Chemistry Basics—Matter and Measurement: Study Guide for Chapter 1

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This section introduces the fundamental classification of matter, which is essential for understanding chemical properties and reactions.

Matter is anything that occupies space and has mass.
There are two main types of matter: pure substances and mixtures.
Pure substances are further classified as elements (made of one type of atom) or compounds (made of two or more elements chemically joined).
Mixtures are combinations of substances and can be separated into their components.
Homogeneous mixtures have uniform composition throughout (e.g., salt water).
Heterogeneous mixtures have non-uniform composition (e.g., salad).
Atoms are the smallest units of matter retaining unique properties.

The periodic table organizes all known elements and provides information about their properties and relationships.

The periodic table lists all elements, each represented by a unique chemical symbol (e.g., H for hydrogen, Na for sodium).
Symbols may derive from English or Latin names (e.g., Na from 'natrium').
Vertical columns are called groups and contain elements with similar chemical behaviors.
Groups are numbered 1–18 (IUPAC system); main-group elements are designated 'A', transition elements 'B'.
Horizontal rows are called periods, numbered 1–7.
The staircase line separates metals from nonmetals; elements bordering the line (except Al) are metalloids.
Essential elements for health include carbon, hydrogen, oxygen, nitrogen, sodium, magnesium, potassium, calcium, chlorine (macronutrients), and iodine, fluorine, iron, zinc (micronutrients).
Compounds are pure substances containing two or more elements in specific ratios, represented by chemical formulas (e.g., for water, for table salt).

Matter can undergo physical or chemical changes, which are fundamental to chemical reactions and processes.

Physical change: Alters the state or appearance of matter without changing its identity (e.g., melting ice).
Chemical change: Changes the chemical identity of a substance, resulting in a chemical reaction (e.g., burning wood).
Chemical equations represent reactions, showing reactants and products, and their physical states: (s)olid, (l)iquid, (g)as, (aq)ueous.
Equations must be balanced so the number of atoms is equal on both sides, illustrating the law of conservation of mass.
Balancing steps:
1. Examine the equation for balance.
2. Add coefficients to balance one element at a time.
3. Check for the smallest set of coefficients.

Mathematical concepts are central to chemistry, enabling precise measurement and calculation.

The Système International d’Unités (SI) is the modern metric system.
Standard units: kilogram (kg) for mass, liter (L) for volume, meter (m) for length.
Prefixes modify units by powers of 10 (e.g., milli-, centi-, kilo-).
Conversion factors and dimensional analysis are used to convert between units.
Steps for dimensional analysis:
1. Determine desired units.
2. Establish given information.
3. Choose appropriate conversion factors.
4. Solve the problem.
Significant figures reflect the precision of measurements; all nonzero digits are significant, zeros may or may not be significant depending on their position.
Rules for calculations:
- Addition/Subtraction: Match the least number of decimal places.
- Multiplication/Division: Match the least number of significant digits.
Scientific notation expresses numbers as , where is the coefficient (1 ≤ C < 10) and is the exponent.
Percent (%) means part per hundred; convert fractions or decimals to percent by multiplying by 100.

Understanding the properties and measurements of matter is crucial for laboratory and clinical applications.

Mass is the amount of material in an object, measured in grams (g).
Volume is the space occupied by matter, measured in milliliters (mL) or cubic centimeters (cm3); 1 mL = 1 cm3.
Density () is the ratio of mass () to volume (): .
Specific gravity is the ratio of a sample's density to water's density; it is unitless and measured with a refractometer.
Temperature is measured in Fahrenheit (°F), Celsius (°C), or Kelvin (K); Kelvin is the SI unit.
Conversion formulas:
Energy is the capacity to do work or supply heat; measured in joules (J) or calories (cal).
Specific heat is the amount of heat needed to raise 1 g of a substance by 1°C.
States of matter:
- Solid: Definite shape and volume; particles tightly packed.
- Liquid: Definite volume, takes shape of container; particles move freely.
- Gas: No definite shape or volume; particles far apart and move rapidly.

Accurate and precise measurement is vital in chemistry and health sciences.

Accuracy means measurements are close to the true value.
Precision means measurements are consistent with each other.
SI/metric units are standard, but U.S. customary units are also used in health care.
Lab reports use metric units; deciliter (dL) = 0.1 L, millimole (mmol) = 0.001 mol, milliequivalent (mEq) for electrolytes.
Dosage calculations involve conversion factors to ensure correct units.
Drop units (gtt) are used for IV medication delivery; drop factor depends on IV tubing diameter.
Percent in health:
- Percent active ingredient in medicines.
- Percent of adult dose for children.
- Percent Daily Value (%DV) in nutrition labeling.

Classify matter as pure substance or mixture; distinguish homogeneous and heterogeneous mixtures.
Understand periodic table organization; identify groups, periods, metals, nonmetals, and metalloids.
Represent changes in matter; distinguish physical and chemical changes; balance chemical equations.
Apply math concepts: unit conversions, significant figures, scientific notation, percent calculations.
Measure mass, volume, density, specific gravity; convert temperatures; understand energy and states of matter.
Apply measurements to health: accuracy, precision, unit conversions, dosage calculations, percent in health.