Metallic Character increases increases increases

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)

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) Memorize Table 4-6 (common WA and A-) & H2S

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, ....

Solubility Rules #2 Rule #1 Rule [Exceptions]

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

(Double Displacement) Metathesis (Double Displacement) AX + BY AY + 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)

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)

1 2 3 4 5 6 7 8

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)

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)

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

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 10 5 11 5 B #n=5 B #n=6

atomic mass Ne = 20.087 amu (20. in sig figs) 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)

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

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

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

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

Electron Configuration Oxygen Orbital Notation Simplified Notation 1s 2s 2p 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

paramagnetic-weakly attracted into a magnetic field due to unpaired 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

Atomic Radii Ionic Radii 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

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

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