Justin Jade Umali - Attachment_ PDF_ Solutions to Exam Review Questions

Unit 1: Matter & Bonding

Definitions

• Atomic Number: Number of protons in an atom.

• Atomic Mass: Total number of protons and neutrons in an atom.

• Reason for Non-Whole Atomic Masses: Average masses based on the abundance of different isotopes.

Chlorine-37 Atom

• Neutrons: 20

• Mass Number: 37

• Atomic Number: 17

• Electrons: 17

Contributions of Key Scientists

• Dalton: Proposed atomic theory; elements consist of atoms.

• Thomson: Discovered electrons; proposed the plum pudding model.

• Rutherford: Discovered the nucleus through gold foil experiment.

• Bohr: Developed model of the atom with quantized energy levels.

• Quantum Mechanical Model: Describes electrons as wave functions.

Periodic Table & Atomic Size Trends

• Developer: Mendeleev.

• Order: Arranged by atomic mass and properties.

• Atomic Size Trends:

  • Down a Group: Increases due to added energy levels.

  • Across a Period: Increases from right to left because the attraction for outer electrons decreases with fewer protons.

Properties of Elements

• Ionization Energy: Energy needed to remove an electron.

• Electron Affinity: Energy released when an electron is added to an atom.

• Electronegativity: Attraction of an atom for electrons.

• Trends:

  • Down a Group: All three properties decrease due to increasing atomic size.

  • Across a Period: All three properties increase due to increasing nuclear charge.

Lewis Dot Diagrams

• Draw for the following:

  • O

  • Al

  • Na

  • I

  • Xe

Unit 1: Physical Properties of Compounds

Ionic Compounds

• Composition: Metal ion + Non-metal ion.

• Properties:

  • High melting point.

  • Conduct electricity when dissolved.

  • Hard and brittle.

  • Usually soluble in water.

  • Solid at room temperature.

Covalent (Molecular) Compounds

• Composition: Non-metal + Non-metal.

• Properties:

  • Do not conduct electricity.

  • Can be gas, liquid, or solid.

  • Lower melting points than ionic compounds.

Bond Classification

• H2: Pure covalent.

• CH4: Non-polar covalent.

• LiF: Ionic.

• H2O: Polar covalent.

Bonding Structures

• Use Lewis dot and structural diagrams for:

  • Mg + P

  • O + Cl

  • P + H

  • Ca + Cl

  • N + N

  • Identify as polar or non-polar.

Intermolecular Forces

Definitions

• Van der Waals Forces: General term, includes London Dispersion, Induced Dipole, and Dipole-Dipole forces.

• London Dispersion Forces: Temporary dipoles due to electron movement.

• Dipole-Dipole Forces: Attraction between molecules with permanent dipoles.

• Hydrogen Bonding: Strong dipole-dipole forces between hydrogen and highly electronegative atoms (N, O, F).

Intermolecular Forces in Compounds

• HCl: Dipole-dipole.

• H2O: Hydrogen bonds.

• NaCl: Ionic bonds.

• CH4: London Dispersion.

Naming Compounds

Formulas & Names

• CuI: Copper(I) iodide

• HI(aq): Hydroiodic acid

• N2O4: Dinitrogen tetroxide

• H3PO3: Hydrogen phosphite (phosphorous acid in aqueous).

• PBr5: Phosphorus pentabromide

• Fe2O3: Iron(III) oxide

• K3N: Potassium nitride

• H2C2O4(aq): Oxalic acid

• HF(aq): Hydrofluoric acid

• NiSO4• 6H2O: Nickel(II) sulfate hexahydrate

• H2S(g): Hydrogen sulfide

Unit 2: Reactions

Balancing Equations

1. N2(g) + O2(g) → N2O5(g)

2. NaOH(aq) + H2SO4(aq) → H2O(l) + Na2SO4(aq)

Types of Reactions

1. Synthesis: A + B → AB

2. Decomposition: AB → A + B

3. Combustion: Fuel + O2 → CO2 + H2O

4. Single Displacement: A + BC → AC + B

5. Double Displacement: AB + CD → AD + CB (neutralization)

Reaction Examples

1. Ag + O2 → Ag2O: Synthesis

2. Zn + NiBr2 → ZnBr2 + Ni: Single displacement

3. C2H6O + O2 → CO2 + H2O: Complete combustion

4. H3PO4 + NaOH → H2O + Na3PO4: Neutralization

5. SO2 + H2O → H2SO3: Synthesis

6. NaHCO3 → Na2O + CO2 + H2O: Decomposition

7. Ba(OH)2 + H2SO4 → H2O + BaSO4: Neutralization

8. Br2 + KI → KBr + I2: Single displacement

Unit 3: The Mole

Molar Calculations

• 1.5g NaCl: 0.026 moles

• 100g KNO3: 0.8 moles

• 1kg NH3: 60 moles

• 5.0 x 10^25 molecules NO: 83 moles

• 100 million Ne atoms: 2 x 10^-16 moles

Percentage Composition in Al(NO3)3

• Al: 12.666%

• N: 19.731%

• O: 67.602%

Empirical and Molecular Formulas

• 0.0250 mol NaF: 1.05g

• Sodium, Carbon, Oxygen Analysis: Na2CO3

• Fe and O from rock: Fe2O3

• C, H, O analysis: C6H14O2

Limiting Reactants

• Definition: Reactant that determines the amount of product formed.

• Significance: Essential for calculating product yield.

Iron Production Example

• 50g iron(III) oxide reacts to produce 35.0g iron.

Unit 4: Solutions

Solubility Principles

• Oil vs. Water: Oil is non-polar; HCl is polar, allowing it to mix with polar water.

• Solubility with Temperature:

  • Solids: More soluble at high temperatures.

  • Gases: Less soluble at high temperatures.

Wine Alcohol Calculation

• 250 mL wine, 12% V/V: 3.0 x 10 mL pure alcohol.

Molar Concentration Calculations

• 15g KCl in 800 mL: 0.3 M

• 250 mL of 6.00 mol/L NaOH: 60g

• 3.00L of 4.00mol/L H2SO4: 7.50 x 10^2 mL

• Combined NaCl Concentration: 0.60 M

Unit 5: Gases

States of Matter

• Gas: Conforms to the shape of the container; compressible.

• Liquid: Conforms to the shape; incompressible.

• Solid: Fixed shape; incompressible.

Molecular Motion Types

• Vibrational Motion: Present in all states.

• Translational Motion: Present in liquids and gases.

• Rotational Motion: Present in gases and liquids.

Kinetic Molecular Theory Explanations

• Evaporation Cooling: Higher kinetic energy particles evaporate, lowering average energy.

• Gas Compressibility: Gases have more empty space, allowing for compression.

• Pressure Increase: Elevated temperature increases molecular speed, resulting in more collisions.

Diffusion and Pressure

• Increased Pressure: Higher temperature means faster diffusion due to increased kinetic energy.

STP and SATP Values

• STP (Standard Temperature and Pressure): 273K, 101.3kPa.

• SATP (Standard Ambient Temperature and Pressure): 298K, 100kPa.

Gas Law Equations

1. Charles’ Law: V1/T1 = V2/T2

2. Boyle’s Law: P1V1 = P2V2

3. Combined Gas Law: P1V1/T1 = P2V2/T2

4. Ideal Gas Law: PV = nRT

Gas Behavior Calculations

• Piston Volume Change: 500 kPa after reducing to 2.0 mL.

• Volume change with heat: 1.75 L for heated balloon.

• Increase in temperature in fixed volume: Pressure increases.

Mole and Molar Mass Calculations

• Gas Sample Analysis: 0.11 mol, 6.67x10^22 molecules, 135 g/mol.

• CO2 Volume at STP: 16.4 L.

• Volume of O2 at SATP for ZnS reaction: 534 mL.

• Volume of SO2 produced: 356 mL.

Unit 1: Matter & Bonding

Definitions

• Atomic Number: The unique number of protons found in the nucleus of an atom, which determines the element’s identity. For example, an atomic number of 1 corresponds to hydrogen.

• Atomic Mass: The weighted average mass of an atom's isotopes, measured in atomic mass units (amu). It reflects the total number of protons and neutrons in an atom's nucleus.

• Reason for Non-Whole Atomic Masses: Atomic masses are not whole numbers due to the presence of isotopes. Each isotope has a different number of neutrons, affecting the overall mass calculation.

Chlorine-37 Atom

• Neutrons: 20 (indicating that this isotope has a mass number of 37 due to 17 protons and 20 neutrons).

• Mass Number: The total number of protons and neutrons, which for Chlorine-37 is 37.

• Atomic Number: 17, which is characteristic of chlorine.

• Electrons: 17, reflecting a neutral atom configuration.

Contributions of Key Scientists

• John Dalton: Formulated a scientific theory of atoms in the early 19th century, proposing that each element consists of specific atoms, leading to the establishment of the law of multiple proportions.

• J.J. Thomson: Discovered electrons through cathode ray experiments and proposed the plum pudding model, where electrons were embedded in a positively charged 'soup.'

• Ernest Rutherford: Conducted the gold foil experiment, leading to the discovery of the atomic nucleus and demonstrating that most of the atom is empty space.

• Niels Bohr: Developed the Bohr model, introducing the concept of quantized energy levels and electron orbits around the nucleus, which explained the emission spectra of hydrogen.

• Quantum Mechanical Model: Provides a mathematical description of the behavior and interactions of subatomic particles, illustrating that electrons exist in probabilistic clouds rather than fixed orbits.

Periodic Table & Atomic Size Trends

• Developer: Dmitri Mendeleev is credited with creating the first widely recognized periodic table, arranging elements by increasing atomic mass and similar chemical properties.

• Order: Elements are ordered by atomic number in the modern periodic table, which accounts for discrepancies when arranged solely by mass.

• Atomic Size Trends:

  • Down a Group: Atomic size increases as additional electron shells are added, leading to greater distance between the nucleus and valence electrons.

  • Across a Period: Atomic size decreases from left to right as increasing nuclear charge pulls electrons closer to the nucleus, despite the addition of electrons.

Properties of Elements

• Ionization Energy: The minimum energy required to remove an electron from a gaseous atom or ion. It generally increases across a period and decreases down a group.

• Electron Affinity: The energy change when an electron is added to a neutral atom; it varies among elements based on their tendency to accept electrons.

• Electronegativity: A measure of an atom’s ability to attract and hold onto electrons when forming a bond. It tends to increase across a period and decrease down a group.

• Trends:

  • Down a Group: Ionization energy, electron affinity, and electronegativity generally decrease because of the increasing distance of valence electrons from the nucleus as well as increased shielding by inner electrons.

  • Across a Period: These properties generally increase due to the stronger nuclear charge affecting the outermost electrons.

Lewis Dot Diagrams

• Draw Lewis dot diagrams to represent valence electrons and bonding patterns for the following:

  • Oxygen (O)

  • Aluminum (Al)

  • Sodium (Na)

  • Iodine (I)

  • Xenon (Xe)

Unit 1: Physical Properties of Compounds

Ionic Compounds

• Composition: Formed from metal ions (cations) and non-metal ions (anions) through the transfer of electrons.

• Properties:

  • Characteristically have high melting points due to strong electrostatic forces between ions.

  • Conduct electricity in aqueous solutions or molten state due to the mobility of ions.

  • Usually have a crystalline structure, making them hard and brittle.

  • Generally soluble in water, forming electrolytes, though some may not be due to lattice energy overcoming solvation energy.

  • Present as solid at room temperature.

Covalent (Molecular) Compounds

• Composition: Formed from non-metal atoms that share electrons to achieve octets (or duplets in the case of helium).

• Properties:

  • Do not conduct electricity due to the lack of charged particles.

  • Can exist in various states: gases (like O2 and Cl2), liquids (like H2O), or solids (like C6H12O6).

  • Typically have lower melting points than ionic compounds due to weaker van der Waals forces

Bond Classification

• H2: Pure covalent bond (equal sharing of electrons).

• CH4: Non-polar covalent bond (electronegativity difference is less than 0.5).

• LiF: Ionic bond (large electronegativity difference, metal to non-metal transfer).

• H2O: Polar covalent bond (unequal sharing due to significant electronegativity difference).

Bonding Structures

• Use Lewis dot and structural diagrams for:

  • Magnesium + Phosphorus (Mg + P)

  • Oxygen + Chlorine (O + Cl)

  • Phosphorus + Hydrogen (P + H)

  • Calcium + Chlorine (Ca + Cl)

  • Nitrogen + Nitrogen (N + N)

• Identify each as polar or non-polar based on symmetry and electronegativity differences.

Intermolecular Forces

Definitions

• Van der Waals Forces: A collective term for intermolecular forces that includes London Dispersion Forces, Induced Dipole, and Dipole-Dipole interactions.

• London Dispersion Forces: Weak forces arising from temporary fluctuations in electron distribution within molecules, affecting non-polar molecules.

• Dipole-Dipole Forces: Attractive forces between molecules that possess permanent dipoles, found in polar molecules.

• Hydrogen Bonding: A strong type of dipole-dipole interaction occurring between a hydrogen atom covalently bonded to highly electronegative atoms (N, O, F) and another electronegative atom.

Intermolecular Forces in Compounds

• HCl: Exhibits dipole-dipole interactions due to its polar nature.

• H2O: Demonstrates hydrogen bonding, which significantly affects its high boiling point and solvent capabilities.

• NaCl: Contains ionic bonds, which define its crystal lattice structure.

• CH4: Demonstrates London Dispersion forces due to its non-polar nature.

Naming Compounds

Formulas & Names

• CuI: Copper(I) iodide

• HI(aq): Hydroiodic acid

• N2O4: Dinitrogen tetroxide

• H3PO3: Hydrogen phosphite (commonly referred to as phosphorous acid in aqueous).

• PBr5: Phosphorus pentabromide

• Fe2O3: Iron(III) oxide

• K3N: Potassium nitride

• H2C2O4(aq): Oxalic acid

• HF(aq): Hydrofluoric acid

• NiSO4• 6H2O: Nickel(II) sulfate hexahydrate

• H2S(g): Hydrogen sulfide

Unit 2: Reactions

Balancing Equations

• N2(g) + O2(g) → N2O5(g): Demonstrates synthesis as two reactants form a product.

• NaOH(aq) + H2SO4(aq) → H2O(l) + Na2SO4(aq): Represents neutralization; balancing acid and base yields water and salt.

Types of Reactions

• Synthesis: A + B → AB; combining elements to form compound.

• Decomposition: AB → A + B; breaking down compounds into constituents.

• Combustion: Fuel + O2 → CO2 + H2O; typical for hydrocarbons reacting with oxygen.

• Single Displacement: A + BC → AC + B; one element displaces another in compound.

• Double Displacement: AB + CD → AD + CB (neutralization); exchange of ions between two compounds.

Reaction Examples

• Ag + O2 → Ag2O: Evidently synthesis; silver combines with oxygen.

• Zn + NiBr2 → ZnBr2 + Ni: A single displacement reaction where zinc displaces nickel.

• C2H6O + O2 → CO2 + H2O: This represents complete combustion.

• H3PO4 + NaOH → H2O + Na3PO4: Shows neutralization of an acid and base.

• SO2 + H2O → H2SO3: A synthesis reaction forming sulfurous acid.

• NaHCO3 → Na2O + CO2 + H2O: Example of decomposition where sodium bicarbonate breaks down.

• Ba(OH)2 + H2SO4 → H2O + BaSO4: Neutralization representing formation of water and barium sulfate.

• Br2 + KI → KBr + I2: Single displacement reaction where bromine displaces iodine.

Unit 3: The Mole

Molar Calculations

• 1.5g NaCl: 0.026 moles (using the molar mass of NaCl: 58.44 g/mol).

• 100g KNO3: 0.8 moles (KNO3's molar mass is approximately 101.1 g/mol).

• 1kg NH3: 60 moles (NH3's molar mass is approximately 17.03 g/mol).

• 5.0 x 10^25 molecules NO: About 83 moles, using Avogadro's number.

• 100 million Ne atoms: Approximately 2 x 10^-16 moles based on Ne's atomic mass of 20.18 g/mol.

Percentage Composition in Al(NO3)3

• Al: 12.666% of total mass.

• N: 19.731% of total mass.

• O: 67.602% of total mass.

Empirical and Molecular Formulas

• 0.0250 mol NaF: corresponds to 1.05g, factoring NaF's molar mass (41.99 g/mol).

• Sodium, Carbon, Oxygen Analysis: Resulting compound is predominantly Na2CO3 indicating stoichiometric proportions.

• Iron and Oxygen from Rock: Identified compound as Fe2O3 based on the ratio of iron to oxygen.

• C, H, O analysis: Identified compound as C6H14O2, indicating the presence of alcohol.

Limiting Reactants

• Definition: The reactant that is entirely consumed first, limiting the extent of the reaction and thus, the amount of product formed.

• Significance: Essential for calculating expected product yield and maximizing efficiency in chemical reactions.

• Iron Production Example: In a reaction with 50g of iron(III) oxide, the calculated efficiency yields 35.0g of iron, showing the importance of identifying limiting reactants for effective production strategies.

Unit 4: Solutions

Solubility Principles

• Oil vs. Water: Oil is a non-polar substance, while HCl is polar, which allows it to mix with water, a polar solvent, highlighting the principle of like dissolving like.

• Solubility with Temperature:

  • Solids: Solubility typically increases with rising temperature, allowing for more solute to dissolve in a given amount of solvent.

  • Gases: Solubility decreases at higher temperatures as gas molecules escape from the liquid phase.

Wine Alcohol Calculation

• Given 250 mL wine at 12% V/V, it can be calculated that the volume of pure alcohol is approximately 3.0 x 10 mL.

Molar Concentration Calculations

• 15g KCl in 800 mL: Results in a molarity of 0.3 M; utilizing KCl's molar mass (74.55 g/mol).

• 250 mL of 6.00 mol/L NaOH: Requires roughly 60g; calculated from NaOH's molar mass (40.00 g/mol).

• 3.00 L of 4.00mol/L H2SO4: Yields approximately 7.50 x 10^2 mL; H2SO4 has a molar mass of about 98.08 g/mol.

• Combined NaCl Concentration: Mixing solutions yields a combined molarity of 0.60 M.

Unit 5: Gases

States of Matter

• Gas: Takes on the shape of its container and is compressible due to the significant amount of empty space between particles, leading to increased molecular motion.

• Liquid: Conforms to the container's shape but remains incompressible, demonstrating resistance to volume change under pressure.

• Solid: Maintains a fixed shape and volume, with tightly packed particles that exhibit minimal molecular movement.

Molecular Motion Types

• Vibrational Motion: Present in all states of matter due to molecular interactions.

• Translational Motion: Observed predominantly in liquids and gases, allowing particles to move freely past one another.

• Rotational Motion: Important in gaseous and liquid states, aiding in the understanding of gas properties.

Kinetic Molecular Theory Explanations

• Evaporation Cooling: Occurs as higher energy molecules escape from a liquid, resulting in a decrease in the overall temperature of the remaining liquid.

• Gas Compressibility: Gases can be compressed because of the large amount of empty space between particles, allowing for packing under pressure.

• Pressure Increase: An increase in temperature leads to an increase in molecular speed, causing a higher frequency of collisions which raises pressure.

Diffusion and Pressure

• Increased Pressure: When temperature rises, kinetic energy increases leading to faster diffusion rates due to greater particle movement.

STP and SATP Values

• STP (Standard Temperature and Pressure): Defined as 273K (0°C) and 101.3 kPa (1 atm), serving as reference conditions for gas behaviors.

• SATP (Standard Ambient Temperature and Pressure): Defined as 298K (25°C) and 100 kPa, offering a common reference for laboratory conditions.

Gas Law Equations

• Charles’ Law: Describes how gas volume increases with temperature: V1/T1 = V2/T2.

• Boyle’s Law: States that pressure and volume are inversely related: P1V1 = P2V2.

• Combined Gas Law: Utilizes multiple variables: P1V1/T1 = P2V2/T2.

• Ideal Gas Law: Expresses relationship between pressure, volume, temperature, and moles: PV = nRT, where R is the gas constant.

Gas Behavior Calculations

• Piston Volume Change: Example where pressure reaches 500 kPa after compressing gas down to 2.0 mL.

• Volume change with heat: Analysis shows heated balloon expands to about 1.75 L.

• Increase in temperature in fixed volume: Results in a rise in pressure as per gas laws.

Mole and Molar Mass Calculations

• Gas Sample Analysis: Calculated mass shows 0.11 mol, a count of 6.67x10^22 molecules with an average molar mass of 135 g/mol.

• CO2 Volume at STP: Example shows a volume of 16.4 L at standard temperature and pressure.

• Volume of O2 at SATP for ZnS reaction: Calculated to be 534 mL under ambient conditions for reaction equations.

• Volume of SO2 produced: Estimated at 356 mL, demonstrating the stoichiometric relationships in reactions.