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Química Geral · Atoms, electrons and the periodic table

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O que você vai aprender

  • Estimate the effective nuclear charge on an outer electron as Z minus the core electrons and use it to explain trends across a period and down a group.
  • Rank atoms and ions by radius, including cations, anions, isoelectronic series and the transition-metal and lanthanide contractions.
  • Explain first and successive ionization energies, including the group 2/13 and 15/16 exceptions, and use ionization data to find the number of valence electrons.
  • Interpret electron affinities with their sign convention, including the fluorine/chlorine and nitrogen exceptions, and calculate simple energy balances from ionization energies and electron affinities.
  • Distinguish electronegativity from electron affinity and rank elements by electronegativity and metallic character.

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Effective nuclear charge and shielding

An outer electron does not feel the full nuclear charge Z, because the electrons between it and the nucleus repel it. The charge it does feel is the effective nuclear charge, Zeff. A simple estimate counts inner electrons as perfect shields and electrons in the same shell as no shield at all: Zeff ≈ Z − (core electrons), where the core electrons are all those in shells below the outermost occupied shell. Sodium (1s² 2s² 2p⁶ 3s¹) has 10 core electrons, so its 3s electron feels about 11 − 10 = +1; sulfur, 16 − 10 = +6. For a main-group atom this number equals the count of valence electrons. Across a period Zeff rises by one at each step, because a proton is added while the new electron joins the same shell. Down a group it stays about the same in this model (+1 for Li, Na and K), because the extra protons are matched by extra core electrons; more careful estimates show it rising slightly. The model also works for ions: in Mg²⁺ (1s² 2s² 2p⁶, drawn) the outermost electrons are in n = 2 and only the 1s² pair lies beneath them, so they feel about 12 − 2 = +10. Real shielding is imperfect, and electrons that penetrate the core feel more than this estimate, but the model gives every trend in the right direction.

Atomic and ionic radii

Atomic radius falls across a period, because the rising Zeff pulls one shell in more tightly, and rises down a group, because each step adds a shell with a larger n. Atoms at the lower left of the table are the largest and those at the upper right the smallest, so N < Si < Mg < Cs. A cation is smaller than its atom: lithium (about 152 pm) loses its only 2s electron, and Li⁺ (76 pm) is left with the 1s shell as its outermost. Losing more electrons shrinks an ion further, so Mn³⁺ (65 pm) is smaller than Mn²⁺ (83 pm). An anion is larger than its atom: S²⁻ (184 pm) has 18 electrons held by the same 16 protons as S (104 pm), so electron repulsion grows and each electron is held less firmly. In an isoelectronic series all members have the same number of electrons and only Z changes, so the radius falls as Z rises: Ca²⁺ 100 pm < K⁺ 138 pm < Cl⁻ 181 pm < S²⁻ 184 pm (drawn), with Z = 20, 19, 17 and 16. Across a transition series radii change little, because each added electron enters the inner (n − 1)d subshell, where it partly screens the outer s electrons from the added proton, so their Zeff rises only slowly. That screening is imperfect, and fourteen 4f electrons screen even more poorly, so after the lanthanides Zeff has risen enough that hafnium (about 159 pm) is no larger than zirconium (about 160 pm), the lanthanide contraction.

First ionization energy

The first ionization energy, IE₁, is the energy needed to remove one electron from each atom of a mole of gaseous atoms: X(g) → X⁺(g) + e⁻. It is always positive. It rises across a period as Zeff rises and falls down a group as the outer electron moves to a larger shell (Li 520, Na 496, K 419, Rb 403, Cs 376 kJ/mol). Two exceptions break the rise across a period. Group 13 is below group 2 (B 801 < Be 899; Al 578 < Mg 738 kJ/mol), because the electron removed from boron or aluminium is a p electron, higher in energy than the s pair beneath it and partly shielded by it. Group 16 is below group 15 (O 1314 < N 1402; S 1000 < P 1012 kJ/mol), because oxygen and sulfur lose an electron from a doubly occupied p orbital, where pairing repulsion makes it easier to remove; nitrogen and phosphorus have one electron in each p orbital. Below the d block the filled (n − 1)d¹⁰ subshell shields poorly, so gallium's IE₁ (578.8 kJ/mol) is slightly above aluminium's (577.5 kJ/mol). Ionization energies can also give a rough Zeff from the one-electron formula IE ≈ 1312 kJ/mol × Zeff²/n²: for potassium (n = 4, IE₁ 419 kJ/mol), Zeff ≈ 4 × √(419/1312) = 2.26, larger than the +1 of the simple model because the 4s electron penetrates the core.

Successive ionization energies

Each further electron is harder to remove, because it leaves an ion of higher charge, so IE₁ < IE₂ < IE₃ always. The rise is gradual while valence electrons are removed and jumps sharply when the first core electron is taken. Magnesium: IE₁ 738, IE₂ 1451, IE₃ 7733 kJ/mol. The first two come from 3s (Zeff about +2); the third comes from the 2p subshell, close to the nucleus, with Zeff about +10. Counting the electrons removed before the big jump gives the number of valence electrons, and so the group of a main-group element. Calcium shows the same pattern (590, 1145, 4912 kJ/mol), and making Ca²⁺(g) from Ca(g) costs IE₁ + IE₂ = 590 + 1145 = 1735 kJ/mol. Comparing second ionization energies across a period follows the same logic: IE₂ of sodium (4562 kJ/mol) is three times that of magnesium (1451 kJ/mol), because Na⁺ must give up a 2p core electron while Mg⁺ still has a 3s electron.

Electron affinity

Electron affinity describes X(g) + e⁻ → X⁻(g). In this topic it is reported as the energy released, so a positive electron affinity means the anion is lower in energy than the atom plus a free electron, and ΔE for the attachment is −EA. Chlorine has EA = 349 kJ/mol: ΔE = −349 kJ/mol for Cl(g) + e⁻ → Cl⁻(g) (drawn). Electron affinity generally rises across a period toward the halogens, but it is irregular. Beryllium, magnesium and the noble gases do not form stable gaseous anions, because the added electron would have to enter a new, well-shielded subshell (3p for Mg) or shell; the heavier group 2 atoms bind one only very weakly (Ca about 2 kJ/mol). Nitrogen (2p³) has essentially no electron affinity (about −7 kJ/mol), because the added electron must pair in a half-filled set, while carbon (2p², 122 kJ/mol) and oxygen (141 kJ/mol) bind one. Chlorine's value exceeds fluorine's (328 kJ/mol), because an electron added to fluorine's small, crowded n = 2 shell meets strong repulsion. Even the largest electron affinity does not pay for ionizing a metal: K(g) + Cl(g) → K⁺(g) + Cl⁻(g) has ΔE = IE₁(K) − EA(Cl) = 419 − 349 = +70 kJ/mol. Ionic compounds are stable because the attraction between the ions in the lattice releases far more energy.

Electronegativity and metallic character

Electronegativity measures how strongly an atom in a bond attracts the shared electron pair; it is not the same as electron affinity, which belongs to an isolated gaseous atom. On the Pauling scale it rises across a period and falls down a group: K 0.82, Mg 1.31, C 2.55, N 3.04, Cl 3.16, O 3.44 and F 3.98, the highest of all. Fluorine is more electronegative than chlorine even though chlorine has the larger electron affinity. Mulliken's definition shows why: it averages the two energies, (IE₁ + EA)/2, which for chlorine is (1251 + 349)/2 = 800 kJ/mol; fluorine's much larger ionization energy gives it the larger average. Metallic character, the tendency to lose electrons and form cations, follows the opposite direction: it rises down a group and from right to left, as ionization energy falls. Group 15 runs from the nonmetals N and P through the metalloids As and Sb to the metal Bi. Across period 3 the oxides change from basic (Na₂O, MgO) through amphoteric Al₂O₃, which dissolves in both strong acid and strong base, to acidic (SiO₂, P₄O₁₀, SO₃). The metalloids, B, Si, Ge, As, Sb and Te, lie along the boundary.

What to carry forward

Most trends follow from Zeff and n. Across a period Zeff rises within one shell, so atoms shrink, ionization energy, electron affinity and electronegativity rise, and metallic character falls; down a group a new shell is added, so atoms grow and the same properties reverse. Cations are smaller and anions larger than their atoms, and an isoelectronic series shrinks as Z rises. Watch the exceptions: group 13 below group 2 and group 16 below group 15 in IE₁, Cl above F and N near zero in electron affinity, and the big jump in successive ionization energies that marks the first core electron.

Problemas de prática (40)

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Preveja o produto 4

Ionization energies

  • Highest first ionization energy in period 4Predict which element in a period has the highest first ionization energy from its position.

Electron affinity and electronegativity

  • Which atom does not bind an extra electronPredict which atom has no favourable electron affinity from its valence configuration.
  • The most electronegative of N, O, S and ClPick the most electronegative element from a set that includes diagonal neighbours.

Metallic character and periodic reasoning

  • The most metallic of four elementsPredict which element has the greatest metallic character from its position in the table.

Classificar 3

Metallic character and periodic reasoning

  • The amphoteric oxide of period 3Classify period 3 oxides as basic, amphoteric or acidic from the metallic character of the element.

Electron affinity and electronegativity

  • The equation for an electron affinityIdentify the process whose energy change defines the electron affinity of an element.

Ionization energies

  • The equation for a first ionization energyIdentify the process whose energy change is the first ionization energy of an element.

Comparar 8

Electron affinity and electronegativity

  • Rank five atoms by electron affinityRank atoms by electron affinity, including the chlorine–fluorine exception.
  • Rank five elements by electronegativityRank elements by Pauling electronegativity from their positions in the periodic table.

Effective nuclear charge and size

  • Iron, iron(II) and iron(III) by sizeRank an atom and its cations of different charge by radius.
  • Rank a neon-like isoelectronic seriesRank the members of an isoelectronic series of ions by radius from their nuclear charges.
  • Rank five atoms by sizeRank main-group atoms from different periods and groups by atomic radius.

Metallic character and periodic reasoning

  • Metallic character down group 15Rank the elements of a group by metallic character.

Ionization energies

  • Period 3 ionization energies with their exceptionsRank period 3 elements by first ionization energy, including the group 2/13 and 15/16 exceptions.
  • Second ionization energies of sodium and magnesiumCompare the second ionization energies of neighbouring elements by asking which electron each ion loses.

Calcular 7

Ionization energies

  • Energy to make gaseous Mg²⁺Calculate the energy to form a gaseous dication by adding successive ionization energies.
  • Effective nuclear charge of sodium from its ionization energyEstimate an effective nuclear charge from a measured ionization energy with the hydrogen-like formula.

Electron affinity and electronegativity

  • Mulliken electronegativity of fluorineCalculate a Mulliken electronegativity as the average of ionization energy and electron affinity and compare two halogens.
  • Energy to make gaseous Na⁺ and Cl⁻Combine an ionization energy and an electron affinity, with correct signs, to find the energy of forming a gaseous ion pair.

Effective nuclear charge and size

  • How much sodium shrinks on ionizationCalculate the percentage decrease in radius when an atom becomes its cation.
  • Effective nuclear charge on a chlorine valence electronEstimate the effective nuclear charge on a valence electron with the simple model Zeff = Z − core electrons.
  • Effective nuclear charge in Na⁺Apply the simple Zeff model to the outermost electrons of a cation that has lost its valence shell.

Encontre o erro 3

Effective nuclear charge and size

  • Find the error: the size of chlorideIdentify the flaw in an argument that an anion is smaller than its parent atom.

Electron affinity and electronegativity

  • Find the error: electron affinity is not electronegativityDiagnose the error of inferring electronegativity order from electron affinities.

Metallic character and periodic reasoning

  • Find the error: gallium and the group trendDiagnose why a strict group trend fails for gallium by the poor shielding of 3d electrons.

Interprete as evidências 1

  • Valence electrons from successive ionization energiesRead the number of valence electrons from a set of successive ionization energies.

Trabalhe de trás para frente 2

Metallic character and periodic reasoning

  • Identify an element from two cluesCombine an ionization-energy exception with oxide behaviour to identify a period 3 element.

Ionization energies

  • Identify a period 4 element from its ionization energiesIdentify an element from the position of the large jump in its successive ionization energies.

Planeje uma síntese 1

  • Choose the comparison that shows pairing repulsionChoose the pair of ionization energies that isolates the effect of electron pairing in a p subshell.

Explique por quê 11

Ionization energies

  • Why boron ionizes more easily than berylliumExplain the group 2/13 dip in first ionization energy from the subshell of the electron removed.
  • Why ionization energy falls down group 1Explain the decrease in first ionization energy down a group.
  • The jump at magnesium's third ionizationExplain a large jump in successive ionization energies by the removal of a core electron.
  • Why oxygen ionizes more easily than nitrogenExplain the group 15/16 dip in first ionization energy by electron pairing in a p orbital.

Electron affinity and electronegativity

  • Why chlorine has the larger electron affinityExplain why chlorine releases more energy than fluorine on gaining an electron.
  • Why nitrogen does not bind an extra electronExplain the near-zero electron affinity of nitrogen from its half-filled 2p subshell.

Metallic character and periodic reasoning

  • Why group 14 becomes metallic toward the bottomExplain the increase in metallic character down a group in terms of ionization energy and atomic size.

Effective nuclear charge and size

  • Why K⁺ is much smaller than KExplain why a cation is much smaller than its parent atom when it loses its whole valence shell.
  • Why atoms shrink across period 3Explain the decrease in atomic radius across a period in terms of nuclear charge and shielding.
  • Why alkali metal atoms grow down the groupExplain why atomic radius increases down a group even though the nuclear charge increases.
  • Why hafnium is no larger than zirconiumExplain the near-equal radii of zirconium and hafnium by the poor shielding of 4f electrons.

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