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Chem test 5
solutions
| Question | Answer |
|---|---|
| dissociation equations for acids and bases dissolved in water | -weak (<---->), strong (--->) -bases give off OH - and acids give off H+ formula --> OH or H+ and compound formed using remaining elements -all products are aq -more than 1 OH or H --> take all |
| strong acids | HCL - hydrochloric HBr - hydrobromic HI - hydriodic HNO3 - nitric HClO4 - perchloric H2SO4 - sulfuric -both weak and strong acids are almost alaways aq |
| strong bases | group I and Sr, Ba, Ca hydroxides Ca(OH)2 Sr(OH)2 Ba(OH)2 -both weak and strong bases are almost alawys aq |
| strong acids and bases VS weak acids and bases | strong- dissociate completely to form ions weak - dissociate partially into ions |
| when solving for dissociation equations... If every compound in the equation is aqueous, then why do you only put coefficients on the product side. | you’re describing ions in solution, not molecules reacting (the reactant doesn't exsist anymore, it was dissolved in solvent) The reactant formula doesn’t represent a real “whole” particle anymore — it’s just showing what used to form those ions. |
| review polyatomics | |
| NH3 | ammonia weak base |
| NH4 + | amonium weak acid |
| dissociation equations for ionic compounds | -split the molecule into indivisual atoms & ions (polyatomics stay together) -keep the reactant a 1 and then balance the equation -write the charges of the reactants -reactants are alawys (aq) |
| dissociation equations for nutral (not acidic or basic) covalent compounds | write the same exact formula of the product as the reactant except change the sign to (aq) -single arrow |
| dissociation | “Dissociation” just means breaking apart into ions in water. But how that happens depends on what kind of substance you have. ionic/ electrolyte --> separate into indivisual ions when dissolved in a solvent covalent --> react with water to form ions |
| how to know when a covalent mc is neutral, a weak base, or an acid | -if a mc has an H that can leave as H⁺ (proton), it’s an acid. -Weak bases don’t have OH⁻ yet, but they have a lone pair of e-. That lone pair can grab H⁺ from water, forming OH⁻. -if a mc has neither it is neutral |
| why are acids and bases only covalent molecules (EXCEPT STRONG BASES) | Covalent comp. (nonmetals sharing e-) have neutral atoms— there are no free ions. To make ions (H⁺, OH⁻), cov. must react w/ H2O — ionic compounds already are ions & just dissociate -However, strong bases incude OH-, making them ionic |
| why can water participate in both acid and base reactions | water itself is amphoteric (it can act as either an acid or a base): Water molecules can: Accept an H⁺ → becoming H₃O⁺ (hydronium) Donate an H⁺ → leaving behind OH⁻ (hydroxide) |
| how acids react in water to produce ions | An acid donates a proton (H⁺) to water. That’s what makes it an acid: it increases the concentration of H₃O⁺ in solution. --> they produce H+ |
| why is the dissociation of weak acids/bases in water equalibrium | |
| how weak bases react in water to produce ions | Weak bases don't already have an OH- in their structure (covalent), so they accept a proton (H⁺) from water, which increases OH⁻ concentration. --> they react with water to produce OH- formula + H2O --> OH- + ___ |
| how strong bases (ionic) react in water | Already contains OH⁻, just separates formula --> ___ + OH- |
| weak acids and strong acids | the only difference is the arrow -both just dissociate into H+ and the other ion formula ----> H+ + ____ |
| what is molarity | molesof solute per one liter solvent |
| colligative properties def | property of a solution that depends only on the number of solute particles per number of solvent, not their identity -not dependent on molarity becuause the number of particles is indepenent from the volume |
| molality | doesn’t change if the solution expands or contracts with temperature. |
| colligatice properties | vapor preassure (of a liquid) (how easily something varporizes) goes down (doen't vaporize as easily) AKA vapor preassure depression boiling point (of a liquid) goes up AKA bp elevation freezing point (of a liquid) goes down AKA fp depression |
| solute | |
| solubility | can mean two things: 1. whether something will dissolve 2. how much of a solute will dissolve in a solvent |
| solvent | |
| solution | |
| saturated | |
| dilutions equation | m1v1 = m2v2 (similar to titration equation --> mava = mamb) m = molarity of solution v = volume of solution |
| how to know if compounds are solids, liquids, or gases | -acids & bases are aq -use ur brain (combustion reactions make water vapor not rain) -small molecules --> assume they are gases -medium mcs --> assume they are liquids large mcs --> assume they are solids |
| combinations of deltaH and deltaS that will never work | negitive deltaS poitive deltaH |
| combinations of deltaH and deltaS that will always work | potivie deltaS negitive deltaH |
| which combination of deltaH and deltaS will be more favorable when the temperature is increased | positive and posiitive |
| which combination of deltaH and deltaS will be less favorable when the temperature is increased | negitive and negitive |
| mole fraction (X) | -no units OR mol solute/mol total moles of A/ total moles in solution |
| molarity | mol solute/ L of solution mole/L |
| finding the concentration: percent | if it doesn't tell you the units, assume it's by mass (g) -could use volume or pressure -part/whole x 100 |
| molality | -modification of molarity -units: m solute OR mol/kg OR mol * kg^-1 moles of solute/kg of solvent |
| density (ρ) | total grams of solution/mL |
| density of water | 1 g/mlL= 1g/ 1 mL |
| parts per million | part/whole x 1,000,000 UNITS: ppm used for chem. w/ VERY low [] -unless said otherwise, assume by mass (could be V or P) |
| parts per billion | part/whole x 1,000,000,000 UNITS: ppb used for chem. w/ VERY low [] -unless said otherwise, assume by mass (could be V or P) |
| parts per trillion | part/whole x 10^12 UNITS: ppt used for chem. w/ VERY low [] -unless said otherwise, assume by mass (could be V or P) |
| Cm^3 is exactly equal to which other unit | mL |
| each conventration value equals two different values, it's not just one number | 15 m = 15 mol solute per 1 kg (100g) solvent density of 15 = 15g solution per 1 cm^3 or mL of solution -basically, take the value given and assums and put it on the top, then assume the bottom is one (since every concentration equation is a fraction) |
| to convert between M and m... | you need density |
| isotonic | a solution having the same osmotic pressure as some other solution |
| osmotic pressure | the pressure needed to prevent osmosis how much pressure you’d need to apply to stop that movement |
| osmosis | movement of solvent particles (usually H2O) through a semipermeable membrane (allows water to pass but not solute particles) from an area of lower solute [] to an area of higher solute [], until reaches equilibrium |
| how is osmosis equalibrium | At first, there’s a net flow of water in one direction (toward the more concentrated side).Equilibrium in osmosis means that the rates of water movement in both directions become equal. This means the concentrations (or osmotic pressures) are balanced. |
| given a percent of a solution (either by mass, volume, or pressure) , what do you do | assume 100g/mL/atm of the solution take off the percent sign to find the grams/pressure/volume of solute subtract the grams of solute from 100 to find grams/v/p of solvent (this works cause [] is an intensive property) |
| how to convert between ppm and percent | divide the ppm by one million, then mulitply by 100 to get the percent |
| finding the molarity of a solute, ions, or specific ions that the solute breaks into | in molarity equation, whatever your moles is based on is the M you get. Write the dissociation equation of the solute & use stoiciometry. -finding ion [] : [] of each ion combined -each ion/ solute mc might have a diff. coefficient and different [] |
| if termpature affects volume, which concentrations are temperature dependent? | -all the ones with volume in the formula -molarity and density -if it doesn't have volume in it then it's constant. that's why molality is useful |
| if solutions become more dilute, what happens to the difference between M and m? | -difference would decrease -at higher values, there's a big difference |
| why does the difference between m and M become smaller was the solution becomes more diluted? | -both have the same value in numerator -concentrated: solute contributes lots to total V of solution diluted solutions (small mol of solute), volume of solution becomes ≈ v of solvent -water’s density ≈ 1.00 g/mL, so L of solution ≈ kg of solvent. |
| given the original conentration, what do you assume/start with If only you know the [] but not the total v or mass ,👉 then choose an amount of solution that makes the math easier | mass %--> 100g solution M --> 1.00 L solution m ---> 1000 g solvent moles fraction ---> 1 mol (solute + solvent) (assume the given value is mol of solute, then assume moles of solution is 1 mole ) -THESE HAVE INFINITE SIG FIGS |
| how to do sig figs for these long concentration calulations with assumed values (when only multiplying and dividing) | the result must have the same number of sig figs as the measued value with the fewest sig figs look at the given value with the lowest sig figs |
| do ionic compounds have LDF | Yes, However, attractions between full positive and negative charges are so powerful that they dominate, making the weaker, temporary dipoles of the London dispersion forces insignificant in comparison. LDF occur between ALL molecules/formula units |
| LDF VS dipole-induced dipole interaction | DIDI: A permanent dipole (polar mc) in one molecule induces a temporary dipole in a neighboring nonpolar molecule. classic LDF: between a nonpolar and nonpolar LDF: temporary dipoles between any molcules: more general term that inclues classic and DIDI |
| do polar molecules have temporary dipoles | yes, Even a polar molecule’s electrons move around |
| how do things dissolve | solvent and solute must interact strongly enough to overcome their own intermolecular attractions. (Stronger attractions between solute and solvent → better solubility) |
| how to know if a chemical is polar just by looking at its formula (assume its a larger organic compound) | Hydrocarbons (C and H only) → mostly nonpolar Organic molecules with electronegative atoms (O, N, halogens) → polar -MOST ORGANIC MCS ARE INSOLUABLE IN WATER CAUSE MOST ARE NP REVIEW: -lewis structures, organic chem |
| how to know how soluable a chemical is in water | 1. polar molecules dissolve better in polar H2O than nonpolar mcs (applies to all sovlents) 2. mcs w/ HB are more soluable in H2O (only for solvents w/ HB) 3. bigger mcs are less soluable in H2O (apply to polar solvents) |
| why does HB making something more soluable in water | H2O is polar (O is δ⁻, H is δ⁺) → water mcs form HB w/ other H2O mcs If solute has HB, H or lone pairs from solute form HB w/ water. This “sticking together” of water and solute molecules allows the solute to mix in water |
| why are bigger molecules less soluable in water | Bigger mcs have long chains of C–H bonds (np cause the difference in electroneg is small), which don’t interact w/ H2O Even if a big mc has a polar group, the large np part dominates bigger = more polar = dissolved in polar solvents better |
| why does HB make a chemical soluable in water but not LDF or DD | To dissolve a solute, H2O breaks some of its HB & forms new interactions w/ the solute. Only solutes w/ of H-bonding form strong enough interactions to replace H-bonds in H2O DD or LDF are too weak to “replace” H2O's internal HB→ poor solubility. |
| types of compounds and their soluability in water | ionic : high --> it has a + & - ions that react w/ H2O covalent network: insoluable: Strong 3D covalent bonds can’t be broken by H2O covalent molecular: they follows the rules from another card metallic: Generally insoluble Metalloids: insoluable |
| is hydrgoen a metal or nonmetal | nonmetal |
| amphipathic | a molecule that has both a hydrophilic (water-loving --> polar) part and a hydrophobic (water-fearing --> np) part. |
| how to identify ampholpathic substances | Does it have a polar “head”? Does it have a nonpolar “tail”? Look at one end of the molecule → check if it’s polar/hydrophilic |
| explain why adding solute to a solvent lower the vp of the solvent | the solvent has a more difficult time escaping, for two reasons, (1) solute particles block the solvent, (2) attractive forces between the solute and solvent (called "Ion Dipole") are stronger than the IMF's the solvent normally has to overcome |
| based on things dissolved in water, is electriacl conductivity a colligative property | the answer depends on how you approach the question, if we base it on things dissolving in water then NO, it is not colligative because it does depend of the TYPE of particle as to whether a solution will conduct electricity |
| when you dilute something, what happens to molarity | M must do down |
| formula for dilutions | Mcon x Lcon = Mdil x Ldil - M times L = mol, so equation is also mol of solute before diluting = mol of solute after Mcon = M of conetrated solu Mdil = M of diluted solu Lcon = V of concentrated solu Ldil = V of diluted solu. after adding H2O |
| when starting with a solution, you are asked how much of this solution you need to make a higher amount of a lower molarity solution. Then find the amount of water needed to silute the solution. | use Mcon x Lcon = Mdil x Ldil to solve for the Lcon (the amount of solution you started with) Water added=Vdil−Vcon |
| What dissolves well in water and what doesnt | sugars (VERY POLAR) and alochols are polar --> dissolve well most organics don't dissolve in water |
| how to affect how much of something dissolves | If something is endo, then increasing the I will allow it to dissolve faster --> solids absorb E to dissolve into a liquid dissolving a gas in a liquid is exo, you need to calm/slow the gas down, cooling it would help it dissolve faster |
| is dissolving something in water a reaction | no a process |
| why does polar O2 dissolve in polar H2O? | because H2O is polar, it has a strong electric feild. THe electric firld of H2O temporarily distorts O2's electron cloud, induscing a tind dipole in O2 (called a dipole induced dipole reaction) |
| do dipole induced dipole reactions happen between every polar and nonpolar mc | yes, but the strength of this interaction depends on: How polar the polar mc is How easily the np mc's electrons can be distorted (polarizability) Larger atoms/molecules → more polarizable → stronger interaction Distance between the molecules |
| effect on pressure on solubility | only works on gases increase in P --> increase of solubility preassure helps push the gas into the solution |
| why does squeezing the air out of a 2L soda bottle make the soda lose carbonation faster | CO2 gas in the bottle & CO2 disosolved in the soda are at equalibrium. Less CO₂ gas above the liquid → less pressure Lower pressure shifts equilibrium CO₂ comes out of the soda to replace the missing gas The soda loses carbonation |
| Henry's law | Cg = KPg k =M/L*atm -there's a direct relationship between P & the [] of the gas -seomtimes CgK x Pg (has different units of k tho) k = constant Cg = M of dissolved gas Pg = partial P of the gas |
| explain the constant's role in henry's law | as long as the T, type of solute, and type of solvent are constant, k will be the same it is NOT dependent on preassure or concentration it is specific to each solvent, gas, and T combination |
| solubility graph | -below the line is an unstaurated solution, on it is saturated, above it supersaturated (assume these don't exsist unless told otherwise) -show how much solute will dissolve at a specific temperature per 100 ml of water |
| reading the curve on a solubility graph | the steeper the curve, the more senitive something is to the T of the water -if the curve/solubility goes up as temperature increases --> it is as solid -if the curve goes down as T increases, it is a gas |
| using a solubility curve... 1. find how many grams of water is needed to dissolve __ grams of solute 2. find how much solute solidifies out of the solution at a given lower termperature | use chart to find a proportion: g H2O/g solute at the given T convert g of solute to g of water convert grams of h2O found ealier to g of solute at the new temp. subtract the new g of solute from the given g from ealier to find how much solidified |
| dissolution occurs in three steps | 1. breaking solute-solute attractions (endo) 2. breaking solvent-solvent attractions (endo) (often IMF in water) 3. forming solute-solvent attractions (exo) -enthalpy of solution is the sum of these individual steps ΔHsolution=ΔH1+ΔH2+ΔH3 |
| solvation | |
| examples of solute-solute attractions | lattice E = E holding the crystal (bonds between cations and anions) together in ionic solids OR a measure of the E contained in a crystal lattus IMF between covalent mcs |
| ion dipole VS dipole dipole | attraction between an ion and a polar mc VS a polar thing and a polar thing |
| dissolving an ionic compound VS a covalent compound into a polar solvent | ionic require more polar solvent mcs to dissolve than covalent -takes more polar solvent mcs to effectively surround & separate ions than it does to separate covalent mcs -overcoming strong ionic attractions requires a highly polar solvent w/ many mcs |
| another way to calculate the deltaH of a solution | ΔHsolution=ΔHhydration−ΔHlattice |
| deltaHlattice | the E change occuring when separted gaseous ions are packed together to form an ionic solid -the E change upon formaion of one moles of a crystalline crystaline compound -to overcome solute interactions, you reverse the lattic E |
| deltaH hydration | combines deltaH2 and deltaH3 in the deltaH solution equation |
| colligative proerpties of solutions deflinition | colligative means that number of particles, not the type of particles colligative properites are basned on the # of dissolved particles, it doesn{t matter the type 'more dissolved particles means that propty is mean intense |
| what affects vapor pressure | 'more solute particles = lower vp -becuase the surface of the liquid is "spotted " w/ solute, the solvent can't escape -ion-dipole bonds are stronger than LDF, so solute-solvent bonds are harder to escape -only applies to non'volitile solutions |
| one solution and one w/ H2O in different containers are placed in a closed system | -H2O moves from a region where vp is high (above water) to in where VP is low (above solution). -some vapor mcs collide w/ solution and enter back into the liquid as solution gains H2O (gets more dilute), its vp goes up -volume of solution increases |
| one solution and one w/ H2O in different containers are placed in a closed system --> how do the vapor preassures above both liquids eventually become the same | system reaches equilibrium when the rate of H2O leaving & entering each liquid balances out in the shared air space. The air above both liquids now has the same concentration of water vapor, which creates one uniform vapor pressure in the closed space. |
| Raoult's law (vp) | VPnew=XH2O(PH2O∘) XH2O = mole fraction of water in the solution PH2O∘ = vapor pressure of pure water at the given temperature VP New = vapor pressure of water above the solution, predicted using Raoult's law (any consistent units) |
| Raoult's law (vp) only works with... | -a liquid solvent -the solute can be either a nonvolitile solid or a nonvolitile liquid |
| how to write the mole fraction for Raoult's law (vp) (it's a little different then normal mole fractions) | moles H2O divided by total moles OR moles H2O divided by moles H2O + i (moles solute) |
| what is i in solutions | # of particles formed per mc or fu -covalent mcs don't dissociate in H2O, so it forms 1 paricle & has an i value of 1 ionic compounds will often have an i of 2 or more -strong acids dissociate completely |
| writing the i value for weak acids | -weak acids partically split, so write acids as an inequality -they wont fully dissociate, so the inequality is this: one number less than a full dissociation < i < full dissociation |
| what makes a solution saturated? Why won't more solute dissolve? | -no more solvent mcs left to interact/bond w/ solute ions -something dissolves & splits into ions (assuming ionic) & H2O mcs (which are polar) sorround the cation or anion w/ the oppsoite charge -higher charge or charge density = more H2O mcs needed |
| charge density | how much charge is concentrated in a small space Large ions = lower cd -Charge is spread over a bigger radius & attract H2O mcs less strongly Small ions have higher cd - strongly attract H2O mcs & more highly hydrated (more tightly bound H2O mcs) |
| when talking about vapor preassure, the observed Vp is the... | the new vapor preassure/the preassure of the solution after the solute is added |
| If a solution obeys raoutes law, it is considered ideal. Why is that? | solute–solvent attractions are just as strong as solvent–solvent or solute–solute attractions. |
| how a stronger solvent solute interaction might cause a deviation from routes law | Stronger solute–solvent interactions: solute & solvent attract each other more than they do themselves. This holds water molecules “tighter.” Fewer water molecules escape as vapor. actual vp is lower than Raoult’s law predicts (negative deviation). |
| how a weaker solvent-solute ineraction might cause a deviation from raoult's law | Weaker solute–solvent interactions: solute and solvent attract each less than in an ideal solution. Wate molecules aren't held as tightly. More water molecules can escape. The actual vp is higher than Raoult’s law predicts (positive deviation). |
| three deviations from raoult's law | 1. higher than predicted vp, liquid-liquid solutions (volatile solute), law assumes the solute is not volatile 2. either stronger or weaker interactions between solute & solvent 3. higher than predicted, incomplete dissociation of ionic substance |
| equation for liquid-liquid solutions | Ptotal=PA+PB=XAPA∘+XBPB∘ |
| what happens to the i value when an ionic compound does not dissociate completely | the i value is lower than expected it gets lower as the concentration goes up -it is ALWAYS lower than predicted cause for it to be higher, you would need more particles than the # produced by full dissociation -not every FU will split |
| why does the vp get lower as solute is added | some surface spots (cause evaporation happens on the surface) are occupied by solute particles & fewer spots are available for solvent particles to escape |
| other henry's law equation | p = c times k p= partical pressure of gas above the solution c = concentration of gas dissolved in the liquid k = L*atm/moles |
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