General Chemistry Fundamentals: Measurement, Matter, Atomic Theory, and Chemical Nomenclature

Major Topics in General Chemistry

Measurement: Precision and Accuracy

Accurate and precise measurements are essential in chemistry for reliable experimental results. Precision refers to how close repeated measurements are to each other, while accuracy describes how close a measurement is to the true value.

• Precision: Consistency of repeated measurements.

• Accuracy: Closeness to the actual or accepted value.

• Example: Measuring the mass of a sample multiple times and obtaining similar results demonstrates precision.

Significant Figures and Scientific Notation

Significant figures indicate the certainty in a measurement. Scientific notation is used to express very large or small numbers efficiently.

• Rules for Significant Figures: All nonzero digits are significant; zeros between significant digits are significant; leading zeros are not significant; trailing zeros in a decimal number are significant.

• Scientific Notation: Expresses numbers as a product of a coefficient and a power of ten, e.g., $6.02 \times 10^{23}$.

• Example: 0.00450 has three significant figures.

Metric Prefixes and SI Units

The metric system uses prefixes to indicate multiples or fractions of base units. The SI system is the standard for scientific measurements.

• Common Prefixes: kilo (k, $10^3$), milli (m, $10^{-3}$), centi (c, $10^{-2}$).

• SI Units: Meter (length), kilogram (mass), second (time), mole (amount of substance).

• Conversion Problems: Use dimensional analysis to convert between units.

Density

Density is a physical property defined as mass per unit volume.

• Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Units: g/cm3 or kg/m3

• Example: A 10 g object with a volume of 2 cm3 has a density of 5 g/cm3.

Matter: Physical and Chemical Properties and Changes

Matter is anything that has mass and occupies space. It can undergo physical or chemical changes.

• Physical Properties: Observable without changing composition (e.g., color, melting point).

• Chemical Properties: Describe reactivity and composition changes (e.g., flammability).

• Physical Change: Change in state or appearance (e.g., melting ice).

• Chemical Change: Formation of new substances (e.g., rusting iron).

Classification of Matter

Matter is classified as substances, elements, or compounds.

• Substance: Pure form of matter with uniform properties.

• Element: Cannot be broken down by chemical means.

• Compound: Composed of two or more elements chemically combined.

Mixtures: Homogeneous and Heterogeneous

Mixtures contain two or more substances physically combined.

• Homogeneous Mixture (Solution): Uniform composition throughout (e.g., saltwater).

• Heterogeneous Mixture: Non-uniform composition (e.g., sand in water).

• Methods of Separation: Filtration, distillation, chromatography.

Compounds and Molecules

Compounds are substances formed from two or more elements. Molecules are groups of atoms bonded together.

• Molecule: Smallest unit of a compound retaining its properties.

• Diatomic Molecules: Molecules composed of two atoms (e.g., H2, O2).

The Periodic Table

The periodic table organizes elements by increasing atomic number and recurring chemical properties.

• Groups: Vertical columns (e.g., alkali metals, halogens).

• Periods: Horizontal rows.

• Main Group Elements: Groups 1, 2, and 13-18.

• Transition Metals: Groups 3-12.

• Special Groups: Noble gases, coinage metals.

Law of Conservation of Mass

This law states that mass is neither created nor destroyed in a chemical reaction.

• Implication: Total mass of reactants equals total mass of products.

• Example: Burning of hydrogen and oxygen to form water preserves total mass.

Temperature Scales and Conversions

Temperature can be measured in Celsius (°C), Fahrenheit (°F), or Kelvin (K).

• Conversion Formulas:

  • $K = ^\circ C + 273.15$

  • $^\circ F = \frac{9}{5}(^\circ C) + 32$

  • $^\circ C = \frac{5}{9}(^\circ F - 32)$

Atomic Theory and Structure

Dalton's Atomic Theory

John Dalton proposed that matter is composed of indivisible atoms, which combine in fixed ratios to form compounds.

• Atoms: Smallest unit of an element.

• Law of Constant Composition: A given compound always contains the same proportion of elements.

• Law of Multiple Proportions: Elements can combine in different ratios to form different compounds.

Subatomic Particles

Atoms consist of protons, neutrons, and electrons.

• Proton: Positively charged, found in nucleus.

• Neutron: Neutral, found in nucleus.

• Electron: Negatively charged, found in electron cloud.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Atomic Number, Mass Number, and Isotopic Abundance

Atomic number is the number of protons; mass number is the sum of protons and neutrons.

• Atomic Number (Z): Number of protons.

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

• Isotopic Abundance: Relative percentage of each isotope in nature.

• Average Atomic Mass Formula:

  • $\text{Average Atomic Mass} = \sum (\text{isotopic mass} \times \text{fractional abundance})$

Ions: Cations and Anions

Ions are atoms or molecules with a net electric charge due to loss or gain of electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

Chemical Nomenclature and Formulas

Naming Compounds and Writing Formulas

Chemical nomenclature provides systematic names for compounds.

• Ionic Compounds: Composed of cations and anions; name cation first, then anion.

• Covalent (Molecular) Compounds: Use prefixes to indicate number of atoms (e.g., CO2 is carbon dioxide).

• Using the Periodic Table: Helps determine charges and names of ions.

Bonding Models

• Ionic Bonding: Transfer of electrons from metal to nonmetal.

• Covalent Bonding: Sharing of electrons between nonmetals.

Molecular Weight, Formula Weight, and Avogadro's Number

• Molecular Weight: Sum of atomic masses in a molecule.

• Formula Weight: Sum of atomic masses in an ionic compound's formula unit.

• Avogadro's Number: $6.022 \times 10^{23}$ particles per mole.

Monatomic and Polyatomic Ions

Monatomic ions consist of a single atom; polyatomic ions contain multiple atoms bonded together.

• Common Monatomic Ions: Fe2+, Fe3+, Cr2+, Cr3+, Cu+, Cu2+, Zn2+, Ag+

• Common Polyatomic Ions: NO3-, ClO4-, SO42-, PO43-, CO32-, HCO3-, NH4+

• Example: ClO3- is chlorate; can be replaced with Br or I for bromate or iodate.

Table: Common Monatomic and Polyatomic Ions

Ion Type	Formula	Name
Monatomic	Fe2+	Iron(II)
Monatomic	Cu2+	Copper(II)
Polyatomic	NO3-	Nitrate
Polyatomic	SO42-	Sulfate
Polyatomic	ClO4-	Perchlorate
Polyatomic	HPO42-	Hydrogen phosphate

Additional info:

• Some content inferred and expanded for completeness, such as detailed definitions, examples, and formulas.

• Polyatomic ions list expanded for clarity.