Bonding and Periodic Table Trends

1. Bonding and Periodic Table Trends Honors Chemistry

2. Electron Configurations • Stable Octet: 8 electrons in the outer level is very stable (includes He) • Ions – gain/lose electrons to achieve a stable octet • Isoelectronic – same electron configuration • Examples: N, O, F, Na, Mg, Al are isoelectronic with Ne – this is called an isoelectronic series • Pseudoisoelectronic – same electron configuration but includes the d orbitals • Fe+2 is pseudoisoelectronic with Ar

3. Periodic Table Trends Introduction • Properties are periodic • An element’s position & its properties are a result of its electrons • The outermost electrons, aka valence electrons, have the greatest influence on the properties of the elements. • Adding an electron to an inner core orbital results in less striking changes in properties than adding an electron to an outer valence orbital (higher energy). • Shielding Effect: electrons in the lower energy levels (inner core electrons), shield electrons in the outer levels from the full effect of the nuclear charge.

4. Trends in the Periodic Table • Atomic Radius • The distance from the center of the nucleus to the outermost electron. • Bond Radius • Atoms get larger going down a group and smaller going across a period. Ex) Na is larger than Mg Na is smaller than K Ga vs. Al

5. Atomic Radii of the Representative Elements

6. Atomic Radii vs Atomic Number

7. Positive Ion Size • When atoms lose electrons, they become (positive) and get smaller. • The sizes of cations increases down a group. • The sizes of cations decreases across a period.

8. Negative Ionic Size • When atoms gain electrons, they become (negative) and get larger. • The sizes of anions increases down a group. • The sizes of anions decreases across a period.

9. Relative Sizes of Positive & Negative Ions The sodium ion lost an electron, and therefore the positive protons in the nucleus exert a stronger pull on the remaining negative electrons, shrinking the orbitals. Thus positive ions are smaller than their atoms. The chloride ion gained an electron, and therefore the fewer positive protons in the nucleus exert a weaker pull on the extra negative electrons, increasing the size of the orbitals. Thus negative ions are larger than their atoms.

10. Within an isoelectronic series, radii decrease with increasing atomic number because of increasing nuclear charge. N-3 O-2 F-1 Na+1 Mg+2 Al+3 How many electrons? Nuclear charge? > > > > >

11. Ionic Radius • Cations & anions decrease in size going across a period • Cations & anions increase in size going down a group

12. Electron Attraction in a Bond & Ion Size

13. Ionization Energy (IE): • The energy needed to remove one electron from an atom. (kJ/mole) • IE measures how tightly electrons are bound to an atom. • Elements that do not want to lose their electrons have high ionization energies. • Elements that easily lose electrons have low ionization energies. X + energy  X+1 + 1 e-

14. 1st Ionization Energy • I.E. decreases down a group. • I.E. increases across a period • Account for deviations across a period. • Metals tend to have low IE1. • Nonmetals tend to have high IE1.

16. Ionization Energy of the 1st 20 Elements

17. Successive Ionization Energies: • Energy required to remove electrons beyond the 1st electron. • Ionization energies will increase for every electron removed. X + IE1 X+1 + 1 e- X+1 + IE2 X+2 + 1 e- X+2 + IE3 X+3 + 1 e-

18. Successive Ionization Energies: • Electron Configuration • Na: [Ne] 3s1 • Mg: [Ne] 3s2 • Al: [Ne] 3s23p1

19. Ionization Energy vs. Atomic Number Notice the dips across the period… why?

20. Period 3 Na - [Ne] 3s1 ___ ___ ___ ___ Mg - [Ne] 3s2 ___ ___ ___ ___ Al - [Ne] 3s23p1 ___ ___ ___ ___ Si - [Ne] 3s23p2___ ___ ___ ___ P - [Ne] 3s23p3___ ___ ___ ___ S - [Ne] 3s23p4 ___ ___ ___ ___ Cl - [Ne] 3s23p5___ ___ ___ ___ Ar - [Ne] 3s23p6___ ___ ___ ___ 3s 3p

21. Electronegativity (EN) • Reflects an atoms ability to attract electrons in a chemical bond. • Up to 4.0 for F • Zero for He, Ne, Ar and Kr • Metals have low EN. • Nonmetals have high EN. • EN decreases down a group. • EN increases across a period.

23. Electron Affinity (EA) • Energy change that occurs when a neutral gaseous atom gains an electron. Units kJ/mol. 1 e- + X X-1 + EA • Most elements have no affinity for an additional electron and have an EA equal to zero. He(g) + e- He- EA = 0 kJ/mol He will not add an electron Cl(g) + e- Cl- + 349 kJ/mol EA = -349 kJ/mol Exothermic!!!

24. Electron Affinity (EA) • Metals have low EA. • Nonmetals have high EA. • EA decreases down a group. • EA increases (becomes more negative) across a period. • EXCLUDES noble gases • Exceptions: Groups IIA (~0) and VA (~0 for N and smallerfor P to Bi) • Why? Filled s and half filled p

26. Metallic Character • Reflected by those elements that can lose electrons easily. • Increases down a group. • Decreases across a period. • The most metallic metal is Cesium. • The most nonmetallic (least metallic) metal is Aluminum.

27. Reactivity • Related to the ability of an element to lose or gain an electron Brainiac Alkali MetalsAlkali Metals Reactivity • Reactivity of metals INCREASE down a group. • Reactivity of metals DECREASE across a period. • Reactivity of nonmetals DECREASE down a group. • Reactivity of nonmetals INCREASE across a period.

28. Bond Chemical 8 Chemical BondingChapter 6Honors Chemistry

29. Introduction • A chemical bond is an attractive force that holds atoms together in elements or compounds (to function as a unit) • intramolecular (within) vs. intermolecular (between) • Bond energy is the energy needed to break or form 1 mole of bonds in a gaseous substance (kJ/mol) • Bonding usually involves only the valence electrons. • In most compounds of the representative elements, the atoms have an electron configuration that is isoelectronic or psuedoisoelectronic with a noble gas • The manner in which atoms are bound together in a given substance has a profound effect on its chemical and physical properties.

30. Octet Rule • Many chemical compounds form such that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. • Very stable • There are exceptions: H, He, B (BF3) • System achieves the lowest possible energy.

31. Types of Chemical Bonds: Metallic Bonds Ionic Bonds Covalent Bonds

32. Metallic Bonds • Simplest crystalline solid – arranged in a very compact and orderly pattern • Sea of electrons – the valence electrons are mobile around metal cations • Electrons are delocalized • Attraction of the metal atoms and the surrounding sea of electrons

33. Metallic Bonds • Explains metallic properties: • High electrical and thermal conductivity (flow of electrons) • Luster (metals absorb wide range of  - excites e- and fall back emitting E in form of light  results in shiny appearance) • Ductility & malleability (mobility of e-, metallic bonding is same in all directions throughout solid) • Metallic bonding visual on Holt

34. Metallic properties due to sea of electrons Ionic compounds are hard but brittle – repulsions result from shift and causes crystal to break Metals vs. Ionic Crystals

35. Ionic bonding • Chemical bonding that results from an electrostatic attraction between cations and anions to form a neutral compound. • “salts” • Octet Rule • Atoms will transfer electrons (e-) to each other in order to have a full set of valence electrons. • When electrons are transferred, ionic bonds are formed.

36. Ionic bonding

37. Covalent Bonding • Sharing one or more electron pairs between 2 atoms

38. Characteristics of Ionic Compounds(hundreds of compounds) • All are high melting solids (>400°C). • Orderly 3D arrangements (pattern) called crystalline solid or crystal lattice. • Simplest arrangement = formula unit • High mp reflects strong bonds – large attractive forces are very stable • Many are white. • Colored compounds usually contain the transition elements (Cu, Cr, Co, Ni, Mn)

39. Characteristics of Ionic Compounds: • Solubility • Many are soluble in polar solvents, such as water (aka aqueous solutions) • Most are insoluble in nonpolar solvents, such as hexane (C6H14) Dissolving Salt Animation

40. Characteristics of Ionic Compounds: • Conductivity • Solids are non conductive – ions cannot move freely • Molten compounds are conductive – ions move freely (NaCl mp ~800°C) • Aqueous solutions are conductive – ions free to move in solution

41. Formation of Ionic Compounds • Formation results from a transfer of electrons and the electrostatic attractions of the closely packed, oppositely charged ions. • Ionic substances are formed when an atom that loses electrons relatively easily reacts with an atoms that has a high affinity for electrons. • Forms between a metal and a nonmetal. (large EN difference) • Metal: low IE, EN, EA • Nonmetal: high IE, EN, EA • Metal is oxidized (loss of e-) and nonmetal is reduced (gain of e-) • The ion pair has lower energy than separated ions.

42. Electron Configuration Distribution of electron density • Na: 1s22s22p63s1 • 186 pm • Cl: 1s22s22p63s23p5 • 99 pm • Na+1: 1s22s22p6 • 95 pm • Cl-1: 1s22s22p63s23p6 • 181 pm

43. Lewis structures or electron-dot structures

44. Lewis structures examples: • Sodium and chlorine • Potassium and phosphorus

45. Ionic, Nonpolar Covalent, Polar Covalent

46. HOLT

48. Nonpolar Covalent Bonds • Electron pair is shared equally between the atoms (ΔEN = 0 to ~0.4) • Diatomic molecules (H O F Br I N Cl); allotropes (S8)

49. Nonpolar Covalent BondsHydrogen

50. Polar Covalent Bonds • Electron pair is shared unequally between atoms (ΔEN = ~0.4 to ~1.9) • Results in an electric dipole (2 poles) • Equal but opposite charges that are separated by a short distance • Separation of charge between 2 covalently bonded atoms • Examples: HF, HBr, H2O