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Unit 3: Properties of Substances and Mixtures flashcards

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3.1 Intermolecular and Interparticle Forces

How do intermolecular forces compare with covalent bonds?
Intermolecular (interparticle) forces act between particles, are Coulombic, and are much weaker than covalent bonds. Boiling a molecular liquid breaks none of its covalent bonds; only the forces between molecules are overcome.
What are London dispersion forces, and what makes them stronger?
Attractions from temporary dipoles that induce dipoles in neighbours. Every molecule has them. They grow with more electrons (larger, more polarizable cloud), π bonds and more contact area (less branching). In large molecules they are often the strongest net force.
What are the requirements for hydrogen bonding?
A donor H bonded directly to N, O or F, and an acceptor lone pair on N, O or F in another molecule (or another part of a large molecule). H on carbon cannot hydrogen bond: CH₃OCH₃ and (CH₃)₂CO cannot hydrogen bond with themselves.
Can a molecule that cannot donate a hydrogen bond still take part in one?
Yes, as an acceptor. Acetone has no H on N, O or F, so it cannot hydrogen bond with itself, but its O lone pairs accept hydrogen bonds from water.
Does a polar molecule always have stronger intermolecular forces than a nonpolar one?
No. Compare electron counts first. CCl₄ (nonpolar, 74 electrons) boils far above CH₃F (polar, 18 electrons) because its dispersion forces are much larger. Let dipoles or hydrogen bonds decide only when electron counts are similar.
Why does pentane boil higher than neopentane?
Both have the same formula and electron count, but the straight chain of pentane has more contact area between molecules, so its London dispersion forces are stronger. Polarizability and contact area, not mass itself, are the causes.
What are ion–dipole forces, and how is water oriented around ions?
Attractions between an ion and the oppositely charged end of a polar molecule; usually stronger than dipole–dipole forces. Stronger with larger ion charge and bigger dipole. Cations face water's O (δ−); anions face water's H (δ+). Ca²⁺ in water: ion–dipole.
What forces act between a polar and a nonpolar molecule?
Dipole–induced dipole forces (plus dispersion). The polar molecule's dipole distorts the nonpolar molecule's electron cloud. They are stronger with a bigger dipole and a more polarizable partner.

3.2 Properties of Solids

Why do boiling point and vapor pressure track force strength better than melting point?
Vaporizing overcomes interparticle forces completely, so boiling point and vapor pressure follow force strength directly. Melting only rearranges particles, and crystal packing also affects melting point, so it can mislead.
What are the properties of ionic solids, and why?
Strong attractions between ions: low vapor pressure, high melting and boiling points. Brittle, because like charges meet when layers slide. Conduct only when molten or dissolved, when the ions become free to move.
What are covalent network solids, and how do they behave?
Nonmetal or metalloid atoms bonded covalently throughout (diamond, graphite, SiO₂, SiC). Very high melting points because covalent bonds must break. 3D networks are hard and rigid because bond angles are fixed. They do not conduct, except graphite.
Why is graphite soft, and why does it conduct?
Graphite's covalent layers are held to each other by weak dispersion forces, so layers slide past each other. It conducts because it has mobile electrons: it is the network-solid exception to "never conducts".
What are the properties of molecular solids?
Separate molecules held by weak intermolecular forces: low melting points and no conductivity, because electrons stay in bonds and lone pairs. Melting or boiling breaks no covalent bonds. Solid CO₂ vaporizes easily, while SiO₂ must break covalent bonds to melt.
How do you classify a solid from its properties?
Conducts as a solid → metallic (or graphite). Conducts only molten or dissolved → ionic. Never conducts, very high melting point → covalent network. Never conducts, low melting point → molecular. Shatters → ionic; bends → metallic. Use several properties together.
Which has the higher vapor pressure at 20 °C, ethanol or diethyl ether, and why?
Diethyl ether. It has no H on O, so it cannot hydrogen bond with itself; its weaker intermolecular forces let more molecules escape at a given temperature.
What mistakes are common when explaining properties of solids?
Saying covalent bonds break when a molecular solid melts; calling a solid ionic only because it melts high; saying ionic solids conduct as solids; explaining diamond's hardness with intermolecular forces; naming the force without describing the particles.

3.3 Solids, Liquids, and Gases

What decides whether a substance is a solid, liquid or gas?
A balance between interparticle forces, which pull particles together, and thermal motion, which pushes them apart. Particles keep their size and identity in every phase.
What is the difference between a crystalline and an amorphous solid?
In a crystalline solid particles form a regular, repeating 3-D pattern (table salt). In an amorphous solid (glass) there is no repeating pattern. In both, particles vibrate about fixed positions but do not move past one another.
How should particles of a liquid and a gas be drawn?
Liquid: particles touching, almost as close as in the solid, but disordered. Gas: particles spread through the whole container, far apart. Same particle size in every phase; molecules stay whole (boiled Br₂ is still Br₂).
Why do melting and boiling change volume so differently?
In solid and liquid the particles are in contact, so their molar volumes are similar and melting changes volume only a little. In a gas the particles separate completely, so boiling increases the volume hundreds or thousands of times.
How do you estimate how much farther apart gas particles are than liquid particles?
Spacing scales as the cube root of the volume ratio. If a gas has 1000 times the molar volume of its liquid, the particles are about ∛1000 = 10 times farther apart. Molar volume = M ÷ density.
What happens to particle spacing when a gas is compressed at constant temperature?
The spacing decreases and collisions become more frequent. Gas spacing and collision frequency depend on temperature, pressure and volume.

3.4 Ideal Gas Law

What is the ideal gas law, and which value of R do you use?
PV = nRT, with T in kelvin. R = 0.08206 L atm mol⁻¹ K⁻¹ for pressure in atm, or 62.36 L torr mol⁻¹ K⁻¹ for torr. 1.00 mol in 22.4 L at 273.15 K exerts 1.00 atm.
How do you solve a gas problem where conditions change and n is fixed?
Use P₁V₁ / T₁ = P₂V₂ / T₂ with T in kelvin. Cancel any variable held constant. At constant P and n, heating from 10 °C to 20 °C multiplies V by 293.15 ÷ 283.15 = 1.035, not by 2.
How do you find the partial pressure of one gas in a mixture?
P_A = X_A × P_total, where X_A = n_A ÷ n_total (mole fraction, not mass fraction). Partial pressures add up: P_total = P_A + P_B + …. A gas with X = 0.25 at 2.0 atm has P = 0.50 atm.
How do you find a molar mass from gas density?
M = dRT / P, with d in g L⁻¹, T in K and R matching the pressure units. It comes from substituting n = m / M into PV = nRT.
What do graphs of the gas laws look like?
V against T (K), P against T (K) and V against n: straight lines through the origin. P against V at fixed n and T: a falling curve that flattens; P against 1/V is a straight line through the origin. V against T in °C extrapolates to zero volume at −273.15 °C.
Does a heavier gas exert more pressure than a lighter one?
Not at the same n, V and T. Pressure depends on the number of moles, not mass. Use mole fractions for partial pressures.
When does the ideal gas law work best?
When particles have negligible volume and no attractions: best at low pressure and high temperature. Also assume no gas leaks and gases in a mixture do not react.

3.5 Kinetic Molecular Theory

What are the assumptions of kinetic molecular theory for an ideal gas?
Particles move in continuous, random straight-line motion; have negligible volume compared with the container; do not attract or repel except during collisions; and collide elastically (total kinetic energy conserved).
How does kinetic molecular theory explain gas pressure?
Pressure comes from particles hitting the walls. It rises if the hits are more frequent or harder. Compressing at constant T does not speed particles up; pressure rises because hits per unit area become more frequent.
How is Kelvin temperature related to particle motion?
Kelvin temperature is proportional to the average kinetic energy of the particles. Doubling T in kelvin doubles average KE, but average speed rises only by √2 (about 1.41 times). Always use kelvin; Celsius ratios are meaningless.
Do two gases at the same temperature have the same average speed?
No. They have the same average kinetic energy (KE = ½mv²), so heavier particles move more slowly: v₁ / v₂ = √(M₂ / M₁). Neon (20.18 g mol⁻¹) moves about 2.04 times faster than krypton (83.80 g mol⁻¹). 4 times heavier means 2 times slower.
What does a Maxwell–Boltzmann distribution look like?
It shows the spread of particle speeds or energies at one temperature: starts at zero, has a single peak and a long tail to the right. It is not a symmetric bell curve. The area under the curve represents the total number of particles.
How does a Maxwell–Boltzmann curve change on heating or cooling?
Heating: peak moves to higher speed and the curve becomes lower and wider, with the same area. Cooling: peak moves to lower speed and becomes taller and narrower. On a speed axis, a heavier gas at the same T peaks at lower speed; on a kinetic-energy axis all gases at one T share one curve.

3.6 Deviation from Ideal Gas Law

Which two assumptions of the ideal gas model do real gases break?
Real particles have volume, and real particles attract each other. Both effects are always present; the conditions decide which dominates.
Why do attractions make the measured pressure lower than ideal?
Attractions pull particles back as they approach the walls, so collisions with the walls are less frequent and less forceful. The measured pressure is below nRT / V, and PV/nRT below 1.
Why does PV/nRT exceed 1 at very high pressure?
The particles' own volume fills a large share of the container, reducing the free space. The measured volume (or pressure) is higher than ideal, so PV/nRT > 1 for every gas at extremely high pressure.
Under what conditions do gases behave most and least ideally?
Most ideal: high temperature, low pressure, small nonpolar particles with few electrons (He). Least ideal: close to condensation (low T, high P), and polar, hydrogen-bonded or large highly polarizable molecules.
Which deviates more from ideal behaviour at 25 °C and 5 atm, N₂ or NH₃?
NH₃. Hydrogen bonding and dipole–dipole forces between NH₃ molecules give stronger attractions than the dispersion forces between N₂ molecules. Name the specific force, and note these are intermolecular, not intramolecular.
Does PV/nRT = 1 at one pressure prove a gas is ideal?
No. The effects of attractions (lowering PV/nRT) and particle volume (raising it) may simply cancel at that pressure. Also, collisions do not cause deviations: ideal particles collide too.

3.7 Solutions and Mixtures

What is the difference between a solution and a heterogeneous mixture?
A solution (homogeneous mixture) has the same macroscopic properties throughout, and can be a gas (air), liquid (sugar in water) or solid (brass). A heterogeneous mixture varies with where you sample it: layers, settling solids, visible grains. Colour is not the test; uniformity is.
What is molarity, and how do you calculate it?
Moles of solute per litre of solution: M = n ÷ V, in mol L⁻¹. Convert mL to L first, and use the volume of solution, not of water added. 50.0 mL of 0.120 M contains 0.00600 mol.
How do you find the concentration of each ion in a solution of an ionic compound?
[ion] = salt concentration × number of that ion in the formula; polyatomic ions stay whole. 0.25 M CaCl₂ gives 0.50 M Cl⁻; 0.10 M Al(NO₃)₃ gives 0.30 M NO₃⁻.
How do you calculate a dilution?
Moles of solute are unchanged, so M₁V₁ = M₂V₂ (any volume unit, same on both sides). 20.0 mL of 1.00 M HCl diluted to 250.0 mL gives 0.0800 M.
How do you find the concentration after mixing two solutions of the same solute?
Add the moles and divide by the total volume: M(mix) = (n₁ + n₂) ÷ (V₁ + V₂). Do not average the concentrations. This assumes the solutions do not react and the volumes add.
Does pouring out part of a solution change its concentration?
No. A homogeneous solution has the same molarity everywhere, so a portion has fewer moles but the same molarity.
How do you prepare a solution of known molarity in the lab?
Calculate the mass needed (m = M × V × molar mass), weigh it, dissolve it in less than the final volume of solvent in a volumetric flask, then add solvent up to the calibration mark and mix. The final volume is the volume of solution.

3.8 Representations of Solutions

What must a particle diagram of a solution show correctly?
The relative numbers of particles (concentration and formula ratio, with total positive charge equal to total negative charge) and the interactions between particles (correct orientation of water). Use labels, a key or distinct shapes, never colour alone.
How should ionic and molecular solutes be drawn in solution?
Ionic solutes as separate ions in the formula ratio, with polyatomic ions whole: Ca(NO₃)₂ → 1 Ca²⁺ : 2 NO₃⁻; Na₃PO₄ → 3 Na⁺ : 1 PO₄³⁻. Molecular solutes (glucose, ethanol) stay as whole molecules, never ions. No ion pairs or lattice pieces.
Which way should water molecules point around ions in a diagram?
Water is polar with δ− O and δ+ H. Ion–dipole attractions make O face cations (O toward Mg²⁺) and H face anions.
How do you compare concentrations from two particle diagrams?
Divide particle count by the volume each box represents. Box X (0.10 L, 5 particles) is 50 per L; box Y (0.20 L, 8 particles) is 40 per L, so X is more concentrated. A diluted solution shows the same number of solute particles in a larger volume.
How do you get a molarity from a particle diagram?
n = (number of particles drawn) × (moles each particle represents); then M = n ÷ V for the volume shown.
What interactions are shown for molecular solutes in water or in nonpolar solvents?
Solutes with O–H or N–H groups form hydrogen bonds with water. Nonpolar solutes in nonpolar solvents interact by London dispersion forces only.

3.9 Separation of Solutions and Mixtures

Why can filtration not separate the parts of a solution?
Filtration traps only undissolved solid. Dissolved ions and molecules are far smaller than the filter's pores, so they pass through with the solvent (sugar stays in sugar water). Separating a solution needs a difference in intermolecular forces.
How does chromatography separate a mixture?
A mobile phase (moving solvent) carries the mixture over a stationary phase (paper, silica gel). Components attracted more strongly to the stationary phase move more slowly. Paper and silica are polar, forming hydrogen bonds and dipole–dipole attractions with polar components.
How do you calculate an Rf value?
Rf = distance moved by the spot ÷ distance moved by the solvent front, both measured from the origin line (to the centre of the spot). It has no unit and 0 ≤ Rf ≤ 1. A spot moving 2.8 cm with a solvent front of 7.0 cm has Rf = 0.40.
On a polar (silica or paper) plate, is the spot that moves furthest the most polar?
No, it is the least polar. More polar components are held back by the polar stationary phase and have lower Rf. If the stationary phase were nonpolar, every polarity conclusion would reverse.
What happens to Rf values if a more polar solvent is used?
All Rf values rise, because the polar solvent competes for the polar stationary surface. Rf values are only comparable on the same plate with the same solvent.
How does distillation separate liquids, and which comes off first?
Heat the mixture and condense the vapor. The component with weaker intermolecular forces has the higher vapor pressure and lower boiling point, so it collects first. Pentane distils before octane (fewer electrons, weaker dispersion forces).
What lab-technique points matter for chromatography and distillation?
Draw the origin line in pencil and keep it above the solvent level. In distillation, read the vapor temperature with the thermometer bulb level with the side arm. Close boiling points need a fractionating column; ethanol–water cannot be fully separated.

3.10 Solubility

What do soluble, miscible and immiscible mean?
Soluble: a noticeable amount dissolves. Miscible: two liquids mix in all proportions. Immiscible: they form two layers. Density decides which layer is on top, not whether liquids mix.
What happens to attractions when a substance dissolves?
Solute–solute and solvent–solvent attractions are overcome and solute–solvent attractions form. Dissolving happens readily when the new attractions are comparable to those lost. Molecular solutes such as glucose stay intact; no covalent bonds break.
What does "like dissolves like" mean, and how should you use it in an answer?
Solutes dissolve best in solvents with similar intermolecular interactions. Do not stop at the slogan: name the forces in the solute, in the solvent, and between them. Example: I₂ dissolves in hexane because dispersion forces replace dispersion forces.
Why do nonpolar substances not dissolve in water?
Water does not repel them; it attracts them weakly. Water molecules hold each other much more strongly by hydrogen bonds, and the weak water–solute attractions cannot replace the hydrogen bonds that would be lost.
Why do alcohols become less soluble in water as the carbon chain grows?
The polar O–H head hydrogen bonds with water, but the nonpolar chain interacts only by dispersion forces. As the chain grows, the nonpolar part dominates. Solubility in water, lowest to highest: 1-hexanol, 1-butanol, ethanol.
Are all ionic compounds soluble in water?
No. "Like dissolves like" ignores strengths: CaCO₃ and AgCl are ionic but barely dissolve because their ion–ion attractions are strong. Always soluble in water: all sodium, potassium, ammonium and nitrate salts.
Why do ethanol and water mix in all proportions?
Ethanol–water hydrogen bonds replace the ethanol–ethanol and water–water hydrogen bonds that are lost, so the new attractions are comparable to the old.

3.11 Spectroscopy and the Electromagnetic Spectrum

What is the order of the electromagnetic spectrum by photon energy?
Increasing photon energy (decreasing wavelength): radio, microwave, infrared, visible, ultraviolet, X-ray, gamma. Long wavelength means low energy.
Which type of molecular transition does each region of radiation cause?
Microwave (1 mm to 1 m): rotational. Infrared (700 nm to 1 mm): vibrational, bonds stretch and bend. Visible (400–700 nm) and ultraviolet (10–400 nm): electronic, an electron moves to a higher level; UV may break bonds.
Why can a molecule absorb only certain photons?
Rotational, vibrational and electronic energies are quantized. A photon is absorbed only if its energy matches a gap between levels. The gaps grow in the order rotational ≪ vibrational ≪ electronic, so the region absorbed reveals the transition type.
What is emission?
The reverse of absorption: the species drops down an energy gap and gives out a photon whose energy matches that gap.
Why does a coloured solution look the colour it does?
It absorbs some visible wavelengths and you see what is left: roughly the complementary colour of the light absorbed. Absorbs orange → looks blue; absorbs violet → looks yellow.
Why does neon gas show no infrared absorption?
Single atoms have no bonds to vibrate, and no molecular rotational spectrum. Atoms show only electronic transitions.
What mistakes are common with the electromagnetic spectrum?
Reversing the energy order; saying infrared excites electrons (it is vibrational) or microwaves cause vibration (they are rotational); comparing wavelengths in different units (convert mm, μm, nm to m); confusing brightness with photon energy.

3.12 Properties of Photons

How are wavelength, frequency and photon energy related?
c = λν with c = 2.998 × 10⁸ m s⁻¹, λ in m, ν in s⁻¹. E = hν with h = 6.626 × 10⁻³⁴ J s, so E = hc/λ. λ and ν are inversely proportional; E is directly proportional to ν and inversely proportional to λ.
How do you find the energy of one photon from its wavelength in nm?
Convert nm to m first (1 nm = 10⁻⁹ m), then E = hc/λ in J per photon. Example: λ = 5.00 × 10⁻⁷ m gives ν = 6.00 × 10¹⁴ s⁻¹; a photon of 1.00 × 10¹⁵ s⁻¹ has E = 6.63 × 10⁻¹⁹ J.
How do you convert a photon energy into kJ mol⁻¹?
E = hν is per photon. Multiply by N_A = 6.022 × 10²³ mol⁻¹ for J mol⁻¹, then divide by 1000 for kJ mol⁻¹. This lets you compare with bond energies.
How is the energy of an emitted or absorbed photon related to energy levels?
E(photon) = E(upper) − E(lower), the gap, not a level energy. A drop from −3.0 × 10⁻¹⁹ J to −7.0 × 10⁻¹⁹ J emits a 4.0 × 10⁻¹⁹ J photon. A negative photon energy means you subtracted the wrong way.
Can brighter light of the same colour cause a transition that dimmer light cannot?
No. Brightness is the number of photons, not the energy of each one. One photon causes one transition, so the energy per photon decides which gaps can be reached.
Why do atomic spectra show sharp lines?
Energy levels are fixed values, so only photons whose energy exactly matches a gap are absorbed or emitted. Each gap gives one sharp wavelength.

3.13 Beer-Lambert Law

What is the Beer-Lambert law, and what are its units?
A = εbc. A (absorbance) has no unit; ε (molar absorptivity) in M⁻¹ cm⁻¹; b (path length) in cm; c (concentration) in M. A is proportional to both b and c.
What does ε depend on?
The absorbing species and the wavelength only, not concentration. It measures how strongly one species absorbs at one wavelength.
What is a calibration curve in spectrophotometry?
A plot of A against c for standard solutions with the same cuvette and wavelength. It is a straight line through the origin with gradient εb, so ε = gradient ÷ b. Read the unknown's c from its A, within the range of the standards.
Why is the spectrophotometer set to λmax?
λmax is the wavelength of maximum absorbance: ε is largest, so the measurement is most sensitive and least affected by small wavelength drift. Zero the instrument with a blank (cuvette plus solvent) first.
How do changes in path length and concentration affect absorbance?
A is proportional to each. Doubling the path length doubles A (0.20 → 0.40). Halving c and doubling b leaves A unchanged, because the beam meets the same number of particles.
How do smudges, bubbles or a diluted sample affect the calculated concentration?
Smudges, bubbles, cloudiness and a longer cuvette raise A, so c is calculated too high. Water droplets that dilute the sample lower A, so c is too low. State whether A goes up or down, then link it to c.
How do you calculate ε from one measurement?
ε = A ÷ (bc). A 3.00 × 10⁻⁵ M solution in a 1.00 cm cuvette with A = 0.450 has ε = 1.50 × 10⁴ M⁻¹ cm⁻¹.
How do you account for a dilution made before a spectrophotometric measurement?
c(original) = c(measured) × dilution factor, where dilution factor = final volume ÷ sample volume. Forgetting this after reading the calibration line is a common error.

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