Acid Base Redox Thermo

  1. What is a solution?
    A solution is a homogeneous mixture in which the molecules or ions of the components freely intermingle.
  2. What is a solvent?
    A solvent is the medium into which the solutes are mixed or dissolved.
  3. What is a solute?
    A solute is any substance dissolved in a solvent.
  4. What makes a compound in solution an electrolyte?
    The compounds ability, once it has dissolved in solution, to conduct electricity. It can either be a strong, weak, or non electrolyte.
  5. What makes a strong electrolyte?
    A substance that makes a strong electrolyte has the ability to complete dissociate into it's ions in solution.

    Examples: NaCl, HCl, NaOH, CaCl2
  6. What makes a weak electrolyte?
    A substance that makes a weak electrolyte will dissociate in solution, but then recombine. At some point this dissociation and recombining of the compound comes to an equilibrium.

    Examples: acetic acid (HC2H3O2)
  7. What makes a substance a non-electrolyte?
    A compound that is a non-electrolyte is a compound that will dissolve in a solute, but only dissolves into the moluecules of the compound, not into it's ions.

    Example: sugar, methanol (CH3OH), ethylene glycol (C2H4(OH)2)
  8. Ionic Equations
    Ionic equations are chemical equations showing the reactants and products in their ionic form in solution.
  9. Rules for ionic equations:
    1. All dissociated ions have to show the charge of that ion.
    2. All elements dissociated or not have to show their state. (ie. g=gas, l= liquid, s=solid, aq=aqueous)
    3. All elements and charges must be balanced.
    • Example:
    • -Chemical equation: Pb(NO3)2(aq) + 2KI(aq) --> PbI2(s) + 2KNO3(aq)
    • -Ionic equation: Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) --> PbI2(s) + 2K+(aq) + 2NO3-(aq)
  10. Net Ionic equation.
    -Are derived from the ionic equation by removing all spectator ions (ions that are on both sides of the equation).
    -Ionic equation: Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) --> PbI2(s) + 2K+(aq) + 2NO3-(aq)

    Remove spectator ions 2NO3-(aq) and 2K+(aq) (they cancel out becuase they are on both sides of the equation).

    -Net ionic equation: Pb2+(aq) + 2I-(aq) --> PbI2(s) (note: PbI2(s) is solid because it's insoluble: see solubility rules, table 4.1, pg 147)
  11. Definition of an acid:
    An acid is a substance that reacts with water to produce hydronium ions, H3O+ (or H+).
  12. Definition of a base:
    A base is a substance that produces hydroxide (OH-) ions in water.
  13. List of Strong Acids (stong electrolytes)
    • hydrochloric acid - HCl(aq)
    • hydrobromic acid - HBr(aq)
    • hydroiodic acid - HI(aq)
    • nitric acid - HNO3(aq)
    • sulfuric acid - H2SO4(aq)
    • chloric acid - HClO3(aq)
    • perchloric acid - HClO4(aq)
  14. Strong Bases (strong electrolytes)
    • Group IA hydroxides
    • -general formula: xOH (x=Li, Na, K, Cs, and Fr)

    • Group IIa hydroxides
    • -general formula: x(OH)2 (x= Ca, Ba, and Sr)
  15. Acid Nomenclature:
    Binary Acids - The binary compounds of hydrogen with many of the nonmetals are acidic, and in their aqueous solutions they are refered to as binary acids. To name you add the prefix hydro- and the suffix -ic to the stem of the nonmetal name, followed by the word acid.
    • Examples:
    • HCl(aq) - hydrochloric acid (as a compound not in solution HCl is a gas and is called hydrogen chloride.)
    • H2S(aq) - hydrosulfuric acid
  16. Acid Nomenclature:
    Oxoacids - Acids that contain hydrogen, oxygen, plus another element are called oxoacids. Oxoacids do not use the prefix hydro-. When there are 2 oxoacids, the one the larger number of Oxygens takes the suffix -ic and the one with the fewer Oxygens takes the suffix -ous.
    • Examples:
    • H2SO4 - sulfuric acid
    • H2SO3 - sulfurous acid

    • HNO3 - nitric acid
    • HNO2 - nitrous acid
  17. Acid Nomenclature:
    Oxoacids con't - The halogens can occur in as many as 4 oxoacids. The oxoacid with the most oxygens has teh prefix per-, and the one with the least has the prefix hypo-. These also follow the same rules for the suffix of oxoacids.
    • Examples:
    • HClO4 - perchloric acid
    • HClO3 - chloric acid
    • HClO2 - chlorous acid
    • HClO - hypochlorous acid
  18. Base Nomenclature - most bases are named a hydroxide of their metal.
    • Example:
    • NaOH - sodium hydroxide
  19. When acids and bases react they produce a salt and water.
    • Example:
    • 2HCl + Ca(OH)2 --> CaCl2 + 2H2O
  20. Molarity (M) = moles of solute/liters of solution

  21. Collectively, in a reaction where there is a transfer of elections it is called a oxidation-reduction, or redox, reaction. Where the substance loosing electron/s (e-) is being oxidized, and the substance gaining electrons is reduced.
    • Examples:
    • Na --> Na+ + e- (oxidation)
    • Cl2 + 2e- --> 2Cl- (reduction)
  22. Rules for assigning oxidation number:
    1. The oxidation number of any free element (and element not combined chemically with a different element) is 0, regardless of how complex its molecules are.
    2. The oxidation number for any simple, monatomic ion (Na+ or Cl-) is equal to the charge of the ion. The charge on a polyatomic ion can be viewed as teh net oxidation number of the ion.
    3. The sum of all the oxidation numbers of teh atoms in a moecule or polyatomic ion must equal the charge on the particle.
    4. In its compounds, flourine has an oxidation number of -1.
    5. In its compounds, hydrogen has an oxidation number of +1.
    6. In its compounds, oxygen has an oxidation number of -2.
  23. In the formation of hydrogen chloride:

    H2 + Cl2 --> 2HCl

    Hydrogen is oxidized to an oxidation number of +1, and chlorine is reduced to an oxidation number of -1.
  24. Balancing Redox Reactions: Ion-Electron Method
    • 1.Identify the half-reactions
    • 2.Balance each atom in the half reaction, saving H and O for last
    • 3.Balance O by adding 1 water molecule for each needed O
    • 4.Balance H by adding 1 H+ ion for each needed H
    • 5.Balance charges by adding electrons to the more positive side
    • 6.Find the least common multiple of electrons for the two half- reactions. Multiply each reaction by the factor needed to achieve the LCM of electronsAdd the half reactions, canceling like substances that appear on both sides
  25. Balancing Basic Reactions
    The simplest way to balance reactions in basic solution is to first balance them as if they were in acidic solution, then “convert” to basic solution
    • Additional Steps for Basic Solutions
    • 8) To both sides of the equation, add the same number of OH- ions as there are H+.
    • 9) Combine H+ and OH- to form H2O
    • 10) Cancel H2O molecules that are on both sides of the reaction.
  26. Activity Series
    An activity series arranges metals in their ease of being oxidized (see table 5.2, pg 191). An activity series can be used to predict reations.
    • Example:
    • Na + H2O --> NaOH + H2
    • Cu + NaOH --> no reaction
  27. 1st law of conservation of energy
    Energy cannot be created or destroy; it can only be changed from one form to another.
  28. Kenetic Energy (KE) - is the energy an object has when it's moving.
    • KE = 1/2mv2
    • m = mass (kg)
    • v = velocity (meter/second; m/s)
  29. Potential Energy (PE) - is the energy an object has that can be changed to kinetic energy, it can be thought of as stored energy.
  30. The SI unit for energy is the Joule (J).

    units of a joule are = kg(m/s)2
    kg - kilograms
    m - meter
    s - seconds
  31. Other units of energy:

    calorie (cal) 1cal = 4.184J
    kilocalorie (kcal) 1kcal = 1000cal = 4.184kJ
    dietary Calorie is actually a kilocalorie 1Cal = 1kcal (note capital "C")
  32. 3 types of systems:
    • 1. Open systems can gain or lose mass and energy across their boundaries.
    • 2. Closed systems can absorb or release energy, but not mass, across the boundary.
    • 3. Isolated systems cannot exchange matter or energy with their surroundings.
  33. Heat gain or loss equation:
    • q=C x delta t
    • q - heat
    • C - heat Capacity (units of heat Capacity are J/oC
    • delta t - temperature change (tfinal - tinitial)
  34. Heat Capacity depends on 2 factors: mass (m) and specific heat (s)
    C = m x s
    so if we substitute m x s for C in the heat gain/loss equation:

    q = m x s x delta t

    • specific heat (s) is an intensive property indendent of sample size
    • units of specific heat (s) are J/goC
  35. State Function;
    internal energy is a state function.
    • •A property whose value depends only on the present state of the system, not on the method or mechanism used to arrive at that state
    • •Position is a state function: both train and car travel to the same locations although their paths vary
    • •The actual distance traveled does vary with path
  36. Heat transfer at constant pressure is called enthalpy (delta H)

    delta H = q (at constant pressure)
    delta H = Hproducts - Hreactants

    (note: enthalpy is a state function)
  37. Endothermic reactions absorb energy. Therefore, the enthalpy change (delta H) has a positive number.

    Exothermic reactions lose energy, Therefore, the enthalpy change (delta H) has a negative number.
  38. An equation that also shows the value of delta Ho is called a thermochemical equation.

    N2(g) + 3H2(g) --> 2NH3(g) delta Ho = -92.38kJ
  39. Hess's Law
    The value of delta Ho for any reaction that can be written in steps equals the sum of the values of delta Ho of each of the individual steps.
  40. Rules for manipulating Thermochemical Equations
    • 1. When an equation is reversed - written in the opposite direction - the sign of delta Ho must also be reversed.
    • 2. Formulas conceled from both sides of an equation must be for the substance in identical physical states.
    • 3. If all the coefficients of an equation are multiplied or divided by the same factor, the value of delta Ho must likewise be multiplied or divided by that same factor.
Card Set
Acid Base Redox Thermo
Acids, Bases, Oxidation/Reduction (Redox) reactions and Thermochemistry