Summary of "2022개정 [중2 오투과학] 1-02 물질의특성(2)-1"

Summary

Dissolution: a substance (solute) disperses into another substance (solvent) and mixes uniformly; the resulting homogeneous mixture is a solution. Solute = substance being dissolved (e.g., sugar). Solvent = substance doing the dissolving (usually a liquid, e.g., water).

1) Dissolution and solutions — definitions and key points

A solution is homogeneous: any sample taken from it has the same composition.
Solute vs. solvent:
- Solute — the substance dissolved (e.g., sugar).
- Solvent — the substance that does the dissolving (typically a liquid, e.g., water).
Solubility states:
- Unsaturated: more solute can still dissolve.
- Saturated: at a given temperature, the maximum amount of solute is dissolved; dynamic equilibrium where rate of dissolution = rate of precipitation.
- Supersaturated: more solute is dissolved than normally possible at that temperature (rare; outside typical middle‑school scope).

2) Solubility (quantitative definition) and solubility curves

Solubility is defined as the maximum grams of solute that dissolve in 100 g of solvent at a specific temperature.
Factors affecting solubility:
- Temperature — for many solids, solubility increases with temperature (extent depends on the substance).
- Pressure — generally negligible effect for solids and liquids.
Solubility curves:
- A point on the curve = saturated at that temperature.
- Points below the curve = unsaturated.
- Slope steepness indicates how strongly solubility changes with temperature (steeper → greater change).
Cooling a saturated hot solution decreases solubility; excess solute precipitates.
- Precipitated mass = initial dissolved amount − solubility at the lower temperature.
Example calculation method:
- Scale solubility to the per‑100 g solvent basis by proportion. For example: 10 g solute in 25 g solvent → x = 10 × (100 / 25) = 40 g per 100 g solvent.

3) Solubility of gases (temperature & pressure effects)

Temperature: gas solubility decreases as temperature increases (cold water holds more dissolved gas).
Pressure: gas solubility increases as pressure increases (qualitative Henry’s law).
Practical consequences:
- Cold, sealed soda retains more CO2; warming or opening lowers pressure and CO2 escapes as bubbles.
- Shaking a sealed carbonated drink speeds release of gas on opening.
- At low external pressure (high altitude or vacuum) water boils at a lower temperature; in a vacuum it can boil at room temperature.

4) Phase changes, melting point, freezing point, boiling point

Phase-change names and heat flow:
- Solid → liquid: melting (endothermic).
- Liquid → gas: vaporization/boiling (endothermic).
- Solid → gas: sublimation.
- Reverse processes (freezing/solidification, condensation/liquefaction, deposition) are exothermic.
Melting point = freezing point for a given substance (temperature where solid and liquid coexist).
Boiling point = temperature where liquid and gas coexist while vaporization occurs; temperature remains constant during the phase change because added heat is used as latent heat.
During a phase‑change plateau, two phases coexist (solid + liquid at melting point; liquid + gas at boiling point).
The amount of substance affects how long it takes to reach a phase change, but not the temperature at which the change occurs.
Boiling point depends on external pressure:
- Higher external pressure → higher boiling point (e.g., pressure cooker).
- Lower external pressure → lower boiling point (e.g., high‑altitude cooking).
Determine a substance’s state at a given temperature by comparing that temperature with its melting and boiling points (examples: butane, liquid nitrogen).

5) Applications and examples

Uses of different melting/boiling points:
- Low‑melting alloys for soldering.
- Glue sticks that melt to bond and solidify on cooling.
- Tungsten as a high‑melting‑point material for light‑bulb filaments.
- Liquid nitrogen for cold storage and demonstrations.
Classroom problem examples:
- Convert solubility for a given mass of solvent to per‑100 g basis by proportion.
- Determine precipitation on cooling: e.g., sodium nitrate with 140 g/100 g at 70°C cooled to 87 g/100 g at 20°C → precipitated = 140 − 87 = 53 g per 100 g solvent.

6) Problem‑solving methodology (step‑by‑step)

Express solubility as grams of solute per 100 g of solvent (standard basis).
If given solute mass for a different solvent mass, scale by proportion:
- x (g per 100 g) = given solute mass × (100 / given solvent mass).
To find precipitated amount when cooling:
- Calculate the initial dissolved amount per 100 g solvent.
- Find the solubility per 100 g at the lower temperature (from a curve or table).
- Precipitated mass = initial dissolved amount − solubility at lower temperature (if positive).
To turn an unsaturated solution into saturated:
- Add more solute until saturation, or
- Lower the temperature until the existing solute reaches saturation capacity.
To compare temperature sensitivity among substances: inspect the slope of their solubility curves — steeper slope = larger change with temperature (e.g., KNO3 vs NaCl).