Foundations of Chemistry: Matter, Measurement, and Atomic Structure

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The Scientific Approach to Knowledge

Empirical Methods and the Scientific Method

The scientific approach relies on empirical evidence, meaning knowledge is gained through observation and experimentation. This process is foundational to chemistry and all sciences.

Observation: Descriptions or measurements about the characteristics or behavior of nature. Also referred to as data.
Example: Antoine Lavoisier observed that the total mass of material in a container did not change during combustion, leading to the law of conservation of mass.
Hypothesis: A tentative interpretation or explanation of observations. Hypotheses are tested through experiments.
Scientific Law: A brief statement summarizing past observations and predicting future ones. Example: Law of Conservation of Mass – "In a chemical reaction, matter is neither created nor destroyed."

The Classification of Matter

States of Matter

Matter is anything that occupies space and has mass. It can be classified by its physical state and composition.

Solid: Atoms or molecules are closely packed in fixed positions, vibrating but not moving past each other. Solids have definite shape and volume.
Liquid: Atoms or molecules are close but can move past each other, giving liquids a definite volume but no fixed shape.
Gas: Atoms or molecules are far apart and move freely, resulting in no definite shape or volume.

Structure Determines Properties: The arrangement of atoms or molecules in each state leads to different physical properties.

Classification by Composition

Matter can also be classified based on its composition:

Pure Substance: Composed of only one type of component with invariant composition.
Mixture: Composed of two or more components in variable proportions.

Types of Pure Substances

Element: Cannot be chemically broken down into simpler substances.
Compound: Composed of two or more elements in fixed, definite proportions.

Types of Mixtures

Heterogeneous Mixture: Composition varies from one region to another; components are visibly distinct. Example: Salt and sand mixture.
Homogeneous Mixture (Solution): Uniform composition throughout; appears as a single substance. Example: Saltwater.

Separation of Mixtures

Mixtures can be separated by exploiting differences in physical or chemical properties.
Decanting: Separates a mixture of sand and water.
Filtration: Separates an insoluble solid from a liquid.
Distillation: Separates components of a homogeneous liquid mixture based on boiling points.

Physical and Chemical Changes

Physical Change: Alters only the state or appearance, not composition. Examples: Melting ice, subliming dry ice, dissolving sugar in water.
Chemical Change: Alters the composition of matter; atoms rearrange to form new substances. Examples: Rusting of iron, burning propane.

Physical and Chemical Properties

Physical Property: Observed without changing composition. Examples: Odor, taste, color, melting point, boiling point, density.
Chemical Property: Observed only by changing composition via a chemical reaction. Examples: Flammability, acidity, toxicity, corrosiveness.

Numbers and Chemistry

Quantitative Aspects

Many chemical concepts are quantitative and require careful measurement and calculation.
Key concepts include units of measurement, uncertainty, significant figures, and dimensional analysis.

Units of Measurement

SI Units: International System of Units, based on the metric system.
Base Units: Meter (m) for length, kilogram (kg) for mass, second (s) for time, kelvin (K) for temperature.
Metric Prefixes: Used to express multiples or fractions of units (e.g., kilo-, centi-, milli-, micro-, nano-, pico-).

Key SI Base Units

Quantity	Unit	Symbol
Length	meter	m
Mass	kilogram	kg
Time	second	s
Temperature	kelvin	K

Temperature Scales and Conversion

Celsius (°C): Water freezes at 0°C and boils at 100°C.
Kelvin (K): Absolute temperature scale; 0 K is absolute zero.
Fahrenheit (°F): Used mainly in the United States.

Temperature Conversion Formulas:

Derived Units: Volume and Density

Volume: Space occupied by a substance; SI derived unit is cubic meter (), but liter (L) and milliliter (mL) are commonly used.
1 L = 1 dm3; 1 mL = 1 cm3
Density: Mass per unit volume.
Density determines whether a substance will sink or float in another.

Problem Solving and Dimensional Analysis

Identify the starting point (given information) and the end point (what you must find).
Devise a conceptual plan to connect the two using known relationships and conversion factors.
When converting units raised to a power, raise both the number and the unit to that power. Example: To convert to , square the conversion factor.

Atomic Structure and the Periodic Table

Discovery of Subatomic Particles

Radioactivity: Discovered by Ernest Rutherford; three types: alpha (α, positive), beta (β, negative), gamma (γ, neutral).
Plum Pudding Model: J.J. Thomson's model with electrons embedded in a positive sphere.
Rutherford's Gold Foil Experiment: Showed that atoms have a small, dense, positively charged nucleus with electrons around it; most of the atom is empty space.

Subatomic Particles

Particle	Symbol	Charge	Approximate Mass (amu)
Proton	p+	+1	1
Neutron	n0	0	1
Electron	e-	-1	0.0005

Protons and neutrons are found in the nucleus; electrons are found outside the nucleus.
Proton and electron charges are equal in magnitude but opposite in sign.

Atomic Number, Mass Number, and Isotopes

Atomic Number (Z): Number of protons in the nucleus; defines the element.
Mass Number (A): Total number of protons and neutrons in the nucleus.
Isotopes: Atoms of the same element with different numbers of neutrons (and thus different mass numbers).

Isotope Notation: where X is the chemical symbol, A is the mass number, and Z is the atomic number.

Relationship:

Atomic Mass and Avogadro's Number

Atomic mass is the weighted average of all isotopes of an element, based on their natural abundance.
Measured using a mass spectrometer.
1 atomic mass unit (amu) = g
1 g = amu (Avogadro's number)

Ions: Cations and Anions

Atoms can gain or lose electrons during chemical changes, forming ions.
Cation: Positively charged ion (e.g., Na+), formed by losing electrons.
Anion: Negatively charged ion (e.g., F-), formed by gaining electrons.

The Periodic Table and Periodicity

The Periodic Law

When elements are arranged in order of increasing mass (now atomic number), certain sets of properties recur periodically.

Reading the Periodic Table

Each element's box lists the atomic number (above the symbol) and atomic weight (below the symbol).
Elements are grouped into periods (rows) and groups (columns).

Classification of Elements

Metals: Left side; shiny, conduct heat and electricity, mostly solids (except mercury).
Nonmetals: Right side; can be solids, liquids, or gases at room temperature.
Metalloids: Border the stair-step line; have properties intermediate between metals and nonmetals (except Al, Po, At).

Groups and Ion Formation

Alkali Metals (Group 1A): Lose one electron to form 1+ ions.
Alkaline Earth Metals (Group 2A): Lose two electrons to form 2+ ions.
Halogens (Group 7A): Gain one electron to form 1- ions.
Oxygen Family Nonmetals (Group 6A): Gain two electrons to form 2- ions.
Main-group metals tend to lose electrons (form cations); main-group nonmetals tend to gain electrons (form anions), achieving the same number of electrons as the nearest noble gas.

Additional info: The notes above cover foundational concepts from the first two chapters of a general chemistry course, including matter, measurement, atomic structure, and the periodic table. These concepts are essential for understanding chemical reactions, bonding, and properties of substances in later chapters.