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Flashcards in MCAT Chemistry 2 Deck (500):
1

ex of a closed syst

steam radiator

2

____________ is a meas of the averag ke of the particles in a syst

temperature

3

G is ?

free energy

4

bond e or bond dissoc e is?

averag of the energy req to break a part type of bond in one mole of gaseous molec

5

give an ex of an isolated system?

insulated bomb reactor

6

T or Fit is not poss to meas H directly

T

7

system

partic part of universe being studied

8

open system is?

can exch matter and e with surroundings.

9

ΔG <0 means a proc can?

occur spont

10

qrev/T =

ΔS

11

conversion of C (diamond) to C (graphite) is __________ but its rate is slow

spont

12

bond ____________ is always endothermic

breakage

13

stand heat of format of a compound is?

enthalpy ch that would occur if one mole of a comp were formed direct from its elem in their standard states

14

T or Fheat and temp are different

t

15

boiling has a ____________ in entropy

increase

16

state functions

prop that dep only on initial and final states of the system

17

E or U is?

internal energy

18

the react quotient Q is=

Q=[C]^c[D]^d/[A]^a[B]^b

19

units of heat?

joules

20

ΔH ΔS Outcome ?- +

spontaneous at all temperatures

21

heat changes at const press, is

enthalpy

22

isothermal process is?

temperat of systm is const

23

a system undergoes a process when?

when one or more of its prop change

24

ΔH ΔS + +

spontaneous only at high temperatures

25

-RTlnKeq =

ΔG

26

system can be?

isolatedclosedopen

27

ΔH ΔS - -

spontaneous only at low temperatures

28

t or fentropy is a state function

T

29

t or frate of a react dep on ΔG

f

30

for hesses law, if you multiply the prod and react by 3 ( or a cert number) then what must you do to ΔH?

also multiply it by 3

31

endothermic react?

absorb e

32

q=mc chTc is?

specific heat

33

q=mc chTq is?

heat

34

T or FΔHreact=Hprod-Hreact

T

35

enviroment

everthing outside system

36

freezing has a ___________ in entropy

decrease

37

a spontaneous react may or may not?

proceed to completion

38

bomb calorimeter, the overall system is?

adiabatic

39

standard free energy of format of a comp is?

free energy ch that ocurs when 1 mol of a comp in its stand form

40

ΔHf of an element in its standard state is?

zero

41

second law of thermodyn says that?

all spont proceed so that entropy of the systm plus its surr increases

42

temp is?

a measure of average ke of particles in a systm

43

name 7 state funct?

tempvolenthalpyentropyfree energyinternal energy E and U

44

H is?

enthalpy

45

what condit are heat changed measured under?

const vol calorimetryconst press calorimetry

46

enthalpy of a process does not depend on the?

path

47

rate of a reac depends on?

activat e

48

standard condit in thermody must not be confused with?

standard temperature and pressure in gas laws STP

49

closed system is?

can exch e but not matter

50

ΔH ΔS + -

nonspontaneous at all temperatures

51

isobaric process is?

press of systm is const

52

q react + q water + q steel=0 in a ?

bomb calorimeter

53

Goose hunters take shotguns reminds you of?

ΔG=ΔH-TΔS

54

discuss when water boils in terms of ch in G

TchS> chH only when T is above 373 K

55

reverse reaction has the same __________ as that of the forward react, but its ________ is oppos

enthalpy,but its sign is oppos

56

for simple diatomic molec such as H2, bond dissoc e can be easily measured?

spectroscopically using react E=hv E=bond dissoc e and v is freq of light absorbed

57

ΔS= ?/T

qrev/T

58

when a react commences ΔG=?

ΔG= ΔG° + RTlnQ

59

calorimetry measures?

heat changes

60

spectroscopy can be used to measure the bond dissoc e of?

simple diatomic molecules

61

standard free energy ΔG is ?

ΔG of a proc occur at 25 C and 1 atm . concentr of sol are 1 M

62

ΔS=Sfin- Sinit

Sfin- Sinit

63

standard heat of react is?

ΔHreact= (sum of ΔHf of prod)-(sum of ΔHf of react)

64

exothermic react?

release e

65

a system undergoes a ____________ when one or more of its properties change

process

66

standard state of a subs is the form a subst takes?

at 25 C and 1 ATM

67

Hess law says?

that enthalpies of react are additive

68

heat absorbed by a system is?

pos

69

Thermodynamics

- Is the study of energy and it's relationship to macroscopic properties of chemical systems. It's functions are based on probabilities and are only valid for systems that are composed of a large number of molecules. * I.e The rules of thermodynamics gov

70

System vs. Surroundings

A system is a part of a universe that is the macroscopic body under study, and the surroundings is everything else. Systems are based on mass and energy exchange with the surroundings. There are three systems: 1) open: exchange both mass and energy with surroundings.2) closed: Exchange energy but not mass3) isolated: Do not exchange energy or mass.

71

State Function

A state is the physical condition of a system described by a specific set of thermodynamic property values. Such properties that describe the state of a system are called State Functions.Two types of properties used to describe the state of a system: 1) Extensive: properties are propotional to the size of the system (eg. V + n)2) Intensive: are independent of the size of the system. ( eg. P + T )Seven State Functions: U (internal Energy), T, P, V, H (Enthalpy), S (entrophy) and G (Gibbs energy).

72

Heat

(q) *Heat has three forms:1) Conduction2) Convection3) Radiation*Always the movement of energy from hot to cold.

73

Work

work = P∆V

74

Convection

- Is the thermal energy transfer via fluid movements. - differences in density or pressure drive warm fluid to cold fluid*Oceans and air currents are commone examples of convection.

75

Radiation

- is thermal energy transfer via electromagnetic waves. (eg. heated metal red, orange, white, blue-white) - All objects above 0K radiate heat.P= σ∈AT⁴

76

The First Law of Thermodynamics

∆E = q + w*Warning: work ON the system is positive for "convention" a passage on MCAT may define work done BY the system as a postive in which case you use this formula.

77

The Second Law of Themodynamics

*See heat engines

78

Internal Energy Types

MCAT may refer to internal energy as "heat energy", "thermal energy," or even "heat.""Heat energy and thermal energy" are really the vibrational, rotational and translational parts of interanl energy. Called this because they effect temperature. "Heat" is a transfer of energy. Don't mistake.

79

Temperature

0 K = 273°C

80

Enthalpy

(H) = is a man- made property that accounts for this extra capacity to do PV work. Unlike functions such as pressure, volume , and temperature, enthalpy is not a measure of some intuitive property. Defined more so as an equation then a property:H≡ U + PV

81

Standard State

* Don't confuse with STP. STP is at 0°C whereas standard state is at 25°C and is arbritrarily assigned an enthalpy value of 0 J/mol

82

Reference Form

Define

83

Standard Enthalpy of Formation

Define

84

Exothermic vs. Endothermic

Define

85

Activation Energy

Define

86

Transition state

Define

87

Intermediates

Define

88

Catalyst

Define

89

Irreversible vs. reversible reactions.

Define

90

Third law of Thermodynamics

Define

91

Spontaneity

Define

92

Gibbs Free Energy

Define

93

Thermodynamics

- Is the study of energy and it's relationship to macroscopic properties of chemical systems. It's functions are based on probabilities and are only valid for systems that are composed of a large number of molecules. * I.e The rules of thermodynamics gov

94

What are two ways to transfer energy between systems

#NAME?

95

Conduction

- is thermal energy transfer via molecular collisions. - Requires direct physical contact- an objects ability to conduct heat is called it's thermal conductivity (k) **On page 46 go over again.Q / t = kA [ T(h) - T(c) ) / L ]

96

Stephan- Boltzman Law

P= σ∈AT⁴Where:A is the surface area of the objectT is the temperature σ is the Stefan- Boltzman constant (5.67 x 10⁻⁸ )∈ is the emissivity of the object's surface.

97

Zeroth Law of Thermodynamics

Discovered after 1st, 2nd, and 3rd Law of Thermodynamics. All three rely on this Law, as it is based off of temperature. States: Two systems in thermal equilibrium with a third system are in thermal equilibrium with each other. The law declares that the two bodies in thermal equilibrium share a termodynamic property, which must be a state function. This is temperature.

98

Standard Temperature and Pressure

(STP)

99

Mean free path

Define

100

Ideal gas

Define

101

Kinetic Molecular theory

Define

102

Ideal gas law

PV = nRTK.E.(avg) = 3/2 (RT)

103

Standard Molar volume

22.4L ---> At STP one mole of any gas (behaving ideally will occupy this amount.

104

Partial Pressure

Pa= Xa Ptotal

105

Dalton's Law

P (total) = P1 + P2 + P3 ...

106

Graham's Law

v1/v2 = √M2 / √M1

107

Effusion

effusion rate 1/ effusion rate 2 = √M2 / √M1

108

Diffusion

Define

109

Chemical Kinetics

Define

110

Collision Model

Define

111

Activation Energy

Define

112

The effect of temperature on the rate of reaction

The rate of a reaction increases with temperature

113

Intermediates

Define

114

Rate Law

Define

115

Order of each respective reactant

Define

116

Overall Order of a reaction

Define

117

Rate Determining Step

Define

118

Catalyst

Define

119

Heterogeneous catalyst

Define

120

Homogeneous Catalyst

Define

121

Chemical Equilibrium

Define

122

Equilibrium Constant

K

123

The law of Mass Action

Define

124

Reaction Quotient

Define

125

Le Chatelier's Principle

Define

126

Metals

- large atoms that tend to lose electrons and form positive ions (cations) to obtain a noble gas configuration. - Can bond with other metal atoms (metallic bonding) - can also form ion bonds with negative ions (anions) *Usually non-metals. -Metals are

127

Ductile

- easily stretched. - occurs in metals- if you stretch or hammer e- the metals can adapt (which is why you can stretch metal without breaking them)

128

Luster

#NAME?

129

Non- metals

- form negative ions (anions) to obtain a noble gas configuration. - form covalent bonds with other non-metals- form ionic bonds with positive ions (usually metals)- poor conductors of heat and electricity- non-malleable, brittle and possess low to moder

130

Metalloids

#NAME?

131

Transition metals

- wont be focused on for MCAT, just know that they exist and that this might be a passage based question. - 1/2 vs fully filled orbitals occur here often.

132

Alkaline Earth Metals

- Harder metallic sollids have 2 valence electrons (can easily form +2 cations)- have a higher melting point than solf metallic solids- Heavier are more reactive than lighter ones.

133

Halogens

F₂: gas at RT and 1 atmCl₂: gas at RT and 1 atmBr₂: liquid at RT and 1 atmI₂: solid at RT and 1 atm* all highly reactive unless in diatomic form. -7 valence e-

134

Noble gases

- 8 valence e-- complete octet so very stable and non-reactive.

135

How do you determine the number of valence electrons in a main group atom?

- look at the group numberif group 1= 1 e- on valence shell (+1)group 2= 2 e- on valence shell (+2)recall: First shell closest to nucleus = ground stateLast shell closest to outside world = Valence e- shell. *Valence e- are involved in chemical bonding

136

Atoms

Define

137

Protons

Define

138

Electrons

Define

139

Nucleus

Define

140

Neutrons

Define

141

Elements

Define

142

Mass Number

(A)

143

Atomic Number

(Z)

144

Isotopes

Define

145

Atomic Weight

Define

146

Molar Mass

(MM or M)

147

Atomic Mass Units

(amu)

148

Mole

Define

149

Avogadro's Number

Define

150

Periodic Table

Define

151

Period

Define

152

Groups or Families

Define

153

Metals

Define

154

Nonmetals

Define

155

Metalloids

Define

156

Transition Metals

Define

157

Inert Gases

Define

158

Shields

Define

159

Effective Nuclear Charge

(Z eff)

160

Periodic trends

Define

161

Atomic Radius

Define

162

Ionization Energy

Define

163

Second Ionization Energy

Define

164

Electronegativity

Define

165

Electron affinity

Define

166

Metallic Character

Define

167

Bonds

Define

168

Covalent Bonds

Define

169

Bond Length

Define

170

Bond Energy

*Or bond dissociation energy

171

Compound

Define

172

Empirical Formula

Define

173

Molecules

Define

174

Molecular formula

Define

175

Ionic Compounds

Define

176

Acids

Define

177

Binary Molecular compounds

Define

178

Physical reaction

Define

179

Chemical reaction

Define

180

Runs to completion

Define

181

Fundamental Reaction types

1) Combination: A+B ---> C2) Decomposition: C ---> A + B3) Single Displacement: A + BC --> B + AC4) Double Displacement: AB +CD ---> AD + CB

182

Principle Quantum Number

(n)*Note: quantum numbers are the equivalent of a mailing address for an electron. - Description: Energy level and average distance from nucleus- Possible values: n= 1, 2, 3 etc.

183

Shell

Define

184

Valence electrons

Define

185

Azimuthal Quantum number

(l)*Also known as angular momentum quantum numberDescription: Orbital shape (s, p, d, f)Possible values: l=0, to ...n-1l=0 is the s-subshelll=1 is the p-subshelll=2 is the d-subshelll= 3 is the f-subshell

186

Subshell

eg. s, p, d and fs= sphericalp= dumbbell shapedd= clover leaf

187

Magnetic Quantum number

m- Description: Orbital orientation (px, py, pz)- Possible values: m= -l to +lfor l=0 (spherical s orbital) there is only one orientation.for l=1 (dumbbell p orbital_ there are three orientations along each of the x, y, and z axes.

188

Atomic Orbital

Space around a nucleus in which e- have the ability to exist.- max 2 e- in each orbital.

189

Electron Spin quantum Number

m (s)- Description: Describes spin of electron- Possible values: m(s) = +½ or -½Either clockwise or counter clockwise.

190

Pauli exclusion Principle

#NAME?

191

Heisenberg Uncertainty Principle

- tells us that the exact position and momentum of such an electron cannot be measure simultaneously. *Ie the more you know about the momentum of an electron, the less you know about it's exact position.

192

Aufbau Prinicple

*Write out on a piece of paper before writing mcat.1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f 6s 6p 6d7s 7p

193

Electron Configuration

Know 5 points on page 29 of lecture manual.

194

Ground State

Define

195

Hund's Rule

Define

196

Planck's Quantum theory

Define

197

Photoelectric effect

- an important experiment demonstrating the particle nature of light. An electron can be ejected from the surface of a metal by shining light of a certain frequency. Incoming light (photon) must have enough energy to eject the electron.E(photon) = h(Plank

198

Wavelength vs. Frequency vs. Velocity

Wavelength = distance btwn two consecutive crests of troughs. measured in meters. Frequency= the number of crests that pass through a given point per unit of time (Hertz)Velocity= the distance a wave travels through space/ unit time (m/s)Frequency= (velocity) / (wavelength)

199

Wave particle duality

all matter, including light can be thought of as both a wave and as a particle.

200

Electromagnetic Radiation

(EMR) is simple energy. as energy travels forward electric and magnetic fields are created perpendicular to each other and to the direction of travel.

201

Order of Colours

RedOrangeYellowGreenBlueViolet

202

Absorption Line spectrum vs. Emmision Line spectrum

ALS= Dark lines on a light backgroundELS= Light lines on a dark back ground.

203

How do you determine the max number of e- an atom can have?

#ERROR!:parse

204

Mass Percent (of a solute)

Mass(solute)/Mass(solution)

205

Archimedes Principle

Any object wholly or partially immersed in a fluid is buoyed up by a force equal to the weight of the fluid displaced by the object

206

Buoyant Force

Buoyant Force=Density(medium)*Volume(object)*g

207

Apparent immersed weight

Weight-Weight of Displaced Fluid

208

Molar Volume of an Ideal Gas at STP

22.4 Liters

209

Empirical Formula

Simplest whole number ratio of atoms in a molecule

210

Mass Percent (of an element)

Mass percent=[(mass atoms)/(mass compound)]*100%

211

Molecular Formula

Requires multiplying the empirical formula by the whole number ratio of the molecular mass:empirical mass; requires knowing the molecular mass of the compound

212

Hygroscopic salt

Absorbs water; examples include calcium chloride and magnesium sulfate

213

KOH

Potassium Hydroxide; undergoes a combustion reaction with carbon dioxide to form solid potassium bicarbonate

214

Molarity

Moles of Solute/Liters of Solution

215

Molality

Moles of Solute/Kilograms of Solvent; does not change with temperature so it's used to calculate the boiling-point elevation and freezing-point depression of solutions containing non-volatile impurities

216

Dilution:Fold vs. Parts

Fold (based on the total volume); parts (based on the volume added)

217

Beer's Law

The specific wavelength of light absorbed and the intensity of the absorbance varies with the solute and its concentration respectively

218

Absorbance

Absorbance=εcl, where ε refers to the absorption constant (a constant for the solution at λ max), c refers to concentration, and l refers to path length (or the width of the cuvette).

219

Limiting Reagents

The reactant with the lowest ratio of actual moles to needed moles

220

Precipitation Reactions (or metathesis or double-displacement reactions)

Involves two aqueous salts that react to form spectator ions and a solid salt precipitate

221

Acid-Base Reactions

A reaction between an acid (proton donor) and a base (proton acceptor) that results in the formation of a neutral salt and water

222

Composition Reaction

The number of reactants exceeds the number of products; entropy decreases and bonds are formed

223

Decomposition Reactions

Reactant(s) decompose to form multiple products; entropy increases and bonds are broken

224

Oxidation-ReductionReactions

Transfer of electrons from one atom to another ; the oxidation state must change in a redox reaction

225

Oxidation

Loss of electrons

226

Reduction

Gain of electrons

227

Reducing Agent

The atom (or compound) that is losing electrons

228

Oxidizing Agent

The atom (or compound) that is gaining electrons

229

Combustion Reactions

A special type of oxidation-reduction reactions where the oxidizing agent is the oxidizing agent and the products are oxides (water and CO2)

230

Molarity

(M) * units: mols/LM= (moles of solute) / (Volume of solution)

231

Mole Fraction

(m) *has no unitsm= (moles of solute) / (kilograms of solvent)

232

Solving for X

X= (moles of solute) / (total moles of all solutes and solvent)

233

Parts per million

PPM= (mass of solute) / (total mass of solution) x10⁶*Note: PPM is NOT the number of solute molecules per million molecules. It is the mass of the solute per mass of solution times on million.

234

Solution Formation

Define

235

Vapor Pressure

Define

236

Nonvolatile Solute

Define

237

Raoult's Law

Define

238

Volatile Solute

Define

239

Solubility

Define

240

Precipitation

Define

241

Saturated

Define

242

Solubility product K(sp)

Define

243

Spectator Ions

Define

244

The common Ion effect

Define

245

Solubility Factors

Define

246

Solution

Define

247

Solute vs. Solvent

Define

248

Colloids

Define

249

London Dispersion Forces

Define

250

Solvation

Define

251

Hydration

Define

252

electrolyte

Define

253

Arrhenius Acid vs. Base

Acid = Is anything that produces hydrogen ions in aqueous solution. [ H+ ]Base = is anything that that produces hydroxide ions in an aqueous solution. [ OH- ]*Only aq solutions.

254

Bronsted and Lowry

DefinesAcid: as anything that DONATES a PROTON.Base: as anything that ACCEPTS a PROTON.

255

Lewis

More general then Bronsted/ Lowry or Arrhenius.Acid: Anything that ACCEPTS a pair of ELECTRONSBase: anything that DONATES a pair of ELECTRONS. *Includes B+L acids and bases and more because it also accepts molecules that have incomplete octets of electrons eg. BF3 or AlCL3. Also includes simple cations (Smaller the cation the higher the charge, the stronger the acid) *Excludes alkali and heavy alkaline earth metals.*

256

Acid + Base vs. Conjugate Acid + Base

Acid + Base= Refers to the Reactants [A]Conjugate Acid + Base= Refers to the Product [HA][HA] is the conjugate ACID of BASE [A-][A-] is the conjugate BASE of ACID [HA]*The stronger the acid/Base the weaker the conjugate base/Acid, HOWEVER, weak acids may have EITHER weak OR strong conjugate bases.

257

Polyprotic

Define

258

Diprotic

Define

259

Amphoteric

Define

260

How does molecular structure affect Acid strength

Define

261

Hydrides

Define

262

Autoionization of water

Define

263

Acid Dissociation Constant

K(a)

264

Finiding the PH

Define

265

Titration

Define

266

Titration Curve

Define

267

Equivalency point

or Stoichiometric point

268

Half equivalence point

Define

269

Buffered

Define

270

Henderson- Hasselbalch Equation

pH= pK(a) + log ( [A-] / [HA] )

271

Indicator

Define

272

Endpoint

Define

273

Polyprotic titrations

Define

274

Strong Acids

- Hydroiodic Acid (HI)- Hydrobromic Acid (HBr)- Hydrochloric Acid (HCl)- Nitric Acid (HNO₃)- Perchloric Acid (HClO₄)- Chloric Acid (HClO₃)- Sulfuric Acid (H₂SO₄)

275

Strong Bases

- Sodium Hydroxide (NaOH)- Potassium Hydroxide (KOH)- Amide Ion (NH₂-)- Hydride Ion ( H-)- Calcium Hydroxide Ca(OH)₂- Sodium Oxide (Na₂O)- Calcium Oxide (CaO)

276

______ occurs at the cathode in an electrochemical cell. Electrons flow _____ the cathode.

reduction, toward

277

_____ occurs at the anode in an electrochemical cell. Electrons flow ____ the anode.

oxidation, from

278

The moles of solute over the kg of solvent

molality

279

Normality is the molar _____ per unit volume

equivalents

280

A ____ function is one in which the result in dependent only on the initial and final measurements and independent of the path needed to get there

state

281

_____ describes the the cagelike assortment of solvent particles around solute particles

solvation

282

Hess's Law

The enthalpy of reaction is equal to the difference between the ∆Hformationproducts-∆Hformationreactants

283

Isobaric process

One that occurs under constant pressure

284

Isochoric process

One that occurs under constant volume and no volume-pressure work is done

285

Constant-volume calorimeters (bomb calorimeters) are used to measure

the total heat absorbed or given off by a reaction

286

Adiabatic process

One that occurs without the transfer of heat

287

Collision theory of chemical kinetics

The rate of reaction is proportional to the number of collisions that occur between reaction molecules per second

288

Disproportionation

When a species is oxidized and reduced in the same reaction

289

Partial pressure equation

Ppartial = Ptotal * X(mole fraction)

290

Kinetic Molecular Theory of Gases

Gases have completely elastic collisions with themselves and the walls of their container. Gas particles have negligible volume, negligible attractive forces and exhibit random motion

291

Ideal gases

Monoatomic, no attractive forces or volume

292

Henry's Law:

The partial pressure of a gas above a solution is directly proportional to the partial pressure of the gas dissolved in the solution

293

Graham's Law:

Particles diffusion and effusion rates are inversely proportional to the squareroots of their molecular weights

294

Effusion

The passage of gas from high to low pressure through a small opening

295

Charles and Gay-Lussac's Law

At constant pressure, the volume of an ideal gas is directly proportional to its temperature

296

Boyle's Law

At constant temperature, the pressure of an ideal gas is inversely proportional to its volume

297

Molecular orbital

A regions where bonding or antibonding orbitals overlap, resulting in a low-energy bonding orbital or high-energy unstable antibonding orbital

298

Avagadros Principle

Different gases at the same temperature, pressure and volume, contain the same number of particles

299

Electrochemical reaction

One which needs or produces electricity

300

Nernst Equation

V=Vo-(0.6/n)logQ

301

VSEPR

Shows the 3D geometrical shape of a molecule that is based on the electronic interactions between bonding and non-bonding electrons

302

Isoelectronic

When two atoms have the same electron configuration

303

Magnetic Quantum number

M(L). Range from -L-L. Determines what orbital within the subshell the electron is likely to reside in.

304

Azimuthal quantum number

L. Ranges from 0-(n-1). Determines what subshell the electron is likely to be found in.

305

Aufbau Principle

Electrons fill an atom in order of increasing energy level

306

Exceptions to the aufbau principle

Cu and Cr. Cu takes an electron from a s orbital and places it in a d-orbital, completing its d-orbital. Cr takes an electron from an s orbital and places it in an d orbital, giving it the maximum number of unpaired electrons.

307

Hunds Rule

Electrons fill orbitals such that a maximum number of unpaired electrons results

308

Pauli Exclusion Principle

No two electrons in an atom can have the same set of 4 quantum numbers

309

Arrhenius Definition

Acid dissociates to form H+ in aqueous solution. Base dissociates to form OH- in aqueous solution.

310

Colligative properties

Those that depend only on the number of solute particles and not on the chemical makeup of the particles. (Boiling point elevation, freezing point depression, osmotic pressure, partial pressure)

311

Raoults Law

The vapor pressure above a solution is directly proportional to the mole fraction in solution. Ppartial = Ptotal*X

312

Energy of emitted electron =

hf- electron binding energy

313

malleability and other metallic characteristics ____ down and to the left on the periodic table

increase

314

most likely to have multiple oxidation states

transition metals

315

ionization energy

: E needed to remove an electron from a gaseous state

316

atomic mass increases ____ and ____ on the periodic table

right and down

317

which ions are typically the largest

negative ones (anions)

318

density _______ with increasing atomic radius

decreases

319

diatomic gases will have _____ densities than monoatomic gases

larger

320

T/F: solutions containing ions of transition metals are frequently colored

T

321

Units of ideal gas law

J/K*mol

322

Kinetic energy of n moles of gas

(3/2)nRT

323

When an outside force of other gases causes the molecule to move

Brownian forces

324

Variations from ideal gas conditions occur at _____

low volumes, low temperatures, high pressures

325

The a and b in the van der Waals equation represent...

a = attractive forces between particles, b=volume of molecules themselves

326

Ways of determining the reaction rate

Measure the inital reaction rate for a variety of reactant concentrations. Graph the concentration of the reactants as a function of time. Find the mechanism of the reaction.

327

N+L rule

determines which electron configuration comes next

328

The rate law constant (k) is affected by ____ and ____

temperature, catalyst

329

At equilibrium, the rate of forward reaction is ______ to/than the backwards reaction

equal

330

If a collision between two reactants does not form product, it can be assumed that....

the positioning was not right and or the energy of collision was not sufficient

331

As temperature increases, the range of reaction kinetic energy at which particles collide _____. The rate of collisions peaks at a higher kinetic energy than at _____ temperatures

increases, lower

332

Do liquids and solids enter into equilibrium equations?

no

333

Rates _____ increase with increasing temperature in endothermic and exothermic reactions

always

334

The emissivity value of blackbodies

1

335

Is it more efficient to heat quickly or slowly?

quickly

336

The area under the graph of a P vs. V graph is equal to _____

the work done on the gas

337

Free adiabatic expansion

No change in temperature since no work is done and no heat is lost (ideal gas)

338

Intensive properties

Not dependent on the amount of substance

339

Extensive properties

Dependent on the amount of substance

340

Percent yield

actual/theoretical * 100

341

Rapid cooling of polymers will result in an _____ _____ whle slow cooling will result in _____ _____

amorphous solid, crystalline solid

342

Diamond is a _____ _____ solid

network covalent

343

Heisenburg uncertainty _____ with increasing mass

decreases

344

Enthalpy will differ from the energy of reaction when...

The change in entropy is high

345

Heat of formation

Amount of heat needed to produce 1 mole of a given product

346

The heat of formation of natural elements is...

zero

347

Gases have ____ entropy than phases of the same substance

higher

348

Example of 2nd law of thermo

Heat always flows from high to low T

349

Bond formation is ____thermic and _____ entropy

exo, decreases

350

Any machine is less than ___% efficient because of ____ loss

50, heat

351

A substance at absolute zero has this entropy value

zero

352

Constant ____ and ____ are necessary in classfying spontaneity of reactions

temperature and pressure

353

When the velocity of a fluid increases, the pressure _____

decreases

354

When temperature decreases in a fluid, velocity _____, and thus pressure _____

increases, decreases

355

When density of a fluid decreases, volume ______ and pressure _____

increases, decreases

356

The pressure of a fluid against a flat surface is equal in momentum/∆t*A

Define

357

Ideal fluid characteristics

Volume flow rate is constant, no viscosity, incompressible

358

The flow of an ideal fluid _____ be determined by pressure changes alone

cannot

359

Possible phases of a solution

gas, liquid, solid

360

In ideal dilute solutions, solute particles ____ interact, and the mole fraction of the solvent approaches ____

never, 1

361

Ideal solutions obey _____ law, where solute and solvent molecules are similar and interact with each other similarily

Roults

362

Non crystalline homogenous solution

Colloid

363

Colloids do not ____ in solution. Cannot be _____ by centrifugation. May be separated by addition of _____, by raising_____ or by this technique.

settle, separated, electrolytes, temperature, dialysis

364

Scattering of light results in the appearance of the lights path

Tyndall effect

365

Colloid particles are ____ than thos of the solvent but not _____ enough to precipitate out

larger, large

366

Ionic compound naming:

-ate, -ite, hypo, per

367

conjugate base of carbonic acid

bicarbonate

368

Hydration involves

breaking of water-water H-bonds, formation of water-solute bonds

369

Ion number

number of water molecules that bind to an ion in an aqueous solution

370

If a compound is composed of two non-metals, it must be a ____ compound

molecular

371

A negative heat of solution indicates that the solute-solvent bonds are ____ than the solute-solute bonds.

stronger

372

Sublimation occurs when the vapor pressure of the solid is _____ than the partial pressure above it

greater

373

Boiling occurs when the vapor pressure of the liquid is greater than the _____ pressure above it

total

374

Vapor pressure ____ with temperature

increases

375

Atmospheric gas will condense when ____ pressure of the atmospheric gas is at least as great as the ____ pressure of the liquid at that temperature

partial, partial

376

Steam refers to

water vapor above 100ºC

377

A non volatile solute has ____ vapor pressure

zero

378

Ionic compounds with greater charges are typically ____ soluble in water

less

379

Deviation from Ksp in real exeriments is usually witness because of ____ ____ and _____ reactions that take ions out of solution

ion pairing, hydrolysis

380

Salt crystals nucleate gas bubbles causing them to ____

coalesce

381

Amorphous solids have ____ melting points

poorly defined

382

Gases nearly always form ____ phase(s)

one

383

Coffe cup calorimeters are ____ systems. They are used to measure ____ change.

open, enthalpy

384

When an impurity is added to a solid, its melting point ____ and ____

decreases, broadens

385

When an impurity is added to a liquid, its boiling point ____

increases

386

vant hoff factor for non-electrolytes

1

387

Strong acids

HCl, HNO3, H2SO4, HBr, HI, HClO4

388

Amphiprotic

Capable of gaining or losing a proton

389

Strong acids have ____ conjugate bases

weak

390

Strong bases have ____ conjugate bases

weak

391

weak conjugate acids can have ____ conjugate bases

weak

392

Strong acids are stronger than ____

H3O+

393

Strong bases are stronger than OH-

Define

394

Strong bases

H-, Na2O, N(3-)

395

Percent ionization of an acid is dependent on

temperature, identity of acid, concentration of acid

396

In living organisms, small pH changes can great ____ rate chagnes

large

397

Atoms with unfilled d orbitals can _____ visible light and move to these orbitals, causing brilliant colors in transition metals

absorb

398

HCl reacts with CO3- to form

CO2

399

Reaction involving the exchange of bonds between the chemical species. Often called double-displacement.

metathesis reaction

400

The strength of oxyacids _____ with the central atoms oxidation state

increases

401

Which acid is stronger HClO or HIO

HClO

402

Which acid is stronger HCl or HI

HI

403

Acidity increases with bond _____ and decreases with bond ____ and and increases with the conjugate base ____

polarity, strength, stability

404

A hydride contains ____ and one other element

H-

405

A substance with a high boiling point will generally have a ____ vapor pressure

low

406

The melting point of water will _____ under low pressure

increase

407

The melting point of most liquids will ____ under low pressure

decrease

408

Osmolarity

Π=iMRT

409

Isobaric

constant pressure

410

Pico (p)

Define

411

Alkaline Earths

Group IIa

412

Ksp of MX

Define

413

Homogenous Catalyst

same phase

414

Heisenberg uncertainty principle

Define

415

Reaction Entropy

Define

416

Anode Galvanic

negative electrode (oxidation)

417

Arrhenius equation

k = Ae^(-Ea/RT)

418

Metals

Elements that are characteristically electropositive, malleable, and ductile. These elements tend to be found on the left side of the periodic table, be lustrous, and have relatively low ionization energies and electron affinities

419

Halogens

Contains nonmetals, 7 valence electrons in it's outermost energy level. Very reactive

420

Equilbrium constant is ______ dependent

temperature

421

Effusion

one gas moves through air

422

Standard delta G =

-RT ln(K)

423

Titration

Analytical procedure in which a solution of known concentrations is slowly added to a solution of unknown concentration to the point of molar equivalency, thereby providing the concentration of the known solution.

424

Valence Electrons

Electrons occupying the outermost electron shell of an atom, participating in chemical bonds. Atoms with the same number of valence electrons tend to have similar properties (families in the Periodic Table).

425

STP

Standard Temperature and Pressure. 273 Kelvin (0 Celsius), 1 atmosphere (760 torr, 760 kPA).

426

Lewis definition

Acids defined as electron-pair acceptors and bases as electron-pair donors.

427

Half-Cell

An electrode immersed in an electrolytic solution that is the site of either oxidation or reduction in a galvanic (voltaic) cell

428

Noble gases

Contains nonmetals that are non-reactive. Full outermost energy level except helium which has 2.

429

Which group is most metallic?

group IA

430

Best insulator has _____ specific heat

highest

431

Speed is dependent on

kinetic energy only

432

autophagy

A process that describes lysosomes using their hydrolytic enzymes to recycle the cell's own organic material

433

Phase Diagram

A pressure vs. temperature plot showing the conditions under which a substance exists in equilibrium between different phases or in which the substance exists in pure phase.

434

Ion

A single or polyatomic particle with an electric charge.

435

Normality

Gram equivalent weight of solute per liter of solution, often denoted by N.

436

Disproportionation

Redox reaction, in which the same species is both oxidized and reduced.

437

Faraday's Constant

Denoted by F, 9.65x10⁴ coulombs/mol e⁻. Commonly used in the formula It = nF (I = Current, t = time (s), n = mol e⁻).

438

Indicator

A chemical species that changes color during dissociation, used to signal the end point of a titration.

439

Redox Half-Reaction

Hypothetical equation showing only the species that is oxidized or reduced in a redox reaction and the correct number of electrons transferred between the species in the complete, balanced equation.

440

solvation

sol, a chemical process in which solvent molecules and molecules or ions of the solute combine to form a compound

441

Amphoteric

having characteristics of both an acid and a base and capable of reacting as either

442

System

The part of the universe under consideration that is separated by some real or imaginary boundary from its surroundings

443

atomic radius

one-half the distance between the nuclei of two atoms of the same element when the atoms are joined

444

mole

the amount of a substance that contains as many particles as there are atoms in exactly 12g of carbon-12

445

bond length

the average distance between the nuclei of two bonded atoms

446

Decomposition reaction

(chemistry) separation of a substance into two or more substances that may differ from each other and from the original substance C>>>>A+B

447

Molecule

The smallest unit of a substance, composed of two or more atoms joined in covalent bonds, that still retains all the chemical properties of that substance

448

What charge to group 7 elements have?

-1

449

Entropy when energy into the system

entropy increases

450

Chemical Similarity between Atoms

stay in same column

451

Money Metals

Au, Pt, Ag, Hg, Cu, Ni

452

Speed of light (c)

3.00 x 10^8 m/s

453

Ion Product

Product of the molar concentrations of dissociated ions in solution at any point in the reaction other than equilibrium or saturation, where each ion is raised to the power of its stoichiometric coefficient. Denoted IP.

454

Molecular Formula

A chemical formula showing the actual number of atoms present in a certain compound.

455

Limiting Reagent

Reactant of a chemical equation that, given nonstoichiometric amounts, determines the amount of product that can form; the reactant that runs out first.

456

Molecular Orbital

Region in a molecule where atomic orbitals overlap, resulting in either a stable low-energy bonding orbital or an unstable high-energy antibonding orbital.

457

Acid Dissociation Constant

An equilibrium expression used to measure weak-acid strength, given by the ratio of the product of the products' molar concentrations to the product of the reactants' molar concentrations, with each term raised to the power of its stoichiometric coefficient. Denoted Ka.

458

Percent Yield

A ratio, calculated as a percentage, of the actual mass of product yielded to the theoretical yield of product mass.

459

Isothermal System

Process in which the system either gains or loses energy to maintain a constant temperature.

460

Isolated System

A system that can exchange neither energy nor matter with its surroundings.

461

Exothermic

A reaction that proceeds with the net release of energy (heat) into the surroundings

462

Magnetic quantum number

specifies the specific orbital in which the electron is most likely to be found., Third quantum number, designated as ml. Describes a particular orbital within a subshell where an electron is very likely to be found. Possible values are integers in the -1 to 1 range, including 0.

463

Chemical Kinetics

the area of chemistry that is concerned with reaction rates and reaction mechanisms

464

Rate-Determining Step

The slowest step in a reaction mechanism that determines the overall rate of the reaction

465

Effective Nuclear Charge

The resulting positive nuclear charge an outer electron senses after accounting for the shielding effect of inner core electrons. Abbreviated Zeff. Increases from left to right and from bottom to top on the periodic table

466

Concentration

tells you how much solute is present compared to the amount of solvent

467

Theoretical Yield

The expected amount of product yielded in a reaction according to reactants' stoichiometry

468

Atomic Absorption Spectrum

The spectrum of certain absorbed wavelengths of light corresponding to an atom's spectrum of emitted frequencies of light

469

Diamagnetic

An atom or a substance that contains no unpaired electrons and is consequently repelled by a magnet

470

Reduction

A reaction in which a species gains electrons

471

Atom

The basic building block of all matter in the universe. An atom is made up of three main components: protons, neutrons, and electrons

472

Azeotrope

A liquid mixture of two or more substances that has a constant boiling point greater than or less than the boiling points of its constituents. The vapor of this unique mixture has the same composition as the liquid state, making difficult to separate the constituents

473

Isochoric Process

A process in which volume remains constant and no net pressure-volume work is done

474

Bronsted-Lowry definition

Common definition of acids as proton (H+) donors and bases as proton acceptors

475

What charge do group 1 elements tend to have?

1

476

From left to right across the periodic table, metallic characteristics (increase, decrease)

decrease

477

Metallic Solid

molecules held in place by delocalized bonding

478

State Variables

conditions that must be specified to establish the state of the system, pressure, volume, temperature, and amounts of substances

479

Glavanic Reduction potentials

higher reduction potential is cathode, lower reduction potential is anode

480

Hund's Rule

orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin

481

Combination Reaction

A reaction in which two or more reactants combine into a single product.

482

Pfund Series

Set of spectral lines resulting when a hydrogen atom undergoes a transition from energy levels n>5 to n=5.

483

Ideal Gas Law

A unification of Boyle's Charles and Gay-Lussac's, and Avogadro's Principle into a formula describing the behavior of ideal gases: PV=nRT. (Pressure*Volume = moles*Avogadro's Number*Temperature).

484

Reducing Agent

A species that is oxidized in the process of reducing another species.

485

Lyman Series

Set of spectral lines appearing in the UV region when a hydrogen atom undergoes a transition from energy levels n>1 to n=1.

486

Specific Heat

The amount of heat required to raise one gram of a substance by 1 degree Celsius;heat capacity

487

Molality

The ratio of the number of moles of solute dissolved in one kilogram of solvent. molality (M = moles solute/kg of solution)

488

Adiabatic Process

A process in which no heat is transferred to or from the system by its surroundings

489

Reaction Quotient

A ratio of the concentrations of the products to the concentrations of the reactants at any point during the reaction aside from equilibrium, where each reactant and product in the expression is raised to the power of its stoichiometric coefficient. Commonly denoted by Q

490

Solution Equilibrium

When a solute is dissolved in a solvent, it will dissociate until reaching an equilibrium point at which the rate of dissociation equals the rate of precipitation of the solute, regardless of any additional solute introduced into the mixture

491

Solubility Product Constant

Product of the molar concentrations of dissociated ions in solution at saturation, where each ion is raised to the power of its stoichiometric coefficient. Denoted Ksp.

492

Aqueous Solution

a solution in which water is the solvent

493

Hydrogen Bonding

involves lone pairs of electron on an electronegative atom of one molecule and a polar bond to hydrogen in another rmolecule. They are confined tomolecules that contain O, N, and F atoms

494

Ideal Gas Postulates

1. molecules are very small compared to the distance between them.2. molecules are constantly moving3. pressure of the gas- collisions of the molecules with container walls4. molecules do not experience intermolecular forces5. KEave proportional to T

495

Osmotic Pressure

increase in pressure due to a solvent crossing a membrane into a more concentrated solution ΠV = nRT

496

Assumptions of ideal gases

No volume, no repulsive forces, elastic collisions, kinetic energy is proportional to temperature

497

Le Châtlier's Principle

When a system in equilibrium is placed under one of several stressors, it will react in order to regain equilibrium. In other words, act on a system, it will work back towards equilibrium.

498

Vapor Pressure

Partial pressure of a vapor when it is in equilibrium with its solid or liquid phase.

499

Single Displacement Reaction

Chemical reaction in which an atom or ion of one compound is replaced by another atom or ion.

500

Balmer Series

A set of spectral lines that appear in the visible light region when a hydrogen atom undergoes a transition from energy levels n>2 to n=2