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

Classifying Matter: Pure Substance or Mixture

Types of Matter

Matter is anything that occupies space and has mass. In chemistry, matter is classified based on its composition and properties.

• Pure substances are materials made up of only one type of substance and can be represented by a single chemical formula or symbol.

• Mixtures consist of two or more substances physically combined and can be separated by physical means.

Classification of Pure Substances

• Elements: The simplest type of matter, composed of only one kind of atom. Each element is represented by a unique chemical symbol.

• Compounds: Pure substances made of two or more elements chemically bonded in fixed ratios. Compounds have unique properties distinct from their constituent elements.

Classification of Mixtures

• Homogeneous mixtures: Also called solutions, these have uniform composition throughout (e.g., salt water).

• Heterogeneous mixtures: Composition varies throughout the sample (e.g., salad, granite).

Example:

• Element: Copper (Cu)

• Compound: Water (H2O)

• Homogeneous mixture: Brass (copper and zinc atoms)

• Heterogeneous mixture: Water with sand

Type	Definition	Example
Element	One type of atom	Oxygen (O2)
Compound	Two or more elements chemically bonded	Water (H2O)
Homogeneous Mixture	Uniform composition	Salt water
Heterogeneous Mixture	Non-uniform composition	Sand and water

Elements, Compounds, and the Periodic Table

The Periodic Table

The periodic table of the elements organizes all known elements based on 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. Main-group elements are labeled with 'A', transition elements with 'B', or by numbers 1–18 (IUPAC system).

• Periods: Horizontal rows numbered 1–7. Elements in the same period have the same number of electron shells.

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

Essential Elements for Life

• Macronutrients: Needed in amounts >100 mg/day (e.g., Na, Mg, K, Ca, Cl).

• Micronutrients: Needed in amounts <100 mg/day (e.g., I, F, Fe, Zn).

Chemical Formulas

• Show the types and numbers of atoms in a compound (e.g., H2O: 2 hydrogen, 1 oxygen).

How Matter Changes

Physical and Chemical Changes

Matter can undergo changes that affect its form or composition.

• Physical change: Alters the state or appearance but not the chemical identity (e.g., melting ice).

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

Chemical Equations

• Represent chemical reactions using formulas and symbols.

• Reactants are transformed into products.

• Physical states are indicated: (s)olid, (l)iquid, (g)as, (aq)ueous.

Balancing Chemical Equations

• Law of conservation of mass: Matter is neither created nor destroyed.

• Equations must have equal numbers of each atom on both sides.

• Add coefficients to balance equations.

Steps to Balance:

1. Examine the equation for balance.

2. Balance one element at a time using coefficients.

3. Check for the smallest set of coefficients.

Math Counts

Measurement Systems

Chemistry uses the metric system (SI units) for measurements.

• Mass: kilogram (kg)

• Volume: liter (L)

• Length: meter (m)

• Prefixes change unit size by powers of 10 (e.g., milli-, centi-, kilo-).

Unit Conversions

• Use conversion factors to change units (e.g., 1 dL = 0.1 L).

• Dimensional analysis: Set up problems so units cancel, leaving the desired unit.

Significant Figures

• Digits known with certainty plus one estimated digit.

• All nonzero digits are significant; zeros may or may not be, depending on position.

• Exact numbers (counted or defined) have infinite significant figures.

Rule	Example
Nonzero digits are significant	123 (3 sig figs)
Zeros between nonzero digits are significant	101 (3 sig figs)
Leading zeros are not significant	0.0025 (2 sig figs)
Trailing zeros with decimal are significant	20.0 (3 sig figs)

Calculations and Rounding

• Addition/Subtraction: Match least decimal places.

• Multiplication/Division: Match least number of significant digits.

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

Scientific Notation

• General form: $C \times 10^n$

• C is the coefficient (1 ≤ C < 10), n is the exponent.

• Positive n: number > 1; negative n: number < 1.

Percent Calculations

• Percent (%) = (part/whole) × 100

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

Matter: The “Stuff” of Chemistry

Mass and Volume

Mass is the amount of material in an object, measured in grams (g). Volume is the space occupied, measured in liters (L), milliliters (mL), or cubic centimeters (cm3).

• 1 mL = 1 cm3

• 5 mL ≈ 1 teaspoon

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 (sp gr): Ratio of sample density to water density. $\text{sp gr} = \frac{\text{density of sample}}{\text{density of water}}$

• Specific gravity is unitless.

Temperature Scales

• Fahrenheit (°F): Used in the US.

• Celsius (°C): Used worldwide.

• Kelvin (K): SI unit for temperature.

• Conversion formulas:

  • $K = ^\circ C + 273$

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

  • $^\circ F = ^\circ C \times \frac{9}{5} + 32$

Body Temperature

• Normal: 98.6°F (37.0°C)

• Hyperthermia: >40.0°C (104°F)

• Hypothermia: <35.0°C (95°F)

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 not created or destroyed.

• SI unit: joule (J); 1 calorie (cal) = 4.184 J; 1 Calorie (Cal) = 1000 cal.

Specific Heat

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

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

• Water has a high specific heat (1.00 cal/g·°C).

Substance	Specific Heat (cal/g·°C)
Water (liquid)	1.00
Iron	0.11
Copper	0.09
Glass	0.20
Aluminum	0.22
Lead	0.03
Gold	0.03
Silver	0.06
Mercury	0.03

States of Matter

• Solid: Definite shape and volume; particles tightly 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	Orderly, close	Less orderly	Random, 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.

Units and Conversions in Health

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

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

• Body weight: pounds (US), kilograms (pharmaceuticals).

Dosage Calculations

1. Determine units for the final answer.

2. Identify given information.

3. Choose conversion factors to cancel units.

4. Set up the equation so only the desired unit remains.

Drop Units and Percent in Health

• gtt/mL: Drops per milliliter; used for IV drip rates.

• Drop factor depends on IV tubing diameter.

• Percent active ingredient: Used in medication formulation.

• Percent of adult dose: Used for pediatric dosing.

• Percent Daily Value (%DV): Used in nutrition labeling.