Chapter Exam #3 Study Guide
∙ Acidbase reactions (where an acid reacts with a base) involve water as a reactant or product.
∙ An acid is a substance that produces H+ ions when dissolved in water. HX H+ (aq) + X (aq)
∙ A base is a substance that produces OH ions when dissolved in water. MOH M+ (aq) + OH (aq)
∙ Acidic solutions arise when certain covalent Hcontaining molecules dissociate into ions in water.
∙ Strong acids AND strong bases dissociate completely into ions when placed in water.
∙ Weak acids AND weak bases dissociate very little into ions.
∙ Strong acids, such as HNO3 and H2SO4, and weak acids, such as HF and H3PO4, have one or more H atoms as part of their structure.
∙ Strong bases have either OH or O2 as part of their structure.
∙ Weak bases, such as ammonia (NH3) do not contain OH ions, but they all have an electron pair on N.
∙ The key event in aqueous reactions between a strong acid and a strong base is that an H+ ion from the acid and an OH ion from the base form a water molecule. ∙ Acidbase reactions occur through the electrostatic attraction of ions and their removal from solution as the product. If you want to learn more check out What is Calorie/Kcal?
∙ The ionic compound that results from the reaction of an acid and base is called a salt.
∙ Acidbase reactions are metathesis (doubledisplacement) reactions. ∙ An acid is a molecule (or ion) that donates a proton.
∙ A base is a molecule (or ion) that accepts a proton.
∙ H3O+ ion acts as the acid and donates a proton to OH ion, which acts as the base and accepts it.
∙ Titration: The known concentration of one solution is used to determine the unknown concentration of another.
∙ Equivalence point: Occurs when the amount (mol) of H+ ions in the original volume of acid has reacted with the same amount (mol) of OH ions from the buret.
∙ End point: Occurs when a tiny excess of OH ions changes the indicator permanently to its basic color.
∙ The amount of base needed to reach the end point is the same as the amount needed to reach the equivalence point.
∙ Oxidationreduction (redox) reaction: Net movement of electrons from one reactant to another.
∙ Ionic compounds: transfer of electrons.
∙ Covalent compounds: shift (sharing) of electrons. We also discuss several other topics like What is cognitive therapy?
∙ Oxidation is the loss of electrons.
∙ Reduction is the gain of electrons.
∙ Example: Formation of MgO
Oxidation (electron loss by Mg): Mg Mg2+ + 2e–
Reduction (electron gain by O2): ½ O2 + 2e– O2–
The oxidizing agent is the species doing the oxidizing (causing electron loss).
The reducing agent is the species doing the reducing (causing electron gain).
In MgO, O2 oxidizes Mg by taking electrons that Mg gives up. Mg reduces O2 by providing the electrons that O2 takes, so Mg is the reducing agent, and O2 is the oxidizing agent.
The oxidizing agent is reduced, the reducing agent is oxidized. ∙ Oxidation numbers:
1. For Group 1(A): O.N. = +1 in all compounds
2. For Group 2(A): O.N. = +2 in all compounds
3. For hydrogen: O.N. = +1 in combination with nonmetals
= –1 in combination with metals and boron
4. For fluorine: O.N. = –1 in all compounds
5. For oxygen: O.N. = –1 in all peroxides
= –2 in all other compounds (except with F).
6. For Group 7(A): O.N. = –1 in combination with metals, nonmetals (except O), and other halogens lower in the group
∙ Transferred electrons are never free because the reducing agent loses electrons and the oxidizing agent gains them simultaneously.
∙ Atoms occur as an element on one side of an equation and as part of a compound on the other. Don't forget about the age old question of what are the difference between atoms and molecules?
∙ In a combination reaction, two or more reactants form a compound. X + Y Z
∙ Metal and nonmetal form an ionic compound. The metal is the reducing agent and the nonmetal is the oxidizing agent.
∙ Two nonmetals form a covalent compound.
∙ Nearly every nonmetal reacts with O2 to form a covalent oxide. ∙ In a decomposition redox reaction, a compound forms two or more products, at least one of which is an element.
Z X + Y
∙ In doubledisplacement (metathesis) reactions, atoms of two compounds exchange places.
AB + CD AD + CB
In solution, singledisplacement reactions occur when an atom of one element displaces the ion of another.
∙ The most reactive metals displace H2 from liquid water. Group 1(A) metals and Ca, Sr, and Ba from Group 2(A) displace H2 from water.
∙ Thermodynamics: The study of energy and its transformations.
∙ Thermochemistry: A branch of thermodynamics that deals with heat in chemical and physical change.
∙ System: The part of the universe we are focusing on (in a reaction). ∙ Surroundings: Everything else (in a reaction). We also discuss several other topics like What is Edgar Allan Poe's theory of the unity of effect?
∙ Internal Energy (E): Sum of all potential and kinetic energy in a system. ∙ ∆E is the difference between internal energy after the change (Efinal) and before the change (Einitial).
∙ Final state minus the initial state:
∆E = Efinal – Einitial = Eproducts – Ereactants
∙ A change in the energy of a system must be accompanied by an equal and opposite change in the energy of the surroundings.
∙ By releasing some energy in a transfer to the surroundings:
Efinal < Einitial so ∆E < 0
∙ By absorbing some energy in a transfer from the surroundings:
Efinal > Einitial so ∆E > 0
∙ Heat: Thermal energy (symbolized by q) is the energy transferred as a result of a difference in temperature.
∙ Work: (symbolized by w) the energy transferred when an object is moved by a force.
∙ The total change in a system’s internal energy is the sum of the energy transferred as heat and/or work: Don't forget about the age old question of what is the first law of thermodynamics?
∆E = q + w
∙ Energy transferred to the system is positive because the system ends up with more energy.
∙ Energy transferred from the system is negative because the system ends up with less energy.
∙ Heat flowing out of a system: heat is released so q is negative and ∆E is negative. ∙ Heat flowing into a system: heat is absorbed so q is positive and ∆E is positive.
∙ Law of Conservation of Energy – First Law of Thermodynamics: The total energy of the universe is constant.
∆Euniverse = ∆Esystem – ∆Esurroundings = 0 We also discuss several other topics like What is the thorndike’s two laws?
∙ Joule (J): The SI unit of energy
1 J = 1 kg • m2/s2
∙ Calorie (cal): Quantity of energy needed to raise the temperature of 1 g of water by 1 ˚C.
1 cal = 4.184 J or 1 J = 1/4.184 cal = 0.2390 cal
∙ British thermal unit: Quantity of energy required to raise the temperature of 1 lb of water by 1 ˚F.
1 BTU = 1055 J
∙ ∆E does not depend on how the change takes place, but only on the difference between the final and initial states.
∙ Pressurevolume work (PV work): The mechanical work done when the volume of the system changes in the presence of an external pressure (P). w = –P∆V
∙ At constant pressure, enthalpy (H) is defined as the internal energy plus the product of the pressure and volume.
H = E + PV
∙ Change in enthalpy (∆H): The change in internal energy plus the product of the pressure, which is constant, and the change in volume (∆V).
∆H = ∆E + P∆V
∙ Exothermic and endothermic process:
∆H = Hfinal – Hinitial = Hproducts – Hreactants
∙ Exothermic: Releases heat and results in a decrease in the enthalpy of a system: Hproducts < Hreactants so ∆H < 0
∙ Endothermic: Absorbs hear and results in an increase in the enthalpy of a system: Hproducts > Hreactants so ∆H > 0
∙ q/∆T = constant
∙ Heat capacity: [refer to above equation] the quantity of heat required to change its temperature by 1 K.
∙ Specific heat capacity: The quantity of heat required to change the temperature of 1 gram of a substance or material by 1 K.
Specific heat capacity (c) = q/mass x ∆T
∙ Molar heat capacity: The quantity of heat required to change the temperature of 1 mole of a substance by 1 K.
Molar heat capacity (C) = q/amount (mol) x ∆T
∙ Calorimeter: A device used to measure the heat released (or absorbed) by a physical or chemical process.
∙ Finding the specific heat capacity of a solid:
csolid = cH2O x massH2O x ∆TH2O/masssolid x ∆Tsolid
∙ A thermochemical equation is a balanced equation that includes the enthalpy change of the reaction (∆H).
∙ Hess’s Law: The enthalpy change if an overall process is the sum of the enthalpy changes of its individual steps:
∆Hoverall = ∆H1 + ∆H2 … + ∆Hn
∙ Standard states:
o For a gas, the standard state is 1 atm* and ideal behavior
o For a substance in an aqueous solution, the standard state is 1 M concentration.
o For a pure substance, the standard state is usually the most stable form of the substance at 1 atm and the temperature of interest (the temperature is usually 25 ˚C (298 K).
∙ Hydronium ion (H3O+) forms H bonds to several other water molecules. ∙ Neutralization: Occurs when an acid and a base react to form H2O. ∙ The stronger the acid, the higher [H3O+] is at equilibrium, and the larger the value of Ka.
∙ Strong acids: HCl, HBr, and HI.
1. Oxoacids in which the number of O atoms equals or exceeds by one the number of ionizable protons, such as HClO, HNO2, and H3PO4
2. Carboxylic acids; such as CH3COOH and C6H5COOH.
∙ Strong bases
1. M2O or MOH
2. MO or M(OH)2
∙ Weak bases
1. Ammonia (NH3)
2. Amines such as CH3CH2NH2
∙ Water dissociates very little into ions: this process is called autoionization. ∙ Higher [H3O+] lower [OH–] and vice versa.
∙ In an acidic solution: [H3O+] > [OH–]
∙ In a basic solution: [H3O+] < [OH–]
∙ In a neutral solution: [H3O+] = [OH–]
∙ pH = –log [H3O+]
∙ The higher the pH, the lower the [H3O+]. An acidic solution has a lower pH (higher [H3O+]) than a basic solution.
∙ pH of an acidic solution < 7.00
∙ pH of a neutral solution = 7.00
∙ pH of a basic solution > 7.00
∙ pOH = –log [OH–]
The halfreaction method divides the overall redox reaction into oxidation and reduction halfreactions.
Example of the halfreaction method using Cr2O72– (aq) + I– (aq) Cr3+ (aq) + I2 (s): Step #1: Divide the reaction into halfreactions.
Step #2: Balance atoms and charges in each halfreaction.
a. Balance atoms other than O and H.
b. Balance O atoms by adding H2O molecules.
Cr2O72– 2Cr3+ + 7H2O
c. Balance H atoms by adding H+ ions.
14H+ + Cr2O72– 2Cr3+ + 7H2O
d. Balance the charge by adding electrons.
6e– + 14H+ + Cr2O72– 2Cr3+ + 7H2O [reduction]
∙ The relative strengths of the bonds in reactants and products determine whether heat is released or absorbed in a chemical reaction.
∙ Kinetic energy: Molecules’ movements through space and their rotations and vibrations.
∙ Potential energy: Phase changes and changes in the attraction between vibrating atoms.
∙ A certain quantity of heat is absorbed (∆H˚ > 0) to break the reactant bonds and form separate atoms.
∙ A different quantity of heat is then released (∆H˚ < 0) when the atoms form product bonds.