full octet) - Noble gases have EAs of about 0
What is electronegativity and what is its periodic trend?
Electronegativity = measure of attraction an atom has for electrons in a chemcial bond - Increases from left to right and bottom to top
Describe some important properties and characteristics of metals.
- Located on left of periodic table - Shiny solids at room temp - High melting points and densities - Malleable and ductile - Good conductors (valence e- can move freely)
Describe the important properties of nonmetals.
- Located on right side - Brittle with little metallic luster - High IEs and EN- - Poor conductors
Describe the important properties of the metalloids.
- Found along line between metals and nonmetals - Properties vary widely - Possess characteristics of both metals and nonmetals
Describe the chemistry of the alkali metals.
- Lower densities than normal metals - Only one valence electron (largest atomic radii) - High reactivity --> low IEs - React readily with nonmetals (halogens, especially)
Describe the properties of the transition elements.
- Very hard - High melting and boiling points - Good conductors
What are the physical properties characteristic of ionic bonds?
- High melting and boiling points (strong electrostatic forces) - Conduct electricity in liquid and aqueous states (NOT in solid state) - Form crystal lattices in which attractive forces between ions of opposite charge are maximized and repulsive forces are minimized
What are the physical properties of covalent bonds?
- Weak intermolecular forces - Low-melting solids - Do not conduct electricity in liquid or aqueous states
Describe the two characteristic features of covalent bonds: bond length and bond energy.
1) Bond length -- average distance between the two nuclei - Decreases with bond order 2) Bond energy -- energy required to SEPARATE two bonded atoms - Bond energy (and strength) increases with bond order
List and describe the three types of intermolecular forces.
1) Dipole-dipole - Attractive arrangement of polar molecules with one another - Occurs in solid and liquid phases, but not gas phase - Causes higher boiling points 2) Hydrogen bonding - Unusually strong dipole-dipole interaction - Unusually high boiling points 3) Dispersion forces - Rapid polarization and counterpolarization of electron cloud, causing formation of short-lived dipoles between molecules - Occur in ALL molecules - Weakest intermolecular force - Strength depends on how easily the electrons can move in a molecule
Define isothermal, adiabatic, and isobaric processes.
Isothermal = process that occurs at constant temperature Adiabatic = process that occurs with no heat exchange Isobaric = process that occurs at constant pressure
Define endothermic and exothermic reactions in terms of heat loss or gain.
Endothermic = reactions that absorb heat energy (dH > 0) Exothermic = reactions that release heat energy (dH < 0)
Define what a "state function" is and list 7 state functions in thermochemistry.
State function = property of a system whose magnitude depends only on the initial and final states, and not on the path of the change Examples: - Pressure - Temperature - Volume - Enthalpy (H) - Entropy (S) - Free energy (G) - Internal energy (E or U)
Define enthalpy.
dH (change in enthalpy) = heat absorbed or evolved by a system at constant pressure
Describe 5 ways by which enthalpy is normally measured.
1) Standard heat of formation - Enthalpy change that would occur if 1 mol of the compound were formed directly from its elements in their standard states 2) Standard heat of reaction - Hypothetical enthalpy change that would occur if the reaction were carried out under standard conditions 3) Hess's Law - Enthalpies of reactions are additive - Important to remember to multiply the dH for each reaction by the number of moles 4) Bond Dissociation Energy - Calculating the sum of the total energy input and the total energy output in breaking and forming bonds of a chemical reaction 5) Heats of Combustion - Enthalpy change that occurs during combustion of a hydrocarbon
Describe how dG relates to the equilibrium constant (K) and reaction quotient (Q) for a reaction.
dH* = -RT ln (K) ^^ This applies to reaction conditions at equilibrium However, once a reaction commences the standard state conditions no longer hold, so K must be replaced by Q: Also, can no longer use dH*, must use dH dH = dH* + RT ln (Q)
What are the assumptions required for the ideal gas law? And under what conditions are these assumptions valid?
An ideal gas is a gas whose molecules: - Have no intermolecular forces - Occupy no volume These assumptions are valid under: - Low pressure - High temperature
Describe the two deviations of real gases from the ideal gas law.
1) Deviations due to pressure - Intermolecular forces become more significant as pressure increases - At moderately high P, volume is less than predicted value - At EXTREMELY high P, particle size becomes large compared to distance between them, so gas takes up LARGER volume than predicted by ideal gas law 2) Deviations due to temperature - As T decreases, V(avg) decreases, and attractive forces become more significant - As T is reduced to boiling point, intermolecular attraction causes gas to have a smaller volume than predicted
Describe the corrections made for the deviations of real gases by the van der Waals equation of state.
The "a" term corrects for the attractive forces between molecules - "a" will be smaller for gas such as helium, and larger for more polarizable gases (Xe or N2) The "b" term corrects for the volume of the molecules themselves - Larger values of "b" are found for larger molecules
State the five assumptions of the Kinetic Molecular Theory of Gases.
1) Gases made up of particles whose volumes are negligible 2) Gas atoms exhibit no intermolecular attractions or repulsions 3) Gas particles are in random motion, undergoing collisions with other particles and the container's walls 4) Collisions are elastic (energy conserved) 5) KE(avg) is proportional only to the absolute temperature of the gas
Why is evaporation a cooling process?
Consider a liquid which is at a particular temperature --> the temp is related to the KE(avg) of the molecules - Few molecules will have enough KE to escape the liquid phase and enter the gaseous phase - Every time the liquid loses a high-energy particle, the KE(avg) and therefore the temp decrease
Describe the regions, components, and lines of a phase diagram.
X-axis = temperature Y-axis = pressure Solid: - Low temp, high pressure Liquid: - High temp, high pressure Gas: - High temp, low pressure
List and describe the four colligative properties. What is special about colligative properties?
Colligative properties --> physical props derived solely from the # of particles present, and not the nature of particles 1) Freezing-point depression - Solute particles interfere with the process of crystal formation during freezing, lowering the temp at which the molecules can align into a crystalline structure - dT(f) = K(f)m 2) Boiling-point elevation - Presence of a solute in a solution can lower the vapor pressure of the solution - Since a liquid's boiling point occurs when vapor pressure equals atmospheric pressure, if the vapor pressure is lower, a higher temperature will be required to get the vapor pressure to equal P(atm) - dT(b) = K(b)m 3) Osmotic pressure - Pressure of liquid in a solution that will counterbalance the osmotic pressure of the solute, preventing further flow in either direction - PI = MRT 4) Vapor-pressure lowering - When solute B is added to pure solvent A, the vapor pressure of A above the solvent decreases
What are the two types of salts that are ALWAYS water soluble?
- Alkali metal salts - Ammonium (NH4+) salts
Define "electrolyte". What are strong and weak electrolytes?
Electrolyte = solutes whose solutions are conductive Strong electrolyte = a solute that dissociates completely into its constituent ions - Includes polar compounds (NaCl, KI, HCl) Weak electrolyte = solute that ionizes or hydrolyzes incompletely in aqueous solution - Weak acids (acetic acid) and weak bases (ammonia)
Give the Arrhenius, Bronsted-Lowry, and Lewis definitions for acids and bases.
Arrhenius: - Acid = produces H+ in aqueous solution - Base = produces OH- in aqueous solution Bronsted-Lowry: - Acid = a species that donates protons - Base = species that accepts protons Lewis: - Acid = electron-pair acceptor - Base = electron-pair donor

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