Chemistry 101 Exam 1 Review: Matter, Measurement, Atomic Structure, and Compounds

Matter and Measurements

Significant Figures

Significant figures are the digits in a measurement that are known with certainty plus one digit that is estimated. They reflect the precision of a measurement.

• 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: 0.00450 has three significant figures.

Scientific Notation

Scientific notation expresses numbers as a product of a coefficient and a power of ten, making it easier to handle very large or small values.

• Format: $a \times 10^n$ where $1 \leq a < 10$ and $n$ is an integer.

• Example: 5,600 = $5.6 \times 10^3$

Units of Measurement

Measurements in chemistry use the International System of Units (SI), including meters (m), grams (g), liters (L), and seconds (s).

• Common units: Length (meter), Mass (gram), Volume (liter), Time (second), Temperature (Kelvin or Celsius).

Conversions Using Dimensional Analysis

Dimensional analysis is a method for converting between units using conversion factors.

• Process: Multiply by conversion factors so units cancel appropriately.

• Example: To convert 25 inches to centimeters: $25 \text{ in} \times \frac{2.54 \text{ cm}}{1 \text{ in}} = 63.5 \text{ cm}$

Density and Specific Gravity

Density is the mass per unit volume of a substance. Specific gravity is the ratio of the density of a substance to the density of water.

• Density formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Specific gravity formula: $\text{Specific gravity} = \frac{\text{Density of substance}}{\text{Density of water}}$

• Example: If a liquid has a density of 1.2 g/mL, its specific gravity is $\frac{1.2}{1.0} = 1.2$

Classification & Properties of Solids, Liquids & Gases

States of Matter

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

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

• Liquids: Definite volume, indefinite shape; particles are less tightly packed.

• Gases: Indefinite shape and volume; particles are far apart and move freely.

Elements, Compounds, and Mixtures

Substances can be classified as elements, compounds, or mixtures.

• Element: Pure substance made of one type of atom (e.g., O2).

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

• Mixture: Physical combination of two or more substances.

Homogeneous vs. Heterogeneous Mixtures

Mixtures can be classified based on uniformity.

• Homogeneous: Uniform composition throughout (e.g., salt water).

• Heterogeneous: Non-uniform composition (e.g., sand and water).

Physical vs. Chemical Changes

Changes in matter can be physical or chemical.

• Physical change: Alters form but not composition (e.g., melting ice).

• Chemical change: Produces new substances (e.g., burning wood).

Energy Units & Conversions

Energy is measured in joules (J) or calories (cal). Conversion between units is often necessary.

• Conversion: $1 \text{ cal} = 4.184 \text{ J}$

Specific Heat Calculations

Specific heat is the amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• Formula: $q = m \times c \times \Delta T$

• Where $q$ = heat (J), $m$ = mass (g), $c$ = specific heat (J/g°C), $\Delta T$ = change in temperature (°C).

Changes of State, Heat of Fusion, Heating & Cooling Curves

Substances change state (solid, liquid, gas) with energy input or removal. Heat of fusion is the energy required to melt a substance.

• Heat of fusion formula: $q = m \times \text{H}_\text{fus}$

• Heating and cooling curves show temperature changes as heat is added or removed.

Atoms and the Periodic Table

Structure of the Atom & Subatomic Particles

Atoms consist of protons, neutrons, and electrons.

• Proton: Positive charge, found in nucleus.

• Neutron: Neutral, found in nucleus.

• Electron: Negative charge, found outside nucleus.

Dalton’s Atomic Theory

Dalton proposed that matter is composed of indivisible atoms, each element has unique atoms, and compounds are formed by combinations of atoms.

• Key points: Atoms cannot be created or destroyed; atoms of same element are identical.

Rutherford’s Gold Foil Experiment

Rutherford discovered the nucleus by observing alpha particles deflected by gold foil.

• Conclusion: Atoms have a small, dense, positively charged nucleus.

Atomic Number & Mass Number

Atomic number is the number of protons; mass number is the sum of protons and neutrons.

• Formula: $\text{Mass number} = \text{Protons} + \text{Neutrons}$

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Periodic Table Organization & Names of Groups

The periodic table arranges elements by increasing atomic number and groups elements with similar properties.

• Groups: Columns (e.g., Alkali metals, Halogens, Noble gases).

• Periods: Rows.

Average Atomic Mass

Average atomic mass is the weighted average of all isotopes of an element.

• Formula: $\text{Average atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

Electron Arrangement & Valence Electrons

Electrons are arranged in energy levels; valence electrons are those in the outermost shell and determine chemical reactivity.

• Example: Sodium has one valence electron.

Metals, Nonmetals, and Metalloids

Elements are classified based on their properties.

• Metals: Conduct electricity, malleable, shiny.

• Nonmetals: Poor conductors, brittle.

• Metalloids: Properties intermediate between metals and nonmetals.

Ionic and Molecular Compounds

Naming & Writing Ionic Compound Formulas

Ionic compounds consist of positive and negative ions. Formulas are written to balance charges.

• Example: Sodium chloride: Na+ and Cl- combine to form NaCl.

Polyatomic Ions

Polyatomic ions are groups of atoms with a charge, acting as a single unit.

• Example: Nitrate: NO3-, Sulfate: SO42-

Naming & Writing Covalent Compound Formulas

Covalent compounds are formed by sharing electrons between nonmetals. Names use prefixes to indicate the number of atoms.

• Example: CO2 is carbon dioxide; N2O is dinitrogen monoxide.

Polar vs. Nonpolar Covalent Compounds

Polarity depends on the difference in electronegativity between atoms.

• Polar: Unequal sharing of electrons (e.g., H2O).

• Nonpolar: Equal sharing (e.g., O2).

Shapes of Molecules

The shape of molecules is determined by the arrangement of atoms and electron pairs (VSEPR theory).

• Examples: Linear (CO2), Bent (H2O), Trigonal planar (BF3), Tetrahedral (CH4).

Type	Example	Key Properties
Element	O2	Pure substance, one type of atom
Compound	H2O	Pure substance, two or more elements
Homogeneous Mixture	Salt water	Uniform composition
Heterogeneous Mixture	Sand and water	Non-uniform composition

State	Shape	Volume	Particle Arrangement
Solid	Definite	Definite	Closely packed
Liquid	Indefinite	Definite	Less tightly packed
Gas	Indefinite	Indefinite	Far apart

Compound Type	Bonding	Example
Ionic	Transfer of electrons	NaCl
Covalent (Polar)	Unequal sharing	H2O
Covalent (Nonpolar)	Equal sharing	O2

Additional info: Academic context and examples were added to expand brief points and make the notes self-contained for exam preparation.