1. Atomic Structure and Properties
* Periodic table
* Alkali metals, alkaline earth metals, transition metals, halogens, noble gases
* Mass number = P + N
* Isotopes - atoms of an element with different numbers of neutrons
* Average atomic mass from weighted average of isotope mass and relative abundance (frequency)
* Moles
* PV = nRT
* Avogadro’s number 6.022*10^23
* AT STP (1 atm, 273K), 22.4 L/mol
* Molarity M = moles/L
* Percent composition - divide the mass of each element/compound by the total molar mass of the substance
* Empirical formula is simplest ratio, molecular formula is actual formula for substance
* Energy
* Electron potential energy increases with distance from nucleus
* Electron energy is quantized - can only exist at specific energy levels at specific intervals, not in between
* Coulomb’s law: F = kq1q2/(r^2) where F is electrostatic force
* Atoms absorb energy in the form of electromagnetic radiation as electrons jump to higher energy levels; when electrons drop levels (closer), atoms give off energy
* Photoelectron spectroscopy
* energy measured in electronvolts (eV)
* Incoming radiation energy = binding energy + kinetic energy of the ejected electron
* Electrons that are further away from nucleus require less energy to eject, thus will move faster
* Photoelectron spectrum
* Each section of peaks represents a different energy level (1, 2, 3, etc.)
* Subshells within each energy level (shape of space electron can be found in orbiting nucleus) are represented by the peaks (1s, 2s, 2p, etc.)
* s(2) - first subshell, p(6) - second subshell
* Height of peaks determines number of electrons in subshell (ex. Peak of p subshell in energy level 2 will be 3x that of s subshell)
* Electron configuration
* Electron configuration - spdf - shorthand with noble gas first
* Configuration rules
* Aufbau principle - electrons fill lowest energy subshells available first
* Pauli exclusion principle - 2 electrons in same orbital cannot have same spin
* Hund’s rule - Electrons occupy empty subshells first
* Zn +2, Ag +1, Al +3, Cd +2, most other transition metal charges vary
* Periodic trends
* Electrons are more attracted if they are closer to the nucleus, or if there are more protons
* Electrons are repelled by other electrons - if there are electrons b/w the valence electrons and nucleus, the e- will be less attracted (shielding)
* Completed shells are very stable, completed subshells are also stable; atoms will add/subtract valence electrons to complete their shell
* INCREASING: atomic radius down left; ionization energy up right; electronegativity up right
* Ionization energy - energy required to remove an electron from an atom
* Electronegativity - how strongly the nucleus of an atom attracts electrons of other atoms in a bond
* Electron affinity - energy change that occurs when an electron is added to an atom in the gas state (usually exothermic - energy is released)
2. Molecular and Ionic Compound Structure and Properties
* Bonds
* Atoms are more stable with full valence shells
* Ionic bonds
* Cation gives up electrons completely
* Electrostatic attractions in a lattice structure
* Metals and nonmetals (salts)
* Coulomb’s law - greater charge leads to a greater bond/lattice energy (higher melting point)
* If both have equal charges, smaller radius will have greater coulombic attraction
* Ionic solid - electrons do not move around lattice; ionic solids are poor conductors of electricity; ionic liquids conduct electricity because ions are free to move around, though e- are still localized around particular atoms
* Metallic bonds
* Sea of electrons model - positively charged core is stationary while valence electrons are very mobile
* Metals bond to form alloys - interstitial alloy w/ metals of different radii; substitutional alloy w/ metals of similar radii
* Molecular covalent bonds
* 2 atoms share electrons - both atoms achieve complete outer shells
* 2 nonmetals
* Creates molecules - 2+ atoms covalently bonded together
* Single has 1 sigma bond - order 1, longest length, least energy; double has 1 sigma and 1 pi bond - order 2, int. length, int. energy; triple has 1 sigma and 2 pi bonds - order 3, shortest length, greatest bond energy
* Bond forms when potential energy is at minimal level
* Too close - potential energy is too high due to repulsive forces
* Too far - potential energy is near 0 because attractive forces are very weak
* Minimul PE occurs when repulsive and attractive forces are balanced
* Network covalent bonds - lattice of covalent bonds - poor conductors, high melting and boiling points
* Conductivity
* Conductivity of different substances in different phases
Solid
Aqueous
Liquid
Gas
Ionic
No
Yes
Yes
No
Molecular Covalent
No
No
No
No
Network Covalent
No
N/A
No
No
Metallic
Yes
N/A
Yes
No
* Lewis dot structures
* Resonance - for bond order calculations, average together all possible orders of a specific bond
* BORON (B) is stable with 6 electrons - only one that does not need a full octet
* Expanded octets - any atom of an element from n=3 or greater (those with a d subshell) can have [8,12] valence electrons on center atom
* Noble gases form bonds by filling empty d orbital with electrons
* Formal charge - number of valence electrons minus assigned electrons (1 e- for each line “shared” bond) - 0 for neutral molecules
* Molecular geometry (VSEPR)
* Double and triple bonds have more repulsive strength than single bonds - occupy more space
* Lone electron pairs have more repulsive strength than bonding pairs, so molecules with lone pairs will have slightly reduced angles between terminal atoms
* Hybridization - how many atoms are attached (sp, sp2, sp3, sp3d, etc.)
3. Intermolecular Forces and Properties
* Polarity
* Covalent bond where electrons are unequally shared - polar covalent
* Dipoles are caused by polar covalent bonds - pair of opposite electric charges separated by some distance, like partial charges on atoms in a polar covalent bond
* If 2 identical atoms bond (ex. Cl-Cl) the electrons are equally shared, creating a nonpolar covalent bond with no dipole
* Bonds can be polar; so can molecules depending on the molecular geometry (and polarity of bonds - secondary)
* In polar molecules, more electronegative atoms will gain negative partial charge
* Usually central atom will be positive - exception is hydrogen (terminal), which is usually positive since it has less electronegativity
* Intermolecular forces
* Forces b/w molecules in a covalently bonded substance - need to be broken apart for covalent substances to change phases
* Changing phase: ionic substances break bonds b/w individual ions; covalent substances keep bonds inside a molecule in place but break bonds b/w molecules
* Dipole-dipole forces
* Polar molecules - positive end of one molecule is attracted to negative end of another molecule
* Greater polarity -> greater dipole dipole attraction -> larger dipole moment -> higher melting/boiling points
* Relatively weak overall - melt and boil at low temps
* Hydrogen bonds
* Special type of dipole-dipole attraction where positively charged hydrogen end of a molecule is attracted to negatively charged end of another molecule containing an extremely electronegative element (F, O, N)
* Much stronger than normal dipole-dipole forces since a hydrogen atom “sharing”/giving up its lone e- to a bond is left w/ no shielding
* Higher melting/boiling points than substances held together only by other types of IMF
* London dispersion forces
* All molecules - very weak attractions due to random motion of electrons on atoms within molecules (instantaneous polarity)
* Molecules w/ more e- experience greater LDF (more random motion)
* Higher molar mass usually means greater LDF (as mass increases, e- increases for the molecule to remain electrically neutral)
* IMF strength
* Ionic substances are generally solids at room temp - melting them requires lattice bonds to be broken - necessary energy determined by Coulombic attraction
* Covalent substances (liquids at room temp) boil when IMF are broken; for molecules of similar size, from strongest to weakest: hydrogen bonds, permanent dipoles, LDF (temporary dipoles - greater for larger molecules)
* Melting/boiling points of covalent substances are LOWER than for ionic substances
* Bonding/Phases
* Substances w/ weak IMF (LDF) tend to be gases at room temp (N2); substances w/ strong IMF (hydrogen bonds) tend to be liquids at room temp (H2O)
* Ionic substances do not experience IMF - since ionic bonds are stronger than IMF, ionic substances are usually solids at room temp
* Vapor pressure
* Molecules in a liquid are in constant motion - if they hit the surface of the liquid with enough kinetic energy, they can escape the IMF holding them to other molecules and transition into the gas phase
* Vaporization (NOT boiling) - no outside energy is added
* Temperature and vapor pressure are directly proportional
* At the same temp, vapor pressure is dependent on strength of IMF (stronger IMF, lower vapor pressure)
* Solution separation
* Solutes and solvents - like dissolves like
* Paper chromatography
* Piece of filter paper with substance on the bottom is dipped in water
* More polar components of substance travel further up the filter paper with the polar water
* Distance substance travels up the paper measured by retention/retardation factor Rf = (distance traveled by solute - substance being separated)/(distance traveled by solvent front - water)
* Stronger attraction - larger Rf
* Column chromatography
* Column is packed with a stationary substance
* separable solution (analyte) is injected, adhering to stationary phase
* another solution (eluent - liquid/gas) is injected into column
* more attracted analyte molecules will move through faster and leave column first
* Distillation
* Takes advantage of different boiling points of substances by boiling a mixture at an intermediate point
* Vapor is collected, cooled, and condensed back to a liquid separate of leftover liquid
* Kinetic molecular theory
* Kinetic energy of a single gas molecule: KE = ½ mv^2
* Average kinetic energy of a gas depends on the temperature (directly proportional), not the identity of the gas (different gases will have same KE at same temp)
* Ideal gases have insignificant volume of molecules, no forces of attraction b/w molecules, and are in constant motion without losing KE
* Deviations occur at low temperatures or high pressures (gas molecules are packed too tightly together)
* Volume of gas molecules becomes significant (less free space for molecules to move around than predicted)
* Gas molecules attract one another and stick together (real pressure is smaller than predicted pressure)
* Maxwell-boltzman diagrams
* Higher temp -> greater KE -> greater range of velocity
* Smaller masses, greater velocities to have same KE
* Effusion
* Rate at which a gas escapes from a container through microscopic holes
* High to low pressure
* Greater speed, greater temp, greater rate of effusion
* If at same temp, gas w/ lower molar mass will effuse first
* Equations
* Ideal gas equation: PV = nRT
* R=0.0821
* Combined gas law: P1V1/T1 = P2V2/T2
* Dalton’s law: P(total) = Pa + Pb + Pc + …
* Partial pressure: Pa = P(total)*(moles of gas A)/(total moles of gas)
* Density: D = m/V
* From ideal gas law: Molar mass = DRT/P
* Electromagnetic spectrum
* E=hv
* E = energy change; h = Planck’s constant 6.626*10^-34; v = frequency
* C = lambda * v
* C = speed of light 2.998*10^8; v = frequency; lambda = wavelength
* Beer’s law: A = abc
* A = absorbance; a = molar absorptivity (constant depending on solution); b = path length of light through solution (constant); c = concentration of solution
* Colorimetry - direct relationship b/w concentration and absorbance
4. Chemical Reactions
* Types of reactions
* Synthesis: everything combines to form one compound
* Decomposition: one compound + heat is split into multiple elements/compounds
* Acid-base rxn: Acid + base -> water + salt
* Oxidation-reduction (redox) rxn: changes the oxidation state of some species
* Combustion: substance w/ H and C + O2 -> CO2 + H2O
* Precipitation: aqueous solutions -> insoluble salt (+ more aq sometimes)
* Can be written as net ionic - Those free ions not in net ionic are spectator ions
* Solubility rules
* Alkali metal cations or ammonium (NH4+) cations are ALWAYS soluble
* Compounds with a nitrate (NO3-) anion are ALWAYS soluble
* Common polyatomic ions
* Calculations
* Percent error: 100 * abs(experimental - expected)/(expected)
* Combustion analysis - use law of conservation of mass (if x g of CO2 is produced, find g of C which will be starting amt)
* Gravimetric analysis - when asked to determine the identity of a certain compound, find g of component produced, then use mass percent (g found / total sample mass) and compare to mass percent of options (molar mass of component / molar mass of entire compound)
* Oxidation states
* Neutral atoms not bonded to other atoms have an oxidation state of 0
* Monoatomic ions have an oxidation state equal to the charge on that ion (ex. Zn2+ will be +2)
* Oxygen is -2 (EXCEPTION: in hydrogen peroxide, H2O2, O is -1)
* Hydrogen is +1 w/ nonmetals, -1 with metals
* In absence of oxygen, most electronegative element in a compound will take an oxidation state equal to its usual charge (ex. F is -1 in CF4)
* IF none of the above rules apply, determine the oxidation state by adding up all elements’ oxidation states to 0/charge on ion
* C, N, S, P frequently vary oxidation states (low electronegativity)
* Redox reactions
* Write full rxn as 2 half reactions (oxidation and reduction; OIL RIG)
* Add H2O to compensate for oxygen on one side
* Add H+ to compensate for H from H2O on other side
* Balance 2 half rxns to have the same number of electrons and add them together to produce one complete reaction
* ACIDIC: stop here
* BASIC: Add OH- to both sides - enough for all H+ on one side to be converted to H2O; then cancel out H2O so it only remains on one side
* Acids and bases (briefly)
* Color change signals the end of a titration (can be redox or acid/base)
* Acids are capable of donating protons (H+); bases are capable of donating electrons
* Species with the H+ ion are acids, same species but without H+ is a base - conjugate acid/base pairs
* Water can act as an acid or base - amphoteric
5. Kinetics
* Rate law
* Rate = k [A]^x [B]^y [C]^z
* Can calculate x, y, z via a table from (concentration factor)^x = (rate factor)
* K is only dependent on temperature (always increases w/ T)
* Keq = K1 (rate constant of forward rxn) / K2 (rate constant of reverse rxn)
* K calculated by dividing any rate in table by the concentrations to their respective powers
* Units for rate are M/s, units for conc are M -> calculate units for k from there
* If A + 2B + C -> D; rate of formation of D = rate of disappearance of A and C = 0.5* rate of disappearance of B
* Orders
* Zero-order
* Rate = k
* Concentration vs time has slope -k
* First-order
* Rate = k[A]
* ln[A] vs time has slope -k
* ln[A]t = -kt + ln[A]0
* Second-order
* Rate = k[A]^2
* 1/[A] vs time has slope k
* 1/[A]t = kt + 1/[A]0
* Half-life
* First order reactions only have a constant half life
* t1/2 = ln(2)/k = 0.693/k
* Collision theory
* Chemical reactions occur because reactants are constantly moving and colliding with one another
* When reactants collide with sufficient energy (activation energy Ea), a reaction occurs
* Gaseous/aqueous: increased concentration increases rate of reaction (more likely to collide)
* Stirring increases reaction rate for heterogeneous mixtures (causing heterogeneous mixture to move around increases collisions; insignificant once the mixture becomes homogeneous due to the number of collisions happening due to inherent motion of aq molecules)
* Greater temp increases reaction rate (greater fraction of reactant molecules has sufficient energy to exceed activation energy barrier - vertical line on Maxwell-Boltzmann w/ multiple temps)
* Reactions only occur if reactants collide with correct orientation to break the right bonds
* Reaction energy profile
* Reaction mechanisms
* Species that are produced in a mechanism but are also fully consumed and do not appear in the balanced equation are intermediates
* Adding up all mechanism steps and canceling out different species leads to the balanced rxn
* Elementary steps w/ 2 reactants (even if they are the same) are bimolecular; elementary steps w/ 1 reactant are unimolecular
* Speed is determined by slow step (rate determining step)
* consistency is determined by slow step and those leading up to it
* Make rate for slow step (ex. If X + B -> Y, rate = k[X][B])
* Substitute in rate for X from above equilibrium reaction
* Compare to actual reaction’s rate equation
* Slow step has highest activation energy
* Catalysts
* Catalysts increase rate of chemical reaction without being consumed in the process
* Catalysts do not appear in balanced equation
* In a reaction mechanism, catalysts enter first, then exit
* Catalysis (reaction with a catalyst)
* Surface catalysis - reaction intermediate is formed
* Enzyme catalysis - catalyst binds to reactants to reduce activation energy
* Acid-base catalysis - reactants lose/gain protons to change reaction rate
6. Thermodynamics
* Temperature/heat
* Temperature is the average amount of kinetic energy due to molecular motion in a given substance
* Heat is the energy flow between 2 different substances at different temperatures
* First law of thermodynamics: energy can be neither created nor destroyed
* When bonds are formed, energy is released; when bonds are broken, energy is absorbed
* Exothermic - energy transferred from system to surroundings (delta H is negative)
* More energy is released when the product bonds form than is necessary to break reactant bonds
* Endothermic - energy transferred from surroundings into system (delta H is positive)
* More energy is required to break reactant bonds than is released when bonds in products form
* Energy diagrams
* Enthalpy
* Enthalpy of formation
* Change in energy when one mole of a compound is formed from its component pure elements under standard conditions (25C/298K)
* Delta Hf = delta Hf for products - delta Hf for reactants
* Multiply delta Hf for each product/reactant by the coefficient
* If delta Hf is negative, energy is released when the compound is formed, so the product is more stable (exothermic)
* If delta Hf is positive, energy is absorbed when the compound is formed, so the product is less stable than its constituent elements (endothermic)
* Heat of formation is 0 when the pure element is in its standard state (ex. H2(g) or F2(g))
* Bond energy
* Delta H (J) = bond energies of reactants - bond energies of products
* Multiply bond energies for each bond by the coefficient
* Hess’s law
* Finding delta H for the overall reaction from knowing delta H for the steps of the reaction
* Flipping the equation flips the sign of delta H
* Multiplying/dividing the equation by a coefficient multiplies/divides delta H by that coefficient
* Adding/subtracting equations adds/subtracts their delta H values
* Enthalpy of solution
* Ionic substances dissolving in water
* 1: Breaking of solute bonds - energy required is equal to the lattice energy (positive delta H since bonds are being broken)
* 2: Separation of solvent molecules - water molecules must spread out to make room for the solute ions (requires energy to weaken the IMF between water molecules - positive delta H)
* 3: Creation new attractions - free floating ions are attracted to the dipoles of water molecules (energy is released - negative delta H)
* Hydration energy = step 2 + step 3 energies
* Coulombic energy - increases with charge magnitude, decreases as size increases
* Enthalpy of solution = step 1 + 2 + 3 energies
* Phase changes
* Solid to gas is sublimation, gas to solid is deposition
* When vapor pressure equals the surrounding atmospheric pressure, the liquid boils - lower atmospheric pressure (high elevation) means a lower boiling point
* Enthalpy of fusion - energy to melt a solid; heat of fusion - heat given off by a substance when it freezes
* Enthalpy of vaporization - energy to turn a liquid into a gas; heat of vaporization - heat given off by a substance condensing
* IMF is stronger for a liquid than a gas, and for a solid than a liquid, and the stronger IMF is more stable, therefore going from a gas to a liquid or a liquid to a solid releases energy (exothermic)
* As heat is added to a substance, the temperature of the substance can increase OR it can change phases, but not both at once
* When a substance is changing phases, the temperature of the substance remains constant
* Calorimetry
* Specific heat - amount of heat required to raise the temperature of one gram of a substance by one degree C/K
* Large specific heat - can absorb much heat without a significant temperature change
* Low specific heat - quickly changes temperature
* Heat added (J or cal) q = mcΔT
* q1 = q2 for mixtures
* Calorimetry - measurement of heat changes during chemical reactions
* Find J from q, find mol from stoich, divide the two to find delta H
* Delta H measured in J/mol
* Heating curves
* For problems where a solid completely melts or the like, add q from mcat to (moles) * (heat of fusion) for the total heat required for process to occur
7. Equilibrium
* Keq
* Reaction is at equilibrium when all concentrations stop changing
* Reaction does not stop - rate of forward and reverse reactions become equal
* All concentrations do NOT sum to initial concentration of reactants
* In reaction 2A -> B, concentration of A will decrease 2x as much as concentration of B increases
* Equilibrium expression/law of mass action
* For the reaction aA + bB -> cC + dD: Keq = ([C]^c * [D]^d) / ([A]^a + [B]^b)
* [A], etc. are molar concentrations/partial pressures at equilibrium
* Products in numerator, reactants in denominator
* Coefficients in balanced equation become exponents in equilibrium expression
* Only gaseous and aqueous species are included in the expression
* Keq has no units
* K>1 favors forward rxn; K<1 favors reverse rxn
* Different equilibrium constants
* Kc for molar concentrations
* Kp for partial pressures
* Ksp is solubility product (no denominator because reactants are solids)
* Ka is acid dissociation constant for weak acids
* Kb is base dissociation constant for weak bases
* Kw describes the ionization of water (Kw = 1*10^-14)
* Manipulating Keq
* Keq for a flipped reaction is the reciprocal of Keq for initial rxn
* Keq for a reaction multiplied by a coefficient is the initial Keq to the power of the coefficient
* Keq for two reactions added together is their respective initial Keq values multiplied together
* Le Chatelier’s principle
* Increasing concentration of reactants shifts rxn to favor products (forward) and vice versa
* Increasing pressure increases partial pressure of all gases in container and shifts rxn to side with fewer gas molecules (moles of gas)
* Increasing volume decreases pressure and vice versa
* Adding a non-reacting gas (noble gas) to a non-rigid container causes the volume to increase while not changing total pressure
* Adding a non-reacting gas to a rigid container would increase the total pressure of the container and not affect the partial pressures of other species - no reaction shift occurs
* Increasing temperature in an endothermic reaction shifts the rxn to favor products (forward); increasing temperature in an exothermic reaction shifts the rxn to favor reactants (reverse)
* Treat “heat” as a reactant (endothermic) or product (exothermic) to see shifts like with concentration change
* Diluting aqueous equilibriums shifts the rxn to favor the side with more aqueous species; removing water (evaporation) shifts the rxn to favor the side with less aqueous species
* Shifts caused by concentration/pressure are temporary shifts and do not change Keq; shifts caused by temperature permanently affects Keq and ratio of products to reactants since it adds/removes energy from the system
* Reaction quotient Q
* Q can be calculated at any point with current concentrations/pressures; Keq can only be calculated with equilibrium values
* For the reaction aA + bB -> cC + dD: Q = ([C]^c * [D]^d) / ([A]^a + [B]^b)
* [A], etc. are initial molar concentrations or partial pressures
* If Q<K, rxn shifts right; if Q>K, rxn shifts left; if Q=K, rxn is at equilibrium
* Solubility
* A salt is considered soluble if more than 1g can be dissolved in 100mL of water
* Soluble salts are assumed to dissociate completely in aqueous solutions
* Most solids become more soluble in a liquid as temp increases
* Solubility product Ksp
* For the reaction AaBb(s) ⇄ aA^b+(aq) + bB^a-(aq): Ksp = [A^b+]^a * [B^a-]^b
* Molar solubility is determined by subbing x, 2x, 3x, etc. in for concentrations in Ksp expression (x if coefficient is 1 in balanced reaction, 2x if coefficient is 2, etc.)
* Molar solubility of a salt is equal to the concentration of any ion that occurs in a 1:1 ration with the salt
* Molar solubility typically increases with temperature since there is more energy available to force water molecules apart to make room for solute ions
* Common ion effect
* Newly added ions from a separate solution affect equilibrium of initial solution if some elements are present in both, even though newly added ions did not come from the initial compound
* ex. Adding NaCl to AgCl affects Cl which affects AgCl equilibrium)
8. Acids and Bases
* pH
* Formulas
* pH = -log([H+])
* pOH = -log([OH-])
* pKa = -log(Ka)
* pKb = -log(Kb)
* pKw = -log(Kw)
* [H+] = [OH-] => neutral, pH = 7
* [H+] > [OH-] => acidic, pH < 7
* [H+] < [OH-] => basic, pH > 7
* Increasing pH means decreasing [H+] (less acidic solution) and vice versa
* Strong acids
* Strong acids dissociate completely in water (rxn goes to completion); no equilibrium, eq constant, or dissociation constant
* Important strong acids/bases
* No tendency for reverse rxn to occur (-> not ⇄) so conjugate base of a strong acid is very weak
* pH of strong acid solution can be found directly from [H+] since it dissociates completely
* Best conductors of electricity
* Weak acids
* Weak acid + water causes a small fraction of its molecules to dissociate into H+ and A- (conjugate base) ions
* Ka and Kb are measures of the strengths of strong/weak acids - equilibrium constants specific to acids/bases
* Acid dissociation constant Ka = [H+]*[A-]/[HA]
* Base dissociation constant Kb = [HB+]*[OH-]/[B]
* Greater Ka means a greater extent of dissociation and a stronger acid
* Greater Kb means a stronger base; base is not dissociating but rather accepting a proton (hydrogen ion) from an acid (protonates/ionizes, not dissociates)
* Set up RICE table w/ values of x for gained/lost concentration to solve for [H+] and pH from Ka or vice versa
* Acid Strength
* Percent dissociation
* The more H+ ions an acid can donate (the easier it is for H+ ions to break free), the stronger the acid is
* Lower concentration -> higher percent dissociation; a greater concentration will lead to more of the conjugate base, making it easier for the reverse rxn to take place -> more HA present in solution and less H3O+ ions (lower percent dissociation)
* Percent ionization: [H3O+]/[HA] * 100
* Acid/base structure
* H is written in front of acids even if H is contained in the conjugate base because that H is attached to a (usually O) atom at the end of the molecule, making it easier for it to detach
* H in a hydroxyl group (-OH) are dissociable due to O being much more electronegative than H
* H bonded to C is almost never dissociable since H and C have similar electronegativity values and share their electrons equally
* Solubility
* Hydroxides dissolve well in solutions with low pH (more H+ ions to react with OH- and speed rxn along)
* Polyprotic acids
* Acids that can give up more than one hydrogen ion (ex. H3PO4)
* More willing to give up first proton than others (after 1st, resulting negative charge attracts remaining protons more strongly)
* H3PO4 is a stronger acid than H2PO4-, HPO42-, etc.
* Amount of each succeeding acid decreases: [H3PO4] > [H2PO4-] > [HPO42-] > [PO43-]
* Kw
* The equilibrium constant of water due to the following reaction: Kw = [H3O+]*[OH-] = [H+]*[OH-] = 1.0*10^-14 at 25 C for any aqueous solution
* pH + pOH = 14
* Kw = 1*10^-14 = Ka*Kb
* pKa + pKb = 14
* Knowing Ka for a weak acid, Kb can be found for its conjugate base
* pH is not limited to a 0-14 scale - very rarely is pH >14 or <0, but it does occur at high concentrations
* Increase in temperature increases Kw (dissociation of water is endothermic) so pKw and pH decrease
* Neutralization reactions
* When an acid and base mix, the acid donates protons to the base in a neutralization rxn
* Strong acid + strong base
* Both substances dissociate completely
* Net ionic is always the creation of water: H+(aq) + OH-(aq) ⇄ H2O(l)
* All other ions are spectator ions
* Strong acid + weak base
* Strong acid (which dissociates completely) will donate a proton to the weak base
* Product is conjugate acid of weak base
* Ex. HCl + NH3: Net ionic is H+(aq) + NH3(aq) ⇄ NH4+(aq)
* Weak acid + strong base
* Strong base will accept protons from weak acid
* Products are conjugate base of weak acid and water
* Ex. HC2H3O2 + NaOH: Net ionic is HC2H3O2(aq) + OH-(aq) ⇄ C2H3O2-(aq) + H2O(l)
* Weak acid + weak base
* Simple proton transfer reaction - acid gives protons to base
* Ex. HC2H3O2 + NH3: Net ionic is HC2H3O2(aq) + NH3(aq) ⇄ C2H3O2-(aq) + NH4+(aq)
* Buffers
* Solution with a very stable pH; acid/base can be added to a buffer solution without greatly affecting pH; gain/loss of water also does not change pH
* Buffers are created by placing large amounts of a weak acid/base into a solution with its conjugate (salt)
* Weak acid and conjugate base remain in solution together without neutralizing each other
* Presence of the conjugate pair makes the buffer effective
* If enough strong acid/base is added that all of the acid or conjugate base is reacted, the buffer breaks
* Higher concentrations of the conjugate pair resist pH change (better buffers) better than lower concentrations
* Henderson-Hasselbalch
* When concentrations of acid and conjugate base in a solution are the same, pH=pKa and pOH=pKb
* Choosing an acid for a buffer solution requires choosing an acid with a pKa close to the desired pH (almost equal amounts of acid and conjugate base; makes buffer flexible in neutralizing both added H+ and OH-)
* Buffers cannot be created from a very strong acid and its conjugate base because the conjugate base will be very weak and will not readily accept protons
* Indicators
* Weak acids which change colors in certain pH ranges due to LeChatelier’s principle
* HIn ⇄ H+ + In-
* Ka = [H+][In-]/[HIn]
* Protonated HIn state must be a different color from deprotonated In- state
* Acidic environment causes excess H+ to drive equilibrium to the left (color of HIn); basic environment causes excess OH- to react with H+ from indicator and drive reaction right (color of In-)
* Color change occurs when [HIn] = [In-]; or pH = pKa
* Choose an indicator whose pKa matches the pH at the titration’s equivalence point
* Titration
* Neutralization reactions are performed by titration, where a base of known concentration is slowly added to an acid or vice versa
* Titration curves
* Midpoint also called half equivalence point occurs when [HA] = [A-] (pH = pKa)
* Equivalence point occurs when just enough base has been added to neutralize all the acid initially present (equimolar)
* HA, A- present before midpoint; A- at midpoint, OH- after midpoint
9. Applications of Thermodynamics
* Entropy
* Measure of randomness or dispersion of the system
* 0 entropy is a solid crystal at 0K (has never been reached experimentally)
* Standard entropies are calculated at 25 C
* Standard entropy change delta S = S for products - S for reactants
* If left side of a reaction has more motion, delta S is negative; if right side has more motion delta S is positive
* Gibbs free energy
* Delta G = Gf of products - Gf of reactants
* Negative delta G is spontaneous (thermodynamically favored); positive delta G is nonspontaneous (thermodynamically unfavored); delta G = 0 means rxn is at equilibrium
* Delta G = delta H - T*delta S (T in K)
* Favorability
* Delta G in phase changes is 0 since at a normal phase transition temp, the substance is equally stable in either phase
* Boiling/melting point can be solved for knowing delta H and S since delta G = 0
* Delta G = -RTlnK (R = 8.314, T in K, Keq)
* The larger the reduction potential on a half reaction, the more likely it is to occur
* Galvanic/voltaic cells
* Favored redox rxn generates a flow of current
* Oxidation at - anode (left); reduction at + cathode (right)
* Electrons flow from anode to cathode (L to R)
* Electrons released from oxidation pass to chamber to be consumed in reduction
* Flow of electrons creates current
* Salt bridge between two cells maintains electrical neutrality
* positive cations from salt bridge solution flow to cathode which is losing positive charge (needs + to balance); negative anions from salt bridge solution flow to anode which is gaining positive charge (needs - to balance)
* Under standard conditions (1M solutions, 1 atm, 25 C), cell voltage is the same as total redox voltage
* Keq >> 1, Q = 1
* Non standard conditions
* If Q=Keq, cell voltage drops to 0; increasing Q decreases cell potential and vice versa
* Overall potential decreases as a reaction progresses (product conc increases, reactant conc decreases)
* Nernst equation: Ecell = E0cell - (RT/nF)*ln(Q) where n is # of electrons transferred (always positive)
* Electrolytic cells
* Outside source of voltage is used to force an unfavored redox reaction to occur
* Occur primarily in aqueous solutions (chemical dissolved in water; ion/water molecule is oxidized/reduced)
* Compare reduction potential of cation with that of water (reduction) to determine which is reduced; compare oxidation potential of anion with that of water (oxidation) to determine which is oxidized
* Then balance the 2 oxidation/reduction half reactions to form one net ionic equation
* value for cell potential from the half reactions should always be negative
* Oxidation at + anode (left); reduction at - cathode (right)
* Signs change from galvanic cell setup
* Electrons flow from anode to cathode (L to R)
* + to - (instead of - to + like in galvanic cells)
* Used for electroplating
* I = q/t
* Moles of electrons = coulombs/ 96500 C/mol
* Moles of metal from moles of electrons (from metal half reaction)
* Moles of metal -> grams
* Voltage and favorability
* Redox reaction is favored if potential is positive
* Delta G = -n*F*E0 (n is positive # of electrons transferred, F is 96500, E0 is standard cell potential in V = J/C)
10. Laboratory Overview
* Weighing hot objects on a scale creates convection currents, making object appear lighter than it truly is
* Not rinsing a buret in a titration leads to it being diluted