1. Last day before the SOL’s cram day
CHEMISTRY
Last day before the SOL’s cram day

2. Most important topics on the SOL
Naming (ionic, covalent, transition metal)
Bonding (ionic, covalent)
Measurement (estimate last digit)
Significant figures (+ and -, x and ÷)
Subatomic particles (p+, n0, e-, mass, atomic #)
Chemical and physical properties and changes
MOLES
Reactions and Equations
Stoichiometry
Molarity and dilutions
Gas Laws
Thermochemistry

3. GROUP or FAMILY
Similar physical
and chemical
properties (same # of
valence electrons
PERIOD

5. Non-metals above the staircase
Metals below the staircase

6. Naming (ionic, covalent, transition metal)
Oxidation number: the charge of a monatomic ion
Indicates the number of electrons an atom has lost(+) or gained(-)
Metals tend to lose electrons (+ oxidation numbers)
Cations with more than one oxidation number:
Indicate the oxidation number with a Roman numeral
Example: Fe 2+ = iron(II) ion
Fe 3+ = iron(III) ion
Ionic: Metal + nonmetal ex: NaCl (sodium chloride)
Covalent: CO2 (carbon dioxide)  use prefixes

7. KMnO4 +1 4 (-2) = -8 What must Mn be? The overall charge is 0
+1 + X + (-8) = 0
X = + 7

8. The first element does not use mono if there’s only one. Examples:
Molecular compounds are composed of two non-metals (above the staircase)
Indicate # of each atom using prefixes (mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca)
The first element does not use mono if there’s only one.
Examples:
OF2 is named oxygen diflouride
N2O is named dinitrogen monoxide

9. Ca 2+ + Cl-  CaCl2 Fe 3+ + O2-  Fe2O3 Fe 2+ + O2-  FeO
Na+ + SO  Na2SO4
Ca 2+ + PO43-  Ca3(PO4)2
calcium chloride
iron(III) oxide
iron(II) oxide
sodium sulfate
calcium phosphate

10. Bonding (ionic, covalent)

13. COVALENT BOND POLARITY
non-polar covalent bonds: bonding electrons shared equally between two atoms
Example: H2
polar covalent bonds (polar bonds): bonding electrons shared unequally.
Example: HCl δ+ δ-
BOND POLARITY BASED ON ATOMS’ ELECTRONEGATIVITY
the more electronegative atom acquires a slight negative charge (δ-).
the less electronegative atom acquires a slight positive charge(δ+ ).

15. Measurement (estimate last digit)

16. Significant figures (+ and -, x and ÷)
Adding/subtracting  lowest number of decimal places
Multiplying/dividing  lowest number of significant figures

17. Subatomic particles (p+, n0, e-, mass, atomic #)
Isotopes – atoms with the same number of protons but different numbers of neutrons
Ions- atoms that have lost or gained electrons making them have a positive or negative charge
ANION: negative ion (gained an e-)
CATION: positive ion (lost an e-)
Atomic number (# of protons) for an atom never changes… that’s how we identify an atom

18. Chemical and physical properties
Physical Properties – descriptions which do not involve a change in composition (color, density, texture, etc.)
Chemical Properties – Descriptions of a change in composition of a substance (reactivity, combustibility, etc.)

19. Chemical and physical change
does not alter the composition of the substance
Includes changes in state of matter, or shape
Chemical Change
Alters the chemical structure of the substance
Cannot get original substance back through physical means
Evolution of a gas
Production of heat / light
Formation of a precipitate
Unexpected color change

20. MOLES Know Your Triangles! 1 mole= 6.02 x 1023 particles
1 mole = molar mass in grams
1 mole = 22.4 L of gas at STP
When in doubt about what to do convert what you are given into moles and see if you can convert moles into what you need.

21. Reactions Synthesis A + B AB Decomposition AB  A+B
Single Replacement A + BX B + AX
Combustion CxHy + O2 CO2+H2O
Double Replacement
AB + XY  AY + BX
HX + MOH MX + H2O
(neutralization)

22. Equations Balance in order to obey Law of Conservation of Mass
Same number of atoms of each type on both sides
Reactants  Products

23. Stoichiometry Change given quantity to moles
Use mole ratio from Equation
Change moles to answer quantity OR
Change the coefficients (if needed) by multiplying by molar mass, molar volume, or avogadro’s number
Set up a proportion and solve.

24. Stoichiometry Cont. Limiting Reactant- The one that runs out first
Calculate twice. The Limiting Reactant
gives you the least amount of product
Percent Yield:
[(Actual Yield) /( Theoretical Yield)] x 100%

25. Molarity and dilutions
Molarity = moles
Liters
Change masses to moles and mL to L
Dilutions: M1V1 = M2V2
M = Molarity V = Volume
1 = concentrated = diluted
solution solution

26. Always use degrees Kelvin
Gas Laws
General Properties of Gases There is a lot of “free” space in a gas. Gases can be expanded infinitely. Gases fill containers uniformly and completely. Gases diffuse and mix rapidly.
Always use degrees Kelvin
°C = K
P1V1 =
P2V2 T1 T2
The Combined Gas Law
Boyle’s Law
Inverse relationship
P1V1 =
P2V2 V1 = V2 T1 T2
Charles Law
Lussac’s Law P1 = P2 T1 T2

27. T= temperature in Kelvin R = universal gas law constant = 8.314
The Ideal Gas Law
Memorize: PV = nRT
P= pressure in kPa
V= liters
n= moles
T= temperature in Kelvin
R = universal gas law constant = 8.314
The SOL test uses
kPa x L
Moles x K
Dalton’s Law of Partial Pressures
Memorize: Ptotal = P1 + P2 + P3 + …

28. Gas pressure is measured in atmospheres, kilopascals (kPa), or mm Hg
One atmosphere = kPa = 760 mm Hg
Assumptions relating to gases:
Gas particles have negligible volume compared to container size
Gas particles do not attract or repel each other*
Gas particle move constantly, rapidly and randomly
All collisions are perfectly elastic (particles collide like billiard balls, not marshmallows)
However, gas particles really do attract each other due to intermolecular forces

29. Catalysts lower the Activation energy, making reactions faster.

30. N2(g) + 3H2(g)  2 NH3(g) + heat
Reversible Reactions
N2(g) + 3H2(g)  2 NH3(g) + heat
The forward reaction takes place at the same rate as the reverse reaction.

31. N2(g) + 3H2(g)  2 NH3(g) + heat
Reversible Reactions
Le Chatelier’s Principle: If a system at equilibrium is stressed, the equilibrium will shift in a direction that relieves that stress.
Equilibrium will shift AWAY from what is added. Here, N2 is added. N2
More “product” made
N2(g) + 3H2(g)  2 NH3(g) + heat

32. N2(g) + 3H2(g)  2 NH3(g) + heat
Reversible Reactions
Le Chatelier’s Principle: If a system at equilibrium is stressed, the equilibrium will shift in a direction that relieves that stress.
Equilibrium will shift AWAY from what is added. Here, NH3 is added.
More “reactants” made
NH3
N2(g) + 3H2(g)  2 NH3(g) + heat

33. N2(g) + 3H2(g)  2 NH3(g) + heat
Reversible Reactions
Le Chatelier’s Principle: If a system at equilibrium is stressed, the equilibrium will shift in a direction that relieves that stress.
Equilibrium will shift TOWARDS what is removed. Here H2 is removed. H2
N2(g) + 3H2(g)  2 NH3(g) + heat
More “reactants” made

34. N2(g) + 3H2(g)  2 NH3(g) + heat
Reversible Reactions
Le Chatelier’s Principle: If a system at equilibrium is stressed, the equilibrium will shift in a direction that relieves that stress.
Equilibrium will shift TOWARDS what is removed. Here heat is removed.
heat
N2(g) + 3H2(g)  2 NH3(g) + heat
More “product” made

35. pH indicates the hydrogen ion molarity [H+] in a solution
What is pH?
pH indicates the hydrogen ion molarity [H+] in a solution
pH = -log[H+]
pOH indicates the hydroxide ion molarity [OH-] in a solution.
pOH = -log[OH-]
Example: A 1.0 x 10-3 molar solution of HCl would have a pH of ___
Example: A 1.0 x 10-4 molar solution of KOH would have a pOH of ___
Memorize: pH + pOH = 14.
Example: A solution with a pH of 8 will have a pOH of: ____. 3 4 6

36. Triple Point—combination of temperature and pressure where all three phases coexist

37. PHASE CHANGES OR CHANGES OF STATE
a to b: solid increases in temperature.
b to c: solid melts to liquid at a constant temperature
c to d: liquid increases in temperature
d to e: liquid vaporizes to gas at a constant temperature
e to f: gas increases in temperature

38. GOOD LUCK ON THE SOL 