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Principles of Chemistry 1 Test 4 Review. Always behave like a duck - keep calm and unruffled on the surface, but paddle like the devil underneath. -Jacob Bravde NOW with corrections !!!. By: Bo Marshall. The Fine Print.

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Principles of Chemistry 1 Test 4 Review

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    1. Principles of Chemistry 1Test 4 Review Always behave like a duck - keep calm and unruffled on the surface, but paddle like the devil underneath. -Jacob Bravde NOW with corrections !!! By: Bo Marshall

    2. The Fine Print This is NOT intended to be your only source of information for the test - this is simply an outline for me to follow to help you review…As usual, I recommend you read your text, attend and pay attention in class, re-do the homework, do suggested homework problems, and take the practice test.

    3. Monatomic Anions/Cations Anions - Add e- to lowest unfilled sub shell Ex, for O2-, Normally, O = [He] 2s22p4 Add 2 more e- for a 2- charge, they will fill the 2p orbital [He] 2s22p6, same as Ne Not the same as Ne though, the NUCLEUS is DIFFERENT!!! So is the anion smaller or larger?

    4. Monatomic Anions/Cations Anion will be larger - more e- for the same nucleus to attract Cations - The last electron in is NOT necessarily the 1st one out… Remove e- from subshell w/ highest n and l values 1st - For Ex… Ca2+, remove from 4s shell Fe2+, remove from 4s shell 1st, leave d shell alone Fe3+, remove from 4s shell 1st, then one from the d shell Will the cation be larger or smaller than the neutral atom?

    5. Para/diamagnetic Paramagnetic - Has unpaired electrons, attracted to magnetic field Diamagnetic - Electrons Paired, repulsion from magnetic field Even if an atom has paramagnetism, there is some diamagnetism b/c some electrons will be paired Ex with Ti2+, paramagnetic, shows that electrons removed from 4s orbital, not 3d to make cation

    6. Periodicity What is the only element that is ONLY paramagnetic? Why are there exceptions to the mass increasing across a period? B/C elements are arranged in groups that behave alike. This gives trends across the periodic table…

    7. Periodic Trends

    8. KNOW THESE TRENDS!!! Atomic Size&Metallic Behavior Increasing Increasing Increasing EA , IE , EN Increasing

    9. Trends Atomic Radius: distance between nuclei of bonded atoms Increases DOWN a group (higher n #) Increases right to left across a period (More nuclear charge w/ same n #)

    10. Trends

    11. Trends Ionization Energy: energy to remove 1 mol of e- from 1 mol of gaseous atoms/ions IE1 = Energy to go from neutral atom to +1 cation IE2 = Energy to go from +1 cation to +2 cation IE will increase as you take more electrons off => Huge jump seen when you get to core electrons Trend opposite of atomic size - easier to remove the electron if it is further from the nucleus

    12. Trends Why is the 1st IE of N larger than that of O? N has 3 unpaired electrons, O has 2…what did that Hund guy say??? 2p orbitals in N 2p orbitals in O

    13. Trends Electron Affinity: Energy for a gaseous atom to gain an electron Negative Value (-) energy given off when electron is gained Positive Value (+) energy needed to form Why would N have a (+) EA?

    14. Trends Metallic Behavior Increases left to right and top to bottom, like atomic size Metals - shiny, solid, etc. Non-metals Metalloids - staircase Metal oxide + Water = metal hydrxide (base) Therefore, metal oxide and an acid gives water and a base (like an acid base Rxn) Nonmetal oxide + water = acid Therefore, nonmetal oxide and base gives salt and water

    15. CO2(g) + H2O -> ? Nonmetal oxide Gives H2CO3 (aq) => An acid! (FYI - this rxn helps keep you alive) CO2 (g) + NaOH (aq) -> ? Nonmetal oxdixe Gives Na2CO3 + H2O Na2O(s) + H2O -> ? Metal oxide, gives a base 2 NaOH (aq) NiO (s) + 2HNO3(aq) -> salt + water Ni(NO3)2 (aq) + H2O

    16. Trends Ionic size Cations:smaller than parent atom => Less electrons for the same nucleus to attract Higher + charge = smaller size Anions: Larger than parent atom => More electrons for the same nucleus to attract  Higher (-) charge = larger size If they have the same charge, the ion size will increase down a group…

    17. So which is larger, P3- or Cl-? Same # of e-, but Cl- has more protons in the nucleus  Cl- attracts the electrons better and is smaller… Isoelectronic: Having the same number of electrons and electron configuration Ex. N3-, O2- and F- all are 1s22s22p6 Are Ca and Ti2+ isoelectronic? Ca: [Ar]4s2 , Ti2+: [Ar] 3d2

    18. Lewis Dot Structures For s and p blocks, group # = electron # and  # of dots Octet Rule (lots of exceptions) Doesn’t apply to TM’s H, He, Li, Be are all OK w/ only 2 e- in outer shell

    19. Lattice Energy/Ionic Bonding If EA is endothermic (+) and IE is exothermic (-), but smaller magnitude than the EA, why do ionic compounds form? The Lattice Energy => energy of the (+) and (-) charges coming together to form ionic solid crystal => Very exothermic, compensates for the energy input needed to form the ions

    20. Lattice Energy Bigger magnitude of LE gives Larger MP Harder solid Decreased solubility Can figure out the LE through the Born-Haber cycle process (Hess’s law application)

    21. Lattice Energy Trends in LE Larger Cation = smaller LE Larger Anion = smaller LE Colums Law The charge has a larger influence than the radius size If charge the same, use radius size

    22. Covalent Bonding non- metal w/ non-metal The nucleus(protons) on one atom attract the electrons in the other and vice versa Distance between them is lowest energy point To separate them will require ENERGY => It will be ENDOTHERMIC to break the bonds!!! (Dr. Donovan said this like 600 times in class!)

    23. Covalent Bonding

    24. Bond Order Single Bond, B.O. = 1 Double and triple bonds Higher Bond order = more bonds = shorter bond and higher bond energy So a double bond takes more energy to break than a single bond…

    25. Covalent Properties Bonds within the molecules (I.e., between the atoms) are strong Bonds between the molecules weak So what does water look like when it is evaporated? Network Covalent solids Don’t exist as individual molecules A bunch of the same atoms covalently bonded together

    26. Electronegativity “Big Important Topic” Quote from Dr. Donovan in class Bonding not evenly shared Need to know that F is most EN (4.0) and Cs least EN (0.7) More electronegative = more polar bond = electrons attracted to the EN atom Super-important for Lewis structure drawing (and for Principles II, and OCHEM, and to sound smart at cocktail parties)

    27. EN and Ox Number Used where the rules you memorized don’t work Done similar to F.C. More EN atom gets all bonded e- Lone pairs to the owner ON = (# valence e- in neutral atom) - (total e- assigned (from rules above)) Ex in class with cyanide ion

    28. Bond Polarity Polar bond occurs whenever 2 atoms with different EN are bonded Shown as partial (+) and (-) charges, or with the arrow We’ll practice this with Lewis Structures Must Understand => There is NO SUCH thing as a perfectly ionic bond, even between Cs and F!!! Still some covalent character to them

    29. Metallic Bonding Valence e- surrounding the nuclei and core electrons of a bunch of metal atoms => all the atoms share the e- Repeating pattern of atoms => Sea of e- allows the atoms to “smush” w/o breaking the repeating pattern

    30. Summary of Bonding Ionic: metal w/ non-metal Covalent: non-metal w/ non-metal Metallic: metal w/ metal No bond is completely ionic, still retains some covalent character

    31. KNOW THESE TRENDS!!! Atomic Size&Metallic Behavior Increasing Increasing Increasing EA , IE , EN Increasing

    32. Lewis Structures To the practice problems… Lets draw some Lewis structures Calculate the FC on each atom Determine the bond polarity Look at some Octet rule exceptions