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Chemistry
Redox
| Question | Answer |
|---|---|
| oxidation definition | the loss of electrons (sm is oxidized when it's charge increases) -the e- appears on the right side of the half equation because it is giving off e- |
| Ionization VS oxidation | oxidation can be considered a type of ionization Since losing electrons creates a positive ion, oxidation often results in ionization. |
| oxidation number | measured charge carried by the atom. These numbers are based on the fact that elementary substances/pure substances are zero |
| in which cases would oxidation NOT result in the formation of a positive ion | -electrons lost are often immediately taken by another substance in the same rxn . -electrons are just transferred within mcs - |
| oxidation number | a measured charge carried by the atom |
| oxidation number VS ion charge | For monatomic ions, they are always equal. In molecules (covalent compounds), atoms don’t have real charges, but we still assign oxidation numbers. -ion charges only apply to ions, not molecules |
| oxidation number VS ion charge polyatomic ion edition | The sum of oxidation numbers of all atoms = the overall ion charge. But individual atoms do NOT usually carry those real charges. |
| oxidation numbers of different atoms | IN THIS ORDER (some things are more insistent on having a praticular # --> assign them 1st & then solve for others) 1. Elementary substances (pure elements) = 0 (unless they're ions) 2. Group I = +1 3. Group 2 = +2 4. F = -1 5. H = +1 6. O = -2 |
| why does group one have an oxidation number of ____ (you should have that blank memorized) | the extra valance e- in the S orbital is given off Atoms tend to become more stable by getting a full outer shell |
| why does H have an oxidation number of ___ | it only has one proton so it's hard for it to hold onto electrons |
| the sum of all ox numbers in a polyatomic is | the charge of that ion |
| the sum of all ox numbers in a mc (or ionic compound) is | zero |
| WARNING: some atoms have a variety of oxidation numbers (use the ones you know to figure it out) | Cr, Mn, Co, Fe, Cu |
| oxygen's oxidation number of -2 except ... | in peroxides or superoxides peroxides: Contains O–O single bond (O2 in formula) Each oxygen = –1 oxidation number superoxides: compound that contains the O₂⁻ ion (instead of O²⁻ or O₂²⁻). -has 1 extra e- than a neutral O₂ -Each O has an ox # of –½. |
| reduction | -reducing the ox number -the gain of e- -in hald-equations, e- are on the left side |
| redox | reduction and ox must occur simultaneously woth no net change in e- -one reaction can't occur without the other |
| what type of reaction is combustion | it is ALWAYS redox Combustion is basically burning a fuel in oxygen, usually producing CO₂ and H₂O if it’s a hydrocarbon: Carbon loses electrons → oxidation Oxygen gains electrons → reduction |
| what type of reaction is precipitate | it is NEVER rexdox Precipitation reactions are double displacement reactions. They swap ions, form a solid, but do not involve oxidation or reduction. |
| balancing equations for redox reactions | 1. split into 2 1/2 equ (1 ox & 1 red) 2. balance atoms for each equ (neutral solutions: can + water if needed, acidic + H+ and/or H2O, basic + OH- & H20) 3. multiply so # of e- in each equ is the same 4. + up equ & cancle |
| when balancing redox equations, how to check ur answer | 1. look at which atoms are losing and gaining electrons -see if this transaction gives one atom a full sheel, full hald shell, etc (lower energy state) 2. your total charge on each side of the arrow must be equal |
| NOTE for balancing equations | if an atom has a coefficent > 1 and an overall charge, then multiply the ox number by the coeficcient to get the overall px number |
| balancing redox equations for basic solutions (staring by pretending its an acid) | 1. if the # of O is uneven, add H2O to that side 2. add the nessesary amount of H+ ions to the other side 3. balance the e- 4. multiply so the # of e- is the same for both sides 5. add and cancle equations |
| balancing redox equations for basic solutions (finish) | 6. the number of H+ ions is the # of OH ions you add to both sides (1OH + 1H = 1H2O) 7. substract the # of H2O mcs (on the side w/ fewest mcs from both sides) NOTE: OH & H+ are alaways aq |
| oxidizing agent | -the substance going through reduction -it causes oxidation in another substance |
| reducing agent | -the substance going through oxidation -causes reuction in another substance |
| redox agents | -causes sm to happen to somthing else -only REACTANTS can redox agents -find them by looking for a change in ox # for each ATOM |
| why are reducing agents ALWAYS reactants | -electron transfer happens during the reaction — so they must be present before the reaction starts. -once the reaction is finished, the oxidizing agent has already gained electrons and the reducing agent has already lost electrons. |
| NOTE; When balancing redox equations using oxidation numbers, you must multiply the oxidation number change by the coefficient and subscript to find the total change. | Total oxidation change = (change in oxidation number per atom) × (number of atoms in the balanced equation) BUT, when asked for the ox #, do the number for each individual atom |
| Electrochemical cells | batterys (fancy fancy) |
| Voltaic (or Galvanic) cells | batteries that produce electricity -a spontaneous rxn used to produce electrical energy -converts chemical E into electrical E (harnesses the flow of electrons) |
| what is electrical energy | In metals: movement of e- through a conductor. In solutions: the movement of ions. In a battery: a chemical reaction pushes electrons from one terminal to another, creating a current. -the effect of moving charges, often electrons |
| cathode | a part of a battery -Cations migrate towards it and consume e- -reduction occurs here |
| anode | anions migrate towards it and liberate electrons -oxidation occurs here |
| do anions have to touch the anode | In a solution (electrolyte), ions don’t have to physically “stick” to the electrode. They just need to get close enough for electron transfer to occur. |
| in a galvanic cell, which electrode is positively charged | cathode |
| if it is accepting electrons shouldn't cathodes be negitive? (voltaic cells) | the electrons that arrive are immediately used up in reduction. They do NOT sit there and pile up. After accepting the e-, the cation becomes neutral |
| What Is Reduction Potential? | How strongly a substance wants to gain e- (be reduced). If something has a high (positive) reduction potential, it means: Electrons are lower in energy when they are attached to it. Electrons move from: Higher energy → lower energy -measured in volts |
| reduction potential VS electronegativity | Electronegativity measures How strongly an atom pulls electrons within a covalent bond. Reduction potential measures: How strongly a substance wants to gain electrons in a redox reaction. |
| salt bridge exmaples | usually KNO3 or KCl |
| salt bridge | allows flow of current but prefents contact between solutions -it completes the circuit and prevents a buildup of charge in each container |
| how does a salt bridge prevent a buildup of charge | Anode: Oxidation happens → e- leave the metal → metal ions go into solution, leaving extra positive ions in solution. Cathode: Reduction happens → electrons arrive → positive ions in solution are removed, leaving extra negative charge in solution. |
| electrodes | cathodes and anodes -1. metal or conductive solides (choose one that corresponds to the half reactions) 2.if nonmetal, aq, gas, etc, use any inert electrode (doesnt have to correstpond) Pt, Nichrome/Graphite work well |
| Reasons the cell stops working | 1. the anode actant is completely consumed 2. Zn²⁺ [] increases & Cu²⁺ decreases. Eventually, rxn reaches equilib 3. bridge runs out of ions 4. Using an inert electrode & a solid coating blocks electrode & doesn't conduct (bubbles block electrode too) |
| Polarization | any factor that slows down the electrode reaction or reduces the cell voltage |
| voltaic cell example I'm supposed to memorize | Zn(s) + Cu+2(aq) --> Zn+2(aq) + Cu(s) -K2SO4 salt bridge |
| what is happening on the negitive side of the ex I'm supposed to memorize | anode: Zn solution: ZnSO4 Zn (s) --> Zn2+(aq) + 2e- anode gets smaller -SO42- comes from salt bridge -e- move out of solution |
| what is happening on the positive side of the ex im suppposed to memorize | cathode: Cu solution: CuSO4 (SO4 is a spectator ion) Cu+2(aq) + 2e- --> Cu(s) cathode grows largers (Cu builds) K+ comes from slat bridge e- come into solution |
| role of the SO42- in the solutions in the ex I memorized | |
| why do positive ions go to the positive cathode | |
| which electrode shrinks in a galvanic cell | anode -the amount it shrinks is usually (the cathode might create gas instead of a solid) the amount the cathode grows |
| porous partition | -can be used instead of a slat bridge -semi-permeable membrane Porous = has small holes or openings. Partition = acts as a divider or barrier. |
| proper notation for writing chemical cells | chemical (both forms, ion and neutral) || (electrode) chemical -order of the two forms matter, the one written first reacts first (this is the order that they are reacting in) |
| electrolytic Cells (electrolysis) | -the opposite direction of a voltaic battery (recharging) -electrical E is pumped into a system to create a non-spontaneous redox reaction -a power source is used to pump e- into the cathode and remove them from the anode -e- go opporite direction |
| why must the power supply for a electrolytic cell involving Cu and Zn be higher than 1.10 V? | |
| in batteries, ions are more stable than neutral atoms so why do they gain e- and become neutra;? | |
| which elements and compounds are solids or liquids at room temp | liquids: Hg, Br, H2O, Gases: H2, N2, O2, F2, Cl, nobel gases, CO, CO2, |
| in electroylitic cells, which electrode is positive | the anode |
| in both type of cells, where do the electrons flow? | the cathode -voltaic-->the cathode is pos -elextrolytic--> the cathode is neg |
| I DONT UNDERSTAND THE FLOW OF ELECTRONS -sometimes it goes towards positive stuff and sometimes negitive and nothing makes sense -so mental note to return here later and figure that out | |
| products of electrolysis in water solutions (cations) | 1.cations reduce to correstponding metal (usually a tran.) Cu+2(aq) + 2e---> Cu(s) 2. H+ reduces to H2 in SA 2H + (aq)+ 2e- --> H2(g) 3. H2O mcs reduce for Group 1, 2, & Al (they are harder to reduce than H2O) 2H2O+ 2e- -->H2(g) + 2OH-(aq) |
| why does H+reduce to H2 in SA? | |
| why does water reduce first istead of Group 1, 2 in water solutions? | -I and II are so reactive reactive w/ h2O, as soon as they form a solid, they react immediately with water 2e- + 2Na+ --> 2Na(s) 2Na(s) + 2H2O--> 2Na + 2OH- +H2 ------------------------------------------ 2H2O+2e- --> H2+2OH- |
| when will Group1, 2 and Al reduce | they can reduce if they are not in water |
| in what cases does group 1, 2, and Al reduce in water (water reduces FIRST but that means these still reduce) | |
| why doesnt Al reduce in water | |
| products of electrolysis in water solutions (anions) ? | 1. anions oxidized to nometal: 2Cl- --> Cl2(g) + 2e- 2.OH- ions oxidized to O2 in SB 2OH-(aq) -->1/2O2(g) +H2O +2e- 3. H2O mcs oxidized istead of NO3-, SO4-2, F- H2O --> 1/2O2(g) +2H+ + 2e- |
| why are OH- ions oxidized to O2 in SB? | |
| why are H2O mcs oxidized istead of NO3-, SO4-2, | they have very high oxidation numbers N+5, S+6 |
| why are H2O mcs oxidized istead of F- | it is VERY electronegative |
| for the electrolysis of NaOH, why does the [] of NaOH go up | |
| when determining the products of electrolysis, what is the similifyied rule (cathode --> cations) | whichever ion is less reactive (either the H+ from water for the metal cation) will form product -the product will udergo reduction at the cathode and become neutral (remember diatomics for the neutral substance) |
| when determining the products of electrolysis, what is the similifyied rule (anode --> anion ) | oxygen is produced (from the OH- in water) unless there is a group 7 element present |
| voltage def | how hard the electricity is pushed its NOT electrons per coulomb. Voltage (V) = energy per coulomb how much energy each coulomb of charge carried measured in volts (v) |
| charge def | ✔ Measured in coulombs (C) ✔ It depends on how many electrons move Charge (Q) = number of electrons (amount of electricity). More precisely: total quantity of charge carried by those electrons. |
| current def | the rate of flow of charge. It tells you the amount of electric charge that passes a point per second. I = current measured in amperes (A) -speed of electricity |
| electroplating | creation of a metal solid from its dissolved ions -add e- to aq ion -metal solids are usually neutral & aq metals are usually ions or are bonded to ions ofopposite charge |
| coulomb (electrolysis) | A coulomb is the amount of charge that flows when a current of 1 amp passes for 1 second. unit = C |
| how does the number of coulombs tell you how many electrons something has gained or lost. | By counting how many electrons passed, you know how many ions got reduced or oxidized you don’t see the electrons directly, but the total charge (coulombs) tells you how many electrons moved, and that tells you how many atoms gained or lost electrons. |
| formula for electrochemical cells | I = q/t t= time (MUST be in sec) I = current (Amps) q = charge(coulombs) |
| faraday constant (F) | the charge of 1 mol of electrons 1F =96485 C/mol e- |
| Why does electrolysis sometimes produce acid? | That usually happens at the anode (oxidation side). Water breaks apart and produces: Oxygen gas & H⁺ ions If those H⁺ ions stay in the solution, the solution becomes acidic. |
| standard voltages (E naught) | voltage measured when all reactants and products are in their standard states:Concentration: 1 M for solutions, Gas pressure: 1 atm, Temperature: 25 °C (298 K) -the temperature can change and the standard voltage might still remain constant |
| how to calculate the standard voltage of a full cell | add the E of the anode & the E of the cathode -a positive E means the cell is spotaneous (voltaic cell: the electrons naturally want to flow from anode → cathode) -a negitive E means the cell is thermodynamically unfavorable (electrolytic) |
| why must red and ox happen at the same time to assign a half-cell voltage | -cause ox & red occur at the same time, a voltage can be assigned to each 1/2 cell. Can’t measure otherwise because voltage is fundamentally a difference in potential. - there's no voltage without a complete circuit, just preassure that goes nowhere |
| half-cell voltage def | A half-cell voltage measures the tendency of a species to gain or lose electrons — its electron “pressure.. |
| voltage lis ike H2O pressure, not H2O itself. 1 side of 1/2-cell wants to lose e- (ox) → high “Preassure.” Other side wants to gain e- (red.) → low “P” -voltage difference is Pressure the pump creates & e- are “water” that flows because of that P | A flow of e- (current) happens when you connect them voltage exists before e- start flowing cauz it measures potential difference, not actual flow -it measures potential difference, its tendency to give or recieve e- |
| what are all standard voltages relative to | the voltage of hydrogen at standard conditions Ered: 2H+(aq) + 2e- --> H2(g) 0.000V |
| the chart in your book is based on the reduction values. How doyou find the oxidation values for those elements? | reverse/flip the reaction and flip the sign of the E |
| voltages dependent and independent on... | dependent: independent: coefficients (V is not dependent on the # of e- transfered) |
| O2 reduction equation | O2 + 4H+ + 4e- --> 2H2O Ered = 1.229V -the O is becomming O-2 but we write it as forming water |
| if we want a themodynamically ufavorable rxn to happen, what do we do | we must use a power source with V > |Ecell | -absolute value of Ecell we have to pump in more electrical energy than the reaction naturally “wants” to give. The reaction has a certain “resistance” to happening — measured by |Ecell| |
| what does it mean if the E = 0 (not E∘) | MOST LIKELY: rxn is @ equilib (assuming that all the chemicals are still present and have the ability to react) 2. sm is wrong with the circuit |
| determining if certain acids can be stored in metal containers (which acids don't react w/ the container/dissolve the metal) -Note: assume that the acid is 1M cause the charr only works for standard conditions. | write equ. (use chart) for the metal & ions from acid involved (ALWAYS H+) add their voltages, if V < 0, it can be stored if it is positive the the rxn is favorable & it can't be stored -exclude any SI: non-hydrogen mc in the acid |
| what if a WA is placed in a container of a metal that reacts with it? How is this different from a SA in a reactive metal? | the only difference is that the acid's low [H+] causes a slower, less vigorous reaction -the non-hydrogen atom is STILL a spectator ion beucase it does not participate in the rxn (even tho it doesnt fully dissociate) |
| if you are a reducing agent, what is happening to you? | you're getting oxidized |
| how to determine the strength of oxidizing/reducing agents | a "Strong" agent has a very "+" E (the rxn likes to go forward/the substace likes to do whatever the equation says) a weak agents are not as positive |
| NOTE: Fe+2, Cr+3, & Sn+2 can be both oxidized and reduced. There are TWO different equations for each on the chart | |
| Are the alkaline metals strong reducing agents or oxidizing agents | VERTY STRONG reducing agents: they like to be oxidized and lose their extra electron in their S orbitals |
| is F2 a strong reducing agents or oxidizing agents? | one of the strongest oxidizing agents -High electronegativity very high electron affinity, meaning that they release a lot of energy when they gain an electron to form F -F–F bond in F₂ is relatively weak & can easily split into two fluorine atom |
| how to calculate E under any conditions (not standard coditions) | nernst equ: E=E∘− (RT)/(∣n∣F) x lnQ R=8.314J/mol*K T=kelvin or C n=moles of e- transfered (from the stoiciometry of the two balanced equations-->rxns must have same mol of e-) F=96485J/molV Q=rxn quotient for overall equ |
| nernst equation shortcut | assuming the T is 25C, E= E∘− (0.0257V)/(n) x lnQ |
| why does [] affect voltage | -V depends on how far sm is from equilib Far → big push → high v Close → low v equilib → v=0 Remove reactants = fewer particles react/collide & closer to equilib Remove reactants or + products = bigger Q (system closer to equilib) |
| why lowering the [reactants] lowers the voltage (galvaic)or slow the rxn/decrease resistance (electrolytic) | If you decrease [A] or [B], the denominator of Q gets smaller. That makes Q larger (because dividing by a smaller number → bigger Q). Subtracting a bigger number → E decreases |
| what factors affect voltage(galvaic) | electrodes/materials Electrolyte concentration (more stuff to react =more rxns) Temp Internal resistance Surface area of electrodes thickness of wire Age of battery |
| how temp affects voltage (galvanic cell) or the speed of the rxn (electrolytic) | Hot temperatures:Reactions inside the battery go faster. Cold temperatures:Reactions slow down. |
| what does a really high K mean | Very high K → usually spontaneous under standard conditions. |
| how delta G realates to voltage formula | deltaG = -nFE∘ detlaG=joules |
| how does the number of electrons change spotaneity of a reaction | no effect on spontaneity since # of electrons doesn't change sign of deltaG∘ (just a different amount of E) -how hard the electricity is being pushed doesn't change but more e- being pushed = more energy |
| calories and joules conversion | 1 cal = 4.184 joules |
| convert between cal and Kcal | cal/1000 = kcal |
| convert kj to j | To convert kJ → J, multiply by 1000. |
| how do we know when to use H2 or H+ | use H+ when balancing a redox equation in acidic solution -use H2 what hydrogen is actually being reduced |
| convert F to charge | 1 f =96485 C/mole e- |
Created by:
Blessed Rectangle 10