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AP Chemistry > Test 1 > Flashcards

Test 1 Flashcards

1
Q

direct relationship:

A

as y increases, x increases

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2
Q

direct relationship formula

A

y=Kx

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3
Q

inverse relationship:

A

as x increases, y decreases

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4
Q

inverse relationship formula

A

y=k/x

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5
Q

exponential relationship:

A

as y increases, x increases a lot more

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6
Q

exponential relationship formula

A

y=Kx^n

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7
Q

ionization energy

A

amount of energy to remove an electron from an atom

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8
Q

Coulomb’s law(inverse square) formula

A

F= (K(Q1Q2))/(r^2)

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9
Q

use Coulomb’s law to…

A

define all trends on the periodic table

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10
Q

photoelectron spectroscopy

A

discipline of identifying unknown pure substances by using x-rays and ionization energy to produce a spectrum

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11
Q

hybridization

A

the process of orbitals hybridizing or fusing together when 2 atoms share electrons(covalent) or when 2 atoms interact with electrostatic force(ionic)

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12
Q

how to determine the hybridization of an atom

A

count number of atoms it’s bonded to and add it to the number of unshared pairs of electrons on it

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13
Q

hybridization 1

A

s

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14
Q

hybridization 2

A

sp

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15
Q

hybridization 3

A

sp^2

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16
Q

hybridization 4

A

sp^3

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17
Q

hybridization 5

A

dsp^3

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18
Q

hybridization 6

A

d^2sp^3

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19
Q

bond order

A

the number of pairs of electrons shared between 2 atoms

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20
Q

bond order single bond

A

1

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21
Q

bond order double bond

A

2

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22
Q

bond order triple bond

A

3

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23
Q

if a structure has resonance

A

take the number of bonds and divide by the number of like atoms sharing a bond

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24
Q

electronegativity

A

an atoms ability to attract electrons to itself

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25
electronegativity trend
up and to the right
26
intramolecular bond
bond between atoms
27
three types of intramolecular bonds
ionic, covalent, polar covalent
28
polar molecule
molecule with uneven charges because the atoms have different electronegativity values
29
dipole
having opposites charged ends
30
if a molecule is diatomic and
the bond between the atoms is polar, then the molecule itself is polar
31
if a molecule has more than 2 atoms
the polarity of the bond and the 3-D shape must be considered
32
if the dipoles go in opposite directions
they cancel each other out forming a non polar molecule
33
intermolecular forces
weak interaction between molecules in liquids and solids
34
intermolecular forces are responsible for
1.solids and liquids staying together 2.boiling point, melting point, solubility, and vapor pressure
35
IMFs from strongest to weakest
ion-dipole, H bonding, dipole-dipole, london dispersion forces(LDFs)
36
ion-dipole
interaction between an ion(dissolved salt) and a polar molecule
37
water molecules surround the ions forming
hydration shells
38
H bonding(hydrogen bonding)
interaction between the hydrogen of one molecule and the F O or N
39
dipole-dipole
the positive end of a polar molecule interacts with the negative end of another molecule
40
london dispersion forces(LDFs)
weakest IMF that both polar and non-polar molecules have
41
if non-polar substances collide
causing a temporary induced dipole that causes more interaction
42
the larger the atom
the stronger the dispersion forces
43
5 types of reactions
synthesis, decomposition, combustion, single displacement, and double displacement
44
neutralization reaction
reaction that is exothermic when an acid is combined with a base producing a salt and water
45
precipitation reaction
reaction that forms an insoluble solid(precipitate) from the mixing of 2 aqueous solutions
46
ap solubility rule
salts made with group 1A elements, ammonium(NH4), and nitrate are soluble
47
ammonium hydroxide
NH4OH -> NH3 + H2O
48
carbonic acid
H2CO3 -> CO2 + H2O
49
dissociation
process of separating positive ions from negative ions in a polar solvent
50
complete ionic equation
balanced chemical equation where soluble salts are expressed in their dissociated state and insoluble salts are expressed in their associated state
51
spectator ions
ions that are present on both sides of a chemical equation that don't participate in a chemical reaction
52
net ionic equation
balanced equation once the spectator ions have been removed
53
1st law of thermodynamics
energy cannot be created nor destroyed; only converted from one form to another
54
enthalpy
change in the heat content of a system from a reaction or phase change
55
+ delta H
endothermic reaction
56
-delta H
exothermic reaction
57
delta H notation
placing the value for enthalpy beside a balanced chemical equation
58
thermochemical notation
the enthalpy value is incorporated into the balanced equation
59
Hess's law(law of heat summation)
the enthalpy value for a chemical reaction is equal to the sum of the enthalpy values of the reactions that make up the reaction
60
which sign of enthalpy is thermodynamically favored
-deltaH
61
standard enthalpy of formation of the reaction=
the standard enthalpy of the products - reactants
62
specific heat equation
q=cxmxdeltaT
63
entropy(s)
measurement of disorder in a system; measured in J/K or J/K(mol)
64
entropy is not a
measurement of energy but how energy is dispersed
65
2nd law of thermodynamics
entropy increases in a closed system
66
which entropy is thermodynamically favored
+deltas
67
microstate
all the possible positions that particles can occupy at a particular time
68
entropy can increase when ___ increases
microstates, electrons, and atoms in cmpd
69
Boltzmann equation
s=Kb(ln(W))
70
Boltzmann's constant
1.38 x 10^-23 J/K
71
3rd law of thermodynamics
the entropy of a pure crystal at absolute zero is zero(only one micro state is possible. the ln(1)=0)
72
positional disorder
based on the position of particles(microstates)
73
thermal disorder
dispersion of heat and energy( increased temperature, increased entropy)
74
delta entropy=
entropy of the system + entropy of the surroundings
75
entropy of surroundings=
-deltaH(1000)/T
76
entropy of system=
delta H/T
77
standard entropy of formation=
standard entropy of formation of the products - reactants
78
gibbs free energy
third thermodynamic quantity measured in kJ/mol that is defined as the energy available to do work in a system
79
thermodynamically favored gibbs free energy
-delta G
80
Gibbs-Hemholtz Equation
deltaG= deltaH - T(deltaS) temp in K covert entropy delta S to kJ
81
bond energy(BE)
energy required to break a bond in 1 mole of a gaseous molecule
82
bond energy requirements
-gaseous products and reactants -values are experimentally determined -units in kJ -known lewis structure
83
bond energy=
sum of reactants - products

AP Chemistry (16 decks)

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