Review Flashcards
(260 cards)
1
Q
Density = ?
A
Mass/ Volume
2
Q
Intensive Property
A
Does NOT depend on the quantity of matter present
Ex: Density, Temp.
3
Q
Extensive Property
A
Depends on the quantity of matter present
Ex: Volume, Mass
4
Q
Law of Conservation of Mass
A
Mass cannot be created nor destroyed
5
Q
Law of Definite Proportions
A
Every sample of the same compound has the same constituent elements present in the same ratio.
Ex: H2O 2 H atoms 1 O atom
6
Q
Law of Multiple Proportions
A
If two compounds contain the same elements, they are present in different ratios in each different compound
Ex: H2O vs H2O2 – both contain H & O but in different ratios
7
Q
Dalton’s Atomic Theory 1.
A
All matter consists of atoms. Atoms= indivisible particles of an element
8
Q
Dalton’s Atomic Theory 2.
A
All atoms of a given element are identical
NOT TRUE NOW –> Isotopes
9
Q
Dalton’s Atomic Theory 3.
A
Atoms of a given element can NOT be converted into atoms of a different element.
Not true now because of nuclear chemistry
10
Q
Dalton’s Atomic Theory 4.
A
Compounds are the result of combinations of atoms of different elements
–remember law of definite proportions and law of multiple proportions
11
Q
Dalton’s Atomic Theory 5.
A
Chemical reactions are the result of separation, combination, or rearrangement of atoms – do NOT result in atom creation or destruction
–remember law of conservation of mass
12
Q
Dalton did NOT
A
did not give a picture of atoms and did not discuss structure of atoms.
13
Q
Isotopes of the same element..
A
have varying number of neutrons
14
Q
Inside the nucleus
A
proton & neutron
15
Q
Outside the nucleus
A
electron
16
Q
Proton
A
+1 charge
Positively charged subatomic particles
Relative mass= 1 amu
17
Q
Neutron
A
no charge
Subatomic particle that has no charge
Relative mass= 1 amu
18
Q
AMU
A
Atomic Mass Unit
19
Q
Electron
A
-1 charge
Negatively charged subatomic particle
Relative mass= 0 amu
20
Q
Atomic #
A
of protons in the nucleus
(for neutral atoms # of protons = # of neutrons
21
Q
Atomic Mass #
A
Sum of # of protons and neutrons
22
Q
Naturally occurring (in nature) percentage of each isotope of each element
A
Natural Abundance
23
Q
Always monoatomic
A
Noble gases
24
Q
Diatomic elements
A
H2, N2, O2, F2, Cl2,Br2, I2
25
If an atom gains or loses electrons, it becomes charged
Ions
26
Cation
result of an atom losing e- (+)
27
Anion
result of an atom gaining e- (-)
28
Main group # indicates ___
of e- in outermost shell (valence e-)
29
Atoms gain or lose e- to form ions with 8 ___ e-
valence
30
Magic # of stability
8
31
Outer shell
Valence shell
32
indicate charge of ion as roman numeral after the element name
Stock Notation
33
____ have the ability to form cations of multiple charges
Transition metals
34
result of electrostatic attraction of metal cation and non-metal anion
Ionic Compounds
35
____ result of e- transfer from metal cation to non-metal anion
ionic compounds
36
_____ combine in such a way that the resulting ionic compound is neutral
metal cations and non-metal anions
37
1
Mono
38
2
Di-
39
3
Tri-
40
4
Tetra-
41
5
Penta-
42
6
Hex-
43
7
Hepta-
44
8
Octa-
45
9
Nona-
46
10
Deca-
47
Covalent compounds
result of 2 non-metal elements combining
48
Polyatomic ions
consist of two or more atoms that collectively have a charge --> act as a unit
49
Tells you exact # of every element in a substance
Chemical Formula
50
Only gives the smallest whole # ration of each element present in a compound (can be useful experimentally)
Empirical Formula
51
Molecular weight -- accounts for the mass of ALL atoms in a molecule/compound
Molar Mass
52
Ionic compound uses ____ method
cross down
53
covalent compound uses ___ method
greek prefixes
54
Transition metals uses ___ method
cross up
55
Molar Mass =
Grams/ Mole
56
There is no direct conversion from grams --> atoms must go through ____.
Moles
57
_____ are the complete sentences of chemistry and outline what happens in a chemical reaction
Balanced chemical equations
58
Reactants
Starting material of the reaction
59
Products
End result of the reaction
60
____ are used to balance chemical equations
Stoichiometric coefficients
61
_____ give mole:mole rations of reactants to products in balanced chemical equations
Stoichiometric coefficients
62
Using balanced chemical equations to do mole mass conversations for reactants and products
Stoichiometry
63
The reactant that LIMITS how much product can be found in a reaction (assuming you did the rxn perfectly)
Limiting reagent
64
Max amount of product that can be formed; based on the amount of limiting reagent used (assuming no errors)
Theoretical yield
65
Actual yield from a reaction when performed in the lab
Experimental yield
66
% yield
experimental yield/theoretical yield X 100
67
Percent yield goal
High
68
Percent error goal
Low
69
To determine amount of excess reagent remaining after the rxn is complete, you first need to determine how much _____ you should have used
Excess reagent
70
Precipitate
insoluble solid
71
Electrolytic solutions
capable of conducting electricity
72
All ionic substances that are soluble in H2O produce ___.
Ex: NaCl in H2O
electrolytic solutions
73
Non-electrolytic solutions
Ex. Sugar in H2O
do NOT conduct electricity
74
Soluble
Will dissolve
75
Insoluble
Will not dissolve
76
ionization AKA
dissociation
77
separation of a substance into ions (cations & anions)
Ex: NaCl ---H2O---> Na+ + Cl-
Ionization
78
Completely ionize/dissociate
Ex: CaCl2 ----H2O---> Ca2+ + 2Cl-
strong electrolytes
79
Incompletely dissociate/ionize
Ex: H20 H+ + OH-
weak electrolytes
80
____ form insoluble products (precipitates)
Precipitation rxns
81
Formation of a _____ is based on solubility of products.
precipitates
82
Spectator ions
- do NOT participate in the real chemistry
| - are NOT chemically changed from reactant --> product side of rxn.
83
Net ionic equation
REAL chemistry
84
Aqueous =
H2O
85
Solution
Homogeneous mixturen of a solute and a solvent
86
Solute
Substance being dissolved by solvent
87
Unsaturated solution
Contains a minimum of solute that the five amount of solvent is capable of dissolving
88
Saturated solution
contains maximum amount of solute that given amount of solvent is capable of dissolving
89
super-saturated solution
contains more solute that the given amount of solvent is capable of dissolving
90
Arrhenius acid
substance that ionizes in H2O to give H+
91
Arrhenius base
substance that ionizes in H2O to give OH-
92
Polyprotic acids
Step-wise dissociation of H+
93
Weak Acids
Incompletely disociate
94
Bronsted acid
proton donor = H+ donor
95
Bronsted base
proton acceptor = H+ acceptor
96
Amphoteric
substance that can act as both an acid or base
97
Strong bases undergo _____
complete -----> dissociation
98
Weak bases undergo _______
incomplete dissociation
99
All Arrhenius acids are also classified as
Bronsted acids
100
Acid-base titrations
analytical method for determining concentration of an unknown acid or base
101
For acid-base titrations you must know _____
the concentration of either the acid or the base, then can experimentally determine the concentration of the other one.
102
Equivalence point
point in titration when stoichiometric equivalence of solution of known concentration has been added to solution of unknown concentration of complete rxn
103
End point
point in titration when indicator changes color
104
Goal in titration
want end point to be as close to equivalence point as possible
105
Dilution calculation
M1V1 = M2V2
106
Oxidation-Reduction rxns
"redox" because oxidation and reduction are always simultaneous processes (1/2 rxns) in the same overall rxn.
107
LEO GER
loss of e- = oxidation
| gain of e- = reduction
108
Redox reactions result in ____
changing of oxidation numbers
109
Oxidizing agent =
reactant that undergoes reduction
110
Reducing agent=
reactant that undergoes oxidation
111
Boyle's law
Pressure-Volume relationship of gases
P1 V1= P2 V2
112
P 1/V @ constant n & T
Boyle's Law
113
Charles Law
Volume-Temp relationship of gas
114
V T @ constant n & P
Charles' Law
115
Variation of Charles Law
P T @ constant n & V
116
Avogadro's Law
Volume - Mole (amount) relationship of gases
117
V n @ constant P & T
Avogadro's Law
118
Ideal Gas Law
Summarizes behavior of ideal gases
119
Pressure ____
most readily measurable property of gases
120
Normal atmospheric pressure
1 atm
121
Normal (lab) conditions
```
P= 1.00 atm
T= 25' C
```
122
Standard conditions (STP)
```
P= 1.00 atm
Temp= 0' C
```
123
Gas stoichometry
Rxn stoichiometry calculations blended w/ gas laws
124
Dalton's Law of Partial Pressures
Total pressure of a mixture of gases= sum of the individual partial pressures of each component gas of the mixture
125
How is the partial pressure of each gas related to the total pressure?
Xi= ni/nt
126
Xi=
Unitless; mole fraction of an individual component of mixture
127
ni=
moles of individual component of mixture
128
nt=
total # moles
129
Pi=Xi Pt
```
Pi= partial pressure of each individual component of mixture
Xi= mole fraction
Pt= total pressure
```
130
The sum of partial pressures should equal ____
the total pressure
131
The kinetic molecular theory of gases
summarizes gas behavior
132
At the molecular level:
increase in temp causes gas molecules to move more because they have more kinetic energy
133
When molecules collide with the sides of a container we get a significant amount of ___
pressure build up (b/c the container does NOT react to the collision)
134
For real gases we need to do a small correction for the press and volume --- luckily we have ____
Vander Waals Equation for Non-Ideal Gases
135
The study of heat associated with chemical reactions
Thermochemistry
136
Law of Conservation of Energy
Energy cannot be created nor destroyed but it can change from one form to another
137
The total amount of ____ in the universe is constant
Energy
138
Almost all reactions absorb or release energy in the form of ____
heat
139
Heat =
transfer of thermal energy between 2 objects that are at different temperatures
140
Heat flows ____
Hot -----> Cold
141
Energy associated with chemical reactions considers
system and surroundings
142
3 types of systems
open system
closed system
isolated system
143
____ allows mass and heat transfer between the system and surroundings
Ex: beaker in the lab
Open system
144
___ does NOT allow mass transfers between system and surroundings but does allow heat transfer
Ex: beaker with a lid in the lab
Closed system
145
____ does NOT allow mass or heat transfer between system and surroundings
Ex: insulated container with lid
Isolated system
146
Endothermic rxns
requires heat to be absorbed in order for the rxn to proceed
147
Exothermic rxns
releases heat as a consequence of the rxn occuring
148
Endothermic rxn diagram
products
^
^
149
Exothermic rxn diagram
reactants
v -----> heat
v
products
150
The study of heat and its interconversions in chemical reactions
Thermodynamics
151
Microscopic properties that we can measure
Ex: composition, energy, temp, pressure, volume
State of a system
152
Property of the system that is defined by the state of the system
State function
153
Changes in state functions are ______. They only depend on initial and final states
independent of pathway
154
Calculating changes in state functions
Delta = final - initial
155
Thermodynamic quantity that allows investigation of the heat changes associated with chemical rxns
Enthalpy (H)
156
_____ is a state function
Enthalpy
157
___ give chemical rxn and associated Delta H for the rxn
Thermochemical rxns
158
Rules for Thermochemical rxns
- States (physical) of substance are important. If the physical states change, the enthalpy changes
- Thermochemical rxns are stoichiometric
159
If you change the direction of a rxn, then you must change ____ of DH (delta H)
the sign
160
Endothermic rxns = ___ DH (delta H)
+ (postive)
161
Exothermic rxns = ___ DH (delta H)
- (negative)
162
measuring heat changes associated with chemical reactions
Calorimetry
163
amount of heat required to raise the temp of 1 gram of a substance by 1 degree C
Specific Heat
164
C =
the speed of light (m/s)
165
E=
energy (J)
166
V=
frequency = 1/s
167
h=
Plank's constant = 6.63X10^ -34
168
Visible region of Electromagnetic Spectrum (EMS)
ROY G BIV
169
de Broglie Wavelength equation
"particle-like properties of light"
wavelength = h/mu
170
particle-like AND wave-like properties
Duality of Light
171
Rydberg's constant
2.18X10^-18
172
n=
energy level of e- (referring to shell)
an increase in n is a increase from distance from nucleus
173
DE (delta E AKA change in E)
E final - E initial
174
Q#
Quantum numbers
175
Principle Q# (n)
"street of the e-"
176
Angular momentum Q# (L)
"shape of house"
| (L) shape of orbital-- identifies orbital
177
"nL"
subshell
178
L=0
s-orbital
179
L=1
p-orbital
180
L=2
d-orbital
181
L=3
f-orbital
182
possible values of L
0....... (n-1)
183
Magnetic Q#=
M sub L
gives orientation of orbital
-L ......0.....+L
184
Any given orbital can only hold a max of ___e-
2
185
Electron Spin Q# =
M sub S-- defines clock-wise or counter clockwise spin of 2e- occupying the same orbital
+1/2 or -1/2
186
Aubauf Principle=
building-up principle of the periodic table
187
As atomic # increases 1 additional proton in nucleus and 1 additional e- outside nucleus
Aubauf Principle
188
Pauli Exclusion Principle
NO two e- can have all four Q#'s identical
*If 2e- are in the same orbital they will have identical "n", "L" and "M sub L" values but values of M sub S will be different (1/2 or -1/2) to designate different spins (clockwise vs counterclockwise)
189
identifies ALL e-
Expanded Electron Configuration
190
the most stable configuration of e- in the same subshell include the greatest number of parallel spins.
Hund's Rule
191
Exceptions to Hund's rule:
- There is slightly greater stability when 1 e- from a "ns" orbital can be "borrowed" from the preceding "(n-1)d" orbital to half-fill (d5) or completely fill (d10) the (n-1)d subshell
- Only effects Cr group and Cu group
- Also observed in the f-block with Sm group and Tm group
192
Cations =
+ ions = result of LOSING electrons from the valence shell
193
____ lose electrons first from the "ns" shell before losing from the "(n-1)d" shell
Transition metals
194
Anions=
- ions = result of GAINING electrons into the valence shell
195
Isoelectric (isoelectronic)
have the same electron configuration (EC)
196
Uses the chemical symbol for elements and adds dots to signify valence electrons
Lewis Dot Structure
197
Ionic bonds=
result of e- TRANSFER from metal cation to non-metal anion; electrostatic attraction
Ex: NaCl
198
Covalent bonds=
result of sharing of electrons between two atoms
199
Polar covalent bonds
result of UNEQUAL sharing of e-
200
Non-polar covalent bonds
result of EQUAL sharing of electrons
201
Electronegativity (EN)
relative measure of the attraction of an atom for electrons
202
How well and atom "pulls" e- density towards itself
Electronegativity
203
___ is the most electronegative atom
F
204
Electronegativity ___ on the periodic table towards F
increases
205
We use the ____ of EN that gives F and EN value of 4.0
Pauling Scale
206
____ can be used to predict types of bonds between atoms by comparing the difference between the EN values of the two atoms in the bond.
Electronegativity
207
> 1.6
ionic bond
208
0.5 - 1.6
polar covalent bond
209
1 - 0.4
non-polar covalent bond
210
__ atoms have ONLY single bonds and are ALWAYS at the end of structures
H ; will never be the central atom
211
Polyatomic molecules and ions often consist of a ___ EN central atom surrounded by ____ EN atoms.
(exception: H)
less; more
212
valence electrons in an isolated atom - # lone electrons (dots) - # bonds (lines)
formal charge
213
structures with the lowest distribution of formal charge
Most plausible Lewis Structures
214
some molecules are very stable with LESS than 8 electrons around the central atom
Ex: SF6; PCl5
Incomplete Octet
215
Some molecules are very stable with MORE than 8 electrons around the central atoms because of possible expansion into d shells of similar energy to valence shells of central atom
Expanded Octet
216
Odd number of electrons-- ____--- contain unpaired electrons -- very unstable
Ex: NO2
Radicals
217
VSEPR
Valence Shell Electron Pair Repulsion
218
___ can be used to predict 3D shape-- if assume that electrons in valence shell of an atom repel one another
Lewis Structure
219
The 3D geometry that a molecule assumes minimizes ____.
repulsions (gets electron pairs as far away from each other as possible)
220
Use DEN (Delta EN) between atoms in each bond to predict whether bonds are ionic, polar, or non-polar
Bond Polarity
221
Use bond polarity and molecular geometry to determine whether a molecule is ionic, polar, or covalent
Molecular Polarity
222
The distance between two consecutive peaks or troughs in a wave
wavelength
223
number of waves (cycles) per second that pass a given point in space
Frequency (v)
224
____ radiation has a higher frequency when compared to _____ radiation
Short-wavelength
Long-wavelength
225
___ is in meters
wavelength
226
___ postulated that energy can be gained or lost only in whole-number multiples of hv
Planck
227
Energy is ____ and can occur in discrete units of hv
quantized
228
A packet of energy
Quantum
229
____ proposed that electromagnetic radiation is a stream of photons
Einstein
230
Phenomenon in which electrons are emitted from the surface of metal when light strikes it
Photoelectric Effect
231
The ____ of the emitted electrons increases linearly with the frequency of the light
kinetic energy (KE)
232
Einstein proposed that ___ has mass.
energy
233
__ ascertained matter that is assumed to be particulate exhibits wave properties
de Broglie
234
Results when white light is passed through a prism
Continuous spectrum
235
Shows only certain discrete wavelengths
Ex: H emission spectrum
Line spectrum
236
Who came up with the Quantum Model for the Hydrogen Atom?
Niels Bohr
237
The electron in a hydrogen atom moves around the nucleus in certain allows circular orbits
Quantum model
238
Single bond
2e- total
239
double bond
4e- total
240
triple bond
6e- total
241
quadruple bond
8e- total
242
an increase in multiplicity increases
bond strength
243
an increase in multiplicity decreases
bond length
244
Summarizes common observations of different systems
Law
245
Withstands the test of time
Law
246
Attempts to explain WHY something happens
Theory
247
Measure of the amount of matter in an object
Mass
248
Force exerted by gravity on an object
Weight
249
Agreement of a particular value with the TRUE value
Accuracy
250
Agreement among several measurements of the same quantity
Precision
251
Anything that occupies space and has mass
Matter
252
Has visibly indistinguishable parts (solutions)
Homogenous mixture
253
Has visibly distinguishable parts
Heterogenous
254
Can be separated into pure substances by physical methods
Mixtures
255
A substance with a constant composition that can be broken down into its elements via chemical processes
Compound
256
Substance that cannot be broken down into simpler substances by physical or chemical means
Element
257
Boiling or freezing water
Physical change
258
Burning logs; chemical rxn bubbling
Chemical change
259
Anything that can be classified as matter
| Ex: compounds, elements
Substance
260
The reactant that is used in an excess amount (more than is really needed) to make the theoretical yield of product.
Excess reagent
