BackGeneral Chemistry: Foundational Concepts, Classification, and Calculations
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Classification of Matter
Physical and Chemical Properties
Understanding the properties of substances is fundamental in chemistry. Properties can be classified as physical (e.g., density, color, melting point, mass) or chemical (e.g., combustibility).
Physical properties: Characteristics that can be observed or measured without changing the substance's identity (e.g., density, color).
Chemical properties: Characteristics that describe a substance's ability to undergo chemical changes (e.g., combustibility).
Example: Melting point is a physical property; combustibility is a chemical property.
Classification of Materials
Materials are classified as elements, compounds, homogeneous mixtures, or heterogeneous mixtures.
Element: Pure substance consisting of one type of atom (e.g., Antimony).
Compound: Substance formed from two or more elements chemically combined (e.g., Steel is not a compound, but an alloy).
Homogeneous mixture: Uniform composition throughout (e.g., Apple juice).
Heterogeneous mixture: Non-uniform composition (e.g., Granite).
Example: Soil is a heterogeneous mixture, not homogeneous.
Units, Conversions, and Calculations
Density and Unit Conversion
Density is a physical property defined as mass per unit volume. Converting units is essential for comparing measurements.
Density formula:
Unit conversions: 1 foot = 30.48 cm; 1 lb = 453.6 g; 1 mL = 1 cm3
Example: Convert 7.86 g/mL to pounds per cubic foot using conversion factors.
Distance and Volume Comparisons
Comparing units of length and volume is important for understanding scale in chemistry.
Length units: 1 Ångström = m; 1 nm = m; 1 mm = m.
Volume units: 1 in = 2.54 cm; 1 cm3 = 1 mL; 1 L = mL.
Example: Which is greater: 5 Ångströms or 1 nm?
Calculating Mass and Molecules
Calculations involving mass, moles, and molecules are central to quantitative chemistry.
Mole concept: 1 mole = particles (Avogadro's number).
Example: Calculate the number of hydrogen atoms in 27.0 g of H2O.
Chemical Compounds and Formulas
Ionic and Molecular Compounds
Chemical compounds are classified as ionic or molecular based on the types of elements involved.
Ionic compounds: Formed from metals and nonmetals; consist of ions (e.g., NaCl, CaO).
Molecular compounds: Formed from nonmetals; consist of molecules (e.g., H2O, CO2).
Example: NH4Cl is ionic; H2O is molecular.
Empirical and Molecular Formulas
The empirical formula shows the simplest whole-number ratio of atoms; the molecular formula shows the actual number of atoms in a molecule.
Empirical formula: Simplest ratio (e.g., CH for hydrocarbons).
Molecular formula: Actual composition (e.g., C2H2 for acetylene).
Example: An alkene with empirical formula CH and molar mass 26.0 amu has molecular formula C2H2.
Formula Weights and Percent Composition
Formula weight is the sum of atomic masses in a compound; percent composition is the percentage by mass of each element.
Formula weight:
Percent composition:
Example: Calculate %O in NaHCO3 or C6H12O6.
Atomic Theory and Laws of Chemistry
Dalton's Atomic Theory
Dalton's theory laid the foundation for modern chemistry, describing atoms and their combinations.
Key ideas:
Elements are composed of small indivisible particles called atoms.
Atoms of the same element are identical; atoms of different elements differ.
Atoms combine in simple whole-number ratios to form compounds.
Atoms are not created or destroyed in chemical reactions.
Example: The Law of Definite Proportions states that a compound always contains the same elements in the same proportion by mass.
Atomic Structure and Isotopes
Atoms consist of protons, neutrons, and electrons. The atomic number is the number of protons; isotopes differ in neutron number.
Atomic number (Z): Number of protons in nucleus.
Mass number (A): Number of protons plus neutrons.
Isotopes: Atoms of the same element with different numbers of neutrons.
Example: Silicon atom has 14 protons.
Classification and Nomenclature
Element Symbols and Compound Names
Each element has a unique symbol; compounds are named according to rules based on their composition.
Element symbols: One or two letters (e.g., Cu for copper, Pb for lead).
Compound formulas: Reflect the ratio of atoms (e.g., Na2SO4 for sodium sulfate).
Example: KCl is potassium chloride, not potassium carbonate.
Ion Charges and Polyatomic Ions
Ions are atoms or groups of atoms with a net charge. Polyatomic ions consist of multiple atoms.
Common ions: Oxide (O2-), fluoride (F-), hydroxide (OH-), nitrate (NO3-), acetate (CH3COO-).
Example: Oxide does not have a -1 charge; it is -2.
Stoichiometry and Chemical Calculations
Mole Calculations and Avogadro's Number
Stoichiometry involves calculations based on the mole concept and Avogadro's number.
Avogadro's number: particles/mol.
Example: Calculate the number of molecules in a given mass or volume.
Empirical Formula Determination
Empirical formulas are determined from percent composition by mass.
Steps:
Convert percentages to grams.
Convert grams to moles using atomic masses.
Divide by smallest number of moles to get ratios.
Write empirical formula.
Example: A compound with 29.1% Na, 40.5% S, 30.4% O yields Na2SO3 as empirical formula.
Fundamental Laws of Chemistry
Law of Definite Proportions
This law states that a chemical compound always contains the same elements in exactly the same proportion by mass.
Example: Water is always 89% oxygen and 11% hydrogen by mass.
Law of Multiple Proportions
When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.
Example: CO and CO2 are compounds of carbon and oxygen with different ratios.
Additional info:
Some questions involve practical applications, such as calculating gasoline costs based on mileage and fuel efficiency.
Atomic theory, classification, and stoichiometry are foundational topics for General Chemistry students.
