Matter and Measurements: Foundations of Chemistry for Healthcare Professionals

Matter and Measurements

Chemistry: The Science of Everyday Experience

Chemistry is the study of matter—its composition, properties, and transformations. Matter is anything that has mass and occupies space. Chemistry is fundamental to understanding the materials and processes encountered in healthcare and daily life.

States of Matter

Matter exists in three primary states: solid, liquid, and gas. Each state has distinct physical properties based on the arrangement and movement of particles.

Solid State

• Definite volume and shape

• Particles are closely packed in a regular pattern

Liquid State

• Definite volume but no definite shape

• Takes the shape of its container

• Particles are close together but can move past one another

Gas State

• No definite shape or volume

• Assumes the shape and volume of its container

• Particles are far apart and move randomly

Chemical and Physical Properties of Matter

Properties of matter are classified as chemical or physical:

• Chemical properties: Observed only when matter is changed into a new substance (e.g., flammability, reactivity).

• Physical properties: Observed without changing the composition (e.g., color, shape, odor, boiling point, melting point, solubility).

Chemical and Physical Changes

• Chemical changes: Accompanied by a change in composition (e.g., burning paper, fizzing of vinegar and baking soda).

• Physical changes: Occur without a change in composition (e.g., freezing, melting, evaporation).

Atoms and Molecules

Dalton’s Atomic Theory (1808)

• All matter is made up of tiny particles called atoms.

• Elements are made up of identical atoms.

• Compounds are combinations of atoms of two or more elements.

• Atoms are rearranged in chemical reactions but are never created nor destroyed.

Molecules vs. Atoms

• Molecule: Smallest particle of a pure substance capable of stable independent existence.

• Atom: Basic particle that makes up molecules.

Classification of Matter

Matter is classified based on its chemical and physical properties:

Pure Substances

• Have constant composition and fixed properties.

• Classified as elements or compounds.

• Example: Pure water always contains the same proportions of hydrogen and oxygen.

Elements

• Made up of homoatomic molecules or individual atoms of the same kind.

• Examples: Oxygen gas (O2), copper metal (Cu).

Compounds

• Made up of heteroatomic molecules or ions of two or more different kinds.

• Examples: Water (H2O), table salt (NaCl).

Mixtures

• Composition and properties can vary.

• Two types: homogeneous (solutions) and heterogeneous.

Heterogeneous Mixtures

• Properties depend on the location of the sample.

• Examples: Blood, oil and vinegar, smog, soda.

Homogeneous Mixtures

• Also called solutions.

• Properties are uniform throughout the mixture.

• Example: Sugar dissolved in water.

Summary Table: Classification of Matter

Type	Definition	Example
Element	Pure substance, one kind of atom	Copper, Oxygen
Compound	Pure substance, two or more kinds of atoms	Water, Table salt
Homogeneous Mixture	Uniform composition	Sugar water, air
Heterogeneous Mixture	Non-uniform composition	Blood, oil and vinegar

Measurement

Measurements are fundamental in chemistry and consist of a number and a unit. Units are standardized and measurements are made using devices such as rulers, balances, and graduated cylinders.

Metric System

• Decimal system; units related by factors of 10.

• Base units: meter (length), gram (mass), liter (volume).

• Prefixes indicate multiples or fractions of base units (e.g., kilo-, centi-, milli-).

Mass vs. Weight

• Mass: Amount of matter in an object; independent of location.

• Weight: Gravitational force acting on an object; depends on location.

Volume

• Base unit: liter (L).

• 1 mL = 1 cm3 = 1 cc.

• Volume formula: $\text{Volume} = \text{Length} \times \text{Width} \times \text{Height}$

Significant Figures

Significant figures reflect the precision of a measurement. Exact numbers result from counting or definitions; inexact numbers result from measurements.

• All non-zero digits are significant.

• Zeros between non-zero digits are significant.

• Leading zeros are not significant.

• Zeros at the end of a decimal are significant.

• Zeros at the end of a non-decimal are not significant.

Rules for Calculations

• Multiplication/Division: Answer has the same number of significant figures as the original number with the fewest significant figures.

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

Scientific Notation

Scientific notation expresses numbers as a coefficient (between 1 and 10) multiplied by a power of 10. Useful for very large or small numbers.

• Example: $2,500 = 2.5 \times 10^3$

• Example: $0.36 = 3.6 \times 10^{-1}$

Problem Solving Using Conversion Factors

Conversion factors are used to convert quantities from one unit to another. They are written as equalities and used as fractions to cancel unwanted units.

• Example: $2.20 \text{ lb} = 1 \text{ kg}$

• To convert 130 lb to kg: $130 \text{ lb} \times \frac{1 \text{ kg}}{2.20 \text{ lb}} = 59 \text{ kg}$

General Steps

1. Identify the original quantity and unit.

2. Choose the appropriate conversion factor.

3. Set up the calculation so unwanted units cancel.

4. Perform the calculation and round to the correct number of significant figures.

Summary

This chapter introduces the foundational concepts of matter, its classification, measurement, and the importance of precision in scientific work. Understanding these basics is essential for further study in chemistry and its applications in healthcare.