General Chemistry: Matter, Atoms, Molecules, and Chemical Quantities

CHAPTER 1 – MATTER AND MEASUREMENTS

Introduction to Chemistry

Chemistry is the study of matter and the changes it undergoes. It is central to many aspects of daily life, from food and medicine to materials and energy.

• Matter: Anything that occupies space and has mass.

• Chemistry: The central science, connecting physical sciences with life sciences and applied sciences.

The Scientific Method

The scientific method is a systematic approach to research and problem-solving in science.

1. Define the problem.

2. Perform experiments, make observations, and record data (qualitative or quantitative).

3. Formulate a hypothesis (tentative explanation).

4. Develop a law (concise statement of a consistent relationship) or a theory (unifying principle explaining facts).

Classification of Matter

• Substance: Matter with definite composition and distinct properties (e.g., copper, water).

• Mixture: Combination of two or more substances retaining their identities. Can be:

  • Homogeneous: Uniform composition (e.g., salt water).

  • Heterogeneous: Non-uniform composition (e.g., sand and iron filings).

• Element: Substance that cannot be separated into simpler substances by chemical means.

• Compound: Substance composed of two or more elements chemically united in fixed proportions.

Physical and Chemical Properties

• Physical Property: Can be measured without changing the substance's composition (e.g., color, melting point).

• Chemical Property: Observed only during a chemical change (e.g., reactivity with acid).

• Extensive Property: Depends on the amount of matter (e.g., mass, volume).

• Intensive Property: Independent of the amount of matter (e.g., density, boiling point).

States of Matter

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes shape of container.

• Gas: No definite shape or volume, expands to fill container.

Ions and Polyatomic Ions

• Ion: Atom or group of atoms with a net charge due to loss or gain of electrons.

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

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

• Polyatomic Ion: Ion composed of more than one atom (e.g., NH4+, SO42−).

Name	Formula/Charge	Name	Formula/Charge
Ammonium	NH4+	Hydronium	H3O+
Phosphate	PO43−	Sulfate	SO42−
Nitrate	NO3−	Hydroxide	OH−
Carbonate	CO32−	Permanganate	MnO4−
Chlorate	ClO3−	Chromate	CrO42−
Acetate	CH3COO−	Oxalate	C2O42−

Measurements and Units

• SI Units: Standard units for scientific measurements (meter, kilogram, second, mole, kelvin).

• Prefixes: kilo (103), milli (10−3), micro (10−6), nano (10−9).

• Temperature: Celsius (°C), Kelvin (K), Fahrenheit (°F).

• Density: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

Scientific Notation and Significant Figures

• Scientific Notation: Expresses numbers as $a \times 10^n$.

• Significant Figures: Digits that convey meaningful information about precision.

• Rules:

  • Nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

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

• Calculations:

  • Addition/Subtraction: Result has as many decimal places as the least precise measurement.

  • Multiplication/Division: Result has as many significant figures as the measurement with the fewest significant figures.

Dimensional Analysis

• Use conversion factors to convert between units.

• Example: $1252\ \text{in} \times \frac{2.54\ \text{cm}}{1\ \text{in}} = 3180\ \text{cm}$

CHAPTER 2 – ELEMENTS, COMPOUNDS, AND THE LAWS OF CHEMICAL COMBINATION

Early Atomic Theory

• Democritus: Matter is composed of indivisible particles called atoms.

• Dalton's Atomic Theory:

  1. Elements are composed of atoms.

  2. Atoms of the same element are identical; different elements have different atoms.

  3. Compounds are combinations of atoms in fixed ratios.

  4. Chemical reactions rearrange atoms; atoms are not created or destroyed.

Laws of Chemical Combination

• Law of Definite Proportions: A compound always contains the same elements in the same proportion 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 small whole-number ratios.

Atomic Structure

• Subatomic Particles:

  • Electron: Negatively charged, very small mass ($9.11 \times 10^{-28}$ g).

  • Proton: Positively charged, mass $1.67 \times 10^{-24}$ g.

  • Neutron: No charge, mass $1.67 \times 10^{-24}$ g.

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

• Mass Number (A): Total number of protons and neutrons.

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

Particle	Mass (g)	Charge (C)	Charge	Mass (amu)
Electron	9.10939 × 10−28	−1.6022 × 10−19	−1	0.00054858
Proton	1.67262 × 10−24	+1.6022 × 10−19	+1	1.00728
Neutron	1.67493 × 10−24	0	0	1.00867

The Periodic Table

• Elements are arranged in periods (rows) and groups (columns) based on properties.

• Metals: Good conductors, malleable.

• Non-metals: Poor conductors.

• Metalloids: Properties intermediate between metals and non-metals.

• Special groups:

  • Group 1A: Alkali metals (+1 ions)

  • Group 2A: Alkaline earth metals (+2 ions)

  • Group 7A: Halogens (−1 ions)

  • Group 8A: Noble gases (unreactive)

Molecules and Ions

• Molecule: Aggregate of two or more atoms held together by chemical bonds.

• Diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2.

• Polyatomic molecules: More than two atoms (e.g., SO2, C8H18).

Chemical Formulas and Naming

• Empirical Formula: Simplest whole-number ratio of elements in a compound.

• Molecular Formula: Actual number of atoms of each element in a molecule.

• Structural Formula: Shows arrangement of atoms.

• Naming Binary Molecular Compounds:

  • First element named first, second element ends with “-ide”.

  • Prefixes indicate number of atoms (mono-, di-, tri-, etc.).

  • Examples: CO = carbon monoxide, CO2 = carbon dioxide.

• Naming Acids:

  • Binary acids: hydro- + root + -ic acid (e.g., HCl = hydrochloric acid).

  • Oxoacids: Based on number of oxygens (per-, -ic, -ous, hypo-).

CHAPTER 3 – SOLUTIONS AND CHEMICAL QUANTITIES

Atomic and Molecular Mass

• Atomic Mass Unit (amu): 1 amu = 1/12 the mass of a 12C atom.

• Average Atomic Mass: Weighted average of isotopic masses.

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

• Formula Mass: Used for ionic compounds (sum of atomic masses in the formula unit).

The Mole and Avogadro’s Number

• Mole (mol): Amount of substance containing $6.022 \times 10^{23}$ entities (Avogadro’s number).

• Molar Mass: Mass of one mole of a substance (g/mol).

Key Equations:

• Number of moles: $n = \frac{\text{mass}}{\text{molar mass}}$

• Number of particles: $N = n \times N_A$

Percent Composition

• Percent by mass of each element in a compound:

$\%\ \text{element} = \frac{\text{(number of moles of element)} \times \text{atomic mass}}{\text{molar mass of compound}} \times 100$

Empirical and Molecular Formulas

• Empirical Formula: Simplest ratio of elements.

• Molecular Formula: Actual number of atoms; a multiple of the empirical formula.

• To determine empirical formula from percent composition:

  1. Assume 100 g sample; convert % to grams.

  2. Convert grams to moles for each element.

  3. Divide by smallest number of moles to get ratio.

  4. Multiply to get whole numbers if necessary.

• To find molecular formula: $n = \frac{\text{molar mass}}{\text{empirical formula mass}}$

Combustion Analysis

• Used to determine empirical formula of compounds containing C, H, and O.

• All C converted to CO2, all H to H2O; O determined by difference.

Stoichiometry and Chemical Equations

• Chemical Equation: Symbolic representation of a chemical reaction.

• Balancing Equations: Ensures conservation of mass; same number of each atom on both sides.

• Stoichiometry: Quantitative relationships between reactants and products.

• Always convert quantities to moles before using mole ratios from the balanced equation.

Limiting Reactant and Theoretical Yield

• Limiting Reactant: Reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical Yield: Maximum amount of product possible from given reactants.

• Actual Yield: Amount of product actually obtained.

• Percent Yield: $\%\ \text{yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100$

Solution Concentrations

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Dilution Equation: $M_1V_1 = M_2V_2$

• Other Units:

  • Mass percentage: $\frac{\text{mass of component}}{\text{mass of solution}} \times 100$

  • Volume percentage: $\frac{\text{volume of component}}{\text{volume of solution}} \times 100$

  • Parts per million (ppm): $\frac{\text{mass of solute}}{\text{mass of solution}} \times 10^6$

  • Parts per billion (ppb): $\frac{\text{mass of solute}}{\text{mass of solution}} \times 10^9$

Example Calculations

• Converting grams to moles: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Finding number of particles: $\text{number of particles} = \text{moles} \times 6.022 \times 10^{23}$

• Calculating percent composition, empirical and molecular formulas, and solution concentrations using the formulas above.

Mass Spectrometry

• Instrument used to determine atomic and molecular masses and isotopic abundances.

• Ions are separated based on mass-to-charge ratio and detected to determine relative abundances.

Additional info: Some content was expanded for clarity and completeness, including the structure of the periodic table, naming conventions, and stepwise procedures for empirical/molecular formula determination.