Chemistry 2

Question 1

compounds

Answer 1

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

Question 2

molecule

Answer 2

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

Question 3

molecular weight

Answer 3

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

Question 4

formula weight

Answer 4

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

Question 5

gram equivalent weight

Answer 5

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).

Question 6

Equivalents

Question 7

Normality

Answer 7

• 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.

Question 8

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

….of H2SO4?

Answer 8

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.

Question 9

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

Answer 9

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

Question 10

law of constant composition

Answer 10

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

Question 11

percent composition by mass

Question 12

combination reaction

Answer 12

two or more reactants forming one product

Ex:

Question 13

decomposition reaction

Answer 13

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

Question 14

single-displacement reaction

Answer 14

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)

Question 15

double-displacement reactions

(metathesis)

Answer 15

• 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

Question 16

Neutralization reactions

Answer 16

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

Question 17

Net ionic equations

Answer 17

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

vs.

Question 18

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

Answer 18

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

2. 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.

Question 19

Percent yield

Question 20

Calculate % Yield

Question 21

Balanced equations (order)

Answer 21

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

Question 22

mechanism

Answer 22

series of steps by which a reaction proceeds

Question 23

What is the rate expression

For the general reaction aA + bB → cC + dD

?

Answer 23

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

Question 24

rate law

Answer 24

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.

Question 25

overall order of the reaction

Answer 25

sum of x + y (+ z… )

Question 26

Qc \< Keq, Δ G \< 0

Answer 26

reaction proceeds in forward direction

Question 27

Qc = Keq, Δ G = 0,

Answer 27

reaction is in dynamic equilibrium

Question 28

Qc \> Keq, Δ G \> 0

Answer 28

reaction proceeds in reverse direction

Question 29

The Law of Mass Action

Answer 29

Question 30

The Reaction Quotient

Answer 30

Question 31

For example, consider the formation of HCl from H2 and Cl2. The overall reaction is ## Footnote  What does the energy diagram look like?

Answer 31

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.

Question 32

What is ΔH for an exothermic reaction?

Answer 32

– Δ H = exothermic = heat given off ## Footnote

Question 33

What is Δ H for an endothermic reaction?

Answer 33

+Δ H = endothermic = heat absorbed

Question 34

How does a catalyst effect the energy of a reaction?

Answer 34


Question 35

Consider the reaction:  It will shift to the RIGHT if...

Answer 35

* A or B added * C removed * pressure increased or volume reduced * temperature reduced

Question 36

Consider the reaction:  It will shift to the LEFT if...

Answer 36

* C added * A or B removed * volume increased or pressure reduced * temperature increased

Question 37

What is an "isolated system"?

Answer 37

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

Question 38

What is a "closed system"?

Answer 38

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

Question 39

What is an "Open System"?

Answer 39

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

Question 40

Adiabatic processes

Answer 40

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

Question 41

isobaric processes

Answer 41

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.

Question 42

isothermal processes

Answer 42

occur when the system's temperature is constant.

Question 43

standard temperature and pressure (STP)

Answer 43

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

Question 44

standard conditions

Answer 44

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

Question 45

standard state

Answer 45

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

Question 46

Joule to cal conversion

Answer 46

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

Question 47

energy transfer to or from a system

Answer 47


Question 48

positive Δ H

Answer 48

Endothermic

Question 49

negative Δ H

Answer 49

Exothermic

Question 50

bomb calorimeter (decompression vessel)

Answer 50

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, Q_calorimeter is 0. So, Δ U_system + Δ U_surroundings = Δ U_calorimeter = Q_calorimeter− W_calorimeter = 0. Therefore, Δ U_system = − Δ U_surroundings. Because no work is done, q_system = − q_surroundings, and m_steelc_steelΔ T + m_oxygenc_oxygenΔ T = − m_waterc_waterΔ T.

Question 51

Specific heat

Answer 51

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

Question 52

Enthalpy change

Answer 52

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

Question 53

standard enthalpy of formation

Answer 53

, 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.

Question 54

standard heat of a reaction

Answer 54

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. 

Question 55

Hess's law

Answer 55

enthalpy changes of reactions are additive 

Question 56

bond dissociation energies

Answer 56

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. 

Question 57

standard heat of combustion

Answer 57

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)

Question 58

second law of thermodynamics:

Answer 58

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

Question 59

Entropy

Answer 59

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

Question 60

Entropy of Universe

Answer 60


Question 61

Is Entropy pathway dependent?

Answer 61

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: ## Footnote 

Question 62

Δ S° _rxn,

Answer 62


Question 63

Gibbs Free Energy

Answer 63

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

Question 64

When is a reaction spontaneous with respect to Δ G ?

Answer 64

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

Question 65

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

Answer 65


Question 66

When is Δ G temperature dependent?

Answer 66

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

Question 67

Gibbs Free Energy

Answer 67


Question 68

Phase Diagram

Answer 68


Question 69

triple point

Answer 69

The point at which the three phase boundaries meet

Question 70

critical point

Answer 70

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

Question 71

The phase diagram for a mixture of two or more components

Answer 71


Question 72

Raoult's Law

Answer 72


Question 73

Boiling Point Elevation

Answer 73


Question 74

Freezing Point Depression

Answer 74


Question 75

Osmotic Pressure formula

Answer 75


Question 76

Osmotic Pressure Diagram

Answer 76

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

Question 77

Dilution

Answer 77

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.

Question 78

Solubility Product Constant

Answer 78

For the dissolution of an ionic solid,  The units of K_sp depend on the solid. The concentration of the ions is generally in units of molarity (mol/L). This formula is used to determine K_sp or molar solubility. It is most often used in the context of the ion product.

Question 79

Ion Product (I.P.)

Answer 79

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.

Question 80

When will a solution precipitate?

Answer 80

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.

Question 81

What is the K_sp expression for the following, respectively * **AgCl** * **ZnF₂ ** * **Fe(OH)₃** *Also, what is the condtion of its use?*

Answer 81

* 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)*

Question 82

Solubility rules

Answer 82

Ksp \>\> 1: soluble Ksp \<\< 1: insoluble **_Always soluble: _** * **Group 1 metal salts** * **Amonium salts** * **Nitrate salts**

Question 83

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

Answer 83

CaCO₃ --\>/\<-- Ca²⁺ + CO₃^2- Careful! Ask yourself: do any of the reactants/products react with an acid or a base... **YES! CO₃^2-** **CO₃^2-** + H₂O --\>/\<-- HCO₃^- + OH^- If OH- increases (bc pH is increased), then CO32- will be increased in effect. If CO₃^2- 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 K_sp does not change, and the molar solubility decreases.

Question 84

Kinetic Molecular Theory of Gases

Answer 84

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.

Question 85

Average Molecular Speeds

Answer 85

*KE= (1/2)mv² = (3/2)kT*

Question 86

Root mean squared speed (*u_rms*)

Answer 86

***u_rms = √(3RT)/M***

Question 87

Graham's Law

Answer 87

**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)

Question 88

Density of an ideal gas

Answer 88

d= (*m*/V₂)

Question 89

Volume of a gas under nonstandard conditions

Answer 89

V₂ = V₁(P₁/P₂)(T₂/T₁)

Question 90

Boyle's Law

Answer 90

P₁V₁=P₂V₂ | (PV=k, k is a contant)

Question 91

Charles' Law

Answer 91

(V/T)=k or (V₁/T₁)=(V₂/T₂) (V/T)=(nR/P)= contant

Question 92

Dalton's Law

Answer 92

(of partial pressures, relating to mole fraction) **P_A=P_T*X*_A** P_T=P_A + P_B + P_C + ... *X_A= n_A/n_T* (moles of A/total moles of all gases)

Question 93

Deviations of Ideal Gases Due to Temperature

Answer 93

* 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](V-nb) = nRT

Question 94

Deviations of Ideal Gases Due to PRESSURE

Answer 94

* 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](V-nb) = nRT

Question 95

Ideal gas law using density and mass

Answer 95

* n=m/M=(mass in grams)/(Molar mass)* * d=m/V* **PV=(mRT)/M** **T=(PVM)/mR** **T=PM/dR**

Question 96

Crystalline

Answer 96

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

Question 97

Ionic solids

Answer 97

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

Question 98

Gibs Function (phase changes)

Answer 98

Themodynamic equilibria: ΔG= G (gas) - G (solid) = 0

Question 99

Colligative properties

Answer 99

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

Question 100

Raoult's Law

Answer 100

(ideal solutions) P_A=X_AP^o_A ![]()

Question 101

Solubility Rules

Answer 101

## Footnote 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.

Question 102

Ionic nomenclature

Answer 102

## Footnote 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.

Question 103

Electrolytes

Answer 103

* 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)

Question 104

Solubility Product Constant

Answer 104

For the dissolution of an ionic solid, A_mB_n ↔ mAⁿ⁺ + nB^m- ***K_sp = [Aⁿ⁺]^m [B^m-]ⁿ***  **\*only dependent on temperature and solvent [K_SP 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. 

Question 105

Ion Product (I.P.)

Answer 105

For the dissolution of an ionic solid, A_mB_n ↔ mAⁿ⁺ + nB^m- ***I.P. = [Aⁿ⁺]^m [B^m-]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

Question 106

Saturation and Precipiation (Q_sp vs. K_sp)

Answer 106

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*.

Question 107

Common K_sp values for slightly soluble salts

Answer 107

* Formula: MX3 Ksp=24x⁴ * Formula: MX2 Ksp=4x³ * Formula: MX Ksp=x²

Question 108

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.

Answer 108

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)*

Question 109

What forces act in solvation?

Answer 109

**_Solvation=t_**he 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

Question 110

K_w

Answer 110

water dissociation equilibrium constant at 298 K_w = 10^-14 = [H₃O+] [OH-] as T increases, so does K_w, as a reult of endothermic nature of the auto-ionization reaction