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Inorg Flashcards

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1
Q

Democritus

A

All matter is made up of indestructive units called atoms

Example sentence: Democritus was an ancient Greek philosopher who proposed the concept of atoms.

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2
Q

Max Planck

A

Proposed the idea of quantization

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3
Q

John Dalton 4

A

Chemical reactions involve reorganization of the atoms

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4
Q

John Dalton 1

A

Each element is made up of atom

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5
Q

John Dalton 2

A

Atoms of a given compound are identical

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6
Q

John Dalton 3

A

Compounds are formed when atoms combine with each other

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7
Q

Frederick Soddy

A

Discovered that there appeared to be more than one element at each position on the periodic table

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8
Q

Margaret Todd

A

Coined the term isotope

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9
Q

Dmitri Mendeleev

A

Created the periodic table based on the periodic functions of their atomic weight

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10
Q

Richard Abegg

A

Found that noble gases have stable electron configurations

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11
Q

James Clerk Maxwell

A

Proposed the theory of electromagnetism and made connection between light and electromagnetic waves

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12
Q

Albert Einstein

A

Created the theories of relativity and hypothesized about the particle nature of light

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13
Q

George Stoney

A

Proposed that electricity was made up of discrete negative particles called electrons

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14
Q

Hans Geiger

A

Invented a device that could detect alpha particles

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15
Q

Sir William Crooke

A

Demonstrated in his experiments that cathode rays have a negative charge

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16
Q

Robert Millikan

A

Determined the charge of an electron through his oil drop experiment

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17
Q

Eugene Goldstein

A

Used cathode ray tube to study canal rays which had electrical and magnetic properties opposite of an electron

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18
Q

Ernest Rutherford

A

Performed alpha particle experiment and established that the nucleus was very dense very small and positively charged

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19
Q

Wilhelm Roentgen

A

Discovered that certain chemicals glowed when exposed to cathode rays called X-rays

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20
Q

Henry Moseley

A

Discovered that the number of protons in an element determines its atomic number

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21
Q

Henri Becquerel

A

Discovered radiation by studying the effects of X-rays on photographic film

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22
Q

Neils Bohr

A

Developed Bohr atomic model with electrons travelling in orbits around the nucleus

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23
Q

Sir Joseph John Thompson

A

Used cathode ray tubes to determine the charge to mass ratio of an electron

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24
Q

Louis de Broglie

A

Proposed that electrons have a wave-particle duality

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25
Q

Ernest Rutherford 2

A

Discovered alpha beta and gamma rays in radiation

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26
Q

Erwin Schrödinger

A

Developed the Schrödinger equation which describes how the quantum state of a system changes with time

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27
Q

Pierre and Marie Curie

A

Theorized that radioactive particles cause atoms to break down releasing radiation that take form in energy and subatomic particles

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28
Q

Pierre and Marie Curie 2

A

Discovered the radioactive elements Polonium and Radium

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29
Q

Antoine Lavoisier

A

Father of Modern Chemistry

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30
Q

Antoine Lavoisier 1

A

Named oxygen and proved that water is a compound of hydrogen and oxygen

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31
Q

Antoine Lavoisier 2

A

Conservation of mass in a chemical reaction

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32
Q

Antoine Lavoisier 3

A

Introduced a new system of nomenclature where each substance was given a single name which described its composition

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33
Q

Amedeo Avogadro

A

Formulated the Avogadro’s law and Avogadro’s number 6.022x10^23

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34
Q

Jons Jakob Berzelius

A

Isolated new elements and developed a chemical notation system using letters and numbers

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35
Q

John Dalton 5

A

Created 36 chemical symbols

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36
Q

Dalton’s Billiard Ball Model

A

First to describe atoms in a modern scientific sense

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37
Q

Thomson’s Plum Pudding model

A

Showed the existence of protons and electrons

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38
Q

Rutherford’s nuclear model

A

Showed the nucleus

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39
Q

Bohr’s planetary model

A

Showed energy levels

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40
Q

Schrödinger’s electron cloud model

A

Showed subshells and shells are actually orbitals

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41
Q

Chadwick

A

Existence of neutrons

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42
Q

Mass number

A

Number of protons and neutrons in an atom

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43
Q

Atomic number

A

Number of protons in an atom

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44
Q

Atomic symbol

A

Abbreviation used to represent atom in chemical formulas

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45
Q

Francis William Aston

A

Discovered isotopes

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46
Q

Johann Dobereiner

A

Proposed the law of Triads where the middle element in certain triads had an atomic weight that was average of the other two members

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47
Q

John Newlands

A

Law of Octaves where every eighth element shared similar properties

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48
Q

Lothar Meyer

A

Studied the relationship of the atomic volume and relative atomic mass of 28 elements

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49
Q

Dmitri Mendeleev 2

A

Formulated Periodic Law and made a periodic table of 63 known elements where their properties are periodic functions of their atomic masses

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50
Q

Antonius van den Broek

A

First suggested that the number of charges in an element’s atomic nucleus is exactly equal to the element’s place on Mendeleev’s table

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51
Q

Henry Moseley

A

Discovered atomic number and its relationship between atomic mass

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52
Q

Glenn Seaborg

A

Discovery of 10 transuranium elements

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53
Q

Glenn Seaborg 1

A

Had an element named after him while he was still alive

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54
Q

Strontium fireworks

A

Red

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55
Q

Calcium fireworks

A

Orange

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56
Q

Sodium Fireworks

A

Yellow

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57
Q

Barium fireworks

A

Green

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58
Q

Copper fireworks

A

Blue

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59
Q

Copper and strontium fireworks

A

Purple

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60
Q

Iron fireworks

A

Gold/light yellow

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61
Q

Aluminum fireworks

A

Silver and white

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62
Q

Magnesium fireworks

A

White

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63
Q

Lithium fireworks

A

Red

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64
Q

Cesium flame color

A

Blue violet

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65
Q

Rubidium flame color

A

Red to violet

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66
Q

Glenn Seaborg

A

Had an element named after him while he was still alive

Example: Seaborgium (Sg)

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67
Q

IUPAC

A

International Union of Pure and Applied Chemistry

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68
Q

Alkali metals

A

Lustrous soft and highly reactive metals ready to form +1 cations and found naturally only in salts

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69
Q

Alkaline earth metals

A

React with water to form alkaline hydroxides readily lose valence to form +2 cations

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70
Q

Transition metals

A

Less reactive than group 1 and 2 metals have higher melting points and densities

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71
Q

Boron group or Icosagens

A

Have low melting points and poor hardness and react with oxygen to form oxides

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72
Q

Carbon group or Crystallogens

A

Has four valence electrons form hydrides with hydrogen tetrahalides with halogens and variety of oxides with oxygen

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73
Q

Nitrogen group or Pnictogens

A

Have five valence electrons all are solid except the first element

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74
Q

Oxygen group or Chalcogens

A

Have six valence electrons and electronegative nonmetals react with metals to form -2 ions

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75
Q

Halogens

A

All are reactive nonmetals have seven electrons and only group that contains solid liquid and gas (I and As are solids Br is liquid and F and Cl are gases)

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76
Q

Noble gases

A

Very low chemical reactivity and colorless gases but exhibit colors when ionized

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77
Q

6 commonly recognized Metalloids

A

B, Si, Ge, As, Sb, Te

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78
Q

Allotrope

A

One of two or more distinct forms of an element

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79
Q

Alfred Stock

A

Stock nomenclature where oxidation states are indicated in parentheses by Roman numerals

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80
Q

Cyanide formula

A

CN-

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81
Q

Cyanate formula

A

OCN-

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82
Q

Carbonate formula

A

CO3^2-

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83
Q

Bicarbonate

A

HCO3 -

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84
Q

Oxalate

A

C2O4 ^2-

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85
Q

Acetate

A

CH3COO-

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86
Q

Borate

A

BO3^3-

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87
Q

Arsenate

A

AsO4^3-

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88
Q

Silicate

A

SiO4^4-

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89
Q

Permanganate

A

MnO4 -

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90
Q

Nitrite

A

NO2 -

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91
Q

Nitrate

A

NO3 -

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92
Q

Hydroxide

A

OH-

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93
Q

Peroxide

A

O2^2-

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94
Q

Thiocyanate

A

SCN-

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95
Q

Sulfite

A

SO3^2-

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96
Q

Bisulfite

A

HSO3 -

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97
Q

Sulfate

A

SO4^2-

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98
Q

Bisulfate

A

HSO4 -

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99
Q

Thiosulfate

A

S2O3^2-

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100
Q

Phosphite

A

PO3^3-

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101
Q

Biphosphite

A

HPO3^2-

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102
Q

Dihydrogen phosphite

A

H2PO3 -

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103
Q

Phosphate

A

PO4^3-

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104
Q

Biphosphate

A

HPO4^2-

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105
Q

Hypochlorite

A

ClO-

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106
Q

Chlorite

A

ClO2 -

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107
Q

Chlorate

A

ClO3 -

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108
Q

Perchlorate

A

ClO4 -

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109
Q

Hypobromite

A

BrO-

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110
Q

Bromite

A

BrO2 -

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111
Q

Bromate

A

BrO3 -

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112
Q

Perbromate

A

BrO4 -

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113
Q

Hypoiodite

A

IO-

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114
Q

Iodite

A

IO2 -

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115
Q

Iodate

A

IO3 -

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116
Q

Periodate

A

IO4 -

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117
Q

Chromate

A

CrO4^2-

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118
Q

Dichromate

A

Cr2O7^2-

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119
Q

Hydrates

A

Compounds that have specific number of water molecules attached to them

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120
Q

Binary acids

A

Contains a hydrogen and an anion

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121
Q

Oxyacids

A

Contains a hydrogen and an oxyanion

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122
Q

Magnesia

A

MgO

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123
Q

Lime

A

CaO

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124
Q

Alumina

A

Al2O3

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125
Q

Silica

A

SiO2

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126
Q

Caustic soda

A

NaOH

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127
Q

Caustic potash

A

KOH

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128
Q

Milk of magnesia

A

Mg(OH)2

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129
Q

Slaked Lime

A

Ca(OH)2

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130
Q

Baking soda

A

NAHCO3

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131
Q

Soda ash

A

Na2CO3

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132
Q

Washing soda

A

Na2CO3 • 10H2O

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133
Q

Pearl ash

A

K2CO3

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134
Q

Magnesite

A

MgCO3

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135
Q

Calcite

A

CaCO3

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136
Q

Dolomite

A

CaMg(CO3)2

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137
Q

Siderite

A

FeCO3

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138
Q

Glauber’s salt

A

Na2SO4 • 10H2O

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139
Q

Epsom salt

A

MgSO4 • 7H2O

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140
Q

Plaster of Paris

A

CaSO4 • 1/2 H2O

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141
Q

Gypsum

A

CaSO4 • 2H2O

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142
Q

Oil of vitriol

A

H2SO4

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143
Q

Blue vitriol

A

CuSO4 • 5H2O

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144
Q

Green vitriol

A

FeSO4 • 7H2O

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145
Q

White vitriol

A

ZnSO4 •7H2O

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146
Q

Diborane

A

B2H6

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147
Q

Silane

A

SiH4

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148
Q

Phosphine

A

PH3

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149
Q

Hydrogen sulfide

A

H2S

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150
Q

Justus Von Liebig

A

Identified the first

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151
Q

Oil of vitriol

A

H2SO4

Sulfuric acid

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152
Q

Blue vitriol

A

CuSO4 • 5H2O

Copper(II) sulfate pentahydrate

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153
Q

Green vitriol

A

FeSO4 • 7H2O

Iron(II) sulfate heptahydrate

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154
Q

White vitriol

A

ZnSO4 • 7H2O

Zinc sulfate heptahydrate

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155
Q

Diborane

A

B2H6

Diborane is a colorless, highly reactive gas

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156
Q

Silane

A

SiH4

Silicon tetrahydride

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157
Q

Phosphine

A

PH3

Phosphine is a colorless, flammable gas

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158
Q

Hydrogen sulfide

A

H2S

Hydrogen sulfide is a colorless gas with a characteristic odor of rotten eggs

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159
Q

Justus Von Liebig

A

Identified the first example of isomerism and that nitrogen is an essential plant nutrient

Chemist and principal founder of organic chemistry

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160
Q

Friedrich Wöhler

A

Accidentally synthesized urea and co-discoverer of Be and Si

German chemist known for his discovery of the synthesis of urea

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161
Q

August Kekulé

A

Structure of benzene’s ring shaped structure

German chemist who proposed the structure of benzene

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162
Q

Kathleen Lonsdale

A

Used X-ray crystallography to prove the benzene ring’s structure

Irish crystallographer and first woman tenured professor at University College London

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163
Q

Linus Pauling

A

Known for his work on chemical bonding and proposed the Pauling electronegativity scale

American chemist and two-time Nobel Prize winner

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164
Q

Michael Faraday

A

Faraday’s constant 96485 C/mol

English scientist who contributed to the fields of electromagnetism and electrochemistry

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165
Q

Edward Frankland

A

Pioneers of organometallic chemistry and pioneered the concept of combining power or valence

English chemist known for his work on valence theory

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166
Q

Jacobus Henricus van’t Hoff

A

First winner of the Nobel prize in Chemistry and one of the founders of physical chemistry laid foundation for stereochemistry

Dutch physical chemist and first winner of the Nobel Prize in Chemistry

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167
Q

Gilbert Lewis

A

Lewis structure

American physical chemist known for his concept of electron pairs

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168
Q

Erick Hückel

A

Developed the Hückel method of approximate molecular orbital calculations on π electron systems

German physical chemist and physicist

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169
Q

Victor Grignard

A

Discovered Grignard reagent and reaction

French chemist and Nobel Prize winner

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170
Q

Emil Fischer

A

Discovered Fischer esterification developed Fischer projection and hypothesized lock and key mechanism of enzyme action

German chemist and Nobel Prize winner

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171
Q

Hybridization

A

Combination of two or more atomic orbitals to form the same number of hybrid orbitals each having the same shape and energy

Concept in chemistry to explain the geometry of molecules

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172
Q

As bond length increases

A

Bond strength decreases

Inverse relationship in chemical bonds

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173
Q

Nonpolar bond

A

Electronegative difference is less than 0.4

Type of covalent bond with equal sharing of electrons

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174
Q

Ionic bond

A

Electronegative difference is 2.0 or more

Type of bond formed between a metal and a non-metal

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175
Q

Polar covalent bonds

A

Electronegative difference between 0.4-1.7

Type of covalent bond with unequal sharing of electrons

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176
Q

Inductive effect

A

The pull of electron density through sigma bonds caused by electronegativity difference of atoms

Effect in organic chemistry that influences the distribution of electrons

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177
Q

Spirocyclic

A

Two rings share one atom, has the prefix spiro[x.y] where x is smaller

Type of bicyclic compound

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178
Q

Fused bicyclic

A

Two rings share two atoms in one bond, bicyclo[x.y.0] where x is bigger

Type of bicyclic compound with shared atoms

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179
Q

Bridged bicyclic

A

Two rings share three or more atoms, separating the two bridgehead atoms by a bridge containing at least one atom, bicyclo[x.y.z] where x is bigger and z is number of bridgehead atoms

Type of bicyclic compound with a bridge

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180
Q

Functional group priority

A

Alkyl halide < ether < alkane < alkyne < alkene < amine < alcohol < ketone < aldehyde < nitrile < amide < ester < carboxylic acid

Order of priority for functional groups in organic chemistry

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181
Q

Wavelength

A

Distance between identical points on consecutive waves

Physical property of a wave

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182
Q

Amplitude

A

Distance between origin and crest or trough

Measure of the height of a wave

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183
Q

Frequency

A

Number of waves that pass per unit time

Measure of the rate of wave oscillation

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184
Q

Speed

A

Wavelength times frequency

Relationship between wavelength and frequency in waves

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185
Q

Speed of light

A

3x10^8 m/s

Constant speed of light in a vacuum

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186
Q

Blackbody radiation

A

Relationship between an object’s temperature and the wavelength of electromagnetic radiation it emits

Thermal radiation from a perfect absorber and emitter of electromagnetic radiation

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187
Q

Planck’s equation

A

E = hv = (hc)/lambda

Equation describing the energy of a photon

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188
Q

Planck’s constant

A

h= 6.626x10^-34 J•s/particle

Physical constant used in quantum mechanics

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189
Q

Indium flame color

A

Blue

Color of flame produced by burning indium

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190
Q

Lead flame color

A

Light blue

Color of flame produced by burning lead

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191
Q

Arsenic flame color

A

Blue

Color of flame produced by burning arsenic

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192
Q

Sulfur flame color

A

Blue

Color of flame produced by burning sulfur

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193
Q

Radium flame color

A

Crimson red

Color of flame produced by burning radium

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194
Q

Antimony flame color

A

Pale green

Color of flame produced by burning antimony

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195
Q

Selenium flame color

A

Azure blue

Color of flame produced by burning selenium

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196
Q

Tin flame color

A

Blue-white

Color of flame produced by burning tin

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197
Q

Tantalum flame color

A

Blue

Color of flame produced by burning tantalum

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198
Q

Zinc flame color

A

Blue-green

Color of flame produced by burning zinc

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199
Q

Tungsten flame color

A

Green

Color of flame produced by burning tungsten

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200
Q

Yttrium flame color

A

Carmine crimson or scarlet red

Color of flame produced by burning yttrium

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201
Q

Zirconium flame color

A

Mid/dull red

Color of flame produced by burning zirconium

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202
Q

Photoelectric effect

A

Irradiating a metal surface causes ejection of electron

Phenomenon where light causes emission of electrons from a material

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203
Q

Work function

A

Minimum energy required to remove electrons from the metal surface

Energy required to remove an electron from a material

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204
Q

Threshold frequency

A

Minimum frequency multiplied by Planck’s constant to obtain work function

Frequency of light required to overcome the work function of a material

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205
Q

Principal quantum number

A

n main energy level and distance if electrons from nucleus

Quantum number in atomic theory indicating main energy levels

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206
Q

Azimuthal quantum number

A

l energy subshells and shape of orbitals

Quantum number indicating energy subshells in an atom

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207
Q

Magnetic quantum number

A

Number of orbitals in subshells and possible orientation of orbitals in space

Quantum number indicating orbital orientations

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208
Q

Spin quantum number

A

Movement of electron around its own axis clockwise and counterclockwise

Quantum number indicating electron spin

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209
Q

Aufbau Principle

A

Building up principle orbital with lower energy is filled up first

Principle in chemistry for filling electron orbitals

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210
Q

Madelung’s rule

A

Energy increases with increasing n + 1

Rule for determining electron configurations

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211
Q

Hund’s rule of Maximum Multiplicity

A

For degenerate orbitals electrons must first occupy them singly with parallel spin before filling with pairs

Rule for filling electron orbitals

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212
Q

Pauli’s exclusion principle

A

No two electrons can have the same set of quantum numbers

Principle in quantum mechanics

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213
Q

Chromium electron configuration

A

[Ar] 4s1 3d5

Electron configuration of Chromium

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214
Q

Molybdenum electron configuration

A

[Kr] 5s1 4d5

Electron configuration of Molybdenum

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215
Q

Copper electron configuration

A

[Ar] 4s1 3d10

Electron configuration of Copper

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216
Q

Silver electron configuration

A

[Kr]

Electron configuration of Silver

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217
Q

Hund’s rule of Maximum Multiplicity

A

For degenerate orbitals electrons must first occupy them singly with parallel spin before filling with pairs

Example: In the 2p subshell, each electron will first occupy a separate orbital before pairing up

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218
Q

Pauli’s exclusion principle

A

No two electrons can have the same set of quantum numbers

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219
Q

Chromium electron configuration

A

[Ar] 4s1 3d5

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220
Q

Molybdenum electron configuration

A

[Kr] 5s1 4d5

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221
Q

Copper electron configuration

A

[Ar] 4s1 3d10

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222
Q

Silver electron configuration

A

[Kr] 5s1 4d10

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223
Q

Slater’s rule

A

Used to calculate the shielding constant

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224
Q

Lanthanide contraction

A

Additional electrons do not add to the atomic size in the 5th and 6th period

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225
Q

Ionization energy

A

Energy required to remove an electron from a gaseous atom in its ground state

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226
Q

Electron affinity

A

Energy change associated with the addition of an electron to a gaseous atom in its ground state

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227
Q

Polarizability

A

Ability to be distorted by an electric field

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228
Q

Polarizability trend

A

Larger species = greater polarizability

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229
Q

AX2E1

A

Bent

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230
Q

AX4E1

A

Seesaw

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231
Q

AX3E2

A

T-shaped

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232
Q

AX5E1

A

Square pyramidal

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233
Q

AX4E2

A

Square Planar

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234
Q

Dipole moment

A

Quantitative measure of bond polarity

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235
Q

Dipole moment equation

A

μ= Q x r where Q is charge r is distance and μ is expressed in (D) Debye units

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236
Q

London dispersion or van der waals

A

Weak forces of attraction as a result of nonsymmetrical electron distribution that created a temporary dipole moment

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237
Q

Dipole-dipole forces

A

Occurs between compounds with permanent dipole moment

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238
Q

Hydrogen bonding

A

Only occurs when H is bonded to N O or F

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239
Q

Ion-ion forces

A

Between compounds with positive and negative charges

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240
Q

Fajan’s rules

A

Small highly charged cations have polarizing ability Large highly charged anions are easily polarized Cations that do not have a noble-gas electron configuration are easily polarized

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241
Q

Metallic bond

A

Attraction between electropositive atoms and delocalized electrons within a metal lattice

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242
Q

Covalent bond

A

Attraction resulting from the sharing of electrons of atoms

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243
Q

Arrhenius acid

A

Produces H+ in aqueous solution

244
Q

Arrhenius base

A

Produces OH- in aqueous solution

245
Q

Brønsted-Lowry acid

A

Proton donor

246
Q

Brønsted-Lowry base

A

Proton acceptor

247
Q

Lewis acid

A

Electron pair acceptor

248
Q

Lewis base

A

Electron pair donor

249
Q

Aqua acid

A

Acidic proton is on a water molecule coordinated to a central metal ion

250
Q

Hydroxoacid

A

Acidic proton is on a hydroxyl group without a neighboring oxo group

251
Q

Oxoacid

A

Acidic proton is on a hydroxyl group with an oxo group attached to the same atom

252
Q

HSAB classification

A

Hard acid bind to hard base and soft acid bind to soft base

253
Q

Hard acid

A

Smaller, high charge and highly polarizing

254
Q

Soft acid

A

Bigger, low charge, and low polarizing

255
Q

Hard base

A

Smaller high charge least polarizable

256
Q

Soft base

A

Big low charge highly polarizable

257
Q

Paramagnetic

A

Molecules with at least one unpaired electron

258
Q

Diamagnetic

A

Molecules with fully paired electrons

259
Q

Ferromagnetic

A

Permanent magnet

260
Q

Principal axis

A

The highest symmetry axis in a molecule

261
Q

Crystalline solids

A

Solids with highly regular arrangements of their components

262
Q

Amorphous solids

A

Solids with considerable disorder in their structures

263
Q

Coordination number

A

Number of nearest nearby atoms in a lattice

264
Q

Conductors

A

When valence band and conduction band overlap

265
Q

Insulators

A

A large band gap between the valence and conduction band which prevents the motion of electrons

266
Q

Semiconductors

A

The band gap is small enough that energy may be inputted to excite valence band electrons to the conduction band

267
Q

Intrinsic semiconductors

A

Elements that exhibit semi-conductive behavior at their pure state

268
Q

n-type dopants

A

Group 15 elements are capable of adding an electron relative to the host semiconductor

269
Q

p-type dopants

A

Group 13 elements provide a positive hole for increasing conductance

270
Q

Alloy

A

A mixture of metals or a mixture of a metal and another elements

271
Q

Substitutional alloy

A

Some of the host metal atoms are replaced by other metal atoms of similar size

272
Q

Interstitial alloy

A

Formed when some of the interstices or holes in the closest packed metal structure are occupied by small atoms

273
Q

Molecular solids

A

Has discrete molecular units at each lattice position

274
Q

Atomic solids

A

Have atoms occupying the lattice points

275
Q

Ionic solids

A

Stable high-melting substance held together by strong electrostatic forces that exist between oppositely charged ions

276
Q

Lattice energy

A

The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid

277
Q

Coordination complex

A

Transition metals (Lewis acid) accept electrons from ligands (Lewis base)

278
Q

Primary valence

A

Oxidation state of metal

279
Q

Secondary valence

A

Coordination number

280
Q

Denticity

A

Number of times a ligands binds to a metal through donor atoms

281
Q

Linkage isomer

A

Only occurs with ambidentate ligands where the same ligand may link through different atoms

282
Q

Ionization isomer

A

Occurs

283
Q

Energy in ionic solid formation

A

the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid

284
Q

Secondary valence

A

coordination number

285
Q

Denticity

A

number of times a ligands binds to a metal through donor atoms

286
Q

Linkage isomer

A

only occurs with ambidentate ligands where the same ligand may link through different atoms

287
Q

Ionization isomer

A

occurs when a ligand and a counterion in one compound exchange places

288
Q

Hydration isomer

A

Ionization isomer but is specific for water

289
Q

Coordination isomer

A

arises when there are different complex ions that can form the same molecular formula, or simply two metals exchanging ligands

290
Q

Polymerization isomer

A

denote complexes which have the same empirical formula but different molar masses

291
Q

Stereoisomers

A

isomers differing only in their three-dimensional arrangement

292
Q

Optical isomers

A

non-superimposable mirror images of each other and can rotate plane-polarized light

293
Q

Crystal field theory

A

electrostatic model that involves interactions between the electrons from the ligands and the electrons in the metal d-orbitals

294
Q

Ligand field splitting parameter, ∆0

A

The separation of the two sets of orbitals

The magnitude of the ligand field splitting parameter depends on the charge on the metal ion principal quantum number and the nature of the ligand

295
Q

Effect of principal quantum number

A

shorter metal-ligand distance, stronger orbital-ligand interaction, ∆0 increases

296
Q

Spectrochemical series of ligands

A

I- < Br- < SCN- < Cl- < NO3- < F- < OH- < H2O < NCS- < NH3 < en < phen < NO2- < PPh3 < CN- < CO

297
Q

Color absorbed and seen relationship for red

A

if color absorbed is red, color seen is green

298
Q

Color absorbed and seen relationship for yellow

A

if color absorbed is yellow, color seen is violet

299
Q

Polyprotic acids

A

acids that yield more than one H3O+ ion

300
Q

Buffer

A

combination of a weak acid/base and its conjugate acid/base in equal concentrations

301
Q

GEROA

A

Gain electrons reduction oxidizing agent

302
Q

LEORA

A

lose electrons oxidation reducing agent

303
Q

Oxidation number

A

total number of electrons removed/added to an element

304
Q

Democritus

A

All matter is made up of indestructive units called atoms

Example sentence: Democritus was an ancient Greek philosopher who proposed the concept of atoms.

305
Q

Max Planck

A

Proposed the idea of quantization

306
Q

John Dalton 4

A

Chemical reactions involve reorganization of the atoms

307
Q

John Dalton 1

A

Each element is made up of atom

308
Q

John Dalton 2

A

Atoms of a given compound are identical

309
Q

John Dalton 3

A

Compounds are formed when atoms combine with each other

310
Q

Frederick Soddy

A

Discovered that there appeared to be more than one element at each position on the periodic table

311
Q

Margaret Todd

A

Coined the term isotope

312
Q

Dmitri Mendeleev

A

Created the periodic table based on the periodic functions of their atomic weight

313
Q

Richard Abegg

A

Found that noble gases have stable electron configurations

314
Q

James Clerk Maxwell

A

Proposed the theory of electromagnetism and made connection between light and electromagnetic waves

315
Q

Albert Einstein

A

Created the theories of relativity and hypothesized about the particle nature of light

316
Q

George Stoney

A

Proposed that electricity was made up of discrete negative particles called electrons

317
Q

Hans Geiger

A

Invented a device that could detect alpha particles

318
Q

Sir William Crooke

A

Demonstrated in his experiments that cathode rays have a negative charge

319
Q

Robert Millikan

A

Determined the charge of an electron through his oil drop experiment

320
Q

Eugene Goldstein

A

Used cathode ray tube to study canal rays which had electrical and magnetic properties opposite of an electron

321
Q

Ernest Rutherford

A

Performed alpha particle experiment and established that the nucleus was very dense very small and positively charged

322
Q

Wilhelm Roentgen

A

Discovered that certain chemicals glowed when exposed to cathode rays called X-rays

323
Q

Henry Moseley

A

Discovered that the number of protons in an element determines its atomic number

324
Q

Henri Becquerel

A

Discovered radiation by studying the effects of X-rays on photographic film

325
Q

Neils Bohr

A

Developed Bohr atomic model with electrons travelling in orbits around the nucleus

326
Q

Sir Joseph John Thompson

A

Used cathode ray tubes to determine the charge to mass ratio of an electron

327
Q

Louis de Broglie

A

Proposed that electrons have a wave-particle duality

328
Q

Ernest Rutherford 2

A

Discovered alpha beta and gamma rays in radiation

329
Q

Erwin Schrödinger

A

Developed the Schrödinger equation which describes how the quantum state of a system changes with time

330
Q

Pierre and Marie Curie

A

Theorized that radioactive particles cause atoms to break down releasing radiation that take form in energy and subatomic particles

331
Q

Pierre and Marie Curie 2

A

Discovered the radioactive elements Polonium and Radium

332
Q

Antoine Lavoisier

A

Father of Modern Chemistry

333
Q

Antoine Lavoisier 1

A

Named oxygen and proved that water is a compound of hydrogen and oxygen

334
Q

Antoine Lavoisier 2

A

Conservation of mass in a chemical reaction

335
Q

Antoine Lavoisier 3

A

Introduced a new system of nomenclature where each substance was given a single name which described its composition

336
Q

Amedeo Avogadro

A

Formulated the Avogadro’s law and Avogadro’s number 6.022x10^23

337
Q

Jons Jakob Berzelius

A

Isolated new elements and developed a chemical notation system using letters and numbers

338
Q

John Dalton 5

A

Created 36 chemical symbols

339
Q

Dalton’s Billiard Ball Model

A

First to describe atoms in a modern scientific sense

340
Q

Thomson’s Plum Pudding model

A

Showed the existence of protons and electrons

341
Q

Rutherford’s nuclear model

A

Showed the nucleus

342
Q

Bohr’s planetary model

A

Showed energy levels

343
Q

Schrödinger’s electron cloud model

A

Showed subshells and shells are actually orbitals

344
Q

Chadwick

A

Existence of neutrons

345
Q

Mass number

A

Number of protons and neutrons in an atom

346
Q

Atomic number

A

Number of protons in an atom

347
Q

Atomic symbol

A

Abbreviation used to represent atom in chemical formulas

348
Q

Francis William Aston

A

Discovered isotopes

349
Q

Johann Dobereiner

A

Proposed the law of Triads where the middle element in certain triads had an atomic weight that was average of the other two members

350
Q

John Newlands

A

Law of Octaves where every eighth element shared similar properties

351
Q

Lothar Meyer

A

Studied the relationship of the atomic volume and relative atomic mass of 28 elements

352
Q

Dmitri Mendeleev 2

A

Formulated Periodic Law and made a periodic table of 63 known elements where their properties are periodic functions of their atomic masses

353
Q

Antonius van den Broek

A

First suggested that the number of charges in an element’s atomic nucleus is exactly equal to the element’s place on Mendeleev’s table

354
Q

Henry Moseley

A

Discovered atomic number and its relationship between atomic mass

355
Q

Glenn Seaborg

A

Discovery of 10 transuranium elements

356
Q

Glenn Seaborg 1

A

Had an element named after him while he was still alive

357
Q

Strontium fireworks

A

Red

358
Q

Calcium fireworks

A

Orange

359
Q

Sodium Fireworks

A

Yellow

360
Q

Barium fireworks

A

Green

361
Q

Copper fireworks

A

Blue

362
Q

Copper and strontium fireworks

A

Purple

363
Q

Iron fireworks

A

Gold/light yellow

364
Q

Aluminum fireworks

A

Silver and white

365
Q

Magnesium fireworks

A

White

366
Q

Lithium fireworks

A

Red

367
Q

Cesium flame color

A

Blue violet

368
Q

Rubidium flame color

A

Red to violet

369
Q

Glenn Seaborg

A

Had an element named after him while he was still alive

Example: Seaborgium (Sg)

370
Q

IUPAC

A

International Union of Pure and Applied Chemistry

371
Q

Alkali metals

A

Lustrous soft and highly reactive metals ready to form +1 cations and found naturally only in salts

372
Q

Alkaline earth metals

A

React with water to form alkaline hydroxides readily lose valence to form +2 cations

373
Q

Transition metals

A

Less reactive than group 1 and 2 metals have higher melting points and densities

374
Q

Boron group or Icosagens

A

Have low melting points and poor hardness and react with oxygen to form oxides

375
Q

Carbon group or Crystallogens

A

Has four valence electrons form hydrides with hydrogen tetrahalides with halogens and variety of oxides with oxygen

376
Q

Nitrogen group or Pnictogens

A

Have five valence electrons all are solid except the first element

377
Q

Oxygen group or Chalcogens

A

Have six valence electrons and electronegative nonmetals react with metals to form -2 ions

378
Q

Halogens

A

All are reactive nonmetals have seven electrons and only group that contains solid liquid and gas (I and As are solids Br is liquid and F and Cl are gases)

379
Q

Noble gases

A

Very low chemical reactivity and colorless gases but exhibit colors when ionized

380
Q

6 commonly recognized Metalloids

A

B, Si, Ge, As, Sb, Te

381
Q

Allotrope

A

One of two or more distinct forms of an element

382
Q

Alfred Stock

A

Stock nomenclature where oxidation states are indicated in parentheses by Roman numerals

383
Q

Cyanide formula

A

CN-

384
Q

Cyanate formula

A

OCN-

385
Q

Carbonate formula

A

CO3^2-

386
Q

Bicarbonate

A

HCO3 -

387
Q

Oxalate

A

C2O4 ^2-

388
Q

Acetate

A

CH3COO-

389
Q

Borate

A

BO3^3-

390
Q

Arsenate

A

AsO4^3-

391
Q

Silicate

A

SiO4^4-

392
Q

Permanganate

A

MnO4 -

393
Q

Nitrite

A

NO2 -

394
Q

Nitrate

A

NO3 -

395
Q

Hydroxide

A

OH-

396
Q

Peroxide

A

O2^2-

397
Q

Thiocyanate

A

SCN-

398
Q

Sulfite

A

SO3^2-

399
Q

Bisulfite

A

HSO3 -

400
Q

Sulfate

A

SO4^2-

401
Q

Bisulfate

A

HSO4 -

402
Q

Thiosulfate

A

S2O3^2-

403
Q

Phosphite

A

PO3^3-

404
Q

Biphosphite

A

HPO3^2-

405
Q

Dihydrogen phosphite

A

H2PO3 -

406
Q

Phosphate

A

PO4^3-

407
Q

Biphosphate

A

HPO4^2-

408
Q

Hypochlorite

A

ClO-

409
Q

Chlorite

A

ClO2 -

410
Q

Chlorate

A

ClO3 -

411
Q

Perchlorate

A

ClO4 -

412
Q

Hypobromite

A

BrO-

413
Q

Bromite

A

BrO2 -

414
Q

Bromate

A

BrO3 -

415
Q

Perbromate

A

BrO4 -

416
Q

Hypoiodite

A

IO-

417
Q

Iodite

A

IO2 -

418
Q

Iodate

A

IO3 -

419
Q

Periodate

A

IO4 -

420
Q

Chromate

A

CrO4^2-

421
Q

Dichromate

A

Cr2O7^2-

422
Q

Hydrates

A

Compounds that have specific number of water molecules attached to them

423
Q

Binary acids

A

Contains a hydrogen and an anion

424
Q

Oxyacids

A

Contains a hydrogen and an oxyanion

425
Q

Magnesia

A

MgO

426
Q

Lime

A

CaO

427
Q

Alumina

A

Al2O3

428
Q

Silica

A

SiO2

429
Q

Caustic soda

A

NaOH

430
Q

Caustic potash

A

KOH

431
Q

Milk of magnesia

A

Mg(OH)2

432
Q

Slaked Lime

A

Ca(OH)2

433
Q

Baking soda

A

NAHCO3

434
Q

Soda ash

A

Na2CO3

435
Q

Washing soda

A

Na2CO3 • 10H2O

436
Q

Pearl ash

A

K2CO3

437
Q

Magnesite

A

MgCO3

438
Q

Calcite

A

CaCO3

439
Q

Dolomite

A

CaMg(CO3)2

440
Q

Siderite

A

FeCO3

441
Q

Glauber’s salt

A

Na2SO4 • 10H2O

442
Q

Epsom salt

A

MgSO4 • 7H2O

443
Q

Plaster of Paris

A

CaSO4 • 1/2 H2O

444
Q

Gypsum

A

CaSO4 • 2H2O

445
Q

Oil of vitriol

A

H2SO4

446
Q

Blue vitriol

A

CuSO4 • 5H2O

447
Q

Green vitriol

A

FeSO4 • 7H2O

448
Q

White vitriol

A

ZnSO4 •7H2O

449
Q

Diborane

A

B2H6

450
Q

Silane

A

SiH4

451
Q

Phosphine

A

PH3

452
Q

Hydrogen sulfide

A

H2S

453
Q

Justus Von Liebig

A

Identified the first

454
Q

Oil of vitriol

A

H2SO4

Sulfuric acid

455
Q

Blue vitriol

A

CuSO4 • 5H2O

Copper(II) sulfate pentahydrate

456
Q

Green vitriol

A

FeSO4 • 7H2O

Iron(II) sulfate heptahydrate

457
Q

White vitriol

A

ZnSO4 • 7H2O

Zinc sulfate heptahydrate

458
Q

Diborane

A

B2H6

Diborane is a colorless, highly reactive gas

459
Q

Silane

A

SiH4

Silicon tetrahydride

460
Q

Phosphine

A

PH3

Phosphine is a colorless, flammable gas

461
Q

Hydrogen sulfide

A

H2S

Hydrogen sulfide is a colorless gas with a characteristic odor of rotten eggs

462
Q

Justus Von Liebig

A

Identified the first example of isomerism and that nitrogen is an essential plant nutrient

Chemist and principal founder of organic chemistry

463
Q

Friedrich Wöhler

A

Accidentally synthesized urea and co-discoverer of Be and Si

German chemist known for his discovery of the synthesis of urea

464
Q

August Kekulé

A

Structure of benzene’s ring shaped structure

German chemist who proposed the structure of benzene

465
Q

Kathleen Lonsdale

A

Used X-ray crystallography to prove the benzene ring’s structure

Irish crystallographer and first woman tenured professor at University College London

466
Q

Linus Pauling

A

Known for his work on chemical bonding and proposed the Pauling electronegativity scale

American chemist and two-time Nobel Prize winner

467
Q

Michael Faraday

A

Faraday’s constant 96485 C/mol

English scientist who contributed to the fields of electromagnetism and electrochemistry

468
Q

Edward Frankland

A

Pioneers of organometallic chemistry and pioneered the concept of combining power or valence

English chemist known for his work on valence theory

469
Q

Jacobus Henricus van’t Hoff

A

First winner of the Nobel prize in Chemistry and one of the founders of physical chemistry laid foundation for stereochemistry

Dutch physical chemist and first winner of the Nobel Prize in Chemistry

470
Q

Gilbert Lewis

A

Lewis structure

American physical chemist known for his concept of electron pairs

471
Q

Erick Hückel

A

Developed the Hückel method of approximate molecular orbital calculations on π electron systems

German physical chemist and physicist

472
Q

Victor Grignard

A

Discovered Grignard reagent and reaction

French chemist and Nobel Prize winner

473
Q

Emil Fischer

A

Discovered Fischer esterification developed Fischer projection and hypothesized lock and key mechanism of enzyme action

German chemist and Nobel Prize winner

474
Q

Hybridization

A

Combination of two or more atomic orbitals to form the same number of hybrid orbitals each having the same shape and energy

Concept in chemistry to explain the geometry of molecules

475
Q

As bond length increases

A

Bond strength decreases

Inverse relationship in chemical bonds

476
Q

Nonpolar bond

A

Electronegative difference is less than 0.4

Type of covalent bond with equal sharing of electrons

477
Q

Ionic bond

A

Electronegative difference is 2.0 or more

Type of bond formed between a metal and a non-metal

478
Q

Polar covalent bonds

A

Electronegative difference between 0.4-1.7

Type of covalent bond with unequal sharing of electrons

479
Q

Inductive effect

A

The pull of electron density through sigma bonds caused by electronegativity difference of atoms

Effect in organic chemistry that influences the distribution of electrons

480
Q

Spirocyclic

A

Two rings share one atom, has the prefix spiro[x.y] where x is smaller

Type of bicyclic compound

481
Q

Fused bicyclic

A

Two rings share two atoms in one bond, bicyclo[x.y.0] where x is bigger

Type of bicyclic compound with shared atoms

482
Q

Bridged bicyclic

A

Two rings share three or more atoms, separating the two bridgehead atoms by a bridge containing at least one atom, bicyclo[x.y.z] where x is bigger and z is number of bridgehead atoms

Type of bicyclic compound with a bridge

483
Q

Functional group priority

A

Alkyl halide < ether < alkane < alkyne < alkene < amine < alcohol < ketone < aldehyde < nitrile < amide < ester < carboxylic acid

Order of priority for functional groups in organic chemistry

484
Q

Wavelength

A

Distance between identical points on consecutive waves

Physical property of a wave

485
Q

Amplitude

A

Distance between origin and crest or trough

Measure of the height of a wave

486
Q

Frequency

A

Number of waves that pass per unit time

Measure of the rate of wave oscillation

487
Q

Speed

A

Wavelength times frequency

Relationship between wavelength and frequency in waves

488
Q

Speed of light

A

3x10^8 m/s

Constant speed of light in a vacuum

489
Q

Blackbody radiation

A

Relationship between an object’s temperature and the wavelength of electromagnetic radiation it emits

Thermal radiation from a perfect absorber and emitter of electromagnetic radiation

490
Q

Planck’s equation

A

E = hv = (hc)/lambda

Equation describing the energy of a photon

491
Q

Planck’s constant

A

h= 6.626x10^-34 J•s/particle

Physical constant used in quantum mechanics

492
Q

Indium flame color

A

Blue

Color of flame produced by burning indium

493
Q

Lead flame color

A

Light blue

Color of flame produced by burning lead

494
Q

Arsenic flame color

A

Blue

Color of flame produced by burning arsenic

495
Q

Sulfur flame color

A

Blue

Color of flame produced by burning sulfur

496
Q

Radium flame color

A

Crimson red

Color of flame produced by burning radium

497
Q

Antimony flame color

A

Pale green

Color of flame produced by burning antimony

498
Q

Selenium flame color

A

Azure blue

Color of flame produced by burning selenium

499
Q

Tin flame color

A

Blue-white

Color of flame produced by burning tin

500
Q

Tantalum flame color

A

Blue

Color of flame produced by burning tantalum

501
Q

Zinc flame color

A

Blue-green

Color of flame produced by burning zinc

502
Q

Tungsten flame color

A

Green

Color of flame produced by burning tungsten

503
Q

Yttrium flame color

A

Carmine crimson or scarlet red

Color of flame produced by burning yttrium

504
Q

Zirconium flame color

A

Mid/dull red

Color of flame produced by burning zirconium

505
Q

Photoelectric effect

A

Irradiating a metal surface causes ejection of electron

Phenomenon where light causes emission of electrons from a material

506
Q

Work function

A

Minimum energy required to remove electrons from the metal surface

Energy required to remove an electron from a material

507
Q

Threshold frequency

A

Minimum frequency multiplied by Planck’s constant to obtain work function

Frequency of light required to overcome the work function of a material

508
Q

Principal quantum number

A

n main energy level and distance if electrons from nucleus

Quantum number in atomic theory indicating main energy levels

509
Q

Azimuthal quantum number

A

l energy subshells and shape of orbitals

Quantum number indicating energy subshells in an atom

510
Q

Magnetic quantum number

A

Number of orbitals in subshells and possible orientation of orbitals in space

Quantum number indicating orbital orientations

511
Q

Spin quantum number

A

Movement of electron around its own axis clockwise and counterclockwise

Quantum number indicating electron spin

512
Q

Aufbau Principle

A

Building up principle orbital with lower energy is filled up first

Principle in chemistry for filling electron orbitals

513
Q

Madelung’s rule

A

Energy increases with increasing n + 1

Rule for determining electron configurations

514
Q

Hund’s rule of Maximum Multiplicity

A

For degenerate orbitals electrons must first occupy them singly with parallel spin before filling with pairs

Rule for filling electron orbitals

515
Q

Pauli’s exclusion principle

A

No two electrons can have the same set of quantum numbers

Principle in quantum mechanics

516
Q

Chromium electron configuration

A

[Ar] 4s1 3d5

Electron configuration of Chromium

517
Q

Molybdenum electron configuration

A

[Kr] 5s1 4d5

Electron configuration of Molybdenum

518
Q

Copper electron configuration

A

[Ar] 4s1 3d10

Electron configuration of Copper

519
Q

Silver electron configuration

A

[Kr]

Electron configuration of Silver

520
Q

Hund’s rule of Maximum Multiplicity

A

For degenerate orbitals electrons must first occupy them singly with parallel spin before filling with pairs

Example: In the 2p subshell, each electron will first occupy a separate orbital before pairing up

521
Q

Pauli’s exclusion principle

A

No two electrons can have the same set of quantum numbers

522
Q

Chromium electron configuration

A

[Ar] 4s1 3d5

523
Q

Molybdenum electron configuration

A

[Kr] 5s1 4d5

524
Q

Copper electron configuration

A

[Ar] 4s1 3d10

525
Q

Silver electron configuration

A

[Kr] 5s1 4d10

526
Q

Slater’s rule

A

Used to calculate the shielding constant

527
Q

Lanthanide contraction

A

Additional electrons do not add to the atomic size in the 5th and 6th period

528
Q

Ionization energy

A

Energy required to remove an electron from a gaseous atom in its ground state

529
Q

Electron affinity

A

Energy change associated with the addition of an electron to a gaseous atom in its ground state

530
Q

Polarizability

A

Ability to be distorted by an electric field

531
Q

Polarizability trend

A

Larger species = greater polarizability

532
Q

AX2E1

A

Bent

533
Q

AX4E1

A

Seesaw

534
Q

AX3E2

A

T-shaped

535
Q

AX5E1

A

Square pyramidal

536
Q

AX4E2

A

Square Planar

537
Q

Dipole moment

A

Quantitative measure of bond polarity

538
Q

Dipole moment equation

A

μ= Q x r where Q is charge r is distance and μ is expressed in (D) Debye units

539
Q

London dispersion or van der waals

A

Weak forces of attraction as a result of nonsymmetrical electron distribution that created a temporary dipole moment

540
Q

Dipole-dipole forces

A

Occurs between compounds with permanent dipole moment

541
Q

Hydrogen bonding

A

Only occurs when H is bonded to N O or F

542
Q

Ion-ion forces

A

Between compounds with positive and negative charges

543
Q

Fajan’s rules

A

Small highly charged cations have polarizing ability Large highly charged anions are easily polarized Cations that do not have a noble-gas electron configuration are easily polarized

544
Q

Metallic bond

A

Attraction between electropositive atoms and delocalized electrons within a metal lattice

545
Q

Covalent bond

A

Attraction resulting from the sharing of electrons of atoms

546
Q

Arrhenius acid

A

Produces H+ in aqueous solution

547
Q

Arrhenius base

A

Produces OH- in aqueous solution

548
Q

Brønsted-Lowry acid

A

Proton donor

549
Q

Brønsted-Lowry base

A

Proton acceptor

550
Q

Lewis acid

A

Electron pair acceptor

551
Q

Lewis base

A

Electron pair donor

552
Q

Aqua acid

A

Acidic proton is on a water molecule coordinated to a central metal ion

553
Q

Hydroxoacid

A

Acidic proton is on a hydroxyl group without a neighboring oxo group

554
Q

Oxoacid

A

Acidic proton is on a hydroxyl group with an oxo group attached to the same atom

555
Q

HSAB classification

A

Hard acid bind to hard base and soft acid bind to soft base

556
Q

Hard acid

A

Smaller, high charge and highly polarizing

557
Q

Soft acid

A

Bigger, low charge, and low polarizing

558
Q

Hard base

A

Smaller high charge least polarizable

559
Q

Soft base

A

Big low charge highly polarizable

560
Q

Paramagnetic

A

Molecules with at least one unpaired electron

561
Q

Diamagnetic

A

Molecules with fully paired electrons

562
Q

Ferromagnetic

A

Permanent magnet

563
Q

Principal axis

A

The highest symmetry axis in a molecule

564
Q

Crystalline solids

A

Solids with highly regular arrangements of their components

565
Q

Amorphous solids

A

Solids with considerable disorder in their structures

566
Q

Coordination number

A

Number of nearest nearby atoms in a lattice

567
Q

Conductors

A

When valence band and conduction band overlap

568
Q

Insulators

A

A large band gap between the valence and conduction band which prevents the motion of electrons

569
Q

Semiconductors

A

The band gap is small enough that energy may be inputted to excite valence band electrons to the conduction band

570
Q

Intrinsic semiconductors

A

Elements that exhibit semi-conductive behavior at their pure state

571
Q

n-type dopants

A

Group 15 elements are capable of adding an electron relative to the host semiconductor

572
Q

p-type dopants

A

Group 13 elements provide a positive hole for increasing conductance

573
Q

Alloy

A

A mixture of metals or a mixture of a metal and another elements

574
Q

Substitutional alloy

A

Some of the host metal atoms are replaced by other metal atoms of similar size

575
Q

Interstitial alloy

A

Formed when some of the interstices or holes in the closest packed metal structure are occupied by small atoms

576
Q

Molecular solids

A

Has discrete molecular units at each lattice position

577
Q

Atomic solids

A

Have atoms occupying the lattice points

578
Q

Ionic solids

A

Stable high-melting substance held together by strong electrostatic forces that exist between oppositely charged ions

579
Q

Lattice energy

A

The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid

580
Q

Coordination complex

A

Transition metals (Lewis acid) accept electrons from ligands (Lewis base)

581
Q

Primary valence

A

Oxidation state of metal

582
Q

Secondary valence

A

Coordination number

583
Q

Denticity

A

Number of times a ligands binds to a metal through donor atoms

584
Q

Linkage isomer

A

Only occurs with ambidentate ligands where the same ligand may link through different atoms

585
Q

Ionization isomer

A

Occurs

586
Q

Energy in ionic solid formation

A

the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid

587
Q

Secondary valence

A

coordination number

588
Q

Denticity

A

number of times a ligands binds to a metal through donor atoms

589
Q

Linkage isomer

A

only occurs with ambidentate ligands where the same ligand may link through different atoms

590
Q

Ionization isomer

A

occurs when a ligand and a counterion in one compound exchange places

591
Q

Hydration isomer

A

Ionization isomer but is specific for water

592
Q

Coordination isomer

A

arises when there are different complex ions that can form the same molecular formula, or simply two metals exchanging ligands

593
Q

Polymerization isomer

A

denote complexes which have the same empirical formula but different molar masses

594
Q

Stereoisomers

A

isomers differing only in their three-dimensional arrangement

595
Q

Optical isomers

A

non-superimposable mirror images of each other and can rotate plane-polarized light

596
Q

Crystal field theory

A

electrostatic model that involves interactions between the electrons from the ligands and the electrons in the metal d-orbitals

597
Q

Ligand field splitting parameter, ∆0

A

The separation of the two sets of orbitals

The magnitude of the ligand field splitting parameter depends on the charge on the metal ion principal quantum number and the nature of the ligand

598
Q

Effect of principal quantum number

A

shorter metal-ligand distance, stronger orbital-ligand interaction, ∆0 increases

599
Q

Spectrochemical series of ligands

A

I- < Br- < SCN- < Cl- < NO3- < F- < OH- < H2O < NCS- < NH3 < en < phen < NO2- < PPh3 < CN- < CO

600
Q

Color absorbed and seen relationship for red

A

if color absorbed is red, color seen is green

601
Q

Color absorbed and seen relationship for yellow

A

if color absorbed is yellow, color seen is violet

602
Q

Polyprotic acids

A

acids that yield more than one H3O+ ion

603
Q

Buffer

A

combination of a weak acid/base and its conjugate acid/base in equal concentrations

604
Q

GEROA

A

Gain electrons reduction oxidizing agent

605
Q

LEORA

A

lose electrons oxidation reducing agent

606
Q

Oxidation number

A

total number of electrons removed/added to an element

Chemistry (1 decks)

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