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Chemistry Class XI
| Question | Answer |
|---|---|
| What is the key distinction between a pure substance and a mixture | A pure substance has constituents of the same chemical nature whereas a mixture contains two or more pure substances whose composition can vary |
| Why can a compound not be separated into its elements by ordinary physical methods | Its constituent elements are chemically combined in a fixed ratio |
| Why can the composition of a mixture vary while that of a compound remains fixed | Mixture components retain their identities and can be present in variable proportions whereas compound constituents are chemically combined in a fixed ratio |
| A substance has definite volume but takes the shape of its container; which state is it | Liquid |
| Why does a gas have neither definite shape nor definite volume | Its particles are relatively far apart and have high freedom of movement |
| What happens to particle arrangement when a solid changes into a liquid | Particles remain close but gain enough freedom to move relative to one another |
| Which SI base unit is used for amount of substance | Mole |
| Which SI base unit is used for thermodynamic temperature | Kelvin |
| Why is weight not an SI base quantity while mass is | Mass measures amount of matter whereas weight is a gravitational force and depends on local gravitational acceleration |
| Convert 1 mL into cubic centimetres | 1 cm³ |
| Convert 1 m³ into litres | 1000 L |
| Convert 1 L into cubic decimetres | 1 dm³ |
| What is the density of a substance whose mass is m and volume is V | ρ = m divided by V |
| If density and volume are known how is mass calculated | m = ρV |
| If mass and density are known how is volume calculated | V = m divided by ρ |
| What happens to the numerical value of density when converting between g mL⁻¹ and kg m⁻³ | 1 g mL⁻¹ = 1000 kg m⁻³ |
| Convert 25 °C into kelvin | 298.15 K |
| Why is absolute temperature required in gas-law calculations | Kelvin represents an absolute temperature scale |
| Why is Kelvin written as K rather than °K | The SI unit kelvin is written without the degree symbol |
| What is the general form of scientific notation | N × 10ⁿ where N is generally between 1 and 10 |
| Before adding or subtracting numbers in scientific notation what must be done | Their powers of ten must first be made compatible |
| Which digits are always significant | All non-zero digits |
| Are zeros between two non-zero digits significant | Yes |
| Are leading zeros significant | No |
| Are trailing zeros after a decimal point significant | Yes |
| Why can trailing zeros in an integer be ambiguous | Without a decimal point or scientific notation their significance may not be clear |
| How many significant figures does 0.0025 have | 2 |
| How many significant figures does 5005 have | 4 |
| How many significant figures does 500.0 have | 4 |
| Why does 500.0 have more clearly defined significant figures than 500 | The decimal point makes the trailing zeros significant |
| How does accuracy differ from precision | Accuracy is closeness to the accepted value whereas precision is closeness of repeated measurements to one another |
| Can measurements be precise but inaccurate | Yes |
| Can measurements be accurate but not precise | Yes |
| For addition and subtraction what determines the number of decimal places in the final answer | The least precise decimal place among the measurements |
| For multiplication and division what determines the significant figures in the final answer | The measurement having the fewest significant figures |
| What happens when a removed digit is greater than 5 during rounding | The preceding digit is increased by 1 |
| What happens when a removed digit is less than 5 during rounding | The preceding digit remains unchanged |
| What special rule must be remembered for an exact 5 in the NCERT rounding convention | The even or odd convention must be applied carefully |
| What is the factor-label method also called | Dimensional analysis or unit-factor method |
| What is the fundamental strategy in dimensional analysis | Multiply by conversion factors equal to one so unwanted units cancel |
| Why are units useful for detecting mistakes in calculations | Incorrect dimensions often reveal an invalid conversion or equation |
| Which scientist is associated with the law of conservation of mass | Antoine Lavoisier |
| What does the law of conservation of mass state | Mass is neither created nor destroyed in a chemical reaction |
| Under what condition is total reactant mass equal to total product mass | In a closed system undergoing the chemical reaction |
| Which scientist is associated with the law of definite proportions | Joseph Proust |
| What does the law of definite proportions state | A given compound always contains the same elements in the same fixed proportion by mass |
| How does the law of multiple proportions differ from the law of definite proportions | Multiple proportions compares different compounds formed by the same two elements whereas definite proportions concerns the fixed composition of one compound |
| Which scientist is associated with the law of multiple proportions | Dalton |
| If two elements form two compounds and masses of one element combining with a fixed mass of the other are 16 g and 32 g what law is illustrated | Law of multiple proportions |
| What ratio do 16 g and 32 g represent | 1:2 |
| What is Gay-Lussac's law of gaseous volumes | Gases react in simple whole-number volume ratios when measured at the same temperature and pressure |
| Under what conditions can gaseous volume ratios be directly compared using Gay-Lussac's law | At the same temperature and pressure |
| What does Avogadro's law state | Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules |
| At constant temperature and pressure how are gas volume and number of molecules related | They are directly proportional |
| At constant temperature and pressure how are gas volume and number of moles related | They are directly proportional |
| Which Dalton postulate is contradicted by the existence of isotopes | The proposal that atoms of the same element are identical in mass |
| Which Dalton postulate is consistent with the law of multiple proportions | Atoms combine in simple whole-number ratios to form compounds |
| What did modern discoveries show about the divisibility of atoms | Atoms contain subatomic particles and are therefore divisible |
| What is the reference standard for atomic mass | Carbon-12 |
| Define one unified atomic mass unit | One twelfth of the mass of one carbon-12 atom |
| How is average atomic mass calculated for naturally occurring isotopes | Sum of fractional abundance multiplied by the corresponding isotopic mass |
| Why must percentage isotopic abundance be converted to a fraction before using the average-mass formula | The weighted average formula uses fractional abundances |
| If an element has two isotopes with equal abundance how is its average atomic mass related to their masses | It is the arithmetic mean of their isotopic masses |
| What is molecular mass | The sum of the atomic masses of all atoms present in one molecule |
| Why is formula mass used for NaCl instead of molecular mass | Solid NaCl is an ionic lattice and does not consist of discrete NaCl molecules |
| What is one mole | The amount of substance containing Avogadro's number of specified entities |
| What types of entities can be counted using the mole concept | Atoms molecules ions electrons formula units and other specified particles |
| What is the value of Avogadro constant given in the source | 6.022 × 10²³ mol⁻¹ |
| How are number of entities and moles related | N = nNₐ |
| How are moles and number of entities related | n = N divided by Nₐ |
| What is molar mass | The mass of one mole of a substance |
| What is the SI-style unit commonly used for molar mass in the chapter | g mol⁻¹ |
| Why do atomic mass in u and molar mass in g mol⁻¹ have the same numerical value | The definitions are constructed so that one mole of atoms has the corresponding numerical mass in grams |
| What is the master conversion pathway between mass and particles | Mass → moles → number of entities or the reverse |
| How do you convert mass directly into number of molecules conceptually | Convert mass to moles using molar mass then moles to molecules using Avogadro constant |
| How do you calculate the number of atoms in a compound from its mass | Convert mass to moles of compound then multiply by Avogadro constant and by the number of the required atoms per formula unit |
| What is the percentage by mass of an element in a compound | Mass contribution of the element in one mole of compound divided by molar mass of compound multiplied by 100 |
| Why is assuming 100 g useful when percentage composition is given | Each percentage can then be treated directly as the corresponding mass in grams |
| What is an empirical formula | The simplest whole-number ratio of atoms of different elements |
| What is the first conversion in a percentage-composition empirical-formula problem | Convert the percentage composition into masses by assuming a 100 g sample |
| What is the second major step after obtaining masses in an empirical-formula problem | Convert each elemental mass into moles |
| Why are all calculated mole values divided by the smallest mole value | To reduce the mole ratio to the simplest ratio |
| What should be done if an empirical-formula ratio is fractional | Multiply all ratios by an appropriate small integer to obtain whole numbers |
| What multiplier is commonly used for a ratio of 1.5 | 2 |
| What multiplier is commonly used for a ratio of about 1.33 | 3 |
| What multiplier is commonly used for a ratio of about 1.25 | 4 |
| What is the relationship between molecular and empirical formula | Molecular formula = empirical formula raised to an integer n |
| How is n calculated when finding a molecular formula | n = molecular molar mass divided by empirical formula mass |
| What must be true about n in a valid molecular-formula calculation | It must be a whole number |
| What is the correct first step in a stoichiometric calculation | Write the correct chemical equation and balance it |
| What do coefficients in a balanced chemical equation represent | Mole ratios between the reacting and product species |
| Why should masses never be used directly as equation coefficients | Balanced-equation coefficients represent mole ratios rather than mass ratios |
| What is the universal stoichiometric pathway | Given quantity → moles → balanced-equation mole ratio → required moles → required quantity |
| How do you solve a mass-to-mass stoichiometry problem | Convert given mass to moles then use the mole ratio and finally convert required moles into mass |
| How do you identify the limiting reagent using product amounts | Calculate the amount of product each reactant could form; the reactant producing the smaller amount is limiting |
| What is the limiting reagent | The reactant consumed first that limits the amount of product formed |
| What happens to a non-limiting reactant | It remains in excess after the limiting reagent is consumed |
| Why can two reactants with equal masses have different limiting behaviour | Their molar masses and stoichiometric requirements can differ |
| What is the key comparison for finding a limiting reagent from moles | Compare available moles with the required stoichiometric ratio |
| After finding the limiting reagent what two quantities can commonly be calculated | The amount of product formed and the amount of excess reagent remaining |
| What is mass percentage of a component in a solution | Mass of component divided by mass of solution multiplied by 100 |
| What is mole fraction of component A | χA = nA divided by total moles |
| What is the sum of mole fractions in a solution | 1 |
| Does mole fraction have a unit | No |
| For a two-component solution what relation connects the two mole fractions | χA + χB = 1 |
| What is molarity | Moles of solute divided by volume of solution in litres |
| What is the unit of molarity | mol L⁻¹ |
| Does molarity use volume of solvent or volume of solution | Volume of solution |
| What is molality | Moles of solute divided by mass of solvent in kilograms |
| What is the unit of molality | mol kg⁻¹ |
| Does molality use mass of solvent or mass of solution | Mass of solvent |
| Why can molarity change with temperature more readily than molality | Molarity depends on solution volume whereas molality depends on solvent mass |
| Which concentration term is based on the mass of solvent | Molality |
| Which concentration term is based on the volume of solution | Molarity |
| If the volume of solution is 2 L and it contains 0.5 mol solute what is the molarity | 0.25 mol L⁻¹ |
| If a solution contains 0.4 mol solute and 200 g solvent what is its molality | 2 mol kg⁻¹ |
| What is the key reason density can be combined with concentration data in numerical problems | Density connects mass and volume |
| What is the main difference between molecular mass and formula mass | Molecular mass applies to discrete molecules whereas formula mass is used for substances represented by formula units such as ionic solids |
| What does matter necessarily possess according to the source | Mass and occupies space |
| What distinguishes homogeneous from heterogeneous mixtures | A homogeneous mixture has uniform composition whereas a heterogeneous mixture does not |
| Why are elements considered pure substances | They contain only one type of atom and cannot be broken into simpler substances by ordinary chemical methods |
| What happens to chemical identity during a physical change | The chemical identity remains unchanged |
| What makes a property chemical rather than physical | It concerns the chemical behaviour or transformation of a substance |
| What was the central experimental basis for discovering the electron | Cathode-ray discharge-tube experiments |
| What did the independence of electron charge-to-mass ratio from gas and electrode material imply | Electrons are universal constituents of matter |
| What did Millikan's oil-drop experiment establish | Electric charge is quantised and the elementary electronic charge can be determined |
| What mathematical relation represents charge quantisation | q = ne |
| What does n represent in q = ne | A positive integer such as 1 2 3 and so on |
| What forces are considered in the Millikan oil-drop experiment | Gravitational force electrostatic force and viscous drag force |
| How is electron mass obtained when its charge and charge-to-mass ratio are known | m = e divided by e over m |
| What is the charge of an electron relative to the proton | Equal magnitude and opposite sign |
| What is the charge of a neutron | Zero |
| Which subatomic particle was discovered by Chadwick | Neutron |
| In what year was the neutron discovered according to the source | 1932 |
| What target was bombarded in Chadwick's neutron experiment | Beryllium |
| What projectile was used to bombard beryllium in Chadwick's experiment | Alpha particles |
| What observation established that cathode rays consist of negatively charged particles | They were deflected by electric and magnetic fields in a manner characteristic of negative charge |
| Why were phosphorescent materials used in cathode-ray experiments | Cathode rays are not directly visible but produce visible effects on such materials |
| What happens to cathode rays in the absence of electric and magnetic fields | They travel in straight lines |
| What did the perforated anode demonstrate in the cathode-ray experiment | The path of cathode rays could be observed beyond the anode |
| What did Thomson's atomic model propose about positive charge | It was uniformly distributed throughout a spherical atom |
| Where were electrons located in Thomson's model | Embedded within the uniformly distributed positive charge |
| What alternative names are associated with Thomson's model | Plum pudding model raisin pudding model and watermelon model |
| What major experimental result disproved Thomson's model | Rutherford's alpha-particle scattering observations |
| Why was gold used in Rutherford's scattering experiment | It could be prepared as an extremely thin foil |
| What did the fact that most alpha particles passed through the gold foil undeflected imply | Most of the atom is empty space |
| What did small-angle alpha-particle deflections imply | Positive charge is concentrated in a small central region |
| What did the very small number of large-angle or backward deflections imply | The nucleus is extremely small dense and positively charged |
| Where is nearly all the mass of an atom concentrated according to Rutherford's model | In the nucleus |
| Why was Rutherford's model unable to explain atomic stability | According to classical electromagnetic theory an orbiting charged electron should continuously lose energy and eventually collapse into the nucleus |
| What other major phenomenon could Rutherford's model not explain | Atomic line spectra |
| What is atomic number Z | The number of protons in the nucleus |
| For a neutral atom what is the relationship between Z and electron number | Z equals the number of electrons |
| What is mass number A | The total number of protons and neutrons |
| How are A Z and neutron number related | A = Z + n |
| How is neutron number calculated from A and Z | n = A − Z |
| What does nuclear notation A over Z X represent | A is mass number Z is atomic number and X is the element symbol |
| How do you calculate electron number for a positive ion | Subtract the positive charge magnitude from the atomic number |
| How do you calculate electron number for a negative ion | Add the negative charge magnitude to the atomic number |
| What are isotopes | Atoms of the same element having the same atomic number but different mass numbers |
| Why do isotopes have different numbers of neutrons | Their mass numbers differ while their atomic numbers are identical |
| Why do isotopes generally show similar chemical behaviour | Chemical behaviour depends mainly on electron number |
| What are isobars | Atoms of different elements having the same mass number but different atomic numbers |
| Why are C-14 and N-14 isobars | They have the same mass number but different atomic numbers |
| What is wavelength | The distance between two successive equivalent points of a wave |
| What is frequency | The number of wave cycles passing a point per second |
| What is the SI unit of frequency | s⁻¹ or hertz |
| What is the relation between speed wavelength and frequency of electromagnetic radiation | c = νλ |
| How is wavenumber defined | The reciprocal of wavelength |
| What is the common unit of wavenumber | cm⁻¹ |
| If wavelength increases what happens to frequency for electromagnetic radiation in vacuum | Frequency decreases |
| If wavelength decreases what happens to photon energy | Photon energy increases |
| What is the correct sequence of electromagnetic regions from longest to shortest wavelength | Radio waves → microwaves → infrared → visible → ultraviolet → X-rays → gamma rays |
| Which region has the highest frequency in the listed electromagnetic spectrum | Gamma rays |
| Which region has the longest wavelength in the listed electromagnetic spectrum | Radio waves |
| What did Planck propose about the emission and absorption of energy | Energy is emitted or absorbed in discrete packets |
| What is a quantum of electromagnetic radiation called | Photon |
| What is the energy of a photon | E = hν |
| How can photon energy be expressed in terms of wavelength | E = hc divided by λ |
| How are photon energy and frequency related | They are directly proportional |
| How are photon energy and wavelength related | They are inversely proportional |
| What is the relation between electron volt and joule given in the source | 1 eV = 1.6020 × 10⁻¹⁹ J |
| What is the photoelectric effect | Emission of electrons from a metal surface when suitable radiation falls on it |
| What are emitted electrons called | Photoelectrons |
| What is work function | The minimum energy associated with removing an electron from the metal surface in the photoelectric-effect framework |
| What is threshold frequency | The minimum frequency of incident radiation required for photoelectron emission |
| What is the Einstein photoelectric equation | hν = φ + KE |
| How is maximum kinetic energy related to photon frequency and work function | KE = hν − φ |
| What happens if photon energy is equal to the work function | The emitted electron has zero maximum kinetic energy |
| What happens to photoelectron kinetic energy when incident frequency increases while the metal remains the same | The maximum kinetic energy increases |
| What distinguishes an emission spectrum from an absorption spectrum | Emission shows discrete wavelengths emitted by excited atoms whereas absorption shows wavelengths absorbed from incident radiation |
| Why does hydrogen produce a line spectrum rather than a continuous spectrum | Its electrons can occupy quantised energy levels and transitions occur only between allowed levels |
| What does the uniqueness of an element's line spectrum imply | Each element has characteristic quantised electronic-energy transitions |
| What does the Rydberg equation describe | The wavelengths or wavenumbers of hydrogen spectral lines |
| What conditions apply to the principal quantum numbers in the hydrogen Rydberg relation | n₁ is 1 2 3 and so on while n₂ is greater than n₁ |
| Which hydrogen spectral series ends at n = 1 | Lyman series |
| Which hydrogen spectral series lies in the visible region | Balmer series |
| Which hydrogen spectral series ends at n = 3 | Paschen series |
| Which hydrogen spectral series ends at n = 4 | Brackett series |
| Which hydrogen spectral series ends at n = 5 | Pfund series |
| In which region does the Lyman series occur | Ultraviolet |
| In which region do the Paschen Brackett and Pfund series occur | Infrared |
| What did Bohr introduce to explain hydrogen's spectrum | Quantised stationary energy states |
| What happens to the energy of an electron while it remains in a permitted stationary state | No energy is emitted |
| What happens when an electron moves from a lower energy level to a higher energy level | The atom absorbs energy |
| What happens when an electron moves from a higher energy level to a lower energy level | The atom emits energy |
| What is Bohr's angular-momentum quantisation condition | mvr = nh divided by 2π |
| What does n represent in Bohr's angular-momentum equation | The principal quantum number identifying the allowed orbit |
| What is the general trend of Bohr orbit radius with increasing n | It increases |
| What is the ground state | The lowest-energy state of an electron with n = 1 |
| What is an excited state | A state with a higher principal quantum number and higher energy than the ground state |
| Why does electron energy approach zero at infinite separation from the nucleus | The electron becomes effectively unbound from the nucleus |
| What determines the energy difference involved in a transition between two levels | The difference between their quantised energy levels |
| What relation connects transition energy with photon frequency | ΔE = hν |
| What relation connects transition energy with photon wavelength | ΔE = hc divided by λ |
| What are hydrogen-like species | One-electron species such as H He⁺ and Li²⁺ |
| Why must nuclear charge be included when applying Bohr formulas to hydrogen-like ions | Their energy and radius depend on nuclear charge as well as principal quantum number |
| Which systems can be treated quantitatively using the Bohr model | Hydrogen-like one-electron species |
| What are major limitations of the Bohr model | It cannot satisfactorily explain multi-electron spectra fine spectral details the Zeeman effect the Stark effect and the wave nature of electrons |
| What did de Broglie propose about matter | Moving matter particles have associated wave-like behaviour |
| What is the de Broglie wavelength equation | λ = h divided by mv |
| How is de Broglie wavelength expressed using momentum | λ = h divided by p |
| What happens to de Broglie wavelength when particle momentum increases | It decreases |
| Why do macroscopic objects generally show negligible observable wave behaviour | Their large mass and momentum give them extremely small de Broglie wavelengths |
| What is Heisenberg's uncertainty principle | The exact position and exact momentum of a particle cannot both be known simultaneously |
| What is the position-momentum uncertainty relation | ΔxΔp ≥ h divided by 4π |
| How can momentum uncertainty be written for a particle of fixed mass | Δp = mΔv |
| What does the uncertainty principle imply about classical electron paths | Electrons cannot be assigned exact classical paths |
| Is Heisenberg uncertainty merely an instrumental limitation | No it is a fundamental quantum limitation |
| What is an atomic orbital | A probability-based region in which the probability of finding an electron is high |
| What is the key difference between an orbit and an orbital | An orbit is a fixed path in the Bohr model whereas an orbital is a probability-based region |
| What does higher probability density mean | There is a greater likelihood of finding the electron in that region |
| How many quantum numbers describe an electron | Four |
| What are the four quantum numbers | Principal azimuthal magnetic and spin quantum numbers |
| What does the principal quantum number n describe | The shell or main energy level and broadly the size and energy of the orbital |
| What are the allowed values of n | 1 2 3 and so on |
| Which shell corresponds to n = 1 | K shell |
| Which shell corresponds to n = 4 | N shell |
| What is the maximum number of electrons in a shell with principal quantum number n | 2n² |
| How many electrons can the n = 3 shell accommodate at maximum | 18 |
| What does the azimuthal quantum number l describe | The subshell and orbital angular-momentum characteristics |
| What are the allowed values of l for a given n | 0 to n − 1 |
| Which subshell corresponds to l = 0 | s |
| Which subshell corresponds to l = 1 | p |
| Which subshell corresponds to l = 2 | d |
| Which subshell corresponds to l = 3 | f |
| How many subshells are present in the nth shell | n |
| Which subshells are possible for n = 3 | 3s 3p and 3d |
| What does the magnetic quantum number m_l describe | The orientation of an orbital |
| What are the allowed values of m_l for a given l | −l through 0 to +l |
| How many orbitals are present in a subshell with azimuthal quantum number l | 2l + 1 |
| How many orbitals are present in an s subshell | 1 |
| How many orbitals are present in a p subshell | 3 |
| How many orbitals are present in a d subshell | 5 |
| How many orbitals are present in an f subshell | 7 |
| What are the possible m_l values for a p subshell | −1 0 and +1 |
| What are the possible m_l values for a d subshell | −2 −1 0 +1 and +2 |
| What does the spin quantum number m_s represent | The two possible spin states of an electron |
| What are the possible values of m_s | +1/2 and −1/2 |
| How many electrons can occupy one orbital | 2 |
| How many electrons can an s subshell hold | 2 |
| How many electrons can a p subshell hold | 6 |
| How many electrons can a d subshell hold | 10 |
| How many electrons can an f subshell hold | 14 |
| What is the shape of an s orbital | Spherical |
| What is the general shape of a p orbital | Dumbbell-shaped |
| How many orientations do the three p orbitals have | Three orientations called pₓ pᵧ and p_z |
| What distinguishes pₓ pᵧ and p_z orbitals within the same subshell | Their spatial orientation |
| What is a node | A region where the probability of finding an electron is zero |
| What is the formula for total nodes | n − 1 |
| What is the formula for angular nodes | l |
| What is the formula for radial nodes | n − l − 1 |
| How many total nodes does a 3p orbital have | 2 |
| How many angular nodes does a 3p orbital have | 1 |
| How many radial nodes does a 3p orbital have | 1 |
| How many total nodes does a 3d orbital have | 2 |
| How many angular nodes does a 3d orbital have | 2 |
| How many radial nodes does a 3d orbital have | 0 |
| How many radial nodes does a 4s orbital have | 3 |
| For hydrogen-like atoms how do orbital energies compare within the same shell | All orbitals with the same n have the same energy |
| How do orbital energies differ in multi-electron atoms | They depend on both n and l |
| What rule compares orbital energies using n + l | The orbital with lower n + l has lower energy |
| If two orbitals have equal n + l which one has lower energy | The orbital with lower n has lower energy |
| Which has lower energy in a multi-electron atom 4s or 3d | 4s because 4s has n + l = 4 while 3d has n + l = 5 |
| Which has lower energy in a multi-electron atom 4p or 5s | 4p because both have n + l = 5 but 4p has lower n |
| Which has lower energy 3p or 3d | 3p because n + l is 4 for 3p and 5 for 3d |
| What is the Aufbau principle | Electrons occupy lower-energy orbitals before higher-energy orbitals |
| What is the standard orbital filling sequence through 7p | 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p |
| What is the Pauli exclusion principle | No two electrons in an atom can have the same set of all four quantum numbers |
| What consequence does Pauli's principle have for electrons in one orbital | At most two electrons can occupy one orbital and they must have opposite spins |
| What is Hund's rule of maximum multiplicity | Electrons occupy degenerate orbitals singly before pairing occurs |
| Why do electrons occupy degenerate orbitals singly before pairing | This arrangement avoids premature pairing and follows Hund's maximum-multiplicity rule |
| For a p⁴ configuration how are electrons distributed among the three p orbitals according to Hund's rule | One orbital contains a pair while the other two contain one electron each |
| What is the main mistake when drawing orbital diagrams for p⁴ if Hund's rule is ignored | Pairing is done before each degenerate orbital receives one electron |
| What three principles together determine ground-state electronic configuration | Aufbau principle Pauli exclusion principle and Hund's rule |
| What is the maximum number of electrons in the nth shell and why | 2n² because the shell contains n subshells and their total orbital capacity gives 2n² |
| Which scientist proposed the law of triads | Dobereiner |
| What characterises a Dobereiner triad | Three chemically similar elements whose middle element has properties intermediate between the other two and whose atomic mass is approximately related to them |
| Which three elements form a classic alkali-metal Dobereiner triad | Li Na K |
| Which three elements form a classic alkaline-earth Dobereiner triad | Ca Sr Ba |
| Which three elements form a classic halogen Dobereiner triad | Cl Br I |
| What was the major limitation of Dobereiner's triads | The pattern worked for only a limited number of groups |
| What was distinctive about de Chancourtois' arrangement | Elements were arranged by increasing atomic weight on a cylindrical arrangement showing recurring periodicity |
| What was Newlands' Law of Octaves | When elements were arranged by increasing atomic weight every eighth element showed properties similar to the first |
| Why was Newlands' arrangement compared with musical octaves | Similarity recurred at every eighth element |
| What was the major limitation of Newlands' Law of Octaves | It worked mainly only up to calcium |
| What periodic physical properties did Lothar Meyer study | Properties such as atomic volume melting point and boiling point |
| What did Lothar Meyer's plots against atomic weight reveal | Periodic repetition of physical properties |
| What was Mendeleev's original periodic law | The properties of elements are periodic functions of their atomic weights |
| How did Mendeleev arrange elements | Generally in increasing atomic mass while placing elements with similar properties together |
| Why were gaps intentionally left in Mendeleev's table | To accommodate elements that had not yet been discovered |
| Why was Mendeleev's prediction of undiscovered elements a major achievement | He could predict their properties before their discovery |
| What was one major limitation of Mendeleev's periodic table involving isotopes | Isotopes could not be given unique positions using atomic mass ordering |
| Why was hydrogen's position problematic in Mendeleev's table | Its properties gave it similarities with more than one group |
| What other issue arose from strict atomic-mass ordering in Mendeleev's table | Some atomic-mass-order anomalies had to be accommodated |
| What is the modern periodic law | The physical and chemical properties of elements are periodic functions of their atomic numbers |
| What is the key distinction between Mendeleev's and modern periodic laws | Mendeleev used atomic mass whereas the modern law uses atomic number |
| How many periods are present in the modern periodic table | 7 |
| How many groups are present in the modern periodic table | 18 |
| What is a period | A horizontal row of the periodic table |
| What is a group | A vertical column or family of elements |
| Why do elements in the same group generally show similar chemical properties | They have related valence-shell electronic configurations |
| Where are metals generally located in the periodic table | Toward the left side |
| Where are non-metals generally located | Toward the upper-right side |
| Where are metalloids generally located | Along the zig-zag boundary between metals and non-metals |
| How many elements are in the first period | 2 |
| How many elements are in the second period | 8 |
| How many elements are in the third period | 8 |
| How many elements are in the fourth period | 18 |
| How many elements are in the fifth period | 18 |
| What is the broad electronic significance of the period number | It corresponds to the principal shell being filled |
| Which group contains the alkali metals | Group 1 |
| Which group contains the alkaline-earth metals | Group 2 |
| Which group contains the halogens | Group 17 |
| Which group contains the noble gases | Group 18 |
| What are the temporary digit roots used for systematic naming of elements above atomic number 100 | 0 nil 1 un 2 bi 3 tri 4 quad 5 pent 6 hex 7 sept 8 oct 9 enn |
| How is a systematic temporary element name constructed from the atomic number | Use the roots corresponding to each digit in order and add ium |
| What should be checked when converting a high atomic number into a temporary systematic name | Each digit must be represented by its correct root and the ending must be formed appropriately |
| How is period generally identified from an electronic configuration | By the highest principal quantum number present |
| What period is represented by 1s²2s²2p⁶ | Period 2 |
| What period is represented by [Ne]3s² | Period 3 |
| What period is represented by [Ar]4s² | Period 4 |
| What is the general valence configuration of a Group 1 s-block element | ns¹ |
| What is the general valence configuration of a Group 2 s-block element | ns² |
| What is the general valence configuration of a Group 17 element | ns²np⁵ |
| What is the general valence configuration of a Group 18 element | ns²np⁶ except helium which is 1s² |
| Why is helium placed in Group 18 despite having a 1s² configuration | It has noble-gas-like stability and belongs with the noble gases |
| What is the general outer configuration of p-block elements | ns²np¹ through ns²np⁶ |
| What differentiates d-block elements electronically | The differentiating electron enters the (n−1)d subshell |
| What differentiates f-block elements electronically | The differentiating electron enters the (n−2)f subshell |
| Which blocks contain Groups 1 and 2 | s-block |
| Which blocks contain Groups 13 through 18 | p-block |
| Where is the d-block located in the periodic table | In the middle region |
| What elements are included in the f-block | Lanthanoids and actinoids |
| What is the general trend of atomic radius across a period | It generally decreases |
| Why does atomic radius generally decrease across a period | Nuclear charge increases while electrons are added to the same principal shell so effective nuclear attraction generally increases |
| What is the general trend of atomic radius down a group | It generally increases |
| Why does atomic radius increase down a group | New electron shells are added and shielding increases |
| What is covalent radius | A radius concept based on the distance between nuclei of covalently bonded atoms |
| What is metallic radius | A radius concept used for atoms in metallic structures |
| What is van der Waals radius | A radius associated with non-bonded atoms in close contact |
| How does a cation's radius compare with its parent atom | It is smaller |
| Why is a cation generally smaller than its parent atom | Electron loss reduces electron-electron repulsion and can remove the outer shell |
| How does an anion's radius compare with its parent atom | It is larger |
| Why is an anion generally larger than its parent atom | Added electrons increase electron-electron repulsion |
| What is the key rule for comparing radii of isoelectronic species | For the same number of electrons higher nuclear charge gives smaller radius |
| Arrange O²⁻ F⁻ Na⁺ Mg²⁺ and Al³⁺ in decreasing ionic radius | O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺ |
| Why is O²⁻ larger than Al³⁺ despite both belonging to the same isoelectronic series | They have the same number of electrons but Al³⁺ has a much greater nuclear charge and attracts them more strongly |
| What is ionisation enthalpy | The enthalpy required to remove the most loosely bound electron from an isolated gaseous atom |
| What is the equation for first ionisation enthalpy | M(g) → M⁺(g) + e⁻ |
| What is the general trend of first ionisation enthalpy across a period | It generally increases |
| What is the general trend of ionisation enthalpy down a group | It generally decreases |
| Why does ionisation enthalpy generally increase across a period | Effective nuclear attraction generally increases and atomic size generally decreases |
| Why does ionisation enthalpy generally decrease down a group | Atomic size and shielding increase making electron removal easier |
| Why can stability of electronic configurations affect ionisation enthalpy | Especially stable filled or half-filled subshells can make electron removal less favourable |
| Why is the ionisation enthalpy of B lower than Be | The electron removed from B is a higher-energy p electron whereas Be has a filled 2s subshell |
| Why is the ionisation enthalpy of Al lower than Mg | The electron removed from Al is a 3p electron whereas Mg has a filled 3s subshell |
| Why is the ionisation enthalpy of O lower than N | O has paired electrons in a p orbital producing extra electron-electron repulsion whereas N has a more stable half-filled p subshell |
| Why is the ionisation enthalpy of S lower than P | S has paired p electrons whereas P has a more stable half-filled p subshell |
| What is the key subshell explanation for the Be to B ionisation anomaly | Removal from a p subshell is easier than removal from a filled s subshell |
| What is the key explanation for the N to O ionisation anomaly | Electron-electron repulsion in the paired p orbital makes removal from O easier |
| What is electron gain enthalpy | The enthalpy change when an electron is added to an isolated gaseous atom |
| What is the general trend of electron gain enthalpy across a period | It tends to become more favourable or more negative with important exceptions |
| Why should fluorine not automatically be assumed to have the most negative electron gain enthalpy among halogens | Its very small size produces greater electron-electron repulsion in the compact valence shell |
| What important comparison involving fluorine and chlorine must be remembered | Chlorine can have a more negative electron gain enthalpy than fluorine |
| Which elements should be checked as important electron-gain-enthalpy exceptions in this chapter | Be Mg N P noble gases and the F versus Cl comparison |
| Why are noble gases exceptional in electron gain enthalpy trends | Their stable filled valence shells make additional electron accommodation unfavourable |
| What is electronegativity | The tendency of an atom in a molecule to attract shared electrons toward itself |
| What is the general trend of electronegativity across a period | It generally increases |
| What is the general trend of electronegativity down a group | It generally decreases |
| Which element has the highest electronegativity | Fluorine |
| Why is electronegativity not identical to electron gain enthalpy | Electronegativity describes attraction for shared electrons in a molecule whereas electron gain enthalpy is an energy change for adding an electron to an isolated gaseous atom |
| How does metallic character change across a period | It decreases |
| How does metallic character change down a group | It increases |
| Why does metallic character decrease across a period | Increasing effective nuclear attraction makes electron loss less easy |
| Why does metallic character increase down a group | Larger atomic size and increased shielding make electron loss easier |
| How does non-metallic character change across a period | It increases |
| How does non-metallic character change down a group | It decreases |
| What are common characteristics of metals listed in the source | They are generally good conductors and usually malleable and ductile |
| Which elements are listed as metalloids in the source | Si Ge As Sb and Te |
| What is the general relationship between valence-shell configuration and common valency for main-group elements | Valency and common oxidation states are related to the number of valence electrons |
| What do metallic elements generally tend to do chemically | Lose electrons |
| What do non-metallic elements generally tend to do chemically | Gain or share electrons |
| Why do second-period elements show anomalous behaviour compared with heavier group members | They have small atomic size high ionisation enthalpy high electronegativity and no d orbitals in the valence shell |
| Which second-period element is compared with other alkali metals for anomalous behaviour | Li |
| Which second-period element is compared with other alkaline-earth metals | Be |
| Which Group 13 second-period element is compared with heavier members | B |
| What are the classic diagonal relationships listed in the source | Li and Mg Be and Al B and Si |
| Why can Li resemble Mg despite belonging to different groups | Their sizes and charge-to-radius characteristics can be similar producing related chemical behaviour |
| What are the main factors to check when explaining a periodic trend | Nuclear charge effective nuclear charge shielding number of shells atomic size electron-electron repulsion subshell stability and s-p differences where relevant |
| Why is effective nuclear charge important for periodic trends | It determines how strongly the nucleus attracts valence electrons after accounting for shielding |
| Why does adding shells generally increase atomic radius | The outer electrons occupy orbitals farther from the nucleus |
| Why does electron-electron repulsion matter for ionisation enthalpy | Repulsion can make a particular electron easier to remove |
| Why does half-filled subshell stability matter in periodic comparisons | A half-filled arrangement can have enhanced stability and make electron removal less favourable |
| Why does penetration matter when comparing s and p electrons | s electrons generally experience stronger nuclear attraction because they penetrate closer to the nucleus |
| What is the main trend of ionic radius down a group | It generally increases |
| Why cannot ionic radius be assigned one simple across-period trend for all ions | The trend depends on charge and electronic structure of the particular ions being compared |
| What is the most important rule for an isoelectronic series | Higher nuclear charge corresponds to smaller ionic radius |
| If two species have the same number of electrons which one is larger | The species with lower nuclear charge |
| If two species have the same nuclear charge but different numbers of electrons which generally has the larger radius | The species with more electrons |
| Why is the second-period series especially important for NEET periodic-trend questions | Its small size and unusual electronic structure create several important anomalies |
| Which second-period elements lack d orbitals in their valence shell | All second-period elements |
| What is the major periodic-table distinction between a period and a group | A period is horizontal while a group is vertical |
| Why do elements in a group tend to have related chemical properties despite increasing atomic size down the group | Their valence-shell electronic configurations remain related |
| What is the key difference between atomic radius and ionic radius | Atomic radius describes an atom whereas ionic radius describes the size of a charged ion |
| Why does losing an electron sometimes cause a dramatic decrease in radius | The outermost shell may be removed entirely leaving a smaller electron shell structure |
| What makes an isoelectronic comparison particularly useful | Electron number is held constant so nuclear charge becomes the major size-determining variable |
| Which periodic trend generally increases both across a period and decreases down a group | Electronegativity and ionisation enthalpy |
| Which periodic trend generally decreases across a period and increases down a group | Metallic character |
| Which periodic property has the strongest classic exception involving N and O | Ionisation enthalpy |
| Which periodic property has the classic F versus Cl exception highlighted in the source | Electron gain enthalpy |
| Why is helium's placement a special case in block/group identification | Its configuration is 1s² which resembles an s-block configuration but its chemical placement is Group 18 because of noble-gas stability |
| What information can electronic configuration provide simultaneously | Period group block and broad element type |
| What is the differentiating-electron criterion for identifying a block | The subshell into which the last differentiating electron enters |
| How can a d-block group number broadly be estimated | By adding electrons in the (n−1)d and ns orbitals |
| What is the general configuration of a d-block element | The differentiating electron enters (n−1)d with ns containing 0 to 2 electrons |
| What is the general filling subshell for f-block elements | The (n−2)f subshell |
| Why are d-block elements called transition elements in the broad classification | They occupy the middle region between the s and p blocks and involve d-subshell filling |
| What should be checked when an electronic configuration question asks for period | The highest principal quantum number present |
| What should be checked when it asks for block | The subshell receiving the differentiating electron |
| What should be checked when it asks for group of an s-block element | The number of valence s electrons |
| What should be checked when it asks for group of a p-block element | The ns²np¹ through ns²np⁶ pattern |
| What is the maximum number of orbitals in a shell with principal quantum number n | n² |
| What is the maximum number of electrons in a shell with principal quantum number n | 2n² |
| How are number of orbitals and maximum electrons in a shell related | Each orbital holds two electrons so a shell with n² orbitals can hold 2n² electrons |
| How many orbitals are in the third shell | 9 |
| How many electrons can the third shell hold | 18 |
| How many subshells are in the fourth shell | 4 |
| Which quantum number determines the number of orbitals in a subshell | Magnetic quantum number through the relation 2l + 1 |
| Which quantum number determines the subshell type | Azimuthal quantum number l |
| Which quantum number identifies the shell | Principal quantum number n |
| Which quantum number identifies electron spin orientation | Spin quantum number m_s |
| For n = 4 what are the possible l values | 0 1 2 and 3 |
| For l = 2 what are the possible m_l values | −2 −1 0 +1 and +2 |
| For l = 3 how many orbitals exist | 7 |
| For l = 3 how many electrons can the subshell hold | 14 |
| For n = 3 what is the maximum possible l value | 2 |
| For n = 1 what is the only possible l value | 0 |
| Can a 2d subshell exist according to the quantum-number rules | No because for n = 2 l can only be 0 or 1 |
| Can a 3f subshell exist according to the quantum-number rules | No because for n = 3 l can only be 0 1 or 2 |
| Can a 4f subshell exist | Yes because n = 4 permits l = 3 |
| How can you quickly test whether a set of n and l values is valid | Check that l lies from 0 through n − 1 |
| How can you quickly test whether a set of l and m_l values is valid | Check that m_l lies from −l through +l |
| How can you quickly test whether a proposed spin quantum number is valid | It must be either +1/2 or −1/2 |
| Why can no two electrons in one orbital have the same spin quantum number | Pauli exclusion principle requires opposite spins if they occupy the same orbital |
| What is the relationship between the number of orbitals in s p d and f subshells | 1 3 5 and 7 respectively |
| What is the relationship between the maximum electrons in s p d and f subshells | 2 6 10 and 14 respectively |
| Why are orbitals in the same subshell degenerate under appropriate conditions | They have the same energy but differ in spatial orientation |
| What does degeneracy mean in the context of orbitals | Different orbitals have the same energy |
| Why do p orbitals differ despite having the same energy within a subshell | They differ in spatial orientation |
| Which has more radial nodes 4s or 3d | 4s has 3 radial nodes whereas 3d has 0 |
| Which has more angular nodes 3d or 4s | 3d |
| What is the total number of nodes in any orbital with n = 4 | 3 |
| What is the total number of nodes in a 5p orbital | 4 |
| What is the number of angular nodes in a 5p orbital | 1 |
| What is the number of radial nodes in a 5p orbital | 3 |
| What is the number of radial nodes in a 4d orbital | 1 |
| What is the number of angular nodes in a 4d orbital | 2 |
| What is the total number of nodes in a 4d orbital | 3 |
| Why are hydrogen-like orbital energies different from multi-electron orbital energies | Hydrogen-like atoms depend mainly on n while multi-electron atoms also experience subshell-dependent shielding and penetration effects |
| What happens to electron energy when n increases in a hydrogen-like atom | The energy becomes higher or less negative |
| What happens to the radius of a hydrogen-like orbit when n increases | It increases |
| What happens to hydrogen-like orbital radius when nuclear charge increases for the same n | The radius decreases |
| What happens to hydrogen-like orbital energy when nuclear charge increases for the same n | The energy becomes more negative and the electron is more strongly bound |
| What is the conceptual relationship between excitation and emission | Excitation requires energy absorption whereas emission occurs when an electron returns to a lower energy state |
| How can the wavelength of emitted radiation be related to the energy difference between levels | λ = hc divided by ΔE |
| Why does a larger energy gap correspond to shorter wavelength radiation | Photon energy is inversely proportional to wavelength |
| What does the visible Balmer series tell us about hydrogen's energy structure | Transitions ending at n = 2 produce visible lines because the allowed energy differences correspond to visible wavelengths |
| Why is the Lyman series in the ultraviolet region | Transitions ending at n = 1 involve relatively large energy differences and therefore higher-energy shorter-wavelength radiation |
| What is the key conceptual reason Rutherford's atom would collapse classically | An accelerating charged electron should radiate energy continuously |
| How did Bohr avoid the classical collapse problem | He postulated stationary states in which electrons do not radiate energy |
| Why was Bohr's model a major improvement over Rutherford's | It introduced quantised energy levels and explained the hydrogen line spectrum |
| Why did quantum mechanics replace the idea of fixed Bohr orbits | The uncertainty principle and wave nature of matter prevent assigning exact classical paths to electrons |
| What is the relationship between momentum and de Broglie wavelength | They are inversely proportional |
| If particle A has twice the momentum of particle B how do their de Broglie wavelengths compare | Particle A has half the wavelength of particle B |
| If two particles have the same momentum what can be said about their de Broglie wavelengths | They have the same de Broglie wavelength |
| If two particles have the same mass and velocity what can be said about their de Broglie wavelengths | They have the same de Broglie wavelength |
| What happens to minimum position uncertainty if momentum uncertainty is reduced | Position uncertainty must increase |
| Why does the uncertainty principle become especially relevant for electrons | Their microscopic scale makes quantum uncertainty fundamental to their description |
| What is the relationship between position and momentum uncertainties | Their product cannot be smaller than h divided by 4π |
| What is the central conceptual difference between probability density and an electron's classical position | Probability density gives likelihood of finding the electron in a region rather than a definite trajectory |
| What is the main conceptual bridge from Bohr's model to the quantum mechanical model | The transition from fixed quantised orbits to probability-based orbitals |
| What is the most important filling-order trap involving 4s and 3d | 4s fills before 3d according to the standard energy ordering listed in the source |
| What is the most important filling-order trap involving 4f and 5d | 4f fills before 5d |
| Why does the n+l rule not simply mean lower n always fills first | Orbital energy in multi-electron atoms depends on both n and l |
| If two orbitals have n+l values 6 and 5 which fills first | The orbital with n+l = 5 |
| If two orbitals both have n+l = 6 but one has n = 4 and the other n = 5 which fills first | The one with n = 4 |
| Why does Hund's rule apply separately to p d and f subshells | Each contains degenerate orbitals that should receive single electrons before pairing |
| What is the correct qualitative distribution for d⁵ under Hund's rule | Five d orbitals each contain one unpaired electron |
| What is the correct qualitative distribution for d⁶ under Hund's rule | Five d orbitals receive one electron each first and the sixth electron pairs in one orbital |
| What is the key test for whether an orbital diagram obeys Pauli's principle | No orbital contains more than two electrons and paired electrons have opposite spins |
| What is the key test for whether an orbital diagram obeys Hund's rule | Degenerate orbitals are singly occupied before any pairing occurs |
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