CHEM 1050 Final Exam Study Guide: Key Topics and Concepts

Chapter 2: Chemistry and Measurements

Measured Numbers, Exact Numbers, and Calculations

Understanding the distinction between measured and exact numbers is fundamental in chemistry, as it affects the precision and accuracy of calculations.

• Measured Numbers: Values obtained by measurement, subject to uncertainty (e.g., length, mass).

• Exact Numbers: Values known with complete certainty, often from counting or defined quantities (e.g., 1 dozen = 12).

• Calculations: Significant figures must be considered when performing mathematical operations with measured numbers.

• Example: Multiplying 2.5 cm (measured) by 3 (exact) gives a result with two significant figures.

Prefixes and Equalities

Metric prefixes are used to express measurements in different scales.

• Common Prefixes: kilo- (k, $10^3$), centi- (c, $10^{-2}$), milli- (m, $10^{-3}$), micro- (μ, $10^{-6}$).

• Equalities: Statements showing the relationship between two units (e.g., 1 m = 100 cm).

Writing Conversion Factors and Unit Conversion

Conversion factors are ratios derived from equalities and are used to convert between units.

• Conversion Factor: $\frac{100\ \text{cm}}{1\ \text{m}}$ or $\frac{1\ \text{m}}{100\ \text{cm}}$

• Unit Conversion: Multiply the given value by the appropriate conversion factor to obtain the desired unit.

• Example: Convert 2.0 m to cm: $2.0\ \text{m} \times \frac{100\ \text{cm}}{1\ \text{m}} = 200\ \text{cm}$

Concept of Bone Density

Bone density is a clinical measurement used to assess bone health.

• Definition: Bone density is the mass of bone per unit volume, typically expressed in g/cm3.

• Application: Used to diagnose osteoporosis and assess fracture risk.

Chapter 4: Atoms and Elements

Periodic Table; Common Elements and Symbols

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

• Element Symbols: One- or two-letter abbreviations (e.g., H for hydrogen, Na for sodium).

• Groups and Periods: Vertical columns are groups; horizontal rows are periods.

Atomic Number and Mass Number, Calculations

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

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

• Calculation: $A = Z + \text{number of neutrons}$

• Example: Carbon-14: Z = 6, A = 14, so neutrons = 8.

Electron Energy Levels and Electronic Configuration

• Energy Levels: Electrons occupy specific energy levels (shells) around the nucleus.

• Electronic Configuration: Distribution of electrons among energy levels and sublevels (e.g., 1s2 2s2 2p6).

Trends/Properties of the Periodic Table

• Atomic Radius: Increases down a group, decreases across a period.

• Ionization Energy: Decreases down a group, increases across a period.

• Electronegativity: Tendency to attract electrons; increases across a period, decreases down a group.

Chapter 6: Ionic and Molecular Compounds

Ions, Naming and Writing Formula of Ionic Compounds

• Ions: Atoms or groups of atoms with a net charge (cations are positive, anions are negative).

• Naming: Name the cation first, then the anion (e.g., NaCl is sodium chloride).

• Writing Formulas: Balance charges to write correct formulas (e.g., Mg2+ and Cl- form MgCl2).

Polyatomic Ions

• Definition: Ions composed of more than one atom (e.g., SO42-, NO3-).

• Common Polyatomic Ions: Hydroxide (OH-), ammonium (NH4+), carbonate (CO32-).

Molecular Compounds

• Definition: Compounds formed by sharing electrons between nonmetals.

• Naming: Use prefixes to indicate the number of atoms (e.g., CO2 is carbon dioxide).

Lewis Structures

• Lewis Structure: Diagram showing valence electrons as dots around atoms.

• Purpose: Predicts bonding and lone pairs in molecules.

Electronegativity

• Definition: The ability of an atom to attract shared electrons in a bond.

• Trend: Increases across a period, decreases down a group.

Shapes and Polarity of Simple Molecules

• Shapes: Determined by VSEPR theory (e.g., linear, bent, tetrahedral).

• Polarity: Molecules are polar if they have polar bonds and an asymmetric shape.

• Example Table:

Molecule	Shape	Polarity
CO2	Linear	Nonpolar
H2O	Bent	Polar
CH4	Tetrahedral	Nonpolar
NH3	Trigonal pyramidal	Polar

Chapter 7: Chemical Reactions and Quantities

Equations for Chemical Reactions

• Chemical Equation: Represents reactants and products using chemical formulas.

• Balancing: Law of conservation of mass requires equal numbers of each atom on both sides.

Types of Chemical Reactions

• Combination: $A + B \rightarrow AB$

• Decomposition: $AB \rightarrow A + B$

• Single Replacement: $A + BC \rightarrow AC + B$

• Double Replacement: $AB + CD \rightarrow AD + CB$

• Combustion: Hydrocarbon + O2 $\rightarrow$ CO2 + H2O

Mole, Molar Mass, Calculations

• Mole: Amount of substance containing $6.022 \times 10^{23}$ particles (Avogadro's number).

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

• Calculation: $\text{Moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

Concept of Limiting Reactant, Calculations

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Calculation: Compare mole ratios of reactants to determine which is limiting.

Chapter 8: Gases

Gas Laws and Calculations

Gas laws describe the relationships between pressure, volume, temperature, and amount of gas.

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

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P and n)

• Gay-Lussac's Law: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$ (at constant V and n)

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Ideal Gas Law: $PV = nRT$

• Example: Calculate the volume of 1.0 mol of gas at STP (0°C, 1 atm): $V = \frac{nRT}{P}$

Chapter 9: Solutions

Types of Solutions

• Definition: Homogeneous mixtures of two or more substances.

• Types: Solid, liquid, or gas solutions (e.g., saltwater, air).

Saturated and Unsaturated Solutions

• Saturated: Contains the maximum amount of solute that can dissolve at a given temperature.

• Unsaturated: Contains less solute than the maximum amount.

Dilutions and Calculations

• Dilution Equation: $C_1V_1 = C_2V_2$

• Application: Used to prepare solutions of lower concentration from a stock solution.

Solution Concentration and Calculations

• Concentration Units: Molarity (M), percent by mass or volume.

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

Chapter 10: Reaction Rates and Chemical Equilibrium

Rates of Chemical Reactions

• Definition: The speed at which reactants are converted to products.

• Factors Affecting Rate: Concentration, temperature, catalysts, surface area.

Chemical Equilibrium

• Definition: State where the rates of forward and reverse reactions are equal.

• Dynamic Equilibrium: Both reactions continue, but concentrations remain constant.

Equilibrium Constants

• Expression: $K = \frac{[\text{products}]}{[\text{reactants}]}$ (for a general reaction)

• Interpretation: Large K favors products; small K favors reactants.

Le Chatelier’s Principle

• Principle: If a system at equilibrium is disturbed, it will shift to counteract the disturbance.

• Disturbances: Changes in concentration, temperature, or pressure.

Chapter 11: Acids and Bases

Definition of Acid and Base

• Arrhenius Definition: Acids produce H+ in water; bases produce OH-.

• Bronsted-Lowry: Acids donate protons (H+); bases accept protons.

Naming of Acids and Bases

• Acids: If the anion ends in -ide, the acid name begins with "hydro-" and ends with "-ic acid" (e.g., HCl: hydrochloric acid).

• Bases: Named as ionic compounds (e.g., NaOH: sodium hydroxide).

Dissociation of Weak Acids and Bases

• Weak Acids/Bases: Partially dissociate in water (e.g., acetic acid, ammonia).

• Equilibrium: Represented by a double arrow ($\rightleftharpoons$).

Dissociation Constant of Weak Acid, Base, and Water

• Acid Dissociation Constant (Ka): $K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}$

• Base Dissociation Constant (Kb): $K_b = \frac{[\text{BH}^+][\text{OH}^-]}{[\text{B}]}$

• Water Dissociation Constant (Kw): $K_w = [\text{H}^+][\text{OH}^-] = 1.0 \times 10^{-14}$ at 25°C

Acid, Base, and Neutral Solution

• Acidic Solution: [H+] > [OH-], pH < 7

• Basic Solution: [OH-] > [H+], pH > 7

• Neutral Solution: [H+] = [OH-], pH = 7