Chapter 2: Measurement and Calculations in Chemistry

Scientific Notation

Scientific notation is a method used to express very large or very small numbers in a compact form, which is essential in chemistry for handling measurements.

• Writing Scientific Notation: A number is written as a × 10n, where 1 ≤ a < 10 and n is an integer.

• Converting Decimal Numbers: Move the decimal point so that only one nonzero digit remains to the left. The number of places moved determines the exponent.

• Example: 0.00056 = ; 123,000 =

Significant Figures (Sig. Figs.)

Significant figures reflect the precision of a measured or calculated quantity. Answers to calculations must be reported with the correct number of significant digits.

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant only if there is a decimal point.

• Example: 0.0405 has three significant figures; 1200 has two (unless specified otherwise).

Unit Conversion

Unit conversion is essential for expressing measurements in different units, especially between the English and metric systems.

• Conversion Factors: Ratios used to express the same quantity in different units (e.g., 1 inch = 2.54 cm).

• Metric Prefixes: Common prefixes include kilo (k = 1000), centi (c = 0.01), milli (m = 0.001).

• Example: To convert 5 km to meters:

Density

Density is a physical property that relates the mass of a substance to its volume.

• Definition: Density () is mass per unit volume.

• Formula: , where is mass and is volume.

• Example: If a block has a mass of 10 g and a volume of 2 cm3, its density is

Temperature Conversions

Temperature can be measured in Celsius, Kelvin, or Fahrenheit. Converting between these units is common in chemistry.

• Celsius to Kelvin:

• Celsius to Fahrenheit:

• Example: 25°C = K

Accuracy vs. Precision

Accuracy and precision are terms used to describe the quality of measurements.

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

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

• Example: If measurements are 2.01, 2.00, and 2.02 g (true value 2.00 g), they are both accurate and precise.

Chapter 3: Matter and Its Properties

Definition of Matter

Matter is anything that has mass and occupies space.

Phases of Matter

Matter exists in three primary phases: solid, liquid, and gas, each with distinct properties.

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

• Liquid: Definite volume but takes the shape of its container; particles are less tightly packed than in solids.

• Gas: No definite shape or volume; particles are far apart and move freely.

Phase Diagram

A phase diagram shows the state of a substance at various temperatures and pressures, indicating the conditions under which phases change (e.g., melting, boiling).

Classification of Matter

Matter can be classified based on its composition.

• Element: A pure substance made of only one kind of atom (e.g., O2).

• Compound: A substance composed of two or more elements chemically combined (e.g., H2O).

• Mixture: A physical blend of two or more substances (e.g., air, saltwater).

Intensive vs. Extensive Properties

Properties of matter can be classified as intensive or extensive.

• Intensive Properties: Do not depend on the amount of matter (e.g., density, boiling point).

• Extensive Properties: Depend on the amount of matter (e.g., mass, volume).

Chemical vs. Physical Properties

Physical properties can be observed without changing the substance's identity, while chemical properties describe a substance's ability to undergo chemical changes.

• Physical Properties: Color, melting point, density.

• Chemical Properties: Flammability, reactivity with acid.

Specific Heat

Specific heat is the amount of heat required to raise the temperature of one gram of a substance by one degree Celsius.

• Formula:

• Where = heat (J), = mass (g), = specific heat (J/g°C), = change in temperature (°C).

• Example: How much heat is needed to raise 10 g of water by 5°C? ( J/g°C): J

Chapter 4: Atomic Theory and Structure

Dalton’s Atomic Theory

Dalton’s Atomic Theory (early 1800s) laid the foundation for modern chemistry by describing the nature of atoms.

• All matter is composed of indivisible atoms.

• Atoms of the same element are identical; atoms of different elements are different.

• Atoms combine in simple whole-number ratios to form compounds.

• Chemical reactions involve rearrangement of atoms; atoms are not created or destroyed.

Atomic Number, Protons, Neutrons, and Electrons

Atoms are composed of protons, neutrons, and electrons. The atomic number and mass number describe their composition.

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Total number of protons and neutrons.

• Number of Neutrons:

• Number of Electrons: In a neutral atom, equals the number of protons.

• Example: For : Z = 11, A = 23, neutrons = 12, electrons = 11

Periodic Table

The periodic table organizes elements by increasing atomic number and similar chemical properties.

• Groups (columns) contain elements with similar properties.

• Periods (rows) show trends in properties across the table.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons (and thus different mass numbers).

• Example: and are isotopes of carbon.

Diatomic Elements

Certain elements naturally exist as molecules composed of two atoms.

• Diatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2

Formation of Ions

Ions are formed when atoms gain or lose electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Example: Na loses one electron to become Na+; Cl gains one electron to become Cl−.