Chemistry 20 Study Guide PDF

What you need to remember or figure out ahead of time: ⭑ How to manipulate formulas ⭑ How to convert numbers to/from scientific notation ⭑ Significant digit rules & rounding ⭑ Balancing equations & writing compounds/chemical formulas using ion charges ⭑ Where & what everything is on your data book + how to use it ⭑ Remember what units need to be used for different calculations (most are implied in the data book but not all) ⭑ Molar Mass and Chemical Amount mean the same thing Unit A: Chemical Bonding 3.1 Structured Formulas - Show how elements are bonded to each other Valence Electrons and Orbitals - Valence electrons: the electrons found on the outermost “shell” of an atom - Orbitals: the spaces around an atom where the electrons orbit. It is a region of space where there is a high likelihood of finding electrons of a particular energy. ○ Bonding Electrons, Bonding Pairs and Lone Pairs - Bonding Electrons: Individual electrons in an atom that are not paired. They are available to bond with another electron - Bonding Pairs: A pair of electrons that is shared between two atoms to form a bond - Lone pairs: Paired electrons that are not currently shared or bonded with anything else. Lewis Dot Diagrams ⬆ - Represents valence electrons using dots - Element that can form the most bonds is most central, then the element with the least amount of atoms Octet Rule - Atoms ideally want to have 8 electrons in their valence shell Electronegativity - The tendency for an atom to draw electrons to itself when chemically attached to another element. The element with a higher electronegativity will attract electrons from the other element. - On the periodic table, elements lower and closer to the left have a lower electronegativity, while elements further up and to the right have a higher electronegativity Chemical Bonding ○ Covalent/Molecular: Between non-metals, where the atoms share electrons ○ Ionic: Electrons are transferred from metal to non-metal to create a positive cation and a negative anion, which attract each other (also has crystal lattice structure) ○ Metallic: Metals typically consist of a closely packed “sea” of electrons which move freely between the cations 3.2 Molecular Elements Polyatomic & Diatomic Elements Molecular Compounds Bonding Capacity Types of Formulas ○ Empirical ○ Molecular ○ Lewis ○ Structural ○ Stereochemical Determining Lewis Formulas ○ Molecular Compounds ○ Polyatomic Ions 3.3 VSEPR Theory (Valence Shell Electron Pair Repulsion) ○ Stereochemistry ○ Determining Molecular Shape Linear Angular Trigonal Planar Trigonal Pyramidal Tetrahedral Dipole Theory ○ Electronegativity and Bond Polarity ○ Bond Polarity and Molecular Polarity Polar Substances and Solubility Unit B: Gases General Study Guide 4.1 4.2 4.3 4.4 Pt 1 4.4 Pt 2 4.1 Introduction to Empirical Properties of Gases Empirical properties are observable and measurable characteristics of gases that help us understand their behavior. These properties are fundamental to understanding how gases interact in various conditions. Key Empirical Properties of Gases Empirical properties describe the macroscopic behavior of gases without reference to the underlying molecular theory. Key properties include: Pressure (P): ○ Pressure is the force exerted by gas molecules when they collide with the walls of a container. ○ Measured in units like kilopascals (kPa) or atmospheres (atm) ○ STP (Standard Temperature and Pressure) on data booklets ○ SATP (Standard Ambient Temperature and Pressure) on data booklets Volume (V): ○ The space that a gas occupies. ○ Gases expand to fill the container they are in. ○ Measured in liters (L) or cubic meters (m³). Temperature (T): ○ Temperature is related to the average kinetic energy of gas molecules. ○ Must always be measured in Kelvin (K) when working with gas laws. Convert Celsius to Kelvin by adding 273.15 Number of Moles of Gas (n): ○ Measured in moles (mol), the amount of gas corresponds to the number of particles. ○ Often derived using the ideal gas law or other gas-related equations. Pressure and the Behavior of Gases Pressure is a crucial empirical property that is influenced by other variables like volume, temperature, and the amount of gas. Barometer: Tool used to measure atmospheric pressure. Gas Laws Gas laws mathematically relate the empirical properties of gases. While 4.1 focuses on the properties of gases, it's important to recall the foundation of gas behavior through these laws: 1. Boyle’s Law: P1V1 = P2V2 ○ At constant temperature, the pressure of a gas is inversely proportional to its volume. 2. Charles’s Law: V1/T1 = V2/T2 ○ At constant pressure, the volume of a gas is directly proportional to its temperature (in Kelvin). 3. Gay-Lussac’s Law: P1/T1 = P2/T2 ○ At constant volume, the pressure of a gas is directly proportional to its temperature (in Kelvin). Units and Conversions It is essential to use consistent units when working with gases: Temperature: Always use Kelvin for gas law calculations. Pressure: Common units include kPa, atm, and mmHg. Volume: Generally in liters (L) or cubic meters (m³). Amount of gas: Measured in moles (mol). Graphical Representation Graphing the relationships between different gas properties helps visualize gas laws: Pressure-Volume Graph (Boyle’s Law): A downward-sloping curve (hyperbolic). Volume-Temperature Graph (Charles’s Law): A straight line passing through the origin when temperature is measured in Kelvin. 4.2 Properties of Gases: Compressibility: Gases can be easily compressed because the particles are far apart. Expansion: Gases expand to fill any container due to the random movement of particles. Low Density: The density of gases is much lower than that of solids and liquids, as particles are spread out. Kinetic Molecular Theory (KMT) Although empirical properties are based on observations, the Kinetic Molecular Theory provides a deeper explanation of gas behavior. KMT explains the following: Particles in constant random motion: Gas molecules are always moving in straight lines until they collide with something (another molecule or the container walls). Elastic collisions: Gas particles collide without losing energy. Negligible particle size: The volume of individual gas particles is very small compared to the total volume of the gas. No intermolecular forces: Gas molecules do not attract or repel each other in ideal gas behavior. Direct relationship between temperature and kinetic energy: As the temperature of a gas increases, the average speed and kinetic energy of the particles increase. Gay-Lussac’s law of combining volumes: When measured at the same temperature and pressure, volumes of gaseous reactants and products of chemical reactions are always in simple ratios of whole numbers. 4.3 Molar Volume of a Gas ○ Molar volume is the volume occupied by one mole of any gas at a specified temperature and pressure. Conditions for Molar Volume ○ The molar volume is dependent on temperature and pressure. When these conditions change, the volume of a gas also changes. 4.4 Ideal Gas Law: The Ideal Gas Law combines several gas laws (Boyle’s, Charles’, and Avogadro’s laws) into a single equation that relates pressure, volume, temperature, and amount of gas. PV = nRT Where: P = pressure (in kPa) V = volume (in L) n = number of moles of gas R = universal gas constant (8.314) T = temperature (in Kelvin) How the Ideal Gas Law Works: It allows us to calculate the missing property of a gas when three of the other variables are known. For example, if you know the pressure, volume, and temperature, you can find the number of moles of gas. Real vs. Ideal Gases Ideal Gas: Follows all gas laws perfectly under all conditions. Real gases approximate ideal behavior under many conditions, but deviations occur at high pressures and low temperatures. Real Gas: Has intermolecular forces and the volume of particles is significant in certain conditions. Deviates from ideal gas laws, especially under high pressure and low temperature. Deviations from Ideal Gas Behaviour: Real gases deviate from the ideal gas law at low temperatures and high pressures. Under these conditions: Gas particles do experience intermolecular forces, which are not accounted for in the ideal gas law. The volume of gas particles becomes significant, especially at high pressure. At low temperatures, gas molecules move more slowly, and attractions between them become more noticeable, causing the gas to deviate from ideal behavior. At high pressures, gas molecules are forced closer together, so their actual volume becomes important, making the ideal gas law less accurate. Unit C: Solutions, Acids & Bases Naming and Identifying Acids and Bases Note: aqueous or the subscript (aq) means that it is dissolved in water. This represents the H2O in an equation, so the formula H2O does not always need to be written. Identifying acids ○ Typically have Hydrogen (H) at the beginning of their formula (HCL, HNO3) or carboxylic acid (COOH). Acids start out molecular, then ionize to form H3O. Naming Acids IUPAC Naming - “-ide” used for an acid that has Hydrogen + One Non-metal - “-ate” used for an acid where the anion is a polyatomic ion ending in “ate” - “ite” used for an acid where the anion is a polyatomic ion ending in “ite” Classical Naming - “Hydro-” used for an acid that has Hydrogen + One Non-metal - “-ic acid” used for an acid where the anion is a polyatomic ion ending in “ate” - “-ous acid” used for an acid where the anion is a polyatomic ion ending in “ite” Identifying bases ○ When dissolved in water, bases dissociate to form Hydroxide (OH) Ex: NH3 + H2O → NH4 + OH NaOH + H2O → Na(aq) + OH(aq) Naming Bases ○ Bases are generally named like any other ionic compound (metal name “hydroxide”) Ex: NaOH ~ Sodium hydroxide KOH ~ Potassium Hydroxide 5.1 Solution: A homogeneous mixture (we can’t separate the components) where one substance is dissolved in the other, comprised of: ○ Solute: The substance being dissolved (Ex. Salt, Sugar) ○ Solvent: The dissolving medium (Typically water) Properties of Aqueous Solutions Aqueous solutions all have water as the solvent. ○ Electrolytes: Substances that dissociate into ions and allow the solution to conduct electricity. A compound is an electrolyte if its aqueous or molten state conducts electricity. Ex: Sodium Chloride (NaCl), Hydrochloric Acid (HCl), Potassium Nitrate (KNO3) ALL acids are electrolytes because they can ionize in water MOST bases are electrolytes, but not all Strong electrolytes fully dissociate into ions in the solution and have high conductivity weak electrolytes only partially dissociate into ions, leading to weak conductivity Electrolytes include: - Soluble ionic compounds - Molecular compounds that are acids - Strong acids and bases - Weak acids and bases (weak electrolytes) ○ Nonelectrolytes: Substances that do NOT dissociate into ions and do NOT conduct electricity Ex: Sucrose (sugar), ethanol, pure water Typically include covalent compounds that do not ionize in water Exhibit no conductivity when dissolved Most molecular compounds that are NOT acids are nonelectrolytes 5.2 When a substance dissolves, particles separate from each other and disperse throughout the solution. Solvation (All solutions): The process of surrounding solute particles with solvent molecules to break up the solute and form a solution. The solvent breaks up the solute. Dissociation (Ionic compounds break up): The process by which an ionic compound separates into its ions when dissolved in water. Ex: NaCl → Na+ + Cl- Note: Charges must be included because the compound is separating into its ions, which are charged, not its molecules/atoms. Refer to data booklet for the charges of each ion. Dissociation Equations: Represent how compound dissociate 1. Write the Ionic compound 2. Include the chemical formula 3. Break up the compound into each element and write the charge for each element to form ions 4. Balance the equation Ionization (Molecular compounds become ionic): The process in which a neutrally charged molecule reacts with water to form ions ○ Acid ionization forms H+ CH3COOH(aq) → H+(aq)+ CH3COO−(aq) ○ Base ionization forms OH- and/or binds to an H+ Ex: NH3 (aq) + H2O(l) → NH4 +(aq)+OH−(aq) - NH3 Gains a hydrogen to form NH4 - The remaining Hydrogen and Oxygen from water form OH Energy Changes ○ Breaking bonds Uses/absorbs energy (endothermic reaction - cold) ○ Forming bonds Releases energy (exothermic reaction - warm) Substances in Water (To represent the equations: A = Anions B,C = Cations) ○ Molecular: disperse as individual molecules (Equation: BC(s/l/g) → BC(aq)) ○ Ionic: dissociate as individual cations and anions (Equation: AC(s) → A+(aq) + C-(aq)) ○ Base (Ionic Hydroxide): dissociate as cations and hydroxide ions (Equation: AOH(s) → A+(aq) + OH-(aq)) ○ Acid: ionize to form new hydrogen ions and anions (Equation: HC(s/l/g) → H+(aq) + C-(aq)) 5.3 Percentage Concentration ○ c = percent concentration (% W/V) ○ m = mass (g) ○ V = Volume (L) c = (Volume of solute/Volume of Solution) 100% c = (mass of solute/Volume of solution) 100% c = (mass of solute/mass of solution) 100% Parts Per Million Concentration ○ 1ppm = 1 mg/L ○ 1ppm = 1g/106mL ○ 1ppm = 1mg/kg (solids) Amount Concentration ○ c = amount concentration (mol/L) ○ n = Chemical amount/number of moles (mol) n = m/M M = m/n m = n M ○ V = Volume (L) c = n/V Concentration of Ions ○ Ionic compound (Strong Electrolyte) 1. Write the dissociation equation 2. Determine the number of ions produced (Balance the Equation) 3. Determine the molar concentration of the solution (mol/L) 4. c (# of moles of one ion/# of moles of whole compound) = ion concentration ○ Strong Acid (Ionize completely) 1. Write the ionization equation 2. Determine the number of ions produced (Balance the Equation) 3. Determine the molar concentration of the solution (mol/L) 4. c (# of moles of one ion/# of moles of whole compound) = ion concentration 5.4 Standard Solution Stock Solution Preparing Standard Solutions Using Dilution (+ calculations) 5.5 Solubility: Solubility is a measure of how much of a substance (solute) can dissolve in a certain amount of solvent (water) to form a saturated solution at a specific temperature. Saturated solution: This is when no more solute can dissolve in the solvent at that specific temperature. Any extra solute will just stay as undissolved solid. Saturation: the point at which a solution can no longer dissolve any more solute at a given temperature. Solubility in Water Generalizations For solids and gasses, the closer the substance is to being a liquid, the easier it mixes with water in liquid form. ○ Solids More soluble in water as the temperature increases (think melting) ○ Gases More soluble in water when temperature is lower and the pressure is higher (condensation) ○ Liquids Varies depending on the liquids Liquids that dissolve into each other are called miscible Liquids that do not are immiscible ○ Elements Generally low solubility in water Dynamic Equilibrium - In a saturated solution, particles dissolve and crystallize (solidify) at the same rate, keeping the saturation level constant. ○ Dynamic: Constantly changing or in motion ○ Equilibrium: Balance within a system While the saturation stays the same, the solution is not static; there is still something happening inside the saturated substance. 6.1 The Nature of Acidic and Basic Solutions Arrhenius Definitions ○ Acids - Substances that ionize in an aqueous solution to form hydrogen ions HCL(aq) → H+(aq) +Cl--(aq) Acid⇧ ○ Bases - Substances that dissociate to form hydroxide ions in an aqueous solution Modified Arrhenius Definitions - A modified modern version of Arrhenius theories 6.2 Neutral water (pH 7) contains trace amounts of both hydronium and hydroxide ions, due to a very slight ionization, so slight that it will usually show no conductivity. Lower pH = More acidic, Higher pH = More basic pH and pOH (pOH is basically ph but opposite) ○ pH + pOH = 14 ○ Hydronium Calculations (Formulas need to be memorized) pH = -log(H3O concentration) H3O concentration = 10-pH ○ Hydroxide Calculations (Formulas need to be memorized) pOH = -log(OH concentration) OH concentration = 10-pOH 6.3 Using acid-base indicators (listed in data booklet) 6.4 Arrhenius and Modified Arrhenius definitions (know the difference) Neutralization reactions 6.5 Empirical Properties of acids and bases Strong vs Weak Acids (strong acids listed in data book) Strong vs Weak Bases Strong and Weak vs Concentrated and Dilute Kind of Confusing-ish Example Questions: ○ What is the difference between a strong dilute acid and a weak concentrated base? ○ Can the concentration of a strong dilute acid and a weak concentrated acid be the same? Polyprotic Substances Unit D: Stoichiometry 7.1 Limitations of Reaction Equations ○ Chemical reaction equations do NOT include: pressure and temperature conditions progress and process measurable quantities Reaction Assumptions ○ Chemical reactions are ASSUMED to be: spontaneous fast quantitative stoichiometric Types of Chemical Equations ○ ★ Net Ionic Equations: An equation that only includes the parts of the equation that react. These are the elements that change state or charge in a reaction. ○ Non-Ionic Equations: Equations where the elements and compounds are written as atoms, molecules and formula units Cu(s) + 2 AgNO3(aq) → 2 Ag(s) + Cu(NO3)2(aq) ○ Total Ionic Equations: The equation that shows all parts of the reaction, including the entities that do not react or change AND the dissociated ions. Cu(s) + 2 Ag+(aq) + 2 NO3-(aq) → 2 Ag(s) + Cu2+(aq) + 2 (NO3)-(aq) ⬆ Dissociated ⬆ ○ Spectator Species: Any entities that do NOT change Cu(s) + 2 Ag+(aq) + 2 NO3-(aq) → 2 Ag(s) + Cu2+(aq) + 2 (NO3)-(aq) Cu starts as a solid with no charge, then becomes aqueous with a 2+ charge. Ag starts out aqueous with a 1+ charge, then becomes an neutrally charged solid. NO3 is aqueous with a 1- charge on both sides of the equation, making it a spectator ion. You would not include it in the net ionic equation. Limiting and Excess Reagents 7.2 Gravimetric Stoichiometry Applications of Stoichiometry Yield - The amount of a substance produced in a chemical reaction Percent Yield = Actual Yield / Predicted Yield ○ Percent Yield - The “efficiency” of a reaction. Amount of product that was actually produced compared to the theoretical amount that should have been produced. ○ Actual Yield - the amount of product that is produced based on the results of a lab ○ Predicted Yield - the amount of product that should be produced based on theoretical stoichiometric calculations 7.3 Gas Stoichiometry ○ Using the Ideal Gas law to find values required for the question (finding the volume after completing stoichiometric calculations, converting from volume to moles, etc) ⬐PV = nRT (know how to manipulate formulas) V = nRT/P n = PV/RT 7.4 Solution Stoichiometry ○ Manipulating concentration formula depending on the question ⬐c = n/V n = cV V = n/c 8.1 Qualitative Analysis Colorimetry - The process of identifying ions by their flame color or the color of their aqueous solution. (Tables in Data Book) ○ Flame Color - Color of flames produced when the ion is exposed to an open flame ○ Solution Color 8.2 Selective Precipitation - The process used to remove a metal from a solution by forming an insoluble compound, which results in the formation of a precipitate. 8.3 Finding the Limiting Reagent ○ Balance the chemical equation ○ Convert given amounts to moles (n = mass/Molar Mass (Chemical Amount) ○ Use mol ratio to determine how much of each is needed for a complete reaction Ex: Ex: Mg (s) + 2HCl (aq) MgCl2 (aq) + H2 (g) For every mole of Mg you will need double the moles of HCl (2:1 ratio - required/given) OR For every mole of HCl you will need half the amount of Mg (1:2 ratio) Remember a ratio is the same as a fraction (1/2 or 2/1) ○ ALL of the limiting reagent will react Calculating Excess Reagent ○ Balance the chemical equation ○ Convert given amounts to moles (n = mass/Molar Mass) ○ Use mol ratio to determine how much of each is needed for a complete reaction Ex: Ex: Mg (s) + 2HCl (aq) MgCl2 (aq) + H2 (g) For every mole of Mg you will need double the moles of HCl (2:1 ratio - required/given) OR For every mole of HCl you will need half the amount of Mg (1:2 ratio) Remember a ratio is the same as a fraction (1/2 or 2/1) ○ Whichever reactant will have leftover is the excess reagent ○ Amount of excess = Total amount - Amount consumed Determining how much excess to use in a reaction ○ Use stoichiometry to determine the minimum mass that you’ll need for a complete reaction (use moles for the calculation, then convert back to mass) ○ Add 10% of the minimum required amount (multiply minimum by 1.10) to ensure the full reaction 8.4 Titration - a method used to determine the amount concentration of substances in a solution. Parts of a Titration ○ Sample - The solution of unknown concentration ○ Titrant - The standard solution (whose concentration is known to a high degree) that is added to the sample If the sample is a base, the titrant will be an acid and vice versa ○ Equivalence point - occurs when the chemical amount of the added titrant is equal to the sample. The solution will become neutral at this point in a strong acid-strong base titration (which is all you deal with in Chem 20) ○ Endpoint - the observable colour change used to stop the titration. Ideally, the endpoint and the equivalence point should be identical, however, the endpoint usually overor underestimates the equivalence point. Appropriate indicators Steps for a Titration 1. Prepare the sample: Use a pipette to measure a specific volume of the sample and transfer it into a clean Erlenmeyer flask. 2. Add an indicator: Add a few drops of a suitable indicator to the sample in the Erlenmeyer flask. 3. Set up the burette: Fill the burette with the titrant, ensuring no air bubbles remain in the burette's tip. 4. Start titration: Slowly open the burette tap to let the titrant drip into the flask while swirling the flask gently. 5. Watch for color change: Stop adding the titrant as soon as the indicator changes color, signaling the reaction's endpoint. 6. Record the volume: Note the burette's reading to determine how much titrant was used. 7. Repeat for accuracy: Perform the titration multiple times until you have 3 trials where the amount of titrant used is within 0.2 mL of each other, then calculate the average of the 3 volumes 8. Use this volume to calculate the concentration of the sample

Chemistry 20 Study Guide PDF

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