AUQ NA Chemistry Notes PDF

1. Physical vs. Chemical Properties Physical Properties: Properties that can be observed or measured without altering the substance’s chemical identity. ○ Examples: Color, density, melting and boiling points, hardness, electrical conductivity, solubility. Chemi...

1. Physical vs. Chemical Properties Physical Properties: Properties that can be observed or measured without altering the substance’s chemical identity. ○ Examples: Color, density, melting and boiling points, hardness, electrical conductivity, solubility. Chemical Properties: Describe a substance's potential to undergo specific chemical changes, transforming it into a different substance. ○ Examples: Reactivity with acids or bases, flammability, oxidation states. Examples of Chemical Changes: ○ Rusting of iron: Iron reacts with oxygen and water to form rust (iron oxide). ○ Burning wood: Combustion reaction with oxygen, producing carbon dioxide and water vapor. ○ Baking a cake: Ingredients chemically react under heat, changing properties and flavor. 2. Matter and its Forms States of Matter: Matter exists in different states based on temperature and pressure. ○ Solid: Defined shape and volume. Particles are tightly packed and only vibrate in place. ○ Liquid: Defined volume but takes the shape of its container. Particles can move past each other. ○ Gas: No fixed shape or volume; fills the container. Particles move freely and are far apart. ○ Plasma: Ionized gas with high energy, where electrons are free from atoms. Found in stars and neon lights. Examples: ○ Solids: Ice, metals, rocks. ○ Liquids: Water, oil, alcohol. ○ Gases: Oxygen, nitrogen, carbon dioxide. ○ Plasma: Lightning, the sun, fluorescent lights. 3. Conservation Laws Law of Conservation of Mass: States that mass cannot be created or destroyed in a closed system during a chemical reaction. ○ Application in Reactions: The total mass of reactants equals the total mass of products.  ○ Evidence in Reactions: In a reaction in a closed container, if we measure mass before and after, it stays the same, proving conservation. Law of Conservation of Energy: In a chemical or physical process, energy is neither created nor destroyed, but it can change forms. 4. Temperature Conversions Temperature Scales: ○ Fahrenheit (°F): Mostly used in the United States. ○ Celsius (°C): Used worldwide and in scientific contexts. ○ Kelvin (K): Absolute temperature scale used in scientific contexts. Conversion Formulas: Critical Temperature: At -40°, the Celsius and Fahrenheit scales intersect, where -40°C = -40°F. 5. Subatomic Particles and Atomic Models Subatomic Particles: ○ Protons: Positively charged particles in the nucleus. ○ Neutrons: Neutral particles in the nucleus. ○ Electrons: Negatively charged particles that orbit the nucleus. Electric Field Influence: ○ Alpha Particles: Positively charged and deflected by electric fields. ○ Beta Particles: Negatively charged (electrons) and also deflected. ○ Gamma Rays: Neutral and not deflected by electric or magnetic fields. Rutherford’s Atomic Model: ○ Proposed after his gold foil experiment. ○ Stated that atoms have a small, dense, positively charged nucleus with electrons orbiting around it, mostly empty space. 6. Atoms, Ions, and Isotopes  Atoms: The smallest unit of an element, consisting of protons, neutrons, and electrons. Ions: ○ Formed when an atom gains or loses electrons. ○ Cations: Positively charged ions (loss of electrons). ○ Anions: Negatively charged ions (gain of electrons). Isotopes: Atoms of the same element with different numbers of neutrons, resulting in different atomic masses. ○ Example: Carbon-12 and Carbon-14 (both are carbon but with different neutron counts). Atomic Mass Calculation: ○ Weighted average of the masses of all naturally occurring isotopes of an element based on their abundance. 7. Chemical Compounds and Formulas Chemical Compounds: ○ Formed when two or more elements chemically combine in fixed proportions. ○ Represented by chemical formulas (e.g., H₂O, CO₂). Formulas: ○ Show the elements and their proportions in a compound. ○ Empirical Formula: Simplest whole-number ratio of elements. ○ Molecular Formula: Exact number of each type of atom in a molecule. Calculating Mass in Reactions: ○ Use molar mass (grams per mole) to determine the mass of reactants and products in a chemical equation. 8. Balancing Chemical Equations Balancing: Ensures that the same number of each atom appears on both sides of a chemical equation. Coefficients: Numbers placed before formulas to balance equations. Stoichiometry: ○ The study of quantitative relationships between amounts of reactants and products. ○ Mole Ratios: Derived from coefficients to calculate amounts in reactions. 9. Oxidation States  Oxidation Number: The charge an atom would have if electrons were completely transferred. ○ Example: In H₂O, oxygen has an oxidation state of -2, and hydrogen has +1. Rules: ○ Pure elements have an oxidation state of 0. ○ Sum of oxidation states in a neutral compound is 0. ○ Sum in an ion equals the ion’s charge. 10. Types of Chemical Reactions Types: ○ Redox Reactions: Involve electron transfer, with one substance oxidized (losing electrons) and another reduced (gaining electrons). ○ Precipitation Reactions: Two solutions combine to form an insoluble solid, known as a precipitate. ○ Neutralization Reactions: An acid and a base react to form water and a salt. ○ Synthesis Reactions: Two or more reactants combine to form a single product. ○ Decomposition Reactions: A single compound breaks down into two or more simpler substances. 11. Chemical Yields and Percentages Theoretical Yield: Maximum possible amount of product calculated based on reactants. Actual Yield: Amount of product actually obtained from a reaction. Percent Yield: 12. Electrolytes and Conductivity Electrolytes: ○ Substances that dissolve in water to produce ions, allowing the solution to conduct electricity. ○ Strong Electrolytes: Fully dissociate into ions (e.g., NaCl). ○ Weak Electrolytes: Partially dissociate (e.g., acetic acid). Conductivity: ○ The ability of a solution to conduct electric current, dependent on ion concentration. ○ Applications include batteries, saltwater conductivity testing, and medical applications (e.g., electrolytes in the human body). For cation: To determine the symbol for this cation, let’s go through the details step-by-step: 1. Number of Protons: ○ Since the atom has 21 electrons and a +3+3+3 charge, it originally had 24 electrons to be neutral. ○ Therefore, it must have 24 protons (since the number of protons in a neutral atom equals the number of electrons). 2. Element Identification: ○ The atomic number (number of protons) identifies the element. An atom with 24 protons is chromium (Cr) on the periodic table. 3. Atomic Mass: ○ The atomic mass of this isotope is the sum of the protons and neutrons. ○ 4. Charge: ○ It has a +3 charge because it has 3 more protons than electrons (21 electrons in total). 5. Symbol: So, the correct symbol for this atom is Cr3+\text{Cr}^{3+}Cr3+ with a mass number of 52. Formula for calculating the moles of an element given the mass of element: Multiply the grams by 1 mole/molar mass of the element To determine whether a compound's name ends with "ate" or "ide," you can follow these guidelines: 1. Understanding the Suffixes: "ate": This suffix typically indicates a polyatomic ion that contains oxygen. It is commonly used for oxyanions (anions that contain oxygen) derived from acids. For example: ○ Nitrate (NO₃⁻) ○ Sulfate (SO₄²⁻) ○ Phosphate (PO₄³⁻) "ide": This suffix is used for simpler anions or non-metal elements in their anionic form. It indicates a binary compound, often consisting of two elements. Examples include: ○ Chloride (Cl⁻) ○ Oxide (O²⁻) ○ Sulfide (S²⁻) 2. Identifying the Type of Compound: Oxyanions vs. Binary Compounds: ○ If the compound involves a polyatomic ion with oxygen, it will likely end in "ate." ○ If the compound consists of just two elements (like a metal and a non-metal), it will often end in "ide." 3. Analyzing the Chemical Composition: For "ate" Compounds: ○ Check if the compound contains an oxygen atom along with another element. For example, calcium sulfate (CaSO₄) contains sulfate (an oxyanion). For "ide" Compounds: ○ Look for compounds that consist of two elements without oxygen. For instance, sodium chloride (NaCl) is a simple ionic compound. 4. Common Patterns: Acids and Their Anions: ○ Acids that end in "ic" (e.g., sulfuric acid) correspond to "ate" ions (e.g., sulfate). ○ Acids that end in "ous" (e.g., sulfurous acid) correspond to "ite" ions (e.g., sulfite). Transition Metals:  ○ In the case of transition metals, the suffix "ide" can indicate a simple anion formed from that metal, while "ate" typically indicates the presence of oxygen. 5. Examples for Clarity: "ate": ○ Nitrate: NO₃⁻ (from nitric acid) ○ Sulfate: SO₄²⁻ (from sulfuric acid) "ide": ○ Chloride: Cl⁻ (from chlorine) ○ Sulfide: S²⁻ (from sulfur) Step 1: Calculate the molar mass of K2Cr2O7K_2Cr_2O_7K2​Cr2​O7​ Multiply each by their molar mass Calculate the mass of K in 1.00×102 g1.00 1.00x10^2 / product of step 1 1. *Chemical Name of SbCl₃:* - Answer: E) antimony trichloride - Explanation: Sb is the symbol for antimony, and Cl is for chlorine. Since there are three chlorine atoms, the correct name includes the prefix "tri" for three, resulting in "antimony trichloride." Antimony has only one oxidation state here, so the (III) in answer D is unnecessary. 2. *Chemical Name of KNO₂:* - Answer: A) potassium nitrite - Explanation: KNO₂ consists of potassium (K) and the polyatomic ion nitrite (NO₂⁻). "Potassium nitrite" is the correct name as the NO₂⁻ ion is called nitrite. The options mentioning "nitrate" or "nitrogen dioxide" are incorrect because those refer to different compounds. 3. *Formula for Barium Nitride:* - Answer: D) Ba₃N₂ - Explanation: Barium (Ba) has a +2 charge, and nitride 👎 has a -3 charge. To balance the charges, we need three barium ions (+6) and two nitride ions (-6) to make a neutral compound, resulting in Ba₃N₂. 4. *Name of Cr(ClO₄)₃·6H₂O:* - Answer: A) chromium(III) perchlorate hexahydrate - Explanation: Cr(ClO₄)₃ indicates that chromium is in the +3 oxidation state (hence chromium(III)), and ClO₄⁻ is the perchlorate ion. The ·6H₂O denotes six water molecules attached, so the name includes "hexahydrate." 5. *Chemical Name of HIO:* - Answer: D) hypoiodous acid - Explanation: HIO is an oxyacid of iodine where iodine is in its lowest oxidation state, forming "hypoiodous acid." The prefix "hypo-" and suffix "-ous" are used to indicate this oxidation state, similar to "hypochlorous acid" for HClO. 6. *Balancing K₂O + H₂O → KOH:* - Answer: C) 3 - Explanation: The balanced equation is K₂O + H₂O → 2 KOH. Adding the coefficients, we get 1 + 1 + 2 = 3. 7. *Balancing Na + HCl → NaCl + H₂:* - Answer: A) 4 - Explanation: The balanced equation is 2 Na + 2 HCl → 2 NaCl + H₂. Adding up the coefficients gives us 2 + 2 + 2 + 1 = 4. 8. *Mass of Lead Sulfate from Lead in a Battery Reaction:* - Answer: A) 57.6 g - Explanation: The balanced equation is Pb + PbO₂ + 2 H₂SO₄ → 2 PbSO₄ + 2 H₂O. The molar mass of Pb is 207.2 g/mol, so 41.4 g of Pb corresponds to 0.2 mol of Pb. Since 1 mol of Pb produces 1 mol of PbSO₄, 0.2 mol will yield 0.2 mol of PbSO₄, with a mass of 57.6 g. 9. *Combustion of C₆H₁₂ in Excess Oxygen:* - Answer: C) 5.23 g CO₂, 2.38 g H₂O - Explanation: In the combustion of C₆H₁₂, each molecule produces 6 CO₂ and 6 H₂O. We calculate the masses based on stoichiometry. For 1.900 g of C₆H₁₂, we obtain 5.23 g of CO₂ and 2.38 g of H₂O based on the molar ratios and molar masses. 10. *Moles of I₂ from Reaction with KMnO₄, KI, and H₂SO₄:* - Answer: A) 0.443 mol - Explanation: We use stoichiometry with the balanced equation. KMnO₄ is the limiting reagent in this case. By calculating moles of each reactant and using the stoichiometric ratio, we find that 0.443 mol of I₂ is produced. 11. *Percent Yield in KClO₃ Reaction:* - Answer: B) 10.0% - Explanation: The theoretical yield is calculated based on the molar mass of KClO₃ and the balanced reaction. Given 14.0 g of KClO₃, the maximum possible yield of KCl is calculated. The actual yield is 1.40 g, so percent yield is (1.40 g / theoretical yield) × 100 = 10%. 12. *Unreacted Potassium Dichromate in Reaction with Zn and H₂SO₄:* - Answer: B) 3.65 g - Explanation: We determine the limiting reagent (zinc) and calculate the moles required for the reaction. Based on stoichiometry, the amount of potassium dichromate left unreacted is calculated to be 3.65 g. 13. *Poorest Electrical Conductor:* - Answer: A) 0.5 M CH₃OH  - Explanation: CH₃OH (methanol) is a covalent compound and does not dissociate in solution, so it is a poor conductor. In contrast, the other options are ionic compounds or weak electrolytes that conduct electricity better than methanol. 14. *Pairs Reacting to Produce Precipitation and Neutralization:* - Answer: A) Mg(OH)₂(aq) and H₂SO₄(aq) - Explanation: This reaction produces MgSO₄ (soluble) and H₂O but can also produce a precipitate of Mg(OH)₂ under certain conditions. The reaction is also an acid-base (neutralization) reaction. 15. *Non-Redox Reaction:* - Answer: C) MgSO₄(s) → MgO(s) + SO₃(g) - Explanation: In this reaction, there is no change in oxidation states for any element. The other options involve changes in oxidation states, which are characteristics of redox reactions. 16. *False Statement about Oxidation State:* - Answer: A) MnO₄²⁻ with +6 oxidation state - Explanation: In MnO₄²⁻, Mn actually has an oxidation state of +7, not +6. The other statements correctly reflect the oxidation states of the underlined elements.