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Chemistry · Ch 1 — Some Basic Concepts of Chemistry

Summary

Summary

  • Matter is anything with mass and volume; classified as mixtures (homogeneous/heterogeneous) or pure substances (elements/compounds).
  • SI base units include: mass (kg), length (m), time (s), amount of substance (mol), temperature (K), electric current (A), luminous intensity (cd).
  • Mass is constant; weight varies with gravity. Volume has SI unit m3m^3; common unit is litre (1 L=10−3 m31 \, L = 10^{-3} \, m^3).
  • Density = massvolume\frac{\text{mass}}{\text{volume}}; SI unit kg/m3kg/m^3, often g/cm3g/cm^3 or g/mLg/mL.
  • Temperature: Kelvin (KK) = ∘C+273.15^\circ C + 273.15; ∘F=95(∘C)+32^\circ F = \frac{9}{5}(^\circ C) + 32.
  • Significant figures: all certain digits plus one uncertain digit. Rules: non-zero digits always significant; zeros between non-zero digits are significant; leading zeros not significant; trailing zeros significant only if decimal is present.
  • Scientific notation: N×10nN \times 10^n, where 1≤N<101 \le N < 10.
  • Precision = closeness of repeated measurements; accuracy = closeness to true value.
  • Laws of chemical combination:
    • Law of conservation of mass: mass is neither created nor destroyed in a chemical reaction.
    • Law of definite proportions: a compound always contains same elements in fixed proportion by mass.
    • Law of multiple proportions: if two elements form more than one compound, masses of one element that combine with fixed mass of the other are in small whole-number ratio.
    • Gay Lussac’s law of gaseous volumes: gases combine in simple whole-number volumes at same TT and PP.
    • Avogadro’s law: equal volumes of gases at same TT and PP contain equal number of molecules.
  • Atomic mass = average mass of an atom relative to 112\frac{1}{12} mass of 12C^{12}C atom; unit amu (or u). Molecular mass = sum of atomic masses of all atoms in a molecule.
  • Mole concept: 1 mol=6.022×10231 \, \text{mol} = 6.022 \times 10^{23} entities (Avogadro’s number, NAN_A). Number of moles n=mass (g)molar mass (g/mol)=number of particlesNAn = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}} = \frac{\text{number of particles}}{N_A}.
  • Molar mass = mass of one mole of a substance in grams; numerically equal to atomic/molecular mass in amu.
  • Percentage composition: mass % of element = mass of element in 1 mol compoundmolar mass of compound×100\frac{\text{mass of element in 1 mol compound}}{\text{molar mass of compound}} \times 100.
  • Empirical formula = simplest whole-number ratio of atoms; molecular formula = n×n \times empirical formula, where n=molecular massempirical formula massn = \frac{\text{molecular mass}}{\text{empirical formula mass}}. …
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