Fundamental Concepts and Measurements in General Chemistry

Study Guide - Smart Notes

Tailored notes based on your materials, expanded with key definitions, examples, and context.

Introduction to Chemistry

Definition and Scope

Chemistry is the study of matter, its properties, and the changes it undergoes. Matter is anything that has mass and occupies space.

Atoms: The basic unit of matter.
Molecules: Two or more atoms bonded together.
Compounds: Molecules composed of two or more different types of atoms.
Mixtures: Physical combinations of elements and/or compounds.

Elements and Compounds

Elements

An element is a unique kind of atom; it may not be broken down into simpler substances by chemical means.
Elements are represented by symbols (one or two letters, first always capitalized).
Oxygen is an important element in many compounds.

Law of Constant Composition

Compounds have a definite composition, meaning the relative number of atoms of each element in a compound is the same in every sample.

Properties of Matter

Physical and Chemical Properties

Physical properties: Can be observed without changing the substance into another substance (e.g., density, boiling point, color).
Chemical properties: Can only be seen when a substance is changed into another substance (e.g., flammability, reactivity).

Intensive and Extensive Properties

Intensive properties: Independent of the amount of substance (e.g., density, boiling point, color).
Extensive properties: Depend on the amount of substance (e.g., mass, volume, energy).

Physical and Chemical Changes

Physical changes: Changes in matter that do not change the substance itself (e.g., changes of state, temperature, volume).
Chemical changes: Result in a new substance (e.g., combustion, oxidation, decomposition).

Separation of Mixtures

Methods of Separation

Filtration: Separates solids from liquids (e.g., sand from water).
Distillation: Uses differences in boiling points to separate a liquid homogeneous mixture into its components.
Chromatography: Separates substances based on their ability to adhere to a solid surface.

Energy and Work

Definitions

Energy: The capacity to do work or transfer heat.
Work: The energy transferred when a force is applied to an object causing movement.
Force: Any push or pull on an object.

Kinetic and Potential Energy

Kinetic energy: Energy of motion, depends on mass and velocity.
Formula:
Potential energy: Stored energy due to position or composition.
Formula:

Units of Measurement

Quantitative Measurements

Measured using SI units (International System of Units).
Examples: meter (length), kilogram (mass), second (time), kelvin (temperature), mole (amount of substance).
Luminous intensity: candela (cd).

SI Prefixes

SI prefixes are used to indicate multiples or fractions of units.

Prefix	Abbreviation	Meaning	Example
Peta	P	1015	1 petawatt (PW) = 1 × 1015 watts
Tera	T	1012	1 terawatt (TW) = 1 × 1012 watts
Giga	G	109	1 gigawatt (GW) = 1 × 109 watts
Mega	M	106	1 megawatt (MW) = 1 × 106 watts
Kilo	k	103	1 kilowatt (kW) = 1 × 103 watts
Deci	d	10-1	1 deciwatt (dW) = 1 × 10-1 watt
Centi	c	10-2	1 centiwatt (cW) = 1 × 10-2 watt
Milli	m	10-3	1 milliwatt (mW) = 1 × 10-3 watt
Micro	μ	10-6	1 microwatt (μW) = 1 × 10-6 watt

Prefix	Abbreviation	Meaning	Example
Nano	n	10-9	1 nanowatt (nW) = 1 × 10-9 watt
Pico	p	10-12	1 picowatt (pW) = 1 × 10-12 watt
Femto	f	10-15	1 femtowatt (fW) = 1 × 10-15 watt
Atto	a	10-18	1 attowatt (aW) = 1 × 10-18 watt
Zepto	z	10-21	1 zeptowatt (zW) = 1 × 10-21 watt

Temperature Scales

Kelvin (K): The SI unit for temperature. Absolute zero (0 K) is the lowest possible temperature.
Conversion formulas:
- Kelvin to Celsius:
- Fahrenheit to Celsius:
- Celsius to Fahrenheit:

Volume and Mass Units

1 liter (L) = 1 cubic decimeter (dm3)
1 milliliter (mL) = 1 cubic centimeter (cm3)
1 kilogram (kg) = 2.20 pounds (lbs)
1 meter (m) = 1.09 yards

Laboratory Measurement Tools

Common Volumetric Equipment

Graduated cylinder: Used to measure liquid volumes accurately.
Syringe: Used to deliver variable volumes.
Burette: Used for titration to deliver precise volumes.
Pipette: Used to deliver a specific volume.
Volumetric flask: Used to prepare solutions of precise volume.

Density

Definition and Formula

Density is the mass per unit volume of a substance.
Formula:

Common Densities

Substance	Density (g/cm3)
Air	0.001
Balsa wood	0.16
Ethanol	0.79
Water	1.00
Ethylene glycol	1.09
Table sugar	1.59
Table salt	2.16
Iron	7.86

Measurement, Precision, and Accuracy

Uncertainty and Significant Figures

All measurements have some degree of uncertainty.
Exact numbers: Known exactly (e.g., counting numbers, defined quantities).
Inexact numbers: Measured using scientific instruments; always have some uncertainty.
Digit of uncertainty: The last digit measured is reliable but not exact.

Precision and Accuracy

Precision: How closely repeated measurements agree with each other.
Accuracy: How close a measurement is to the true value.

Significant Figures (Sig Figs)

Rules for determining significant figures:
1. All nonzero digits are significant.
2. Zeros between nonzero digits are significant.
3. Zeros at the beginning are not significant.
4. Zeros at the end are significant if there is a decimal point.
For addition/subtraction: The result has the same number of decimal places as the measurement with the fewest decimal places.
For multiplication/division: The result has the same number of significant figures as the measurement with the fewest significant figures.

Energy Units and Conversions

Joule (J): SI unit of energy. 1 J = energy of 2 kg moving at 1 m/s.
Calorie (cal): 1 cal = 4.18 J.
1 nutritional Calorie (Cal) = 1 kcal = 1000 cal.

Examples

Example of density calculation: If a sample has a mass of 10 g and a volume of 2 cm3, its density is .
Example of significant figures: 0.00450 has three significant figures (the leading zeros are not significant).

Additional info: Some context and explanations have been expanded for clarity and completeness based on standard General Chemistry curriculum.