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AP Chem Repractice > All Units > Flashcards

All Units Flashcards

1
Q

Molecules vs Ions

Molecules

A
  • Covalent
  • All states
  • Simplest = molecules
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2
Q

Molecules vs Ions

Ions

A
  • Electrostatic Forces of Attraction
  • Solids with 3d crystal lattices
  • No prefixes with exceptions
  • can become hydrates
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3
Q

Types of Reactions

Decomposition

A
  • Breaking down compound to form simpler compounds/elements
  • Endothermic
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4
Q

Types of Rxns

Ionization

A
  • Rxn with polar covalent ions to ions in H2O
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5
Q

Types of Rxns

Dissociation

A
  • Separation of ions in H2O
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6
Q

Types of Rxns

Melting

A
  • Solid ==> liquid
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7
Q

Solubility

Soluability Rule

A
  • Salts with Na+, K+, NH4+, NO3-
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8
Q

Solubility

Insolubility Rule of Thumb

A
  • max concentration less than .01 M, insoluable
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9
Q

Conduction and Electrolytes

Strong Electrolytes

A
  • More ions = More Strength
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10
Q

Conduction and Electrolytes

Weak Electrolytes

A
  • Weak acids and Bases
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11
Q

Conduction and Electrolytes

Non-Electrolytes

A
  • nonpolar componds/ org compounds except carboxylic acids and amines
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12
Q

Thermodynamics Introduction

Heat

A
  • speed or Energy of particles
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13
Q

Thermodynamics Introduction

Energy

A

Capacity to do work

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14
Q

Thermodynamics Introduction

Work

A

Action of Force through ∆x

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15
Q

Thermodynamics Introduction

Total Energy of object

A
  • KE: E associated with motion ==> Thermal Energy
  • PE: (bond energy); energy associated with position or composition
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16
Q

Energy Units

1 L*atm

A

101.325 J

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17
Q

Energy Transfer

Energy Transfer between system and surroundings

A
  • Sys decrease(-) = Increase surroundings(+)
  • Sys increase(+) = Decrease surroundings(-)
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18
Q

Definition of Specific Heat(c)

A
  • Amount of energy required to raise 1 gram by 1 C
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19
Q

Exothermic vs Endothermic

Exothermic

A
  • Graph increases then ends lower than starting level(∆ PE = -)
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20
Q

Exothermic vs Endothermic

Endothermic

A
  • Graph increases then ends higher than starting level(∆ PE = +)
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21
Q

Types of Systems

Open

A
  • matter and Energy exchanged with surroundings
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22
Q

Types of Systems

Closed

A
  • Only Energy may exchange with surroundings but not matter
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23
Q

Type of System

Insulated

A
  • No energy or matter exchanged with surroudings
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24
Q

Transfer between system and surroundings

Energy Transfer with work and heat

A
  • E = q + w
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25
Q

Transfer between system and surroundings

Heat/Q

A
  • Driving Force = ∆T
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26
Q

Transfer between system and surroundings

Work/W

A
  • Driving Force = - P * ∆V
  • W= Fd
  • Compression = (+)
  • Expansion = (-)
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27
Q

Heat

Heat Capacity

A
  • C
  • q=C∆T
  • slope of Q and ∆T
  • units = J/C
  • Extensive: Depends on mass
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28
Q

Heat

Specific Heat

A
  • c
  • c = Q/m/∆T
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29
Q

Calorimeter

Bomb

A
  • Constant Volume
  • qcal=Ccal* ∆T
  • qcal = qrxn = ∆Erxn
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30
Q

Calorimeter

Coffee-Cup

A
  • Constant Pressure
  • qcal=Ccal* ∆T
  • qcal = qrxn = ∆Erxn
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31
Q

Standard Enthalpy of Formation

Standard State

A
  • State of Pure Substance at l atm P and Temp of Interest(25 C)
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32
Q

Standard Enthalpy of Formation

Standard Enthalpy of change

A
  • ∆H^o
  • ∆H of reactants and products
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33
Q

Standard Enthalpy of Formation

Standard Enthalpy of Formation

A
  • ∆H when 1 mol substance formed from compund standard state elements
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34
Q

Standard Enthalpy of Formation

Standard Enthalpy of Formation Equation

A
  • ΔHoreaction=ΣΔHof(p)−ΣΔHof(r)
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35
Q

Hess’ Law

Hess’ Law

A
  • If equation can be explained as sum of 2+ equations, ∆ H for desired equation = ∆ H sum of other equation
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36
Q

Phase Changes

Heat absorbed into system

A
  • q > 0
  • melting, vaporization, sublimation
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37
Q

Phase Changes

Heat released into system

A
  • q < 0
  • freezing, condensation, deposition
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38
Q

Gas Laws

Interaction of External and Internal Pressure

A
  • No change: Internal Pressure = External Pressure
  • Compression: Internal Pressure < External Pressure
  • Expansion: Internal Pressure > External Pressure
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39
Q

Gas Laws

Calculating Pressure under liquid

Equations

A
  • With Atmosphere: P(H2O)[“Column Pressure”] + P(atm)[atmospheric/ barometric pressure]
  • In Vaccum: P(H2O)[“Column Pressure”]
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40
Q

Properties of Gases

What is a Vapor?

A
  • Gaseous state of substance usually a liquid or solid at room temp and pressure
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41
Q

Properties of Gases

Gas Elements

that are gases at Room Temp

A
  • Noble gases(are isolated atoms)
  • Diatomic Gases: H2, N2, F2, O2 and Cl2
  • HCl, NH3, CO2, N2O, CH4, HCN
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42
Q

Gas Characteristics

Gas Characteristics

A
  1. Gases assume shape and volume of container
  2. Move in constant, random motion but in straight line
  3. Gas Density less than Liquid Density or Solid Density
    • It is highly variable[*Increased T –> Decreased D vs Increased P –> Increased D] *
  4. Gases form Homogeneous Mixture with each other in any proportion [0 rxn = mutually miscible]
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43
Q

Pressure

Pressure equation

A
  • F/a
  • dgh[in Pascals]
  • mmHg is a measure as well as a pressure unit
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44
Q

Monometer

What is a Monometer

A
  • Measures gas P
  • P = difference of liquid height
  • 2 types
    1. Close-end(vaccum)
    2. Open-end(atm)
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45
Q

Pressure

Conversions

A

1 atm = 760 mmHg = 760 Torr = 101.325 kPa = 1.0325 bar = 14.7 lb/in^2 = 101,325 N/m^2

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46
Q

Monometer

Monometer Equations

A
  • For Close-Ended: Pgas = Δh(liquid)
  • For Open Ended:
    1. Pgas = Δh(liquid) + Patm[Pgas>Patm]
    2. Pgas = Δh(liquid) - Patm[Pgas
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47
Q

Ideal Gas Law Equation

What is STP?

A
  • Standard Temperature and Pressure
  • 273.15 K and 1 atm
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48
Q

Ideal Gas Law Equation

What is an Ideal Gas?

A
  • @ higher Temperature: More gases
  • @ Lower pressure: More space
  • Rules:
    1. Gases move randomly
    2. No attraction between particles
    3. “Infinite” volume and and Volume of gas not important
    4. Obeys simple gas laws: Boyle, Charles, Avagadro

MCQ?

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49
Q

Ideal Gas Law Equation

Boyle, Charles, Avagadro Equations

A
  • Boyle: V=1/P
  • Charles: V=T
  • Avagadro: V=n

MCQ?

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50
Q

Ideal Gas Law Equation

Combined Gas Law

A

P1xV1/(n1xT1)=P2xV2/(n2xT2)

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51
Q

Ideal Gas Law Equation

Units

A
  • P = any
  • V = any
  • n= mol
  • T = K
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52
Q

Ideal Gas Law Equation

Types of R

A
  • R = 0.08206 liter·atm/mol·K
  • R = 62.36 L·Torr/mol·K or L·mmHg/mol·K
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53
Q

Density equation

A

D=MP/RT

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54
Q

Molarity Equation

A

M=mRT/(PV)=DRT/P

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55
Q

Stoichiometry with Gases

Law of Combining Gas Volumes

A
  • Ratios from balanced equations can be in L instead of mol when Same Temp and P
  • If not use ideal gas law as part of stoic(Ex. 1/P to get rid/add Pressure)
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56
Q

Dalton’s Law of Partial Pressure

Equation for P.P.

A
  • Ptot=P1+P2+…Pi
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57
Q

Mole Fractions of Gases

Equations

A
  • Xi=ni/ntot=Pi/Ptot
  • Mole % = Mole Fraction * 100%
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58
Q

Kinetic Energy Molecular Theory

Postulates

A
  1. Gas particles so small and distance so large that individual volume is negligible
  2. Particles in Constant Motion
  3. No Forces between Particles(No Attraction or Repulsion)
  4. Average Ke proportional to T(in K)
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59
Q

Kinetic Energy Molecular Theory

Deviation of Ideal Gas Law Equation

A
  • P=1/3(N/V)mū^2
  • N= # of molecules
  • m = mass of 1 mole
  • ū^2 = Average of squared velocities
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60
Q

Kinetic Energy Molecular Theory

Proportions

A
  • KE proportional to T
  • Urms proportional to √1/M
  • Urms proportional to √T
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61
Q

Kinetic Energy Molecular Theory

Equalities

A
  • U1/U2 = √(M2/M1)
  • t1/t2 = √(M1/M2)
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62
Q

Kinetic Energy Molecular Theory

Root Mean Square Velocity Equation

A
  • √(ū^2) = √(3RT/M)
    1. R = 8.314 J/(mol·K)
    2. M= kg/mol
  • OR sum of speeds squared divided by number of molecules, then sqrted
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63
Q

Kinetic Energy Molecular Theory

What does Volatile mean?

A
  • Molecules that have weaker IMFS; easier to go from liquid to gas
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64
Q

Kinetic Energy Molecular Theory

What does Effusion Mean?

A
  • Gas escapes container through hole into an evacuated chamber
  • Diffusion Properties still in play(Go from high to low pressure)
  • Effusion inversely proportional to √M
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65
Q

Non-Ideal Gas Behavior

When does Non-Ideal Gas Behavior happen?

A
  • Low Temp, High Pressure
  • As Temp increases, deviation decreases: Less IMFs because of less interaction
  • As Pressure increases, deviation increases: Actual volume available will be greater than predicted because gas molecules take up more space
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66
Q

Non-Ideal Gas Behavior

Van der Waals Equation

A
  • P=nRT/(V-nb) - (n^2)(a)/(V^2)
    1. a and b gas constants
    2. nb = accounting for size of gas molecules
    3. n^2a = accounting for IMFs
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67
Q

Electrons

How Light Energy is Determined?

A
  • wavelength
  • frequency
  • energy
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68
Q

Electrons

When are wavelengths seen?

A

When electrons return to ground level

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69
Q

Electrons

Planck’s Constant

A
  • h
  • 6.022E-34 Jᐧs
  • Slope of E∝V
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70
Q

Properties of Light

What is Electromagnetic Radiation

A
  • EMR
  • Properties of Light
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71
Q

Properties of Light

Wavelength

A
  • Length of one wave
  • Meters
  • 𝝀
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72
Q

Properties of Light

Frequency

A
  • Number of wavelengths(or cycles) per sec passing a point
  • 1/s or s^-1 or Hz
  • 𝛎
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73
Q

Properties of Light

Speed of Light

A
  • 2.998E8 m/s
  • c
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74
Q

Properties of light

Amplitude

A
  • Higher is Brighter while Lower is darker
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75
Q

Properties of Light

Equations for Energy, and Light Parts

A
  • c = 𝝀 ᐧ 𝛎
  • E = h𝛎
  • E = hᐧc/𝝀
    * Is the Energy of 1 photon
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76
Q

What are Photons?

A

Particle side of electrons

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77
Q

Planck’s Work

What does delocalized mean?

A

the electron is not with a certain atom or nucleus

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78
Q

Planck’s Work

What is a Quantum

A

A packet of Energy for an electron to jump to the next electron level

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79
Q

Einstein’s work

What is Photon Energy related to?

A
  • Ephoton ∝ 𝛎
  • Ephoton ∝ 1/𝝀
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80
Q

PES equation

A

Ephoton = KEelectron + BEelectron
- BE = binding energy

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81
Q

Bohr’s Work

Energy Equation to move electron in Hydrogen

NOT IMPORTANT FOR TEST

A

ΔE = (-2.178E-18 J )((1/nf^2) - (1/ni^2))

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82
Q

de Broglie’s Work

Mass Equation

A

h/(𝝀v)

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83
Q

Quantum Numbers

Types

A
  • n = Priniciple quantum number
  • l = Angular quantum number
  • m = Magnetic quantum number
  • ms = Magnetic Spin
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84
Q

Quantum Numbers

Nodes

A

No possibility for electron(white shell)

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85
Q

Quantum Numbers

Principle Quantum Number

A
  • n
  • Integral Number
  • Related to size and energy of orbital
  • Corresponds to Bohr’s energy level
  • More n
    1. Increased Orbital
    2. Increased distance of electron from orbital
    3. Increased Energy
    4. Decreased energy between orbitals
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86
Q

Quantum Numbers

Angular Momentum Quantum Number

A
  • Shape
  • l = 0–> (n-1) for each n val
  • if n=3, possible orbitals s(0)–>d(2)
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87
Q

Quantum Numbers

Magnetic Quantum Number

A
  • m
  • integers specifying orbital orinetation
  • Values are from -l←→+ l
  • Includes 0
  • Example: l = 2: m=-2,-1,0,1,2 ⇒ 5 orbitals
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88
Q

Quantum Numbers

Electron Spin

A
  • Up = + .5 spin
  • Down = - .5 spin
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89
Q

Quantum Numbers

Difference between H and Multielectron atoms

A
  • H = subshell E levels that are degenerate(same n-int at same level)
  • Multi = lower orbital energies
    * Subshell of prinicple shell at different energies
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90
Q

Quantum Numbers

Rules for Electron Placement

A
  1. Pauli Exclusion Principle
    * No 2 electrons in atom has same 4 quantum numbers
    * electrons in 1/2 filled orbitals have parallel spins
  2. Hund’s Rule
    * One electron for each orbital before doubling up
  3. Aufbau Principle
    * Electron occupy lowest energy level possible
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91
Q

Electron Configuration

How do you write Electron Configuration

A
  • Removal Order
  • Ex) Se: [Ar] 3d^10 4s^2 4p^4
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92
Q

Electron Configuration

Valence Electrons

A
  • Outermost principle shell
  • Usually S or S and P
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93
Q

Electron Configuration

Shortcut using Noble Gases

A
  • Can use closest previous noble gas in brackets then build the rest of the electron config from there
  • Cannot use this for noble gas in ground state
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94
Q

Electron Configurations

Electron Configuration Exceptions

A
  • Cr: [Ar] 3d^5 4s^1 (Same with Mo)
  • Cu: [Ar]3d^10 4s^1(Same rule with Au and Ag)
95
Q

Ionization Energy

What is Ionization Energy

A
  • Minimum Energy neede to remove electron from atom or ion
96
Q

Ionization Energy

Ionization Energy Requirements

A
  1. Gas State
  2. Endothermic
  3. Valence electron first
  4. Successive removed for 2nd etc IEs
97
Q

PES

Why are X-rays used?

A
  • Can dislodge electrons
98
Q

PES

Relation between BE and KE

A
  • Inversely Related
99
Q

Magnetic Properties of electrons

Paramagnetic

A
  • one or more unpaired electron
  • attracted by a magnetic field
100
Q

Magnetic Properties of electrons

Diamagnetic

A
  • Paired electrons
  • opposite spins that cancel out their magnetic fields
  • Are not attracted to outside magnetic field
101
Q

Magnetic Properties of electrons

How are they detected

A
  • weighing a substance in the presence of a magnetic field
102
Q

Periodic Trends

Atomic Radius

A
  • Is the distance form the nucleus to the valence electrons
  • Across Period: decreased radius due to higher effective nuclear charge
  • Down Group: Increased radius due to higher n and a greater distance from the nucleus(higher n ==> higher V)
103
Q

Periodic Trends

Cations

A
  • Decreased sized compared to original
  • decreased electrons with same proton number
104
Q

Periodic Trends

Anions

A
  • Increased size
  • Increased electrons with same proton number
  • Electron - Electron Repulsion
105
Q

Ionization Energy

What is the first IE equal to?

A
  • Binding Energy
106
Q

Electron Affinity

Electron Affinity

A
  • Energy released when neutral atoms gain electrons
    1. Needs to be in gas state
    2. M(g) + 1 electron –> M^1-(g) + EA etc(Successive)
  • Exothermic
  • Increased energy leads to Increased negative EA
  • Outer electrons are delocalized: can move around which leads to sea of electrons
107
Q

Periodic Trends

Metallic

A
  • Most = Bottom Left
  • Least = Top Right
108
Q

Coulomb’s law

Force

A
  • (k x q1 x q2)/r^2
  • q = charges
  • Repulsion Force = (+)
  • Attraction Force = (-)
109
Q

Coulomb’s law

Energy

A
  • (k x q1 x q2)/r
  • q = charges
  • E released or required to make bond
110
Q

Lattice Energy

Lattice Energy

A
  • Energy required to seperate 1 mole of an ionic solid into gaseous ions
  • Higher LE ==> Higher Ionic Bond Strength
111
Q

Lattice Energy

Lattice Energy Calculations

A
  • Born-Haber Cycle = Applying Hess’ Law to its maximum
112
Q

Lewis Structures

Lewis Structures

A
  • Used to explain that atoms combined to acheieve more stable electrong configuration
113
Q

Lewis Structures

Ions

A
  • = [ ] + charge
  • Ex) .Cä. –> [Ca] 2+
114
Q

Lewis Structures

Coordinate Covalent Bond

A
  • Electrons donated from one atom to another
115
Q

Formal Charge

What are Formal Charges used for?

A
  • Determines most possible lewis structure
116
Q

Formal Charge

Formal Charge Equation

A
  • FC= V.E. - (non-bonding e-) - (# of bonds)
117
Q

Formal Charge

Formal Charge Rules

A
  1. FC= 0 for neutral or Ion FC= Ion charge
  2. FC = 0 more favorable than FC >/< 0
  3. Less FC > Greater FC
  4. Best strucure = Lewis Structure with FCs similar to ENs
    * Increased EN = Increased - FC
118
Q

Resonance

Resonance Theory

A
  • Molecule/ Ions with 2+ plausible Lewis Structures with diff e- distribution
  • Hybrids joined with double-headed arrows
119
Q

Delocalized Electrons (Again)

A
  • bonding electrons spread out over several atoms
120
Q

Exceptions to the Octet Rule

Expanded Octet

A
  • More than 8 electrons
  • Happens in elelments that have n > 2
  • Have s–>d sublevels
121
Q

Valence Bond Theory

Sigma (𝞂) Bond

A
  • 1 covalent bond where orbitals overlap
122
Q

Valence Bond Theory

Domain of Double Bond

A
  • 1
123
Q

Valence Bond Theory

Hybridization

A
  • Promotion of electrons and mixing of orbitals
  • Only central atoms
124
Q

Valence Bond Theory

Pi (𝛑) bond

A
  • For bonds after single bonds
  • ex) Double or Triple
  • Bonding pairs are parallel to each other
125
Q

Resonance Strutures

Relation to delocalized atoms

A
  • Having them makes a structure more stable and the more you have the more stable because the lectrons are able to spread out over a larger volume
126
Q

Polarizability

What is Polarizability?

A
  • Ability to induce another nonpolar molecule to have a momentary dipole between instantaneuous and indused dipole
  • Happens with Covalent Compounds
127
Q

IMF Trends

LDF Trends

A
  • Increased = Increased size
  • Increased = Increased polarizability
  • Increased = Greater Charge Separation
128
Q

Dipole-Dipole forces

What are they?

A
  • Polar molecules with positive “end” and negative “end” which have a permanent dipole which attracts other dipoles when closer
  • Stronger than LDF
129
Q

Important Note!

What to do with LDF on Test

A
  • Write full form: London Dispersion Forces
130
Q

Hydrogen bonds

Conditions

A
  1. H has to be attached to F,O,N
  2. Other end is also F,O,N with/without H
131
Q

Hydrogen bonds

What are they?

A
  • Strong and special d-d
  • Have to write both d-d + h-b
  • bet H proton and negative side of F,O,N
132
Q

Surface Tension

Definition

A
  • Amount of energy required to stretch
  • Increased surface of liquid by 1 unit area
  • Higher IMFs = Higher ST
133
Q

Surface Tension

Adhesion

A

IMFs between unlike molecules

134
Q

Surface Tension

Cohesion

A

IMFs between like molecules

135
Q

Viscocity

Definition

A
  • Measure of a fluid’s resistance to flow
  • Increased T = Decreased V
  • Increased V = Increased IMFs
  • Tangling ==> Increased V
136
Q

H2O

Properties

A
  • Excellent Solvent
  • High Specific heat
  • Ice Density is less than Water because struct for spread out
137
Q

Liquid-Vapor Equilibrium

Vapor Pressure

A
  • Partial pressure of liquids once equilbrium is established between gases and liquids
  • Book Definition: Gaseous molecules pressure form evap. liquid
138
Q

Liquid-Vapor Equilibrium

Dynamic Equilbrium

A
  • Rate of forward process(l–>g) = Rate of backward process(g–>l)
  • Constant Pressure
139
Q

Liquid-Vapor Equilibrium

When is Boiling Point?

A
  • When Vapor Pressure is equal to External pressure
140
Q

Phase Changes

Critical Point

A
  • Temperature that gas cant become a liquid no matter the pressure
141
Q

IMF Trends

How are physical characteristics affected by IMFs?(only writing abt. high)

BP, VP, Viscocity, ST, Cohesive Forces

A
  • Higher BP
  • Lower VP
  • Higher viscocity
  • Higher ST
  • High Cohesive Forces
142
Q

Hydrogen Bonding

Salicylic Acid

A
  • H-Bonding within molecule
143
Q

Chromatography Lab

Rf

A

solute Distance/ Solvent Distance

144
Q

VP Curves

Equation

A
  • ln(Pvap1, T1/Pvap2, T2)=(∆Hvap/R)((1/T2)-(1/T1))
145
Q

Alloys

Definition

A
  • Solid Soln consisting of 2+ metals or metals with 1 + nm
146
Q

Covalent Network Solids

Definition

A
  • Network of covalently-bonded atoms into 2D and 3D network, holding it firmly together
147
Q

Covalent Network Solids

Allotropes

A
  • Elements exist in 2 diff formes in same physical state
  • Ex) diamond and graphite for carbon
148
Q

Semiconductors

Definitions

A
  • Elements that are normally not conductors but are at high Temperatures or when coombined with other elements
149
Q

Semicoductors

N-type

A
  • Have an extra electron from bonding which can be used to create a voltage
  • Donor Impurities
150
Q

Semiconductors

P-type

A
  • Have 1 less electron which creates “positive holes” that constatnly shift and allow current to flow through holes
  • Acceptor Impurities
151
Q

Physical Properties of Solns

Solvents

A
  • determine soln’s state of matter
  • usually majority component
152
Q

Physical Properties of Solns

Solutes

A
  • Substance dissolbed/dispersed in the solvent
153
Q

Physical Properties of Solns

Solution Concentration

A
  • Amnt of solute/ amount of solvent/soln
154
Q

Physical Properties of Solns

Molality(m)

A
  • moles of solute/ kg of solvent
  • varies with mass of solvent and is independent of temperature
155
Q

Physical Properties of Solns

Enthalpy of Soln

A
  • Combination of Heat required breaking IMFs of solute and solvent plus the heat released for letting them mix
156
Q

Physical Property of Solutions

Non-Ideal Solutions

A
  • Not additive solutions(adding two solutions doesn’t result in a solutions with both that is not equal than combined amount: Ex) 50mL +50mL≠100mL)
  • Happens due to unequal IMFS
  • example of Ideal Mixture = Benzene and Methylbenzene
157
Q

Physical Properties of Solutions

How Heat of Solution impacts ideality

A
  • If = 0, Ideal
  • If < 0, stronger solute-solvent forces; exothermic
  • If > 0, weaker solute-solvent forces; endothermic
158
Q

Formation of a Saturated Solution

What is a Saturated Solution

A
  • Maximum amount of solid or gas allowed in a liquid with constant Temperature
159
Q

Formation of Saturated Solution

How temperature affects saturation

A
  • Increased ==> Increased in solids
  • Opposite in gases
160
Q

Formation of Saturated Solution

Henry’s Law

A
  • Increased Solubility = Increased Pressure
  • s = P * k
161
Q

Formation of Saturated Solution

K in Henry’s Law

A
  • Units: mg[gas]/100g[H2O]/atm
162
Q

Colligative vs Noncolligative Properties

What are Colligative Properties?

A
  • Physical Properties that depends only on concentration of solute particles, not their identity
  • Decreased solvent vapor pressure, Freezing Point Depression, Boiling Point Elevation
163
Q

Colligative vs Noncolligative Properties

Examples of Noncolligative Properties

A
  • Color, odor, etc.
164
Q

Volatile Solvents

What if solute is non-volatile?

A
  • Lower solvent volatility because less on surface
  • Raoult’s Law: Vapor Psolv = Xsolv ·P°solv
165
Q

Volatile Solvents

What if solute is volatile?

A
  • Have to consider both Vapor Pressures
  • Pressure[Tot] = P[solvent A] + P[solvent B]
    = (XsolvA · P°solvA) + (XsolvB · P°solvB)
166
Q

Fractional Distillation

How does Fractional Distillation work?

A
  1. Boil 2 liquids
  2. Lower Boiling Point will boil out
  3. Rise to the condenser
  4. distillate
    * Example: Water(50mL) + Ethanol(50mL) = mixture (< 100mL)
    • Distillation = 50 mL each again
167
Q

Chemical Kinetics

What is it?

A

How fast rxns take place, facts that affect rates, and mechanisms

168
Q

Chemical Kinetics

Factors affecting Kinetics

A
  1. Concentration
  2. Temp
  3. Surface
  4. Catalyst
169
Q

Chemical Kinetics

what does Rate of Rxn equal?

A

-(1/a)(Rate of Reactant A) = -(1/b)(Rate of Reactant B) = (1/c)(Rate of Reactant C) = (1/d)(Rate of Reactant D)

170
Q

Rate Law

What does Rate Law equal?

A

= k[A]^m[B]^n
* for aA + bB –> cC + dD
* m≠a and n≠b
* m = order of rxn wrt A
* n = order of rxn wrt B

171
Q

Rate Law

Overall Order of Rxn

A

m+n

172
Q

Rate Law

K

A
  • Proportionality constant
  • rate constant
  • function of T
173
Q

Reaction Rate

Methods to monitor change in reactant/product concentration

A
  1. Change in color using spectrophotometer
  2. Change in Pressure with manometer
  3. Change in Electrical Conductance
174
Q

Reaction rates

Zero Order

A
  • Rxn Rate: r=k
  • Integrated Rxn Rate: [A]t = -kt + [A]0
  • Half Life: ([A]0)/2k
175
Q

Reaction Rate

First Order

A
  • Rxn Rate: r=k[A]
  • Integrated Rxn Rate: ln[A]t= -kt + ln[A]0
  • Half Life: .693/k
176
Q

Reaction Rate

Second Order

A
  • Reaction Rate: r=k[A]^2
  • Integrated Reaction Rate: 1/[A]t = kt + 1/[A]0
  • Half Life: 1/([A]0* k)
177
Q

Activation Energy

What is it?

A

Energy barrier that prevents less energetic molecules from reacting
Separatesineffective and effective collisions wrt energy

178
Q

Activation Energy

Area under Maxwell-Boltzmann

A
  • Represents fraction of collisions that are effective with correct orientation
  • A= e ^(-Ea/RT)
  • R = 8.314 J/mol/k
  • T = K
179
Q

Activation Energy

Rate constant equation with Activation Energy

A
  • k = A * e ^(-Ea/RT)
  • A = frequency factor = constant of soln to find rate constant
  • R = 8.314 J/mol/k
  • T = K
180
Q

Activation Energy

To compare K values at different temperatures…

A

ln(k1/k2) = Ea/R(1/(T1) - 1/(T2))

181
Q

Reaction Mechanism

What is it?

A
  • Sequence of steps(each step = elementary step) and sum of all steps
  • In the ES, rxn order for each reactant = stoic coeff.
182
Q

Reaction Mechanism

Rate determining step

A
  • Slowest ES = Speed of overall rxn
183
Q

Reaction Mechanism

How to determine plausible mechanism

A
  1. ES must sum up to overall rxn
  2. Rate-determining step must have same rate law as overall exp.aly-det. rate law
184
Q

Reaction Mechanism

Intermediate

A
  • Produced in ES and consumed in later ES
185
Q

Reaction Mechanism

Rate determining step senarios

A
  • If the 1st step slow: 1st = rate of rxn and rest fast
  • If 2 + steps slow: Step right before slow step = equilibrium and slow step = RDS
186
Q

Activation Energy - Catalyst

What do they do?

A
  • Increases rate of rxn(faster rxn) by decreasing Ea
  • Present at beginning and end of rxn
187
Q

Activation Energy - Catalyst

RDS relation to Ea

A

Highest Ea = RDS

188
Q

Equilibrium

Chemical Equilibrium

A

Occurs when opposing rxn proceeds at equal rates
* Concentrations dont change with time making rxn appear stopped

189
Q

Equilibrium

What is K?

A
  • Equilibrium Constant
  • = [A]eq/[B]eq = Pressure of A eq/ Pressure of B eq
  • At Constant Temperature
  • K&raquo_space;1(>10): Equilibrium favors product formation
  • K «1( <.10): Equilibrium favors reactant formation
190
Q

Equilibrium

Law of Mass Action for general rxn

A
  • Kc = [C]^c * [D]^d/([A]^a * [B]^b)
  • Kp same just with P instead of [ ]
191
Q

Equilibrium

Rules

A
  • Solids not considered
  • Liquid water when in excess not considered
192
Q

Equilbrium

Converting from Kc–>Kp

A
  • Kp=Kc(RT)^∆n
  • R = .08206
  • delta n = moles of product gas - moles of reactant gas
193
Q

Modifying Equilibrium Expressions

Types of Modification

A
  1. Reversing Equation: K’c = 1/Kc
  2. Multiplying Equation by x: K’c = Kc^x
  3. Adding 2 equations: K’c = Kc1 * Kc2
194
Q

Rxn Quotient

What is it?

A
  • Qc/ Qp
  • Concentration ratio for Kc/Kp at non equilibrium conditions
  • Not constant but allows us to predict direection of net change
195
Q

Rxn Quotient

How it helps us predict

A
  • Q > K ==> backward rxn favored
  • Q < K ==> forward rxn favored
196
Q

ICE Table

What are they?

A
  1. I: Initial [ ] or P
  2. C: Change over time
  3. E: When at Equilibrium
197
Q

Le Chat’s principle

What is it?

A
  • Predicts how eq’m restored
  • System will react to stress
198
Q

Le Chat’s Principle

Adding or Removing a Substance in rxn

A
  • Adding: system will shift and increased rate of rxn that decreases [substance]
  • Removing: system will shift and decrease rate of rxn that decreases [substance]
199
Q

Le Chat’s Principle

Stress sys by changing T

A
  • Exo:
    1. Increased T = Increased R and Decreased P
    2. Decreased T = Increased P and Decreased R
  • Endo:
    1. Increased T = Increased P and Decreased R
    2. Decreased T = Increased R and Decreased P
200
Q

Le Chat’s Principle

Change P or V of Sys

A
  • As P increases or V decreases(decrease space)
    • eq net shift to side with less number of moles of gas
  • As P decreases or V increases(increase space)
    • Shift to side with more moles of gas
  • If number of moles reactant = number of moles product, then there is no shift
201
Q

Le Chat’s Priniciple

Special Cases

A
  1. Partial Pressures of reacting changes does change with adding or removing inert gases
  2. In a rigid container: Total Pressure will increase but Partial Pressures don’t(no shift)
  3. Constant External Pressure[balloon/ moveable piston]
    1. Same as ∆Pressure
202
Q

Le Chat’s Principle

Changes that will not cause Eq. Shift

A
  1. Adding Inert Gas to constant volume container
  2. Catalyst
  3. Pure solids or Liquids
203
Q

pKw, pH, pOH

Equations

A
  • pH= -log([H3O+])
  • Also H+ = 10^-pH
  • pOH= -log([OH-])
  • OH- = 10^-pOH
  • pKw = pH + pOH = -log([OH-] [H3O+])
  • Ka * Kb = Kw
204
Q

pH, pOH

How to decide concentration of H3O+ or OH-

A
  • if in strong acid just that from the acid because self-ionization of water is negligible
205
Q

Conjugate Base/Acid

Hydrolysis Rxns

A
  • Acid - base rxns between ions(from salts) and H2O molecules
  • May react with water in a/b rxn
206
Q

Conjugate Base/Acid

pH vs H3O+ / OH- concentration

A
  • 7: [H30+] = [OH-]
  • <7: [H30+] > [OH-]
  • > 7: [H30+] < [OH-]
207
Q

Conjugate Base/Acid

Determining Conjugate Base/Acid

A
  • If acid, then find with extra electron (goes from donator to acceptor) ⇒conj base
  • If base, then find with less electron (goes from acceptor to donator) ⇒ conj acid
208
Q

Acid, Bases, and Acid-Base Equilibria

Acid

A
  • Proton Donor
209
Q

Acid, Bases, and Acid-Base Equilibria

Base

A
  • Proton Acceptor
210
Q

Acid, Bases, and Acid-Base Equilibria

Strong Acids

A
  1. HCl
  2. HBr
  3. HI
  4. HNO3
  5. H2SO4
  6. HCLO4
    * Complete Ionization: →
211
Q

Acid, Bases, and Acid-Base Equilibria

Weak Acids

A
  1. CH3COOH
  2. NH4+
    * Partial Ionization: ⇌
212
Q

Acid, Bases, and Acid-Base Equilibria

Strong Bases

A
  • Group 1A and 2A with Hydroxides
  • Complete Ionization: →
213
Q

Acid, Bases, and Acid-Base Equilibria

Weak Base

A
  • NH3
  • Partial Ionization: ⇌
214
Q

Acid, Bases, and Acid-Base Equilibria

Ka/ Kb

A
  • Acid/Base Ionization Constant
215
Q

Acid, Bases, and Acid-Base Equilibria

Relation bet a/b and conj a/b

A
  • Stronger the original, weaker the conjugate and vice versa
  • Favored in the direction = stronger to weaker
216
Q

Acid, Bases, and Acid-Base Equilibria

Kw

A
  • Ion product of H2O
  • = [H3O+][OH-] = 1.0 * 10^-14 M
  • both H3O+ and OH- have concentrations of 1.0 * 10^-7
217
Q

Acid, Bases, and Acid-Base Equilibria

How to deal with weak Acid/ Base equilibria

A
  • Use ICE Table
  • If M(a/b)/K(a/b) > 100, then [A/B]≈ Eq. value
218
Q

Acid, Bases, and Acid-Base Equilibria

Which ions hydrolyze

A
  • Weak acids or bases hydrolyze appreciably
  • Strong a/b form neutral solutions
219
Q

Acid, Bases, and Acid-Base Equilibria

Relationship between weak and strong a/b

A
  • Weak Acids and Strong Bases = basic
  • Strong Acids and Weak Bases = acidic
  • Weak a + b = any of them
220
Q

Common Ion Effect

Common Ion Effect

ASK ABT THIS

A
  • Supression of ionization of weak acids or weak bases by presence of common ion from a strong electrolyte
221
Q

Buffer Solutions

Definition

A

Solution that changes pH only slightly when small amounts of strong acid or base added
* Usually contain: weak acid/base with salt

222
Q

Buffer Solutions

Buffer Rules

A
  1. ratio of [conjugate base/ acid] / [weak acid/ base] between .10 and 10
  2. both [conjugate base/ acid] and [weak acid/ base] exceed Ka by factor of 100+
223
Q

Buffer Solutions

Henderson–Hasselbalch equation

A

pH = pKa + log([conj base]/ [weak acid])

224
Q

Buffer Solutions

Heylman equation

A

pOH = pKb + log([conj acid]/ [weak base])

225
Q

Buffer Solutions

Buffer Action

A
  • best buffer is when [conj acid / base] = [weak base / acid] → best buffer!
226
Q

Buffer Solutions

Effective Buffer Range

A
  • Approximately when pH = pKa ± 1 pH-unit
227
Q

Acid-Base Titrations

Analyte

A
  • solution of unknown concetration
228
Q

Acid-Base Titrations

Titrant

A

Solution with known concentration

229
Q

Acid-Base Titrations

Equivalence point

AKA Stoichiometric pt

A
  • # of moles of H3O+ = # of moles of OH-
230
Q

Indicators

What are they

A

A weak organic acid that has a different color than conjugate base

231
Q

Indicators

Relation bet. [HIn] and [In-]

A
  • equal means intermediate color
  • [In-]/[HIn] >10: solution will be color of In-
  • [In-]/[HIn] < .10: solution will be color of HIn
232
Q

Indicators

Endpoint

A
  • Point where indicator changes color
233
Q

Indicators

Equation

A
  • Ka/[H3O+] = [In-]/[HIn]

AP Chem Repractice (1 decks)

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