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increases

Metallic Character. increases. increases. increases. Electrolytes. Strong Electrolyte -Substances that conduct electricity well in dilute aqueous solution (completely ionize in water). KMnO 4(s)  K + (aq) + MnO 4 - (aq).

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  1. Metallic Character increases increases increases

  2. Electrolytes Strong Electrolyte-Substances that conduct electricity well in dilute aqueous solution (completely ionize in water) KMnO4(s) K+(aq) + MnO4-(aq) Weak Electrolyte-Substances that conduct electricity poorly in dilute aqueous solution (very few molecules ionize in water) CH3COOH(aq) CH3COO-(aq) + H+(aq) nonelectrolyte-Substances that either do not or poorly conduct electricity in aqueous solution (none of the molecules ionize in water) CH3CH2OH(aq) CH3CH2O-(aq) + H+(aq)

  3. Strong Acids (SA)-Acids that are strong electrolytes: There are 7 common SA: HCl, HBr, HI, HNO3, HClO4, HClO3 & H2SO4 Memorize Table 4-5 (SA and A-) HCl(aq) H+(aq)+ Cl-(aq) Weak Acids-Acids that are weak electrolytes. All acids that aren’t Strong Acids are Weak Acids (WA) Note: HF is a weak acid. HF(aq) H+(aq)+ F -(aq) & H2S Memorize Table 4-6 (common WA and A-)

  4. Strong Bases (SB)-Bases that are strong electrolytes: Group IA hydroxides & oxides. Heavier Group IIA hydroxides & oxides (Ca(OH)2, CaO, Sr(OH)2, SrO, Ba(OH)2 & BaO) Na2O + H2O 2NaOH 2Na+ + 2OH- BaO + H2O Ba(OH)2(s) Ba2+ + 2OH- Weak Bases-Bases that are weak electrolytes. All soluble bases that aren’t Strong Bases are Weak Bases (WB) NH3 + H2O NH4+ + OH- Memorize Table 4-7 Also, CH3NH2, (CH3)2NH & (CH3)3N CH3NH2 + H2O CH3NH3 + + OH- Insoluble Bases-Bases that are sparingly soluble to insoluble. All hydroxides except those of Group IA and heavier Group IIA. For example: Cu(OH)2, Zn(OH)2, Fe(OH)2, Fe(OH)3, ....

  5. Solubility Rules #2 Rule #1 Rule [Exceptions]

  6. Acid-Base Reactions (Neutralization Rxns) Acid + Base Salt + (Water) + heat HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) Formula Unit Equation H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) Na+ (aq) + Cl - (aq)+ H2O (l) Total Ionic Equation (Spectator Ions) H+ (aq) + OH- (aq) H2O (l) SA+SB Net Ionic Equation

  7. Metathesis (Double Displacement) AX + BYAY + BX No Change in Oxidation Number • 1) Acid-Base neutralization HCl + NaOH  H2O + NaCl • 2) Precipitation reaction AgNO3 + NaCl AgCl(s) + NaNO3 • 3) Gas Formation 2HCl(aq) + MnS(s) H2S(g)+ MnCl2(aq) • oxidation numbers (+1)(-1) (+2)(-2) (+1)(-2) (+2)(-1)

  8. reducing agent-compound that is oxidized (Na(s)) oxidizing agent-compound that is reduced (H2O) oxidation-loss of electrons reduction-gain of electrons • The oxidation state of any atom in a free uncombined element is zero (Cl2, H2, O2, P4, Na(s) ... ) • The oxidation state of any monatomic ion is equal to the charge (Na+ ox. state = +1 ; Cl- ox. state = -1 ; Fe2+ ox. state = +2) • The sum of oxidation numbers of all atoms in a compound is zero • The sum of oxidation numbers of all atoms in an ion is equal to the charge of the ion (SO42- sum of ox. numbers = -2)

  9. 1 2 3 4 5 6 7 8

  10. Combination Rxn: 2 or more substances combine to form a compound • Element + Element  Compound • Metal + Nonmetal  Binary Ionic Compound • 2Na(s) + Cl2(g)NaCl(s) (also Redox) • Nonmetal + Nonmetal  Binary Covalent Compound • P4(s) + 10Cl2(g) PCl5(s) (also Redox) • Compound + Element  Compound • PCl3(l) + Cl2(g)PCl5(s) (also Redox) • Compound + Compound  Compound • CaO(s) + H2O(l) Ca(OH)2(aq) (NOT a Redox)

  11. Decomposition Rxn: a compound decomposes to form products • Compound  Element + Element • 2HgO(s) Hg(l) + O2(g) • Compound Compound + Element • 2H2O2(l) 2H2O(l) + O2(g) • (Disproportionation rxn-H2O2 is the reducing agent and the oxidizing agent) • Compound  Compound + Compound • CaCO3(s) CaO(s) + CO2(g)

  12. Displacement Rxn: one element displaces another from a compound Table 4-12 (pg. 148) Activity series of SOME Elements The more active element displaces the less active element • more active metal + salt of less active metal less active metal +salt of more active metal • Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s) • active metal + nonoxidizing acid  salt of acid + H2(g) • Zn(s) + H2SO4(aq) ZnSO4 + H2(g) • more active nonmetal + salt of less active nonmetal less active nonmetal +salt of more active nonmetal • Br2(l) + 2NaI(aq) I2(s) + 2NaBr(aq) • F2 > Cl2 > Br2 > I2

  13. Isotopes - Atoms of the same element (have the same # of protons) that have different masses (have differing number of neutrons) • Mass number A = Z (#p) + #n Nuclide Symbol E A Z 11 5 10 5 #n=6 #n=5 B B

  14. Mass Spectrometry - measures the charge-to-mass (e/m) ratio of charged particles. If we put Neon in a mass spectrometer we could determine how many natural isotopes there are for neon and the percent abundance of each isotope 20Ne 90.48% 21Ne 0.27% 22Ne 9.25% atomic mass of Ne = 19.99244*0.9048 + 20.99384*0.0027 + 20.99384*0.0925 atomic mass Ne = 20.087 amu (20. in sig figs)

  15. Electromagnetic Radiation hc  E = c/ Intensity of light - # of photons striking a given area of the plate per second

  16. Quantum Mechanics • Atoms and molecules can exist only in certain energy states (quantized energy levels) • When they change their state they absorb or emit radiation (light/photons) E = hc/ • Allowed energy states are described by 4 quantum numbers

  17. Principle Quantum Number - n = 1, 2, 3 .... • Describes the main energy level and the extent of the orbital • Angular Momentum Quantum Number - l = 0, 1, 2, ... , (n - 1) • Describes the shape of the orbital • l = 0 s orbital, l = 1 p orbital, l = 2 d orbital, l = 3 f orbital, l = 4 g orbital, l = 5 h orbital ..... • Magnetic Quantum Number - ml = -l, -l+1 , ... , 0 , ... , l-1 , l • Describes the orientation of each orbital • Spin Quantum Number - ms - ±(1/2) • Describes the spin of the electron in each orbital

  18. n 1 2 3 4 5 6 7 d (l=2) 3 4 5 6 p (l=1) 4 5 s (l=0) f (l=3) Saunders 7.13

  19. Electron Configuration Oxygen Orbital Notation Simplified Notation 1s2s2p O 1s22s22p4 [He]2s22p4 Unpaired Ground State Spin Paired Excited State [He]2s22p33s1 Anion - O2- [Ne] Aufbau Principle-add electrons to give the lowest energy Hund’s Rule - electrons fill all orbitals of a subshell before pairing and they will have parallel spin Pauli Exclusion Principle - no two electrons can have the same four quantum numbers pg. 216

  20. 3d 4s saunders 8.3sb for magnetism Mn [Ar] 3d54s2 K [Ar] 4s1 K+ [Ar] paramagnetic-weakly attracted into a magnetic field due to unpaired electrons 3d 4s Zn [Ar] 3d104s2 diamagnetic-very weakly repelled by a magnetic field due to all electrons being spin coupled Exceptions Cr [Ar] 3d54s1 Cu [Ar] 3d104s1 pg. 224

  21. Effective Nuclear Charge - Nuclear charge experienced by the outer shell electrons Ionic Radii Atomic Radii Neutral Atom  Anion size Increase Increase Neutral Atom  Cation  size Increase Isoelectronic - species that have the same number of electrons

  22. First Ionization Energy - The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form an ion with a 1+ charge Increase Increase Increase

  23. Electron Affinity - the amount of energy absorbed when an electron is added to an isolated gaseous atom to form an ion with a 1- charge. Increase Increases = less negative or more positive Increase Increase Electronegativity - measure of the relative tendency of an atom to attract electrons to itself when it is chemically combined with another atom. (qualitative) Increase Increase Increase

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