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

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 NaOH Na+ + 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]

7. 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

8. (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 HCl(aq) + MnS(s)  H2S(g) + MnCl2(aq)
oxidation numbers (+1)(-1) (+2)(-2) (+1)(-2) (+2)(-1)

9. 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)

10. 1 2 3 4 5 6 7 8

11. 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)

12. 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)

13. 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

15. 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

16. 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 %
21Ne %
22Ne %
atomic mass of
Ne = * * *0.0925
atomic mass Ne = amu (20. in sig figs)

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

18. 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

19. 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

20. 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)

21. Electron Configuration
Oxygen Orbital Notation Simplified Notation
1s s p
O s22s22p [He]2s22p4
Unpaired
Ground State
Spin Paired
Excited State [He]2s22p33s1
Anion - O [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

22. 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] 3d54s Cu [Ar] 3d104s1
pg. 224

23. 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

24. 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

25. 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