Periodic Properties of Atoms Glenn V. Lo Department of Physical Sciences Nicholls State University

Ionization energy First ionization energy = energy needed to remove an electron from a neutral atom. Second ionization energy = energy needed to remove a second electron Third ionization energy = ? What do you call the energy associated with the following: Na(g)  Na + (g) + e - Na + (g)  Na 2+ (g) + e -

Ionization energy Unless otherwise specified, the term “ionization energy” refers to the first ionization energy and is minimum energy needed to remove an electron. The energy needed to remove an electron depends on the orbital occupied by the electron. Valence electrons are easier to remove than core electrons. Example: Na, 1s 2 2s 2 2p 6 3s 1 Which electron is easiest to remove?

Looking up ionization energy Values tabulated in kJ/mol. Ex. Na: kJ/mol NOTE: 1 mole of atoms = 6.02x10 23 atoms NIST atomic spectral database Values tabulated in eV (per atom)

Trends in first IE What is the trend? Same period, left-to- right? Same group, top-to- bottom? In any given row, which element has highest IE? Which has a higher IE, Na or K? Note: there are exceptions.

Trends in atomic sizes Based on the figure on the right, what is the general trend in (first) atomic size? Same row, left-to-right? Same group, top-to- bottom? Arrange the followoing in order of increasing atomic size: Na, Rb, Cl

Test Yourself The ionization energy of bromine atoms is the energy associated with which of the following? A. Br(g) → Br + (g) + e - B. Br 2 (l) → 2 Br + (g) + 2e - C. Br(g) + e - → Br - (g) D. Br 2 (l) + 2e - → 2 Br - (aq)

Test Yourself Which of the following has the highest ionization energy in the period where it belongs? A. Ne B. Na C. Br, D. Mn

Explaining Trend in IE Examine valence electrons and use Coulomb’s law: Farther from nucleus: weaker inward pull from nucleus Nuclear charge: more protons in nucleus – stronger inward pull Other electrons: partially shields outer electrons from pull of nucleus

Explaining Trend in IE Left-to-right. Outermost electrons are in the same valence shell: distance from nucleus is more or less the same More protons increases inward pull More electrons increases repulsions But… force due to additional proton is always inward, force due to other electrons is partly directed inward. In other words… one additional proton in the nucleus and one additional electron in the valence shell means a net increase in effective nuclear charge.

Explaining Trend in IE Top to bottom Outermost electrons significantly farther away: weaker pull towards the nucleus (Coulomb’s Law: force is inversely proportional to square of distance.) Effective nuclear charge: more or less the same

Explaining Trend in Size Same explanation as IE trend. Top to bottom: size of valence orbital increases significantly Left to right: increasing effective nuclear charge; electrons “pulled in” more strongly

Test Yourself Which of the following has the highest ionization energy? A. Na B. Rb C. N

Test Yourself Which of the following has the smallest atomic radius? A. O, B. S, C. Li

Test Yourself The increasing trend in ionization energy across a period, from left to right, is due to... A. Increasing repulsions among electrons in the valence shell B. Increasing effective nuclear charge C. Increasing number of core electrons D. Increasing size of the valence shell

Test Yourself The decreasing trend in ionization energy going down a group is, is due to... A. Increasing repulsions among electrons in the valence shell B. Increasing effective nuclear charge C. Decreasing number of core electrons D. Increasing size of valence shell

Explain ion formation Why do atoms belonging to the same group tend to form the same type of ions? group IA  ions with +1 charge group IIA  ions with +2 charge group VIIA  ions with –1 charge Answer: behavior of atoms depend primarily on valence configuration: same group  same valence configuration Ex. group IA: s 1, group IIA: s 2, group VIIA: s 2 p 5

Explaining cation formation It is extremely hard to remove an electron from a noble-gas-like configuration. Stable, naturally-occurring cations are formed when loss of electrons lead to noble-gas-like configuration. Example: Na: [Ne] 3s 1 (neutral atom) Na + : [Ne]; noble-gas-like Mg: [Ne] 3s 2 (neutral atom) Mg + : [Ne] 3s 1 ; not noble-gas-like Mg 2+ : [Ne]; noble-gas-like; Na + and Mg 2+ are isoelectronic with Ne

Example What is likely charge of beryllium ion based on the following data (from webelements.com)

Explaining anion formation Why does F tend to become F - ? F: [He] 2s 2 2p 5 (neutral atom) F 7+ : [He], but will require too much energy! F - : [He] 2s 2 2p 6 or [Ne] F 2- : [Ne] 3s 1 (valence electron very easily lost)

Electron Affinity EA = energy released when atom gains an electron Same in magnitude as energy needed to remove electron from ion with -1 charge. Trend similar to ionization energy: except for atoms with completely filled subshells (which have very low EA); also increases from row 2 to row 3

Ionic Bonding IE of Na = 5.1 eV, EA of F = 3.4 eV To transfer electron from Na to F requires 1.7 eV Atoms don’t like to form noble-gas-like ion???? Stability is achieved when cation and anion move closer together. Na + (g) + Cl - (g)  NaCl(g), 5.3 eV Na + (g) + Cl - (g)  NaCl(s), huge energy release 9.5 eV

Explaining transition metal ion formation Transition metals cations: usually +2. Valence s electrons are lost first. Example: Fe: [Ar] 4s 2 3d 6 Fe 2+ : [Ar] 3d 6 Other charges due to loss of inner d electrons. Fe 3+ : [Ar] 3d 5

Test Yourself What is the electron configuration of naturally occurring aluminum ion? A. [Ne] 3s 2 3p 1 B. [Ne] 3s 2 C. [Ne] 3s 1 D. [Ne]

Test Yourself Which atom's first five ionization energies best resembles the pattern shown in the graph? A. Na, B. Mg, C. Al, D. K

Test Yourself What is the electron configuration of Mn 2+ ? A. [Ar] 4s 2 3d 6 B. [Ar] 4s 2 3d 3 C. [Ar] 4s 1 3d 4 D. [Ar] 3d 5

Cations vs. Atoms Compared to atom it came from, a cation is smaller and has higher IE. Why? Higher effective nuclear charge on the remaining electrons. Compare: Mg, Mg +, Mg 2+

Example Compare: Mg 2+ vs. Ne

Anions vs. Atoms Compared to atom it came from, an anion is larger and has lower IE. Why? Smaller effective nuclear charge (more electrons; same #protons) If valence shell is filled, additional electron to to next shell; additional electron much farther away (easier to remove) Example: Compare: F, F -, F 2-

Example Compare: F 2- vs. Na

Test Yourself Which of the following has the largest size? A. K B. K + C. Se D. Li

Test Yourself Which of the following has the highest ionization energy? A. Ne B. Mg 2+ C. Mg D. F

Test Yourself Which of the following has the largest size? A. Cl B. Cl - C. F D. F -

Electronegativity Electronegativity: ability of an atom to attract shared electrons. Trend generally follows ionization energy. F is the most electronegative element. Why?

Test Yourself Which of the following has the highest electronegativity? A. Na B. Br C. F

Test Yourself Which type of element has the lowest electronegativity? A. a metal B. a nonmetal C. a metalloid