Review Flashcards

1
Q

Density = ?

A

Mass/ Volume

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2
Q

Intensive Property

A

Does NOT depend on the quantity of matter present

Ex: Density, Temp.

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3
Q

Extensive Property

A

Depends on the quantity of matter present

Ex: Volume, Mass

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4
Q

Law of Conservation of Mass

A

Mass cannot be created nor destroyed

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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

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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

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7
Q

Dalton’s Atomic Theory 1.

A

All matter consists of atoms. Atoms= indivisible particles of an element

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8
Q

Dalton’s Atomic Theory 2.

A

All atoms of a given element are identical

NOT TRUE NOW –> Isotopes

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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

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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

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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

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12
Q

Dalton did NOT

A

did not give a picture of atoms and did not discuss structure of atoms.

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13
Q

Isotopes of the same element..

A

have varying number of neutrons

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14
Q

Inside the nucleus

A

proton & neutron

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15
Q

Outside the nucleus

A

electron

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16
Q

Proton

A

+1 charge
Positively charged subatomic particles
Relative mass= 1 amu

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17
Q

Neutron

A

no charge
Subatomic particle that has no charge
Relative mass= 1 amu

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18
Q

AMU

A

Atomic Mass Unit

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19
Q

Electron

A

-1 charge
Negatively charged subatomic particle
Relative mass= 0 amu

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20
Q

Atomic #

A

of protons in the nucleus

(for neutral atoms # of protons = # of neutrons

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21
Q

Atomic Mass #

A

Sum of # of protons and neutrons

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22
Q

Naturally occurring (in nature) percentage of each isotope of each element

A

Natural Abundance

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23
Q

Always monoatomic

A

Noble gases

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24
Q

Diatomic elements

A

H2, N2, O2, F2, Cl2,Br2, I2

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25
Q

If an atom gains or loses electrons, it becomes charged

A

Ions

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26
Q

Cation

A

result of an atom losing e- (+)

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27
Q

Anion

A

result of an atom gaining e- (-)

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28
Q

Main group # indicates ___

A

of e- in outermost shell (valence e-)

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29
Q

Atoms gain or lose e- to form ions with 8 ___ e-

A

valence

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30
Q

Magic # of stability

A

8

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31
Q

Outer shell

A

Valence shell

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32
Q

indicate charge of ion as roman numeral after the element name

A

Stock Notation

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33
Q

____ have the ability to form cations of multiple charges

A

Transition metals

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34
Q

result of electrostatic attraction of metal cation and non-metal anion

A

Ionic Compounds

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35
Q

____ result of e- transfer from metal cation to non-metal anion

A

ionic compounds

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36
Q

_____ combine in such a way that the resulting ionic compound is neutral

A

metal cations and non-metal anions

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37
Q

1

A

Mono

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38
Q

2

A

Di-

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39
Q

3

A

Tri-

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40
Q

4

A

Tetra-

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41
Q

5

A

Penta-

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42
Q

6

A

Hex-

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43
Q

7

A

Hepta-

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44
Q

8

A

Octa-

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45
Q

9

A

Nona-

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46
Q

10

A

Deca-

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47
Q

Covalent compounds

A

result of 2 non-metal elements combining

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48
Q

Polyatomic ions

A

consist of two or more atoms that collectively have a charge –> act as a unit

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49
Q

Tells you exact # of every element in a substance

A

Chemical Formula

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50
Q

Only gives the smallest whole # ration of each element present in a compound (can be useful experimentally)

A

Empirical Formula

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51
Q

Molecular weight – accounts for the mass of ALL atoms in a molecule/compound

A

Molar Mass

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52
Q

Ionic compound uses ____ method

A

cross down

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53
Q

covalent compound uses ___ method

A

greek prefixes

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54
Q

Transition metals uses ___ method

A

cross up

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55
Q

Molar Mass =

A

Grams/ Mole

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56
Q

There is no direct conversion from grams –> atoms must go through ____.

A

Moles

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57
Q

_____ are the complete sentences of chemistry and outline what happens in a chemical reaction

A

Balanced chemical equations

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58
Q

Reactants

A

Starting material of the reaction

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59
Q

Products

A

End result of the reaction

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60
Q

____ are used to balance chemical equations

A

Stoichiometric coefficients

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61
Q

_____ give mole:mole rations of reactants to products in balanced chemical equations

A

Stoichiometric coefficients

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62
Q

Using balanced chemical equations to do mole mass conversations for reactants and products

A

Stoichiometry

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63
Q

The reactant that LIMITS how much product can be found in a reaction (assuming you did the rxn perfectly)

A

Limiting reagent

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64
Q

Max amount of product that can be formed; based on the amount of limiting reagent used (assuming no errors)

A

Theoretical yield

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65
Q

Actual yield from a reaction when performed in the lab

A

Experimental yield

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66
Q

% yield

A

experimental yield/theoretical yield X 100

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67
Q

Percent yield goal

A

High

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68
Q

Percent error goal

A

Low

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69
Q

To determine amount of excess reagent remaining after the rxn is complete, you first need to determine how much _____ you should have used

A

Excess reagent

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70
Q

Precipitate

A

insoluble solid

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71
Q

Electrolytic solutions

A

capable of conducting electricity

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72
Q

All ionic substances that are soluble in H2O produce ___.

Ex: NaCl in H2O

A

electrolytic solutions

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73
Q

Non-electrolytic solutions

Ex. Sugar in H2O

A

do NOT conduct electricity

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74
Q

Soluble

A

Will dissolve

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75
Q

Insoluble

A

Will not dissolve

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76
Q

ionization AKA

A

dissociation

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77
Q

separation of a substance into ions (cations & anions)

Ex: NaCl —H2O—> Na+ + Cl-

A

Ionization

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78
Q

Completely ionize/dissociate

Ex: CaCl2 —-H2O—> Ca2+ + 2Cl-

A

strong electrolytes

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79
Q

Incompletely dissociate/ionize

Ex: H20 H+ + OH-

A

weak electrolytes

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80
Q

____ form insoluble products (precipitates)

A

Precipitation rxns

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81
Q

Formation of a _____ is based on solubility of products.

A

precipitates

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82
Q

Spectator ions

A
  • do NOT participate in the real chemistry

- are NOT chemically changed from reactant –> product side of rxn.

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83
Q

Net ionic equation

A

REAL chemistry

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84
Q

Aqueous =

A

H2O

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85
Q

Solution

A

Homogeneous mixturen of a solute and a solvent

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86
Q

Solute

A

Substance being dissolved by solvent

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87
Q

Unsaturated solution

A

Contains a minimum of solute that the five amount of solvent is capable of dissolving

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88
Q

Saturated solution

A

contains maximum amount of solute that given amount of solvent is capable of dissolving

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89
Q

super-saturated solution

A

contains more solute that the given amount of solvent is capable of dissolving

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90
Q

Arrhenius acid

A

substance that ionizes in H2O to give H+

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91
Q

Arrhenius base

A

substance that ionizes in H2O to give OH-

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92
Q

Polyprotic acids

A

Step-wise dissociation of H+

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93
Q

Weak Acids

A

Incompletely disociate

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94
Q

Bronsted acid

A

proton donor = H+ donor

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95
Q

Bronsted base

A

proton acceptor = H+ acceptor

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96
Q

Amphoteric

A

substance that can act as both an acid or base

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97
Q

Strong bases undergo _____

A

complete —–> dissociation

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98
Q

Weak bases undergo _______

A

incomplete dissociation

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99
Q

All Arrhenius acids are also classified as

A

Bronsted acids

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100
Q

Acid-base titrations

A

analytical method for determining concentration of an unknown acid or base

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101
Q

For acid-base titrations you must know _____

A

the concentration of either the acid or the base, then can experimentally determine the concentration of the other one.

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102
Q

Equivalence point

A

point in titration when stoichiometric equivalence of solution of known concentration has been added to solution of unknown concentration of complete rxn

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103
Q

End point

A

point in titration when indicator changes color

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104
Q

Goal in titration

A

want end point to be as close to equivalence point as possible

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105
Q

Dilution calculation

A

M1V1 = M2V2

106
Q

Oxidation-Reduction rxns

A

“redox” because oxidation and reduction are always simultaneous processes (1/2 rxns) in the same overall rxn.

107
Q

LEO GER

A

loss of e- = oxidation

gain of e- = reduction

108
Q

Redox reactions result in ____

A

changing of oxidation numbers

109
Q

Oxidizing agent =

A

reactant that undergoes reduction

110
Q

Reducing agent=

A

reactant that undergoes oxidation

111
Q

Boyle’s law

A

Pressure-Volume relationship of gases

P1 V1= P2 V2

112
Q

P 1/V @ constant n & T

A

Boyle’s Law

113
Q

Charles Law

A

Volume-Temp relationship of gas

114
Q

V T @ constant n & P

A

Charles’ Law

115
Q

Variation of Charles Law

A

P T @ constant n & V

116
Q

Avogadro’s Law

A

Volume - Mole (amount) relationship of gases

117
Q

V n @ constant P & T

A

Avogadro’s Law

118
Q

Ideal Gas Law

A

Summarizes behavior of ideal gases

119
Q

Pressure ____

A

most readily measurable property of gases

120
Q

Normal atmospheric pressure

A

1 atm

121
Q

Normal (lab) conditions

A
P= 1.00 atm
T= 25' C
122
Q

Standard conditions (STP)

A
P= 1.00 atm
Temp= 0' C
123
Q

Gas stoichometry

A

Rxn stoichiometry calculations blended w/ gas laws

124
Q

Dalton’s Law of Partial Pressures

A

Total pressure of a mixture of gases= sum of the individual partial pressures of each component gas of the mixture

125
Q

How is the partial pressure of each gas related to the total pressure?

A

Xi= ni/nt

126
Q

Xi=

A

Unitless; mole fraction of an individual component of mixture

127
Q

ni=

A

moles of individual component of mixture

128
Q

nt=

A

total # moles

129
Q

Pi=Xi Pt

A
Pi= partial pressure of each individual component of mixture
Xi= mole fraction
Pt= total pressure
130
Q

The sum of partial pressures should equal ____

A

the total pressure

131
Q

The kinetic molecular theory of gases

A

summarizes gas behavior

132
Q

At the molecular level:

A

increase in temp causes gas molecules to move more because they have more kinetic energy

133
Q

When molecules collide with the sides of a container we get a significant amount of ___

A

pressure build up (b/c the container does NOT react to the collision)

134
Q

For real gases we need to do a small correction for the press and volume — luckily we have ____

A

Vander Waals Equation for Non-Ideal Gases

135
Q

The study of heat associated with chemical reactions

A

Thermochemistry

136
Q

Law of Conservation of Energy

A

Energy cannot be created nor destroyed but it can change from one form to another

137
Q

The total amount of ____ in the universe is constant

A

Energy

138
Q

Almost all reactions absorb or release energy in the form of ____

A

heat

139
Q

Heat =

A

transfer of thermal energy between 2 objects that are at different temperatures

140
Q

Heat flows ____

A

Hot —–> Cold

141
Q

Energy associated with chemical reactions considers

A

system and surroundings

142
Q

3 types of systems

A

open system
closed system
isolated system

143
Q

____ allows mass and heat transfer between the system and surroundings

Ex: beaker in the lab

A

Open system

144
Q

___ does NOT allow mass transfers between system and surroundings but does allow heat transfer

Ex: beaker with a lid in the lab

A

Closed system

145
Q

____ does NOT allow mass or heat transfer between system and surroundings

Ex: insulated container with lid

A

Isolated system

146
Q

Endothermic rxns

A

requires heat to be absorbed in order for the rxn to proceed

147
Q

Exothermic rxns

A

releases heat as a consequence of the rxn occuring

148
Q

Endothermic rxn diagram

A

products

149
Q

Exothermic rxn diagram

A

reactants
v —–> heat
v
products

150
Q

The study of heat and its interconversions in chemical reactions

A

Thermodynamics

151
Q

Microscopic properties that we can measure

Ex: composition, energy, temp, pressure, volume

A

State of a system

152
Q

Property of the system that is defined by the state of the system

A

State function

153
Q

Changes in state functions are ______. They only depend on initial and final states

A

independent of pathway

154
Q

Calculating changes in state functions

A

Delta = final - initial

155
Q

Thermodynamic quantity that allows investigation of the heat changes associated with chemical rxns

A

Enthalpy (H)

156
Q

_____ is a state function

A

Enthalpy

157
Q

___ give chemical rxn and associated Delta H for the rxn

A

Thermochemical rxns

158
Q

Rules for Thermochemical rxns

A
  • States (physical) of substance are important. If the physical states change, the enthalpy changes
  • Thermochemical rxns are stoichiometric
159
Q

If you change the direction of a rxn, then you must change ____ of DH (delta H)

A

the sign

160
Q

Endothermic rxns = ___ DH (delta H)

A

+ (postive)

161
Q

Exothermic rxns = ___ DH (delta H)

A
  • (negative)
162
Q

measuring heat changes associated with chemical reactions

A

Calorimetry

163
Q

amount of heat required to raise the temp of 1 gram of a substance by 1 degree C

A

Specific Heat

164
Q

C =

A

the speed of light (m/s)

165
Q

E=

A

energy (J)

166
Q

V=

A

frequency = 1/s

167
Q

h=

A

Plank’s constant = 6.63X10^ -34

168
Q

Visible region of Electromagnetic Spectrum (EMS)

A

ROY G BIV

169
Q

de Broglie Wavelength equation

A

“particle-like properties of light”

wavelength = h/mu

170
Q

particle-like AND wave-like properties

A

Duality of Light

171
Q

Rydberg’s constant

A

2.18X10^-18

172
Q

n=

A

energy level of e- (referring to shell)

an increase in n is a increase from distance from nucleus

173
Q

DE (delta E AKA change in E)

A

E final - E initial

174
Q

Q#

A

Quantum numbers

175
Q

Principle Q# (n)

A

“street of the e-“

176
Q

Angular momentum Q# (L)

A

“shape of house”

(L) shape of orbital– identifies orbital

177
Q

“nL”

A

subshell

178
Q

L=0

A

s-orbital

179
Q

L=1

A

p-orbital

180
Q

L=2

A

d-orbital

181
Q

L=3

A

f-orbital

182
Q

possible values of L

A

0……. (n-1)

183
Q

Magnetic Q#=

A

M sub L
gives orientation of orbital

-L ……0…..+L

184
Q

Any given orbital can only hold a max of ___e-

A

2

185
Q

Electron Spin Q# =

A

M sub S– defines clock-wise or counter clockwise spin of 2e- occupying the same orbital

+1/2 or -1/2

186
Q

Aubauf Principle=

A

building-up principle of the periodic table

187
Q

As atomic # increases 1 additional proton in nucleus and 1 additional e- outside nucleus

A

Aubauf Principle

188
Q

Pauli Exclusion Principle

A

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
Q

identifies ALL e-

A

Expanded Electron Configuration

190
Q

the most stable configuration of e- in the same subshell include the greatest number of parallel spins.

A

Hund’s Rule

191
Q

Exceptions to Hund’s rule:

A
  • 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
Q

Cations =

A

+ ions = result of LOSING electrons from the valence shell

193
Q

____ lose electrons first from the “ns” shell before losing from the “(n-1)d” shell

A

Transition metals

194
Q

Anions=

A
  • ions = result of GAINING electrons into the valence shell
195
Q

Isoelectric (isoelectronic)

A

have the same electron configuration (EC)

196
Q

Uses the chemical symbol for elements and adds dots to signify valence electrons

A

Lewis Dot Structure

197
Q

Ionic bonds=

A

result of e- TRANSFER from metal cation to non-metal anion; electrostatic attraction

Ex: NaCl

198
Q

Covalent bonds=

A

result of sharing of electrons between two atoms

199
Q

Polar covalent bonds

A

result of UNEQUAL sharing of e-

200
Q

Non-polar covalent bonds

A

result of EQUAL sharing of electrons

201
Q

Electronegativity (EN)

A

relative measure of the attraction of an atom for electrons

202
Q

How well and atom “pulls” e- density towards itself

A

Electronegativity

203
Q

___ is the most electronegative atom

A

F

204
Q

Electronegativity ___ on the periodic table towards F

A

increases

205
Q

We use the ____ of EN that gives F and EN value of 4.0

A

Pauling Scale

206
Q

____ 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.

A

Electronegativity

207
Q

> 1.6

A

ionic bond

208
Q

0.5 - 1.6

A

polar covalent bond

209
Q

1 - 0.4

A

non-polar covalent bond

210
Q

__ atoms have ONLY single bonds and are ALWAYS at the end of structures

A

H ; will never be the central atom

211
Q

Polyatomic molecules and ions often consist of a ___ EN central atom surrounded by ____ EN atoms.
(exception: H)

A

less; more

212
Q

valence electrons in an isolated atom - # lone electrons (dots) - # bonds (lines)

A

formal charge

213
Q

structures with the lowest distribution of formal charge

A

Most plausible Lewis Structures

214
Q

some molecules are very stable with LESS than 8 electrons around the central atom
Ex: SF6; PCl5

A

Incomplete Octet

215
Q

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

A

Expanded Octet

216
Q

Odd number of electrons– ____— contain unpaired electrons – very unstable

Ex: NO2

A

Radicals

217
Q

VSEPR

A

Valence Shell Electron Pair Repulsion

218
Q

___ can be used to predict 3D shape– if assume that electrons in valence shell of an atom repel one another

A

Lewis Structure

219
Q

The 3D geometry that a molecule assumes minimizes ____.

A

repulsions (gets electron pairs as far away from each other as possible)

220
Q

Use DEN (Delta EN) between atoms in each bond to predict whether bonds are ionic, polar, or non-polar

A

Bond Polarity

221
Q

Use bond polarity and molecular geometry to determine whether a molecule is ionic, polar, or covalent

A

Molecular Polarity

222
Q

The distance between two consecutive peaks or troughs in a wave

A

wavelength

223
Q

number of waves (cycles) per second that pass a given point in space

A

Frequency (v)

224
Q

____ radiation has a higher frequency when compared to _____ radiation

A

Short-wavelength

Long-wavelength

225
Q

___ is in meters

A

wavelength

226
Q

___ postulated that energy can be gained or lost only in whole-number multiples of hv

A

Planck

227
Q

Energy is ____ and can occur in discrete units of hv

A

quantized

228
Q

A packet of energy

A

Quantum

229
Q

____ proposed that electromagnetic radiation is a stream of photons

A

Einstein

230
Q

Phenomenon in which electrons are emitted from the surface of metal when light strikes it

A

Photoelectric Effect

231
Q

The ____ of the emitted electrons increases linearly with the frequency of the light

A

kinetic energy (KE)

232
Q

Einstein proposed that ___ has mass.

A

energy

233
Q

__ ascertained matter that is assumed to be particulate exhibits wave properties

A

de Broglie

234
Q

Results when white light is passed through a prism

A

Continuous spectrum

235
Q

Shows only certain discrete wavelengths

Ex: H emission spectrum

A

Line spectrum

236
Q

Who came up with the Quantum Model for the Hydrogen Atom?

A

Niels Bohr

237
Q

The electron in a hydrogen atom moves around the nucleus in certain allows circular orbits

A

Quantum model

238
Q

Single bond

A

2e- total

239
Q

double bond

A

4e- total

240
Q

triple bond

A

6e- total

241
Q

quadruple bond

A

8e- total

242
Q

an increase in multiplicity increases

A

bond strength

243
Q

an increase in multiplicity decreases

A

bond length

244
Q

Summarizes common observations of different systems

A

Law

245
Q

Withstands the test of time

A

Law

246
Q

Attempts to explain WHY something happens

A

Theory

247
Q

Measure of the amount of matter in an object

A

Mass

248
Q

Force exerted by gravity on an object

A

Weight

249
Q

Agreement of a particular value with the TRUE value

A

Accuracy

250
Q

Agreement among several measurements of the same quantity

A

Precision

251
Q

Anything that occupies space and has mass

A

Matter

252
Q

Has visibly indistinguishable parts (solutions)

A

Homogenous mixture

253
Q

Has visibly distinguishable parts

A

Heterogenous

254
Q

Can be separated into pure substances by physical methods

A

Mixtures

255
Q

A substance with a constant composition that can be broken down into its elements via chemical processes

A

Compound

256
Q

Substance that cannot be broken down into simpler substances by physical or chemical means

A

Element

257
Q

Boiling or freezing water

A

Physical change

258
Q

Burning logs; chemical rxn bubbling

A

Chemical change

259
Q

Anything that can be classified as matter

Ex: compounds, elements

A

Substance

260
Q

The reactant that is used in an excess amount (more than is really needed) to make the theoretical yield of product.

A

Excess reagent