Electron Configuration and Periodic Properties

1. Electron Configuration and Periodic Properties • Atomic Radii • The size of an atom is defined by the edge of its orbital • Since this boundary is fuzzy, the radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together

2. Atoms tend to get smaller as you move across a period due to the increased positive charge • They get larger as you move down a group due to the increasing energy levels occupied

5. Ionization Energy • Ionization energy is the energy required to remove one electron from a neutral atom • Made on isolated atoms in the gas phase • In general, ionization energies of the main group elements (s&p) increase across a period • Generally decrease down a group

8. With sufficient energy, electrons can be removed from positive ions as well as from neutral atoms • The energies are referred to as the second ionization energy, third ionization energy, and so on • These energies generally increase due to the stronger effective nuclear charge • There are large jumps in energies when stable arrangements are ionized (in particular- the noble gas configurations)

10. Electron Affinity • The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity • Atoms that release energy have a negative affinity (they want the electron) • Atoms that require energy to “force” the electron on them have a positive affinity (they will lose the electron spontaneously)

11. The halogens gain electrons most readily • The p group elements generally become more negative as you move across a period (again exceptions caused by stable electron arrangements) • The trends in groups are not as regular (competing increased nuclear charge and atomic radius) • Generally the size predominates

12. For an isolated ion in the gas phase, it is always more difficult to add a second electron to an already negatively charged ion • Second affinities are therefore always positive • Ions like Cl-2 never occur

14. Ionic Radii • A positive ion is known as a cation • Caused by the loss of electrons • The remaining electrons are drawn closer to the nucleus by the unbalanced charge • A negative ion is known as an anion • Formed from the addition of extra electrons • The electrons are not drawn as tightly as they were before the addition

15. The metals on the left tend to form cations, while the nonmetals on the upper right tend to form anions • Cationic radii decrease across a period due to increasing nuclear charge • Anionic radii (starting w/ group 15) decrease across a period • Ionic radii tend to increase down a group

17. Valence Electrons • Chemical compounds form because electrons are lost, gained, or shared between the outermost energy levels of atoms (the inner electrons are too tightly held • These available electrons are called the valence electrons • For the main group elements these are in the s & p shells

19. Electronegativity • Valence electrons hold atoms together • In many compounds, the negative charge of the valence electrons is concentrated closer to one atom than to another • Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons • Fluorine is assigned a number of 4.0

20. Electronegativities tend to increase across each period • Electronegativities tend to either decrease down a group or remain about the same • Noble gases do not form many compounds and may not have values

24. Properties of the d and f block elements • The properties of the d block elements vary less and with less regularity than those of the main group elements • Both the outer s and the d electrons are available to interact with their surroundings • The atomic radii of the d block elements generally decrease across a period • The d electrons shield the outer electrons • The electrons repel each other

25. The f block elements behave in a similar way • Ionization energies generally increase across a period for d & f block elements • In contrast, they generally increase down a group because the electrons available for ionization in the outer s level are less shielded (incomplete d shell) from the increasing nuclear charge

26. Ion formation in the d & f block elements follows the reverse order of electron configuration • For d block, although electrons are being added to the d, they are removed from the outer s first (most d block elements therefore form +2 ions) • The d & f block elements all have similar electronegativities • Follow general trend

28. Chemical Bonding

29. Introduction to Chemical Bonding • A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds atoms together • Atoms bond because it decreases their potential energy, creating more stable arrangements of matter

30. Chemical bonding that results from the electrical attraction between large numbers of cations and anions is called ionic bonding • Covalent bonding results from the sharing of electrons pairs between two atoms • In a purely covalent bond, the shared electrons are “owned” equally by the two bonded atoms

31. Bonding is rarely purely ionic or covalent • Electronegativity is a measure of an atom’s ability to attract electrons • The degree of ionic or covalent character is determined by calculating the difference in electronegativity

32. The d indicates a partial charge

33. Covalent Bonding and Molecular Compounds • A molecule is a neutral group of atoms that are held together by covalent bonds • Individual unit capable of existing on its own • May consist of two or more atoms • A chemical compound whose simplest units are molecules is called a molecular compound

34. A chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts • A diatomic molecule is a molecule containing only two atoms

35. A balance is reached between the attractive forces and the repulsive forces between the nuclei and electrons. This results in the most energetically stable arrangement.

36. In a covalent bond, the electrons orbitals can be pictured as overlapping (the electrons are free to move in either orbital) • The distance between two bonded atoms at their minimum potential energy is the bond length • The atoms will vibrate a bit

37. The difference between the potential energy zero level (separate atoms) and the bottom of the valley (bonded atoms) is the bond energy that is released when the bond is formed • It is also the energy required to break a chemical bond and form neutral isolated atoms • Atoms tend to acquire noble gas configurations when bonding

39. Octet rule: Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level

40. There are exceptions to the octet rule • Boron: In BF3 , boron will share its three valence electrons and acquire a total of 6 • When some elements combine with the very electronegative atoms of F, O, and Cl, an expanded valence that involves electrons in the d orbitals occurs

41. Electron-dot notation is an electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol

42. Electron dot notations can also be used to represent molecules • A shared pair of electrons is drawn between two atoms, an unshared pair is a pair of valence electrons that belongs exclusively to one atom and is not involved in bonding H:H

43. A shared pair of electrons is often represented with a dash • The are called Lewis structures • A structural formula indicates the kind, number, arrangement, and bonds, but not the unshared pairs of atoms in a molecule

44. A single bond is a covalent bond produced by the sharing of one pair of electrons between two atoms • A double covalent bond is produced by the sharing of two pairs of electrons between two atoms • A triple covalent bond is a bond produced by the sharing of three pairs of electrons between two atoms • Double and triple bonds are referred to as multiple bonds

46. C, N, and O can have multiple bonds • H can have only one bond

47. Resonance structures cannot be correctly represented by a single Lewis structure • Ozone • Once thought to split time between two structures • Experiments show that bonds are equivalent ( average of two bonds)

48. Not all covalent compounds are molecular • Some are continuous 3 dimensional networks of covalently bonded atoms • Called covalent-network bonding

49. Ionic Bonding and Ionic Compounds • An ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal • Most are crystalline solids • The formula simply represents the simplest ratio of ions that give neutrality of charge – called a formula unit