Foundations of Chemistry: Matter, Measurement, and Atomic Structure

The Scientific Approach to Knowledge

Empirical Nature of Science

The scientific method relies on empirical evidence—knowledge gained through observation and experimentation. This approach ensures that scientific knowledge is based on reproducible and objective data.

• Observation: Descriptions about the characteristics or behavior of nature, also known as data.

• Example: Antoine Lavoisier observed that the total mass of material remains unchanged during combustion, leading to the law of conservation of mass.

• Hypothesis: A tentative interpretation or explanation of observations.

• 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 locked in a fixed arrangement but vibrate in place.

• Liquid: Atoms or molecules are close together but can move past one another; liquids take the shape of their container.

• Gas: Atoms or molecules are far apart and move rapidly; gases fill the entire volume of their container.

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 with variable composition.

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 (e.g., salt and sand mixture).

• Homogeneous Mixture (Solution): Uniform composition throughout; appears as a single substance (e.g., saltwater).

Separation of Mixtures

• Decanting: Separates mixtures like sand and water by pouring off the liquid.

• Filtration: Separates an insoluble solid from a liquid.

• Distillation: Separates homogeneous mixtures based on differences in boiling points.

Physical and Chemical Changes

Physical Changes

• Alter only the state or appearance, not the composition.

• Examples: Melting ice, sublimation of dry ice, dissolving sugar in water.

Chemical Changes

• Alter the composition of matter; atoms rearrange to form new substances.

• Examples: Rusting of iron, burning propane.

Physical vs. Chemical Properties

• Physical Property: Observed without changing composition (e.g., odor, color, melting point, density).

• Chemical Property: Observed only by changing composition via a chemical reaction (e.g., flammability, acidity, toxicity).

Numbers and Chemistry

Quantitative Aspects

• Many chemical concepts are quantitative and require measurement.

• Key concepts: units of measurement, uncertainty, significant figures, dimensional analysis.

Units of Measurement

• Metric System: Used worldwide.

• English System: Used in the United States.

• SI Units (Système International d’Unités): Standard system for scientific measurements.

Quantity	SI Unit	Symbol
Length	meter	m
Mass	kilogram	kg
Time	second	s
Temperature	kelvin	K
Amount of substance	mole	mol

Important SI Units

• Meter (m): Base unit of length; 1 m = 39.3 inches.

• Kilogram (kg): Base unit of mass; 1 kg = 2.205 lb.

• Gram (g): 1 g = 1/1000 kg.

• Second (s): Base unit of time.

• Kelvin (K): Base unit of temperature; absolute zero is 0 K.

Temperature Scales and Conversions

• Celsius (°C): Water freezes at 0°C and boils at 100°C.

• Kelvin (K): Absolute temperature scale; K = °C + 273.15.

• Fahrenheit (°F): Used in the US; °F = 1.8 × °C + 32.

• Conversion Formulas:

$C = \frac{F - 32}{1.8}$ $K = C + 273.15$

SI Prefixes

• Prefixes indicate powers of ten (e.g., kilo = $10^3$, milli = $10^{-3}$, pico = $10^{-12}$).

Derived Units: Volume and Density

• Volume: Space occupied by a substance; SI derived unit is cubic meter ($m^3$), but liter (L) and milliliter (mL) are commonly used.

• 1 L = 1 dm3; 1 mL = 1 cm3 = 1 cc (cubic centimeter).

• Density: Mass per unit volume; $\text{Density} = \frac{\text{mass}}{\text{volume}}$

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

• Density determines whether a substance will sink or float in another substance.

Problem Solving and Dimensional Analysis

General Problem Solving Strategy

• Identify the starting point (given information).

• Identify the end point (what you must find).

• Devise a conceptual plan to connect the two using known relationships.

Dimensional Analysis

• Technique for converting between units using conversion factors.

• When converting units raised to a power, raise both the number and the unit to that power (e.g., $1\ \text{in}^2 = (2.54\ \text{cm})^2$).

Atomic Structure and Subatomic Particles

Discovery of Radioactivity

• Ernest Rutherford discovered three types of radiation:

  • Alpha (α) particles: positively charged

  • Beta (β) particles: negatively charged (like electrons)

  • Gamma (γ) rays: uncharged

Early Atomic Models

• Plum Pudding Model (J.J. Thomson): Atom as a positive sphere with embedded electrons.

• Rutherford's Gold Foil Experiment: Showed that atoms have a small, dense, positively charged nucleus with electrons around the outside; most of the atom is empty space.

Subatomic Particles

• Proton: Positive charge, mass ≈ 1 amu.

• Neutron: No charge, mass ≈ 1 amu.

• Electron: Negative charge, mass ≈ $5.486 \times 10^{-4}$ amu (negligible compared to protons and neutrons).

Particle	Symbol	Charge	Mass (amu)
Proton	p+	+1	1.0073
Neutron	n0	0	1.0087
Electron	e-	-1	0.00055

Atomic Mass Unit (amu)

• 1 amu = $1.66054 \times 10^{-24}$ g

• 1 g = $6.02214 \times 10^{23}$ amu$\ (\text{Avogadro's number})$

Atomic Number, Mass Number, and Isotopes

Atomic Number (Z)

• Number of protons in the nucleus; defines the element.

• In a neutral atom, number of protons = number of electrons.

Mass Number (A)

• Total number of protons and neutrons in the nucleus.

• Symbolized as A; $A = \text{number of protons} + \text{number of neutrons}$

Isotopes

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

• Isotopes are represented as $^{A}_{Z}X$ where X is the chemical symbol, A is the mass number, and Z is the atomic number.

Atomic Mass

• Weighted average of all naturally occurring isotopes of an element.

• Measured using a mass spectrometer.

• Calculation:

$\text{Atomic mass} = \sum (\text{fractional abundance} \times \text{isotope mass})$

Ions

Formation of Ions

• Atoms can gain or lose electrons during chemical changes, forming ions.

• Cation: Positively charged ion (e.g., Na+).

• Anion: Negatively charged ion (e.g., F-).

The Periodic Table and Periodicity

The Periodic Law

• When elements are arranged in order of increasing mass, certain sets of properties recur periodically (Mendeleev's periodic law).

Reading the Periodic Table

• Each element's box lists the atomic number (above the symbol) and atomic weight (below the symbol).

• Elements are organized into periods (rows) and groups (columns).

Classification of Elements

• Metals: Left side; shiny, conduct heat/electricity, mostly solids.

• Nonmetals: Right side; can be solids, liquids, or gases.

• Metalloids: Border the stair-step line; properties intermediate between metals and nonmetals (except Al, Po, At).

Groups and Periodic Trends

• 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 (Group 6A): Gain two electrons to form 2- ions.

Periodicity

• Elements show repeating patterns of chemical properties and reactivity across periods and groups.

Additional info: Some explanations and examples were expanded for clarity and completeness, including the inclusion of tables and formulas for SI units, subatomic particles, and atomic mass calculation.