AP Chemistry

Pages
- Class
  Laboratory report rubric
  Notes
  1. Chemical foundations
  2. Atoms, molecules, and ions
  3. Stoichiometry
  4. Types of chemical reactions and solution stoichiometry
  5. Gases
  6. Thermochemistry
  7. Atomic structure and periodicity
  8. Bonding: general concepts
  9. Covalent bonding: orbitals
  10. Liquids and solids
  11. Properties of solutions
  12. Chemical kinetics
  13. Chemical equilibrium
  14. Acids and bases
  15. Acid-base equilibria
  16. Solubility and complex ion equilibria
  17. Spontaneity, entropy, free energy
  18. Electrochemistry
  Drug unit
  Basics
  Analgesics
  Antacids
  Anesthetics
  Depressants
  Stimulants
  Antibiotics
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  Mind-altering drugs
- Textbook (incomplete)
  1. Chemical foundations
  2. Atoms, molecules, and ions
  3. Stoichiometry
  4. Types of chemical reactions and solution stoichiometry
  5. Gases
  6. Thermochemistry
  7. Atomic structure and periodicity
  8. Bonding: general concepts
- CED
  1. Atomic structure and properties
  2. Compound structure and properties
  3. Properties of substances and mixtures
  4. Chemical reactions
  5. Kinetics
  6. Thermochemistry
  7. Equilibrium
  8. Acids and bases
  9. Thermodynamics and electrochemistry

AP Chemistry

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4. Types of chemical reactions and solution stoichiometry

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4. Types of chemical reactions and solution stoichiometry

Water, the common solvent

aqueous solutions: solutions in which water is the solvent (dissolving medium)

water can dissolve many different substances

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H₂O molecules

O—H bonds are covalent bonds (electron sharing between oxygen and hydrogen atoms

electrons not shared equally

oxygen has greater attraction for electrons than hydrogen

oxygen atoms gain slight excess of negative charge; hydrogen atoms become slightly positive

polar molecule: unequal charge distribution

δ (delta) indicates partial charge (+ or -)

water dissolves ionic substances

hydration: when “positive ends” of water molecules are attracted to negatively-charged anions and “negative ends” are attracted to positively-charged cations

hydration of ions tends to cause salt to “fall apart” or dissolve into the water

strong forces among the ions are replaced by strong water-ion interactions

when ionic substances (salts) dissolve in water, they break up into the individual cations and aniosn

e.g. ammonium nitrate dissolves in water:

NH₄NO₃ —(H₂Ol)→ NH₄⁺ (aq) + NO₃⁻ (aq)

solubility of ionic substances in water differs greatly

differences typically depend on:

relative attractions of the ions for each other (holding the solids together)

attractions of the ions for water molecules (which cause the solid to disperse in water)

when an ionic compound does dissolve in water, the ions become hydrated and are dispersed

water also dissolves many nonionic substances

e.g. ethanol (C₂H₅OH)

contains a polar O—H bond like in water; very compatible

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many substances do not dissolve in water

e.g. animal fat: fat molecules are nonpolar and do not interact effectively with polar water molecules

“like dissolves like”: polar and ionic substances are expected to be more soluble in water than nonpolar substances

Strong and weak electrolytes

solute: substance dissolved in liquid to form a solution

solvent: dissolving medium in a solution

electrical conductivity: ability of a solution to conduct an electric current

electrolyte: substance that when dissolved in water produces a solution that can conduct electricity

solutions with strong electrolytes can conduct a current very efficiently

solutions with weak electrolytes conduct only a small current

solutions with nonelectrolytes permit no current to flow

basis for conductivity properties of solutions

first identified by Svante Arrhenius (1859-1927)

then a Swedish graduate student in physics

came to believe that the conductivity of solutions arose from the presence of ions

initially scorned by majority of scientific establishment

late 1890s: atoms found to contain charged particles; ionic theory became widely accepted

extent to which a solution can conduct an electric current depends directly on the number of ions present

Strong electrolytes

substances that are completely ionized when they are dissolved in water

soluble salts: contain array of cations and ions that separate and become hydrated when the salt dissolves

strong acids: strong electrolytes that dissociate (ionize) completely in aqueous solution

Arrhenius: found that when HCl, HNO₃, and H₂SO₄ were dissolved in water, they behaved as strong electrolytes

acid: substance that produces H⁺ ions (protons) when it is dissolved in water

ionization of an acid: HA (aq) + H₂O (l) → H₃O⁺ (aq) + A⁻ (aq)

hydrochloric acid, nitric acid, sulfuric acid are aqueous and should be written with (aq) (HCl (aq), HNO₃ (aq), H₂SO₄ (aq)) although they often appear without them

a strong acid is one that dissolves completely into its ions; virtually no HCl molecules exist in aqueous solutions

sulfuric acid: formula indicates that it can produce two H⁺ ions per molecule, but only the first H⁺ ion is completely dissociated

second can be pulled off under certain conditions

aqueous solution of H₂SO₄ contains mostly H⁺ ions and HSO₄⁻ ions

strong bases: soluble ionic compounds contain the hydroxide ion (OH⁻)

when dissolved in water, cations and OH⁻ ions separate and move independently

most common basic solutions: when sodium hydroxide (NaOH) or potassium hydroxide (KOH) is dissolved in water to produce ions

Weak electrolytes

substances that exhibit a small degree of ionization in water

produce relatively few ions when dissolved in water

most common: weak acids and weak bases

weak acids: any acid that dissociates (ionizes) only to a slight extent in aqueous solutions

formulas

often written with acidic hydrogen atom(s) (any that will produce H⁺ ions in solution) listed first

any nonacidic hydrogens written later

example: HC₂H₃O₂ (acetic acid) indicates one acidic and three nonacidic hydrogen atoms

disassociation reaction for acetic acid in water: HC₂H₃O₂ (aq) + H₂O (l) ⇌ H₃O⁺ (aq) + C₂H₃O₂⁻ (aq)

unlike strong acids, very small numbers of molecules (~1% for acetic acid) dissociate in aqueous solutions at typical concentrations

weak bases: resulting solution is a weak electrolyte

most common weak base: ammonia (NH₃)

ammonia dissolved in water: NH₃ (aq) + H₂O (l) ⇌ NH₄⁺ (aq) + OH⁻ (aq)

basic because OH⁻ ions produced

Nonelectrolytes

substances that dissolve in water but do not produce any ions

e.g. ethanol: entire C₂H₅OH molecules dispersed into water

molecules do not break up into ions

resulting solution does not conduct electric current

e.g. table sugar/sucrose (C₁₂H₂₂O₁₁): very soluble in water but produces no ions when dissolved

Composition of solutions

to perform stoichiometric calculations when two solutions are mixed, you must know:

the nature of the reaction (depends on the exact forms the chemicals take when dissolved)

the amounts of chemicals present in the solutions (usually expressed as concentrations)

most commonly used expression of concentration: molarity (M)

moles of solute per volume of solution in liters

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Dilution

dilution: adding water to achieve the molarity desired for a particular solution

common acids are purchased as concentrated solutions and diluted as needed

calculating

determining how much water must be added to an amount of stock solution to achieve a solution of the desired concentration

moles of solute after dilution = moles of solute before dilution

steps

determine the number of moles of substance in the final solution by multiplying the volume by the molarity

use a volume with a known molarity and known amount of substance to solve for the volume

example: 500. mL of 1.00 M acetic acid (HC₂H₃O₂) from what volume of a 17.4-M stock solution of acetic acid?

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to make 500 mL of a 1.00-M acetic acid solution, take 28.7 mL of 17.4 M acetic acid and dilute to a total volume of 500 mL with distilled water

dilution procedure

pipet: device for accurately measuring and transferring a given volume of solution

volumetric (transfer) pipets: come in specific sizes (e.g. 5 mL, 10 mL, 25 mL)

measuring pipets: used to measure volumes for which a volumetric pipet is not available

volumetric flask

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M: molarity

V: volume

Types of chemical reactions

millions of possible chemical reactions

grouped into classes

precipitation reactions

acid-base reactions

oxidation-reduction reactions

Precipitation reactions

precipitation reaction (double displacement reaction): when two solutions are mixed and a precipitate forms (insoluble substance/solid)

predicting identity of precipitate

example: K₂CrO₄(aq) added to Ba(NO₃)₂(aq) form yellow solid

think about what products are possible

in virtually every case, when a solid containing ions dissolves in water, the ions separate

think about the nature of each reactant solution

Ba(NO₃)₂(aq)

barium nitrate (white solid) dissolved in water

contains Ba²⁺ and NO₃⁻ ions (not Ba(NO₃)₂ units)

K₂CrO₄(aq)

potassium chromate (solid) dissolved in water

contains K⁺ and CrO₄²⁻ ions

two ways to represent the mixing of Ba(NO₃)₂(aq) and K₂CrO₄(aq)

K₂CrO₄(aq) + Ba(NO₃)₂(aq) → products

2K⁺(aq) + CrO₄²⁻(aq) + Ba²⁺(aq) + 2NO₃⁻(aq) → products

solution contains:

K⁺

CrO₄²⁻

Ba²⁺

NO₃⁻

know that:

when ions form a solid compound, the compound must have a zero net charge

products contain both anions and cations

K⁺ and Ba²⁺ could not combine

CrO₄²⁻ and NO₃⁻ could not combine

most ionic materials contain only two types of ions: one type of cation and one type of anion

determine possibilities for the solid

possible combinations of a given cation and anion from the list of ions in the solution:

K₂CrO₄

KNO₃

BaCrO₄

Ba(NO₃)₂

K₂CrO₄ and Ba(NO₃)₂ were the reactants, so the solid could be KNO₃ or BaCrO₄

need more facts

K⁺ ion and NO₃⁻ ion are both colorless, so KNO₃ would be white, not yellow

CrO₄²⁻ ion is yellow (K₂CrO₄ is yellow) so the solid is most likely BaCrO₄

K⁺ and NO₃⁻ were left dissolved in the solution

K₂CrO₄(aq) + Ba(NO₃)₂(aq) → BaCrO₄(s) + 2KNO₃(aq)

as long as water is present, the KNO₃ remains dissolved as separate ions

example: AgNO₃(aq) and KCl(aq) form white solid

Ag⁺, NO₃⁻ + K⁺, Cl⁻ → Ag⁺, NO₃⁻, K⁺, Cl⁻ → white solid

possible compounds: AgNO₃, KCl, AgCl, KNO₃

cannot be AgNO₃ or KCl, so either AgCl or KNO₃

KNO₃ is quite soluble in water, so solid KNO₃ will not form; solid must be AgCl

AgNO₃(aq) + KCl(aq) → AgCl(s) + KNO₃(aq)

doing chemistry requires both understanding ideas and remembering key information

Solubility of salts in water

Solubility of salts in water

most nitrate (NO₃⁻) salts are soluble

most salts containing the alkali metal ions (Li⁺, Na⁺, Cs⁺, Rb⁺) and the ammonium ion (NH₄⁺) are soluble

most chloride, bromide, and iodide salts are soluble (notable exceptions: salts containing Ag⁺, Pb²⁺, Hg₂²⁺)

most sulfate salts are soluble (notable exceptions: BaSO₄, PbSO₄, Hg₂SO₄, CaSO₄)

most hydroxides (e.g. NaOH, KOH) are only slightly soluble; Ba(OH)₂, Sr(OH)₂, Ca(OH)₂ are marginally soluble

most sulfide (S²⁻), carbonate (CO₃²⁻), chromate (CrO₄²⁻), and phosphate (PO₄³⁻) salts are only slightly soluble, except those containing the cations in Rule 2

There are no rows in this table

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when solutions containing ionic substances are mixed, determine the products by thinking in terms of ion interchange

take the cation from one reactant and combine it with the anion of the other reactant

use solubility rules to predict whether each product forms as a solid

focus on the actual components of the solution before any reaction occurs and then figure out how these components will react with each other

Describing reactions in solution

formula equation: overall reaction stoichiometry but not necessarily the actual forms of the reactants and products in solution

e.g. K₂CrO₄(aq) + Ba(NO₃)₂(aq) → BaCrO₄(s) + 2KNO₃(aq)

complete ionic equation: represents as ions all reactants and products that are strong electrolytes

e.g. 2K⁺(aq) + CrO₄²⁻(aq) + Ba²⁺(aq) + 2NO₃⁻(aq) → BaCrO₄(s) + 2K⁺(aq) + 2NO₃⁻(aq)

net ionic equation: only those solution components undergoing a change

spectator ions: ions that do not participate directly in the reaction

e.g. Ba²⁺(aq) + CrO₄²⁻(aq) → BaCrO₄(s)

Stoichiometry of precipitation reactions

the same principles of chemical stoichiometry apply to reactions that take place in solutions

convert all quantities to moles

use coefficients of the balanced equation to assemble mole ratios

determine which reactant is limiting

in solution reactions:

it is sometimes difficult to tell immediately what reaction will occur when two solutions are mixed

always first write the species that are actually present in the solution

to obtain the moles of reactants, use the volume of the solution and its molarity

solving stoichiometry problems for reactions in solution

identify the species present in the combined solution, and determine what reaction occurs

write the balanced net ionic equation for the reaction

calculate the moles of reactants

determine which reactant is limiting

calculate the moles of product(s), as required

convert to grams or other units, as required

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Acid-base reactions

acid: proton donor

base: proton acceptor

predicting acid-base reaction:

example: HCl(aq) and NaOH(aq)

focus on the species present in the mixed solution

contains H⁺, Cl⁻, Na⁺, OH⁻

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