General Chemistry I Final Exam Review Notes

Final Exam Review Session

Exam Structure and Preparation

• The final exam is cumulative and covers Chapters 1-11 of General Chemistry I.

• Format: 40 questions in 55 minutes; completion is required to pass the course.

• Materials: Pink Scantron, writing utensil, non-programmable calculator.

• Curve Equation: $\text{Score} = \sqrt{\text{raw score}/40} \times 100$

• Periodic table provided will have only symbols, not names.

• Recommended resources: ACS study guide, back-of-chapter problems.

Ch. 1: Matter, Measurement & Problem Solving

1.1 Matter and Its Composition

• All matter is composed of atoms and molecules.

• Chemistry investigates the properties of matter by examining its atomic and molecular composition.

• Example: Water molecule (H2O) consists of two hydrogen atoms and one oxygen atom bonded together.

1.2 Scientific Method

• Science begins with observations of the physical world.

• Generalizations from observations form a law.

• A hypothesis is a tentative explanation of observations.

• Multiple well-established hypotheses may lead to a theory, which explains the underlying reasons for laws and observations.

• Laws, hypotheses, and theories lead to predictions that can be tested by experiments.

1.3 States and Classification of Matter

• States of matter: Solid, liquid, gas.

• Classification: Pure substances (elements, compounds) and mixtures (homogeneous, heterogeneous).

• Example: Helium (element), pure water (compound), wet sand (heterogeneous mixture), tea with sugar (homogeneous mixture).

1.4 Physical and Chemical Changes

• Physical changes alter only the state or appearance, not composition (e.g., water boiling).

• Chemical changes alter the composition of matter; atoms rearrange to form new substances (e.g., rusting of iron).

1.6 Significant Figures and SI Units

• Significant Figure Rules:

  1. All nonzero digits are significant.

  2. Interior zeroes (between nonzero digits) are significant.

  3. Leading zeroes are not significant.

  4. Trailing zeroes are significant if after a decimal point.

• SI Base Units:

  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:

  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

2.2-2.3 Atomic Theory

• Each element is composed of indestructible particles called atoms.

• Atoms of a given element have the same mass and properties.

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

• Atoms of one element cannot change into atoms of another element.

• In chemical reactions, atoms change the way they are bound together.

2.4 Discovery of Subatomic Particles

• J.J. Thomson discovered the electron via cathode ray experiments; electrons are negatively charged.

• Millikan measured the charge of the electron, enabling calculation of its mass.

2.5 Rutherford's Nuclear Model

• Rutherford's gold foil experiment showed that atoms are mostly empty space with a dense nucleus.

• Protons and neutrons reside in the nucleus; electrons occupy the surrounding cloud.

2.6 Subatomic Particles and Isotopes

• Atoms are composed of protons, neutrons, and electrons.

• Protons and neutrons have nearly identical masses:

  • Proton: $1.67262 \times 10^{-27}$ kg

  • Neutron: $1.67493 \times 10^{-27}$ kg

  • Electron: $0.00091 \times 10^{-27}$ kg

• Mass number (A) = number of protons + number of neutrons.

• Atomic number (Z) = number of protons.

• Isotope notation: $^{A}_{Z}X$ or X-A (chemical symbol or name followed by mass number).

2.7 Periodic Table Divisions

• Elements are classified as metals, nonmetals, and metalloids.

• Periodic table groups elements by similar properties.

2.8-2.9 Atomic Mass and Avogadro's Number

• Atomic mass (atomic weight) is the average mass of an element's isotopes, weighted by natural abundance.

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

• Avogadro's number: $6.02214 \times 10^{23}$

• 1 mole of any substance contains $6.02214 \times 10^{23}$ units (atoms, molecules, etc.).

• Conversion relationships:

  • grams → moles → atoms (using molar mass and Avogadro's number)

Ch. 3: Molecules and Compounds

3.3-3.4 Chemical Formulas and Classification

• Types of chemical formulas:

  • Empirical formula: simplest whole-number ratio of elements.

  • Molecular formula: actual number of atoms of each element.

  • Structural formula: shows arrangement of atoms.

• Elements: atomic (e.g., Na) or molecular (e.g., O2).

• Compounds: molecular (e.g., H2O) or ionic (e.g., NaCl).

3.5-3.7 Naming Compounds and Molar Mass

• Ionic compounds: metal + nonmetal; molecular compounds: nonmetals only.

• Acids: binary acids (H + nonmetal), oxyacids (H + polyatomic ion).

• Molar mass = mass (g) of 1 mole of molecules or formula units. Example: CO2 molar mass = 12.01 + 2(16.00) = 44.01 g/mol

3.9-3.10 Mass Percent and Empirical/Molecular Formulas

• Chemical formula and molar masses indicate relative quantities of elements.

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

• If mass percent composition and molar mass are known, empirical and molecular formulas can be determined.

• Functional groups (alcohols, ethers, aldehydes, etc.) are important in organic chemistry.

Ch. 4: Chemical Reactions and Chemical Quantities

4.2-4.5 Chemical Equations and Stoichiometry

• Chemical reactions: Reactants → Products

• Equations provide formulas, states, and relative numbers of reactants/products.

• Coefficients specify relative amounts in moles; used as conversion factors.

• Limiting reactant: restricts amount of product formed.

• Excess reactant: not completely consumed.

• Theoretical yield: maximum product from limiting reactant.

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

• Be able to balance combustion, alkali metal, and halogen reactions.

Ch. 5: Introduction to Solutions and Aqueous Solutions

5.2-5.3 Solutions and Molarity

• Aqueous solution: homogeneous mixture of water (solvent) and another substance (solute).

• Molarity (M): $M = \frac{\text{amount of solute (mol)}}{\text{volume of solution (L)}}$

• Stock solutions can be diluted: $M_1V_1 = M_2V_2$

• Stoichiometry in aqueous reactions uses volume, concentration, and coefficients to calculate moles.

5.4-5.5 Electrolytes and Solubility Rules

• Strong electrolytes: completely dissociate/ionize; conduct electricity well.

• Weak electrolytes: partially dissociate/ionize; weakly conduct electricity.

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

• Solubility rules:

  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 Sr2+, Ba2+, Pb2+, Ca2+

• Equations:

  • Molecular equation: shows all reactants/products as compounds.

  • Complete ionic equation: shows all strong electrolytes as ions.

  • Net ionic equation: shows only species that change during reaction.

5.4-5.5 Acids and Bases

• Acids produce H+ ions; Bases produce OH- ions in solution.

• Acid-base reactions produce a salt and water.

• Strong acids/bases dissociate completely; weak acids/bases do not.

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

*Ammonia is a weak base.

5.7-5.8 Titration and Gas-Evolving Reactions

• Titration: procedure to determine concentration of an unknown solution using a solution of known concentration (titrant).

• Equivalence point: point at which stoichiometric amounts of reactants have reacted.

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

• Gas-evolving reactions produce gases such as H2S, CO2, SO2, NH3 depending on reactants.

Reactant Type	Intermediate Product	Gas Evolved	Example
Sulfides	None	H2S	2 HCl(aq) + K2S(aq) → H2S(g) + 2 KCl(aq)
Carbonates/Bicarbonates	H2CO3	CO2	2 HCl(aq) + K2CO3(aq) → CO2(g) + H2O(l) + 2 KCl(aq)
Sulfites/Bisulfites	H2SO3	SO2	2 HCl(aq) + K2SO3(aq) → SO2(g) + H2O(l) + 2 KCl(aq)
Ammonium	NH4OH	NH3	NH4Cl(aq) + KOH(aq) → NH3(g) + H2O(l) + KCl(aq)

5.9 Redox Reactions

• Redox reaction: any reaction involving 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.

• OIL RIG: Oxidation Is Loss; Reduction Is Gain (of electrons).

Ch. 6: Gases

6.1-6.2 Pressure Units and STP

• Common pressure units:

  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

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

6.3 Gas Laws

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

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

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

• Combined Gas Law: $P_1V_1/n_1T_1 = P_2V_2/n_2T_2$

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

6.4-6.5 Applications of Ideal Gas Law

• Units must be in liters, atmospheres, moles, and Kelvin.

• Can determine molar volume, density, and molar mass of an ideal gas.

• Example: 1 mol of any ideal gas at STP occupies 22.4 L.