8. Chemical Reactions in Aqueous Solutions

Dissolved particles get in the way of water molecules joining the ice crystal, whatever they are, and the more there are, the lower the temperature must fall before ice can form. A property that depends on the count of particles and not their nature is called colligative. Boiling-point rise and osmotic pressure are the other familiar ones.

Oppositely charged ions in solution sometimes stick together for a moment, and then count as one particle. The effect is largest for ions with several charges, such as the sulfate ion or aluminum ion. Real van 't Hoff factors are slightly below the whole numbers here, and nearer to them for dilute solutions.

Yes. Salt dissolves in the film of water on the ice and lowers its freezing point, so the ice melts at temperatures below 0 °C. Calcium chloride, which gives three particles per formula unit, works at lower temperatures than the same amount of table salt.

K_f = 1.853 K kg/mol for water is read from Wikipedia's article on the cryoscopic constant, which cites SI Chemical Data. The temperatures on the thermometer are computed from it: nothing has been measured here, and the page says so.

Because the solute takes up room: 9 g of salt and a litre of water make slightly more than a litre of solution. Molarity is defined by the final volume, which is why you dissolve the solute and then fill to the mark of a volumetric flask, rather than adding it to a measured litre of water.

Molarity is moles per litre of solution. Molality, used in the freezing-point experiment, is moles per kilogram of solvent. Molality does not change with temperature because masses do not, so it is used for colligative properties. For dilute solutions in water the two are almost equal.

It has the same concentration of dissolved particles as blood plasma, so water does not move in or out of blood cells (it is isotonic). A solution that is too dilute makes cells swell, and one too concentrated makes them shrink.

In outline, yes: Thomas Latta described giving saline intravenously to cholera patients in 1832, and some recovered. The method was then largely forgotten for decades, and returned once solutions could be made reliably. The dosage in his reports is the part that was a matter of trial and error.

The usual ones: compounds of sodium, potassium and ammonium are soluble; so are all nitrates, chlorates, perchlorates and acetates. Chlorides, bromides and iodides are soluble except with silver, lead and mercury(I). Sulfates are soluble except with barium, lead, silver and (slightly) calcium. Hydroxides are insoluble except with group 1 and barium; carbonates, phosphates and sulfides are insoluble except with group 1 and ammonium. They are summaries of experiment, not laws of nature, so there are borderline cases.

In the solid, each ion is surrounded by ions of the opposite charge, as in the crystal you built in Unit 4, and the attraction is strong. In solution each ion is surrounded by water. Whichever gives the lower energy wins: for lead iodide the lattice wins, and for sodium chloride water does.

As far as this reaction goes, yes. They keep the solution electrically neutral, and they are there at the end exactly as at the start. Evaporate the water and they will crystallize as potassium nitrate. But they are not part of the chemical change, so the net ionic equation leaves them out.

Then nothing has happened. Mixing sodium chloride and potassium nitrate solutions gives a solution that contains sodium, chloride, potassium and nitrate ions, and it is the same solution you would get from dissolving sodium nitrate and potassium chloride together. Writing "NaCl + KNO₃ → NaNO₃ + KCl" would suggest a reaction, but no reaction has occurred.

The equivalence point is where exactly the right amounts have been mixed: here, pH 7. The endpoint is where the indicator changes colour, which for phenolphthalein is a little later, around pH 8.2 to 10. Because the pH rises so steeply at the equivalence point, the difference is a fraction of a drop, and a good indicator makes the two agree to within the precision of the burette.

Far from the equivalence point, a whole millilitre barely changes the pH. Near it, one drop does. The curve is nearly vertical there, so to locate the volume to a tenth of a millilitre you have to go drop by drop.

Svante Arrhenius, in 1884, defined an acid as something that gives H⁺ in water and a base as something that gives OH⁻. In 1923 Johannes Brønsted and Thomas Lowry independently widened the idea: an acid is a proton donor and a base a proton acceptor, which includes ammonia, which has no OH to give. Strong acids and bases, such as HCl and NaOH, are completely split into ions in water. Weak ones, such as acetic acid, are only partly, and their titrations have a different shape.

It is a simulation of one. The pH at each point is calculated from the exact equations for a strong acid with a strong base, and the colour is that of phenolphthalein, which is colourless below about pH 8.2 and pink above. The unknown concentrations were chosen for the exercise.

No. They are a bookkeeping convention, assigned as if every bond were fully ionic. In NaCl the +1 and −1 are close to the truth. In CO₂ the carbon is +4 by the rules but is not really carrying four positive charges. The numbers are useful because the total change in them across a reaction counts the electrons that moved.

Because the three iron atoms are not all alike: Fe₃O₄ contains one Fe²⁺ and two Fe³⁺ for every four oxygens. Their average is (2 + 3 + 3) / 3 = 8/3. A fraction in an oxidation number is a sign that atoms of the same element in the species have different numbers.

Yes. In burning methane, carbon goes from −4 to +4 (it is oxidized) and each oxygen from 0 to −2 (reduced). Burning a fuel is always an electron transfer from the fuel to oxygen, and the large amount of energy released is the reason fuels burn.

From metallurgy: heating a metal ore with charcoal "reduces" it to the metal. Iron(III) oxide is reduced to iron, in which iron goes from +3 to 0, gaining electrons. The same word now covers every gain of electrons, whether there is a furnace or not.

E° for Ag⁺/Ag is +0.800 V and for Cu²⁺/Cu is +0.337 V. The difference, +0.463 V, is positive, so copper loses electrons to silver ions: Cu + 2Ag⁺ → Cu²⁺ + 2Ag. The solution turns blue as copper goes in and silver needles grow on the strip.

Only differences of potential can be measured, never a single one. Chemists chose the hydrogen electrode, H⁺ in solution with hydrogen gas over platinum, as the zero of the scale. Metals with a negative E° (zinc, iron) will displace hydrogen from acid and bubble; those with a positive E° (copper, silver) will not.

The potential says what is possible, not how fast. Aluminium coats itself instantly with a thin, tight layer of oxide that seals the metal underneath. Iron's oxide, rust, is flaky and lets the attack continue. The series tells you who would win, not how quickly.