CHEM 1106

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Chemistry
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Introductory chemistry
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1

Giga (G)

109

2

Mega (M)

106

3

Kilo (k)

103

4

Deci (d)

10-1

5

Centi (c)

10-2

6

Milli (m)

10-3

7

Micro (μ)

10-6

8

Nano (n)

10-9

9

Pico (p)

10-12

10

Femto (f)

10-15

11

Density equation

Mass (m)
Volume (v)

12

Fahrenheit to Celsius equation

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13

Celsius to Fahrenheit equation

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14

Celsius to Kelvin equation

x + 273

15

Another name for a homogeneous mix of metals

Alloy or solution

16

Definition of homogeneous

Same appearance and texture throughout

17

Definition of heterogeneous

Variable in texture and appearance

18

Definition of physical change

Alters the appearance of a material only

19

Definition of chemical change

Creates a new product or substance

20

Definition of a pure element

One type of element

21

Definition of a pure compound

Two or more elements in a fixed ratio

22

Properties of gas particles

Particles in rapid motion experiencing many collisions.

23

Properties of liquid particles

Particles close together without a regular arrangement.

24

Properties of solid particles

Particles tight together with regular arrangement.

25

Separation techniques for gases and liquids

Distillation, chromatography, membrane filteration

26

Separation techniques for solids

Sublimation, paper/gel chromatography, solubility diferences

27

Dalton theorized

Each element s made of atoms

28

Lavoisier stated

atoms are neither created or destroyed (Law of conservation of mass)

29

Proust stated

New compounds can be formed and compounds consist of specific types and numbers of atoms (Law of constant composition)

30

Dalton stated

Manipulating reaction parameters can create different compounds from the same atoms (Law of multiple proportions)

31

Discovered by J.J. Thompson

Electrons (e-) found using magnets and cathode-ray tubes

32

Discovered by R. Millikan

Electron (e-) mass found by exposing oil to e- and observing the variable rate at which they fell (or rose).

33

Discovered by E. Rutherford

Protons found by exposing gold foil to alpha-decay and measuring the angles of deflection.

34

Charge of proton

(+) positive

35

Discovered by J. Chadwick

Neutrons found by exposing beryllium foil to alpha-decay which knocked out a neutron that collided with a wax layer and knocked out a detectable proton.

36

Definition of an isotoope

The same element with differing numbers of neutrons.

37

Periodic orientation of groups (families)

Vertical

38

Periodic orientation of periods

Horizontal

39

The metalloids

Boron, silicon, geranium, arsenic, antimony, tellurium, and astatine

40

Have No Fear Of Ice Cold Beer

Mnemonic used to remember diatomic elements H2, N2, F2, O2, I2, Cl2, Br2

41

Properties of metals

Ductility, malleability, conductivity, luster. Forms Ionic bonds with non-metals (electrolytes).

42

Properties of non-metals

Naturally occur as solid, liquid, and gas. Form molecular compounds with non-metals.

43

Example of isotopes

13C, 14C, 12C

44

Definition of beta (β)-minus decay

The decay of a neutron into a proton. Releases an electron (e-).

45

Definition of beta (β)-plus decay

The decay of a proton into a neutron. Releases a positron (e+)

46

Example of beta (β)-minus decay

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47

Example of beta ( β)-plus decay

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48

( β)-decay tissue penetration

5mm

49

Released in alpha (α)-decay

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50

Example of alpha (α)-decay

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51

Tissue penetration alpha (α)-decay

Approximately 3cm

52

Result from decay of meta-stable nuclei or when positrons (e+) collide with electrons e-

Gamma (γ)-radiation

53

Example decay of meta-stable nuclei

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54

Tissue penetration of gamma (γ)-radiation

50cm

55

Definition of half-life

Amount of time it takes 1/2 of nuclei to decay

56

Symbol for half-life

t1/2

57

Characteristics of ionic (metal to non-metal) compounds

Brittle and conductive when aqueous

58

Characteristics of molecular (non-metal to non-metal) compounds

hard or soft, not conductive

59

How to find an elements charge, not including transition metals.

+1 for each group moving right until carbon, then -1 until helium

60

Hydronium

H₃O⁺

61

Ammonium

NH₄⁺

62

Hydroxide

OH⁻

63

Cyanide

CN⁻

64

Nitrate

NO₃⁻

65

Nitrite

NO₂⁻

66

Carbonate

CO₃²⁻

67

Hydrogen Carbonate

HCO₃⁻ (g)

68

Sulfite

SO₃²⁻

69

Sulfate

SO₄²⁻

70

Phosphate

PO₄³⁻

71

Hydrogen Phosphate

HPO₄²⁻ (g)

72

Dihydrogen Phosphate

H₂PO₄⁻ (g)

73

Acetate

C₂H₃O₂⁻

74

Permanganate

MnO₄⁻

75

Chromate

CrO₄²⁻

76

Dichromate

Cr₂O₇²⁻

77

Perchlorate

ClO₄⁻

78

Chlorate

ClO₃⁻

79

Chlorite

ClO₂⁻

80

Hydrosulfuric Acid

H₂S (aq)

81

Hypochlorite

ClO⁻

82

Hydrofluoric Acid

HF (aq)

83

Hydrobromic Acid

HBr (aq)

84

Hydroiodic Acid

HI (aq)

85

Nitric Acid

HNO₃ (aq)

86

Hypochlorous Acid

HClO (aq)

87

Nitrous Acid

HNO₂ (aq)

88

Chlorous Acid

HClO₂ (aq)

89

Chloric Acid

HClO₃ (aq)

90

Sulfuric Acid

H₂SO₄ (aq)

91

Perchloric Acid

HClO₄ (aq)

92

Sulfurous Acid

H₂SO₃ (aq)

93

Phosphoric Acid

H₃PO₄ (aq)

94

Phosphorous Acid

H₃PO₃ (aq)

95

hydrochloric acid

HCl (aq)

96

hydrogen chloride

HCl (g)

97

hydrogen fluoride

HF (g)

98

hydrogen iodide

HI (g)

99

hydrogen bromide

HBr (g)

100

hydrogen sulfide

H2S (g)

101

hydroselenic acid

H2Se (aq)

102

hydrogen selenide

H2Se (g)

103

bromic acid

HBrO3 (aq)

104

iodic acid

HIO3 (aq)

105

carbonic acid

H2CO3 (aq)

106

bromous acid

HBrO2 (aq)

107

iodous acid

HIO2 (aq)

108

hypobromous acid

HBrO (aq)

109

hypoiodous acid

HIO (aq)

110

hypophosphorous acid

H3PO2 (aq)

111

perbromic acid

HBrO4 (aq)

112

periodic acid

HIO4 (aq)

113

bromate

BrO3 -

114

iodate

IO3 -

115

bromite

BrO2 -

116

iodite

IO2 -

117

phosphite

PO3 -3

118

hypobromite

BrO-

119

hypoiodite

IO-

120

hypophosphite

PO2 -3

121

perbromate

BrO4 -

122

periodate

IO4 -

123

Mnemonic to remember most oxyanions

Nick the Camel ate Clam for Supper in Phoenix

124

Equal to # of consenants

# of oxygen (nick the camel...)

125

Equal to # of vowels

# of charge (nick the camel...)

126

Occurs when oxyanion gains an oxygen

Gains prefix "per"

127

Occurs when a gaseous oxyanion loses an oxygen

Gains suffix "ite" (g)

128

Occurs when an oxyanion loses 2 oxygen

Gains prefix "hypo"

129

Occurs when an aqueous oxyanion loses an oxygen

Gains suffix "ous"

130

Occurs when an aqeuous diatomic element ("have no fear") has all charge replaced by hydrogen

Gains prefix "hydro"

131

Occurs when some charge of a gaseous oxyanion is replaced by hydrogen

gains prefix "hydrogen"

132

The ratio moles in H2O

2 mol Hydrogen
1 mol Oxygen

133

The ratio of moles in H2O2

2 mol Hydrogen
2 mol Oxygen

134

Equals molar mass of CH4 (MMCH4)

MMC + 4 x MMH (12.011 g/mol + 4 x 1.0079g/mol)

135

The unit for molar mass (MM)

grams (g)
moles (mol)

136

The symbol for molar mass

MM

137

What you do when given grams

CONVERT TO MOLES

138

Find the empirical formula given grams or moles of the elements in a compound

Divide all mol by lowest mol to get the number of each element in the compound. Multiply to as necessary to get whole number.

139

Find molecular formula

Multiply each empirical element by quotient of MMmolec. / MMemp.

140

The limiting reagent

The reactant completely consumed, leaving other reactants partially consumed.

141

Water in a solution.

The solvent

142

Sugar in a solution

The solute

143

The molarity at which a solute is considered insoluble in a solvent.

<0.01 M (mol/L)

144

Symbol for molarity

M

145

Equation for molarity (M)

moles (mol)
liter (L)

146

Example of alternate concentration concerning occupation of space

% by volume (vol/vol)

147

Example of alternate concentration concerning mass

% by mass (m/m or w/w)

148

Given moles of CO, equation to find grams of C

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149

Given grams of CO, find moles of O

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150

Given grams of Mg(OH)2, find grams of Mg

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151

Find percent composition of Mg in Mg(OH)2

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152

Find empirical formula given 13.65% carbon and 86.35% fluorine

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153

Di

2

154

Tri

3

155

Tetra

4

156

Penta

5

157

Hexa

6

158

Definition of percent yield

The amount of theoretical yield that was actually obtained.

159

Equation for percent yield

Actual Yield
Theoretical Yield

160

Commonly referred to as Redox

Reduction - oxidation reactions

161

Reduction Involves Gain (RIG)

Refers to a reactant gaining an electron

162

Oxidation Involves Loss (OIL)

Refers to a reactant losing an electron

163

True about the reactant that oxidizes another reactant

Is the reactant being reduced

164

True about the reactant that reduces another reactant

Is the reactant being oxidized

165

The charge of a reducing agent

Increases (+) because it loses an electron

166

The charge of an oxidizing agent

Decreases (-) because it gains an electron

167

Standard pressure (atmospheres)

1 atm

168

Standard pressure supports how much water

10.33m

169

Standard pressure (mm Hg)

760 mm Hg

170

Standard pressure (kPa)

101.3 kPa

171

Standard pressure (torr)

760 torr

172

Standard pressure (psi)

14.7 psi

173

Equation for Boyles Law

P1V1 = P2V2

174

Equation for Charles law

V1 / T1 = V2 / T2

175

Equation for Avogadros law

V1 / n1 = V2 / n2

176

Avogadros number

6.022 x 10^23

177

Standard temperature and pressure (STP)

0oC and 1 atm

178

Ideal gas constant (R)

0.08206 (Lxatm)/(molxK)

179

Amontons Law

P1 / T1 = P2 / T2

180

Ideal gas law

PV = nRT

181

The units for pressure and temperature in the combined gas law

atm, oK

182

Combined gas law

P1V1 / T1 = P2V2 / T2

183

Equation given pressure (atm), moles (mol), and temperature (K)

P1 / N1T1 = P2 / N2T2

184

Equation for Daltons law of partial pressure

Ptotal = P1 + P2

185

Related to partial pressure ratios

Mole (mol) ratio of gases

186

Equation for gas density with 2 variables

Mass (g) / volume (L)

187

Equation for gas density (g/L) with 4 variables

Pressure (atm) x Molar mass (g/mol)
constant R (Latm/molK) x temperature (K)

P x MM
R x T

188

All gas density relationships

d = m / v = PMM / RT

189

Diameter of nucleus

10-14m

190

Volume occupied by electrons

10-10m to 5 x 10-10m

191

Acronym for wave energies (low to high)

ROY G BIV

192

Lowest wave energy

Radio waves

193

Highest wave energy

Gamma waves

194

Plancks constant

h = 6.626 x 10-34J-s

195

Speed of light

3.0 x 108

196

Relationship between wavelength and energy

Inversely proportional v = 1/λ

197

Equation for energy

E = h x c / λ

198

Equation for frequency (v) without energy

v = c / λ

199

Equation for wavelength (λ) without energy

λ = c / v

200

Units for frequency (v)

1 / seconds(s) or s-

201

What electrons do when emitting light

Jump down an energy level or orbital

202

Line spectra Bohr couldn't explain

>He

203

Found electrons act like waves

Louis De Broglie

204

De Broglie found electrons act like waves through

Passing β-radiation through a crystal causes diffraction.

205

Heisenburgs uncertainty principal

Cannot know both electron speed and location.

206

Idea Heisenburgs uncertainty principal was founded on

The energy required to view an electrons location changes its speed.

207

Developed by Schroedinger

Wave mechanics

208

Wave mechanics concerns

Electron energy, mass, and space occupied in probabilities

209

Definition of an orbital

Area an electron has 90% probability of being found

210

Orbital with 2 electrons

s orbital

211

Orbital with 6 electrons

p orbital

212

Orbital with 10 electrons

d orbital

213

Orbital with 14 electrons

f orbital

214

Energy level the s orbital starts at

1

215

Energy level the p orbital starts at

2, after 2s2

216

Energy level the d orbital starts at

3, after 4s2

217

Energy level the f orbital starts at

4, after 6s2

218

Hund's rule

Each available orbital in an energy level is filled with 1 electron before doubling up

219

Pauli exclusion principal

No electron in an atom has the same 4 quantum numbers (energy level [n], orbital [l], spatial orientation [m1], spin [ms])

220

Shorthand for core electrons configurations

Written as last noble gas ([Ar]4s1 for K)

221

Shared by columns in periodic table

# of valence electrons

222

Core electrons shield valence electrons from

Nuclear force of the nucleus

223

Definition of ionization energy

Energy to remove an electron

224

Rubidium I.E., in relation to hydrogen

Lower ionization energy

225

I.E increases

From francium to helium

226

Definition of electron affinity (EA)

Ability of an atom to add electrons