SCH3U · Unit 4 test review

Solutions & Solubility: Test Review

Everything for the Unit 4 test on one page: the key facts from all 12 lessons, the mistakes that cost the most marks, and a practice test that points you back to the lessons you need.

Key facts

Characteristics of Solutions

Solute = dissolved substance (smaller amount); solvent = dissolving medium (larger amount).
Saturated solution: dissolving and recrystallizing occur at equal rates (dynamic equilibrium).
Supersaturated solutions are unstable and will crystallize out excess solute if disturbed.
Strong electrolytes produce many ions and conduct well; non-electrolytes produce no ions and do not conduct.
Solutions can be gaseous (air), liquid (salt water) or solid (alloys such as brass); miscible liquids mix in all proportions.
Predict conductivity from the formula: soluble ionic compounds and strong acids/bases → strong electrolytes; weak acids and $NH_{3}$ → weak electrolytes; sugars and alcohols → non-electrolytes.

How Dissolving Happens

"Like dissolves like": polar/ionic solutes dissolve in polar solvents; non-polar solutes dissolve in non-polar solvents.
Dissociation: an ionic compound separates into ions that already existed, e.g. $NaCl (s) \to Na^{+} (aq) + Cl^{-} (aq) .$
Ionization: a molecular compound reacts with water to form new ions, e.g. $HCl (g) + H_{2} O (l) \to H_{3} O^{+} (aq) + Cl^{-} (aq) .$
Water is polar because oxygen is more electronegative than hydrogen, giving it a bent shape with $δ^{-}$ at O and $δ^{+}$ at each H.
Hydration is the process of polar water molecules surrounding a freed ion or polar molecule in solution.
Stirring, crushing and heating speed up dissolving; only temperature (and pressure, for gases) changes how much can dissolve.

Solubility Curves

A point ON the solubility curve = saturated; BELOW the curve = unsaturated; ABOVE the curve = supersaturated.
Most solids: solubility increases as temperature increases (though the rate of increase varies a lot between compounds).
Gases: solubility decreases as temperature increases, because faster-moving gas particles escape the liquid more easily.
Mass of solute that crystallizes on cooling = (solubility at higher T − solubility at lower T) scaled to the actual mass of water present.
Scale solubility to the real mass of water: $m_{max} = S \times \frac{m _{water}}{100 g}$ .
Gas solubility increases with pressure — which is why an opened soft drink goes flat.

Concentration

Molar concentration: $c = \frac{n}{V}$ , where $n$ is moles of solute and $V$ is volume of solution in litres.
Rearranged forms: $n = c V$ and $V = \frac{n}{c}$ .
$1 ppm \approx 1 mg solute per L of dilute aqueous solution .$
% w/v = (grams of solute ÷ mL of solution) × 100%, which is simply the grams of solute in every 100 mL of solution.
A standard solution's volume (V) is always the TOTAL solution volume, fixed by a volumetric flask's calibration line — never just the solvent added.
$ppm = \frac{m _{solute}}{m _{solution}} \times 1 0^{6};$ in water 1 ppm ≈ 1 mg/L and 1 ppb ≈ 1 µg/L. Convert mol/L to g/L by multiplying by molar mass.

Dilutions

Dilution equation: $C_{1} V_{1} = C_{2} V_{2}$ (moles of solute are conserved during dilution).
$C_{1}, V_{1}$ = concentration and volume BEFORE dilution (the stock); $C_{2}, V_{2}$ = concentration and volume AFTER dilution.
Diluting always means adding solvent, never adding solute.
Rearranged to solve for any unknown: e.g. $V_{1} = \frac{C _{2} V _{2}}{C _{1}}$ .
Dilution factor $= \frac{V _{2}}{V _{1}} = \frac{C _{1}}{C _{2}}$ ; in a serial dilution the overall factor is the PRODUCT of the step factors.
Mixing two solutions of the same solute: $C_{final} = \frac{C _{a} V _{a} + C _{b} V _{b}}{V _{a} + V _{b}}$ (add moles, divide by total volume).

Ions in Solution & Precipitation Reactions

All Group 1 and NH4+ compounds are soluble; all nitrates and acetates are soluble — no exceptions.
AgCl, PbI2, PbCl2, Hg2Cl2 type halides are insoluble; BaSO4, PbSO4, SrSO4 are insoluble sulfates.
Carbonates, phosphates, hydroxides, and sulfides are generally insoluble (except with Group 1/NH4+).
Total ionic equation: split every aqueous (soluble) compound into ions; keep solids/liquids/gases whole.
Net ionic equation: cancel spectator ions (identical ions unchanged on both sides) from the total ionic equation.

Qualitative Analysis: Identifying Unknown Ions

Selective precipitation: choose a reagent that precipitates only some possible ions, narrowing down the unknown.
Flame test colours: Li+ = crimson red, Na+ = bright yellow, K+ = lilac, Ca2+ = orange-red, Sr2+ = red, Ba2+ = pale green, Cu2+ = blue-green.
Flame colour arises from electrons absorbing heat energy then releasing it as light of specific wavelengths on falling back down.
A single test is rarely conclusive — combine flame tests, selective precipitation, and solubility rules for confident identification.
In a separation scheme, add the MOST selective reagent first (e.g. Cl- before SO4 2- before OH-) and filter after each step.
Anion tests: Cl- + Ag+ → white AgCl; SO4 2- + Ba2+ → white BaSO4; CO3 2- + acid → CO2 bubbles that turn limewater milky. Cu2+ solutions are blue.

Solution Stoichiometry

First step for any solution stoichiometry problem: $n = c V$ to get moles from concentration and volume.
Mole ratio comes from the coefficients of the balanced chemical equation.
Limiting reagent = whichever reactant produces the LESS product when each is assumed to react completely.
Excess reagent's leftover amount = initial moles − moles actually consumed (found using the mole ratio with the limiting reagent).
To find an unknown concentration, get moles of the unknown from the reaction, then $c = \frac{n}{V}$ using the volume of the unknown sample.
After two solutions are mixed, concentrations of leftover ions use the TOTAL volume $V_{1} + V_{2}$ .

Properties & Naming of Acids and Bases

Acids: sour, turn litmus red, react with metals/carbonates, conduct electricity.
Bases: bitter, slippery, turn litmus blue, conduct electricity.
Binary acid naming: hydro- + nonmetal root + -ic acid (e.g. HCl = hydrochloric acid).
Oxyacid naming: -ate ion → -ic acid; -ite ion → -ous acid (e.g. sulfate → sulfuric acid, sulfite → sulfurous acid).
Bases (metal hydroxides) are named like standard ionic compounds: metal name + hydroxide.
In an acid formula, the number of H atoms equals the size of the anion's charge (sulfate $SO_{4}^{2 -}$ → $H_{2} SO_{4}$ ); acid names apply to the (aq) solution.

Arrhenius Theory of Acids and Bases

Arrhenius acid: produces H+(aq) / H3O+(aq) in water. Arrhenius base: produces OH-(aq) in water.
Bases (ionic hydroxides) dissociate: existing OH- ions separate from the solid lattice.
Acids (molecular) ionize: they react with water to form new H3O+ ions.
Strong acid/base = ionizes/dissociates (nearly) completely. Weak acid/base = only partially ionizes/dissociates.
Strong/weak (degree of ionization) is completely independent from concentrated/dilute (amount of solute present).

The pH Scale

$pH = - lo g [H^{+}]$ and $pOH = - lo g [OH^{-}]$ .
$pH + pOH = 14$ at 25°C (from $K_{w} = [H^{+}] [OH^{-}] = 1.0 \times 1 0^{- 14}).$
$[H^{+}] = 1 0^{- pH}$ — use this to reverse a pH value back into a concentration.
pH < 7 acidic, pH = 7 neutral, pH > 7 basic (at 25°C).
Each 1-unit drop in pH means a 10× increase in [H+].

Titration

Neutralization: acid + base → salt + water.
Equivalence point = the theoretical stoichiometric match point (moles acid = moles base, scaled by mole ratio); endpoint = the observed indicator colour change.
A well-chosen indicator makes the endpoint occur very close to the equivalence point.
Core calculation: $n_{a c i d} = c_{a c i d} V_{a c i d}$ , then apply the mole ratio, then solve for the unknown $c$ or $V$ .
Phenolphthalein: colourless below ~pH 8.2, pink above ~pH 10. Litmus: red below ~pH 4.5, blue above ~pH 8.3.

Where marks get lost

Saying a saturated solution "can't dissolve any more solute at all," when really the solid keeps dissolving — it's just balanced by an equal rate of recrystallization (dynamic equilibrium), not a dead stop. (Characteristics of Solutions)
Confusing supersaturated with concentrated — a solution can be dilute yet supersaturated for a solute with very low solubility, and concentrated yet still unsaturated for a highly soluble solute. (Characteristics of Solutions)
Assuming any solution that dissolves fully must conduct electricity — molecular (covalent) solutes like sugar dissolve completely but produce no ions, so they don't conduct. (Characteristics of Solutions)
Mixing up solute and solvent when the amounts are close, e.g. rubbing alcohol (70% isopropanol, 30% water) — the solvent is whichever component is in the larger amount, here the alcohol. (Characteristics of Solutions)
Thinking the solute's mass disappears when it dissolves — the mass of a solution always equals the mass of solute plus the mass of solvent (20 g of salt in 100 g of water makes 120 g of solution), even though the volumes do not always add exactly. (Characteristics of Solutions)
Saying water 'breaks' the ionic bonds inside the crystal — really it out-competes them: many weaker ion-dipole attractions collectively overcome the fewer, stronger internal ionic attractions. (How Dissolving Happens)
Calling every dissolving process 'ionization' — only molecular compounds that chemically react with water to form new ions undergo ionization; ionic solids simply dissociate. (How Dissolving Happens)
Assuming all molecular compounds are non-electrolytes — this is only true if the molecule doesn't ionize; acids like HCl are molecular but ionize to form ions and are strong electrolytes. (How Dissolving Happens)
Forgetting that 'like dissolves like' refers to polarity of the solvent and solute, not just to whether both are liquids or solids. (How Dissolving Happens)
Believing that stirring, shaking or crushing increases solubility — they only make the solute dissolve faster. The maximum amount that can dissolve at a given temperature is unchanged. (How Dissolving Happens)
Applying the 'solubility increases with temperature' rule to gases — gases are the opposite, and this is one of the most common mix-ups on solubility curve questions. (Solubility Curves)
Reading the solubility value at the wrong temperature on the x-axis, or misreading the y-axis scale (especially on curves like KNO3 that rise very steeply). (Solubility Curves)
Forgetting to scale solubility data (given per 100 g water) to the actual mass of water in the problem before comparing to the amount of solute given. (Solubility Curves)
Assuming a point above the curve is impossible — it represents a supersaturated (unstable but real) solution, not an error. (Solubility Curves)
Calculating 'grams that crystallize' as the solubility at the new temperature (forgetting to subtract it from the amount originally dissolved). (Solubility Curves)
Using the volume of water added instead of the total final volume of solution in the $c = n / V$ formula. (Concentration)
Forgetting to convert volume from mL to L before using the concentration formula (a factor-of-1000 error). (Concentration)
Adding water 'to a rough level' instead of exactly to the calibration line on a volumetric flask, which invalidates the 'standard' concentration. (Concentration)
Confusing % w/v (grams per 100 mL) with % v/v (mL per 100 mL) when the solute is a liquid. (Concentration)
Treating ppm as a mass fraction of the solvent alone rather than of the whole solution. (Concentration)
Mixing up which volume is V1 (the smaller volume of concentrated stock taken) and which is V2 (the larger final total volume). (Dilutions)
Adding solute instead of solvent when trying to 'dilute' — this actually increases concentration. (Dilutions)
Forgetting that V2 must be the TOTAL final volume, not just the volume of water added on top of V1 (V2 = V1 + water added, not just the water). (Dilutions)
Using inconsistent units, e.g. mixing mL and L on opposite sides of the equation without converting. (Dilutions)
Using $C_{1} V_{1} = C_{2} V_{2}$ when two DIFFERENT solutes are mixed and react (e.g. HCl with NaOH) — that is solution stoichiometry, not a dilution; and when two solutions of the SAME solute are mixed, averaging the two concentrations instead of adding moles and dividing by the total volume. (Dilutions)
Writing the precipitate as split into ions in the total ionic equation — only AQUEOUS (soluble) compounds get split; solids, liquids, and gases stay as whole formulas. (Ions in Solution & Precipitation Reactions)
Forgetting to balance charge and atoms when splitting compounds into ions (e.g. writing Ba2+ but forgetting the 2 in front of Cl- when balancing BaCl2). (Ions in Solution & Precipitation Reactions)
Cancelling ions that aren't actually identical on both sides (different charges or different states don't count as the same spectator ion). (Ions in Solution & Precipitation Reactions)
Assuming any double-displacement reaction produces a precipitate — many combinations (like KNO3 + NaCl) produce only soluble products and no reaction of interest occurs. (Ions in Solution & Precipitation Reactions)
Mixing up which ion is the cation and which is the anion when predicting the new formula's subscripts (e.g. forgetting Pb2+ needs two I- ions to balance charge in PbI2). (Ions in Solution & Precipitation Reactions)
Relying on flame test colour alone without considering that similar-looking colours (e.g. Ca2+ orange-red vs Sr2+ red) can be confused, especially by an inexperienced observer. (Qualitative Analysis: Identifying Unknown Ions)
Forgetting that a strong sodium contamination can mask other flame colours, since sodium's bright yellow is very intense even in trace amounts. (Qualitative Analysis: Identifying Unknown Ions)
Assuming a selective precipitation test that produces NO precipitate proves an ion's absence with total certainty — it only rules out ions whose compounds with that reagent are known to be insoluble. (Qualitative Analysis: Identifying Unknown Ions)
Using a dirty wire loop for a flame test, which can carry contamination from a previous sample and give a false colour. (Qualitative Analysis: Identifying Unknown Ions)
Adding reagents in the wrong order (e.g. NaOH first), so several cations precipitate together and none can be identified — or forgetting to filter between steps, so a later reagent is tested on a mixture that still contains the earlier precipitate. (Qualitative Analysis: Identifying Unknown Ions)
Jumping straight to the mole ratio with concentration or volume numbers, instead of first converting to moles using n = cV. (Solution Stoichiometry)
Forgetting to convert mL to L before using n = cV. (Solution Stoichiometry)
In limiting reagent problems, assuming the reactant with the smaller volume or smaller concentration is automatically the limiting reagent — you must actually compare moles and apply the mole ratio. (Solution Stoichiometry)
Using the wrong reactant's moles to calculate product amount (must use the LIMITING reagent's moles, not the excess reagent's). (Solution Stoichiometry)
Forgetting to account for the mole ratio (not 1:1) when converting between two different substances' moles. (Solution Stoichiometry)
Naming HNO3 as 'hydronitric acid' — it's an oxyacid (contains the nitrate polyatomic ion), not a binary acid, so it should be 'nitric acid' with no 'hydro-' prefix. (Properties & Naming of Acids and Bases)
Mixing up -ate/-ic and -ite/-ous: forgetting that the ion ending in -ite becomes an acid ending in -ous, not -ic. (Properties & Naming of Acids and Bases)
Forgetting the Roman numeral when naming a hydroxide base of a variable-charge metal (e.g. writing 'iron hydroxide' instead of specifying Iron(II) or Iron(III)). (Properties & Naming of Acids and Bases)
Assuming all acids taste sour is a safe way to identify unknown lab chemicals — this is a historical property, never a real lab safety test. (Properties & Naming of Acids and Bases)
Writing the wrong number of H atoms in an oxyacid formula (e.g. HSO4 for sulfuric acid) — the number of H atoms must equal the size of the ion's negative charge, so sulfate (2−) gives H2SO4 and phosphate (3−) gives H3PO4. (Properties & Naming of Acids and Bases)
Confusing 'strong' with 'concentrated' and 'weak' with 'dilute' — these describe two completely different, independent properties. (Arrhenius Theory of Acids and Bases)
Assuming all acids ionize completely — only STRONG acids do; weak acids like acetic acid mostly stay as intact molecules in solution. (Arrhenius Theory of Acids and Bases)
Forgetting that bases usually dissociate (separating pre-existing ions) while acids usually ionize (forming new ions via reaction with water) — the underlying particle-level process is different even though both are often loosely called 'ionizing.' (Arrhenius Theory of Acids and Bases)
Writing an Arrhenius acid's defining ion as just 'H+' without recognizing it's conventionally shown attached to water as H3O+ (hydronium) in an aqueous solution. (Arrhenius Theory of Acids and Bases)
Splitting a weak acid such as CH3COOH into ions in a total ionic equation, or forgetting that a base like Ba(OH)2 releases TWO hydroxide ions per formula unit (so its OH- concentration is double the base concentration). (Arrhenius Theory of Acids and Bases)
Forgetting the negative sign in $pH = - lo g [H^{+}]$ , which would give a negative pH for an acidic solution instead of a positive one. (The pH Scale)
Thinking a LOWER pH means a LOWER [H+] — it's the opposite: pH and [H+] move in opposite directions because of the negative log. (The pH Scale)
Forgetting that each 1-unit change in pH is a 10-FOLD change in concentration, not a linear (additive) change — e.g. pH 4 does not have 'twice' the acidity of pH 8, it has 10,000 times the [H+]. (The pH Scale)
Using $1 0^{- pH}$ incorrectly, e.g. computing $1 0^{pH}$ instead (dropping the negative sign) when converting pH back to [H+]. (The pH Scale)
Forgetting that pH + pOH = 14 only holds at 25°C, since it depends on the value of Kw at that specific temperature. (The pH Scale)
Confusing the equivalence point (theoretical, calculated) with the endpoint (observed indicator colour change) — they are related but not identical concepts. (Titration)
Forgetting to apply the mole ratio from the balanced equation when the acid and base aren't in a 1:1 ratio (e.g. H2SO4 with NaOH is 1:2). (Titration)
Forgetting to convert volumes from mL to L before using n = cV. (Titration)
Overshooting the endpoint by adding titrant too quickly near the colour change, leading to an inaccurate (too-large) volume reading. (Titration)
Choosing an indicator whose colour-change range doesn't match where the actual equivalence point falls, causing the endpoint to occur well before or after the true equivalence point. (Titration)

Practice test

24 questions from across the unit. Each one shows the lesson it comes from.

Characteristics of Solutions
1
Which best describes a supersaturated solution?
Characteristics of Solutions
2
What word describes two liquids that mix in all proportions to form a single solution, like ethanol and water?
How Dissolving Happens
3
Which best explains why NaCl dissolves in water?
How Dissolving Happens
4
Rank these alcohols from MOST to LEAST soluble in water.
1. 1Butanol (4 carbons)
2. 2Methanol (1 carbon)
3. 3Octanol (8 carbons)
Solubility Curves
5
A solution's data point falls exactly on a solubility curve. The solution is:
Solubility Curves
6
What term describes the release of warm cooling water into a lake or river, lowering its dissolved oxygen?
Concentration
7
In the formula $c = n / V$ , what does V represent?
Concentration
8
Match each situation to the most suitable concentration unit.
Arsenic in a well-water test
Hospital saline IV bag
Rubbing alcohol label
Acid used in a titration calculation
Dilutions
9
Why does the dilution equation $C_{1} V_{1} = C_{2} V_{2}$ hold true?
Dilutions
10
Match each symbol or term to its meaning in a dilution.
$C_{1}$
$V_{1}$
$V_{2}$
$V_{2} / V_{1}$
Ions in Solution & Precipitation Reactions
11
Which pair of solutions will produce a precipitate when mixed?
Ions in Solution & Precipitation Reactions
12
Aqueous silver nitrate is added to tap water and a white precipitate forms. Write the formula of the precipitate, assuming the water contains dissolved chloride.
Qualitative Analysis: Identifying Unknown Ions
13
Which metal ion produces a lilac (pale violet) flame test colour?
Qualitative Analysis: Identifying Unknown Ions
14
Match each observation to the ion it indicates.
White precipitate with acidified AgNO3
Fizzing with dilute acid
Blue solution
Crimson flame
Solution Stoichiometry
15
What is the correct first step when solving a stoichiometry problem given concentration and volume?
Solution Stoichiometry
16
Match each quantity you want to the final conversion you use.
1. A. $c = n / V$
2. B. $m = n M$
3. C. $n = c V$
4. D. $V = n / c$
Mass of product
Volume of a solution needed
Concentration of an unknown
Moles from a solution
Properties & Naming of Acids and Bases
17
What is the correct name for HCl(aq)?
Properties & Naming of Acids and Bases
18
What is the correct name for $HClO_{4} (aq)$ ?
Arrhenius Theory of Acids and Bases
19
According to the Arrhenius definition, what makes a substance a base?
Arrhenius Theory of Acids and Bases
20
Equal volumes of 0.1 mol/L HCl and 0.1 mol/L CH3COOH are tested with a conductivity tester. What is observed, and why?
The pH Scale
21
A solution has pH = 2. Compared to a solution with pH = 5, its [H+] is:
The pH Scale
22
A lake's pH falls from 6.5 to 4.5 because of acid rain. How has its hydrogen ion concentration changed?
Titration
23
What is the key difference between the equivalence point and the endpoint of a titration?
Titration
24
Why is it acceptable for the Erlenmeyer flask to contain a little distilled water before the sample is pipetted in?

Endless drill

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Solutions & Solubility: Test Review

Key facts

Where marks get lost

Practice test

Endless drill

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