Chapter 1: Essential Ideas – Matter and Measurement (General Chemistry)

Basic Concepts in Chemistry

What is Chemistry?

Chemistry is the field of science concerned with the characterization, composition, and transformation of matter. It studies anything that has mass and occupies space.

• Matter: Anything that has mass and occupies space.

• Mass: The amount of matter present in a sample.

• Matter includes both living and nonliving things found in nature (e.g., plants, rocks, air, bacteria) as well as man-made synthetic materials (e.g., tables, chairs, chalk).

• Forms of energy (heat, light, electricity) are not considered matter.

The Nature of Matter: Atoms

Atomic Theory and the Atom

The concept of atoms was first proposed by John Dalton in the early 19th century. An atom is the smallest unit of an element that cannot be chemically or physically divided into smaller particles by ordinary means.

• Element: A pure substance that cannot be broken down into simpler substances by chemical or physical means. Examples: silver, diamond, copper, iron.

• Compound: A pure substance composed of two or more elements chemically bonded in a fixed proportion. Compounds can be broken down into simpler substances by chemical means. Examples: CO2, H2O, rust.

Dalton's Atomic Theory (Summary)

• All matter is composed of extremely small particles called atoms.

• All atoms of a given element are identical in size, mass, and properties.

• Atoms of different elements are different.

• Compounds are composed of atoms of more than one element. In any given compound, the same types of atoms are always present in the same relative numbers (Law of Definite Proportions).

• Example: CO2 and H2O always have the same ratio of their constituent atoms.

Law of Definite Proportions

A chemical compound always contains the same proportion of its component elements by mass.

• Example: Water (H2O) always contains hydrogen and oxygen in a 2:1 ratio by number of atoms.

Law of Multiple Proportions

If two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in small whole number ratios.

• Example: CO and CO2 both contain carbon and oxygen, but in different ratios.

Law of Conservation of Mass

Atoms cannot be created or destroyed in a chemical reaction; they are only rearranged and transformed.

States and Classification of Matter

Physical States of Matter

Matter exists in three primary physical states, classified by shape and volume:

• Solid: Definite shape and definite volume.

• Liquid: Indefinite shape and definite volume.

• Gas: Indefinite shape and indefinite volume.

Water is a common substance that exists in all three states.

Classification of Matter

• Pure Substance: A single kind of matter that cannot be separated into other kinds of matter by physical means. Has a definite and constant composition. Examples: elements (gold, iron), compounds (water, salt).

• Mixture: A combination of two or more substances in which each substance retains its own identity. Mixtures can be separated by physical means.

• Homogeneous Mixture: Uniform composition and appearance throughout (e.g., saltwater, air).

• Heterogeneous Mixture: Non-uniform composition; different regions have different properties (e.g., sand in water, salad).

Distinguishing Between Classes of Matter

• Identify a material as an element, compound, homogeneous mixture, or heterogeneous mixture based on its composition and uniformity.

• Examples: Air in a scuba tank (homogeneous mixture), gases inside a party balloon (mixture), solid carbon dioxide (compound).

Properties of Matter

Physical and Chemical Properties

A property is a distinguishing characteristic of a substance that is used for identification and description.

• Intensive Properties: Independent of the amount of substance (e.g., temperature, density, color).

• Extensive Properties: Dependent on the amount of substance (e.g., mass, volume, length).

• Physical Properties: Can be measured without changing the substance into another substance (e.g., color, melting point, boiling point, physical state).

• Chemical Properties: Can be observed only by reacting the substance to form something else (e.g., reactivity, flammability, oxidation).

Density

Density is the ratio of the mass of a substance to its volume.

• Formula: $d = \frac{m}{V}$

• Common units: g/mL, g/cm3 (for solids and liquids), g/L (for gases)

Compounds: Ionic Compounds

Ionic compounds contain ions (charged particles) held together by the attraction of opposite charges. Example: sodium chloride (table salt).

Measurement in Chemistry

Measurement and Units

Measurement is the determination of the dimensions, capacity, or extent of something. Common measurements include mass, volume, length, time, temperature, and concentration.

• Two main systems of measurement: English system (inch, foot, pound, quart) and Metric (SI) system (gram, meter, liter).

SI Base Units

Quantity	Unit Name	Symbol
Mass	kilogram	kg
Length	meter	m
Temperature	kelvin	K
Time	second	s
Electric current	ampere	A
Amount of substance	mole	mol

Temperature Scales

• Celsius (°C)

• Fahrenheit (°F)

• Kelvin (K)

Conversion formulas:

• $T(\text{K}) = T(\degree\text{C}) + 273.15$

• $T(\degree\text{C}) = \frac{5}{9}(T(\degree\text{F}) - 32)$

• $T(\degree\text{F}) = \frac{9}{5}T(\degree\text{C}) + 32$

Exact and Inexact Numbers

• Exact Numbers: Have no uncertainty (e.g., counted values, defined conversions such as 1 in = 2.54 cm exactly).

• Inexact Numbers: Have a degree of uncertainty (e.g., measurements).

Precision and Accuracy

• Precision: How closely repeated measurements agree with each other.

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

Significant Figures

• All nonzero numbers are significant.

• Zeros at the beginning of a value are not significant.

• Zeros after a decimal point and after a nonzero number are significant.

• Zeros at the end of a value with no decimal point may or may not be significant; use scientific notation to clarify.

• Zeros between nonzero numbers are always significant.

Scientific Notation

Scientific notation expresses very large or very small numbers in the form $A \times 10^n$, where A is the coefficient (a number with a single nonzero digit to the left of the decimal place) and n is an integer.

• Example: $3.14 \times 10^2$

Significant Figures in Calculations

• Multiplication & Division: The result should have no more significant figures than the measurement with the fewest significant figures.

• Addition & Subtraction: The result should have no more digits to the right of the decimal point than the measurement with the fewest such digits.

Density Calculation Example

• Density ($d$) is calculated as $d = \frac{m}{V}$.

• If a nugget has a mass of 163 g and displaces water from 50.0 mL to 58.5 mL, the volume is $58.5 - 50.0 = 8.5$ mL.

• Density = $\frac{163\ \text{g}}{8.5\ \text{mL}} = 19.2\ \text{g/mL}$ (rounded to 3 significant figures).

Unit Conversions and Dimensional Analysis

Unit Conversions

• Conversion factors are ratios of equivalent quantities used to convert from one unit to another.

• Common equivalencies:

Quantity	Equivalent Units
1 kg	2.205 pounds (lb) = 35.27 ounces (oz)
1 m	39.37 inches (in) = 3.281 feet (ft)
1 km	0.6214 mile (mi)
1 in	2.54 cm (exactly)
1 L	0.2642 gallon (gal) = 1.057 quarts (qt)

Dimensional Analysis

Dimensional analysis is a method for converting quantities using conversion factors. It allows the initial units to be canceled and replaced with the desired units.

• Example: To convert inches to centimeters, use the conversion factor $1\ \text{in} = 2.54\ \text{cm}$ (exactly).

General steps:

1. Identify the starting quantity and units.

2. Multiply by appropriate conversion factors so that units cancel as needed.

3. Continue until the desired units are obtained.