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Chem Test
acids/bases
| Question | Answer |
|---|---|
| arrhenius def | acid contains H+ ions and bases contain OH- ions -acid is a substance that increases the concentration of hydrogen ions in an aqueous solution (produces H+ when dissolved) (OH- for bases) -ONLY WORKS FOR H2O SOLUTIONS |
| bronstead def | acids donate protons (H+) and bases accept protons acid is a substance that can donate a proton to another substance base is a substance that can accept a proton |
| the color of phenolphthalein in an acid or base | acid - colorless neutral- colorless base - pink |
| he color of bromthymol blue in an acid or base | acid - yellow neutral - green base - blue |
| acid properties | sour taste corrosive to metals (react with metals to forming H2 gas, process damages metal) react with indicators electrolytes react with bases |
| base properties | bitter taste caustic (react with skin, ahir, protein ect.) slippery react with indicators electrolytes react with acids |
| strong acids list | HCL - hydrochloric HBr - hydrobromic HI - hydriodic HNO3 - nitric HClO4 - perchloric, and HClO3, H2SO4 - sulfuric |
| strong bases list | group I and Sr, Ba, Ca hydroxides Ca(OH)2 Sr(OH)2 Ba(OH)2 |
| NH4 + | amonium weak acid |
| NH3 | ammonia weak base |
| formula for pH and pOH | pH = -log [H+] pOH- = -log [OH-] M of a base is the [OH-] and M of an acid is the [H+] -when we use {H} or {OH}, we are assuming that the {H} = the initil acid concentration (AKA this only works for strong acids/bases cause they fully dissociate) |
| how to convert [H+ ] into [OH] and vise versa | Kw = [H+] [OH-} Kw = 1 x 10 ^-14 (no for sig figs) [H+] vs [OH-], whichever is bigger determined whether its a base , acid, or neutral substance if there is more than 1 H+ or OH in a compound (Ca(OH-)2) then multiply the [] by the amound of H or OH |
| relationship between pOH and pH | pOH + pH = 14 |
| To convert pH into hydrogen ion concentration | [H] = 10^-pH |
| how to write net neutralizationequations for acid/base rxns | -all aq -always a single arrow -leave out spectator ions (SA & SB: ion that is not H or OH is a spectator ion) -for ionic compounds, anything but H or OH is a sp. ion -if you can't make HOH then just stick them together |
| after the equivalence point of a titration is reached, how do you know if something is acidic or basic | strong acid + strong base = neutral weak acid + strong base = weak base strong acid + weak base = weak acid weak acid + weak base = unpredicable (dependent on Ka vs Kb) |
| what is the formula for a hydronium ion | H3O+ |
| is naturally occuring rainwater acidic, basic, or neautral | -normal is slightly acidic because of CO2 in the atmosphere dissolving in water to form H2CO2 acid if SULFUR and NITROGEN MONOXIDE form, reactions occur producing harmful levels of acidity |
| in a titration, what do you do | you add base to an acid until it changes colors -the strength of the acid/base does NOT determine a titration, only the moles of acid VS base |
| amphoteric | the substance can act as an acid or a base - it becomes the opposite of what you put it with EX: water -when by itself, it might want to be one thing, but mixing it with an acid or base forces it the other way |
| alkaline | descibes soltuions with a pH greater than 7 |
| how to calculate the pH when the Ka and initial M is provided WRITE THESE STEPS IN YOUR WORK | 1. disossiation equa. 2. expression 3. ICE chart and either OH- or H+ is the variable (X) 4. in E row for the reactant "assume X is small" (WRITE THIS) 5. plug into expression & solve 6. find pH or pOH use quadratic equation if u dont assume x is small |
| new sig fig rule for logarithms | add one sig fig when taking the log (applies to natural and negitive logs too) drop one sig fig when taking the inverse log |
| monoprotic | the acid only gives off 1 H+ -assume eveything is this unless given a reason to think otherwise |
| diprotic | the acid gives off 2 H+ |
| triprotic | the acid gives off 3 H+ |
| polyprotic | acid can give off more than 1 H -diprotic might give off one H in certain situation, it doesn't ALAWYS give off 2 it just has the ability to -each equation has a differ. K (they get smaller each time) - |
| strong acids def | -ionize completely into ions -strong electrolytes (substances that dissociate or ionize into positive and negative ions. These mobile ions carry electric charge through the solution) -usually polar because polarity helps them form ions in water. |
| when do strong acids or bases NOT ionize completely | when there isn't enough H2O in the solution to ionize every molecule |
| what makes something a strong acid/why do they dissociate completely | Hydrogen is bonded to a very electronegative atom That atom pulls electron density away from H The H–X bond becomes very polarized (the bond is easier to break) their conjugate bases are stable. |
| weak acids (description) | the HX is is alawys greater in conc. than H+ -the original [] of H bonded is alawys greater than the [] of H+ given off |
| types of weak acids (1) | molecular weak acids (including organic acids and indicatiors (often written has HIn) -NH4+ -covalently bonded mcs that give off an H+ |
| types of weak acids (2) | anions containing ionizable H atoms H3PO4 <--> H+ + H2PO4-1 |
| types of weak acids (3) | cations (except Group 1, Ca+2, Cr +2, Ba+2) become hydrated. +1 ions get 2 H20 , +2 ions get 4, +3 ion get 6 1. show the dissociation eq of the compound 2. cation reacting with water to form a complex ion 3. complex ion giving off one OH |
| types of weak bases (1) | 1. Mg, Al, and transition hydroxides 2. molecular (produces OH in water, includes NH3) (include amines --> type of organic mcs) 3. Anions (just an anion with nothing attached to it) (anions derived from WA is itself a WB) |
| lewis definition | acids accept electron pairs, bases donate them (you have to draw the lewis structures for this method) |
| why will CuCl2 corrode the metal balance tops | -the Cu+2 will absorb the H2O from the air and form the complex ion Cu(H2O)6. That will react with the metal top |
| Ka | the ionization constant or acid dissociation constant -if the rxn starts w/ a weak acid use Ka, vice vera for weak base |
| why don't strongs have Ka or Kb | strongs always go forward (they dissociate completely) -if they had one, it would be huge, since the Ka is based on how easily it dissociates (bigger Ka = more dissociation) |
| why does the Ka get smaller as an acid loses H | After losing H⁺: the species becomes more negatively charged Pulling another positive H⁺ off a negative ion is harder |
| the larger the Ka value... | the stronger the WA -the rxn goes forward (with a small Ka the rxn reverses) |
| relationship of Ka and Kb (of conjugates) | Ka (Kb) = 1 x 10^-14 -the constant only applies to a system @ 25 degrees C -the value changes with T -conjugates differ only by 1 proton (H+) |
| how to find the K overall of a polyprotic acid | 1. write the equ for each H 2. write the K expressions for each 3. multiply and cross cancle the K expressions 4. this should equal the overall K equation 5. write a K expression for the overall equ |
| K overall (polyprotic acids have multiple equations) equals | Ka1 (Ka2) (Ka3) |
| most common buffer | Ch3COOH and CH3COO -1 |
| 5% rule | -used only for weak acid or weak base dissociation -if less than 5% of acid/base conc. dissociate, then the intial conc. of acids/Base is about the same at equilibrium -assume x is small if [H+]/initial conc. x 100 < 5% |
| why should there always be a strong in a net neutralization equation | -the strong pushed the rxn forward -a weak + weak could move backwards |
| if you do not assume x is small, you can either use the quadratic formula or successive approximation | replace the bottom x in expression (the one we assumed was small) with the final [] calculated by assuming x is small -replace the new x value you found into the original equation, and repeat this cycle until you get the same x -value multiple times |
| if, when asked to write a net ionic equation, they say that the M and V are equal, then | the moles of acid = moles of base you can safely assume 1 mol of each to keep it simple |
| salts defintion | something ionic -a cation + anion -can be acidic, basic, or neutral |
| how to determine if a salt is acidic, neutral, or basic | classify each ion separately, then add the results if there is a WA and WB combo, then compare the Ka to the Kb Ka > Kb = acidic kb > Ka = basic -a SA + SB = neutral (neutral ions don't have neutralization equations) |
| NBA chart: neutral anions (anything neutral is a spectator ion) | Cl-, Br-, I-, NO3-, ClO4- -all the anions from the strong acids (except SO4) |
| NBA chart: neutral cations | Li+, Na, K, Rb, Cs, Ca, Ba, Sr (group 1 + SrBaCa) -excpet francium cause they only appear during radioactive decay and don't last long |
| NBA chart: basic anions | HCO3- HS F- CO3-2 S-2 PO4-3 CN- NO2- HPO4 SO4-2 CH3COO |
| NBA chart: basic cations | none -positive ions can't consume a positive ion (bases accept H+) -they can't produce OH- on their own -NOTE: amines are usually basic |
| NBA chart: acidic anions | HSO4 H2PO4 |
| NBA chart: acidic cations | Mg+2 Al+3 NH4+ transitions |
| when asked to write an equation if a salt if acidic or basic, | write equa showing that it is basic/acidic 1. write dissociation equ, then a separate eq to show the basic/acidic ion reacting to either accept or give off H+ (add water if needed) -might have to form complex ion then show that giving off H (3 equ) |
| good systems to memorize (acting as acids or bases) | NH4+/ NH3 HF/F- H3PO4/H2PO4-/HPO4-/PO4-3 CH3COOh/CH3Coo- |
| when finding the pH of solution mixtures, and you are mixing a SA and WA, then | the pH is due only to the SA -treat the WA like water (it is so weak compared to the strong acid you are just diluting it) -use only the strong H+ moles and divide by the total volume |
| normality | -accounts for the [] of polyprotic acids and bases -equivalence/liter -M x Equ = N |
| equivalence | # of H+ or OH- given off per mole of substance - |
| for polyprotic acids., although each Ka is different, | the volume of base needed to reach the equivalence point is the same for each |
| eqivalence point | the point in a titration where 1 H+ had been completely removed from the acid (there can me multiple points) -when one H+ has been neutralized with one OH- |
| endpoint | the point in a titration where every H+ has been removed from the acid -where all H+ have been neutralized with OH- -It does NOT always mean the moles of acid = moles of base, unless the reaction is 1:1. |
| how to find the normality or volume of an acid/base | Na (Va) = Nb (Vb) |
| titration curves | -endpoint is 1/2 way up the vertical jump -buffer is 1/2 way through the titration (titration starts w/ the first drop of base & ends @ endpoint) |
| what happens at the half way point of a titration | the pH of the solution = pKa of the weak acid being titrated |
| titration curve for diprotic acids | the distance (ml of base) to the 1st equivalence point is equal to the distance to the second equivalence point (AKA endpoint) |
| protonated state | the form of a molecule (or ion) after it has gained a proton (H⁺) In other words, the molecule has accepted a proton, which usually gives it a positive charge or makes it more acidic |
| deprotonated state | the state of a molecule after it has lost a proton (H⁺). So instead of gaining H⁺, it donates one—often giving the molecule a negative charge or making it more basic. |
| relationship between Ka and Kb (not the equation) | When Ka, is high, it means the acid is strong and it donates protons easily. Consequently, the conjugate base has a much lower tendency to accept protons, which means its Kb is lower. It’s basically an inverse relationship between Ka and Kb. |
| if the weak acid and conjugate base versions of an indicator are equal, then | the pH that the titration ends at is = the pKa of the indicator -we select indicators based on their pH at the endpoint |
| what changes in pH are visible to the human eye | can only detect about a one-unit shift. So, for example, if the pH goes from 7 to 8, that's noticeable, but a change from, say, 7.1 to 7.2 might not really be perceived visually |
| -SA & SB titration curve | - has largest vertical jump cause once all SA is neutralized, you're adding base to H2O -any indicator with a Ka within this large jump will work, so any indicator anywhere near the endpoint will work |
| what is present in significant amounts if strong base is added to a buffer (the buffer is basic) yep, a buffer can be either acidic or basic. | -SB it's not reacting, you are just diluting the strong base |
| calculate the Ph before any indicator is added | -write the dissociation equation for the acid and use an ice chart to solve for [H] |
| find the pH after a random amount of base is added | -subtract the moles of base from the moles of acid -the difference is the excess acid -the moles of OH is the amount of weak base produced -divide both by the total volume -plug into the K expression for the dissociation of the weak acid |
| calculate pH at the endpoint | at the endpoint of a weak acid and strong base titration, you’ll have more of the conjugate base than acid, so change the Ka into Kb -also flip the net equation for the titration so that the base is first |
| when adding an acid or base at the endpoint of a titration... | we only deal with the extra stuff added -the solution is neutral at the endpoint so its basically water |
| Do we ruin buffers at the endpoint? | At the endpoint of a titration, the buffer capacity is essentially lost because the acid is fully neutralized, leaving mostly the conjugate base. |
| buffer def | a solution that resists changes in pH when small amounts of acid or base are added -it is made by combining a weak acid and conjugate base -where the pH is not changing |
| first way to make a buffer | you can mix both a WA and a CB together |
| second way to make a buffer | 'they are made through a reaction 'you can start with one component (WA or CB) and react some of it to it's conjugate -what you add must be a lower quanity than what you started with (too much will neutralize all of the other) |
| because a weak acid dissociates in water, why isn't it always a buffer | -it does not make a significant amount of CB -it dissociates less than 5% -a buffer requires a significant amount of each component |
| how to choose chemicals to make a buffer at a desired pH | if the concentration of the weak acid = the concentration of conjugate base: pH of the solution = pKa of the weak acid -once you have chosen your acids and bases, start with an equal amount of both and adjust |
| The pH of a buffer solution is determined by | the pKa of the weak acid and the ratio of the concentrations of the conjugate base and the weak acid. T |
| note: if a base or acid is solid, then it doesn't affect the volume much. Just use the volume of the liquid | |
| adding a strong acid to a buffer | -the acid reacts with/consumes the conjugate base -(a base would react with/consume the weak acid |
| Why does adding a strong acid to a buffer not affect the pH very much? (same for adding a SB) | the added hydrogen ions are neutralized by the conjugate base in the buffer. This forms the weak acid, effectively using up the added H⁺ ions while leaving some CB left. The ratio of the conjugate base to the weak acid doesn’t shift dramatically |
| how is pH affected by pouring two acids together | they dilute each other. The pH ends up between the acidities of the 2 separate acids |
| The total moles of WA (weak acid) and CB (conjugate base) in a buffer is essentially | constant (as long as you’re only adding small amounts of acid or base). -only the ratio of CB to WA changes The total moles stay the same because you’re just converting one form into the other, not removing them from the system. |
| how to tell if it's the half-way neutralization point | the moles of strong base added = 1/2 the initial moles of weak acid this means that the moles of CB and WA are equal |
| how to find the K for the neutralization equation of a strong base and strong acid | this equation is the reverse for the formation equation of water (the constant for that is 1 x 10^-14) -therefore, reverse the constant (1/1 x 10^-14 OR 1 x 10^14) |
| at the equivalence point of a SA and SB titration, | the cation and anion from the acid and the base come together to form a neutral compound (the ions that are not H or OH) -also the pH is 7 (a SA and WB combo would have a pH of below 7 at the equivalence point) |
| when, during a titration, does the pH increase or decrease the most | right after the equivalence point |
| if the neutralization equation of an acid and base is the same as an acid dissociation equation for, lets say, NH4, then to find the K of the neutralization rxn | divide 1 by the Ka of NH4 |
| amphiprotic | A substance that can both donate and accept a proton (H⁺). Key point: This is a subset of amphoteric substances: all amphiprotic substances are amphoteric, but not all amphoteric substances are amphiprotic. |
| what is the concentration of a [H] and [OH] in a neutral solution is | 1 x 10-7 -water dissociated into H and OH in a 1 to 1 ratio -a neutral aq aolution has an equal M of H and OH |
| relationship between [H] and pH values | as the pH increases by one unit, the [h+] decreases by a power of 10 -the power of the [H] is the opposite of the pH 10^-10 = 10 |
| as the [H+] of the acid increases, explain the affect on the strength of the acid | none -the strength of an acid does not depend on its molarity |
| percent ionization | (how much of an acid or base actually ionizes in water) |
| as the [H+] of the acid increases, explain the affect on the percent ionization of the acid | even though we have more hydrogen ions, the fraction of acid mcs that accutally ionize becomes smaller -a smaller % of the acid is acctually dissociating -adding H or subtracting water shifts equilib towards the ions |
| if an acid is diluted, the Ka | stays the same -depends only on temperature, not starting concentrations -they all dilute at the same rate |
| how to determine how an amphoteric compound will behave in a soltuion | compare the Ka to the Kb Ka > Kb = acidic kb > Ka = basic |
| what happens when you add water to a buffer | it does not change the pH because it dilutes both parts of the buffer, so the ratio stays the same |
| are indicators basic, neutral, or acidic | WA -if drop it into a base, it will grab the proton off -if drop it in an acid, it will accept a proton OR stay as is -it stays as is in a neutral solution OR both forms (protanated and deprotanated exsist at once) -most indicators are amphoteric |
| how to choose an indcator | the pKa of the indicator should fall within the vertical jump on the titration curve 'at the middle of an indicator's color change (equal amounts of both forms), pH = pKa |
Created by:
Blessed Rectangle 10