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Flashcards in Chapter 4 Deck (46):
1

Limiting Reagent

- When chemical reaction involves 2 or more reactants, often 1 reactant is of limited quantity and the others are in excess
- the limiting reagent determines or limits the amount of product formed in a reaction
- to determine the limiting reagent, calculate the amount of product (in g. or mol.) produced from each reactant -> the lower quantity is the limiting reagent
- amount not reacted = total amount - amount reacted

2

Theoretical Yield

- amount of product that would be formed if the reaction went to completion (value with the limiting reagent)
- based on stoichiometry of the reaction and the ideal conditions in which starting material is completely consumed and no losses in work-up procedure
- calculated based on balanced chemical equation

3

Percent Yield

- under experimental conditions, actual yield is almost always less than amount predicted by theory (theoretical yield)
- equation = (actual yield ÷ theoretical yield) x 100
- yields above about 90% = very good
- yields above about 75% = good
- yields above about 60% = modest
- yields above about 30% = poor

*yields can also be above 100% when an extraneous chemical from outside of the reaction has found its way into the yield

4

Solution

- homogeneous mixture of two or more substances
- solutions can be solids, liquids, gases, or gas/liquid mixtures

examples:
gas solution = dry air
solution of ions = sea water
gas/liquid solution = carbonated water
solid solution = brass (homogeneous metal alloy with different metal components)

5

Solute

- does not react with solvent
ie.) dissolved table salt (NaCl ions) does not react chemically with water
- present in smaller amount
- uniformly dispersed in another substance/ solvent
- aqueous solutions are transparent

6

Solvent

- present in larger amount

example:
25 g. sugar dissolved in 100 g. water
- sugar is smaller amount (=solute)
- water is larger amount (=solvent)

7

Aqueous solutions

- solvent = water
- solutes = dissolved ionic or polar substances
- homogeneous mixture of dissolved ions in water
- water binds to both positive and negative ions in an aqueous solution. when ionic substance dissolves in water/becomes solvated, its cation and anion are surrounded by a sheath of water molecules. Dissolved ions can move about freely.

8

Molarity (M)

- concentration of solution= amount of solute dissolved in a given volume of solution
- Molarity = mol. solute ÷ vol. solution (solute + solvent)
- molarity refers to amount of solute per liter of solution, NOT per liter of solvent
- When making solutions of a given molarity, dissolve the solute in a volume of solvent smaller than the desired volume of solution. Then add solvent until the final volume of solution is reached.

9

Unimolecular Solutes vs. Ion Solutes

- unimolecular = do not form ions when dissolved in water; number of solute particles they form in a solution remain unchanged before and after the solvation process

example: glucose and urea

- ionic solutes: soluble in water; produce ions whose concentrations can be a multiple of the compounds form which they were derived
- to calculate the molarity of a specific ion solute, one must take into account the ionic formula of the compound from which the ion was derived

example:
1 mol, CaCl2 (s) --(H2O)--> 1 mol. CaCl2 (aq)
or
1 mol Ca 2+ (aq)
2 mol Cl - (aq)
= 3 mol. solute particles total

10

Dilution

- made when a solution of lower concentration is prepared form a more concentration one
- solution concentration inversely proportional to volume of solution
- C1 V1 = C2 V2

example:
when water (solvent) is added to a can of orange juice concentrate (solute), the amount of orange juice present does not change BUT its concentration decreased, because total volume of solution increases

11

Dilution problem:
1.) a stock solution is prepared by dissolving 10.8 g. (NH4)2 SO4 (molar mass: 132.15 g/mol) in enough water to make 100 mL (solution 1). A 10.0-mL sample of this stock solution (solution 1) is then place in a 50-mL volumetric flask and diluted to the mark with water to make solution 2.
a.) What is the molarity of solution 2?
b.) (NH4)2 SO4 is an ionic compound which dissolves in water to form an aqueous solution of ions. What is the molarity of the NH4+ ions in solution 2?

a.) 0.163 M (NH4)2 SO4
b.) 0.326 M NH4+

12

Electrolytes

- ionic or molecular substances which, when dissolved in water to form solution, produce ions that can conduct an electric current
- commonly exists as solutions of acids, bases, or salts

13

Strong Electrolytes

- ionic compounds that dissociate 100% in solution to form ions only
- conduct electricity strongly

example:
HCl (aq) -> H+ (aq) + Cl- (aq)
0% 100% 100%
(strong acid = dissociates completely)

14

Weak Electrolytes

- Ionic compounds that dissociate only partially (< 100%) in solution to form some ions
- conduct electricity weakly

example:
HF (aq) -> H+ (aq) + F- (aq)

15

Non-Electrolytes

- polar, covalent compounds that dissolve in aqueous solution BUT do not form ions (i.e. no dissociation)
- DO NOT conduct electricity
- many non electrolyte compounds have Oxygen atoms linked to Hydrogen atoms, forming hydroxyl group (-OH). Presence of PH groups in nonionic substances allow H-bonding with water to occur, thereby increasing the solubility of these substances in aqueous solution.

examples:
Methanol (CH3OH), Ethanol (CH3CH2OH)

16

Selective Precipitation Problem:
A solution contains one or more of the following ions: Ag+, Ca 2+, and Cu 2+.
a.) When sodium chloride is added to the solution, no precipitate occurs.
b.) When sodium sulfate is added to the solution, a white precipitate occurs, which is then filtered off.
c.) To the remaining solution, sodium carbonate is added, producing a precipitate.

A.) Which ions were present in the original solution?
B.) Write a balanced net ionic equation for the formation of each of the precipitates observed?
C.) Write a balanced molecular equation with NO3- for the formation of each of the precipitates observed.

A.) Ca 2+ and Cu 2+

B.) Ca 2+ (aq) + SO4 2- (aq) -> CaSO4 (s)
Cu 2+ (aq) + CO3 2- (aq) -> CuCO3 (s)

C.) Ca(NO3)2 (aq) + Na2SO4 (aq) -> CaSO4 (s) + 2NaNO3 (aq)
Cu(NO3)2 (aq) + Na2CO3 (aq) -> 2NaNO3 (aq) + CuCO3 (s)

17

Molecular Equation

- equation that shows the complete neutral formulas for every compound in the reaction

example:
AgNO3 (aq) + KCl (aq) -> AgCl (s) + KNO3 (aq)

18

Net ionic equation

- equation that shows only the species that actually participate in the reaction

example:
Ag+ (aq) + Cl- (aq) -> AgCl (s)

19

Spectator Ions

- ions that do not participate in the reaction

example:
K+ and NO3-

20

Acid

- substance that, when dissolved in water, increases the concentration of hydrogen ions (H+), in the solution
- H+ ions give acids a sour taste
- turns blue litmus paper red
- acid acts as an ionic compound that dissociates in solution to form a cation and an anion

examples:
vinegar, citrus fruits, wine, asprin
HCl (aq) -> H+ (aq) + Cl- (aq)

*H+ ions that are formed in aqueous solutions DO NOT exist as free H+ ions. Rather, they are attracted to polar H2O molecules to form hydronium ions (H3O+)

example:
HCl (aq) + H2O (l) -> Cl- (aq) + H3O (aq)

21

Base

- substance that, when dissolved in water, increases the concentration of hydroxide ions, OH-, in the solution
- OH- ion give base a bitter taste
- soapy and slippery feel
- turns red litmus paper blue

examples:
baking soda, antacids, soap, tonic water
KOH (aq) -> K+ (aq) + OH- (aq)

*some bases do not carry a hydroxide (OH-) in their formula but nevertheless produce OH- ions upon reacting with H2) like ammonia and their derivative compounds

examples:
NH3 (aq) + H2O (l) -> NH4+ (aq) + OH- (aq)
CH3--NH3+ (aq) + OH-(aq)

22

Strong Acids

HCl (aq) Hydrochloric acid
HBr (aq) Hydrobromic acid
HI (aq) Hydroiodic acid
HNO3 Nitric acid
HClO4 Perchloric acid
H2SO4 Sulfuric acid (first H+ only)

23

Strong Bases

LiOH Lithium hydroxide
NaOH Sodium hydroxide
KOH Potassium hydroxide
Ba(OH)2 Barium hydroxide
Sr(OH)2 Strontium hydroxide

24

Acid-Base Neutralization Reaction

acid + base -> salt + water

example:
2HCl (aq) + Ba(OH)2 (aq) -> BaCl2 (aq) + 2H2O (l)

*reaction between weak acid with strong base = must keep reactant in molecular form in net equation because it does not dissociate 100%

example:
ME: CH3COOH (aq) + NaOH (aq) -> CH3COO-Na+ (aq) + H2O (l)
NIE: CH3COOH (aq) + OH- (aq) -> CH3COO- (aq) + H2O (l)

*reaction between strong acid with weak base

example:
ME: 3 HNO3 (aq) + Fe(OH)3(s) -> Fe(NO3)3 (aq) + 3 H2O (l)
NIE: 3 HNO3 (aq) + Fe(OH)3 (s) -> Fe3+ (aq) + 3 H2O (l)

25

Acid and Base Concentration Formulas

[H3O+] = [OH-] = 1.0 × 10^-7 M
Kw = [H3O+] [OH-] = 1.0 × 10^-14
from equation get:
[H3O+] = (1.0 × 10^-14) ÷ [OH-]
[OH-] = (1.0 × 10^-14) ÷ [H3O+]

26

pH

pH = -log [H3O+]
pOH = -log [OH-]

[H3O+] = 10^-pH
[OH-] = 10^-pOH

pH + pOH = 14

*sig figs?: the number of decimal places in the pH value is the same as the number of sig figs in the coefficient of [H3O+]. The number on the left of the decimal point in the pH value is the power of 10.

example:
[H3O+] = 1.0 × 10^-2 M pH = 2.00

27

Titration

- performed to determine the concentration of an unknown acid or base or some other substance in a solution
- titrant = acid/base of known concentration
- analyte = acid/base of unknown concentration
- analyte (say acid) with acid-base color indicator like phenolphthalein and titrated with measured volume (say base) of known concentration
- in example, acid is gradually neutralized by added base
- at titration endpoint, an approximately equivalent amount of base had been added to neutralize all of the acid
- indicator colorless at acidic pH (pH < 7) but turns pink when the solution PH is > 8, which signals the titration endpoint
- at endpoint of titration: mol. OH- ions added = mol. H+ ions present in sample

28

Standardizing an Acid Solution by Titration Problem
1.) A 0.263-g sample of Na2CO3, a base, requires 28.35 mL of aqueous HCl for titration to the endpoint. What is the molarity of the HCl solution?

0.175 M HCl

29

Determining the Molar Mass by Titration Problem
2.) A monoprotic acid (HA) reacts with NaOH according to the following equation:
HA (aq) + NaOH (aq) -> NaA (aq) + H2O (l)
Calculate the molar mass of HA if 0.856 g of the acid requires 30.08 mL of 0.323 M NaOH.

88.1 g/mol HA

30

Reactions of Acids with Carbonates/Bicarbonates

when acid added to carbonate (CO3 2-) or bicarbonate (HCO3-), products are CO2 gas, water, and salt

31

Reactions of Acids with Sulfites/Bisulfites

when acid added to sulfites (SO3 2-) or bisulfite (HSO3 -), the products are SO2 gas, water, and salt

32

Reactions of Acids with Sulfides/Bisulfides

when acid added to sulfide (S 2-) or bisulfide (HS-), the products are H2S gas and salt

33

Reactions of Bases with Ammonium Compounds

when base added to solution of NH4+ containing compound, the products are water, salt, and NH3 gas/ammonia gas

34

Reactions of Acids with Metals

acids react with certain metals to produce a salt and H2 gas

examples of metals that react with acids include:
group 1A metals (Na, K)
group 2A metals (Mg, Ca)
group 3A metals (Al)
group 4A metals (Gn)
transition metals (Zn, Fe)

35

Oxidation-Reduction (Redox) Reactions

- reaction in which electrons transfer from one atom to another

36

Oxidation

- loss of one or more electrons by an atom (Lose electron, oxidized)
- becomes more positive/ less negative
- loss of Hydrogen
- gain of Oxygen

37

Reduction

- gain of one or more electrons by an atom (gain electron, reduced)
- becomes more negative/ less positive
- gain of Hydrogen
- lose of Oxygen

38

Guidelines for Determining Oxidation Numbers

1.) atoms in their pure elemental state has oxidation number of 0
example: Fe^0 (s), Mg^0 (s), Cu ^0 (s)

2.) Naturally occurring diatomic molecules have oxidation number equal to 0
example: H2, O2, N2, Cl2, I2, Br2, F2

3.) Oxidation number of mono atomic ion is equal to charge
example: Na+ = +1 oxidation number, Cl- = -1 oxidation number

4.) oxidation number of Oxygen is -2 in most compounds EXCEPT peroxide (O2 2- = -1 oxidation number)

5.) Oxidation number of Hydrogen is +1 in most compounds EXCEPT hydride (oxidation number of -2)

6.) Oxidation number of Fluorine is -1 for ALL CASES

7.) Oxidation number of other halogens is usually -1 EXCEPT when combined with Oxygen or Fluorine. In these cases, oxidation number is variable.

8.) Oxidation number of atoms in neutral compound must add up to zero. For polyatomic ion, oxidation numbers of atoms add up to charge on the ion

39

Oxidizing Agent

- substance that causes oxidation to occur and is itself reduced

40

Reducing Agent

- substance that causes reduction to take place and is itself oxidized

41

Oxidation Reduction Problems
1.) What reagent can perform the following conversions?
X + CH4 -> CO2
a. acid b. base c. oxidizing agent d. reducing agent e. H20
2.) Which are redox reactions?
a. HCl (aq) + KOH (aq) -> KCl (aq) + H2O (l)
b. AgNO3 (aq) + HCl (aq) -> HNO3 (aq) + AgCl (s)
c. 3 C2H6O + 2 Cr2O7 2- + 16 H3O + -> 3 C2H4O2 + 4 Cr 3+ + 27H2O
d. 2 C4H10 (g) + 13 O2 (g) -> 8 CO2 (g) + 10H2O (l)

1.) C. Oxidizing agent
2.) C. and D.

42

Synthesis/ Combination Reaction

- two or more reactants combine to yield a single product
A + B -> AB

43

Decomposition

- one reactant splits into two or more products
AB -> A + B

44

Single Replacement/ Single Displacement Reaction

- one element replaces another element
A + BC -> AC + B

45

Double Replacement/ Double Displacement Reaction

- two elements replace each other
AB + CD -> AD + CB

46

Combustion Reaction

- reaction of a substance with oxygen, usually with the rapid release of heat to produce a flame. When a hydrocarbon is involved, the products are always CO2 and H2O (+energy)