Chapter 1: Chemistry Basics—Matter and Measurement (GOB Chemistry Study Notes)

Classifying Matter

Pure Substances and Mixtures

Matter is anything that occupies space and has mass. It can be classified based on its composition into pure substances and mixtures.

• Pure Substance: Matter made up of only one type of substance, represented by a single chemical formula or symbol.

• Element: The simplest type of pure substance, consisting of only one type of atom.

• Compound: A pure substance composed of two or more elements chemically combined in fixed ratios.

• Mixture: A physical combination of two or more substances that can be separated by physical means.

• Homogeneous Mixture: Has a uniform composition throughout (e.g., salt water).

• Heterogeneous Mixture: Composition is not uniform and varies throughout (e.g., salad, sand in water).

Example: Air is a homogeneous mixture, while a salad is a heterogeneous mixture.

Classification Table

Type	Subcategory	Example
Pure Substance	Element	Copper (Cu)
Pure Substance	Compound	Water (H2O)
Mixture	Homogeneous	Brass (copper and zinc atoms)
Mixture	Heterogeneous	Water with sand

Elements, Compounds, and the Periodic Table

The Periodic Table

The periodic table organizes all known elements according to their properties. Each element is represented by a chemical symbol, often derived from its English or Latin name (e.g., Na for sodium, Au for gold).

• Groups: Vertical columns with similar chemical properties. Labeled 1–18 (IUPAC) or with A/B designations (main-group/transition elements).

• Periods: Horizontal rows numbered 1–7.

• Metals, Nonmetals, Metalloids: The staircase line separates metals (left) from nonmetals (right); elements bordering the line (except Al) are metalloids.

Example: Carbon (C), hydrogen (H), oxygen (O), and nitrogen (N) are essential elements for life.

Essential Elements for Health

Type	Examples	Daily Requirement
Macronutrients	Sodium, Magnesium, Potassium, Calcium, Chlorine	>100 mg
Micronutrients	Iodine, Fluorine, Iron, Zinc	<100 mg

Chemical Formulas

Chemical formulas indicate the elements present in a compound and the number of atoms of each. For example, water is H2O (2 hydrogen, 1 oxygen), and sodium chloride is NaCl (1 sodium, 1 chlorine).

How Matter Changes

Physical and Chemical Changes

• Physical Change: Alters the form or appearance of matter without changing its identity (e.g., melting ice).

• Chemical Change: Alters the chemical identity of a substance; a chemical reaction occurs (e.g., burning wood).

Chemical Equations

Chemical equations represent chemical reactions, showing reactants and products. The arrow ($\rightarrow$) means "react to form." Physical states are indicated: (s) solid, (l) liquid, (g) gas, (aq) aqueous.

Example:

$\mathrm{C} (s) + \mathrm{O}_2 (g) \rightarrow \mathrm{CO}_2 (g)$

Balancing Chemical Equations

• Atoms must be balanced on both sides (law of conservation of mass).

• Add coefficients to balance elements.

• Check that coefficients are the smallest possible whole numbers.

Math Counts

SI Units and Metric System

• Mass: Kilogram (kg)

• Volume: Liter (L)

• Length: Meter (m)

• Prefixes modify units by powers of 10 (e.g., milli-, centi-, kilo-).

Unit Conversions and Dimensional Analysis

• Use conversion factors to change units (e.g., $1\ \mathrm{dL} = 0.1\ \mathrm{L}$).

• Dimensional analysis involves multiplying by conversion factors to cancel units.

Example: To convert 10 dL to L: $10\ \mathrm{dL} \times \frac{0.1\ \mathrm{L}}{1\ \mathrm{dL}} = 1\ \mathrm{L}$

Significant Figures

• All nonzero digits are significant.

• Zeros are significant if between nonzero digits or after a decimal point.

• Exact numbers (from counting or definitions) have infinite significant figures.

Rules for Calculations

• Addition/Subtraction: Result has as many decimal places as the measurement with the fewest decimal places.

• Multiplication/Division: Result has as many significant digits as the measurement with the fewest significant digits.

Rounding

• If the first digit to be dropped is 4 or less, leave the last retained digit unchanged.

• If it is 5 or greater, increase the last retained digit by 1.

• Round only at the end of multi-step calculations.

Scientific Notation

Scientific notation expresses numbers as $C \times 10^n$, where $1 \leq C < 10$ and $n$ is an integer.

• Positive exponent: number > 1

• Negative exponent: number between 0 and 1

Percentages

Percent means "per hundred."

• $\text{percent} = \frac{\text{part}}{\text{whole}} \times 100$

• Convert fractions or decimals to percent by multiplying by 100.

Matter: The "Stuff" of Chemistry

Mass and Weight

• Mass: Amount of matter in an object, measured in grams (g) or kilograms (kg).

• Weight: Force of gravity on an object; varies with location but is proportional to mass on Earth's surface.

Volume

• Volume is the amount of space occupied by matter.

• Common units: milliliter (mL), liter (L), cubic centimeter (cm3 or cc).

• 1 mL = 1 cm3

Density and Specific Gravity

• Density (d): Ratio of mass to volume.

$d = \frac{m}{V}$

• Density of water at 4°C is 1.00 g/mL.

• Specific Gravity: Ratio of the density of a sample to the density of water (unitless).

$\text{Specific gravity} = \frac{\text{density of sample}}{\text{density of water}}$

Temperature Scales

• Celsius (°C): Used worldwide and in science.

• Fahrenheit (°F): Used in the United States.

• Kelvin (K): SI unit for temperature; absolute scale.

Conversions:

• $\text{Kelvin} = \text{Celsius} + 273$

• $\text{Celsius} = (\text{Fahrenheit} - 32) \times \frac{5}{9}$

• $\text{Fahrenheit} = (\text{Celsius} \times \frac{9}{5}) + 32$

Body Temperature and Health

• Normal body temperature: 98.6°F (37.0°C).

• Hyperthermia: >40.0°C (104°F); can cause serious health issues.

• Hypothermia: <35°C (95°F); can be life-threatening.

Energy and Specific Heat

• Energy: Capacity to do work or supply heat.

• Potential Energy: Stored energy.

• Kinetic Energy: Energy of motion.

• Law of Conservation of Energy: Energy is neither created nor destroyed.

• Joule (J): SI unit of energy.

• Calorie (cal): Energy to raise 1 g of water by 1°C; $1\ \text{cal} = 4.184\ \text{J}$.

• Calorie (Cal): Nutritional unit; $1\ \text{Cal} = 1000\ \text{cal}$.

Specific Heat

• Amount of heat needed to raise the temperature of 1 g of a substance by 1°C.

$\text{Specific heat (SH)} = \frac{\text{heat}}{\text{grams} \times \Delta T}$

• Metals have low specific heat; water has high specific heat.

States of Matter

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

• Liquid: Definite volume, takes shape of container; particles less orderly, move freely.

• Gas: No definite shape or volume; particles far apart, move rapidly.

Property	Solid	Liquid	Gas
Shape	Definite	Adopts container	Adopts container
Volume	Definite	Definite	Fills container
Particle Arrangement	Closely packed	Less packed	Far apart
Energy	Lowest	Intermediate	Highest

Measuring Matter

Accuracy and Precision

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

• Precision: How close repeated measurements are to each other.

• Best practice: Take multiple measurements and average them.

Health and Measurement Units

• SI/metric units are standard, but U.S. customary units are also used in healthcare.

• Common conversions: 1 dL = 0.1 L = 100 mL; 1 mmol = 0.001 mol; 1 mEq for electrolytes.

• Body weight: pounds (U.S.), kilograms (pharmaceutical dosing).

Dosage Calculations

• Determine final units needed.

• Identify given information and conversion factors.

• Set up the equation so all units cancel except the desired unit.

Units in Medication Delivery

• Medications may be measured in drops per milliliter (gtt/mL).

• Drop factor depends on IV tubing diameter.

Percent in Health

• Active ingredient percent: Indicates concentration in medications.

• Percent of adult dose: Used for pediatric dosing.

• Nutrition labels: Show % Daily Value (%DV) for nutrients.