Topics / Learning Objectives / Challenge Index

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concepts of mol, Avogadro's constant, molar mass, number of particles, etc

conceptual understanding of counting by weighing

simple computation with mole

mass percent and empirical formula

fundamental stoichiometry based on balanced equation

mass percent, molarity of solution and related computation

solution stoichiometry

computation of solution mixing and dilution

gas stoichiometry

limiting reactant and percent yield

impurity analysis

combustion analysis

mixture analysis

concept of molality

vapor pressure of solutions with nonvolatile solutes

colligative properties of solutions (boiling point elevation, freezing point depression, *osmotic pressure, etc.)

colligative properties of solutions with electrolytes as solutes (van't Hoff factor of strong/weak/nonelectrolytes) (check T2)

stoichiometry with more complicated computation

stoichiometry based on graph analysis

color of flame test (Na, K, Ca, Ba, Li, Cu, etc.)

color of common ions (MnO4-, CrO42-, Zn2+, Co2+, *Cr3+, *Ni2+, *Fe3+, main group cations, etc.) 

solubility rules, common precipitates of chlorides, fluorides, sulfates, etc.

color of precipitates

abundance and states of elements

heat of reactions/dissolution/dilution

properties of reactive metals (alkali/alkali earth metals)

*properties of amphoteric metals (Al)

properties of nonreactive metals (Cu/Ag/Au)

properties of representative nonmetals (halogens)

common strong acids (HCl, HNO3, H2SO4, HBr, HI, HClO4)

common weak acids (HF, H2CO3, H3PO4, CH3COOH, etc.)

common strong bases (NaOH, KOH, Ba(OH)2)

common soluble weak bases (NH3, amines)

acid-base properties of salts

types of reactions (neutralization, precipitation, redox, etc.)

typical gas evolution reactions (H2, CO2, SO2, NH3)

typical redox titrations (acidified KMnO4 + H2C2O4/Fe2+/H2O2, *iodometry, etc.) 

names of common used glassware

volumetric glassware and uncertainty (graduated cylinder, volumetric pipette, graduated pipette, burette, etc.)

prepare a certain molarity solution using a volumetric flask 

transfer a certain volume of solution using pipette and bulb/finger

other common laboratory (filtration, distillation, titration, chromatography, gravimetric analysis, etc.)

concepts of spectroscopy, spectrum, absorbance, etc.

color of solution vs light absorbed (complementary color)

Beer's law (Abs=abc)

concepts of standard solution, blank, calibration curve

selection of an appropriate wavelength for the calibration curve

error analysis in spectroscopy

error analysis in acid-base titrations (standardization, indicator, etc.)

error analysis in gravimetric analysis (precipitation)

*error analysis in measurement of the molar mass of a volatile liquid

error analysis in measurement of hydrated water in a hydrate

application of crystal field theory (magnetism, color, etc.) (check T9)

organic laboratory (melting point measurement, vacuum filtration, reflux, etc.)

more complicated error analysis

iodometry and error analysis in iodometry

solubility of sulfide precipitates and hard-soft acid-base theory

concepts of IMFs (London dispersion force, dipole-dipole interaction, hydrogen bonds, ion-dipole interaction, etc.)

Hydrogen bonds (N/O/F-H)

comparison of IMFs and mp/bp/vapor pressure/viscosity/solubility/etc.

gas, liquid, and solid states of matter and their molecular models

phase changes

factors to determine of state - IMFs vs kinetics energy (T, P)

phase diagram (triple point, sublimation, thermal expansion/contraction (water), supercritical fluid, etc.)

concept of ideal gas

conditions for gases to behave ideally (high T, low P)

ideal gas law (PV=nRT)

gas density vs molar mass (d=PM/RT)

kinetic molecular theory (origin of pressure, KE vs T, etc.) and its application in explaining the gas laws

Graham's effusion law (rate∝sqrt[molar mass])

real gas vs ideal gas

concept of dynamic equilibrium of vapor pressure

conceptual understanding of equilibrium vapor pressure (constant at the same temperature)

vapor pressure vs T

normal boiling point

*computation of vapor pressure with volume change 

four types of solids (ionic, metallic, molecular, covalent network)

concept of lattice energy

comparison of lattice energy and mp of ionic solid

typical physical properties of each type of solids

common covalent network solids (diamond, Si, SiO2, SiC, etc.)

unit cell models (preliminary, body-centered, face-centered cubic) and unit formula

*density calculation based on unit cell model

quantitive comparison of molecules with different types of IMFs (CH3Cl vs CCl4, etc.)

computation of vapor pressure vs T (Clapeyron–Clausius equation)

more complicated computation of vapor pressure (two substances, temperature change, etc.)

concepts of calorimetry (thermometer, calorimeter, insulation, etc.)

setup of a coffee-cup calorimeter

measurement of heat of dissolution/dissolution/etc and error analysis 

computation of calorimetry (q=cmΔT) with and without phase change

first law of thermodynamics (U=q+w)

concept of enthalpy (H=U+PV)

*ΔH vs ΔU (internal energy)

concept of ΔH of formation

computation based on ΔH of formation

computation based on Hess's law

estimation of ΔH(rxn) using bond energies

concept of entropy (S)

concept of standard entropy

sign of entropy change based on states and structures of matter

concept of Gibbs free energy (G)

concept of standard free energy change (ΔG) and ΔG of formation

ΔG(standard) and spontaneity

computation based on ΔG=ΔH-TΔS

sign analysis based on ΔG=ΔH-TΔS

dependence of ΔG(standard) with T 

ΔG(standard) vs K (ΔG(standard)=-RTlnK)

*van't Hoff equation (lnK vs 1/T)

*computation of vapor pressure/boiling point at different temperature (Clapeyron–Clausius equation, check T3)

internal energy vs enthalpy

calculation of volume work

second law of thermodynamics

ΔH at different temperature vs heat capacity

graph analysis of lnK vs 1/T

ΔG vs ΔG(standard) - ΔG=ΔG(standard)+RTlnQ

concept of reaction rate and relative rates

concept of rate laws and rate constant, unit of rate constant

integrated rate laws of 0th, 1st, 2nd order reaction

rate constant and half-life

kinetics of nuclear decay (first-order process)

collision model, activation energy (Ea)

factors to affect reaction rate

transition-state model

transition state vs intermediate

dependence of rate on temperature - Arrhenius equation (k vs T)

plot of lnk vs 1/T (measurement of Ea)

catalyst (homogeneous vs heterogeneous)

(multistep) reaction energy profile in the presence and absence of catalyst

measurement of rate law using initial rate method

measurement of rate law using pseudo kinetic analysis

([A] >> [B]) - fading of crystal violet

plot of [A], ln[A], 1/[A] vs t

kinetics analysis of consecutive reactions (A->B->C)

*derivation of rate law using approximations (steady-state approximation, pre-equilibrium approximation, etc.

rate law determination based on graph analysis ([A] vs t, rate vs [A], etc.)

conceptual understanding of activation energy (forward vs reverse reaction rate)

conceptual understanding of Arrhenius equation (pre-exponential factor, etc.)

temperature-dependence of the rate constants with different Ea

half-life of parallel reaction (A->B, A->C)

computation of concentration change of [A] or [B] in A+B->C with a certain initial concentrations

kinetics of enzyme-catalyzed reactions

kinetics analysis of reversible, parallel, and consecutive reactions

concepts of reversible reactions and equilibrium state

concepts of K and Q

reaction prediction based on Q/K comparison

computation with K - RICE table

*Kc vs Kp

Le Chatelier principle (Cobalt equilibrium, NO2/N2O4 equilibrium, etc.)

Brønsted–Lowry acid–base theory and conjugate acid-base pair

strong acids vs weak acids

auto-ionization of water, concept of Kw

concept of pH

pH computation based on Ka or Kb

percent of ionization and its relation with molarity of weak acids

conjugate acid-base pair seesaw (pKa+pKb=pKw)

pH calculation of the conjugate bases of weak acids (CN- or *CO32-, etc.)

concepts of buffer and common ions

*computation based on Henderson-Hasselbalch equation

*preparation of buffers by reactions (NaHCO3 + NaOH, etc.)

acid-based titrations

pH curves of strong acid/strong base titrations

pH curves of strong acid/weak base (equivalence point, end point, buffer region, half equivalence point, etc.)

common used indicators (phenolphthalein, pH=8-10)

principle of indicator (pKIn±1)

selection of indicator based on pH range of color change

concepts of solubility (mass and molar)

solubility equilibrium and concept of Ksp

Ksp vs solubility

prediction of precipitation using Q/K comparison

*solubility vs pH

*complex equilibrium

*dissolution of precipitates by complexation

prediction of reaction shift with pressure change (constant volume or pressure with inert gas added)

titration curve of mixture (Na2CO3 and NaHCO3)

titration curve of polyprotic acids

equilibrium based on graph analysis

error analysis of titrations

computation of multiple equilibria involving dissolution and complex, etc.

concepts of oxidation number, oxidation, and reduction

redox titrations (permanganate)

half-reaction and balance of redox reaction equations

Galvanic cell model, cathode, anode, salt bridge, line notation

standard reduction potential and standard hydrogen electrode

*interpretation of the sign of standard reduction potential

standard cell potential (emf) - Ecell = Ecathod - Eanode

*metal reactivity series and relation to standard reduction potentials

concept of Faraday constant

ΔG(standard)=-nFE(standard)=-RTlnK

measurement of Ksp using electrochemical methods

*potential calculation from related half reactions

application of LCP in redox equilibrium

Nernst equation - E = E(standard)-RT/nF*lnQ

*E-pH diagram

electrolytical cell vs Galvanic cell

Faraday's law of electrolysis

*electrolysis of aqueous solutions (AgF, NaCl, KI, etc.)

E(standard) vs T

oxidation with complex formed (dissolution of Au in aqua regia)

electrolysis of mixed ions

computation of concentration cell

complicated computation based on Nernst equation

structure of atoms and subatomic particles (nucleus, protons, neutrons)

concept of isotopes

name, scientist(s), experiment, and key points of each atomic model (Dalton, Thomson, Rutherford, Bohr, electron cloud)

different modes of nuclear decay (alpha, beta, positron, electron capture, etc.

nuclear fusion vs nuclear fission (U-235 + neutron)

prediction of the mode of nuclear decay

self-practice

light is wave (λf=c, E=hf)

key points of Bohr's atomic model (orbit, energy level, ground state vs excited state, etc.)

Hydrogen's emission spectrum (Lyman series - UV with n[final]=1, Balmer series - Vis with a n[final]=2, Paschen series - IR with a n[final]=3)

*computational understanding of Hydrogen's emission spectrum (energy gap between two neighboring energy levels decreases dramatically with increasing of n)

Self-practice

orbit vs orbital (Heisenberg's uncertainty principle)

square of the value of the wavefunction (Ψ) equals to the probability density of electrons at a certain region

symbol, value, and meaning of the four quantum numbers (n, l, ml, ms)

*radial wavefunction (Ψ) and nodes (Ψ=0)

Self-practice

Aufbau principle, Pauli exclusion principle, Hund's rule

subshells vs orbitals

electron configuration of atoms at ground/excited state

exception of electron configuration of Cr and Cu

photoelectron spectroscopy

electron configuration of ions at ground state

concept of paramagnetic and diamagnetic

counting of unpaired electrons

Self-practice

history and structure (period, group, block) of the periodic table

concept of valence electrons  

concepts of alkali metals, alkaline earth metals, halogens, transition metals, lanthanide, etc.

general trends and justification of atomic/ionic size, first ionization energy (IE), electron affinity, electronegativity, etc.

interpretation of successive IE

trends and justification of mp/bp of alkali metals

*trends and justification of binary acids

Self-practice 

exceptions of first IE1 of group III and VI elements and the justifications

exceptions of electron affinity (O/F are smaller than S/Cl due to their smaller sizes and the justification

*explanation of exceptions in electron affinity across the period

*concepts and affects of lanthanide contraction  (Zr vs Hf) and 3d contraction (Al vs Ga)

Self-practice

number of nodes (n-1) and relationship with quantum number

computational understanding of Hydrogen's emission spectrum based on Rydberg's formula (energy gap between two neighboring energy levels decreases dramatically with increasing of n)

photoelectric effect

explanation of exceptions in electron affinity across the period

concepts and affects of lanthanide contraction  (Zr vs Hf) and 3d contraction (Al vs Ga)

trends of transition elements in the same groups (MnO4- vs TcO4-)

concepts and models of covalent, ionic, and metallic bonds

relationship between bond types and average electronegativity (EN) and ΔEN

polar vs nonopolar bonds (dipole moment)

energy profile of diatomic molecules (bond energy and bond length)

comparison of bond strength and bond length

exceptions in bond strength (Cl2 vs F2)

Self-practice

strategy to draw Lewis structures (octet rule, lone pairs)

label formal charges to confirm the Lewis structures

exceptions to octet rule (odd-electrons [NO/NO2], electron deficient [BF3], hypervalent [PCl5])

concept of resonance structures (atom positions are fixed)

weighted average explanation of bond lengths using multiple resonances (three N-O bonds are the same in NO3-)

comparison of relative stability of resonance structures (octet or not, formal charge, etc.)

Self-practice

VSEPR models (name of each model - AEnXm)

application of VSEPR model to analyze the geometry of molecules/ions

concept of dipole moment

molecular polarity (polar vs nonpolar molecules)

comparison of bond angles

concept of electron-sea model (delocalization)

relationship between metallic bond model and conductivity/malleability/etc.

explanation of melting points of metals in the same group/period using the electron-sea model

melting points vs #valence electrons

valence bond theory (overlap of atomic orbitals)

sigma bonds vs pi bonds and bond counting

concept and analysis of hybridization of central atoms (sp3, sp2, sp)

*violence of VSPER model due to conjugation (delocalized pi bonds, eg. HCONH2)

*uniqueness of 2nd period elements in making multiple bonds (P4 vs N2, S8 vs O2, CO2 vs SiO2, etc.)

Self-practice

concepts of molecular orbital theory (bonding vs anti-bonding, bond order)

bond order of common diatomic molecules (N2, O2, NO, etc.)

explanation of unpaired electrons using MO theory (O2, O2-, NO, etc.)

*application of MO theory in explaining IE1 (Cl vs Cl2)

Self-practice

concepts of complex (coordination covalent bonds, ligands)

isomerism of square planar and octahedral complexes (cis/trans, fac/mer, optical)

*isomer counting of square planar and octahedral complex (MA2B2, MA2BC, MA3B3, MA2B4, etc.)

Self-practice

advanced MO theory (four rules)

s-p mixing in MO (O2 and F2)

explanation of unpaired electrons using MO theory - heteronuclear diatomic species

application of MO theory in explaining IE -  heteronuclear diatomic species

geometric vs stereoisomers

isomer counting (with bidentate ligands)

crystal field theory (splitting of d orbitals in octahedral)

concepts of strong ligand and weak ligands

color of complex ion and splitting energy

hydrocarbons (alkanes/alkenes/alkynes, benzene), alkyl halides, alcohols, aldehydes/ketones, acids, esters, amino acids, etc.

concepts of double bond equivalence and molecular formula of organic compounds

number of sigma and pi bonds in a organic molecule

Note: naming of alkanes, alkenes, and esters were commonly assessed

Self-practice

planar structure of alkenes (double bonds can't rotate freely)

cis vs trans isomers of alkenes

Note: the naming of cis/trans of ring structures were NOT assessed

a sp3-hybridized carbon with four different groups is called a chiral carbon

a chiral molecule is NOT superimposable with its mirror image, and can rotate the plane-polarized light (optical isomers)

Note: assign of R/S of a chiral carbon is NOT assessed

Extension: a chiral molecule has neither plane of symmetry and nor inversion center.

concepts of structural isomers (functional, chain, positional) and stereoisomers (geometric [cis/trans] and optical)

Note: advanced concepts such as enantiomers, diastereomers were assessed in the recent years

Conceptual: geometric isomers are diastereomers (stereoisomers but not enantiomers)

count the number of isomers with a certain molecular formula with a systematic strategy

Note: isomer counting for alkanes, alkenes, alcohols, haloalkanes were assessed

Note: alkenes may have cis/trans isomers, and others may have chiral carbons with optical isomers

Note: isomers with rings were NOT assessed yet 

formation of esters by reacting carboxylic acids and alcohols under acid catalysis (reversible)

hydrolysis of esters under acid (reversible) and base (saponification, irreversible) catalysis

Extension: basic hydrolysis of triglycerides (fat and oil) into salts of fatty acids (soap) and glycerol

primary, secondary, and tertiary alcohols

primary alcohols can be oxidized into aldehydes [intermediate] and further into acids

secondary alcohols can be oxidized into ketones (no further reaction)

tertiary alcohols can NOT be oxidized

Extension: strong oxidants (permanganate/chromate) oxidize primary alcohols into acids, while mild oxidants produces aldehydes

concepts of nucleophiles, electrophiles, leaving groups, etc.

intermediate/transition state of nucleophilic substitutions

*relative rates of primary/secondary/tertiary substrates and different halides as leaving groups

Note: stereochemistry of SN1/SN2 was NOT assessed in the past USNCO exams

benzene goes substitution instead of addition despite its high degree of unsaturation (benzene doesn't decolorize bromine like alkenes do)

benzene can undergo substitution reactions such as halogenation, nitration, alkylation, etc. under Lewis Acids catalysis

Note: the unique properties of benzene is caused by delocalization of pi electrons (aromatic)

Extension: the activation/deactivation effect of different groups on benzene was assessed in the National Exam

acid-catalyzed dehydration of alcohols to produces alkenes

base-catalyzed dehydrohalogenation (removal of HX) of haloalkanes produces alkenes

Note: the beta-hydrogen (the hydrogen connected to the carbon adjacent to the central carbon with -OH or -X) is removed in both reactions

Note: tendency of elimination: tertiary  > secondary > primary alcohols

Extension: the regioselectivity (Zaitsev rule) and stereoselectivity (cis/trans alkenes) of eliminations were assessed in the National Exam Part II?

alkenes can undergo addition reactions to produce substituted alkanes, such as the addition of *Br2, HX, *H2O (acid-catalyzed), H2, etc.

Note: the decolorization of bromine is used to test alkenes/alkynes

Note: the addition rate of substituted alkenes and alkynes was assessed in the National Exam Part I, though the stability comparison of tertiary, secondary, primary carbocations was NOT directly assessed.

Note: Markovnikov rule and radical-based addition were assessed in the National Exam Part 

*chair and boat conformation of cyclohexanes and their relative stability

*concept of staggered and eclipsed bonds

Note: the relative stability of the axial and equatorial substituted cyclohexanes were assessed in the National Exam

4n+2 rule - comparison of benzene and cyclooctatetraene

hydration of aldehydes/ketones

formation and reactions of hemiacetals/acetals

aldol condensation of aldehydes/ketones

reactivity of alkenes vs alkynes

reactive of alkenes vs conjugated alkenes

reactivity of allylic vs regular positions

concepts and structural explanation of reducing sugar

competition of nucleophilic substitutions and eliminations

synthesis of carboxylic acid derivatives, such as amides

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