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 joined together in fixed ratios.

• Mixture: A combination of two or more substances that are not chemically bonded and can be separated by physical means.

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

• Heterogeneous Mixture: Has a non-uniform composition with visibly different parts (e.g., salad, sand and iron filings).

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

Classification Table

The following table summarizes the classification of matter:

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

Elements, Compounds, and the Periodic Table

The Periodic Table

The periodic table of the elements is an organized chart listing all known elements. Each element is represented by a unique chemical symbol, often derived from its English or Latin name (e.g., Na for sodium, Au for gold).

• Each block on the table contains the element's symbol and numbers indicating atomic number and atomic mass.

• Vertical columns are called groups (or families) and contain elements with similar chemical properties.

• Groups are labeled either with numbers and letters (A for main-group, B for transition elements) or with numbers 1–18 (IUPAC system).

• Horizontal rows are called periods and are numbered 1–7.

• A staircase line separates metals (left) from nonmetals (right); elements bordering the line (except Al) are metalloids.

Example: The symbol for iron is Fe (from Latin ferrum).

Elements Essential for Life

• Macronutrients: Elements needed in amounts greater than 100 mg/day (e.g., sodium, magnesium, potassium, calcium, chlorine).

• Micronutrients: Elements needed in amounts less than 100 mg/day (e.g., iodine, fluorine, iron, zinc).

Chemical Formulas

A chemical formula shows 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 but does not change its identity (e.g., melting ice).

• Chemical Change: Alters the chemical identity of a substance; a new substance is formed (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 as (s)olid, (l)iquid, (g)as, or (aq)ueous.

• Example: $C + O_2 \rightarrow CO_2$

Balancing Chemical Equations

• The number of atoms of each element must be the same on both sides (law of conservation of mass).

• Balance by adding coefficients in front of formulas.

• Steps: (1) Examine the equation, (2) Balance one element at a time, (3) Check for smallest whole-number coefficients.

Math Counts

SI Units and Metric System

• 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-).

Unit Conversions and Dimensional Analysis

• Equivalent units (e.g., 1 dL = 0.1 L) are used as conversion factors to change units.

• Dimensional analysis involves multiplying by conversion factors to cancel units and obtain the desired unit.

• Steps: (1) Determine final units, (2) Identify given information, (3) Choose conversion factors, (4) Solve.

Significant Figures

• All nonzero digits are significant.

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

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

Significant Figures in Calculations

• Addition/Subtraction: Result should have the same number of decimal places as the measurement with the fewest decimal places.

• Multiplication/Division: Result should have the same number of significant digits as the measurement with the fewest significant digits.

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

Scientific Notation

• General form: $C \times 10^n$, where $1 \leq C < 10$ and $n$ is an integer.

• Positive $n$ for numbers greater than 1; negative $n$ for numbers less than 1.

• Only significant figures are shown in the coefficient.

Percentages

• Percent (%) means "per hundred."

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

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

Matter: The “Stuff” of Chemistry

Mass and Weight

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

• Weight: The force of gravity on an object; varies with location but is often used interchangeably with mass on Earth.

Volume

• Volume is the amount of space occupied by matter.

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

• 1 mL = 1 cm3; 5 mL ≈ 1 teaspoon.

Density and Specific Gravity

• Density (d): The ratio of mass to volume: $d = \frac{m}{V}$

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

• Specific Gravity (sp gr): The ratio of the density of a sample to the density of water: $\text{sp gr} = \frac{\text{density of sample}}{\text{density of water}}$

• Specific gravity is unitless.

Temperature Scales

• Fahrenheit (°F): Used in the U.S.

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

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

• Conversions:

  • $K = °C + 273$

  • $°C = (°F - 32) \times \frac{5}{9}$

  • $°F = (°C \times \frac{9}{5}) + 32$

Example: Normal body temperature is 98.6°F or 37.0°C.

Energy and Specific Heat

• Energy: The capacity to do work or supply heat. Measured in joules (J) or calories (cal).

• 1 calorie = 4.184 joules; 1 Calorie (nutritional) = 1000 calories.

• Specific Heat (SH): The amount of heat needed to raise the temperature of 1 g of a substance by 1°C: $\text{SH} = \frac{\text{heat}}{\text{grams} \times \Delta T}$

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

States of Matter

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

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

• Gas: No definite shape or volume; particles are far apart and 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 several measurements and average them.

Health and Dosage Calculations

• 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 (milliequivalent) is used for electrolytes.

• Body weight is often measured in pounds (U.S.) or kilograms (pharmaceuticals).

• Dosage calculation steps: (1) Determine final units, (2) Identify given information, (3) Choose conversion factors, (4) Set up and solve.

• Medication delivery may be measured in drops per milliliter (gtt/mL), where 'gtt' is from Latin gutta (drop).

• Percentages are used in active ingredient labeling, pediatric dosing, and nutrition labeling (e.g., % Daily Value).

Study Guide Summary

• Classify matter as pure substance or mixture; further as element, compound, homogeneous, or heterogeneous.

• Understand the organization of the periodic table, groups, periods, and element symbols.

• Distinguish between physical and chemical changes; balance chemical equations.

• Convert between metric units, apply significant figures, use scientific notation, and calculate percentages.

• Define and measure mass, volume, density, and specific gravity; convert temperatures; understand energy and specific heat; compare states of matter.

• Apply measurement concepts to health, including accuracy, precision, unit conversions, and dosage calculations.