General Chemistry Fundamentals: Scientific Method, Measurement, Atomic Structure, and Chemical Nomenclature

The Scientific Method and Branches of Chemistry

Introduction to Chemistry

Chemistry is the study of matter, its properties, and the changes it undergoes. The scientific method is a systematic approach used to investigate natural phenomena and acquire new knowledge.

• Scientific Method Steps:

  1. Make an observation (qualitative or quantitative).

  2. Formulate a hypothesis – a possible explanation that must be testable.

  3. Test the hypothesis/perform experiments.

• Theory: An explanation or model that holds up over time and may change with new observations.

• Law: A summary of observed behavior; tells us what happens, not why.

Branches of Chemistry

• Biochemistry

• Organic Chemistry

• Inorganic Chemistry

• Analytical Chemistry

• Environmental Chemistry

• Geochemistry

• Physical Chemistry

Measurement: Scientific Notation and Significant Figures

Scientific Notation

Scientific notation is a way to express very large or very small numbers using powers of ten.

• Format: $a \times 10^n$

• Example: $3.72 \times 10^{11}$

• Positive exponent: large number; negative exponent: small number.

Significant Figures (Sig Figs)

Significant figures indicate the precision of a measured value.

• Rules:

  1. All non-zero digits are significant.

  2. Zeros between non-zero digits are significant.

  3. Leading zeros are not significant (placeholders).

  4. Trailing zeros are significant if there is a decimal point.

• Adding/Subtracting: Answer has the same precision as the measurement with the least sig figs.

• Multiplying/Dividing: Answer has no more total sig figs than the measurement with the least.

• Exact numbers: Do not influence sig figs.

Example: $21.35$ rounds to $21.4$ (3 sig figs).

Accuracy: How close a measurement is to the true value. Precision: How close measurements are to each other.

Unit Conversions and the Metric System

Basic SI Quantities and Conversion Factors

The metric system uses standard units for measurement. Conversion factors allow for changing between units.

Quantity	SI Unit	English Unit	Extra
Length	Meter (m)	Inch (in)
Mass	Kilogram (kg)	Pound (lb)
Time	Second (s)	Second(s)
Temperature	Celsius (°C)	Fahrenheit (°F)	Kelvin (K)
Luminous Intensity	Candela (cd)
Amount of Substance	Mole (mol)

• Distance: 1 in = 2.54 cm, 1 m = 100 cm, 1 km = 1000 m

• Mass: 1 kg = 1000 g, 1 lb = 454 g

• Volume: 1 L = 1000 mL, 1 gal = 3.785 L

• Time: 1 hr = 60 min, 1 yr = 365.25 days

Metric Prefixes

• Larger: Tera ($10^{12}$), Giga ($10^9$), Mega ($10^6$), Kilo ($10^3$), Hecto ($10^2$), Deca ($10^1$)

• Smaller: Deci ($10^{-1}$), Centi ($10^{-2}$), Milli ($10^{-3}$), Micro ($10^{-6}$), Nano ($10^{-9}$), Pico ($10^{-12}$)

Example: $45\,\text{kg} \times 1000\,\text{g/kg} = 45000\,\text{g} = 4.5 \times 10^4\,\text{g}$

Temperature Conversions

• Celsius to Fahrenheit: $F = 1.8C + 32$

• Celsius to Kelvin: $K = C + 273$

Example: Water boils at $373.15\,K$ ($100\,°C$), freezes at $273.15\,K$ ($0\,°C$).

Energy Units

• Joule (J): SI unit of energy

• Calorie (cal): 1 cal = 4.184 J

• Food Calorie (Cal): 1 Cal = 1000 cal

Classification of Matter

Types of Matter

• Elements: Simplest form of matter, cannot be broken down into simpler substances.

• Compounds: Substances composed of two or more elements chemically combined.

• Mixtures: Physical blend of two or more substances.

  • Heterogeneous: Not the same throughout.

  • Homogeneous: Same composition throughout.

• Pure Substances: Elements and compounds with uniform composition.

The Periodic Table

Organization and Properties

• Rows are called periods.

• Columns are called groups or families (similar chemical properties).

• Group 1A: Alkali metals (Li, Na, K)

• Group 2A: Alkaline earth metals (Be, Mg, Ca)

• Group 7A: Halogens (F, Cl, Br, I)

• Group 8A: Noble gases (He, Ne, Ar, Kr)

Properties of Metals, Nonmetals, and Metalloids

• Metals: Shiny, conduct electricity and heat, malleable, ductile.

• Nonmetals: Any color, solid/liquid/gas, not malleable or ductile, do not conduct electricity.

• Metalloids: Properties intermediate between metals and nonmetals; conduct electricity under certain conditions.

States of Elements at Room Temperature

• Liquids: Mercury (Hg), Bromine (Br)

• Gases: Hydrogen (H), Helium (He), Nitrogen (N), Oxygen (O), Fluorine (F), Chlorine (Cl)

• Solids: All other elements

Symbols of Elements

• One or two-letter symbol; first letter capitalized, second letter lowercase.

• Examples: H (Hydrogen), He (Helium), Li (Lithium), C (Carbon), O (Oxygen), Na (Sodium), Fe (Iron)

Atomic Structure: Protons, Neutrons, Electrons, and Isotopes

Atomic Number and Mass Number

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Number of protons plus neutrons.

• Number of Neutrons: $\text{Mass Number} - \text{Atomic Number}$

Isotopes

• Atoms of the same element with different numbers of neutrons (different mass numbers).

• Isotopes can be written as H-1, H-2, etc.

Atomic Mass

• The average mass of an atom of a given element, weighted by the abundance of its isotopes.

• Weighted Average Formula: $\text{Atomic Mass} = (\text{Fraction}_1 \times \text{Mass}_1) + (\text{Fraction}_2 \times \text{Mass}_2) + \ldots$

Example: Carbon has two naturally occurring isotopes: C-12 and C-13. $\text{Atomic Mass} = (0.9889 \times 12.00) + (0.0111 \times 13.00) = 12.01$

Chemical Compounds and Nomenclature

Molecular and Structural Formulas

• Molecular Formula: Shows the types and numbers of atoms in a molecule (e.g., $\text{CH}_3\text{OH}$).

• Structural Formula: Shows how atoms are connected (e.g., H–C–O–C–H).

Ionic and Molecular Compounds

• Ionic Compounds: Formed between metals and nonmetals; involve transfer of electrons.

• Molecular (Covalent) Compounds: Formed between two nonmetals; involve sharing of electrons.

• Ionic Bond: Attraction between positive and negative ions.

Naming Ionic Compounds

• Metal + Nonmetal → Ionic

• Sodium and chlorine combine to form sodium chloride: $\text{Na}^+ + \text{Cl}^- \rightarrow \text{NaCl}$

• Name the nonmetal with an -ide ending (e.g., Magnesium chloride: $\text{MgCl}_2$)

• Transition metals use Roman numerals to indicate charge (e.g., Fe(III) chloride: $\text{FeCl}_3$)

Polyatomic Ions

• Groups of atoms with an overall charge.

• Examples:

  • $\text{CO}_3^{2-}$ = carbonate

  • $\text{NO}_3^-$ = nitrate

  • $\text{NH}_4^+$ = ammonium

  • $\text{SO}_4^{2-}$ = sulfate

Naming Molecular Compounds

• Use prefixes to indicate the number of each atom (e.g., CO = carbon monoxide, CO2 = carbon dioxide).

• Diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2

Additional info: For a complete list of polyatomic ions, refer to a standard chemistry reference table.