Unit 1: Structure of Matter Sections: 2.4, 8 Copyright © The McGraw-Hill Companies, Inc.  Permission required for reproduction or display.

When the Elements Were Discovered

The Modern Periodic Table Group Period Metal: good conductor of heat and electricity Nonmetal: usually a poor conductor of heat and electricity Metalloid: has properties that are intermediate between those of metals and nonmetals

Classification of the Elements

The Modern Periodic Table Alkali Earth Metal The Modern Periodic Table Halogen Noble Gas Alkali Metal

Ground State Electron Configurations of the Elements ns2np6 Ground State Electron Configurations of the Elements ns1 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 d10 d1 d5 Valence electrons: the outermost electrons 4f 5f Core electrons: all nonvalence electrons in an atom

Review of Concepts In viewing the periodic table, do chemical properties change more markedly across a period or down a group?

Group 1A Elements (ns1, n  2) Alkali Metals M M+1 + 1e- 2M(s) + 2H2O(l) 2MOH(aq) + H2(g) Low IE = very reactive Never found in the pure state 4M(s) + O2(g) 2M2O(s) Increasing reactivity

Group 1A Elements (ns1, n  2)

Group 2A Elements (ns2, n  2) Alkaline Earth Metals M M+2 + 2e- Be(s) + 2H2O(l) No Reaction Mg(s) + 2H2O(g) Mg(OH)2(aq) + H2(g) M(s) + 2H2O(l) M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba Increasing reactivity

Group 2A Elements (ns2, n  2)

Comparison of Group 1A and 1B The metals in these two groups have similar outer electron configurations, with one electron in the outermost s orbital. Chemical properties are quite different due to difference in the ionization energy. Lower IE1, more reactive

Group 3A Elements (ns2np1, n  2)

Group 3A Elements (ns2np1, n  2)

Group 4A Elements (ns2np2, n  2)

Group 4A Elements (ns2np2, n  2)

Group 5A Elements (ns2np3, n  2)

Group 5A Elements (ns2np3, n  2)

Group 6A Elements (ns2np4, n  2)

Group 6A Elements (ns2np4, n  2)

Group 7A Elements (ns2np5, n  2) Halogens X + 1e- X-1 X2(g) + H2(g) 2HX(g) Increasing reactivity

Group 7A Elements (ns2np5, n  2)

Group 8A Elements (ns2np6, n  2) Noble Gas Completely filled ns and np subshells. Highest ionization energy of all elements. No tendency to accept extra electrons.

Compounds of the Noble Gases A number of xenon compounds XeF4, XeO3, XeO4, XeOF4 exist. A few krypton compounds (KrF2, for example) have been prepared.

Practice Exercise 8.1 An atom of a certain element has 20 electrons. Without consulting a periodic table, answer the following questions: What is the ground-state electron configuration of the element? (b) How should the element be classified? (c) Is the element diamagnetic or paramagnetic?

Representing Free Elements in Chemical Equations Metals Nonmetals Do not exist in discrete molecular units Use empirical formulas in chemical equation Empirical formulas are the same as the symbols that represent the elements No single rule Complex three-dimensional network of atoms, use empirical formula Carbon Metalloids Diatomic and polyatomic molecules use molecular formulas HONClBrIF P4, S8 Noble gases (monatomic species) use their symbols

Electron Configurations of Cations and Anions Of Representative Elements Na [Ne]3s1 Na+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. Ca [Ar]4s2 Ca2+ [Ar] Al [Ne]3s23p1 Al3+ [Ne] H 1s1 H- 1s2 or [He] Atoms gain electrons so that anion has a noble-gas outer electron configuration. F 1s22s22p5 F- 1s22s22p6 or [Ne] O 1s22s22p4 O2- 1s22s22p6 or [Ne] N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Cations and Anions Of Representative Elements +1 +2 +3 -3 -2 -1

Isoelectronic: have the same number of electrons, and hence the same ground-state electron configuration Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne] O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

Electron Configurations of Cations of Transition Metals When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals. Fe: [Ar]4s23d6 Mn: [Ar]4s23d5 Fe2+: [Ar]4s03d6 or [Ar]3d6 Mn2+: [Ar]4s03d5 or [Ar]3d5 Fe3+: [Ar]4s03d5 or [Ar]3d5 *Most transition metals can form more than one cation and frequently the cations are not isoelectronic with the preceding noble gas *The order of electron filling does not determine or predict the order of electron removal for transition metals.

Identify the elements that fit the following descriptions: Review of Concepts Identify the elements that fit the following descriptions: An alkaline earth metal ion that is isoelectronic with Kr An anion with a -3 charge that is isoelectronic with K+ An ion with a +2 charge that is isoelectronic with Co3+

Effective nuclear charge (Zeff) is the “positive charge” felt by an electron. Zeff = Z - s 0 < s < Z (s = shielding constant) Zeff  Z – number of inner or core electrons Zeff Core Z Radius (pm) Na Mg Al Si 11 12 13 14 10 1 2 3 4 186 160 143 132

Trend in Effective Nuclear Charge (Zeff) increasing Zeff increasing Zeff *core electrons shield valence electrons much more than valence electrons shield one another

Atomic Radii covalent radius metallic radius Atomic radius is one-half the distance between the two nuclei in two adjacent metal atoms or in a diatomic molecule (or one-half the distance between the nuclei in two neighboring atoms).

Trends in Atomic Radii

Practice Exercise 8.2 Referring to a periodic table, arrange the following atoms in order of decreasing atomic radius: C, Li, Be. N < P < Si

Compare the size of each pair of atoms listed here: Review of Concepts Compare the size of each pair of atoms listed here: Be, Ba Al, S 12C, 13C

Video

Comparison of Atomic Radii with Ionic Radii

Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed.

The Radii (in pm) of Ions of Familiar Elements

Practice Exercise 8.3 For each of the following pairs, indicate which one of the two species is smaller: K+ or Li+ Au+ or Au3+ P3− or N3−

Review of Concepts Identify the spheres shown here with each of the following: S2−, Mg2+, F−, Na+

IE1 first ionization energy Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state. *the magnitude of ionization energy is a measure of how “tightly” the electron is held in the atom IE1 + X (g) X+(g) + e- IE1 first ionization energy IE2 + X+(g) X2+(g) + e- IE2 second ionization energy IE3 + X2+(g) X3+(g) + e- IE3 third ionization energy IE1 < IE2 < IE3 *Chemical properties of any atom are determined by the configuration of the atom’s valence electrons. The stability of these outermost electrons is reflected directly in the atom’s ionization energies.

*Note: while valence electrons are relatively easy to remove from the atom, core electrons are much harder to remove. Thus, there is a large jump in ionization energy between the last valence electron and the first core electron.

Variation of the First Ionization Energy with Atomic Number Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell

General Trends in First Ionization Energies Increasing First Ionization Energy Increasing First Ionization Energy

Exceptions Between Group 2A and 3A (ex: Be and B) Group 3A is lower b/c of the single electron in the outermost p subshell which is shielded Between Group 5A and 6A (ex: N and O) Group 6A doubles up one electron, the proximity of two electrons in the same orbital results in greater electrostatic repulsion and lower energy

Practice Exercise 8.4 Which atom should have a larger first ionization energy: nitrogen or phosphorus? Which atom should have a smaller second ionization energy: sodium or magnesium?

Review of Concepts Label the plots shown here for the firs, second, and third ionization energies for Mg, Al, and K.

*lg. pos. electron affinity means that the negative ion is very stable Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. *or think of it as the energy that must be supplied to remove an electron from the anion. X (g) + e− X− (g) F (g) + e− F− (g) DH = -328 kJ/mol EA = +328 kJ/mol *lg. pos. electron affinity means that the negative ion is very stable O (g) + e− O− (g) DH = -141 kJ/mol EA = +141 kJ/mol *The more positive the electron affinity is of an element, the greater the affinity of an atom of the element to accept an electron.

Variation of Electron Affinity With Atomic Number (H – Ba)

Practice Exercise 8.5 Why are the electron affinities of the alkaline earth metals, shown in Table 8.3, either negative or small positive values? Is it likely that Ar will form the anion Ar−

Review of Concepts Why is it possible to measure the successive ionization energies of an atom until all the electrons are removed, but it becomes increasingly difficult and often impossible to measure the electron affinity of an atom beyond the first stage?

Diagonal Relationships on the Periodic Table *similarities between pairs of elements in different groups and periods of the periodic table