1. Periodicity of Atomic Properties Elements in the same group have the same number of valence electrons and related electron configurations; hence have similar chemical properties. The ground state electron configuration of the elements vary periodically with atomic number; all properties that depend on electron configuration tend to vary periodically with atomic number

3. Valence electrons - involved in bonding, determine reactivity Bi (Atomic number 83): [Xe] 4f 14 5d 10 6s 2 6p 3 6 valence electrons Os (Atomic number 76): [Xe] 4f 14 5d 6 6s 2 8 valence electrons

4. The variation of effective nuclear charge through the periodic table plays an important role in determining periodic trends Z eff increases from left to right across a period, but drops every time an outer electron occupies a new shell.

5. Atomic Radius Electron clouds do not have sharp boundaries, so do not really define the radius of an atom. Atomic radius - half the distance between the nuclei of neighboring atoms In solids centers of atoms are found at definite distances from one another.

6. For metals: atomic radius is half the distance between the centers of the atoms in the solid. For nonmetals the atomic radius is half the distance between the nuclei of atoms joined by a chemical bond; the covalent radius. van der Waals radius (used for noble gases): half the distance between the centers of the atoms in the solidified gas

7. Trend: Atomic radius generally decreases from left to right across a period and increases down a group Across a period: new electron added to the same shell, and nuclear charge is increasing

9. Ionic radius: effective radius of an ion in an ionic solid

10. Cations are smaller than the neutral atom; anions are larger Ionic radii generally increase down a group and decrease from left to right across a period.

11. Ionization Energies Ionization energy: energy needed to remove an electron from an atom in the gas phase. First ionization energy (I 1 ): energy required to remove an electron from the neutral gas phase atom Cu(g) -> Cu + (g) + e - energy = I 1 For Cu I 1 = 785 kJ/mol Second ionization energy (I 2 ): energy required to remove an electron from the singly charged gas phase atom Cu + (g) -> Cu 2+ (g) + e - energy = I 2 For Cu I 2 = 1955 kJ/mol

12. Trend: Ionization energy generally decreases down a group, and tends to increase moving across a period from left to right Departures from these trends can usually be traced to repulsion between electrons, particularly electrons occupying the same orbitals Elements with low ionization energies can be expected to form cations readily and to conduct electricity in their solid form.

14. Elements in the lower left of the periodic table tend to have lower ionization energies than those in the upper right. These are the elements in the s block, d block, f block and the lower left of the p block - metallic solids

15. Electron Affinity Electron attachment energy: energy change when a gas phase atom in its ground state gains a single electron X(g) + e - -> X - (g)  E = electron attachment energy Electron affinity = -  E(electron attachment) Electron affinity higher in the upper right

16. Main Group Elements (s and p) An s-block element has low ionization energy; outer electrons can easily be lost Group 1 form +1 ions; Group 2 form +2 ions An s block element is likely to be a reactive metal Since ionization energies are lowest at the bottom of each group, these elements are most reactive in the s-block p- block Elements on the left have low enough ionization energies to be metallic, but higher than the s-block elements and so are less reactive Elements on the right have high electron affinities (tend to gain electrons to form closed shell ions)

17. Transition Metals (d block) All d-block elements are metals; properties are transitional to s and p block Many d-block elements exhibit more than one oxidation state since the d-electrons have similar energies (inner shell) and a variable number can be lost

19. Chemical Bonds Chemical Bond: arise from sharing or transfer of electrons between two or more atoms Covalent bond: when electron density is mainly shared between two atoms Ionic bond: electron density mainly transferred from one atom to another.

20. Potential energy (kJ/mol)

21. Electronegativity: ability of an atom to attract electrons shared in a bond toward itself Elements to the lower left have low ionization energies and low electron affinities; tend to act as electron donors Elements to the upper right have high ionization energies and high electron affinities; tend to act as electron acceptors

22. R. Mulliken  = 0.5 (IE 1 + EA) L. Pauling: quantitative measure of electronegativity,  |  A -  B | = 0.102{D(A-B) - 0.5 [ D(A-A) + D(B-B)]} 0.5 A polar colavent bond is a bond with partial electric charges arising from their difference in electronegativity. The partial charges give rise to an electric dipole moment Pauling Scale for 

23. Bond dissociation energy  E d : energy required to separate the bonded atoms H-Cl(g) -> H(g) + Cl(g) Greater the  E d, stronger the bond Bond enthalpy: enthalpy change (heat absorbed at constant pressure) when bond is broken  E d =  H d - R T Since bond enthalpy of a bond X-Y between two atoms X and Y does not vary much between compounds, define an “average” bond enthalpy or “average” bond energy

24. Variations in atomic nuclei contribute to trends in bond energy

25. Number of bonds between two atoms also influence bond energy single < double < triple

26. Bond length: distance between the centers of two atoms joined by a covalent bond Bonds between larger atoms tend to be longer Multiple bonds are shorter than single bonds A bond length is approximately the sum of the covalent radii of the two atoms Covalent radius: effective radius of an atom in a covalent bond

27. Bond Order: number of shared electron pairs Single bond: bond order = 1 Double bond: bond order = 2 Triple bond: bond order = 3 Bond length decrease with bond order Bond strength increases with bond order

28. Chemical Bonding The Lewis model of the chemical bond assumes that each bonding electron is located between two bonded atoms - localized electron model. However, location of an electron in an atom cannot be described in terms of a precise position; same is true of electrons in a molecule. Valence Bond Theory - first quantum mechanical description of distribution of electrons in bonds.

29. Valence Bond Theory Minimum requirement for the formation of a covalent bond is two electrons and two valence orbitals that can overlap. The overlapping of the two orbitals concentrate the electron probability between the two nuclei, creating a chemical bond.

30. H 2 : two H atoms, each in the ground state 1s 1 In VB theory: the two H atoms come together and their atomic orbitals merge or overlap Greater the extent of overlap, greater the strength of the bond

31. Overlap of the two 1s orbitals form a sigma (  ) bond. In a sigma bond - electron density is distributed along the bond axis Note that the electron spin is paired

32. HF H: 1s 1 F: 1s 2 2s 2 2p 5 Overlap between the valence orbital of H (1s) and valence orbital of F (2p) to form a  bonds Note: electron spin is paired in the  orbital By definition: z is the direction along the internuclear axis