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Chapter 1. Fundamentals of Chemistry

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1. What is Chemistry?

2. Grammar of Chemical Formulas

3. Chemical Reaction Equations

a. Fundamentals of Organic Chemistry

1. What is Chemistry?

⑴ Chemistry: The study of matter

① Composition, properties, structure of matter

② Changes in matter

⑵ States of Matter

① Solid

○ Has a definite shape and fixed volume

② Liquid

○ More irregular than a solid

○ Because its movements are relatively free, it has flowing properties

○ It doesn’t have a fixed shape and takes the shape of its container, though its volume remains constant

③ Gas

○ More disordered than liquids

○ Molecules are far apart, negligible molecular interactions

○ Much lighter compared to solids or liquids of the same volume, compressible, occupies space

⑶ Classification of Matter

① Pure Substances: Distinguished between elements and compounds

② Mixtures: Classified into homogeneous and heterogeneous mixtures

③ Elements

○ Substances composed of only one type of atom

○ Examples: Copper (Cu), Nitrogen (N₂), Iron (Fe), Diamond (C), Aluminum (Al)

④ Compounds

○ Substances composed of two or more different elements in a fixed ratio

○ Examples: Carbon Dioxide (CO₂), Copper Sulfate (CuSO₄), Water (H₂O)

⑤ Homogeneous Mixtures (Solutions)

○ Mixtures where multiple pure substances are uniformly distributed, composition is the same throughout

○ Examples: Air, Sugar Solution, Saltwater, Alloys

⑥ Heterogeneous Mixtures

○ Mixtures where multiple pure substances are not uniformly distributed, composition varies in different parts

○ Examples: Muddy Water, Milk, Granite, Blood, Smoke

⑷ Properties of Matter

① Physical Properties

○ Changes in state or form of matter without altering its chemical composition

○ Examples: Color, Melting Point, Boiling Point, Density, Solubility, etc.

② Chemical Properties

○ Transformation of a substance into new substances through chemical reactions

2. Grammar of Chemical Formulas

⑴ Atoms, Elements, Molecules, Ions

① Atom: Fundamental unit in most chemical reactions. Units smaller than atoms were observed after the 1900s

② Element: Collection of atoms with similar characteristics

③ Molecule: Functional unit composed of bonded atoms

④ Ion: Atom with a charge due to gaining or losing electrons

⑤ Representation of Atoms or Ions

○ a: Number of protons + Number of neutrons

○ b: Number of protons

○ c: Indicates charge state of X. -2, -, +, +2, etc.

○ d: Indicates the number of X present

⑥ Isotope: Elements with the same atomic number but different neutron numbers, thus having different mass numbers

⑵ Mole

① Mole: Unit for counting atoms or molecules

② Avogadro’s Number (N_A): Number of atoms in one mole, defined as the number of atoms in 12 g of carbon-12

○ N_A = 6.02 × 10²³

③ Gas Volume of 1 Mole is Constant for All Gases, at STP (0 ℃, 1 atm) it’s 22.4 L

○ Derived from the Ideal Gas Law

④ Humans are composed of around 10,000 moles

⑶ Chemical Formulas, Chemical Reaction Equations, and Reaction Yield

① Chemical Formula: Representation of a substance

○ Empirical Formula: Simplest whole number ratio of atoms

○ Example: Empirical formula of glucose is CH₂O

○ Molecular Formula: Actual number of atoms in a molecule

○ Example: Molecular formula of glucose is C₆H₁₂O₆

○ Determining Empirical to Molecular Formula: Mass percentage composition analysis → Empirical formula determination → Molecular formula determination (using mass spectrometry)

② Structural Formula

⑷ Stoichiometry

① Mass

○ Atomic Weight: Relative mass of one atom. Based on ¹²C = 12.00 as the standard (no units)

○ Molecular Weight: Relative mass of one molecule. Sum of atomic weights of atoms in the molecule (no units)

○ Experimental Weight: Sum of atomic weights in the experimental formula

○ Average Atomic Weight: Weighted average considering the isotopic distribution. Calculation of carbon’s average atomic weight is shown

② Concentration

○ Molar Concentration: Moles of solute (mol) ÷ Volume of solution (L)

○ Molal Concentration: Moles of solute (mol) ÷ Mass of solvent (kg)

○ Boiling point elevation, freezing point depression, etc.

○ Mole Fraction: Moles of a specific substance (mol) ÷ Total moles in the mixture (mol), dimensionless quantity

○ Mass Percent: Mass of solute (kg) ÷ Mass of solution (kg) × 100 (%)

○ ppm (parts per million), ppb (parts per billion)

○ 1 ppm X = 10^-6 kg X / 1 kg solution

○ 1 ppb X = 10^-9 kg X / 1 kg solution

○ ppmv (parts per million by volume), ppbv (parts per billion by volume)

○ 1 ppmv X = 10^-6 L X / 1 L solution

○ 1 ppbv X = 10^-9 L X / 1 L solution

③ Gram Equivalent

○ The mass or molecular weight divided by the equivalent

○ Represents the number of grams corresponding to 1 mole equivalent

3. Chemical Reaction Equations

⑴ Chemical Reaction Equation: Representation of chemical reactions using chemical formulas

① Reactant

② Product

③ Reagent: Chemical substance available for use in the laboratory

④ Spectator Ion: Ions indicated in the equation that do not participate in the reaction but are present

○ The equation excluding spectator ions is called the net ionic equation

⑤ Method for Writing Chemical Reaction Equations

○ Step 1: Represent reactants and products with chemical formulas

○ Step 2: Skeletal equation: Write down which chemical species react

○ Step 3: Balanced chemical equation: Calculate coefficients in the skeletal equation

○ Step 4: Detailed chemical equation: Include states of each chemical species (e.g., g , l , s , aq ) and conditions (e.g., Δ, hν)

○ Gas state: (g)

○ Liquid state: (l)

○ Solid state: (s)

○ Aqueous solution state: (aq)

⑵ Laws of Chemical Reaction Equations

① Law of Conservation of Mass: Number of atoms must be conserved for all atoms, except in nuclear reactions

② Law of Conservation of Charge: Sum of charges in reactants equals sum of charges in products

③ Law of Constant Composition: Elements combine in fixed ratios

④ Law of Multiple Proportions: Different elements combine in whole number ratios to a fixed mass of one element

○ Except when two elements form more than one compound

⑶ Coefficient Calculation in Chemical Reaction Equations

① For redox reactions, the reciprocal ratio of oxidation number change between oxidized and reduced species gives the coefficient ratio

○ Usually, the redox reaction involves determining the coefficients for species with changing oxidation numbers, then balancing the coefficients of H2O and H+

② Example

○ C₆H₁₄ + O₂ → CO₂ + H₂O

○ 1C₆H₁₄ + O₂ → CO₂ + H₂O

○ 1C₆H₁₄ + O₂ → 6CO₂ + H₂O (∵ C)

○ 1C₆H₁₄ + O₂ → 6CO₂ + 7H₂O (∵ H)

○ 1C₆H₁₄ + 19/2O₂ → 6CO₂ + 7H₂O (∵ O)

○ 2C₆H₁₄ + 19O₂ → 12CO₂ + 14H₂O

⑷ Limiting Reactant

① Substance consumed entirely in a reaction

② Amount of products formed is proportional to the amount of limiting reactant

⑸ Percentage Reaction Yield (%) = Actual Yield / Theoretical Yield × 100

① Reactions are not always complete due to chemical equilibrium and reverse reactions

⑹ Classification of Reactions

① All reactions involve electron movement, breaking and forming bonds

② Reactions with electron movement and change in oxidation number: Redox Reactions

③ Reactions with electron movement but no change in oxidation number: Acid-Base Reactions. Further categorized into precipitation reactions and non-precipitation reactions

Input: 2019.01.08 21:46