General Chemistry Foundations: Matter, Measurement, Atomic Theory, and Chemical Bonding

Intro to Chemistry

What is Chemistry?

Chemistry is the scientific study of matter, its properties, composition, and the changes it undergoes. It is often called the "central science" because it connects physical sciences with life sciences and applied sciences.

• Matter: Anything that has mass and occupies space.

• Chemical Change: A process where substances are transformed into new substances.

• Physical Change: A change that affects the form of a chemical substance, but not its chemical composition.

• Application: Chemistry is essential in medicine, engineering, biology, and environmental science.

Matter & Change

Classification and Properties of Matter

Matter can be classified by its physical state and composition. Changes in matter can be physical or chemical.

• States of Matter: Solid, liquid, gas, and plasma.

• Physical Properties: Characteristics observed without changing the substance (e.g., color, melting point).

• Chemical Properties: Characteristics that determine how a substance reacts (e.g., flammability).

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

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

Measurement & Units

Quantitative Description in Chemistry

Measurements are fundamental in chemistry for expressing quantities and making comparisons.

• SI Units: Standard units used in science (meter, kilogram, second, mole, kelvin, ampere, candela).

• Common Units: Length (m), mass (kg), volume (L), temperature (K or °C).

• Precision: How close repeated measurements are to each other.

• Accuracy: How close a measurement is to the true value.

Conversion Factors

Changing Units in Calculations

Conversion factors are ratios used to express a quantity in different units.

• Definition: A conversion factor is a fraction equal to one, used to convert between units.

• Example: $1~\text{inch} = 2.54~\text{cm}$, so $\frac{2.54~\text{cm}}{1~\text{inch}}$ is a conversion factor.

• Application: Used in dimensional analysis to solve problems.

Uncertainty & Precision

Limits of Measurement

All measurements have some degree of uncertainty due to limitations of instruments and techniques.

• Uncertainty: The doubt that exists about the result of any measurement.

• Precision: The reproducibility of measurements.

• Accuracy: The closeness of a measurement to the true value.

• Reporting: Measurements should be reported with the correct number of significant figures.

Dimensional Analysis

Problem Solving with Units

Dimensional analysis is a method for converting between units using conversion factors.

• Process: Multiply by conversion factors so that units cancel appropriately.

• Example: To convert 10 inches to centimeters: $10~\text{inches} \times \frac{2.54~\text{cm}}{1~\text{inch}} = 25.4~\text{cm}$

• Application: Used in stoichiometry, concentration calculations, and more.

Significant Figures

Expressing Measurement Reliability

Significant figures indicate the precision of a measured or calculated quantity.

• Rules:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros in a decimal number are significant.

• Example: 0.00450 has three significant figures.

• Application: Used in reporting measurements and calculations.

Classifications & States of Matter

Types and Forms of Matter

Matter can be classified by its composition and state.

• Pure Substances: Elements and compounds with fixed composition.

• Mixtures: Physical combinations of substances (homogeneous and heterogeneous).

• States: Solid (fixed shape/volume), liquid (fixed volume, variable shape), gas (variable shape/volume).

Classification	Description	Example
Element	Pure substance, one type of atom	Oxygen (O2)
Compound	Pure substance, two or more types of atoms	Water (H2O)
Mixture	Physical blend of substances	Air, salt water

Atomic Structure

Components of the Atom

Atoms are the basic units of matter, composed of protons, neutrons, and electrons.

• Proton: Positively charged particle in the nucleus.

• Neutron: Neutral particle in the nucleus.

• Electron: Negatively charged particle outside the nucleus.

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

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

Example: Carbon-12 has 6 protons, 6 neutrons, and 6 electrons.

Dalton’s Atomic Theory

Foundations of Modern Atomic Theory

John Dalton proposed the first modern atomic theory in the early 19th century.

• All matter is composed of indivisible atoms.

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

• Atoms combine in simple whole-number ratios to form compounds.

• Chemical reactions involve rearrangement of atoms; atoms are not created or destroyed.

Example: Water (H2O) is formed by combining two hydrogen atoms and one oxygen atom.

Isotopes & the Nucleus

Variations in Atomic Structure

Isotopes are atoms of the same element with different numbers of neutrons.

• Isotope: Same atomic number, different mass number.

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

• Nucleus: Central part of atom containing protons and neutrons.

Application: Isotopes are used in radiometric dating and medical imaging.

Periodic Table

Organization of Elements

The periodic table arranges elements by increasing atomic number and groups elements with similar properties.

• Groups: Vertical columns with similar chemical properties.

• Periods: Horizontal rows.

• Metals, Nonmetals, Metalloids: Classification based on properties.

• Example: Group 1: Alkali metals; Group 17: Halogens.

Compounds

Formation and Properties

Compounds are substances formed from two or more elements chemically combined in fixed ratios.

• Types: Ionic compounds (metal + nonmetal), molecular compounds (nonmetal + nonmetal).

• Properties: Compounds have properties different from their constituent elements.

• Example: Sodium chloride (NaCl) is formed from sodium and chlorine.

Chemical Formulas

Representation of Compounds

Chemical formulas show the types and numbers of atoms in a compound.

• Empirical Formula: Simplest whole-number ratio of atoms.

• Molecular Formula: Actual number of atoms of each element.

• Example: Glucose: Empirical formula CH2O, molecular formula C6H12O6.

Ionic & Molecular Bonding

Types of Chemical Bonds

Chemical bonds hold atoms together in compounds. The two main types are ionic and covalent (molecular) bonds.

• Ionic Bond: Transfer of electrons from metal to nonmetal, forming ions.

• Covalent Bond: Sharing of electrons between nonmetals.

• Example: NaCl (ionic), H2O (covalent).

Bond Type	Formation	Properties	Example
Ionic	Transfer of electrons	High melting point, conducts electricity when dissolved	NaCl
Covalent	Sharing of electrons	Low melting point, does not conduct electricity	H2O

Nomenclature (Naming Compounds)

Systematic Naming of Chemical Compounds

Nomenclature is the system of naming chemical compounds according to specific rules.

• Ionic Compounds: Name the cation first, then the anion (e.g., NaCl: sodium chloride).

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

• Acids: Special rules for naming acids (e.g., HCl: hydrochloric acid).

Example: N2O5 is dinitrogen pentoxide.