Periodic Trends in Atomic Radii Ionic Radii Paramagnetism Ionization Energy Electron Affinity Metallic Character

Transcription

1 Periodic Trends in Atomic Radii Ionic Radii Paramagnetism Ionization Energy Electron Affinity Metallic Character

2 Shielding In a multi-electron system, electrons are simultaneously attracted to the nucleus and repelled by each other. Outer electrons are shielded from nucleus by the core electrons. The shielding causes the outer electrons to not experience the full strength of the nuclear charge.

3 Effective Nuclear Charge The effective nuclear charge is the net positive charge that is attracting a particular electron. Z is the nuclear charge S is the number of electrons in lower energy levels Zeffective = Z S Electrons in same energy level contribute to screening, but very little.

4 Screening & Effective Nuclear Charge

5 Trend in Atomic Radius Main Group There are several methods for measuring the radius of an atom, and they give slightly different numbers. Van der Waals radius = nonbonding Covalent radius = bonding radius Atomic radius as an average radius There is a General Trend: Atomic Radius Increases down group (top to bottom) Atomic Radius Decreases across period (left to right)

6 Trend in Atomic Radius Main Group

7 Quantum-Mechanical Explanation for the Group Trend in Atomic Radius The size of an atom is related to the distance the valence electrons are from the nucleus. The larger the orbital an electron is in, the farther its most probable distance will be from the nucleus and the less attraction it will have for the nucleus. Traversing down a group adds a principal energy level. The larger the principal energy level an orbital is in, the larger its volume. Therefore, quantum-mechanics predicts the atoms should get larger down a column.

8 Trend in Atomic Radius

9 Quantum-Mechanical Explanation for the Period Trend in Atomic Radius The larger the effective nuclear charge an electron experiences, the stronger the attraction it will have for the nucleus. The stronger the attraction the valence electrons have for the nucleus, the closer their average distance will be to the nucleus. Traversing across a period increases the effective nuclear charge on the valence electrons. Therefore, quantum-mechanics predicts the atoms should get smaller across a period.

10 Practice Choose the Larger Atom in Each Pair N or F C or Ge N or Al Al or Ge C or O Li or K C or Al Se or I

11 Trends in Atomic Radius Transition Metals Atoms in the same group increase in size down the column. Atomic radii of transition metals roughly the same size across the d block. Effective nuclear charge by the ns 2 electrons approximately the same.

12 Trend in Atomic Radius Main Group

13 Electron Configurations of Main Group Cations in Their Ground State Cations form when the atom loses electrons from the valence shell. Al atom = 1s 2 2s 2 2p 6 3s 2 3p 1 Al 3+ ion = 1s 2 2s 2 2p 6 [Ne]

14 Electron Configurations of Transition Metal Cations in Their Ground State The first electrons removed are the valence electrons. Electrons may also be removed from the sublevel closest to the valence shell after the valence electrons. The iron atom has two valence electrons: Fe atom = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 When iron forms a cation, it first loses its valence electrons: Fe 2+ cation = 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 However, it can then lose 3d electrons: Fe 3+ cation = 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5

15 Magnetic Properties of Transition Metal Atoms & Ions Electron configurations that result in unpaired electrons mean that the atom or ion will have a net magnetic field this is called paramagnetism. Electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field this is called diamagnetism.

16 Paramagnetism Illustrated (a) A sample (left side of balance) is weighed in the absence of a magnetic field. (b) When the field is turned on, the balanced condition is upset. The sample appears to gain weight. This is because it is now subjected to two attractive forces, the force of gravity and the force exerted by the magnetic field.

17 Transition Metal Atoms and Ions: Electron Configuration & Magnetic Properties Both Zn atoms and Zn 2+ ions are diamagnetic: Zn atoms [Ar]4s 2 3d 10 Zn 2+ ions [Ar]4s 0 3d 10 Ag forms both Ag + ions and, rarely, Ag 2+ Ag atoms [Kr]5s 1 4d 10 are paramagnetic Ag + ions [Kr]4d 10 are diamagnetic Ag 2+ ions [Kr]4d 9 are paramagnetic

18 Write the electron configuration and determine whether the Fe atom and Fe 3+ ion are paramagnetic or diamagnetic. Fe Z = 26 Previous noble gas = Ar (18 electrons) Fe atom = [Ar]4s 2 3d 6 Unpaired electrons Paramagnetic 4s 3d Fe 3+ ion = [Ar]4s 0 3d 5 Unpaired electrons Paramagnetic 4s 3d

19 Determine whether the following are paramagnetic or diamagnetic. Mn Mn = [Ar] 4s 2 3d 5 paramagnetic Sc 3+ Sc = [Ar]4s 2 3d 1 Sc 3+ = [Ar] diamagnetic

20 Trends in Ionic Radius Ions from elements in same group have same charge. Ion size increases down the column. Cations smaller than neutral atoms. Anions larger than neutral atoms. For isoelectric species, cations smaller than anions.

21 Radii of the Group VIIA atoms (Halogens) and Their Anions

22 Radii of Group IA Atoms and Their Cations Radii of Group IIA Atoms and Their Cations

23 Trends in Ionic Radius

24 Trends in Ionic Radius

25 Radii of Metal Atoms and Their Cations

26 Quantum-Mechanical Explanation for the Trends in Cation Radius When atoms form cations, the valence electrons are removed. The farthest electrons from the nucleus then are the p or d electrons in the (n 1) energy level. This results in the cation being smaller than the atom. The remaining electrons also experience a larger effective nuclear charge, shrinking the ion even more. Moving down a group, adds a new primary level, causing the cations to get larger. Moving to the right across a period, increases the effective nuclear charge for isoelectronic cations, causing the cations to get smaller.

27 Radii of Nonmetal Atoms and Their Anions

28 Quantum-Mechanical Explanation for the Trends in Anion Radius When atoms form anions, electrons are added to the valence shell. The new outer electrons experience a smaller effective nuclear charge than the old valence electrons, increasing the size. The result is that the anion is larger than the atom. Moving down a group, increases the n level, causing the anions to get larger. Moving to the right across a period, increases the effective nuclear charge for isoelectronic anions, causing the anions to get smaller.

29 S or S 2 Choose the larger of each pair S 2 is larger because there are more electrons (18 e ) for the 16 protons to hold Ca or Ca 2+ Ca is larger because its valence shell has been lost from Ca 2+ Br or Kr the Br is larger because it has fewer protons (35 p + ) to hold the 36 electrons than does Kr (36 p + )

30 Order the following sets by size (smallest to largest) Zr 4+, Ti 4+, Hf 4+ Na +, Mg 2+, F, Ne Ti 4+ < Zr 4+ < Hf 4+ Mg 2+ < Na + < Ne < F I, Br, Ga 3+, In + Ga 3+ < In + < Br < I

31 Ionization Energy Minimum energy needed to remove an electron from an atom or ion

32 Ionization Energy Measured in the gas state An endothermic process Valence electron easiest to remove, lowest IE M(g) + IE1 M 1+ (g) + 1 e - M +1 (g) + IE2 M 2+ (g) + 1 e - 1st IE energy = energy to remove electron from neutral atom 2nd IE = energy to remove from 1+ ion; etc.

33 General Trends in 1 st Ionization Energy The larger the effective nuclear charge on the electron, the more energy it takes to remove it. The farther the most probable distance the electron is from the nucleus, the less energy it takes to remove it. 1st IE decreases down the group. 1st IE generally increases across the period.

34 General Trends in 1 st Ionization Energy

35 Quantum-Mechanical Explanation for the Trends in First Ionization Energy The strength of attraction is related to the most probable distance the valence electrons are from the nucleus and the effective nuclear charge the valence electrons experience. The larger the orbital an electron is in, the farther its most probable distance will be from the nucleus and the less attraction it will have for the nucleus. Therefore, quantum-mechanics predicts the atom s first ionization energy should get lower down a column. Traversing across a period increases the effective nuclear charge on the valence electrons. Therefore, quantum-mechanics predicts the atom s first ionization energy should get larger across a period.

36 Choose the atom in each pair with the larger first ionization energy Al or S As or Sb N or Si O or Cl Mg or P Ag or Cu Ca or Rb P or Se

37 Trends in Successive Ionization Energies Removal of each successive electron costs more energy. 1)Shrinkage in size due to having more protons than electrons. 2) Outer electrons closer to the nucleus, therefore harder to remove. Regular increase in energy for each successive valence electron. Large increase in energy when start removing core electrons.

38 Trends in Successive Ionization Energies

39 Electron Affinity Energy released when an neutral atom gains an electron in the gas state: M(g) + 1e M 1 (g) + EA Defined as exothermic ( ), but may actually be endothermic (+) (some alkali earth metals & all noble gases) The more energy that is released, the larger the electron affinity. The more negative the number, the larger the EA.

40 Trends in Electron Affinity Generally decreases down a column. Generally increases across period. Highest EA in any period = halogen

41 Electron affinities of the main-group elements.

42 Trends in Electron Affinity Measured electron affinities for elements 1-57 and A negative value means that energy is released when an electron adds to an atom, while a value of zero means that energy is absorbed but the exact amount can t be measured experimentally.

43 Properties of Metals & Nonmetals Metals malleable & ductile shiny, lustrous, reflect light Nonmetals brittle in solid state dull, non-reflective solid surface conduct heat and electricity most oxides basic and ionic form cations in solution lose electrons in reactions oxidized electrical and thermal insulators most oxides are acidic form anions in solution gain electrons in reactions reduced

44 Metallic Character Metallic character is how closely an element s properties match the ideal properties of a metal. Metallic character decreases left-to-right across a period. Metallic character increases down the column.

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46 Quantum-Mechanical Explanation for the Trends in Metallic Character Metals generally have smaller first ionization energies and nonmetals generally have larger electron affinities except for the noble gases. Therefore, quantum-mechanics predicts the atom s metallic character should increase down a column because the valence electrons are not held as strongly. Therefore, quantum-mechanics predicts the atom s metallic character should decrease across a period because the valence electrons are held more strongly and the electron affinity increases.

47 Choose the more metallic element in each pair Sn or Te P or Sb Ge or In S or Br Mg or Al Si or Sn Br or Te Se or I

48 Atomic Properties- Summary of Trends in the Periodic Table

49 Atomic Properties- Summary of Trends in the Periodic Table

50 Atomic Properties- Summary of Trends in the Periodic Table

51 Atomic Properties- Summary of Trends in the Periodic Table

52 Trends in the Alkali Metals Atomic radius increases down the column. Ionization energy decreases down the column. Very low ionization energies. good reducing agents, easy to oxidize very reactive, not found uncombined in nature react with nonmetals to form salts compounds generally soluble in water Electron affinity decreases down the column. Melting point decreases down the column. All very low MP for metals. Density generally increases down the column.

53 Trends in the Halogens Atomic radius increases down the column. Ionization energy decreases down the column. Very high electron affinities. good oxidizing agents, easy to reduce very reactive, not found uncombined in nature react with metals to form salts compounds generally soluble in water found in seawater Reactivity increases down the column. React with hydrogen to form HX, acids. Melting point and boiling point increase down the column. Density increases down the column.

54 Reaction between potassium metal and bromine gas: K(s) + Br2(g)? K(s) + Br2(g) KBr 2 K(s) + Br2(g) 2 KBr(s) (ionic compounds are all solids at room temperature) Reaction between rubidium metal and liquid water: Rb(s) + H2O(l)? Rb(s) + H2O(l) Rb + (aq) + OH-(aq) + H2(g) 2 Rb(s) + 2 H2O(l) 2 Rb + (aq) + 2 OH-(aq) + H2(g) (alkali metal ionic compounds are soluble in water)

55 Trends in the Noble Gases Atomic radius increases down the column. Ionization energy decreases down the column. Very high IE Very unreactive Melting point and boiling point increase down the column. All gases at room temperature Very low boiling points Density increases down the column.