Unit 5: Trends and Bonding

1. Unit 5: Trends and Bonding

2. Valance electrons: available electrons to be lost, gained, or shared in the formation of compounds Ionic Bond: losing or gaining electrons to form a bond between ions, creating a neutral compound Ion: an atom or group of bonded atoms that has a positive or negative charge. Cation: When metals lose electrons to form positive ions Anion: When non-metals gain electrons to form negative ions. Covalent Bond: sharing of electrons to form neutral compounds

3. Charges on Periodic Table +1 0 +2 +3 -3 -2 -1 +/-4

4. The Periodic Law Mendeleev: created the first periodic table to relate the properties of elements and arranged them according to atomic mass • Problems: (1) Most elements could be arranged in order of increasing atomic mass but a few could not?(2) What was the reason for chemical periodicity? Mosely:arranged elements based on atomic number Periodic Law: the physical and chemical properties of the elements are periodic functions of their atomic number Periodic Table: an arrangement of elements in order of atomic number so that elements with similar properties fall into the same group or column.

5. s-Block Elements: Groups 1 &2 • group 1 elements are more reactive than group 2 because it is easier to remove 1 electron rather than 2 • Alkali metals: elements of group 1, highly reactive • Alkaline-Earth metals: elements of group 2, very reactive d-Block Elements: Group 3-12. • Transition elements that are typically less reactive than s-block elements

6. p-block elements: Groups 13-18 • properties vary greatly, metals, non-metals, metalloids • Halogens: group 17 elements, most reactive non-metals • Ability to react is based on them having 7 electrons in outer shell, (they want 8 to be stable) • Noble Gases: group 18 elements, stable and unreactive. • They already have 8 valence e- f-block elements: Lanthanides and Actinides Lanthanides: shiny metals similar in reactivity to group 2 Actinides: all but first four are man made in laboratories

7. Trends in Periodic Table (1) Atomic radius - one-half the distance between the nuclei of identical atoms that are bonded together. r = d/2 • Period: AR decreases as you go across a period. (leaving noble gases out of it) • More e- in an energy level, the smaller it is, because of the increased attraction of the negative to the positive nucleus. • Group: AR increases as you go down a group. - adding extra layers of electrons- more energy levels

8. Atomic Radius M S L

9. (2) Ionic Radius: Ion: an atom that has a positive or negative charge Cation: a positive ion, from loss of electrons • decreases atomic radius b/c less electrons (energy level) Anion: negative ion, from addition of electrons • increase atomic radius, increases electron repulsion • Metals tend to form cations • Non-metals tend to form anions • Period: decreases across a period (like atomic radius) • Group: increases down a group (like atomic radius)

10. Ionic Radius M S L

11. (3)Ionization Energy - amount of energy required to remove the outermost electron from a neutral atom -to make positive ions The energy you need to put into an atom to take away an electron. (give and take) They give you an e- because you gave them energy • Ionization: the formation of an ion • Period: IE increases as you go across a period • When elements are closer to having a full octet, they do not want to give up an electron, so IE is much higher • Group: IE decreases as you go down a group. Or increases as you go up a group • it is easier to remove the electrons from outer energy levels because they are farther from the positive nucleus and thus less pull by the nucleus

12. Ionization Energy L M S

13. (4) Electron Affinity: the energy change that occurs when an electron is acquired by a neutral atom-to make negative ions • The energy given/ released when you add an electron to an atom (give and take) • They give you energy because you give them an e- • “love of electrons” • Most atoms release energy when they acquire an electron, A + e- A- + energy • Period: increases as you go across, but it is negative energy because the atom is releasing it. The more they want electrons the more they will “pay” for it in energy, thus giving more energy

14. Group: decrease as you go down because electrons add with greater difficulty going down a group, because of increased atomic radius. They don’t want e- so they won’t “pay” much for it. • Non-metals: Some atoms want electrons, so it is easy to give them one and they are grateful and give you energy in return. So they release energy (negative E) • Metals: Some atoms do not want electrons, so you have to force the electron onto the atom by also using/giving energy. So they absorb energy (positive E)

15. Electron Affinity L M S

16. (5) Electronegativity: an atom's ability to grab another atom's electrons (ability to attract electrons) • Occurs in covalent bonding, when electrons are shared • Period - increases as you go across a period • Because atoms want electrons more as you go across • Group - decreases as you go down a group or remain the same. • Since there are more energy levels the positively charged nucleus cannot attract the electrons well enough

17. Electronegativity 4.0 0.7

18. Electronegativity

19. Electronegativity

20. Reactivity (6) Reactivity - refers to how likely an atom is to react with other substancesMetals: the more likely to lose an electron…the more reactiveNon-metals: the more likely to gain an electron….the more reactive. L Non-metals Metals L

21. Melting Point and Boiling Point • Metals: Decreases as you go down a groupand increases as you go across a period • Non-metals: Increases as you go down a group.Generally decreases as you go across a period M L S S L

22. BONDING: Electron Dot Notation: 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. Remember- orbitals fill with one electron before any one fills with two electron (Hund’s Rule) Octet Rule: compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has a full octet in its highest occupied energy level When electrons are shared, orbitals are overlapped. • To form a bond: release energy • To break a bond: absorb energy, it requires energy

23. Electron-dot can be used to represent molecules: Unshared Pair (lone pair): pair of electrons that is not involved in bonding and that belongs exclusively to one atom. (not bonded) Lewis Structures: ( F F ) • dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds • dots adjacent to only one atomic symbol represent unshared electrons

24. Two Types of Compounds • There are two ways to achieve a full octet to become stable- Ionic bonding or Covalent Bonding • Ionic compounds: result from ionic bonding • Metals and Non-metals combine • composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal • group 1or 2 want to give electrons away and group 16 or 17 want to take electrons, so valence electrons are transferred between atoms upon collision. • All metals can combine with a non-metal to form Ionic bonds through electron transfer.

25. Show the transfer with charges! • Positive ions- lose electrons • Negative ions- gain electrons • Forces that hold ions together is very strong • Ca (+2) and F (-1) give you CaF2 (1-to-2 ratio) Electron-Dot: Ca F2: Na2O:

26. What are Ionic compounds like?: • salt, like ocean water. • Strong and brittle like egg shells, once you crack the shell it all falls to pieces. • 3-D, tightly bound structure, crystal structure • most ionic compounds exist as crystalline solids • Number line: -3, -2, -1, 0, 1, 2, 3, 4 N O F Ne Li Be B C • Ionic bonds are stronger than covalent bonds.

27. 2)Covalent Compounds: when electrons are shared • Bonding between Non-metals and Non-metals • when neither atom wants to give up electrons fully • Hydrogen always forms covalent bonds because a Hydrogen can’t lose its only electron by transfer so it must share • H2O: • covalent bond: atoms share electrons (no ions) • molecule: neutral group of atoms held together by covalent bonds

28. What are Covalent Compounds like?: • In a purely covalent bond, the shared electrons are “owned” equally. • Bonds within covalent compounds are strong, but attraction between molecules are weak, so they will break apart into molecules easily. • Small difference in electronegativity, so they must be close together on the periodic chart. • Electrons are shared, so don’t write charges! • CO2 • [NH4]+1 • More than two electrons can be shared CO2

29. Steps for Covalent Structures • Cross structure (carbon in middle-symmetrical) • Give them what they want (8 e-, except H) • Count what they want • Determine how many valence electrons they will bring to the compound (allowed) • Remove the difference in pairs of e-, then move a pair from the atom next to it, to be shared between them • Check- count total, and make sure that none have more than 8

30. Ionic: MgCl2 • Covalent: CH2S Want: 14e- C: 4e- H:1e- x 2= 2e- S: 6e- Allowed: 12e-

31. Single Bond: covalent bond produced by the sharing of one pair of electrons between two atoms. CH3I Double Bond: covalent bond produced by sharing of two pairs of electrons between two atoms. Triple Bond: covalent bond produced by sharing of three pairs of electrons between two atoms.

32. Multiple Bonds: double or triple bonds. • Bond Strength: Triple > Double > Single Bond Energies: Triple > Double > Single • Bond Length: Single > Double > Triple • So Triple bonds are strongest, so they require more energy to break, but they are the shortest. Resonance Structures: • Some molecules cannot be represented by a single Lewis Structure. One such molecule is Ozone (O3) • This is called Resonance and the different Lewis structures are called resonance structures. To indicate resonance a double-headed arrow is placed in between

33. Polar Covalent: have partial charges d+ represents a partial positive charge d- represents a partial negative charge HCl: Hydrogen and Chlorine. • Chlorine is more electronegative so when sharing electrons is attracts electrons more, so it is partially negative • Hydrogen wants electrons less so partial positive Non Polar Covalent : two elements with similar desire for electrons so has NO partial charges, because of equal sharing, usually gases

34. Overall….. • Like dissolves Like: so charged ions and molecules will dissolve each other and not those without charges. • Ionic- give and take- transfer of e-, thus full charges • Polar Covalent- must share, but desire for e- is slightly different, so unequal sharing and partial charges • Non-polar Covalent- same desire for e-, so equal sharing and no charges

35. Ionic vs. Covalent Elements combine to form either ions or molecules. Properties of Ionic Compounds: - Physical: strong, hard, brittle, well organized, tightly bound, 3-D crystal structures- takes a lot of energy to break bonds- tend to dissolve in water (like salt)- electrolytes: any compound that conducts electricity- large difference in electronegativity- stronger compounds- higher melting points- Ex: salt, egg shells

36. Ionic vs. Covalent Properties of Covalent Compounds:- Molecules held together by strong covalent bonds- strong bonds within molecules, but a weak bond that holds one molecule to another.- Polar : liquids or soft solids that don’t conduct electricity (sugar) - small difference in electronegativity - weak compounds - low melting points -Non-Polar: gases that don’t conduct electricity - almost no difference in electronegativity - weak compounds - lowest melting points-Ex: Candles, plastics, crayons, diamonds

37. Molecular Geometry • Properties of molecules depend on bonding & geometry Molecular polarity: uneven distribution of molecular charges VSEPR Theory: (used to predict geometry) states that repulsion between valance electrons surrounding an atom causes them to be oriented as far apart as possible. Bond Angles : AB2(2 or 3 atoms) AB3 (4 atoms) AB4 (5 atoms)

38. Linear: - AB or AB2 - Example: HCl or CO2 - 0 lone pairs of e- - 2 atoms bonded to the central atom Bent or Angular: - AB2(E) - Example: SnCl2 - 1 lone pair of e- - 2 atoms bonded to the central atom

39. Trigonal Planar: - AB3 - Example: BF3 - 0 lone pairs of e- - 3 atoms bonded to the central atom Tetrahedral: - AB4 - Example: CH4 - 0 lone pair of e- - 4 atoms bonded to the central atom

40. Trigonal Pyramidal: - AB3(E) - Example: NH3 - 1 lone pairs of e- - 3 atoms bonded to the central atom Bent or Angular: - AB2(E2) - Example: H2O - 2 lone pair of e- - 2 atoms bonded to the central atom

41. Trigonal Bipyramidal: - AB5 - Example: PCl5 - 0 lone pairs of e- - 5 atoms bonded to the central atom Octehedral: - AB6 - Example: SF6 - 0 lone pair of e- - 6 atoms bonded to the central atom

42. Bond Energies and Heats of Reaction (DH) Bond Energy is the energy required to break a chemical bond.  Tabulated values (Table 6-1) are average bond energies in units of kJ / mole. Bond-breaking is endothermic, bond-making is exothermic. DH for a reaction can be estimated from bond energies as follows. (Counting ALL bond energies as positive values!) H BE (bonds broken) -  BE (bonds formed) Problem: Use data in Table 6-1 to estimate DH° for the reaction. CH2=CH2 + H2O CH3-CH2-OH Bonds BrokenBonds Formed C=C 612 C-C 348 H-O 463 C-H 412  = 1,075 C-O 360  = 1,120 DH° 1,075 - 1,120  - 45 kJ/mole