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Lydia Bo
Period 4 (IMF (of liquids and solids) (Intermolecular…
Lydia Bo
Period 4
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Solutions
Molarity: mole/ L
molality: mole solute / kg solvent
Mole Fraction: mole of one component in the total moles of all components of the solution
- total mole fraction = 1 and unitless
Mole Percentage: % of moles of one component in total moles of all components in solution
Heat of Solution:
- amount of heat energy absorbed or released when specific amount of solute dissolves in a solvent
*exothermic heat of solution = favored
Energetics of Solution Formation
- ∆H1 > 0 overcoming ALL attractions between solute particles
- ∆H2 > 0 overcome some attractions between solvent molecules
- ∆H3 < 0 form new attractions between solute particles and solvent molecules
∆Hsol'n = ∆Hsolute +∆Hsolvent + ∆Hmix
Raoult's Law: presence of an nonvolatile solute (solids) lowers the vapor pressure of the solvent
Psolution = χ solvent + Psolvent
Liquid-Liquid Solutions (both volatile = nonideal)
Ptotal = Pa + Pb = χaPA + χbPb
Ideal Solution
- liquid-liquid solution follow Raoult's law
- no solution = perfectly ideal
- Negatvie deviations from Raoult's Law (lower than predicted vapor pressure for the solution)
- solute + solvent = similar with strong force of attraction
- ∆H solution is large and (-)
- Positive deviations from Raoult's Law (higher than predicted vapor pressure for solution)
- solute + solvent = dissimilar with only weak forces of attraction
- particles easily escape attractions in solution to enter vapor phase
Colligative Properties: properties whose value depends only on the # of solute particles, not what they are
- vapor pressure lowering
- boiling point elevation
- freezing point depression
- osmotic pressure
Solubility Trends and Limits
- solubility of most solids :arrow_up: with temperature
- rate at which solids dissolves :arrow_up: with :arrow_up: surface area of solid
- solubility of gases :arrow_down: with :arrow_up: in temperature
- solubility of gases :arrow_up: with pressure above solution (force in gas)
- saturated = solution in dynamic equilibrium
- unsaturated = less solute
- supersaturated = more solute than saturated
Acids and Bases
Bronsted-Lowry Theory
- acids: proton donors
- bases: proton acceptors
- defines acid-base rxn as any rxn where H+ is transferred
Arrhenius Model
- acids produce H+ in aq solution
- bases produce OH- in aq solution
Problems with Theory
- does not explain why molecular substance dissolves in H2O to form a basic solution
- does not explain why molecular substances dissolves in H2O to form an acidic solution
Amphoteric Substances: act as either acid or base ---> both transferable H+ and an atom with lone pair electrons
Conjugate Acid-Base Pairs
- Bronsted-Lowry Acid-base rxn
- original base ---> acid in reverse rxn
strong acid/base ---> weak conjugates
- original acid ---> base in reverse rxn
- each reactant and product becomes its conjugate pair
Acid Dissociation
- strong acids dissociate completely in solution
- weak acids assume dissociate slightly in solution
Self-Ionization of water
Acid-Base Properties of Salts
Aqueous Equilibria
Buffered Solutions:
- solution that resists a change in pH when either OH- or protons added
- Buffer solutions contain
weak acid +salt (conjugate base)
weak base +salt (conjugate acid)
Kinetics
Integrated Rate Laws: For the reaction of A ----> products, the rate law depends on the concentration of A
More than one Reactant: Observe the rate with one reactant in very low concentration and the others in much higher concentrations
R=k[A][B][C]
If [B]>>[A] and [C]>>[A] then [B] and [C] do not change as greatly relative to [A]
(Pseudo first order)
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Reaction Mechanisms:
- chemical rxns with an equation listing all reactant and product molecules
- most reactions occur in a series of small rxns (series of elementary steps)
- elementary steps must be the sum to overall rxn
- rate law predicted by mechanism must be consistent with experimentally observed rate law
Rate determining step= slowest step
Intermediates: a product that also shows up as a reactant
Catalyst: first a reactant then becomes a product
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Molecularity: # of species that must collide to produce the rxns indicated by that step
- unimolecular step: rxn involving one molecule
- bimolecular step: rxn with collision of 2 molecules
- termolecular step: rxn with collision of 3 molecules
Factors Affecting Rate Rxn
- mixing gives more particle contact
- particles separated allowing more effective collisions per second
- forming some solutions breaks bonds that need to be broken
Effect of Temperature on Rate
change in temperature=change in rate constant of rate law
k=A[-Ea/e^RT]
Arrhenius Equation: Ln(k)=-Ea/R(1/T)+Ln(A)
- :arrow_up: temperature raises average KE of reactant molecules
- there is a minimum amount of KE needed for collision to be converted into enough PE to form the activated complex
- :arrow_up: temperature is :arrow_up: # of molecules with sufficient KE to overcome Ea
- rxn rate :arrow_up: as concentration or partial pressure of reactant molecules :arrow_up:
- more molecules lead to more molecules with sufficient KE for affective collision
Atomic Structure
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Periodic Trends
Atomic Radius:
- average radius of atom based on measuring large # of element and compounds
- :arrow_down: across period
-adding e- to same valence shell
-Zeff charge :arrow_up:
- :arrow_up: down a group
-valence shell further from nuclues
First Ionization Energy:
- strength of attraction related to most probable distance the valence e- are from the nucleus and Zeff charge valence e- experience
- larger level = farther from nuclues and less attraction
- :arrow_up: across period
- :arrow_down: down group
Electron Affinity:
- energy released when an neutral atom gains electrons ---> CHANGE in energy (gas state)
- defined as exothermic, but may be endothermic
-alkali earth metals + noble gases = endo ---> metals loose e- to get full orbital
- more energy released = larger electron affinity
- :arrow_down: alkali metals
-irregular :arrow_up: in 2nd and 3rd period
- "Generally" :arrow_up: across period
-Group 5A generally lower ---> extra e- needs pair
-Group 2A & 8A generally low ---> added e- to higher energy level
- highest electron affinity in any period = HALOGENS
Ionic Radius:
- :arrow_down: across period
- :arrow_up: down group
- ions in same group = same charge
- cations < anions and neutral atoms (except Rb+ and Cs+ = F- and O2-)
- Larger (+) charge = smaller cation
- Larger (-) charge = larger anion (isoelectronic species)
Electronegativity:
- ability of an atom to attract bonding e- to itself
- :arrow_up: across period
- :arrow_down: down group
- larger difference in electronegativity = more polar the bond
Photoelectron Spectroscopy (PES): method of analyzing matter using electromagnetic radiation
- peaks = e- in single sublevel in atom
- x-axis increases going left
Gases
Nature of Gases:
- expand to fill container
- fluid
- low density
- compressible with space in between molecules
effuse and diffuse
Pressure is cause by collisions of molecules with the walls of the container
1 atm = 101,325 kPa = 760 mm Hg = 760 torr
STP
P = 1 atm
T = 273 K
molar V = 22.42 L
Ideal Gases:
- imaginary gases perfectly fit KINETIC MOLECULAR THEORY
- collision between gas particles and between particles and wall = elastic collisions
- no forces of attraction between gas particles
PV = nRT
P = atm
V = L
n = moles ---> mass/molar mass (m/M)
R = gas constant
= 0.08206 L atm/mol K
= 8.314 L kPa/mol K
= 62.4 L mm Hg/ mol K
T = Kelvins
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Gas Density = mass / volume = molar mass / volume
STP: Density = molar mass/ 22.4 L
Not at STP: Density = MP / RT
[M = molar mass]
Diffusion vs Effusion:
- Diffusion = mixing of gases
- Effusion = gas escaping into an evacuated chamber
Graham's Law
Effusion: Rate of lighter gas 1 / Rate of heavier gas 2 = √M2 / √M1
Diffusion: Distance traveled by gas 1 / Distance traveled by gas 2 = √M2 / √M1
Electrochemistry
Assigning Oxidation #s
Oxidation vs. Reduction
- oxidation: looses e-
- reduction: gains e-
- oxidizing agent = substance that is reduced
- reducing agent = substance that is oxidized
- anode: electrode where oxidation occurs
- cathode: electrode where reduction occurs
Voltage
- potential difference: difference in potential energy between reactants and products
- the amount of force pushing e- through wire = electromotive force (emf)
- current: # of e- that flow through system per second
- electrode surface area dictates the # of e- that can flow
Cell Potential
- difference in potential energy between the anode and cathode in voltaic cell
- depends on relative ease with which oxidizing agent is reduced at the cathode and the reducing agent oxidized at the anode
- standard emf - cell potential under certain standard
Standard Reduction Potential
- standard hydrogen electrode - H+ reduction to H2 under standard conditions assigned potential difference
Half-Cell Potentials
- standard reduction [potentials compare the tendency for a particular reduction half rxn to occur relative of H+ to H2
- half rxn with stronger tend. toward reduction than the SHE have a positive value
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Nernst Equation
Chemical Equilibrium
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Equilibrium Constant (K): ratio of concentrations of products to reactants raised to their stoichiometric coefficients
K=[C][D]/[A][B]
- Amounts and Concentrations DO NOT effect K
- Catalysts = NO EFFECT b/c changes rate, not concentrations
- K>1 product favored b/c [products]>[reactants]
- K<1 reactant favored b/c [reactants]>[products]
- K=1 rxn is at equilibrium
- Inverse: Kbackward= 1/ Kforward
- multiply coefficents by n factor: K^n
- Adding 2 or more chemical equations: K1•K2=Koverall
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Heterogeneous Equilibria
- no solids or liquids included in equilibrium expression
- their concentrations DO NOT change
Reaction Quotient(Q): concentration ratio of products to reactants [any condition not in equilibrium]
Q=[C][D]/[A][B]
- Q=K system is at equilibrium
- Q>K system shift left consuming products (reverse)
- Q<K system shift right consuming reactants (forward)
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Le Chatelier's Principle: if a system at equilibrium is disturbed, the position of equilibrium will shift to minimize the disturbance
When Volume :arrow_up:, Pressure :arrow_down:
Volume :arrow_down:, Pressure :arrow_up:
Pressure goes towards the side with less gas moles
Bonding
Coulomb's Law
describes the attractions and repulsions between charged particles
- strength of interactions :arrow_up: as size of charges :arrow_up:
- electrons are more strongly attracted to a nucleus with a +2 charge than a nucleus with a +1 charge
- ionic compounds form solids at ordinary temperatures
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LIKE charges---> potential energy = (+) and :arrow_down: as particles get further as radius :arrow_up:
OPPOSITE charges ---> potential energy = (-) and becomes more (-) as particles get closer together
Lattice Energy
energy of gas atoms forming solid compound
g+g ---> s (exothermic)
- sublimation (+)
- Dissociation (+)
- Ionization (+)
- Electron Affinity (-)
- Lattice Energy - Ionic crystal (-)
:arrow_up: Q :arrow_up: LE
:arrow_down: r :arrow_up: LE
- Greatest charge
- Smallest radius
- more lattice energy means stronger attraction between ions
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Formal Charges
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- atoms in molecules try to achieve FC as close to ZERO as possible
- and (-) FCs ---> expected on most electronegative atoms
- Σ of FC of all atoms must equal to the overall charge of the ion/molecule
Localized Electron Model (LEM): electron pairs can be thought as belonging to pairs of atoms when bonding using atomic orbitals ---> lone pairs belong to one atom
:arrow_right_hook: Resonance = delocalization/moving
Bonds
Polar vs. Nonpolar
- polar bonds = uneven electron distribution
- nonpolar bonds = even/symmetric electron distribution
Covalent bonds: only with nonmetals
Ionic bonds: nonmetal and metal
- breaking bonds requires energy (+)
- forming bonds release energy (-)
Sigma (σ) = single bonds
Pi (π) = double bonds
- single bonds: 1σ
- double bonds: 1σ, 1π
- triple bonds: 1σ, 2π
Sigma bonds stronger than Pi bonds
Thermochemistry
Standard Enthalpy Change (ΔHº): enthalpy change when all reactants and products are in their standard states
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Standard Enthalpy of Formation (ΔHº f): the enthalpy change for reaction forming 1 mole of a pure compound from its constituent elements
Standard Conditions:
- Standard state:the state of a material at a defined set of conditions
- pure gas= 1 atm pressure
- pure solid and liquid in most stable form= 1 atm pressure and temperature of interest (usually 25ºC)
- substance in a solution with concentration= 1 M
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Hess's Law: going from a particular set of reactants to products, change in enthalpy is the same whether reaction takes place in one step or a series of steps
Calculating ΔH: any reaction can be written as the sum of formation reaction for reactants and products
ΔHºf=∑nΔHº(products)-∑nΔHº(reactants)
Thermodynamics
Standard Entropy Change (ΔS): difference in the sum of the internal energy and PV work energy of reactants and products
ΔSºrxn= ∑nΔSº(products)-∑nΔSº(reactants)
Entropy: the more freedom of motion, increases the randomness of a system ----> more randomness= energy released
-ΔS is not thermodynamically favorable
- random systems require less energy than ordered systems
- driving force for a spontaneous process is an increase in entropy of the universe
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Positional Entropy: probability of occurrence of a particular state depends on the number of ways in which that arrangement can be achieved
- MICROSTATES: solid < liquid << gas
ΔS is (+): entropy change is favorable when result is a more random system
- rxns whose products are in a more random state
- rxns that have a larger number of product molecules than reactant molecules
- increase in temp
- solids dissociating into ions upon dissolving
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Laws of Thermodynamics:
- First Law: Energy is neither created nor destroyed and the energy of the universe is constant
- Second Law: For any spontaneous process, the entropy of the universe increases (ΔS>0)
- Third Law: Entropy of a perfect crystal at absolute zero (0 K) is zero
Gibbs Free Energy (ΔG): max amount of work energy that can be released to the surroundings by a system for a constant temp and pressure system
ΔGºf=∑nΔGº(products)-∑nΔGº(reactants)
ΔG=ΔH-TΔS
Process will be Spontaneous when ΔG is (-)
ΔH (-) ΔS (+) ΔG (-) at any temperature
ΔH (-) ΔS (-) ΔG (-) only at low temperature
ΔH (+) ΔS (+) ΔG (-) only at high temperature
Non-Spontaneous ΔG is (+)
ΔH (+) ΔS (-) ΔG (+) at any temperature
ΔH (-) ΔS (-) ΔG (+) at high temperature
ΔH (+) ΔS (+) ΔG (+) at low temperature
Free Energy and Equilibrium:
- equilibrium point occurs at lowest value of the free energy available to reaction system
- at equilibrium ΔG=0 and Q=K (equilibrium constant)
ΔGº=-RTLn(K)
ΔGº=0 K=1
ΔGº<0 K>1
ΔGº>0 K<1