TRENDS IN ATOMIC PROPERTIES: THE PERIODIC TABLE

TRENDS IN ATOMIC PROPERTIES: THE PERIODIC TABLE Electron configurations determine organization of the periodic table Next properties of elements and their periodic behavior Elemental properties determined by: size (n) and shape (l) of orbitals atomic number (nuclear charge) Atomic sizes Ionization energies Electron affinities

IONIZATION ENERGY Definition: energy needed to remove electron 1st I.E. Remove 1st electron 2nd I.E. Remove 2nd electron etc. Outer electrons more easily removed EXAMPLE: Mg (Ne) 3s 2 1s 2 2s 2 2p 6 3s 2 1st I.E. = 735 kj/mol 2nd I.E. = 1445 kj/mol 3rd I.E. = 7730 kj/mol

IONIZATION ENERGY DOWN A FAMILY I. E. (kj/mol) Li 510 Na 490 K 418 Rb 403 Cs 375 smallest largest SIZE Electron further from nucleus easier to remove Cs electron furthest from nucleus

IONIZATION ENERGY RELATED TO RADIUS Coulomb s Law QQ 1 2 E = C d I. E. ~ inversely proportional to radius atomic radii

IONIZATION ENERGY ACROSS A PERIOD 1st I.E. (kj/mol) Na Mg Al Si P S Cl Ar 490 735 580 780 1060 1005 1225 1550 I.E. increases left to right (exceptions) Atomic size decreases left to right Electron further from nucleus easier to remove Exceptions: it takes extra energy to remove electrons from filled subshell (Mg, Ar) or from half-filled subshell (P)

ALKALI METALS REACTIVITY 2 Li(s) + 2 H 2 O(l) 2 Li + (aq) + 2 OH (aq) + H 2 (g) 2 Na(s) + 2 H 2 O(l) 2 Na + (aq) + 2 OH (aq) + H 2 (g) 2 K(s) + 2 H 2 O(l) 2 K + (aq) + 2 OH (aq) + H 2 (g) Li Na K Rb Be Mg Ca Sr IE decreases Reactivity increases Reactivity of the alkali metals increases as the ionization energy decreases

ELECTRON AFFINITY Definition: energy needed to add an electron to an atom or ion in gas phase Cl + e Cl (Ne)3s 2 3p 5 (Ne)3s 2 3p 6 E.A. = 349 kj/mol energy released E.A. s have periodic variation Halogens have large negative values: want to fill subshell p 6 Be, Mg have positive values: do not want to start new subshell p 1 to make negative ions

ELECTRON AFFINITY PERIODIC TRENDS X + e X

HALOGEN REACTIVITY 2 KBr(aq) + Cl 2 (aq) 2 KCl(aq) + Br 2 (aq) 2 KI(aq) + Cl 2 (aq) 2 KCl(aq) + I 2 (aq) 2 KI(aq) + Br 2 (aq) 2 KBr(aq) + I 2 (aq) F 2 Cl 2 Br 2 I 2 EA increases Reactivity increases Reactivity of the halogens increases as the electron affinity increases

CHEMICAL BONDS What is a chemical bond? force holding atoms or ions together Types of chemical bonds: ionic: electrostatic force between ions of opposite charge covalent: sharing of electrons between two bonded atoms metallic: found in metals

SOME COMMON CATIONS (Table 2.4) 1+ charge H +, Na +, K +, Ag +, NH 4 + 2+ charge Mg 2+, Ca 2+, Ba 2+, Fe 2+, Mn 2+ 3+ charge Al 3+, Cr 3+, Fe 3+ SOME COMMON ANIONS (Table 2.5) 1- charge Cl -, Br -, I -, CN -, OH -, NO 3-, MnO - 4 2- charge O 2-, S 2-, CO 2-3, CrO 2-4, SO 2-4 3- charge PO 4 3-

IONIZATION ENERGY AND ELECTRON AFFINITY Na Cl 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 5 valence electron e Na + Cl 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 6 Valence electrons Ionization Energy & Electron Affinity Na Na + + e I.E. low 496 kj/mol Cl + e Cl E.A. 349 kj/mol Source of electrons to form ions? Other atoms

Properties of elements are related to valence e Valence e are shown with Chlorine Lewis Symbols. Cl: : : [Ne]3s 2 3p 5 Sodium Na. [Ne]3s 1 Octet Rule: atoms gain or lose electrons to have octets (noble gas configuration) Each now has a noble gas configuration Na + smaller than Na Cl larger than Cl Ionic Bond electrostatic attraction

Electronic Configuration Of Some Common Ions Group Ion El. Config. Na IA Na + [Ne] Mg IIA Mg 2+ [Ne] Al IIIA Al3 + [Ne] F VIIA F - [Ne] O VIA O 2- [Ne] N VA N 3- [Ne] Composition of Ionic Compounds Al 3+ + O 2- Al 2 O 3 Al 2 O 3

Bonding Examples Using the periodic table, predict the chemical formula of the compound formed by the following elements. a) B and F b) Al and O c) Ca and N d) Li and P

TRANSITION METAL IONS s electrons are part of the valence electrons Fe [Ar] 4s 2 3d 6 Ag [Kr] 5s 1 4d 10 When forming an ion: s electrons are lost first then maybe d electrons Transition metal ions can have various charges EXAMPLES: Ag [Kr] 5s 1 4d 10 Ag + [Kr] 4d 10 Fe [Ar] 4s 2 3d 6 Fe 2+ [Ar] 3d 6 Fe 3+ [Ar] 3d 5 V [Ar] 4s 2 3d 3 has 5 oxidation states

VANADIUM electron configuration [Ar]4s 2 3d 3 what oxidation states are possible? start with VO 2 + V 5+ is reduced V 2+ yellow green blue blue purple green VO + 2 VO 2+ V 3+ V 2+ Point: transition metals have many oxidation states WHY?

BORN-HABER CYCLE Na(s) + 1 2 Cl 2(g) 1 Na(g) 2 Cl(g) 3 4 NaCl(s) 5 Cl (g) + Na + (g) kj 1 Na(s) Na(g) 111 2 ½Cl 2 (g) Cl(g) 120 3 Na(g) Na + (g) + e 495 4 Cl(g) + e Cl (g) 349 5 Na + (g) + Cl (g) NaCl(s) U Na(s) + ½ Cl 2 (g) NaCl(s) ΔH f = 377 kj U ΔH f (expt) = 411 kj (from Table) U = (377 + 411) = 788 kj Lattice energy

LATTICE ENERGY Lattice energy = the change in energy when an ionic solid is separated into isolated ions in the gas phase Na + (g) + Cl (g) NaCl(s) 788 kj/mol very endothermic

LATTICE ENERGY E el = κ Q 1 Q 2 d Different salts have different Q s and d s NaCl 788 kj MgO 3795 kj NaCl Q 1 Q 2 d = (1)(1) d why the huge difference (X5)? MgO Q 1 Q 2 d = (2)(2) d factor of 4 due to charges Na + Cl radii 0.97 A 1.81 A Mg 2+ O 2 radii 0.66 A 1.40 A 2.78 A 2.06 A d smaller for MgO

ION SIZES & IONIC BONDS ion sizes are important in ionic bonds size effects lattice energy lattice energy effects melting point MP NaF 943 C NaCl 801 C NaBr 747 C NaI 661 C larger anion MgO 2800 C charges +2 and -2

Ionic Sizes Predict the largest atom or ion in the following compounds and specify its charge a) LiBr or Br 2 b) RbF or RbBr c) BeO or BaO d) RbI or Xe

LATTICE ENERGY Rank the series of ionic compounds in order of increasing lattice energy. a) NaCl : RbI : LiF : KBr b) AlN : NaF : MgO : GaP c) CaO : KCl : KBr : MgO

CHEMICAL BONDS What is a chemical bond? force holding atoms or ions together Types of chemical bonds: ionic: electrostatic force between ions of opposite charge covalent: sharing of electrons between two bonded atoms metallic: found in metals

REPRESENTING MOLECULES CH 4 Molecular formula

COVALENT BONDING Electrons are shared by atoms Each atom in covalent molecule has a noble gas configuration Lewis structures & Octet Rule 2p 2 CH 4 H H: C: H H : : shared electrons are bonds 2p 3 NH 3 H H: N: H : : unshared electrons pairs shown 2p 4 H 2 O H:O : H : :

COVALENT MOLECULES and the octet rule methane CH 4 H H: C: H H : : H H C H H chlorine Cl 2 : Cl : : Cl: : : lone pairs Octet rule: most elements bond in covalent molecules so as to have 8 electrons in outer shell (hydrogen has 2) (are some exceptions) Covalent bonds are strong

LEWIS STRUCTURE RULES count number of valence electrons draw structure and connect atoms with single bonds distribute lone pairs to get octets add extra electrons to central atom use multiple bonds if necessary

MULTIPLE BONDS Single bond = sharing 1 pair of electrons Double bond = sharing 2 pairs of electrons Triple bond = sharing 3 pairs of electrons DOUBLE BOND EXAMPLES O=C=O O= CH 3 C OH acetic acid O=O O= CH 3 C CH 3 acetone H 2 C=CH 2 alkenes TRIPLE BOND EXAMPLES N N HC CH alkynes

LEWIS STRUCTURE EXAMPLE CCl 4 carbon tetrachloride C Cl Cl Cl Cl VE 4 7 7 7 7 = 32 Cl Cl C Cl Cl Used 4 pairs Have 12 more pairs Cl Cl C Cl OK Cl

LEWIS STRUCTURE EXAMPLE carbonyl chloride COCl 2 C O Cl Cl No. of VE 4 6 7 7 = 24

LEWIS STRUCTURE sulfite ion SO 3 2- No. of VE EXAMPLE S O O O

BOND POLARITY, ELECTRONEGATIVITY Electron sharing in covalent bonds depends on electronegativity of atoms Nonpolar bond: equal sharing Cl 2 Polar covalent bond: unequal sharing BrCl Br Cl δ+ δ Cl attracts shared electrons more than Br does Electronegativity: tendency of an atom to attract electrons to itself in a bond Is a relative scale F most electronegative atom 4.0

PERIODIC TRENDS IN ELECTRONEGATIVITY Fig. 8.6

BOND POLARITY EXERCISE Most to Least Polar? Least to Most Polar HI HF HCl HBr H 2 HF CsF RbF LiF NaF KF

FORMAL CHARGE How to decide between alternative Lewis structures? Use formal charge Formal Charge = # of valence e s on atom # of lone pair e s + ½# of bonding e s Rule 1: minimize formal charge EXAMPLE o c s o c s -1 0 +1 o c s o= c= s +1 0-1 0 0 0 BEST formal charges neutral

FORMAL CHARGE Rule 2: put negative charge on most electronegative atom EXAMPLE N C O -1 A) N C O -2 0 +1-1 B) N C O 0 0-1 -1 C) N= C= O -1 0 0-1