General Chemistry I Final Exam Review: Key Concepts and Formulas

Ch. 1: Matter, Measurement & Problem Solving

Matter and Its Composition

Matter is anything that has mass and occupies space. It is composed of atoms and molecules. Chemistry investigates the properties of matter by examining the atoms and molecules that compose it.

• Atoms: The smallest unit of an element that retains its chemical properties.

• Molecules: Groups of atoms bonded together, representing the smallest fundamental unit of a chemical compound.

• Example: Water (H2O) is a molecule composed of two hydrogen atoms and one oxygen atom.

Scientific Method: Observations, Laws, Hypotheses, and Theories

Science begins with observations of the physical world. Multiple observations can be generalized into a law. A hypothesis is a tentative explanation of the observations. Multiple well-established hypotheses may prompt the formation of a theory, which is a model that explains the underlying reasons for laws and observations. Laws, hypotheses, and theories lead to predictions that can be tested by experiments.

• Law: A statement that summarizes past observations and predicts future ones.

• Theory: A well-substantiated explanation of some aspect of the natural world.

States and Classification of Matter

Matter exists in three primary states: solid, liquid, and gas. It can be classified as a pure substance (element or compound) or a mixture (homogeneous or heterogeneous).

• Solid: Definite shape and volume.

• Liquid: Definite volume, indefinite shape.

• Gas: Indefinite shape and volume.

• Pure Substance: Element or compound with uniform composition.

• Mixture: Combination of two or more substances; can be homogeneous (uniform) or heterogeneous (non-uniform).

Physical and Chemical Changes

Changes that alter only the state or appearance of a substance, but not its composition, are physical changes. Changes that alter the composition of matter are chemical changes.

• Physical Change: No change in chemical identity (e.g., water boiling).

• Chemical Change: Atoms rearrange, transforming substances (e.g., rusting of iron).

Significant Figures and Measurement

Significant figures reflect the precision of a measured quantity. Rules for significant figures:

• All nonzero digits are significant.

• Interior zeroes (between nonzero digits) are significant.

• Leading zeroes are not significant.

• Trailing zeroes are significant if after a decimal point.

SI Units Table

Quantity	Unit	Symbol
Length	Meter	m
Mass	Kilogram	kg
Time	Second	s
Temperature	Kelvin	K
Amount of substance	Mole	mol
Electric current	Ampere	A
Luminous intensity	Candela	cd

SI Prefixes Table

Prefix	Symbol	Multiplier
exa	E	1018
peta	P	1015
tera	T	1012
giga	G	109
mega	M	106
kilo	k	103
centi	c	10-2
milli	m	10-3
micro	μ	10-6
nano	n	10-9
pico	p	10-12
femto	f	10-15
atto	a	10-18

Ch. 2: Atoms & Elements

Atomic Theory

Each element is composed of indestructible particles called atoms. All atoms of a given element have the same mass and other properties. Atoms combine in simple, whole-number ratios to form compounds. In chemical reactions, atoms change the way they are bound together to form new substances.

Discovery of Subatomic Particles

• Electron: Discovered by J.J. Thomson using cathode rays; negatively charged.

• Millikan's Oil Drop Experiment: Measured the charge of the electron, allowing calculation of its mass.

• Rutherford's Gold Foil Experiment: Revealed the nuclear structure of the atom; atoms are mostly empty space with a dense nucleus.

Subatomic Particles and Atomic Structure

• Protons: Positive charge, mass = $1.67262 \times 10^{-27}$ kg

• Neutrons: Neutral, mass = $1.67493 \times 10^{-27}$ kg

• Electrons: Negative charge, mass = $0.00091 \times 10^{-27}$ kg

Atomic Number, Mass Number, and Isotopes

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

• Mass Number (A): Number of protons plus neutrons: $A = p + n$

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Notation: $^{A}_{Z}X$ or X-A (e.g., C-12)

Periodic Table Organization

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

Atomic Mass and Avogadro's Number

• Atomic Mass: Weighted average of the masses of an element's isotopes.

• Formula: $\text{Atomic Mass} = \sum (\text{fraction of isotope n} \times \text{mass of isotope n})$

• Avogadro's Number: $6.02214 \times 10^{23}$ particles/mol

• Conversions: grams $\leftrightarrow$ moles $\leftrightarrow$ atoms

Ch. 3: Molecules and Compounds

Chemical Formulas

Chemical formulas represent the composition of compounds and can be categorized as:

• Empirical Formula: Simplest whole-number ratio of elements.

• Molecular Formula: Actual number of atoms of each element in a molecule.

• Structural Formula: Shows how atoms are connected.

Classification of Elements and Compounds

• Elements: Atomic (e.g., Ne) or Molecular (e.g., O2).

• Compounds: Molecular (e.g., H2O) or Ionic (e.g., NaCl).

Naming Compounds

Compounds are named according to specific rules depending on whether they are ionic, molecular, or acids.

• Ionic Compounds: Metal + nonmetal; use charge balancing.

• Molecular Compounds: Nonmetals; use prefixes (mono-, di-, tri-, etc.).

• Acids: Binary acids (hydro- prefix) and oxyacids (based on polyatomic ions).

Molar Mass and Percent Composition

• Molar Mass: Mass of one mole of a compound (g/mol).

• Formula: $\text{Molar Mass} = \sum (\text{number of atoms} \times \text{atomic mass})$

• Percent Composition: $\text{mass \% of element X} = \frac{\text{mass of element X in 1 mol of compound}}{\text{molar mass of compound}} \times 100\%$

Organic Functional Groups Table

Family	Name Ending	General Formula	Example	Name	Occurrence/Use
Alcohols	-ol	R-OH	CH3CH2OH	Ethyl alcohol	Alcoholic beverages
Ethers	ether	R-O-R'	CH3OCH3	Dimethyl ether	Anesthetic, solvent
Aldehydes	-al	R-CHO	CH3CHO	Ethanal	Perfume, flavors
Ketones	-one	R-CO-R'	CH3COCH3	Propanone	Paint remover
Carboxylic acids	-oic acid	R-COOH	CH3COOH	Acetic acid	Vinegar
Esters	-ate	R-COO-R'	CH3COOCH3	Methyl acetate	Laboratory solvent
Amines	amine	R-NH2	CH3NH2	Ethyl amine	Smell of rotten fish

Ch. 4: Chemical Reactions and Chemical Quantities

Chemical Equations and Stoichiometry

Chemical reactions are represented by equations showing reactants and products. The coefficients specify the relative amounts in moles of each substance.

• Reactants: Substances consumed in the reaction.

• Products: Substances formed in the reaction.

• Balancing Equations: Ensures conservation of mass.

• Conversion Factors: Use coefficients to relate moles of reactants and products.

Limiting Reactant, Theoretical Yield, and Percent Yield

• Limiting Reactant: Reactant that limits the amount of product formed.

• Excess Reactant: Reactant not completely consumed.

• Theoretical Yield: Maximum amount of product possible.

• Percent Yield Formula: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100$

Types of Reactions

• Combustion: Reaction with O2 producing CO2 and H2O.

• Alkali Metal: Metal reacts with water or halogens.

• Halogen: Halogen reacts with metals or hydrogen.

Ch. 5: Introduction to Solutions and Aqueous Solutions

Solutions and Molarity

An aqueous solution is a homogeneous mixture of water (solvent) and another substance (solute). Concentration is expressed as molarity (M):

• Formula: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Dilution Formula: $M_1V_1 = M_2V_2$

Electrolytes

• Strong Electrolytes: Completely dissociate/ionize; conduct electricity well.

• Weak Electrolytes: Partially dissociate/ionize; conduct electricity weakly.

• Non-Electrolytes: Do not dissociate/ionize; do not conduct electricity.

Solubility Rules Table

Ions That Are Generally Soluble	Exceptions	Ions That Are Generally Insoluble	Exceptions
Li+, Na+, K+, NH4+	None	OH-, S2-	With Li+, Na+, K+, NH4+
NO3-, C2H3O2-	None	CO32-, PO43-	With Li+, Na+, K+, NH4+
Cl-, Br-, I-	With Ag+, Hg22+, Pb2+
SO42-	With Ba2+, Pb2+, Ca2+, Sr2+

Types of Equations in Solution Chemistry

• Molecular Equation: Shows all reactants and products as compounds.

• Complete Ionic Equation: Shows all strong electrolytes as ions.

• Net Ionic Equation: Shows only species that change during the reaction.

Acids, Bases, and Titrations

• Acids: Produce H+ ions in solution.

• Bases: Produce OH- ions in solution.

• Acid-Base Reaction: Produces a salt and water.

• Titration: Procedure to determine concentration using a reaction with a known solution (titrant) and an unknown (analyte).

• Equivalence Point: Point at which stoichiometric amounts of acid and base have reacted.

• Indicator: Dye that changes color depending on acidity/basicity.

Acids and Bases Table

Name of Acid	Formula	Name of Base	Formula
Hydrochloric acid	HCl	Sodium hydroxide	NaOH
Hydrobromic acid	HBr	Potassium hydroxide	KOH
Nitric acid	HNO3	Calcium hydroxide	Ca(OH)2
Sulfuric acid	H2SO4	Barium hydroxide	Ba(OH)2
Acetic acid	HC2H3O2	Ammonia*	NH3

*Ammonia is a weak base.

Redox Reactions

• Redox Reaction: Any reaction in which there is a change in oxidation states of atoms.

• Oxidizing Agent: Oxidizes another substance and is itself reduced.

• Reducing Agent: Reduces another substance and is itself oxidized.

• Mnemonic: OIL RIG = Oxidation Is Loss; Reduction Is Gain

Ch. 6: Gases

Pressure Units and STP

• Pressure Units: Pascal (Pa), atmosphere (atm), torr, psi, mmHg.

• STP: Standard Temperature and Pressure = 273.15 K (0°C), 1 atm

Pressure Units Table

Unit	Abbreviation	Average Air Pressure at Sea Level
Pascal	Pa	101,325 Pa
Pounds per square inch	psi	14.7 psi
Torr (1 mmHg)	torr	760 torr
Inches of mercury	in Hg	29.92 in Hg
Atmosphere	atm	1 atm

Gas Laws

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T and n)

• Charles's Law: $V_1/T_1 = V_2/T_2$ (at constant P and n)

• Avogadro's Law: $V_1/n_1 = V_2/n_2$ (at constant P and T)

• Ideal Gas Law: $PV = nRT$ where $R = 0.08206 \frac{L \cdot atm}{mol \cdot K}$

Applications of the Ideal Gas Law

• Calculate molar volume, density, and molar mass of gases.

• Units must be consistent: Liters, atm, moles, Kelvin.

*Additional info: These notes cover the foundational topics for a General Chemistry I final exam, including atomic theory, chemical reactions, stoichiometry, solutions, and gas laws. For further study, refer to the ACS study guide and practice problems as recommended in the course materials.*