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AP Chemistry

Instructions:

Answer 50 questions in 15 minutes.
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Match each statement with the correct term.
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1. q cal = ?
are not
Effusion
Counterions
CAT

2. Neutralization Reaction (general format) (net ionic of which is always (H+) + (OH-) --> (H2O))
Acid + Base --> Salt + Water
nitrite
0.0826Latm/Kmol
blue

3. Electron (symbol)
condensation
STP
e-
p

4. All cations are soluble with bromide - chloride and iodide EXCEPT
Polar Covalent
Ag+ - Pb2+ - Hg2+
nitrate
strong acid strong base rxn

5. Kinetic Energy of an individual particle formula
M = square root (3RT/mm)
strong acids
Valence Electrons(assigned)
Molar Heat Capacity

6. CN¹?
cyanide
no precipitate forms
London dispersion forces
seesaw

7. The rest of the universe (in thermodynamics)
lambda
s
surroundings
Work

8. Moles of solute/volume of soln(L)
State Functions
Molarity
Closed System
strong acids

9. 2 or more covalently bonded atoms
Anion
wavelength
insoluble
Molecule

10. Kinetic Energy per mol
nitrate
Alkali metals
3/2RT
Theoretical yield

11. AX3E
trigonal pyramidal
Alkali metals
voltaic cells
charge

12. % yield
see-saw
sulfide
acid
actual yield/theoretical yield x 100%

13. Freezing point depression formula
Molecular
?Tf= kf x molality
permanganate
Cation

14. AX5
First-Order Rate Law
trigonal bipyramidal
Oxidizing Agent
0.512°C

15. Newton's Second Law
P of a =(X of a)(total pressure)
AH rxn = (Sum of AH of formation of products) - (Sum of AH of formation of reactants)
group
Force = mass x acceleration

16. Arrhenius equation
p+
geometric isomers
Joule
k=Ae^(-Ea/RT)

17. R in ideal gas law
Law of Conservation of Energy
0.0821 atm L/mol K
standard solution
Pi Bond

18. The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles
London Dispersion Forces
increasing
+1 - except if bonded to Alkali Metal -1
electrolyte

19. AX6
octahedral
third
Valence Electrons(assigned)
Beta Particles-

20. Ecell= E°cell -RT/nF x lnQ
Ideal Gas Law
Nernst Equation
1) Assume percentages are g (where you have 100 g) 2) convert to moles 3) divide by lowest value 4) plug into compound
adiabatic

21. Connects the 2 half cells in a voltaic cell
non-
salt bridge
insoluble
hydroxide

22. If anion ends in -ide - acid name ends in
hydro-ic acid
Polar Covalent
oxidation
Entropy (S)

23. STP
0 degrees C - 1 atm
Allotrope
Specific Heat (s)
conjugate acid

24. Similar to atomic orbitals - except between molecules
Limiting reactant
Molecular Orbitals (MOs)
but-
trigonal bipyramidal

25. Dalton's Law of Partial Pressures (to find partial pressure formula)
2nd law of thermodynamics
P of a =(X of a)(total pressure)
0.0826Latm/Kmol
-ous acid

26. When gas compresses ...
1) Convert to moles 2) divide by lowest moles 3) if any halfs - double all values 4) plug into Compound
Positive work value; work done on system
non-
Isotopes

27. Change in moles (An) =?
(moles of products that are gasses) - (moles of reactants that are gasses)
heat capacity
Ligand
allotrope

28. Aldehyde suffix
Hydrogen bonding
Gamma Ray-
-al
Antibonding Molecular Orbital

29. (organics) seven carbons
hept-
Molar Mass of Element/ Total Molar Mass %
square planar
hydrolysis

30. (organics) eight carbons
LeChatelier's Principle
insoluble
0.0821 atm L/mol K
oct-

31. Group 1 and heavier Group 2 bases
first
Monoprotic
oxide gas and water
strong bases

32. 6.022x10^23
mol
4.184
Isotopes
0

33. Each orbital can hold two e?s each w/ opposite spins
Pauli Exclusion Principle
-ous acid
zero
Amino-

34. S²?
sulfide
Effusion
?Tf= kf x molality
p+

35. r=k[A]
moles of solute/ L of solution
First-Order Rate Law
k=Ae^(-Ea/RT)
oxalate

36. For colligative properties for electrolytes - the # of mols of ions/ mols of solute

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37. q H2O = ?
third
1 atm
msAT
1) multiply each AMU by the percentage that represents that isotopes occurrence in nature 2) then add all the AMUs together.

38. Group 1 metals
Calorimeter
Electronegativity
Alkali metals
no precipitate forms

39. Consists of a complex ion - a transition metal with attached ligands - and counterions
Law of Conservation of Mass
Coordination Compound
oxide gas and water
Acid Dissociation Constant

40. Atoms with the same number of protons but a different number of neutrons
deposition
permanent gases
Isotopes
Pi Bond

41. Kinetic Energy is proportional to ______
iodide
Temperature
End Point
P1V1/N1T1=P2V2/N2T2

42. Molality =
moles solute/kg solvent
0
Molecular
s

43. Proton (symbol)
reduction
viscosity
p+
ether

44. Electron pairs found in the space between the atoms
Bonding Pairs
group
methoxy-
T-shape

45. H?+OH??H2O
Decrease Volume and Increase Temperature
strong acid strong base rxn
activation energy
freezing

46. frequency symbol
supercritical fluid
nu
?Hvap
Ionic Compounds

47. (organics) nine carbons
Volume Metric Flask & Pipet
log[H+]
Molal FP Depression Constant
non-

48. Molecules' tendency to stick to the container
Calorimetry
adhesion
Naming Binary Ionic Compounds
condensation

49. H + donor
Covalently w/in themselves - Ionicly bonded w/ each other
Law of Multiple Proportions
Bronsted-Lowry Acid
M1V1=M2V2

50. Diatomic Molecules
mol Fraction
Average KE = 1/2(mass)(average speed of all particles)
Acid + Base --> Salt + Water
H2 - F2 - N2 - O2 - Cl2 - Br2 - I2