CHM1045 Exam Study Sheet (Chapters 1 & 2) - Vocabulary Flashcards

Vocabulary flashcards covering key terms and definitions from Chapters 1 and 2 (CHM1045).

Scientific Method

The systematic process of observing, hypothesizing, experimenting, and modifying a hypothesis based on results.

Theory

A well-supported explanation of why phenomena occur in nature.

Law

A concise statement describing what happens in nature, often expressed as a general rule.

Macroscopic

Relating to the large-scale properties of matter observable without a microscope.

Microscopic

Relating to the tiny, atomic-level details not visible to the naked eye.

Symbolic Chemistry

Representation of substances and reactions using symbols, formulas, and equations.

States of Matter

The physical forms of matter: solid, liquid, gas, and plasma.

Solid

A state with definite shape and volume; particles vibrate in place.

Liquid

A state with definite volume that takes the shape of its container; particles slide past one another.

Gas

A state with indefinite shape and volume; particles move freely and are highly compressible.

Plasma

An ionized, high-energy state of matter with free electrons.

Pure Substance

Matter with uniform composition; either an element or a compound.

Element

A pure substance that cannot be chemically broken down into simpler substances.

Compound

A pure substance formed from two or more elements chemically bonded in fixed proportions.

Mixture

Matter composed of two or more substances that can be separated by physical means.

Homogeneous

Mixture with uniform composition throughout (one phase).

Heterogeneous

Mixture with visibly different parts or phases (more than one substance).

Physical Property

A characteristic observed or measured without changing the substance’s composition.

Chemical Property

A characteristic observed during a chemical change (reactivity or composition change).

Intensive Property

Property that does not depend on the amount of substance (e.g., density, temperature).

Extensive Property

Property that scales with the amount of substance (e.g., mass, volume).

Physical Change

Change that does not produce a new substance; composition remains the same.

Chemical Change

Change that produces one or more new substances with different properties.

SI Base Units

Fundamental units for base quantities (e.g., meter, kilogram, second, Kelvin).

Kelvin (K)

SI unit of absolute temperature; 0 K is absolute zero.

Volume (V)

Amount of space occupied; 1 m³ = 1000 L; 1 L = 1000 mL = 1000 cm³.

Density (ρ)

Mass per unit volume; ρ = m / V.

Significant Figures

Digits that carry meaning about precision in a measurement.

Leading Zeros

Zeros to the left of the first nonzero digit; not significant.

Sandwiched Zeros

Zeros between nonzero digits; significant.

Trailing Zeros with Decimal

Trailing zeros after a decimal point are significant.

Trailing Zeros without Decimal

Trailing zeros without a decimal point are not significant.

Multiply/Divide Rule (sig figs)

Result has as many sig figs as the operand with the fewest sig figs.

Add/Subtract Rule (decimals)

Result has the least number of decimal places among operands.

Fahrenheit to Celsius

°C = (°F − 32) × 5/9.

Celsius to Fahrenheit

°F = (°C × 9/5) + 32.

Celsius to Kelvin

K = °C + 273.15.

Dimensional Analysis

Method to convert units using conversion factors to reach a desired unit.

Inch to Centimeter

1 inch = 2.54 cm.

Kilometer to Mile

1 km = 0.62137 miles.

Ounce to Gram

1 oz = 28.349 g.

Dalton’s Atomic Theory

Five postulates about atoms: all matter is made of atoms; atoms combine in fixed whole-number ratios; atoms are rearranged in reactions; atoms of same element have same properties.

Law of Conservation of Mass

Mass is conserved in chemical reactions; matter is neither created nor destroyed.

Law of Definite Proportions

A compound contains elements in fixed, definite mass ratios.

Law of Multiple Proportions

When two elements form more than one compound, the mass ratios are small whole numbers.

Proton

Positively charged subatomic particle; +1; mass ~1 amu.

Neutron

Electrically neutral subatomic particle; 0 charge; mass ~1 amu.

Electron

Negatively charged subatomic particle; −1; negligible mass.

Atomic Number (Z)

Number of protons in the nucleus.

Mass Number (A)

Total number of protons and neutrons in the nucleus.

Isotopes

Atoms with the same Z but different A.

Ions

Charged atoms; cations are positive, anions are negative.

Avg Atomic Mass

Weighted average of isotopic masses based on natural abundances.

Isotopic Mass

Mass of a particular isotope.

Molecular Formula

Actual number and type of atoms in a molecule.

Empirical Formula

Simplest whole-number ratio of atoms in a compound.

Structural Formula

Shows connectivity and bonding between atoms in a molecule.

Avogadro’s Number

6.022 × 10^23 particles per mole.

Molar Mass

Mass per mole of a substance (g/mol).

The Mole

Amount of substance containing 6.022 × 10^23 particles.

Grams ↔ Moles ↔ Particles

Conversions between mass, amount (moles), and number of particles using molar mass and Avogadro’s number.