Matter, Measurement, and Problem Solving: Foundations of General Chemistry

Matter, Measurement, and Problem Solving

Properties of Matter

The properties of matter are determined by the properties and arrangement of molecules and atoms. Understanding these properties is fundamental to the study of chemistry.

• Atoms: The smallest unit of matter that retains the properties of an element. Atoms are the fundamental building blocks of ordinary matter.

• Molecules: Groups of atoms bonded together in specific geometric arrangements.

• Free atoms: Rare in nature; most atoms exist as part of molecules or compounds.

Chemists seek to understand the behavior of matter by studying the behavior of atoms and molecules. Adjustments on the small scale can impact large-scale properties.

The Scientific Approach

The scientific approach is based on observation and experiment, providing a systematic method for understanding nature and its behavior.

• Observation: The starting point, often driven by curiosity.

• Hypothesis: A tentative interpretation or explanation of observations. Hypotheses must be testable and falsifiable.

• Experimentation: Used to test hypotheses and establish cause-and-effect relationships.

• Law: A summary of past observations that predicts future ones. Laws cannot be broken and are always true within their domain.

• Theory: A general explanation for characteristics and behavior of nature, validated by experiments. Theories are models for how nature works and can never be conclusively proven.

• Data: Can be qualitative (descriptive) or quantitative (numerical).

Example: The Law of Conservation of Mass states that mass is neither created nor destroyed in a chemical reaction.

Classification of Matter

Matter is anything that occupies space (has volume) and has mass. It can be classified by its physical form and composition.

• Physical Form:

  • Solid: Definite shape and volume; atoms are closely packed and vibrate in place.

  • Liquid: Definite volume but indefinite shape; atoms are close but can move past each other.

  • Gas: Indefinite shape and volume; atoms are far apart and move freely.

• Composition:

  • Pure Substance: Made up of only one component. Can be an element (cannot be broken down) or a compound (composed of two or more elements in fixed proportions).

  • Mixture: Composed of two or more components in varying proportions. Can be:

    • Heterogeneous: Composition varies from one region to another (e.g., cereal, trail mix).

    • Homogeneous: Uniform composition throughout (e.g., coffee, air).

Separating Mixtures

Mixtures can be separated based on differences in their physical properties.

• Heterogeneous Mixtures: Filtration (pouring a liquid through filter paper).

• Homogeneous Mixtures: Distillation (boiling off the more volatile liquid and condensing it).

Physical and Chemical Changes

Changes in matter can be classified as physical or chemical.

• Physical Change: Alters the state or appearance of a substance without changing its composition (e.g., melting, boiling).

• Chemical Change: Alters the composition, resulting in new substances (e.g., rusting, combustion).

Example of Chemical Change:

$\mathrm{CH_4(g) + 2\ O_2(g) \rightarrow CO_2(g) + 2\ H_2O(g)}$

Properties of Matter

• Physical Property: Displayed without changing composition (e.g., smell, color, melting point, boiling point).

• Chemical Property: Displayed only by changing composition via chemical change (e.g., flammability, corrosiveness, acidity, toxicity).

Energy in Chemistry

Energy is the capacity to do work. It is involved in all physical and chemical changes.

• Work: Action of a force through a distance ($W = F \times d$).

• Kinetic Energy: Associated with motion of objects.

• Potential Energy: Associated with position or composition.

• Thermal Energy: Associated with temperature (type of kinetic energy).

Systems with high potential energy tend to change in ways that lower their potential energy, transferring energy to their surroundings.

Measurement in Chemistry

Measurements are essential for quantifying observations in chemistry. The International System of Units (SI) is used for standardization.

• Length: Meter (m)

• Mass: Kilogram (kg)

• Time: Second (s)

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

Temperature Conversions:

• $^\circ C = \frac{(^\circ F - 32)}{1.8}$

• $K = ^\circ C + 273.15$

Metric Prefix Multipliers

Metric prefixes are used to express multiples or fractions of units.

Prefix	Symbol	Multiplier
kilo	k	103
centi	c	10-2
milli	m	10-3
micro	μ	10-6
nano	n	10-9
pico	p	10-12
femto	f	10-15
deci	d	10-1
deca	da	101
hecto	h	102
mega	M	106
giga	G	109
tera	T	1012

Derived Units and Density

• Volume: Liter (L), cubic centimeters (cm3), or milliliters (mL).

• Density: Ratio of mass to volume of a substance.

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

  • Intensive property: independent of amount of substance.

  • Extensive property: dependent on amount of substance (e.g., mass).

Significant Figures and Measurement Reliability

Significant figures (sig figs) reflect the precision of a measurement. Every digit is certain except the last, which is estimated.

• All nonzero digits are significant.

• Interior zeros (between nonzero digits) are significant.

• Leading zeros are not significant.

• Trailing zeros after a decimal point are significant.

• Trailing zeros before a decimal point are not significant.

• Trailing zeros before a decimal point and after a nonzero digit are significant.

• Use scientific notation to clarify significant figures.

Exact Numbers

• Have unlimited significant figures (e.g., counting discrete objects, defined conversions).

• Examples: 1 L = 1000 mL, 1 kg = 1000 g.

Calculations with Significant Figures

• Multiplication/Division: Result has the same number of sig figs as the value with the fewest sig figs.

• Addition/Subtraction: Result has the same number of decimal places as the value with the fewest decimal places.

• Rounding: If the last digit dropped is 4 or less, round down; if 5 or more, round up.

Accuracy and Precision

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

• Precision: How close a series of measurements are to one another.

• Reproducibility: How reproducible the measurements are.

Solving Chemical Problems: Unit Conversion

Unit conversion is essential for solving chemical problems. Dimensional analysis is used to convert between units.

• Always include units in calculations.

• Cancel out units you don't want.

• General formula: Given unit × (desired unit / given unit) = desired unit

Example: To convert 10 cm to meters: $10\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.1\ \text{m}$