Units 1-4

Chapter 1:  Chemical Foundations

Scientific Method:

  • Therory- a set of tested hypotheses that gives an overall explanation of some natural phenomenon. (explains what happened)
  • Natural law- a statement that expresses generally observed behavior. (summarized what happened)
  • Law of Conservation of Mass- the observation that the total mass of materials is not affected by the chemical change in those materials

SI Units:

  • Precision- the degree of agreement among several measurements of the same quantity
  • Accuracy- the agreement of a particular value with the true value

Significant Figures:

5. Exact #s never limit the number of sig figs in a calculation

Density:

D=M/V

Classification of Matter:

Matter (element, compound, mixture)

  • Matter- anything occupying space and having mass
  • Mixtures- has a variable composition
    • Homogeneous- having visibly indistinguishable parts
      • Solution
  • Heterogeneous- having  visibly distinguishable parts
  • Pure substance- one with constant composition
  • Compound- a substance with constant composition that can be broken down into elements by chemical processes
  • Element- a substance that cannot be decomposed into simpler substances by chemical or physical means

Physical/chemical properties and changes

  • Physical change- a change in the form of a substance, but not in its chemical composition; chemical bonds are not broken in a physical change.
  • Chemical change- the change of substances through reorganization of the atoms; a chemical reaction

Physical separation techniques

  • Distillation- a method for separating the components of a liquid mixture that depends on differences in the ease of vaporization of the components
  • Filtration- a method for separating the components of a liquid mixture containing a solid and a liquid
  • Chromatography- the general name for a series of methods for separating mixtures by employing a system with a mobile phase (liquid or gas) and a stationary phase (solid)

Chapter 2: Atoms, Molecules, and Ions

Law of definite proportions (aka Proust’s Law) – a given compound always contains exactly the proportions of element by mass.

Law of multiple proportions- when two elements form a series of compounds, the ratios of the masses of the second element combine with one gram of the first element can always be reduced to small whole numbers.

Dalton’s atomic theory

  1. Each element is made up of tiny particles called atoms
  2. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways
  3. Chemical compounds are formed when atoms or different elements combine with each other. A given compound always has the same relative number and types of atoms
  4. Chemical reactions involve reorganization of the atoms- changes in the way they are bounded together. The atoms themselves are not changed in a chemical reaction.

Atomic history/scientists

Molecules and ions

Naming compounds (ionic, covalent, acids, roman numerals, polyatomic ions, etc.)

Chapter 3: Stoichiometry

Average atomic mass- the average mass of the natural occurring isotopes of an element

Mole – molar mass

  • 1 mole = 6.022×10^23 units

% composition of compound

(#atoms of the element*atomic mass of element)/

(molar mass of the compound)                         *100

Empirical/molecular formula

  • Empirical formula:
  1. Determine the percentage of each element in your compound
  2. Treat % as grams, and convert grams of each element to moles of
  3. each element
  4. Find the smallest whole number ratio of atoms
  5. If the ratio is not all whole number, multiply each by an integer so that all elements are in whole number ratios
  • Molecular formula:
  1. Find the empirical formula mass
  2. Divide the known molecular mass by the empirical formula mass,deriving a whole number, n
  3. Multiply the empirical formula by n to derive the molecular formula

Balancing equations – predicting products

  • Determine what reaction is occurring
  • Write the unbalance chemical formula
  • Balance the equation by inspection
  • Include phase information

General stoichiometry – mass to mass

  1. Balance the chemical equation
  2. Convert grams of reactant or product to moles
  3. Compare moles of the known moles to the desired substance
    1. A ration derived from the coefficients in the balanced equation
    2. Convert from the moles back to grams, if required

Limiting reactant

  • It controls the amount of product that can be formed

Solving limiting reaction problems

  • Convert grams of reactant to moles
  • Use stoichiometric ratios to determine the limiting reactant
  • Solve as before, beginning the stoichiometric calculation with the grams of the limiting reactant

Percent

  1. Actual yield- what you got by actually performing the reaction
  2. Theoretical yield- what stoichiometric calculation says the reaction SHOULD have produced

Chapter 4: Types of Chemical Reactions and Solution Stoichiometry

Solubility (dissociation, electrolytes, solute/solvent)

  • Solute- substance
  • Solvent- what substance is dissolved in
  • Strong electrolyte- solution that conducts electricity
    • Substances that are completely ionized when they are dissolved in water
    • Completely ionized in water
      • Ionic compounds
      • Strong acids (HNO3(aq), H2SO4(aq), HCl(aq))
      • Strong bases (KOH and NaOH)
      • Weak electrolyte- solutions that conduct a weak electric current
        • Slightly ionized in water
          • Weak acids (organic acids- acetic, citric, butyric, and malic)
          • Weak bases (ammonia)
          • Nonelectrolytes- solutions that conduct no electric current
            • No ions present in solution
              • Alcohols and sugars

Molarity – dilution

  • Molarity- moles of solute per liters of solution
  • Dilution- the process of adding solvent to lower the concentration of solute in a solution
    • M1V1=M2V2
      • M=moles of solute before or after dilution

Precipitation reaction – solubility rules

  • Precipitation reaction- when two solutions are mixed, an insoluble substance sometimes forms; that is, a solid (precipitate) forms and separates from solution
    • Determine what reaction takes place
    • Write the balanced net ionic equation for the reaction
    • Calculate the moles of reaction
    • Determine which reactant is limiting
    • Calculate the moles of product or products
    • Convert grams or other units as required

Solubility Rules

  1. Most nitrate (NO3-) are soluble.
  2. Most salts containing the alkali metal ions (Li, Na, K, Cs, Rb) and the ammonium (NH4+) ions are soluble.
  3. Most chloride, bromide, and iodide salts are soluble. Exceptions are salts containing the ions Ag+, Pb2+, and Hg2+.
  4. Most sulfate salts are soluble. Notable exceptions are BaSO4, PbSO4. HgSO4, and CaSO4.
  5. Most hydroxide dalts are only slightly soluble. The important soluble hydroxides are NaOH and KOH. The hydroxides of barium, strontium, and calcium are marginally soluble.
  6. Most sulfide (S2-), carbonate (CO3^2-), chromate (CrO4^2-), and phosphate (PO4^3-)salts are only slightly soluble

Acid/base reaction – neutralization

  • Acids are proton donors (H+ ions)
  • Bases are proton acceptors (OH- ions)

Stoichiometry Calculations for Acid-Base reactions

  1. List the species present in the combined solution before any reaction occurs; decide what reaction will occur
  2. Write the balanced net ionic equation for this reaction
  3. Calculate the mole of reactants
    1. For reactions in solution, use volumes of the original solutions and their molarities
    2. Determine the limiting reactant where appropriate
    3. Calculate the moles of the required reactant or product
    4. Convert to grams or volume of solution as required

Redox – oxidation #, balancing ½ reactions

  • Electron transfer (LEO and GER)
    • Lose Electrons= Oxidation
    • Gain Electron= Reduction

Rules for Assigning Oxidation Numbers

  1. The oxidation number of the atom of a free element is zero
  2. The oxidation number of a monotonic ion equals its charge
  3. In compounds, oxygen has an oxidation number of -2, except in peroxides, where it is -1
  4. In compounds, fluorine is ALWAYS assigned an oxidation number of -1
  5. The sum of the oxidation states for an electrically neutral compound must be zero
  • Characteristics of oxidation-reduction reactions
    • The oxidized substance:
      • Loses electrons
      • Increases oxidation state
      • Is the reducing agent
  • The reduced substance:
    • Gains electrons
    • Decreases oxidation state
    • Is the oxidizing agent
    • Neutralization- an acid-base reaction
      • H+(aq) + OH-(aq) –> H2O(l)
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