Advanced High School Chemistry Flashcards: Complete 9-Part Course Review
Review a rigorous nine-part high school chemistry course with 450 concise cards covering concepts, models, equations, calculation setup, and lab reasoning.
About this deck
Study advanced high school chemistry through 450 independently written English flashcards arranged as one nine-part sequence. The course moves from atomic structure and bonding to properties of substances and mixtures, reactions, kinetics, thermochemistry, equilibrium, acids and bases, thermodynamics, and electrochemistry. Prerequisites appear before dependent models and calculations.
What the cards practice
The cards practice five useful recall paths: concept to explanation; model or representation to interpretation; equation to meaning and use; short setup to a result with units and reasoning; and laboratory observation to a chemical conclusion. This includes definitions, relationships, conditions, contrasts, particle and energy models, focused calculation steps, measurements, errors, and visible changes.
Selected reverse and contrast prompts appear only when the reverse direction has one clear standalone target. Mechanical permutations, graph-dependent prompts that need a missing figure, copied test formats, long multipart derivations, and visual recall tied to third-party figures are excluded. The review scheduler handles long-term spacing after installation.
The prompts, answers, examples, organization, metadata, and cover were created independently from common chemistry knowledge and original work. No protected questions, answer choices, scoring materials, curriculum prose, commercial card text, source figures, or third-party media were copied.
The Common knowledge · CC0 1.0 label applies only to the original prompts, answers, examples, organization, metadata, and cover, to the extent applicable rights exist. It does not claim ownership of scientific facts, equations, or third-party material.
Cards in this deck
Card 1
Question
What does one mole count?
Answer
Exactly 6.02214076 × 10^23 representative particles.
Card 2
Question
What does a peak in an element's mass spectrum represent?
Answer
An isotope with a particular mass-to-charge ratio; for singly charged monatomic ions, the position tracks isotopic mass.
Card 3
Question
What does an empirical formula show?
Answer
The lowest whole-number ratio of the elements' atoms in a compound.
Card 4
Question
How does a mixture differ from a pure substance at the particle level?
Answer
A mixture contains chemically distinct representative units in variable proportions; a pure substance contains one element or compound with fixed composition. Different isotopes do not make an elemental sample a mixture.
Card 5
Question
Which particles make up an atom's nucleus?
Answer
Protons and neutrons. Electrons occupy the space outside the nucleus.
Card 6
Question
What does a larger binding energy on a PES spectrum mean?
Answer
More energy is required to remove that electron, so it is held more strongly by the nucleus.
Card 7
Question
How does atomic radius generally change across a period and down a group?
Answer
It decreases from left to right as effective nuclear charge rises, and it increases down a group as additional electron shells increase distance and shielding.
Card 8
Question
What typical ion charge do Group 1 metals form?
Answer
+1, by losing their one valence electron.
Card 9
Question
How do you convert moles to particles?
Answer
Multiply by Avogadro's number: particles = moles × 6.022 × 10^23 mol^-1.
Card 10
Question
How is average atomic mass estimated from isotope data?
Answer
Add each isotopic mass multiplied by its fractional abundance.
Card 11
Question
How is an element's mass percent in a compound calculated?
Answer
Divide the mass contributed by that element by the compound's molar mass, then multiply by 100%.
Card 12
Question
How can measured elemental composition reveal a sample's purity?
Answer
Compare the measured mass fraction with the fraction expected for the pure compound; a mismatch indicates another component.
Card 13
Question
How do you build a ground-state electron configuration with the Aufbau principle?
Answer
For ordinary ground states, move through the periodic table in atomic-number order, filling each s, p, d, or f block as it appears. The subshell capacities are s², p⁶, d¹⁰, and f¹⁴. For example, Br is [Ar] 4s² 3d¹⁰ 4p⁵.
Card 14
Question
What does the relative area or height of an ideal PES peak indicate?
Answer
The relative number of electrons in the corresponding subshell.
Card 15
Question
How does first ionization energy generally change across a period and down a group?
Answer
It increases from left to right as effective nuclear charge rises, and it decreases down a group as distance and shielding make a valence electron easier to remove.
Card 16
Question
Why do elements in the same group form similar compounds?
Answer
Their ground-state valence patterns repeat, including which outer subshells are full or partly full. That leads to similar bonding and typical ion charges.
Card 17
Question
How do you convert a sample's mass to moles?
Answer
Divide its mass by its molar mass: n = m/M.
Card 18
Question
Which mass-spectrum interpretation lies outside the usual single-element model used in this deck?
Answer
Assigning peaks in mixtures or peaks from multiply charged or polyatomic species; the standard model uses singly charged monatomic ions of one element.
Card 19
Question
What does the law of definite proportions state?
Answer
Every pure sample of a given compound has the same element mass ratios.
Card 20
Question
Why can two samples of the same mixture have different compositions?
Answer
Mixture components are physically combined, so their relative amounts are not fixed by a chemical formula.
Card 21
Question
How does Coulomb's law connect charge and separation to attraction?
Answer
Attraction grows with the magnitude of the charge product and decreases with the square of the separation distance.
Card 22
Question
Which PES electrons usually appear at the highest binding energy?
Answer
Core electrons closest to the nucleus, because they feel the strongest nuclear attraction.
Card 23
Question
How does electron affinity generally change across a period and down a group?
Answer
Electron gain generally becomes more favorable from left to right across a period and less favorable down a group as distance and shielding increase. Stable subshell patterns create substantial exceptions.
Card 24
Question
Why are alkali metals generally more reactive down the group?
Answer
Their valence electron is farther from the nucleus and easier to remove.
Card 25
Question
How many moles are in 18.0 g of H₂O?
Answer
About 0.999 mol. Use 18.0 g ÷ 18.02 g mol^-1.
Card 26
Question
An element is 75% isotope 10 and 25% isotope 11; what is its average atomic mass?
Answer
10.25 u. Calculate (0.75 × 10) + (0.25 × 11).
Card 27
Question
A compound is 40.0% C, 6.7% H, and 53.3% O by mass; what is its empirical formula?
Answer
CH₂O. For a 100 g sample, convert each mass to moles and divide by the smallest amount.
Card 28
Question
A 10.0 g impure sample contains 8.5 g of the target compound; what is its mass-percent purity?
Answer
85%. Calculate (8.5 g ÷ 10.0 g) × 100%.
Card 29
Question
Which electrons are removed first when a transition metal forms a cation?
Answer
Electrons in the occupied orbital with the highest principal quantum number: 4s before 3d. For example, Fe²⁺ is [Ar] 3d⁶.
Card 30
Question
A PES spectrum has peaks proportional to 2, 2, and 6 electrons; which configuration fits?
Answer
1s² 2s² 2p⁶, the configuration of Ne.
Card 31
Question
How does electronegativity generally change across a period and down a group?
Answer
It increases from left to right across a period and decreases down a group as atomic size and shielding increase.
Card 32
Question
What empirical formula results from Al³⁺ and O²⁻?
Answer
Al₂O₃, because two Al³⁺ ions balance three O²⁻ ions.
Card 33
Question
How does a particle's mass in atomic mass units relate to its molar mass?
Answer
The numerical value is the same: a molecular or formula-unit mass of x u corresponds to a molar mass of x g mol^-1.
Card 34
Question
What does the tallest isotope peak usually indicate in a simple mass spectrum?
Answer
The most abundant isotope, assuming comparable detection response and singly charged ions.
Card 35
Question
How much oxygen is present in 25.0 g of a compound that is 32.0% oxygen by mass?
Answer
8.00 g O. Multiply 25.0 g by 0.320.
Card 36
Question
What does a particle diagram with two unbonded species in changing ratios represent?
Answer
A mixture, because more than one particle type is present and the ratio is not fixed in a formula unit.
Card 37
Question
What distinguishes valence electrons from core electrons?
Answer
Valence electrons are available for bonding or ion formation; main-group valence electrons occupy the outermost shell, while transition metals may also use (n−1)d electrons. Core electrons mainly shield nuclear charge.
Card 38
Question
Why can PES peak groups reveal an atom's occupied subshells?
Answer
Electrons in different subshells require distinct removal energies, producing separate binding-energy groups.
Card 39
Question
How do ion radii compare with neutral atoms and within an isoelectronic series?
Answer
Cations are smaller than their neutral atoms, while anions are larger. Among species with the same electron count, more protons pull the electrons closer and produce the smaller radius.
Card 40
Question
What formula is expected for a compound between a Group 2 metal M and a Group 17 nonmetal X?
Answer
MX₂, because M forms M²⁺ and X forms X⁻.
Card 41
Question
When is a covalent bond considered nonpolar?
Answer
When the bonded atoms have identical or very similar electronegativities, so the shared electron density is distributed approximately evenly.
Card 42
Question
Why does a bonded pair of atoms have an equilibrium bond length?
Answer
At that separation, attractive and repulsive interactions balance at minimum potential energy.
Card 43
Question
How are particles arranged in an ionic solid?
Answer
Cations and anions occupy a repeating three-dimensional lattice held by electrostatic attraction.
Card 44
Question
What model explains bonding in a metal?
Answer
Positive metal cores are held together by attraction to mobile, delocalized valence electrons.
Card 45
Question
How do you construct a Lewis diagram?
Answer
Count total valence electrons, adding electrons for a negative charge and subtracting them for a positive charge. Choose a skeleton, connect atoms with single bonds, complete terminal duets or octets, and place remaining electrons on the central atom. Add multiple bonds if needed, then check the electron total and formal charges.
Card 46
Question
What does resonance mean in a molecule or ion?
Answer
Resonance uses two or more valid Lewis diagrams with the same atom arrangement but different electron placement. The actual electron distribution is a hybrid; equivalent contributors have equal weight.
Card 47
Question
What determines molecular shape in VSEPR theory?
Answer
Electron domains around the central atom arrange to minimize repulsions.
Card 48
Question
How does an ionic bond differ from a covalent bond?
Answer
Ionic bonding is attraction among oppositely charged ions in an extended structure; covalent bonding uses shared electron density between atoms.
Card 49
Question
What happens to potential energy when bonded atoms are pushed much closer than equilibrium?
Answer
Potential energy rises sharply because nucleus–nucleus and electron–electron repulsions dominate.
Card 50
Question
Why are many ionic solids brittle?
Answer
A shifted lattice can align like charges, creating strong repulsion that splits the crystal.
Card 51
Question
What molecular shapes arise from two electron domains with no lone pairs and from three domains with zero or one lone pair?
Answer
Two bonding domains give linear with a 180° angle. Three domains with no lone pairs give trigonal planar with 120° angles; replacing one bond with a lone pair gives bent with an angle slightly below 120°.
Card 52
Question
Why are metals electrically conductive as solids?
Answer
Their delocalized electrons can move through the solid when an electric field is applied.
Card 53
Question
How is formal charge calculated for an atom in a Lewis diagram?
Answer
Formal charge = valence electrons − nonbonding electrons − half the bonding electrons.
Card 54
Question
Why can't electronegativity difference alone classify a bond as ionic or covalent?
Answer
Bonding lies on a continuum. A larger difference means more ionic character, but the element types and especially the compound's properties give the best classification.
Card 55
Question
Which shapes and bond-angle trends arise as lone pairs replace bonds in four electron domains?
Answer
Four bonds give tetrahedral with ideal 109.5° angles. One lone pair gives trigonal pyramidal with smaller angles; two lone pairs give bent with typically smaller angles again because lone pairs repel more strongly than bonding pairs.
Card 56
Question
What feature of a potential-energy curve represents bond dissociation energy?
Answer
The energy difference from the curve's minimum to the separated-atoms limit.
Card 57
Question
When does an ionic compound conduct electricity?
Answer
When molten or dissolved so its ions can move; not as a rigid solid lattice.
Card 58
Question
What is a substitutional alloy?
Answer
An alloy in which atoms of a similar size replace some host-metal atoms in the lattice.
Card 59
Question
How do two, three, and four electron domains map to hybridization?
Answer
Two domains map to sp, three to sp², and four to sp³, with ideal angles of 180°, 120°, and 109.5°. Hybridization involving d orbitals is outside this deck’s scope.
Card 60
Question
What usually makes one resonance contributor more favorable than another?
Answer
Smaller formal-charge magnitudes, appropriate negative charge on more electronegative atoms, and complete valence shells where applicable.
Card 61
Question
How many sigma and pi bonds are in single, double, and triple bonds?
Answer
A single bond has one sigma bond; a double has one sigma and one pi bond; a triple has one sigma and two pi bonds. Head-on sigma overlap is stronger than side-by-side pi overlap.
Card 62
Question
Why is a polar covalent bond polar?
Answer
Unequal electronegativity creates an uneven sharing of electron density and partial charges.
Card 63
Question
Which molecular shapes arise as lone pairs replace bonds in five electron domains?
Answer
Five bonds give trigonal bipyramidal; four bonds and one lone pair give seesaw; three bonds and two lone pairs give T-shaped; two bonds and three lone pairs give linear.
Card 64
Question
How do ionic charge and ionic radius affect attraction between ions?
Answer
Larger charge magnitudes and smaller ionic radii produce stronger attraction because the charge product increases and the ion centers are closer.
Card 65
Question
Why do ionic solids often have high melting points?
Answer
Many strong Coulombic attractions throughout the lattice must be overcome to free the ions.
Card 66
Question
What is an interstitial alloy?
Answer
A smaller atom occupies holes between host-metal atoms, often making lattice layers harder to slide.
Card 67
Question
What shape has six bonding domains and no lone pairs on the central atom?
Answer
Octahedral.
Card 68
Question
Which elements commonly form incomplete octets in stable Lewis diagrams?
Answer
Hydrogen forms a duet, and electron-deficient central atoms such as boron or beryllium can have fewer than eight electrons.
Card 69
Question
How do bond order and atomic size affect covalent bond length and strength?
Answer
Within a comparable bond family, higher bond order gives shorter, stronger bonds. Larger bonded atoms generally give longer bonds, which are often weaker because their orbitals overlap less effectively.
Card 70
Question
What bonding model best fits a sample that is malleable and conducts as a solid?
Answer
Metallic bonding with mobile, delocalized electrons and nondirectional attractions.
Card 71
Question
What shape has six electron domains, five bonds, and one lone pair?
Answer
Square pyramidal.
Card 72
Question
Which lattice should have stronger attractions: MgO or NaCl, assuming similar separations?
Answer
MgO, because the charge product for Mg²⁺ and O²⁻ is larger than for Na⁺ and Cl⁻.
Card 73
Question
Why are pure metals often malleable?
Answer
Metal cores can shift while the mobile electron sea maintains nondirectional attraction instead of exposing fixed like-charge planes.
Card 74
Question
What is the best Lewis structure for CO₂?
Answer
O=C=O, with two lone pairs on each oxygen and no formal charges.
Card 75
Question
What shape has six electron domains, four bonds, and two opposite lone pairs?
Answer
Square planar.
Card 76
Question
What limitation does an odd total number of valence electrons create for a Lewis diagram?
Answer
At least one electron must remain unpaired, so not every atom can have a complete paired-electron octet.
Card 77
Question
What does a higher bond order do to a bond's potential-energy curve?
Answer
It generally places the minimum at a shorter internuclear distance and makes the well deeper, corresponding to a shorter bond and a larger bond-dissociation energy.
Card 78
Question
When can a carbon–carbon double bond produce geometric isomers?
Answer
When each carbon has two different substituents. The pi bond restricts rotation, so distinct spatial arrangements can persist.
Card 79
Question
When may a third-period central atom exceed an octet in a Lewis diagram?
Answer
When the valid electron count and lower formal charges favor an expanded valence shell, as in species such as SF₆.
Card 80
Question
How do you decide whether a molecule with polar bonds is polar overall?
Answer
Add the bond-dipole vectors using the molecular shape; symmetry may cancel them, while an asymmetric arrangement leaves a net dipole.
Card 81
Question
Which interparticle forces act between all atoms and molecules?
Answer
London dispersion forces, caused by temporary and induced dipoles.
Card 82
Question
What four broad solid types does this deck compare?
Answer
Ionic, metallic, molecular, and covalent-network solids.
Card 83
Question
How do gas particles differ from liquid particles?
Answer
Gas particles are much farther apart and move independently; liquid particles stay close but can move past one another.
Card 84
Question
What relationship connects pressure, volume, amount, and temperature for an ideal gas?
Answer
PV = nRT, with absolute temperature in kelvins and units consistent with R.
Card 85
Question
What does temperature measure in kinetic molecular theory?
Answer
The particles' average translational kinetic energy.
Card 86
Question
What two ideal-gas assumptions fail most clearly for real gases?
Answer
Particles have nonzero volume and experience intermolecular attractions.
Card 87
Question
How is molarity defined?
Answer
Moles of solute per liter of solution: M = n/V.
Card 88
Question
What must a correct particulate diagram of NaCl(aq) show?
Answer
Separated Na⁺ and Cl⁻ ions in a 1:1 ratio, each surrounded by oriented water molecules.
Card 89
Question
Which separation method removes an insoluble solid from a liquid?
Answer
Filtration: the solid stays as residue while the liquid passes as filtrate.
Card 90
Question
What does “like dissolves like” mean at the particle level?
Answer
A solute tends to dissolve when new solute–solvent attractions can compete with the attractions disrupted in the pure substances.
Card 91
Question
What happens when matter absorbs electromagnetic radiation?
Answer
Its particles move to an allowed higher-energy state when the photon energy matches the energy gap.
Card 92
Question
Which equations connect photon energy, frequency, and wavelength?
Answer
E = hν and c = λν.
Card 93
Question
What is the Beer–Lambert law?
Answer
A = εbc: absorbance equals molar absorptivity at the chosen wavelength times path length times concentration.
Card 94
Question
What molecular features generally strengthen London dispersion forces?
Answer
More electrons and a more polarizable cloud strengthen temporary dipoles; greater contact area and accessible π-electron density can also strengthen the attraction.
Card 95
Question
Why do molecular solids usually have low melting points and fail to conduct electricity?
Answer
Distinct molecules are held together by relatively weak intermolecular forces, while their valence electrons stay localized in bonds and lone pairs.
Card 96
Question
How do particles move in a solid?
Answer
They vibrate about fixed positions and do not translate past one another.
Card 97
Question
What graph shapes connect V or P with T(K) or n for an ideal gas?
Answer
All four are straight lines through the origin: V versus T(K) at fixed n and P; P versus T(K) at fixed n and V; V versus n at fixed P and T; and P versus n at fixed V and T.
Card 98
Question
At the same temperature, which gas has the greater average molecular speed: He or Xe?
Answer
He. Both have the same average kinetic energy, but KE = ½mv² means the lower-mass particles move faster.
Card 99
Question
Why do real gases deviate more at high pressure?
Answer
Particles are crowded, so their own volume is no longer negligible compared with the container volume.
Card 100
Question
Which relationship describes dilution when solute amount is conserved?
Answer
M₁V₁ = M₂V₂.
Card 101
Question
Why does an aqueous ionic solution conduct electricity?
Answer
Dissolved ions are mobile and carry charge through the solution.
Card 102
Question
Which property lets simple distillation separate two liquids?
Answer
A sufficient difference in volatility or boiling point, so the vapor is enriched in the more volatile component.
Card 103
Question
Why are many ionic compounds soluble in water but poorly soluble in a nonpolar solvent?
Answer
Water can form strong ion–dipole attractions that stabilize separated ions; a nonpolar solvent cannot provide comparable attractions.
Card 104
Question
Which molecular transition is commonly associated with microwave absorption?
Answer
A transition between quantized rotational energy levels.
Card 105
Question
What frequency corresponds to a 600. nm photon?
Answer
5.00 × 10^14 s^-1. Use ν = c/λ with 600. nm = 6.00 × 10^-7 m.
Card 106
Question
What is the absorbance to two significant figures when ε = 2.0 × 10² L mol^-1 cm^-1, b = 1.00 cm, and c = 0.0020 M?
Answer
0.40. Use A = εbc.
Card 107
Question
What conditions allow hydrogen bonding between two molecules?
Answer
One molecule must donate an H covalently bonded to N, O, or F, and the other must provide a lone pair on N, O, or F. A molecule can be a donor, an acceptor, or both.
Card 108
Question
Why are covalent-network solids often very hard with high melting points?
Answer
A continuous network of strong covalent bonds must be disrupted to deform or melt the solid.
Card 109
Question
Why do a substance's solid and liquid phases usually have similar molar volumes?
Answer
Their particles remain in close contact in both phases, even though liquid particles can move past one another.
Card 110
Question
How is a gas mixture's total pressure related to its component pressures?
Answer
Ptotal = ΣPi; each partial pressure is the pressure that component would exert alone in the same volume and temperature.
Card 111
Question
What microscopic events create gas pressure?
Answer
Gas particles collide with container walls and transfer momentum.
Card 112
Question
Why do intermolecular attractions matter more for gases at low temperature?
Answer
Particles move more slowly, so attractions can alter their paths and promote condensation.
Card 113
Question
What is the final concentration after 50.0 mL of 2.00 M solution is diluted to 200.0 mL?
Answer
0.500 M. Use M₂ = M₁V₁/V₂.
Card 114
Question
What must a particulate representation of a solution communicate?
Answer
The relative concentrations of its components and the particle-level interactions among those components.
Card 115
Question
What causes components to separate in chromatography?
Answer
They differ in attraction to the stationary phase and the mobile phase, so they travel at different rates.
Card 116
Question
Why are many polar molecular solutes soluble in water?
Answer
Dipole attractions or hydrogen bonds with water can replace the solute–solute and water–water attractions disrupted during mixing.
Card 117
Question
Why does an atom produce discrete spectral lines?
Answer
Its electrons can occupy only quantized energy levels, so only photons matching allowed energy differences are absorbed or emitted.
Card 118
Question
How does photon energy change when frequency doubles?
Answer
It doubles because E = hν.
Card 119
Question
Why is a calibration curve useful in spectrophotometry?
Answer
It relates measured absorbance to known concentrations, letting an unknown concentration be read by interpolation within the linear range.
Card 120
Question
How does an ion–dipole attraction form, and how does it compare with dipole–dipole attraction?
Answer
An ion attracts the oppositely charged end of a polar molecule. Ion–dipole attractions tend to be stronger than dipole–dipole attractions.
Card 121
Question
Which solid type is usually both conductive and malleable?
Answer
A metallic solid, because its delocalized electrons move and its nondirectional bonding tolerates layer shifts.
Card 122
Question
How does a crystalline solid differ from an amorphous solid?
Answer
A crystalline solid has long-range repeating order; an amorphous solid lacks that long-range periodic arrangement.
Card 123
Question
How is a gas component's partial pressure found from mole fraction?
Answer
Pi = XiPtotal.
Card 124
Question
How does heating a fixed-volume gas affect its pressure in the ideal model?
Answer
Pressure rises because faster particles collide with the walls more forcefully and frequently.
Card 125
Question
Why can attractions make a real gas's measured pressure lower than the ideal prediction?
Answer
Attractions pull approaching particles away from the walls, reducing momentum transfer during wall collisions.
Card 126
Question
How many moles of ions result from complete dissolution of 0.20 mol CaCl₂?
Answer
0.60 mol ions: 0.20 mol Ca²⁺ plus 0.40 mol Cl⁻.
Card 127
Question
How should water orient around Cl⁻ in a particle model?
Answer
Its partially positive hydrogen ends point toward Cl⁻.
Card 128
Question
Can filtration separate dissolved components of a liquid solution?
Answer
No. Dissolved particles pass through the filter with the solvent; filtration only retains an insoluble solid.
Card 129
Question
Why do nonpolar molecular solutes often dissolve in nonpolar solvents?
Answer
Both rely mainly on compatible London dispersion forces, so mixing can replace the attractions disrupted in the separate substances.
Card 130
Question
What does a shorter absorbed wavelength imply about an energy transition?
Answer
A larger energy gap because E = hc/λ.
Card 131
Question
What is the energy of a photon with frequency 5.0 × 10^14 s^-1?
Answer
3.3 × 10^-19 J. Multiply by Planck's constant: E = (6.626 × 10^-34 J·s)(5.0 × 10^14 s^-1).
Card 132
Question
How does doubling cuvette path length affect absorbance in the linear Beer–Lambert range?
Answer
Absorbance doubles if concentration and molar absorptivity stay constant.
Card 133
Question
How can noncovalent interactions affect a large biomolecule?
Answer
Attractions between molecules or between different regions of the same molecule help set its shape, which strongly affects its properties and function.
Card 134
Question
Why does an ionic solid usually fail to conduct as a solid?
Answer
Its ions are fixed in lattice positions. The same substance conducts when molten or dissolved because the ions can then move.
Card 135
Question
Why does a gas have no definite shape or volume?
Answer
Its widely spaced particles move constantly and experience minimal interparticle attraction, so they spread through the available container.
Card 136
Question
What graph shapes show the inverse pressure–volume relationship for a fixed amount of ideal gas at constant temperature?
Answer
A plot of P against V is a decreasing curve, while P against 1/V is a straight line through the origin.
Card 137
Question
At the same temperature, do different ideal gases have different average kinetic energies?
Answer
No. Average translational kinetic energy depends only on absolute temperature.
Card 138
Question
Under which conditions is ideal-gas behavior most accurate?
Answer
Low pressure and high temperature, where particles are far apart and attractions matter least.
Card 139
Question
What particle-level feature distinguishes a solution from a heterogeneous mixture?
Answer
A solution—whether solid, liquid, or gas—is uniform throughout; a heterogeneous mixture has regions or phases with different compositions.
Card 140
Question
How should water orient around Na⁺ in a particulate model?
Answer
Its partially negative oxygen end points toward Na⁺.
Card 141
Question
In paper chromatography, why does one solute spot travel farther than another?
Answer
It interacts more strongly with the mobile phase or more weakly with the stationary phase. With known phase polarities, that travel difference can reveal relative solute polarity.
Card 142
Question
What energy competition helps explain whether an ionic solid dissolves?
Answer
The energy needed to separate lattice ions competes with the energy released when ion–solvent attractions form.
Card 143
Question
Which molecular motions commonly absorb infrared radiation?
Answer
Bond vibrations whose changing dipole can interact with the radiation.
Card 144
Question
Why must wavelength be converted to meters in c = λν when c is in m s^-1?
Answer
Consistent units are required so meters cancel correctly and frequency comes out in s^-1.
Card 145
Question
How can fingerprints on a cuvette affect a visible-light absorbance reading?
Answer
They can absorb or scatter extra light, making measured absorbance too high and the inferred concentration too high.
Card 146
Question
What causes and controls the strength of dipole–dipole attractions?
Answer
Opposite partial charges on neighboring polar molecules attract. Strength increases with larger molecular dipoles and depends on how favorably the dipoles are oriented.
Card 147
Question
Why is graphite conductive and soft while diamond is insulating and hard?
Answer
Graphite has delocalized electrons within its sheets, so it conducts, and its layers can slide, so it is soft. Diamond has a rigid three-dimensional network of localized covalent bonds, making it hard and insulating.
Card 148
Question
Why are gases much more compressible than liquids?
Answer
Gas particles have large empty spaces between them; liquid particles are already close together.
Card 149
Question
What volume does 0.500 mol CO₂ occupy at 1.00 atm and 300. K if it behaves ideally?
Answer
12.3 L. Use V = nRT/P = (0.500 mol)(0.08206 L atm mol^-1 K^-1)(300. K)/(1.00 atm).
Card 150
Question
Why does a lighter gas effuse faster than a heavier gas at the same temperature?
Answer
Its particles have a higher average speed because equal average kinetic energy is shared by less mass.
Card 151
Question
How does finite particle volume affect a real gas at very high pressure?
Answer
The free volume available for particle motion is smaller than the container volume assumed by the ideal model.
Card 152
Question
How should 250.0 mL of 0.100 M NaCl be prepared from solid NaCl?
Answer
Dissolve 0.0250 mol NaCl, or 1.46 g, then dilute to exactly 250.0 mL in a volumetric flask.
Card 153
Question
What changes in a particle diagram when a solution is diluted without losing solute?
Answer
The solute-particle count stays constant while solvent volume and particle spacing increase.
Card 154
Question
Why is fractional distillation better than simple distillation for liquids with close boiling points?
Answer
Repeated vaporization–condensation steps enrich the vapor in the more volatile component more effectively.
Card 155
Question
Why are oil and water usually immiscible?
Answer
Water's strong hydrogen-bond network isn't replaced by equally strong water–oil attractions, so the substances separate into phases.
Card 156
Question
Which molecular transition is commonly associated with ultraviolet or visible absorption?
Answer
A transition between electronic energy levels.
Card 157
Question
Which photon carries more energy, blue light or red light?
Answer
Blue light, because it has shorter wavelength and higher frequency.
Card 158
Question
Why is absorbance often measured at the wavelength of maximum absorbance in Beer–Lambert analysis?
Answer
It gives the largest concentration-sensitive signal, and the flat top near the maximum makes small wavelength-setting errors less influential.
Card 159
Question
What creates a dipole–induced-dipole attraction, and what controls its strength?
Answer
A permanent dipole distorts a nearby nonpolar particle's electron cloud and creates an attractive temporary dipole. A larger permanent dipole and a more polarizable nonpolar partner make the attraction stronger.
Card 160
Question
How do stronger intermolecular forces affect vapor pressure, boiling point, and melting point?
Answer
They lower vapor pressure and raise boiling point. Melting point often rises too, but the trend is less direct because melting rearranges rather than fully separates particles.
Card 161
Question
How do particles behave in a liquid?
Answer
They stay in close contact while moving and colliding continuously. Temperature and interparticle attractions affect their arrangement and motion.
Card 162
Question
Why must Celsius temperature be converted to kelvins in gas-law calculations?
Answer
Gas-law proportionalities require an absolute temperature scale whose zero corresponds to zero extrapolated thermal motion.
Card 163
Question
How does raising temperature change a Maxwell–Boltzmann speed distribution?
Answer
The distribution broadens, its peak lowers and shifts right, and a larger fraction of particles have high speed.
Card 164
Question
Why does the ideal-gas model treat collisions as elastic?
Answer
It assumes total kinetic energy is conserved in particle–particle and particle–wall collisions.
Card 165
Question
How many moles of solute are in 75.0 mL of a 0.400 M solution?
Answer
0.0300 mol. Multiply 0.400 mol L^-1 by 0.0750 L.
Card 166
Question
For equal solution volumes drawn at the same scale, what shows which solution is more concentrated?
Answer
The more concentrated diagram contains more solute particles in that equal volume.
Card 167
Question
How do differences in intermolecular attractions let distillation separate a liquid solution?
Answer
They give the components different vapor pressures, so the vapor is enriched in the more volatile component.
Card 168
Question
What comparison helps predict whether two liquids will be miscible?
Answer
Liquids with similar types and strengths of intermolecular attractions are more likely to mix uniformly.
Card 169
Question
How can an absorption spectrum help identify a substance?
Answer
Its allowed energy gaps produce a characteristic pattern of absorbed wavelengths that can be compared with known spectra.
Card 170
Question
How does absorbing or emitting a photon change an atom's or molecule's energy?
Answer
Absorption raises the species' energy by exactly the photon energy; emission lowers it by the same amount.
Card 171
Question
What macroscopic evidence can support that a chemical reaction occurred?
Answer
Evidence can include gas formation, precipitate formation, a persistent color change, or an energy change, interpreted with particle-level changes.
Card 172
Question
What does a net ionic equation include?
Answer
Only the dissolved or reacting species that undergo chemical change; spectator ions are omitted.
Card 173
Question
What must a correct particulate reaction diagram conserve?
Answer
The number of atoms of every element and the total charge.
Card 174
Question
What distinguishes a chemical change from a physical change?
Answer
A chemical change rearranges bonds into new substances; a physical change alters state or arrangement without changing chemical identity.
Card 175
Question
What does a balanced equation's coefficient ratio provide?
Answer
The mole ratio among reacting and produced species.
Card 176
Question
What is the equivalence point of a titration?
Answer
The point where titrant and analyte have reacted in the stoichiometric ratio given by the balanced equation.
Card 177
Question
What defines a precipitation reaction?
Answer
Aqueous ions combine to form a sparingly soluble solid.
Card 178
Question
What happens in a Brønsted–Lowry acid–base reaction?
Answer
A proton transfers from the acid (donor) to the base (acceptor). In aqueous solution, H₂O can play either role.
Card 179
Question
What does oxidation mean in a redox reaction?
Answer
Loss of electrons and an increase in oxidation number.
Card 180
Question
What particle-level change confirms that a process is chemical?
Answer
Atoms rearrange into new combinations, producing substances with different compositions.
Card 181
Question
Which ions are spectators when AgNO₃(aq) reacts with NaCl(aq)?
Answer
Na⁺ and NO₃⁻. The net ionic reaction is Ag⁺(aq) + Cl⁻(aq) → AgCl(s).
Card 182
Question
How does a particulate diagram reveal the limiting reactant?
Answer
After forming the maximum product allowed by the ratio, none of the limiting reactant remains while excess reactant particles do.
Card 183
Question
Is melting ice a chemical or physical change?
Answer
A physical change. H₂O molecules remain H₂O while their arrangement and motion change.
Card 184
Question
How is the limiting reactant identified from given amounts?
Answer
Convert each reactant to the same product amount using the balanced equation; the smaller product amount identifies the limiting reactant.
Card 185
Question
How does an endpoint differ from an equivalence point?
Answer
The endpoint is an observed signal such as indicator color change; the equivalence point is the exact stoichiometric condition.
Card 186
Question
How is complete combustion of a hydrocarbon in excess oxygen classified, and what products form?
Answer
It is a redox combustion reaction that forms CO₂ and H₂O.
Card 187
Question
What are the conjugate acid and conjugate base in NH₃ + H₂O ⇌ NH₄⁺ + OH⁻?
Answer
NH₄⁺ is the conjugate acid of NH₃, and OH⁻ is the conjugate base of H₂O.
Card 188
Question
What does reduction mean in a redox reaction?
Answer
Gain of electrons and a decrease in oxidation number.
Card 189
Question
Which common changes are physical rather than chemical?
Answer
Phase changes and the formation or separation of mixtures are physical when each substance keeps its composition.
Card 190
Question
How are strong soluble electrolytes written in a complete ionic equation?
Answer
As separated aqueous ions; solids, liquids, gases, and weak electrolytes stay intact.
Card 191
Question
A diagram starts with six A particles and four B₂ particles for 2A + B₂ → 2AB; what remains after completion?
Answer
One B₂ remains. Six A consume three B₂ and form six AB.
Card 192
Question
Why is dissolving NaCl in water normally classified as a physical change?
Answer
Na⁺ and Cl⁻ separate and become hydrated, but retain their chemical identities. Removing the water recovers NaCl; the shift from ion–ion to ion–dipole attractions does not by itself form a new substance.
Card 193
Question
What mass of AgCl can form from 25.0 mL of 0.200 M AgNO₃ mixed with excess Cl⁻?
Answer
0.717 g AgCl. The 1:1 reaction gives 0.00500 mol AgCl; multiply by 143.32 g mol^-1.
Card 194
Question
What calculation finds unknown analyte moles at equivalence?
Answer
Use titrant moles, n = MV, then apply the balanced-reaction mole ratio.
Card 195
Question
Which feature identifies an acid–base, redox, or precipitation reaction?
Answer
Acid–base reactions transfer protons, redox reactions change oxidation numbers through electron transfer, and precipitation reactions form a sparingly soluble solid.
Card 196
Question
What is the net ionic equation for strong acid–strong base neutralization?
Answer
H⁺(aq) + OH⁻(aq) → H₂O(l).
Card 197
Question
What is the oxidation number of sulfur in SO₄²⁻?
Answer
+6. Four oxygens contribute -8 total, so sulfur must be +6 to give -2 overall.
Card 198
Question
Why can gas bubbles alone be ambiguous evidence of reaction?
Answer
Bubbles may also come from boiling or dissolved gas escaping, so the context and particle identities must support a chemical change.
Card 199
Question
How is melting ice represented as a balanced physical-change equation?
Answer
H₂O(s) → H₂O(l). The formula and atom count stay the same because only the physical state changes.
Card 200
Question
What does a particle diagram show when no reaction occurs after two aqueous ionic solutions mix?
Answer
All ions remain separated and solvated, with no new bonded particles, precipitate, or gas.
Card 201
Question
Why is rusting iron a chemical change?
Answer
Iron atoms form new iron-oxide substances through electron transfer and new bonding.
Card 202
Question
For 2H₂O₂(aq) → 2H₂O(l) + O₂(g), what volume of O₂ forms from 0.100 mol H₂O₂ at 298 K and 1.00 atm?
Answer
1.22 L O₂. The mole ratio gives 0.0500 mol O₂, then V = nRT/P.
Card 203
Question
A 25.0 mL monoprotic acid sample requires 20.0 mL of 0.150 M NaOH; what is the acid concentration?
Answer
0.120 M. At 1:1 equivalence, moles acid = 0.0200 L × 0.150 M, then divide by 0.0250 L.
Card 204
Question
Which salts does the minimum solubility rule in this deck treat as soluble?
Answer
All salts containing Na⁺, K⁺, NH₄⁺, or NO₃⁻ are treated as soluble in water.
Card 205
Question
How are the strengths of a conjugate acid and its conjugate base related?
Answer
A stronger acid has a weaker conjugate base, and a stronger base has a weaker conjugate acid.
Card 206
Question
How are oxidation and reduction half-reactions combined into one balanced equation?
Answer
Multiply them so electrons lost equal electrons gained, add the half-reactions, then cancel electrons and any identical species on both sides.
Card 207
Question
How do molecular, complete ionic, and net ionic equations differ?
Answer
Molecular equations keep compounds intact, complete ionic equations split strong soluble electrolytes, and net ionic equations remove spectators. All three conserve atoms and charge.
Card 208
Question
How should coefficients change particle counts in a reaction diagram?
Answer
They set whole-particle ratios while preserving each particle's chemical formula.
Card 209
Question
Is separating a mixture by distillation a chemical or physical change?
Answer
A physical change. Components change phase and location but keep their chemical identities.
Card 210
Question
What equation results from Cu → Cu²⁺ + 2e⁻ and Ag⁺ + e⁻ → Ag?
Answer
Cu + 2Ag⁺ → Cu²⁺ + 2Ag. Multiply the silver half-reaction by 2 and cancel 2e⁻; both atom counts and net charge then match.
Card 211
Question
How is average reaction rate found from a reactant concentration?
Answer
Use the negative concentration change divided by elapsed time, adjusted by its stoichiometric coefficient when comparing species rates.
Card 212
Question
What does a rate law express?
Answer
It shows how the measured rate depends on reactant concentrations. In rate = k[A]^m[B]^n, m and n are the orders in A and B, and m + n is the overall order.
Card 213
Question
A plot of ln[A] versus time is linear; what is the order in A and its integrated rate law?
Answer
First order: ln[A]t = ln[A]0 − kt, so the plot's slope is −k.
Card 214
Question
What is an elementary reaction?
Answer
A single step in a mechanism whose rate law follows directly from its reactant molecularity.
Card 215
Question
What two collision conditions are needed for reaction?
Answer
Sufficient collision energy and a productive molecular orientation.
Card 216
Question
What does activation energy represent on a reaction-energy profile?
Answer
The energy difference from the reactants to the transition state. The reaction coordinate tracks the step's structural progress, not elapsed time.
Card 217
Question
What must the elementary steps of a valid mechanism do when added?
Answer
Cancel intermediates and reproduce the overall balanced reaction.
Card 218
Question
How is a proposed mechanism tested against kinetics?
Answer
Its derived rate law must agree with the experimentally measured rate law.
Card 219
Question
What does a pre-equilibrium approximation assume?
Answer
A fast reversible step reaches equilibrium before a later slow step consumes its intermediate.
Card 220
Question
What does each peak on a multistep energy profile represent?
Answer
A transition state for one elementary step.
Card 221
Question
How does a catalyst increase reaction rate?
Answer
It provides an alternate mechanism with a lower activation-energy pathway.
Card 222
Question
Why does crushing a solid reactant usually increase its reaction rate?
Answer
Crushing increases exposed surface area, so more reactant particles can collide with the other reactant each second.
Card 223
Question
How is reaction order found from initial-rate data?
Answer
Compare trials where one reactant concentration changes while the others stay constant, then match the rate factor to the concentration factor.
Card 224
Question
A plot of [A] versus time is linear; what is the order in A and its integrated rate law?
Answer
Zero order: [A]t = [A]0 − kt, so the plot's slope is −k.
Card 225
Question
What is the rate law for the elementary step 2A + B → products?
Answer
rate = k[A]²[B]. This inference is valid because the step is elementary.
450 cards
Advanced High School Chemistry Flashcards: Complete 9-Part Course Review
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Card 226
Question
How does raising temperature change a Maxwell–Boltzmann energy distribution and reaction rate?
Answer
The distribution shifts and broadens toward higher energies, so a larger fraction of collisions exceeds the activation-energy threshold and can react.
Card 227
Question
How is ΔH read from a reaction-energy profile?
Answer
ΔH = energy of products − energy of reactants.
Card 228
Question
What is a reaction intermediate?
Answer
A species formed in one mechanism step and consumed in a later step, so it cancels from the overall equation.
Card 229
Question
Why can't overall reaction coefficients usually supply rate-law exponents?
Answer
The overall equation hides the mechanism; exponents come from experiment unless the reaction is a stated elementary step.
Card 230
Question
How does pre-equilibrium remove an intermediate from a rate law?
Answer
Use the fast-step equilibrium relation to express the intermediate concentration in terms of stable reactants.
Card 231
Question
What does each valley between peaks represent on a multistep profile?
Answer
A reaction intermediate.
Card 232
Question
Does a catalyst change ΔH or the equilibrium constant?
Answer
No. It changes the pathway and rates, not reactant/product energies or the equilibrium composition.
Card 233
Question
For 2A → B, how are disappearance of A and appearance of B related?
Answer
Reaction rate = -(1/2)Δ[A]/Δt = Δ[B]/Δt.
Card 234
Question
How do the units of k depend on a rate law's overall order?
Answer
They must make the rate unit M s^-1: zero order uses M s^-1, first order s^-1, and second order M^-1 s^-1.
Card 235
Question
A plot of 1/[A] versus time is linear; what is the order in A and its integrated rate law?
Answer
Second order: 1/[A]t = 1/[A]0 + kt, so the plot's slope is +k.
Card 236
Question
What is molecularity?
Answer
The number of reacting particles in an elementary step, such as unimolecular or bimolecular.
Card 237
Question
How does raising temperature affect k in the qualitative Arrhenius model?
Answer
k increases, often sharply, because a larger fraction of collisions can reach the transition state. Arrhenius-equation calculations are outside this deck’s scope.
Card 238
Question
A reactant falls from 0.80 M to 0.50 M in 30. s; what is its average disappearance rate to two significant figures?
Answer
0.010 M s^-1. Use -(0.50 − 0.80) M ÷ 30. s.
Card 239
Question
How does a catalyst differ from an intermediate in a mechanism?
Answer
A catalyst is consumed early and regenerated later; an intermediate is formed early and consumed later.
Card 240
Question
For 2NO₂ → NO₃ + NO (slow), followed by NO₃ + CO → NO₂ + CO₂ (fast), what rate law is predicted?
Answer
rate = k[NO₂]². The first step is elementary and rate-limiting, so its molecularity sets the observed rate law.
Card 241
Question
On a multistep reaction-energy profile, which feature often identifies the rate-determining step?
Answer
The step with the largest activation barrier measured from its preceding valley to its peak.
Card 242
Question
If changing [B] leaves rate unchanged, what is the order in B?
Answer
Zero order, so [B]^0 = 1 in the measured rate law.
Card 243
Question
What mechanism changes can binding, acid–base, or surface catalysis introduce?
Answer
They can orient reactants, lower activation barriers, or create new bound, protonated, or deprotonated intermediates and elementary steps; the catalyst is regenerated.
Card 244
Question
What is special about a first-order reaction's half-life?
Answer
It is constant and independent of starting concentration: t1/2 = ln 2/k. Radioactive decay is a common first-order example.
Card 245
Question
Why is a termolecular elementary collision uncommon?
Answer
Three particles must collide simultaneously with suitable energy and orientation, which is much less probable than one- or two-particle events.
Card 246
Question
On a reaction-energy profile, how are reverse activation energy, forward activation energy, and ΔH related?
Answer
Ea,reverse = Ea,forward − ΔH. The reverse barrier is measured from products to the same transition state.
Card 247
Question
Why can correct orientation matter even above the activation energy?
Answer
The colliding reactive sites must align so old bonds can break and new bonds can form along the reaction pathway.
Card 248
Question
How does detecting a proposed reaction intermediate affect a mechanism claim?
Answer
It supports a mechanism that contains that intermediate, but it doesn't prove that mechanism is unique.
Card 249
Question
For 2NO ⇌ N₂O₂ (fast equilibrium), followed by N₂O₂ + O₂ → 2NO₂ (slow), what observed rate law results?
Answer
rate = kobs[NO]²[O₂]. Start with rate = k₂[N₂O₂][O₂], use [N₂O₂] = K[NO]² from the fast equilibrium, then substitute.
Card 250
Question
What does the highest point of a one-step energy profile represent?
Answer
The transition state, an unstable arrangement at the top of the activation barrier.
Card 251
Question
What sign does q have for an endothermic system?
Answer
Positive, because the system absorbs heat from the surroundings.
Card 252
Question
How does an exothermic reaction appear on an enthalpy diagram?
Answer
Products lie below reactants, so ΔH is negative.
Card 253
Question
What condition defines thermal equilibrium?
Answer
Objects in contact have the same temperature, so there is no net heat transfer.
Card 254
Question
What equations relate heat capacity and temperature change to heat transfer?
Answer
Use q = mcΔT with specific heat capacity, or q = nCₘΔT with molar heat capacity.
Card 255
Question
Why is temperature constant during a phase-change plateau?
Answer
Added or removed energy changes interparticle potential energy instead of average kinetic energy.
Card 256
Question
What does ΔHrxn describe?
Answer
The heat absorbed or released at constant pressure for the reaction exactly as written under the stated conditions.
Card 257
Question
How is reaction enthalpy estimated from average bond enthalpies?
Answer
ΔHrxn ≈ Σ(bonds broken) − Σ(bonds formed).
Card 258
Question
What is the standard enthalpy of formation of an element in its standard state?
Answer
Zero by definition.
Card 259
Question
In a Hess’s law calculation, how should a step change when the target needs twice its reverse?
Answer
Reverse the equation, double every coefficient, and multiply its ΔH by -2.
Card 260
Question
How can energy cross a system boundary during a process?
Answer
As heat or work. Heat transferred to or work done on the system increases its energy; heat transferred from or work done by the system decreases it.
Card 261
Question
How does an endothermic reaction appear on an enthalpy diagram?
Answer
Products lie above reactants, so ΔH is positive.
Card 262
Question
How are heat gained by a system and heat lost by its surroundings related in an isolated setup?
Answer
qsystem = -qsurroundings.
Card 263
Question
In coffee-cup calorimetry, how is reaction heat related to solution heat?
Answer
qrxn = -qsolution when calorimeter heat is negligible and pressure is constant.
Card 264
Question
What heat is required to melt n moles at the melting point?
Answer
q = nΔHfus.
Card 265
Question
How does reversing a reaction change ΔH?
Answer
It reverses the sign of ΔH.
Card 266
Question
Why is breaking a bond endothermic?
Answer
Energy must be supplied to separate atoms against their bonding attraction.
Card 267
Question
How is ΔH°rxn calculated from standard enthalpies of formation?
Answer
ΣνΔHf°(products) − ΣνΔHf°(reactants).
Card 268
Question
How does multiplying an equation by 3 affect its ΔH?
Answer
Multiply ΔH by 3 because enthalpy change scales with reaction amount.
Card 269
Question
Why can an exothermic dissolution warm the solution?
Answer
The solution warms because forming solute–solvent attractions releases more energy than is absorbed in separating the original particles. The net potential-energy decrease raises particle kinetic energy and temperature.
Card 270
Question
Does an energy diagram's activation barrier determine ΔH?
Answer
No. ΔH depends on reactant and product energy levels, while the barrier controls kinetics.
Card 271
Question
Why does heat flow from a warmer object to a cooler object?
Answer
Energy transfers through collisions until their average kinetic energies, and therefore temperatures, equalize.
Card 272
Question
How much heat warms 100.0 g of water by 5.0°C?
Answer
2.1 kJ. Use q = (100.0 g)(4.184 J g^-1 °C^-1)(5.0°C).
Card 273
Question
How are the molar enthalpies of a phase change and its reverse related?
Answer
They have equal magnitudes and opposite signs, such as ΔHcond = -ΔHvap and ΔHfreeze = -ΔHfus.
Card 274
Question
How does doubling every coefficient in a thermochemical equation affect ΔH?
Answer
It doubles ΔH.
Card 275
Question
Why is forming a bond exothermic?
Answer
Atoms move to a lower-potential-energy bonded arrangement and release energy.
Card 276
Question
What formation equation defines ΔHf° for CO₂(g)?
Answer
C(s, graphite) + O₂(g) → CO₂(g), forming exactly one mole from elements in standard states.
Card 277
Question
What should happen to intermediate species when equations in a Hess’s law calculation are added?
Answer
They cancel, leaving the target overall reaction.
Card 278
Question
If the surroundings warm during a process, what is the likely sign of qsystem?
Answer
Negative; the system likely released heat to the surroundings.
Card 279
Question
For a profile with reactants at 40 kJ and products at 10 kJ, what is ΔH?
Answer
-30 kJ for the reaction as drawn.
Card 280
Question
Assuming no phase change, what determines the final temperature when two substances exchange heat in an insulated container?
Answer
Energy conservation: q_warm + q_cool = 0. Use each substance's mass, heat capacity, and initial temperature to solve for the common final temperature.
Card 281
Question
How would heat loss to the room affect an exothermic calorimetry result?
Answer
The observed temperature rise is too small, so the calculated magnitude of released heat is too low.
Card 282
Question
What heat expression covers warming a liquid without a phase change?
Answer
q = mcΔT, not nΔHphase.
Card 283
Question
If forming 1 mol of product has ΔH = -50 kJ mol^-1, what is q when 2 mol forms?
Answer
-100 kJ. Use q = nΔH = (2 mol)(-50 kJ mol^-1).
Card 284
Question
Breaking reactant bonds requires 500 kJ, and forming product bonds releases 650 kJ; what is the estimated ΔH?
Answer
-150 kJ, from 500 − 650.
Card 285
Question
For CO(g) + ½O₂(g) → CO₂(g), what is ΔH°rxn if ΔHf°[CO] = -110.5 and ΔHf°[CO₂] = -393.5 kJ mol^-1?
Answer
-283.0 kJ. Use -393.5 - [-110.5 + ½(0)], since ΔHf°[O₂(g)] = 0.
Card 286
Question
In a Hess’s law calculation, two valid steps have ΔH values +25 kJ and -60 kJ; what is the combined ΔH?
Answer
-35 kJ, provided the equations add to the target reaction.
Card 287
Question
Why is “bonds breaking releases energy” incorrect?
Answer
Bond breaking absorbs energy; the overall reaction releases energy only when forming new bonds releases more than breaking old bonds requires.
Card 288
Question
How would melting appear on an energy diagram?
Answer
The liquid lies above the solid, so ΔHfus is positive; the diagram represents a physical, endothermic change.
Card 289
Question
Can two objects at the same temperature exchange energy microscopically?
Answer
Yes, but their energy transfers balance, so there is no net heat flow.
Card 290
Question
Why must the calorimeter's heat capacity be included when it isn't negligible?
Answer
The apparatus can absorb or release heat, so include q_cal = C_calΔT in the energy balance: q_process + q_solution + q_cal = 0.
Card 291
Question
What makes chemical equilibrium dynamic?
Answer
Forward and reverse reactions continue at equal rates even though macroscopic concentrations stay constant.
Card 292
Question
For aA + bB ⇌ cC, what is the concentration-form expression for Q?
Answer
Q = [C]^c / ([A]^a[B]^b), using current rather than necessarily equilibrium concentrations.
Card 293
Question
What does K much greater than 1 indicate?
Answer
Products predominate at equilibrium, though K says nothing about reaction speed.
Card 294
Question
How does reversing a reaction change its equilibrium constant?
Answer
K becomes 1/K.
Card 295
Question
Can a reversible system reach equilibrium when it starts with only products?
Answer
Yes, if the reverse reaction is possible. The equilibrium composition depends on temperature, initial amounts, and volume or pressure.
Card 296
Question
How do Q and K predict reaction direction?
Answer
Q < K shifts forward, Q > K shifts reverse, and Q = K means equilibrium.
Card 297
Question
Which species are omitted from a heterogeneous equilibrium expression?
Answer
Pure solids and pure liquids because their activities are effectively constant.
Card 298
Question
How does increasing a dissolved reactant's concentration or a gaseous reactant's partial pressure affect equilibrium at constant temperature when other Q terms are initially unchanged?
Answer
It lowers Q relative to K, so the system shifts toward products until Q = K again. Changing the amount of a pure solid or liquid omitted from Q does not cause this shift while that pure phase remains present.
Card 299
Question
What does a flat concentration-time graph mean at equilibrium?
Answer
Each concentration is constant, not necessarily equal to the others.
Card 300
Question
For A ⇌ B in one fixed volume, a particulate model shows 16 A and 0 B initially, then 4 A and 12 B at equilibrium. What changed, what predominates, and what is Kc?
Answer
The net change was forward: 12 A particles became 12 B particles. B predominates at equilibrium, and Kc = [B]/[A] = 12/4 = 3.0 because both counts come from the same fixed volume.
Card 301
Question
What can Ksp tell you about a salt's solubility, and when can two Ksp values be compared directly?
Answer
Ksp > 1 indicates a soluble salt. For salts with the same dissolution stoichiometry, a larger Ksp generally means greater molar solubility; across different stoichiometries, calculate molar solubility before comparing.
Card 302
Question
What is the common-ion effect on solubility?
Answer
Adding an ion already in the dissolution equilibrium usually decreases the solid's molar solubility.
Card 303
Question
How does uniform dilution shift an aqueous equilibrium based on the stoichiometric powers in Q?
Answer
It shifts toward the side with the larger sum of stoichiometric coefficients for dissolved species included in Q. If the sums are equal, dilution causes no shift by this effect; pure solids and liquids remain omitted.
Card 304
Question
What happens if a reversible reaction starts with reactants only?
Answer
The forward rate is initially largest; products form, the reverse rate grows, and the rates eventually become equal.
Card 305
Question
What is the purpose of an ICE table?
Answer
To organize initial, change, and equilibrium concentrations using reaction stoichiometry.
Card 306
Question
Can a reaction with a very large K be slow?
Answer
Yes. K describes thermodynamic equilibrium position, while rate depends on kinetics and activation energy.
Card 307
Question
What happens to Q immediately after product concentration increases?
Answer
Q increases; if it rises above K, the reaction shifts toward reactants.
Card 308
Question
How does multiplying every reaction coefficient by 2 affect K?
Answer
The new equilibrium constant is K².
Card 309
Question
For A ⇌ B, Kc = 4.0 and initially [A] = 1.0 M and [B] = 0, what are the equilibrium concentrations?
Answer
[A] = 0.20 M and [B] = 0.80 M. Let x form: Kc = x/(1.0 − x) = 4.0, so x = 0.80 M.
Card 310
Question
What macroscopic properties stay constant at equilibrium?
Answer
Properties such as concentration, color, and pressure remain constant when external conditions are fixed.
Card 311
Question
How does decreasing volume shift a gaseous equilibrium?
Answer
Toward the side with fewer moles of gas, if the two sides have different gaseous mole counts.
Card 312
Question
For N₂ + 3H₂ ⇌ 2NH₃, what is Kc?
Answer
Kc = [NH₃]² / ([N₂][H₂]³).
Card 313
Question
For CaF₂(s) ⇌ Ca²⁺ + 2F⁻, how is Ksp written in terms of molar solubility s in pure water?
Answer
Ksp = s(2s)² = 4s³ because [Ca²⁺] = s and [F⁻] = 2s.
Card 314
Question
What does K much less than 1 indicate?
Answer
Reactants predominate at equilibrium.
Card 315
Question
How does decreasing a dissolved product's concentration or a gaseous product's partial pressure affect equilibrium when other Q terms are initially unchanged?
Answer
It lowers Q and drives a net forward reaction until equilibrium returns. Changing the amount of a pure solid or liquid omitted from Q does not cause this shift while that phase remains.
Card 316
Question
Does equilibrium mean the reaction has stopped?
Answer
No. Both directions continue, but equal rates produce no net macroscopic change.
Card 317
Question
Why does adding NaF reduce CaF₂ solubility?
Answer
The added F⁻ raises Qsp, shifting the dissolution equilibrium toward solid CaF₂.
Card 318
Question
For N₂ + 3H₂ ⇌ 2NH₃, what is Kp when P_N₂ = 0.50 atm, P_H₂ = 1.50 atm, and P_NH₃ = 0.25 atm?
Answer
0.037. Use Kp = (P_NH₃)²/[(P_N₂)(P_H₂)³] = (0.25)²/[(0.50)(1.50)³]. Use equilibrium partial pressures directly; Kc↔Kp conversion isn't assessed.
Card 319
Question
What happens to Q when a gaseous equilibrium mixture is compressed at constant temperature if products have fewer gas moles?
Answer
Q falls relative to K, so the reaction shifts toward products.
Card 320
Question
How do K and Q transform when a reaction is reversed, its coefficients are multiplied, or reactions are added?
Answer
They follow the same algebra: reversing takes the reciprocal, multiplying every coefficient by c raises the value to the power c, and adding reactions multiplies their K or Q values.
Card 321
Question
When is the small-x approximation acceptable?
Answer
When x is small relative to the initial concentration and the final result confirms the neglected change is suitably small.
Card 322
Question
What graph feature shows a disturbance followed by re-equilibration?
Answer
A sudden or gradual concentration change followed by new constant plateaus while rates return to equality.
Card 323
Question
If Q = 0.20 and K = 5.0, which direction is favored next?
Answer
Forward, because Q < K.
Card 324
Question
At equilibrium, are reactant and product concentrations equal?
Answer
Not necessarily. They are constant, while forward and reverse rates are equal.
Card 325
Question
CaF₂ has Ksp = 3.2 × 10^-11 in pure water; what is its molar solubility?
Answer
2.0 × 10^-4 M. If the molar solubility is s, then [Ca²⁺] = s, [F⁻] = 2s, and Ksp = 4s³.
Card 326
Question
For N₂ + 3H₂ ⇌ 2NH₃, how is Qp written?
Answer
Qp = (P_NH₃)²/[(P_N₂)(P_H₂)³], using the current partial pressures rather than necessarily equilibrium values.
Card 327
Question
How does heating shift an endothermic forward reaction?
Answer
Toward products, and K increases because temperature changes the equilibrium constant.
Card 328
Question
Why do both forward and reverse rates change as equilibrium is approached?
Answer
As reactant and product concentrations change, the collision frequencies for the two directions change until their rates match.
Card 329
Question
CaF₂ has Ksp = 3.2 × 10^-11. What is its molar solubility in 0.10 M NaF?
Answer
About 3.2 × 10^-9 M. With [F⁻] ≈ 0.10 M, Ksp = [Ca²⁺][F⁻]² gives s = (3.2 × 10^-11)/(0.10)². The common ion lowers solubility but does not change Ksp at constant temperature.
Card 330
Question
What concentration data must be used to calculate Kc?
Answer
Equilibrium concentrations, each raised to its stoichiometric coefficient and excluding pure solids and liquids.
Card 331
Question
What is a Brønsted–Lowry acid?
Answer
A proton donor.
Card 332
Question
How is pH defined?
Answer
pH = -log[H₃O⁺].
Card 333
Question
What is Ka for HA + H₂O ⇌ H₃O⁺ + A⁻?
Answer
Ka = [H₃O⁺][A⁻]/[HA].
Card 334
Question
How does stabilizing a base affect its basicity and the strength of its conjugate acid?
Answer
It makes the base weaker and its conjugate acid stronger. A more stable base is less willing to accept H⁺.
Card 335
Question
What is a Brønsted–Lowry base?
Answer
A proton acceptor.
Card 336
Question
At 25°C, what are Kw and the relationship between pH and pOH?
Answer
Kw = [H₃O⁺][OH⁻] = 1.0 × 10^-14. Taking negative logarithms gives pH + pOH = 14.00.
Card 337
Question
What is Kb for B + H₂O ⇌ BH⁺ + OH⁻?
Answer
Kb = [BH⁺][OH⁻]/[B].
Card 338
Question
Why can lowering pH increase the solubility of a salt containing a basic anion?
Answer
H₃O⁺ consumes the anion, pulling the dissolution equilibrium toward more dissolved ions.
Card 339
Question
What are conjugate acid–base pairs?
Answer
Species that differ by exactly one proton.
Card 340
Question
What is the pH of 1.0 × 10^-3 M HCl?
Answer
3.00, assuming complete dissociation and negligible water contribution.
Card 341
Question
How are pKa and pKb defined?
Answer
pKa = -log Ka, and pKb = -log Kb.
Card 342
Question
Why does acid strength increase across a row of comparable hydrides?
Answer
Increasing electronegativity stabilizes the conjugate base and polarizes the H–A bond.
Card 343
Question
What is an amphiprotic species?
Answer
A species that can donate or accept a proton, such as HCO₃⁻.
Card 344
Question
What amounts remain after a limited amount of strong base partially neutralizes weak acid HA?
Answer
Subtract the reacted moles from HA and form the same number of moles of A⁻. The result gives the remaining HA and formed A⁻ amounts before any equilibrium or buffer-pH calculation.
Card 345
Question
How are Ka, Kb, pKa, and pKb related for a conjugate pair at 25°C?
Answer
KaKb = Kw = 1.0 × 10^-14, and pKa + pKb = pKw = 14.00.
Card 346
Question
When does pH have little effect on a salt's solubility?
Answer
When neither dissolved ion reacts appreciably with H₃O⁺ or OH⁻.
Card 347
Question
How does H₂O act in HCl + H₂O → H₃O⁺ + Cl⁻ and in NH₃ + H₂O ⇌ NH₄⁺ + OH⁻?
Answer
It acts as a base in the first reaction by accepting H⁺, and as an acid in the second by donating H⁺.
Card 348
Question
After mixing weak base B with strong acid, what controls the final solution in the three stoichiometric regimes?
Answer
Excess B leaves a B/BH⁺ buffer; equimolar amounts leave BH⁺, so the solution is acidic; excess strong acid sets the pH from the remaining H₃O⁺.
Card 349
Question
What two components make a typical weak-acid buffer?
Answer
A weak acid and a significant amount of its conjugate base.
Card 350
Question
What do the successive half-equivalence pH values approximate in a diprotic weak-acid titration?
Answer
The first approximates pKa₁ and the second approximates pKa₂ because each conjugate pair has equal concentrations at its half-equivalence point.
Card 351
Question
Which acid is stronger, one with pKa 2 or pKa 5?
Answer
The acid with pKa 2; lower pKa means larger Ka.
Card 352
Question
What is the Henderson–Hasselbalch equation?
Answer
pH = pKa + log([A⁻]/[HA]).
Card 353
Question
Why are larger binary hydrides down a group often stronger acids?
Answer
The H–A bond becomes weaker as the central atom grows, so proton release is easier.
Card 354
Question
What mainly determines buffer capacity?
Answer
The concentrations of both members of the conjugate acid–base pair. Increasing both concentrations at a fixed ratio increases capacity without changing pH; capacity is best balanced for added acid and base when their concentrations are similar.
Card 355
Question
Why does acid increase CaCO₃ solubility?
Answer
H₃O⁺ converts CO₃²⁻ to HCO₃⁻ or carbonic acid species, reducing free carbonate and driving more CaCO₃ to dissolve.
Card 356
Question
What does pH < pKa imply for a weak-acid pair?
Answer
The protonated form HA predominates over A⁻.
Card 357
Question
What happens when stoichiometrically equal amounts of a monoprotic weak acid and strong base are mixed?
Answer
The weak acid is consumed to its conjugate base; at equivalence, the solution isn't a buffer containing both forms.
Card 358
Question
What is [H₃O⁺] when pH = 4.50?
Answer
3.2 × 10^-5 M, from [H₃O⁺] = 10^-pH.
Card 359
Question
What is the pH of 0.010 M Ba(OH)₂ at 25°C?
Answer
About 12.30. Complete dissociation gives [OH⁻] = 0.020 M, so pOH = 1.70. At 25°C, pH + pOH = 14.00, so pH = 12.30.
Card 360
Question
How does a buffer respond to a small amount of added strong acid?
Answer
Its conjugate base consumes H⁺, converting to the weak acid and limiting the pH change.
Card 361
Question
Why is the equivalence-point solution basic in a monoprotic weak-acid–strong-base titration?
Answer
The conjugate base produced at equivalence reacts with water to form OH⁻, so the pH is above neutral—above 7.00 at 25°C.
Card 362
Question
How is percent ionization calculated for a weak acid or weak base?
Answer
For HA, use ([H₃O⁺]equilibrium ÷ [HA]initial) × 100%. For B, use ([BH⁺]equilibrium ÷ [B]initial) × 100%, under the usual monoprotic setup.
Card 363
Question
When is Henderson–Hasselbalch useful for an initial buffer-pH calculation?
Answer
Use it when both members of a conjugate acid–base pair are present in meaningful amounts, including after in-scope stoichiometry creates a buffer. Calculating the pH change after acid or base is added to an existing buffer is outside this deck’s scope.
Card 364
Question
Why does adding oxygen atoms usually strengthen oxyacids with the same central atom?
Answer
Extra oxygens withdraw electron density and delocalize negative charge in the conjugate base.
Card 365
Question
A prepared buffer is accidentally diluted to twice its intended volume; what happens to its pH and capacity?
Answer
Its pH stays nearly the same, and its capacity per liter is halved because both component concentrations halve. The total neutralizing moles in the sample remain unchanged.
Card 366
Question
How does adding OH⁻ affect Mg(OH)₂ solubility?
Answer
It decreases solubility through the common-ion effect, shifting Mg(OH)₂(s) ⇌ Mg²⁺ + 2OH⁻ toward the solid.
Card 367
Question
A buffer has equal [A⁻] and [HA]; what is its pH?
Answer
pH = pKa because log(1) = 0.
Card 368
Question
How should a weak acid–strong base mixture be solved before equivalence?
Answer
First use mole stoichiometry; if both HA and A⁻ remain, use the resulting buffer relation.
Card 369
Question
Why can pure neutral water have a pH other than 7.00?
Answer
Kw changes with temperature. Neutrality means [H₃O⁺] = [OH⁻], while pH = 7.00 only when Kw = 1.0 × 10^-14 at 25°C.
Card 370
Question
25.0 mL of 0.200 M HCl is diluted to 100.0 mL; what is the pH?
Answer
1.301. Dilution gives [H₃O⁺] = (0.200 M)(25.0 mL)/(100.0 mL) = 0.0500 M, so pH = -log(0.0500).
Card 371
Question
How does a buffer respond to a small amount of added strong base?
Answer
The weak acid consumes OH⁻, forming conjugate base and water.
Card 372
Question
How do you find the final pH after mixing a strong acid and strong base at 25°C?
Answer
Use H₃O⁺ + OH⁻ → 2H₂O and compare their moles. Divide excess H₃O⁺ or OH⁻ by the total volume, then calculate pH or pOH from that excess concentration. Equal moles give pH 7.00 at 25°C.
Card 373
Question
What distinguishes acid strength from acid concentration?
Answer
Strength is the equilibrium tendency to donate H⁺, reflected by Ka or pKa; concentration is the amount of acid per solution volume.
Card 374
Question
If [A⁻]/[HA] = 10, how does pH compare with pKa?
Answer
pH = pKa + 1 because log 10 = 1.
Card 375
Question
Which conjugate base is more stable, one with localized or resonance-delocalized charge?
Answer
The resonance-delocalized conjugate base, which generally corresponds to the stronger acid.
Card 376
Question
Which 1.0 L buffer has greater capacity: 1.0 mol each of HA/A⁻ or 0.10 mol each at the same ratio?
Answer
The 1.0 mol pair; both have the same initial pH, but the larger amounts neutralize more added acid or base.
Card 377
Question
For BHX(s) ⇌ BH⁺ + X⁻, why can raising pH increase the salt's solubility?
Answer
OH⁻ consumes BH⁺ to form B and H₂O, so dissolution shifts right to replace BH⁺. This is a qualitative prediction, not a pH-dependent solubility calculation.
Card 378
Question
What does pH > pKa imply for a weak-acid pair?
Answer
The deprotonated form A⁻ predominates over HA.
Card 379
Question
For HA + B ⇌ A⁻ + BH⁺, which side is favored when pKa(HA) = 4 and pKa(BH⁺) = 9?
Answer
Products are favored. Proton transfer moves toward the weaker acid–base pair, and K ≈ 10^(9 − 4) = 10^5.
Card 380
Question
What is the pOH when [OH⁻] = 2.5 × 10^-4 M?
Answer
3.60, from -log(2.5 × 10^-4).
Card 381
Question
What is the pH of 0.100 M HA when Ka = 1.0 × 10^-5?
Answer
About 3.00. The ICE setup gives Ka = x²/(0.100 − x); x ≈ 1.0 × 10^-3 M, and the 1.0% change validates the approximation.
Card 382
Question
Why does a buffer fail after too much strong acid is added?
Answer
Its conjugate base is depleted, so added H⁺ is no longer consumed effectively.
Card 383
Question
What do two clear equivalence regions on an acid titration curve suggest?
Answer
At least two distinguishable titratable protons; on a clean ideal curve with exactly two equivalence regions, this is consistent with a diprotic acid.
Card 384
Question
A buffer has pKa 4.8 and [A⁻]/[HA] = 0.10; what is pH?
Answer
3.8, from 4.8 + log(0.10).
Card 385
Question
Why is HCl stronger than HF in water despite F being more electronegative?
Answer
The H–F bond is much stronger; bond strength dominates this down-group binary-acid comparison.
Card 386
Question
Why does percent ionization increase when a weak acid is diluted?
Answer
Dilution shifts ionization toward more particles, so a larger fraction ionizes even though [H₃O⁺] decreases.
Card 387
Question
A buffer contains more HA than A⁻. Which addition can it neutralize in greater amount: strong acid or strong base?
Answer
Strong base. The larger HA reserve consumes more added OH⁻; a buffer with more A⁻ than HA instead has greater capacity for added strong acid.
Card 388
Question
Why can removing a basic anion increase a salt's molar solubility without changing Ksp?
Answer
The equilibrium shifts to replace the consumed ion; Ksp remains fixed at the same temperature.
Card 389
Question
Why can an acid–base indicator change color as pH changes?
Answer
Its protonated and deprotonated forms have different colors or other observable properties, and their relative amounts change with pH.
Card 390
Question
What buffer results from mixing 1.0 mol HA with 0.40 mol OH⁻?
Answer
0.60 mol HA and 0.40 mol A⁻ remain, forming a buffer before any equilibrium calculation.
Card 391
Question
What is the pH of 0.200 M weak base B when Kb = 2.0 × 10^-5 at 25°C?
Answer
About 11.30. The ICE setup gives Kb = x²/(0.200 − x); x ≈ 2.0 × 10^-3 M OH⁻, and the 1.0% change validates the approximation.
Card 392
Question
Why does a weak acid alone not make an effective buffer?
Answer
It lacks a substantial conjugate-base reserve to consume added strong acid.
Card 393
Question
What controls pH after excess strong base passes equivalence?
Answer
The concentration of excess OH⁻ after accounting for reaction stoichiometry and total volume.
Card 394
Question
How should an indicator be chosen for a titration?
Answer
Its color-change range should fall within the steep pH change near the equivalence point.
Card 395
Question
How can a measured pH and known pKa give a conjugate-base/acid ratio?
Answer
Rearrange Henderson–Hasselbalch: [A⁻]/[HA] = 10^(pH − pKa).
Card 396
Question
Can a weak base and its conjugate acid form a buffer?
Answer
Yes, when both are present in significant amounts.
Card 397
Question
For equal-volume buffers with the same conjugate-base/acid ratio, how does adding the same amount of strong acid affect a more concentrated versus less concentrated buffer?
Answer
The concentrated buffer changes pH less because it has greater capacity.
Card 398
Question
How does equivalence-point pH compare for strong acid–strong base, weak acid–strong base, and weak base–strong acid titrations at 25°C?
Answer
Strong acid–strong base: pH 7.00. Weak acid–strong base: above 7.00 because the conjugate base reacts with water. Weak base–strong acid: below 7.00 because the conjugate acid reacts with water.
Card 399
Question
Why should mole ratios replace concentration ratios after mixing buffer solutions?
Answer
Both components share the same final volume, so that volume cancels in [A⁻]/[HA].
Card 400
Question
How does adding a little strong acid change a buffer's conjugate-base and conjugate-acid amounts?
Answer
The conjugate base decreases and its conjugate acid increases by the amount of strong acid consumed.
Card 401
Question
What does entropy measure qualitatively?
Answer
The dispersal of matter and energy among available microstates.
Card 402
Question
How is standard reaction entropy calculated?
Answer
ΔS°rxn = ΣνS°(products) − ΣνS°(reactants).
Card 403
Question
What equation gives ΔG° from ΔH° and ΔS°, and what standard states do the degree symbols assume?
Answer
ΔG° = ΔH° − TΔS°. The standard states are pure substances, 1.0 M solutions, and gases at 1 atm or 1 bar; T is in kelvins and energy units must match.
Card 404
Question
Does thermodynamic favorability guarantee a fast reaction?
Answer
No. A favorable reaction can be slow when its activation barrier is large.
Card 405
Question
What is ΔG at equilibrium?
Answer
Zero under the current conditions because there is no net driving force.
Card 406
Question
Why can an endothermic dissolution still be thermodynamically favorable?
Answer
A sufficiently positive entropy change can make TΔS exceed positive ΔH, giving negative ΔG.
Card 407
Question
How can an unfavorable reaction be driven by a favorable one?
Answer
Couple them so their equations and ΔG values add to a negative overall ΔG.
Card 408
Question
Where does oxidation occur in every electrochemical cell?
Answer
At the anode.
Card 409
Question
How are standard cell potential and standard free energy related?
Answer
ΔG° = -nFE°cell.
Card 410
Question
What equation gives cell potential under nonstandard conditions?
Answer
E = E° − (RT/nF) ln Q. When Q = 1, ln Q = 0, so E = E°.
Card 411
Question
How is electrical charge related to current and time?
Answer
q = It.
Card 412
Question
Which phase has greater molar entropy, liquid water or ice at the same temperature?
Answer
Liquid water because its molecules have more accessible arrangements and motion.
Card 413
Question
Do elements in their standard states have zero standard molar entropy?
Answer
No. Their ΔHf° is zero, but their absolute S° values are positive above 0 K.
Card 414
Question
How do the four ΔH° and ΔS° sign combinations determine thermodynamic favorability across temperature?
Answer
ΔH° < 0 and ΔS° > 0 is favorable at every temperature; ΔH° > 0 and ΔS° < 0 is thermodynamically unfavored at every temperature. If both are positive, favorability requires high temperature; if both are negative, it requires low temperature.
Card 415
Question
What does it indicate when a thermodynamically favored process does not occur at a measurable rate?
Answer
It is under kinetic control, commonly because of a high activation energy; no measurable reaction does not mean the system is at equilibrium.
Card 416
Question
How are ΔG° and K related?
Answer
ΔG° = -RT ln K.
Card 417
Question
What two contributions compete in dissolving an ionic solid?
Answer
Enthalpy changes from separating and solvating particles, and entropy changes from their new dispersal and solvent organization.
Card 418
Question
What must cancel when coupled reactions are added?
Answer
Shared intermediates, leaving the desired net reaction.
Card 419
Question
Where does reduction occur in every electrochemical cell?
Answer
At the cathode.
Card 420
Question
What sign of E°cell indicates a favorable standard galvanic reaction?
Answer
Positive E°cell, corresponding to negative ΔG°.
Card 421
Question
If Q increases for a galvanic reaction, how does E change at fixed temperature?
Answer
E decreases according to the Nernst equation. Le Châtelier's principle does not apply to an operating cell away from equilibrium; use Q and Nernst reasoning instead.
Card 422
Question
How are moles of electrons found from charge?
Answer
Moles e⁻ = q/F, where F ≈ 96485 C mol^-1 e⁻.
Card 423
Question
How does producing more gas particles usually affect system entropy?
Answer
It increases entropy because the particles have more positional microstates.
Card 424
Question
Can a dissolution with negative ΔH be unfavorable?
Answer
Yes. A sufficiently negative entropy change at the stated temperature can make ΔG positive.
Card 425
Question
When can a process with ΔH > 0 and ΔS > 0 become favorable?
Answer
At sufficiently high temperature, when TΔS exceeds ΔH.
Card 426
Question
How does a catalyst affect ΔG?
Answer
It does not change ΔG; it lowers the activation barrier for both directions.
Card 427
Question
For A → B, ΔGf°(A) = -50 kJ mol^-1 and ΔGf°(B) = -80 kJ mol^-1. What is ΔG°rxn?
Answer
-30 kJ mol^-1. Use ΣνΔGf°(products) − ΣνΔGf°(reactants) = -80 − (-50).
Card 428
Question
Why can dissolving a gas in a liquid have a negative entropy change?
Answer
Gas particles lose much of their translational freedom when confined and solvated in the liquid.
Card 429
Question
If coupled steps have ΔG values +20 kJ and -35 kJ, what is overall ΔG?
Answer
-15 kJ, so the combined process is thermodynamically favorable under those conditions.
Card 430
Question
What role does each half-cell solution play in an electrochemical cell?
Answer
It supplies dissolved redox species at an electrode interface and carries ions within its compartment. Separate compartments prevent direct mixing while the external circuit and salt bridge connect the half-cells.
Card 431
Question
How is E°cell found from standard reduction potentials?
Answer
E°cell = E°cathode − E°anode, using both tabulated values as reductions.
Card 432
Question
How does a cell's potential magnitude change as Q approaches or moves away from K, and what is E at equilibrium?
Answer
|E| falls toward zero as Q approaches K and grows as the system moves farther from equilibrium. At equilibrium, Q = K and E = 0.
Card 433
Question
How many moles of electrons pass when 1.93 × 10^5 C flows?
Answer
2.00 mol e⁻, from q/F.
Card 434
Question
How does a salt bridge maintain charge balance in a galvanic cell?
Answer
Anions migrate toward the anode compartment and cations toward the cathode compartment, countering the net charge imbalances created by the two half-reactions.
Card 435
Question
Why does raising a substance's temperature generally increase its entropy?
Answer
Energy spreads across more accessible particle energy states, increasing the number of possible microscopic arrangements.
Card 436
Question
When can a process with ΔH < 0 and ΔS < 0 be favorable?
Answer
At sufficiently low temperature, where the unfavorable -TΔS term is small.
Card 437
Question
Why can diamond persist even though graphite is more stable at standard conditions?
Answer
Conversion has a large activation barrier, so diamond is kinetically persistent.
Card 438
Question
What do the external circuit and measuring device do in an electrochemical cell?
Answer
The circuit carries electrons from anode to cathode; a voltmeter measures potential difference, while an ammeter in series measures current.
Card 439
Question
At constant temperature, how does increasing the volume available to a gas affect its entropy?
Answer
Entropy increases because the gas particles can occupy more positions in the larger space, so more microstates are accessible.
Card 440
Question
How does reversing one coupled reaction affect its ΔG?
Answer
It reverses the sign of that reaction's ΔG.
Card 441
Question
Why is n required in ΔG° = -nFE°?
Answer
It is the moles of electrons transferred per balanced reaction, linking charge flow to reaction extent.
Card 442
Question
What makes an electrolytic cell operate?
Answer
An external power source drives a thermodynamically unfavorable redox reaction; oxidation still occurs at the anode and reduction at the cathode.
Card 443
Question
In an Mⁿ⁺/M concentration cell, which half-cell is the anode: the dilute or concentrated ion solution?
Answer
The dilute half-cell. Oxidation produces Mⁿ⁺ there, while reduction consumes Mⁿ⁺ in the concentrated half-cell, so electrons flow from dilute to concentrated as the concentrations move toward equality.
Card 444
Question
How is deposited metal mass found from current and time?
Answer
Find q = It, convert q/F to moles e⁻, use the half-reaction ratio to moles metal, then multiply by molar mass.
Card 445
Question
Given product S° total 500 J mol^-1 K^-1 and reactant total 420 J mol^-1 K^-1, what is ΔS°?
Answer
+80 J mol^-1 K^-1.
Card 446
Question
How do electrode masses change in a Zn–Cu galvanic cell?
Answer
The Zn anode loses mass as Zn → Zn²⁺ + 2e⁻, while the Cu cathode gains mass as Cu²⁺ + 2e⁻ → Cu.
Card 447
Question
What is ΔG° when ΔH° = 50 kJ mol^-1, ΔS° = 0.200 kJ mol^-1 K^-1, and T = 300 K?
Answer
-10 kJ mol^-1, from ΔG° = 50 − (300)(0.200).
Card 448
Question
Why can temperature change a solid's solubility?
Answer
Temperature changes the balance of ΔH and TΔS, so it changes the free energy of dissolution and the equilibrium constant.
Card 449
Question
What does the size of ΔG° relative to RT imply about K?
Answer
ΔG° near zero gives K near 1. When |ΔG°| is much larger than RT, K is far from 1: negative ΔG° gives K ≫ 1, while positive ΔG° gives K ≪ 1.
Card 450
Question
Bubbles form at an inert cathode in acidic solution; which half-reaction can explain them?
Answer
2H⁺ + 2e⁻ → H₂(g). Gas evolution at the cathode can be direct evidence of reduction.
450 cards
Advanced High School Chemistry Flashcards: Complete 9-Part Course Review
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