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Chemistry 2 Flashcards

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1
Q

compounds

A

pure substances composed of two or more elements in a fixed proportion

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2
Q

molecule

A
  • combination of two or more atoms held together by covalent bonds.
  • smallest unit of a compound that displays the identifying properties of that compound.
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3
Q

molecular weight

A

sum of the atomic weights of all the atoms in a molecule, and its units are atomic mass units (amu).

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4
Q

formula weight

A

adding up the atomic weights of the constituent ions according to its empirical formula

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5
Q

gram equivalent weight

A

where n is the number of protons, hydroxide ions, electrons, or monovalent ions “ produced” or “ consumed” per molecule of the compound in the reaction.

For example, you would need 49 grams of H2SO4 (molar mass = 98 g/mol) to produce one equivalent of hydrogen ions, because each molecule of H2SO4 can donate two hydrogen ions (n = 2).

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6
Q

Equivalents

A
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7
Q

Normality

A
  • measure of concentration.
  • units for normality are equivalents/liter
  • where n is the number of protons, hydroxide ions, electrons, or monovalent ions “ produced” or “ consumed” per molecule of the compound in the reaction.
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8
Q

What is the molarity of a 1 N acid solution consisting of dissolved HCl?

….of H2SO4?

A

In a 1 N acid solution consisting of dissolved HCl, the molarity of HCl is 1 M because HCl is a monoprotic acid,

but if the dissolved acid is H2SO4, then the molarity of H2SO4 in a 1 N acid solution is 0.5 M, because H2SO4 is a diprotic acid.

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9
Q

Will one mole of HCl will not completely neutralize one mole of Ca(OH)2?

A

NO

one mole of HCl will not completely neutralize one mole of Ca(OH)2, because one mole of HCl will donate one equivalent of acid but Ca(OH)2 will donate two equivalents of base

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10
Q

law of constant composition

A

any pure sample of a given compound will contain the same elements in an identical mass ratio

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11
Q

percent composition by mass

A
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12
Q

combination reaction

A

two or more reactants forming one product

Ex:

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13
Q

decomposition reaction

A

A single compound reactant breaks down into two or more products, usually as a result of heating or electrolysis.

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14
Q

single-displacement reaction

A

an atom (or ion) of one compound is replaced by an atom of another element

(Single-displacement reactions are often further classified as redox reactions)

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15
Q

double-displacement reactions

(metathesis)

A
  • elements from two different compounds swap places with each other (hence, the name double-displacement) to form two new compounds
  • occurs when one of the products is removed from the solution as a precipitate or gas or when two of the original species combine to form a weak electrolyte that remains undissociated in solution
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16
Q

Neutralization reactions

A

specific type of double-displacement reaction in which an acid reacts with a base to produce a salt

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17
Q

Net ionic equations

A

Net ionic equations list only the elements important for demonstrating the actual reaction that occurs during a displacement reaction

vs.

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18
Q

For problems involving the determination of the limiting reactant, you must keep in mind two principles

A
  1. All comparisons of reactants must be done in units of moles. Gram-to-gram comparisons will be useless and maybe even misleading.

  1. It is not the absolute mole quantities of the reactants that determine which reactant is the limiting reactant. Rather, the rate at which the reactants are consumed (the stoichiometric ratios of the reactants) combined with the absolute mole quantities determines which reactant is the limiting reactant.
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19
Q

Percent yield

A
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20
Q

Calculate % Yield

A
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21
Q

Balanced equations (order)

A

Balanced equations are determined using the following steps in order:

  1. Balancing the non-hydrogen and non-oxygen atoms
  2. Balancing the oxygen atoms, if present
  3. Balancing the hydrogen atoms, if present
  4. Balancing charge when necessary
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22
Q

mechanism

A

series of steps by which a reaction proceeds

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23
Q

What is the rate expression

For the general reaction aA + bB → cC + dD

?

A

Rate is expressed in the units of moles per liter per second (mol/L· s) or molarity per second (M/s).

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24
Q

rate law

A

where k is the reaction rate coefficient or rate constant. Rate is always measured in units of concentration over time; that is, molarity/second. The exponents x and y (or x, y, and z, if there are three reactants, etc.) are called the orders of the reaction: x is the order with respect to reactant A, and y is order with respect to reactant B.

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25
Q

overall order of the reaction

A

sum of x + y (+ z… )

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26
Q

Qc < Keq, Δ G < 0

A

reaction proceeds in forward direction

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27
Q

Qc = Keq, Δ G = 0,

A

reaction is in dynamic equilibrium

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28
Q

Qc > Keq, Δ G > 0

A

reaction proceeds in reverse direction

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29
Q

The Law of Mass Action

A
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30
Q

The Reaction Quotient

A
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31
Q

For example, consider the formation of HCl from H2 and Cl2. The overall reaction is

What does the energy diagram look like?

A

Figure 5.2 shows that the reaction is exothermic. The potential energy of the products is less than the potential energy of the reactants; heat is evolved, and the enthalpy change of the reaction is negative.

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32
Q

What is ΔH for an exothermic reaction?

A

– Δ H = exothermic = heat given off

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33
Q

What is Δ H for an endothermic reaction?

A

+Δ H = endothermic = heat absorbed

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34
Q

How does a catalyst effect the energy of a reaction?

A
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35
Q

Consider the reaction:

It will shift to the RIGHT if…

A
  • A or B added
  • C removed
  • pressure increased or volume reduced
  • temperature reduced
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36
Q

Consider the reaction:

It will shift to the LEFT if…

A
  • C added
  • A or B removed
  • volume increased or pressure reduced
  • temperature increased
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37
Q

What is an “isolated system”?

A

The system cannot exchange energy (heat and work) or matter with the surroundings; for example, an insulated bomb calorimeter.

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38
Q

What is a “closed system”?

A

The system can exchange energy (heat and work) but not matter with the surroundings; for example, a steam radiator.

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39
Q

What is an “Open System”?

A

The system can exchange both energy (heat and work) and matter with the surroundings; for example, a pot of boiling water.

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40
Q

Adiabatic processes

A

occur when no heat is exchanged between the system and the environment; thus, the heat content of the system is constant throughout the process.

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41
Q

isobaric processes

A

occur when the pressure of the system is constant. Isothermal and isobaric processes are common, because it is usually easy to control temperature and pressure.

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42
Q

isothermal processes

A

occur when the system’s temperature is constant.

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43
Q

standard temperature and pressure (STP)

A

temperature is 0° C (273 K) and pressure is 1 atm.

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44
Q

standard conditions

A

The standard conditions are defined as 25° C (298 K) and 1 atm.

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45
Q

standard state

A

Under standard conditions, the most stable form of a substance

Common: H2 (g), H2O (l), NaCl (s), O2 (g), and C (s) (graphite) are the most stable forms of these substances under standard conditions

46
Q

Joule to cal conversion

A

The unit of heat is the unit of energy: joule (J) or calorie (cal), for which 1 cal =4.184 J.

47
Q

energy transfer to or from a system

A
48
Q

positive Δ H

A

Endothermic

49
Q

negative Δ H

A

Exothermic

50
Q

bomb calorimeter (decompression vessel)

A

A sample of matter, typically a hydrocarbon, is placed in the steel decomposition vessel, which is then filled with almost pure O2 gas. The decomposition vessel is then placed in an insulated container holding a known mass of water. The contents of the decomposition vessel are ignited by an electric ignition mechanism. The material combusts (burns) in the presence of the oxygen, and the heat that evolves in the combustion is the heat of the reaction. Because W = PΔ V, no work is done in an isovolumic (Δ V = 0) process, so Wcalorimeter = 0. Furthermore, because of the insulation, the whole calorimeter can be considered isolated from the rest of the universe, so we can identify the “ system” as the sample plus the oxygen and steel vessel, and the surroundings as the water. Because no heat is exchanged between the calorimeter and the rest of the universe, Qcalorimeter is 0.

So, Δ Usystem + Δ Usurroundings = Δ Ucalorimeter = Qcalorimeter− Wcalorimeter = 0.

Therefore, Δ Usystem = − Δ Usurroundings.

Because no work is done, qsystem = − qsurroundings, and msteelcsteelΔ T + moxygencoxygenΔ T = − mwatercwaterΔ T.

51
Q

Specific heat

A

This equation is used to determine temperature after heat transfer. It is most commonly encountered in a calorimetry problem

52
Q

Enthalpy change

A

This equation is used to determine enthalpy change. It is most commonly used to determine whether a reaction is endothermic or exothermic.

53
Q

standard enthalpy of formation

A

, Δ H° f, is the enthalpy change that would occur if one mole of a compound in its standard state were formed directly from its elements in their respective standard states. Remember that standard state is the most stable physical state of an element or compound at 298 K and 1 atm. Note that Δ H° f of an element in its standard state, by definition, is zero. The Δ H° f of most known substances are tabulated.

54
Q

standard heat of a reaction

A

Δ H° rxn, is the hypothetical enthalpy change that would occur if the reaction were carried out under standard conditions. What this means is that all reactants must be in their standard states and all products must be in their standard states.

55
Q

Hess’s law

A

enthalpy changes of reactions are additive

56
Q

bond dissociation energies

A

Bond dissociation energy is the average energy that is required to break a particular type of bond between atoms in the gas phase (remember, bond dissociation is an endothermic process). Bond dissociation energy is given as kJ/mol of bonds broken.

57
Q

standard heat of combustion

A

Δ H° comb. Because measurements of enthalpy change require a reaction to be spontaneous and fast, combustion reactions are the ideal process for such measurements.

usually in precense of O2 but keep in mind there are others: F2 (for example)

58
Q

second law of thermodynamics:

A

Energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so

59
Q

Entropy

A

measure of the spontaneous dispersal of energy at a specific temperature: how much energy is spread out, or how widely spread out energy becomes

60
Q

Entropy of Universe

A
61
Q

Is Entropy pathway dependent?

A

Entropy is a state function, so a change in entropy from one equilibrium state to another is pathway-independent and only depends upon the difference in entropies of the final and initial states:

62
Q

Δ S° rxn,

A
63
Q

Gibbs Free Energy

A

This state function is a combination of the two that we’ve just examined: enthalpy and entropy.

64
Q

When is a reaction spontaneous with respect to Δ G ?

A
  1. If Δ G is negative, the reaction is spontaneous.
  2. If Δ G is positive, the reaction is nonspontaneous.
  3. If Δ G is zero, the system is in a state of equilibrium; thus Δ H = TΔ S
65
Q

How do the signs on Δ H and Δ S and temperature affect the spontaneity?

A
66
Q

When is Δ G temperature dependent?

A

Δ G is temperature-dependent when Δ H and Δ S have the same sign.

67
Q

Gibbs Free Energy

A
68
Q

Phase Diagram

A
69
Q

triple point

A

The point at which the three phase boundaries meet

70
Q

critical point

A

The phase boundary between the liquid and gas phases, however, terminates at this point

71
Q

The phase diagram for a mixture of two or more components

A
72
Q

Raoult’s Law

A
73
Q

Boiling Point Elevation

A
74
Q

Freezing Point Depression

A
75
Q

Osmotic Pressure formula

A
76
Q

Osmotic Pressure Diagram

A

One compartment contains pure water, while the other contains water with dissolved solute. The membrane allows water but not solute to pass through. Because substances tend to flow, or diffuse, from higher to lower concentration (which results in an increase in entropy), water will diffuse from the compartment containing pure water into the compartment containing the water-solute mixture. This net flow will cause the water level in the compartment containing the solution to rise above the level in the compartment containing pure water

77
Q

Dilution

A

A solution is diluted when solvent is added to a solution of high concentration to produce a solution of lower concentration. The concentration of a solution after dilution can be conveniently determined using the equation:

where M is molarity, V is volume, and the subscripts i and f refer to the initial and final values, respectively.

78
Q

Solubility Product Constant

A

For the dissolution of an ionic solid,

The units of Ksp depend on the solid. The concentration of the ions is generally in units of molarity (mol/L).

This formula is used to determine Ksp or molar solubility. It is most often used in the context of the ion product.

79
Q

Ion Product (I.P.)

A

For the dissolution of an ionic solid,


The units of I.P. depend on the solid. The concentration of the ions is generally in units of molarity (mol/L).

I.P. is compared to Ksp to determine the behavior of a solution.

80
Q

When will a solution precipitate?

A

If the solution is supersaturated, Qsp > Ksp, precipitation will occur.

If the solution is unsaturated, Qsp < Ksp, the solute will continue to dissolve.

If the solution is saturated, Qsp = Ksp, then the solution is at equilibrium.

81
Q

What is the Ksp expression for the following, respectively

  • AgCl
  • ZnF2
  • Fe(OH)3

Also, what is the condtion of its use?

A
  • AgCl: Ksp = x2
  • ZnF2: Ksp = 4x3
  • Fe(OH)3: Ksp = 27x4

Condition: can only use if ther is not a common ion (ie: Cl- is in solution already from a different souce)

82
Q

Solubility rules

A

Ksp >> 1: soluble

Ksp << 1: insoluble

_Always soluble: _

  • Group 1 metal salts
  • Amonium salts
  • Nitrate salts
83
Q

If the pH is increased, how dos the Ksp and molar solublity of CaCO2 (s) change?

A

CaCO3 –>/<– Ca2+ + CO32-

Careful! Ask yourself: do any of the reactants/products react with an acid or a base…

YES! CO32-

CO32- + H2O –>/<– HCO3- + OH-

If OH- increases (bc pH is increased), then CO32- will be increased in effect.

If CO32- concentration is increased, then then the molar solubility (S, x) will be decreased in effect.

NOTE: Ksp does not change (constant for a given rxn at a given temp). So answer: The Ksp does not change, and the molar solubility decreases.

84
Q

Kinetic Molecular Theory of Gases

A

Assumptions:

  1. Gases are made up of particles whose volumes are negligible compared ot the container volume.
  2. Gas aoms or molcules exhibit no intermolecular attractions or repulsions.
  3. Gas particles are in coninuous, random motion, undergoing collisions with other particles and the container walls.
  4. Collisions between any 2 gas particles are elastic, (conservation of kinetic energy and momentum)
  5. Average kinetic energy of gas partilces is proportional to the absolute temperature (in Kelvin) of the gas; it is the same for all gases at a given temperature, irrespective of their chemical identity or atomic mass.
85
Q

Average Molecular Speeds

A

KE= (1/2)mv2 = (3/2)kT

86
Q

Root mean squared speed (urms)

A

urms = √(3RT)/M

87
Q

Graham’s Law

A

r1/r2 = √(M2/M1)

*under isothermal and isobaric conditions, the rates at which 2 gases diffuse is inversely proportional to the square root of their molar masses

*(ie: a gas that has a molar mass four times that of another gas will travel half as fast as the lighter gas)

88
Q

Density of an ideal gas

A

d= (m/V2)

89
Q

Volume of a gas under nonstandard conditions

A

V2 = V1(P1/P2)(T2/T1)

90
Q

Boyle’s Law

A

P1V1=P2V2

(PV=k, k is a contant)

91
Q

Charles’ Law

A

(V/T)=k or

(V1/T1)=(V2/T2)

(V/T)=(nR/P)= contant

92
Q

Dalton’s Law

A

(of partial pressures, relating to mole fraction)

PA=PTXA

PT=PA + PB + PC + …

XA= nA/nT (moles of A/total moles of all gases)

93
Q

Deviations of Ideal Gases Due to Temperature

A
  • intermolecular forces become more significant as temperatures decrease
  • reduced Temp: gas is closer to condensation point (same as boiling point, volume is less than ideal gas bc of intermolecular attraction
  • closer to gas’s bp: less ideal
  • Van Der Waals Correction: a=attraction
  • P+(n2a)*/V2 = nRT
94
Q

Deviations of Ideal Gases Due to PRESSURE

A
  • intermolecular forces become more significant as pressure increase
  • moderately high P (a few hundred atmospheres): gas’s volume is less than ideal gas bc of intermolecular attraction
  • extremely high P: size of particles is relatively large compared to distance between them, causing gas to take up larger volume than predicted by ideal
  • Van Der Waals Correction: a=attraction
  • P+(n2a)*/V2 = nRT
95
Q

Ideal gas law using density and mass

A
  • n=m/M=(mass in grams)/(Molar mass)*
  • d=m/V*

PV=(mRT)/M

T=(PVM)/mR

T=PM/dR

96
Q

Crystalline

A

Crystalline structures allow for a balance of both attractive and repulsive forces to minimize energy.

97
Q

Ionic solids

A

aggregates of positively and negatively charged ions that repeat according to defined patterns of alternating cations and anions

high melting points, high boiling points, and poor electrical conductivity in the solid state but high conductivity in the molten state or in aqueous solution

98
Q

Gibs Function (phase changes)

A

Themodynamic equilibria:

ΔG= G (gas) - G (solid) = 0

99
Q

Colligative properties

A

physical properties of solutions that are dependent on the concentration of dissolved particles but not on the chemical identity of the dissolve particles

  • vapor pressure depression
  • boiling point elevation
  • freezing point depression
  • osmotic pressure
100
Q

Raoult’s Law

A

(ideal solutions)

PA=XAPoA

101
Q

Solubility Rules

A

one infallible solubility rule:
All sodium salts are completely soluble,
and all nitrate salts are completely soluble.

  1. All salts of alkali metals are water soluble.
  2. All salts of the ammonium ion (NH4+) are water soluble.
  3. All chlorides, bromides, and iodides are water soluble, with the exceptions of those formed with Ag+, Pb2+, and Hg22+.
  4. All salts of the sulfate ion (SO42− ) are water soluble, with the exceptions of those formed with Ca2+, Sr2+, Ba2+, and Pb2+.
  5. All metal oxides are insoluble, with the exception of those formed with the alkali metals and CaO, SrO, and BaO, all of which hydrolyze to form solutions of the corresponding metal hydroxides.
  6. All hydroxides are insoluble, with the exception of those formed with the alkali metals and Ca2+, Sr2+, and Ba2+.
  7. All carbonates (CO32− ), phosphates (PO43− ), sulfides (S2− ), and sulfites (SO32− ) are insoluble, with the exception of those formed with the alkali metals and ammonium.
102
Q

Ionic nomenclature

A

  1. For elements (usually metals) that can form more than one positive ion, the charge is indicated by a Roman numeral in parentheses following the name of the element.
  2. An older but still commonly used method is to add the endings -ous or -ic to the root of the Latin name of the element to represent the ions with lesser or greater charge, respectively.
  3. Monatomic anions are named by dropping the ending of the name of the element and adding -ide.
  4. Many polyatomic anions contain oxygen and are therefore called oxyanions. When an element forms two oxyanions, the name of the one with less oxygen ends in -ite and the one with more oxygen ends in -ate.
  5. When the series of oxyanions contains four oxyanions, prefixes are also used. Hypo- and per- are used to indicate less oxygen and more oxygen, respectively.
  6. Polyatomic anions often gain one or more H+ ions to form anions of lower charge. The resulting ions are named by adding the word hydrogen or dihydrogen to the front of the anion’s name. An older method uses the prefix bi- to indicate the addition of a single hydrogen ion.
103
Q

Electrolytes

A
  • A solute is considered a strong electrolyte if it dissociates completely into its constituent ions. (NaCl)
  • A weak electrolyte, on the other hand, ionizes or hydrolyzes incompletely in aqueous solution, and only some of the solute is dissolved into its ion constituents. (weak acids, weak bases)
  • Non-electrolytes: Many compounds do not ionize at all in aqueous solution, retaining their molecular structure in solution, which usually limits their solubility. (ie: glucose–dissolves but ring structure does not let it dissociate)
104
Q

Solubility Product Constant

A

For the dissolution of an ionic solid,

AmBn ↔ mAn+ + nBm-

Ksp = [An+]m [Bm-]n

*only dependent on temperature and solvent [KSP ISNOT AFFECTED BY COMMON ION EFFECT]

The units of Ksp depend on the solid. The concentration of the ions is generally in units of molarity (mol/L).

This formula is used to determine Ksp or molar solubility. It is most often used in the context of the ion product.

105
Q

Ion Product (I.P.)

A

For the dissolution of an ionic solid,

AmBn ↔ mAn+ + nBm-

I.P. = [An+]m [Bm-]n

  • is analogous to the reaction quotient Q for chemical reactions.
  • The ion product equation has the same form as the equation for the solubility product constant.
  • The difference: concentrations are of the ionic constituents at that given moment in time.
  • IS affected by common ions (acts as a stress on the system and can shift equilibrium without changing equilibrium constant
106
Q

Saturation and Precipiation (Qsp vs. Ksp)

A

If the solution is supersaturated, Qsp > Ksp, precipitation will occur.

If the solution is unsaturated, Qsp < Ksp, the solute will continue to dissolve.

If the solution is saturated, Qsp = Ksp, then the solution is at equilibrium.

107
Q

Common Ksp values for slightly soluble salts

A
  • Formula: MX3
    Ksp=24x4
  • Formula: MX2
    Ksp=4x3
  • Formula: MX
    Ksp=x2
108
Q

What types of mixtures would have a higher vapor pressure than predicted by Raoults Law?

  • a) Two liquds of that have different properties
  • b) Two liquds of that have similar properties.
A

Two liquds of that have different properties.

Mixtures that have stronger solute-solute an solvent-solvent interactions than solute-solbent interactions are less likely to want to be together –> more likely to evaporate

ie: hexane (hydrophobic) and ethannol (hydrophilic)

109
Q

What forces act in solvation?

A

Solvation=the formation of a solvent shell around a solute molecule through electrostatic interactions.

Forces: ionic bonds, dipole-dipole interactions, van er Waals, hydrogen bonds

  • endothermic process (+∆H), +∆S generally

_*special case:_ enthalpy change tends to be negative (-∆H) with dissolution of gases bc gases have no intermolecular attractions; thus E is released when new intermolecular attractions are formed favored by low temps

110
Q

Kw

A

water dissociation equilibrium constant

at 298 Kw = 10-14 = [H3O+] [OH-]

as T increases, so does Kw, as a reult of endothermic nature of the auto-ionization reaction

111
Q
A

mcat (24 decks)

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