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Chemistry Flashcards

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1
Q

xPrincipal quantum number

A

n (1,2,3…)

Defines what shell the electron is in. Higher n shells are higher in energy.

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2
Q

How many orbitals per shell?

A

n2

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3
Q

How many electrons per orbital?

A

2

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4
Q

How many electrons per shell?

A

2n2

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5
Q

Excited state electrons come back down to the ground state via ___ of energy.

A

release

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6
Q

The ___ ___ shows what wavelenghts of light are absorbed.

It looks like ___ lines on a ___ background.

A

The absorption spectrum shows what wavelenghts of light are absorbed.

It looks like black lines on a rainbow background.

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7
Q

The ___ ___ shows what wavelenghts of light are emitted.

It looks like ___ lines on a ___ background.

A

The emission spectrum shows what wavelenghts of light are emitted.

It looks like colored lines on a black background.

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8
Q

What are the quantum numbers

A
  • l: angular momentum (range from 0 to n-1
    • spdf: l=0,1,2,3 respectively
  • m: magnetic quantum number (range from -1 to 1 including zero)
  • s: spin quantum number (either +½ or -½)
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9
Q

spdf subshells

A
  • s holds 1 orbital, p holds 3, d holds 5, f holds 7
  • Each orbital holds a max of 2 electrons
  • higher subshells have higher energy
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10
Q

How are subshells filled in increasing energy?

A

Going across rows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 5s, 4f, 5d, 7s, 5f, 6d

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11
Q

A ___ on the period table represents the number of electrons in that subshell.

A

column

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12
Q

COnventional notation for electronic structure?

A
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13
Q

Aufbau principle

A

Shells/subshells of lower energy gets filled first.

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14
Q

Hund’s rule

A

When you fill a subshell with more than 1 orbital (p, d, f) you first fill each orbital with a single electron and with the same spin. Electron-electron repulsion in doubly occupied orbitals make them higher in energy than singly occupied orbitals.

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15
Q

Pauli Exclusion Principle

A

2 electrons in the same orbital must be of different spins

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16
Q

What is special about d4 and d9 elements?

A

Instead of s2d4 its s1d5 and s1d10 because they want to achieve a half-full or full d subshell.

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17
Q

Equation for effective nuclear charge?

A

Effective Nuclear charge = Nuclear charge - shielding electrons

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18
Q

What are shielding electrons?

A
  • Stand between the nucleus and the electron we are intrested in
  • In subshells closer to the nucleus (lower in energy) than the electron we are interested in
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19
Q

Higher the effective nuclear charge for an electron means it is ___ stable.

A

Higher effective nuclear charge = more stable (higher ionization energy, not easily knocked off)

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20
Q

Classification of elements into groups.

A
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21
Q

Alkali Metals

A
  • Single valence electron- low ionization energy, very reactive
  • Wants to lose an electron
  • More reactive as you go down because if increasing radii
  • Reacts with oxygen to form oxides
  • Reacts with water to form hydroxides and releases hydrogen
  • Reacts with acids to form salts and releases hydrogen
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22
Q

Alkaline earth metals

A
  • 2 valence electrons - relatively low in ionization energy, quite reactive
  • Wants to lose both electrons
  • More reactive as you go down because of increasing radii
  • Reacts with oxygen to form oxides
  • Reacts with water to form hydroxides and releases hydrogen
  • Reacts with acids to form salts and releases hydrogen
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23
Q

Halogens

A
  • 7 valence electrons (2 from s subshell and 5 from p subshell)
  • Wants to gain one electron to achirve full valence shell
  • More reactive as you go up because of deceasing radii
  • Reacts with alkali metals and alkaline earth metals to form salts
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24
Q

Noble gases

A
  • Full valence shell of 8 - high ionization energy couple with low electron affinity
  • Don’t react
  • Found in the oxidation state of 0
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25
Q

Transition metals

A
  • High conductivity due to free flowing (loosely bound) outer d electrons
  • In the presence of ligands (when in a chemical complex) the d orbitals become nondegenerate (different in energy)
  • Electron transitions between nondegenerate d orbitals gives transition metal complexes vivid colors
  • Varied oxidation states, always positive
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26
Q

Physical properties of Metals vs. Nonmetals

A
  • Metals: good conductor of heat and electricity, malleable, ductile, luster, solid at room temp
  • Nonmetals: poor conductor of heat and electricity, solid liq or gas at room temp, brittle if solid and without luster
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27
Q

Chemical Properties of Metals vs. Non-metals

A
  • Metals: good reducing agent, likes to lose electrons to gain a + oxidation state, lower electronegativity, partially positive in covalent bond, forms basic oxides
  • Nonmetals: Likes to gain electrons to form a - oxidation state, good oxidizing agent, higher electronegativity, partially negative in a covalent bond with metal, forms acidic oxides
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28
Q

Representative elements

A
  • s and p block
  • No free flowing outer d electrons
  • Valence fills from left to right
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29
Q

Definition of first and second ionization energies:

A
  • First ionization energy- energy needed to knock of first valence electron
  • Second ionization energy- energy needed to knock off second valence electron
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30
Q

Ionization Energy Trend

A
  • Decreases as you go down (increasing radii)
  • Inceases as you go right (decreasing radii)
  • Highest peaks (noble gases)
  • Lowest troughs (alkali metals)
  • Local maxima for filled subshells and half-filled p subshells
  • Second ionization energy is always higher than the first
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31
Q

What is electron affinity?

A

The amount of energy released when something gains an electron. (How easily it can gain an electron)

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32
Q

Electron affinity trends

A
  • Decreases down a group (larger radii)
  • Increases from L to R
  • Highest = Halogens
  • Lowest = noble gases
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33
Q

Definition of electronegative

A

How much something hordes electron in a covalent bond

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34
Q

Electronegativity trends

A
  • Increases towards the top right
  • F is most electronegative (N, O, F, Cl, Br are all v electroneg)
  • Noble gases can be electronegative if they participate in bond formatino (Kr and Xe)
  • More electroneg atom in covalent bond gets partial negative charge
  • If electroneg difference is too great- ionic bond will form
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35
Q

Atomic radii trend

A
  • size increases down a column
  • size decreases as you go across a row
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36
Q

What is an ionic bond?

A

electron transfer completely from one atom to another. Oppositely charged species attract each other via electrostatic interaction.

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37
Q

Formula for electrostatic energy?

A

Electrostatic energy=kq1q2/r

(Negative because charges are opposite)

(Larger indicates stronger bond)

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38
Q

What is lattice energy?

A
  • Lattice energy measures ionic bond strength.
  • It is the energy required to break the ionic bond.
  • Larger magnitude = harder to break
  • Proportional to the electrostatic attraction between ions
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39
Q

Coulomb’s Law

A

F=kq1q2/r2

  • k=9E9
  • Like law of gravitation, but G is tiny compared to k
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40
Q

What is a covalent bond?

A

Bond that results from sharing of electrons between two atoms, resulting in the overlap of their electron orbitals.

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41
Q

sigma vs pi bonds?

A
  • σ: single bonds, make up the first bond of double and triple bonds
  • π: double and triple bonds, make up the second bond in a double bond and both the second and third bond in a triple bond
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42
Q

What is the VSEPR number?

A

Total # bonds + unbonded electron pairs

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43
Q

Structure of borane?

A
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44
Q

Structure of borohydride ion?

A
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45
Q

Structure of Amine/Ammonia?

A
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46
Q

Structure of ammonium?

A
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47
Q

Structure of an Imine

A
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48
Q

What are resonance structures?

A

Multiple satisfactory Lewis structures for a molecule.

Shifts are fast, spending more time in stable resonance structures.

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49
Q

Characteristics of a stable resonance structure?

A
  • Octet rule satisfied in every atom (except Boron group and H)
  • No formal charges
  • If there must be formal charge, like charges are apart and unlike charges are close together
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50
Q

Equation for formal charge?

A

Formal charge = (valence electron # in unbonded atom)-(electron # in bonded atom)

Formal charge=(valence electron #)-(# dots + lines)

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51
Q

What is a Lewis acid?

A
  • Acceptor of electron pairs. They don’t have lone pairs on central atom.
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52
Q

Lewis base?

A

Lewis bases donate electron pairs. They have lone pairs on their central atom.

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53
Q

Dipole moment

(Is it greater or smaller with more electroneg difference?)

A
  • Depends on charge and distance
  • Greater electronegativity difference = greater charge = greater dipole moment
  • Greater distance separating charge = greater dipole moment
  • Things with a dipole moment are polar
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54
Q

Molarity vs. Molality?

A

Molarity = M = mol/L

Molality = m = mol/kg

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55
Q

Avogadro’s Number

A

6.02x1023

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56
Q

Density of water

A

1 g/mL

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57
Q

Density of lead

A

11 g/mL

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58
Q

Common Oxidizing Agents

A
  • Oxygen and Ozone
  • Permanganates- MnO4-
  • Chromates- CrO42-, Dichromates- Cr2O72-
  • Peroxides- H2O2
  • Lewis acids
  • Stuff with oxygens
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59
Q

Common reducing agents

A
  • Hydrogen
  • Metals (ex. K)
  • Zn/HCl
  • Sn/HCl
  • LAH (lithium aluminum hydride)
  • NaBH4 (Sodium borohydride)
  • Lewis Bases
  • Stuff with lots of hydrogens
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60
Q

What is disproportionation?

A

When an element in a single oxidation state reacts to form 2 different oxidation states. Disproportionation can occur when a species undergo both oxidation and reduction.

ex. 2Cu+ → Cu + Cu2+

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61
Q

Redox Titration Terms:

  • Analyte (A)
  • Titrant (T)
  • Standard (S)
  • Intermediate (X)
A
  • Analyte (A) = stuff with unknown concentration that you want to find out by titration
  • Titrant (T) = stuff you add drip by drip to determine how much is needed to complete titration
  • Standard (S)= something with an accurately known amount or concentration (produces a known amount/conc of I2)
  • Intermediate (X)= species that is not present in the net equation of the overall reaction
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62
Q

A color changes occurs at what point of a redox reaction?

A

Equivalence point

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63
Q

What is absolute zero?

A

0 K = -273 °C

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64
Q
  1. Freezing point of water?
  2. Room temp?
  3. Body Temp?
  4. Boiling point of water?
A
  1. Freezing point of water? 0 °C
  2. Room temp? 25 °C
  3. Body Temp? 37 °C
  4. Boiling point of water? 100 °C
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65
Q

Convert from C to K

Convert from C to F

A

K=C+273

F=1.8C+32

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66
Q

Atmospheric pressure in atm? mm Hg? torr? kPa?

A

1 atm = 760 mm Hg = 760 torr = 101 kPa

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67
Q

SI unit for pressure?

A

Pascals

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68
Q

Molar volume of a gas?

A

At 0 °C and 1 atm 22.4 L/mol

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69
Q

Ideal gas

A
  • Moves randomly
  • No intermolecular forces
  • No molecular volume
  • Perfectly elastic collisions (conservation of total kinetic energy)
  • Low P, High T
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70
Q

Deviations from ideal gas occur at ___ P and ___ T.

A

High P, Low T

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71
Q

Combined Gas Law

A

P1V1/T1=P2V2/T2

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72
Q

Boyle’s Law

Charles’ Law

A

Boyle (Constant T): P1V1=P2V2

Charles’ (Constant P): V1/T1=V2/T2

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73
Q

Kinetic theory of gases

A
  • Random molecular motion
  • No intermolecular forces
  • Negligible molecular volume
  • Perfectly elastic collisions (conservation of KE)
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74
Q

___ is a measure of the average KE of a gas.

A

Temperature

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75
Q

Pressure of a gas is due to:

A

molecules constantly colliding with the walls of its container.

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76
Q

What is diffusion?

A

Random molecular motion- causing a substance to move from an area of higher conc to an area of lower concentration (diffusion down concentration gradient)

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77
Q

What is effusion?

A

Random molecular motion, causing a substance to escape a container through a very small opening.

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78
Q

Graham’s law of effusion/diffusion?

A

Rate1/Rate2=(M2/M1)½

(Lighter molecules travel faster)

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79
Q

Van der Waals equation

A
  • b for bounce. repulsion term. Larger b = more reuplsion = greater pressure
  • a for attraction. Greater a, more attraction, less pressure.
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80
Q

Dalton’s law of partial pressures

A

Pi=xiPtotal

Ptotal=ΣPi

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81
Q

Hydrogen bonding

A
  • Weak interaction between a partially positive H and a partially negative atom (F/O/N)
  • Incrases boiling point
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82
Q

Van der Waals’ forces

A

(London dispersion forces)

  • Only significant for non-polar molecules
  • Result from induced and instantaneous dipoles
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83
Q

Induced vs. Instantaneous dipoles

A

Induced: when a polar molecule interacts with a non-polar molecule, then polar molecule induces a dipole in the non-polar molecule

Instantenous: Non-polar molecules have randomly fluctuating dipoles that tend to align with one another from one instant to the next

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84
Q

Dispersion forces are stronger or weaker for a larger molecule?

A

Stronger

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85
Q

Phase diagram

A
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86
Q

On a phase Diagram, what is the:

  • Triple Point?
  • Critical Point?
  • Critical Temp?
A
  • Triple Point: T and P at which all 3 phases of matter coexist at equilibrium
  • Critical Point: T and P at which liquids and gases become indistinguishable
  • Critical Temp: T above which you can no longer get a liquid no matter how much pressure you apply
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87
Q

What is a colligative property?

A

Depends on # of solute particles

88
Q

Van’t Hoff Factor? (for glucose, NaCl)

A

(i)

Convert concentration to reflect total + particles in solution.

(Glucose: i=1, NaCl: i=2)

89
Q

Raoult’s Law?

A

P = Xsolvent·P°solvent

ΔP = Xsolute·P°solvent

90
Q

Boiling point elevation equation

A

ΔTb = kb·m·i

  • kb = molal boiling point constant
  • m = molality
  • i = van’t hoff factor
91
Q

Freezing point depression equation

A

ΔTf = -kf·m·i

  • kf= molal freezing point constant
  • m = molality (mol solute/kg solvent)
  • i = van’t hoff factor
92
Q

Osmotic pressure equation

A

Π=MRT*i

  • Π = osmotic pressure
  • M = molarity
  • R = ideal gas constant
  • T = temp in K
  • Determines whether and in what direction osmosis will occur
  • Solvent goes from low to high Π
93
Q

What is a colloid?

A

Things are mixed at a “semi-molecular level” with solute aggregates that are tiny. Colloids will stay mixed until centrifuged.

94
Q

What is a suspension?

A

Mixed at a particle level, will not stay mixed.

95
Q

Henry’s Law

A

Psolute=k[solute]

  • Psolute= partial pressure of solute at surface
  • k = constant
  • [solute] is the solute concentration
96
Q
A
97
Q

What is the reaction rate?

A

Rate = -ΔReactant/ΔTime =ΔProduct/ΔTime

98
Q

What is the rate law for the following single-step reaction?

aA + bB → cC + dD

What if it is multi-step?

A

Single Step: Rate = k[A]a[B]b

Multi-Step: Rate =k[A]x[B]y

99
Q

What is the rate constant?

A

k in the rate law is the rate constant

The rate constant is an empirically determined value that changes with different reactions and reaction conditions

100
Q

Reaction order?

A

sum of all exponents of the concentration variables in the rate law

101
Q

Unimolecular? Bimolecular? Termolecular? Zero Order?

A

Unimolecular (1st order): r=k[A]

Bimolecular (2nd order): r=k[A]2, r=k[A][B]

Termolecular (3rd order): r=k[A]3, r=k[A]2[B], r=k[A][B][C]

Zero Order: r=k

102
Q

What is the rate-determining step?

A
  • Slowest step of a multi-step reaction
  • Rate of the whole rxn = rate of RDS
  • Rate law corresponds to components of RDS
103
Q

What is the activated complex?

A
  • What is present at the transition state
  • Bonds are just beginning to form and break
  • Peak of energy profile
  • Can’t be isolated
104
Q

Negative ΔH = ____

Positive ΔH = ____

A

Negative ΔH = exothermic

Positive ΔH = endothermic

105
Q

Arrhenius Equation

A

k=Ae-Ea/RT

  • Ea = activation energy
  • T = temp (K)
  • A = constant
106
Q

Thermodynamic vs kinetic product?

A
  • Kinetic product: lower activation energy, formed preferentially at a lower T
  • Thermodynamic product: lower/more favorable ΔG, higher T
107
Q

How do you know if a reaction is spontaneous?

A

ΔG is negative

108
Q

Equation relating G, H, T and S?

A

ΔG = ΔH - TΔS

109
Q

What factors favors a reaction? disfavor?

A
  • Favor: exothermic (-ΔH), increase entropy (+ΔS)
  • Disfavor: endothermic (+ΔH), decrease in entropy (-ΔS)
110
Q

Does a rxn occur faster with higher or lower activation energy?

A

lower activation energy

111
Q

What do catalysts do?

A
  • They speed up a reaction without getting itself used up
  • Enzymes are biological catalysts
  • Lower activation energy, speeds up forward and re verse reaction
  • Alter kinetics, not thermodynamics
  • Help system achieve equilibrium faster, does not alter position of equilibrium
  • Increase k, does not alter Keq
112
Q

What is the law of mass action?

A
  • Basis for equilibrium constant
  • Rate of a reaction depends only on the concentratino of the pertinent substances participating in the reaction
  • rforward=rreverse
113
Q

Keq

A
  • Keq = [C]c[D]d/[A]a[B]b
  • ΔG° = -RT ln (Keq)
114
Q

If Keq is greater than 1? equal to 1? less than 1?

A
  • > 1: more products are present
  • = 1: equilibrium is in the center
  • < 1: more reactants are present
115
Q

What is the reaction quotient, Q?

A
  • Same as Keq but can be used at any point, not just at quilibrium
  • If Q < Keq, rxn is still moving to the right
  • If Q = Keq, rxn is at equilibrium
  • If Q > Keq, rxn is moving back left
116
Q

LeChatelier’s principle?

A

If you knock a system off its equilibrium, it will readjust itself to reachieve equilibrium

117
Q

Relationship of the equilibrium constant and the standard free energy change

A

ΔG = ΔG° + RT ln Q

  • ΔG = 0 at equilibrium (Q = Keq)
  • ΔG° = -RT ln(Keq)
118
Q
A
119
Q

Hydroxide

A

OH-

120
Q

Chloride

A

Cl-

121
Q

Hypochlorite

A

ClO-

122
Q

Chlorite

A

ClO2-

123
Q

Chlorate

A

ClO3-

124
Q

Perchlorate

A

ClO4-

125
Q

Halide ion, hypohalide, etc.

A

X-, XO-, etc

126
Q

Carbonate

A

CO32-

127
Q

Hydrogen carbonate (Bicarbonate)

A

HCO3-

128
Q

Sulfate

A

SO42-

129
Q

Hydrogen sulfate (Bisulfate)

A

HSO4-

130
Q

Sulfite

A

SO32-

131
Q

Thiosulfate

A

S2O32-

132
Q

Nitrate

A

NO2-

133
Q

Nitrite

A

NO2-

134
Q

Phosphate

A

PO33-

135
Q

Hydrogen Phosphate

Dihydrogen Phosphate

A

HPO42-

H2PO4-

136
Q

Phosphite

A

PO33-

137
Q

Cyanide

A

CN-

138
Q

Thiocyanate

A

SCN-

139
Q

Peroxide

A

O22-

140
Q

Oxalate

A

C2O42-

141
Q

Acetate

A

C2H3O2-

142
Q

Chromate

Dichromate

A

CrO42-

Cr2O72-

143
Q

Permanganate

A

MnO4-

144
Q

Hydronium

A

H3O+

145
Q

Ammonium

A

NH4+

146
Q

What is hydration?

A
  • Solvation
  • Hydration is where water forms a shell around ions in solution
  • Oxygen atoms surround cations, Hydrogen atoms surround anions
147
Q

What is normality?

A

N = molarity of species that matter

1 M HCl = 1 N HCl

1 M H2SO4 = 2 N H2SO4

1 M H3PO4 = 3 N H3PO4

148
Q

What is Ksp?

A
  • Solubility product constant, the equilibrium expression
    • ex. AgCl (s) ↔ Ag+ (aq) + Cl- (aq)
    • Ksp for AgCl = [Ag+][Cl-]
  • ​Ksp value are found in a table
  • Higher values = more reaction products dominate in a saturated solution
149
Q

What is the solubility of MX2 for a given Ksp?

A
  1. MX2 ↔ M2+ + 2X-
  2. Ksp = [M2+][X-]2= [M2+][2M2+]2 = 4[M2+]3
  3. Solve for [M2+]. Solubility is the same thing as [M2+], because you used Q=Ksp for a saturated solution
150
Q

What is the common-ion effect?

A
  • AgCl (s) ↔ Ag+ (aq) + Cl- (aq)
  • If you add Cl- to the solution above, then less AgCL will dissolve
  • If you add NaCl to a saturated solution of AgCl, some AgCl would crash out of solution
151
Q

Complex ion formation

A
  • Metal+ + Lewis base: → Complex ion
  • M+ + L → M-Ln+
  • Keq=Kf (formation constant)
152
Q

Complex ion effect

A
  • Opposite of common ion effect
  • AgCl (s) ↔ Ag+ (aq) + Cl- (aq)
    M+ + Cl- ↔ M-Clncomplex ion.
  • When complex ion forms, Cl- is taken out, so more AgCl will dissolve
153
Q

Are acids more or less soluble in bases?

A
  • Acids are more soluble in bases.
  • HA → H+ + A-
  • Putting the above in a base will take out the H+, thus more HA will dissolve
154
Q

Are bases more soluble in other bases or acids?

A
  • Bases are more soluble in acids.
  • B + H+ → BH+
  • Putting this equation in an acid will add more H+, and thus drive more B to dissolve according to Le Chatelier’s principle
155
Q

Brondsted acid and base? (conj acid and base?)

A

H-Acid + Base-↔ Acid- + H-Base

  • Acid = proton donor
  • Base = proton acceptor
  • Conjugate base = acid after losing proton
  • Conjugate acid = base after gaining proton
156
Q

Ionization of water

A
  • Kw=[H+][OH-] = 10-14
  • At standard conditions, pure water has [H+]=10-7M and [OH-]=10-7
157
Q

Equations for pH, pOH

(What is acidic, basic, neutral?)

A

pH = -log[H+]

pOH = -log[OH-]

  • Acidic: pH < 7
  • Neutral: pH = 7
  • Basic: pH > 7
  • pH + pOH = 14
158
Q

Strong acids?

A
  • completely dissociate in solution (conj. base anion is highly stable)
159
Q

Strong Bases?

A
160
Q

Weak acids?

A
161
Q

Weak Bases?

A
162
Q

Do weak acids and bases dissociate more or less in a solution with added salt?

A

They will dissociate less in a solution with a salt.

163
Q

How to calculate the pH of the solutions of salt of weak acids?

A
  • CH3COO- + H2O ↔ CH3COOH + OH-
  • For M molar CH3COO-, start to abstract protons
    • [CH3COO-]= M-x
    • [CH3COOH] = [OH-] = x
  • Kb=Kw/Ka = [CH3COOH][OH-] / [CH3COO-] = x2/(M - x)
    • x is very small: Kb=x2/M (Solve for x)
    • pOH = -log(x)
    • pH = 14 - pOH
164
Q

Equilibrium constants Ka, Kb, pKa, pKb

A
  • H-Acid ↔ H+ + Acid-
    • Ka=[H+][Acid-]/[HAcid]
  • Base + H2O ↔ H-Base+ + OH-
    • Kb=[HBase+][OH-]/[Base]
  • KaKb=Kw=10-14
  • pKa=-logKa
  • pKb=-logKb
  • pKa+pKb=14
165
Q

pKa and pKb?

A
  • Ka*Kb=Kw=10-14
  • pKa = -logKa
  • pKb = -logKb
  • pKa + pKb = 14
166
Q

What is a buffer?

A
  • Solutions that resist changes in pH
  • Salts of weak acids and bases form buffer systems
  • Equilibrium between acidic species and a basic species
  • Acidic species donates protons to resist increase in pH
  • Basic species accepts protons to resist decreases in pH
  • Weak acid buffers: buffering capacity at pH = pKa
167
Q

Buffering capacity of weak acid?

A

pH = pKa

[Acid] = [conj base]

168
Q

Buffering capacity of a weak base?

A

pH = 14 - pKb

[base] = [conj acid]

169
Q

How do buffers influence titration curves?

A
  • Buffers make titration curves “flat” at the region where buffering occurs (point of inflection)
  • Inflection point at pH = pKa = 14 - pKb
  • Area around inflection point is the region where the solution has buffering capacity (usually pKa +/- 1)
170
Q

How do titration indicators work?

A

H-In ↔ H+ + In-

  • Behave like weak acids/bases
  • Present in small amount, doesn’t affect solution’s pH
  • When solution has a low pH (high [H+]), indicator is mostly H-In (one color)
  • When solution has a high pH, (low [H+]), indicator is mostly In-, another color
171
Q

Titration curve adding NaOH to HCl

A
172
Q

Titration Curve adding HCl to NaOH

A
173
Q

Titration curve adding NaOH to Acetic Acid

A
174
Q

Titration curve adding HCl to NH3

A
175
Q

What occurs at the point of inflection when titrating with a strong acid/weak base or strong base/ weak acid?

A
  • [acid] = [conj base]
  • [base] = [conj acid]
  • pH = pKa
  • [titrant] = 1/2 [weak acid/base]
176
Q

Titration curve adding NaOH to H2CO3

A
177
Q

What is reduction?

A
  • reduction in charge
  • decreased oxidation number
  • gain of electrons
178
Q

What is oxidation?

A
  • Increase in charge
  • Increased oxidation number
  • Losing electrons
179
Q

Galvanic/Voltaic Charge Flow

A
180
Q

Electrolytic Charge Flow

A
181
Q

Salt Electrolysis Charge Flow

A
182
Q

Electrolytic Cell

A
  • Requires potential/voltage input (battery)
    • Galvanic cell has either a resistor or voltmeter
  • potential/voltage inpot + the cell potential must be > 0 for rxn to occur
  • Cell potential is negative for electroytic cells
    • Cell potential is positive for galvanic/voltaic cell
183
Q

T/F: Electrolytic cell requires a battery?

A

T

184
Q

In both electrolytic and galvanic/voltaic cells:

___ is always the place where oxidation happens.

___ is always the place where reduction happens.

A

Anode is always the place where oxidation happens.

Cathode is always the place where reduction happens.

  • An Ox = ANode OXidation*
  • Red Cat = REDuction CAThode*
185
Q

___ shoots out electrons, ___ takes in electrons.

A

Anode shoots out electrons, Cathode takes in electrons.

186
Q

What are electrolytes?

A
  • Ions
  • Conduct electricity by motion of ions
  • Necessary for a circuit
187
Q

Faraday’s law relating amount of elements deposited (or gas liberated) at an electrode to current.

A

I=q/t

Faraday’s constant = coulombs of charge per mol of electron

F=q/n

188
Q

Faraday’s constant

A

F=q/n

F = total charge over total mols of electrons

189
Q

I*t = ?

A

It = nF

(Combine q = It and q = nF)

190
Q

Oxidation is ___ electrons while reduction is ___ electrons.

A

Oxidation is losing electrons, Reduction is gaining electrons.

Oil Rig: Oxidation Is Losing, Reduction is Gaining

191
Q

Galvanic/Voltaic cell half reactions

A
  • Oxidation ½ rxn: species loses e-
  • Reduction ½ rxn: species gains e-
192
Q

Reduction potentials; cell potentials

A
  • Reduction potential = potential of the reduction half reaction
  • Oxidation potential = reduction of the oxidation half rection (= neg reduciton potential)
  • Cell potential = Reduction potential + oxidation Potential
  • Want to make cell potential positive
193
Q

What direction do electrons flow in a galvanic/voltaic cell? In an electrolytic cell?

A
194
Q
  • Electrons always flow from anode to cathode (A to C)
  • Galvanic cell:
    • Species with highest oxidation potential (lowest reduction potential) will be anode
    • Species with highest reduction potential will be cathode
    • Cu is anode, Ag is cathode
    • Electrolytic cell is opposite (Cu is cathode, Ag is anode)
A
195
Q
  1. Isolated system?
  2. Closed system?
  3. Open system?
A
  1. Isolated system = no exchange of heat, work, or matter with surroundings
  2. Closed system = exchange of heat and work, but not matter with surroundings
  3. Open system = exchange of heat, work and matter with surroundings
196
Q

What is a state function?

A
  • Path-independent
  • ΔH (enthalpy), ΔS (entropy), ΔG (free energy change), ΔU (internal energy change)
  • Also called state quantity, or functions of state
197
Q

Endothermic vs Exothermic?

A
  • Endothermic = energy is taken up by rxn (+ ΔH)
  • Exothermic = energy is released by rxn (- ΔH)
198
Q

What is the standard heat of reaction, ΔHrxn?

Standard heat of formation, ΔHf?

A
  • ΔHrxn: The change in heat content for any reaction
  • ΔHf: The change in heat content in a formation reaction
199
Q

Hess’ law of heat summation?

A

ΔHrxn=Δ(ΔHf)=Σ(ΔHf)products-Σ(ΔHf)reactants​

200
Q

What is bond dissociation energy?

A

Energy required to break bonds

  • ΔHrxn= (bond dissociation energy for reactants) - (bond dissociation energy of bonds in products)
201
Q

What is heat capacity?

A
  • Amount of heat required to raise the temperature of something by 1 °C
  • Molar heat capacity = J/mol°C
  • Specific heat capacity = J/g°C
202
Q

1 calorie = ___ J ?

1 Calorie = ___ calorie?

A

1 calorie = 4.2 J

1 Calorie = 1000 calorie

203
Q

What is entropy?

A

A measurement of “disorder” in J/K

  • Entropy of gas > liquid > crystal states
204
Q

Zeroth Law of thermodynamics

A
  • Heat flows from hot objects to cold objects to achieve thermal equilibrium
205
Q

First Law of Thermodynamics

A
  • ΔE = q + w
  • Energy is conserved
  • Q is positive when heat is absorbed by system, and negative when heat leaks out of system
  • W is positive when work is done on the system (compression), negative when work is done by the system (expansion)
206
Q

Second law of thermodynamics

A

Things like to be in a state of higher entropy and disorder

ΔS ≥ q / T

207
Q

The universe as a whole is ___ in entropy

A

increasing

208
Q

If ΔS ≥ q / T

What is true for reversible processes?

Irreversible processes?

A

ΔS = q / T (reversible)

ΔS > q / T (irreversible)

209
Q
  1. Conduction
  2. Convection
  3. Radiation
A
  1. Heat transfer by direct contact
  2. Heat transfer by flowing current. Needs physical flow of matter
  3. Heat transfer by electromagnetic radiation. Does not need a medium
210
Q

Heat of fusion?

A

Energy input needed to melt something from solid to liquid at constant T

211
Q

On a PV diagram, what represents work?

A

Area

212
Q

What is an adiabatic process?

A

No heat exchange (Q=0, ΔE = W)

213
Q

What is an isothermal process?

A

No change in temperature. ΔT = 0

214
Q

What is isovolumetric (isochoric) process?

A

W=0, ΔE = q

215
Q

Calorimetry diagram

A
216
Q

Equation for q

A

q=mcΔT

(only works if no phase change is involved)

217
Q
A

MCAT Physical Sciences (2 decks)

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