Chapter 6

Starts with the ability to understand

1. Ionic Nomenclature & Symbols: 📝

* Recognizing common monatomic ions (cations and anions)

Monatomic ions are single atoms that have gained or lost electrons. For instance, Sodium ion (Na⁺) and Chloride ion (Cl⁻) are monatomic ions. This skill requires understanding the periodic table and how elements lose or gain electrons to form ions.

🎯Quiz Time! 🎯

1. 🔵 What do atoms gain or lose to become ions? 🔵 A) NeutronsB) ProtonsC) Electrons

2. 💠 Which of these is a monatomic ion? 💠A) H2OB) Na⁺C) CO2

3. 🟣 How many electrons does a Sodium atom (Na) lose to become an ion? 🟣A) Gains 1 electronB) Loses 2 electronsC) Loses 1 electron

4. ⚪ Which of the following is a monatomic ion? ⚪A) H2OB) Na⁺C) CO2

5. 🔶 Which element loses 2 electrons to form a monatomic ion? 🔶A) SodiumB) CarbonC) Magnesium

🎯 ANSWER KEY 🎯

1. 🅲️ Electrons (Atoms gain or lose electrons to become ions!)

2. 🅱️ Na⁺ (Only Na⁺ is a monatomic ion, which means it's a single atom that has gained or lost electrons.)

3. 🅲️ Loses 1 electron (Sodium is in Group 1 of the periodic table, so it loses 1 electron to have a full outer shell, becoming Na⁺.)

4. 🅱️ Na⁺ (Only Na⁺ is a monatomic ion, which means it's a single atom that has gained or lost electrons.)

5. 🅲️ Magnesium (Magnesium is in Group 2 of the periodic table, so it loses 2 electrons to have a full outer shell, becoming Mg2+.)

* Recognizing common polyatomic ions

🎯 Quiz Time! 🎯

1. 🔵 What is the charge of the nitrate ion (NO₃⁻)? 🔵 A) 1-B) 2-C) 2+

2. 💠 How many atoms are there in the sulfate ion (SO₄²⁻)? 💠A) 4B) 5C) 6

3. 🟣 Which of these is a polyatomic ion? 🟣A) Na⁺B) Cl⁻C) OH⁻

4. ⚪ What two elements make up the ammonium ion (NH₄⁺)? ⚪A) Nitrogen and HydrogenB) Nitrogen and HeliumC) Hydrogen and Helium

5. 🔶 Which of the following is not a polyatomic ion? 🔶A) NO₃⁻B) SO₄²⁻C) Na⁺

🎯 ANSWER KEY 🎯

1. 🅰️ 1- (The nitrate ion (NO₃⁻) has a charge of 1-!)

2. 🅱️ 5 (The sulfate ion (SO₄²⁻) consists of one sulfur atom and four oxygen atoms, for a total of 5 atoms.)

3. 🅲️ OH⁻ (OH⁻ is a polyatomic ion, as it's composed of more than one atom, unlike Na⁺ and Cl⁻ which are monatomic ions.)

4. 🅰️ Nitrogen and Hydrogen (The ammonium ion (NH₄⁺) is made up of nitrogen and hydrogen atoms.)

5. 🅲️ Na⁺ (Na⁺ is a monatomic ion, which means it's a single atom that has gained or lost electrons, whereas NO₃⁻ and SO₄²⁻ are polyatomic ions.)

* Understanding charge neutrality in ionic compounds

This involves understanding that the total positive charge from cations must balance with the total negative charge from anions, resulting in a net charge of zero. For instance, in sodium chloride (NaCl), the +1 charge of sodium (Na⁺) balances with the -1 charge of chloride (Cl⁻).

🎯 Quiz Time! 🎯

🟢 What is the net charge of sodium chloride (NaCl)? 🟢 A) -1B) +1C) 0

🔵 What is the net charge of calcium chloride (CaCl₂)? 🔵A) -1B) +1C) 0

🟡 If an ion has a 2- charge, how many ions with a 1+ charge are needed to balance it? 🟡A) 1B) 2C) 3

🟣 In magnesium oxide (MgO), what is the charge on the oxygen ion? 🟣A) -1B) -2C) +2

🔶 What is the net charge of aluminum sulfate (Al₂(SO₄)₃)? 🔶A) -2B) 0C) +2🎯

ANSWER KEY 🎯🅲️ 0 (The net charge of sodium chloride (NaCl) is zero because the +1 charge from sodium (Na⁺) balances with the -1 charge from chloride (Cl⁻).)🅲️ 0 (The net charge of calcium chloride (CaCl₂) is zero because the +2 charge from calcium (Ca²⁺) balances with the -1 charges from the two chloride (Cl⁻) ions.)🅱️ 2 (To balance a 2- charge, two 1+ charges are needed.)🅱️ -2 (In magnesium oxide (MgO), the oxygen ion (O²⁻) has a -2 charge.)🅱️ 0 (The net charge of aluminum sulfate (Al₂(SO₄)₃) is zero. Each aluminum ion (Al³⁺) has a +3 charge and there are two of them, for a total positive charge of +6. Each sulfate ion (SO₄²⁻) has a -2 charge, and there are three of them, for a total negative charge of -6. Therefore, the net charge is 0.)

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• Section 6.1.2

• (18) Name the following: AgI, KOH, PbSO4, BaCr2O7, Li2CO3

• (19) Write the formulae of the following:

• a. ammonium nitrate,

• b. lead chromate,

• c. hydrogen fluoride,

• d. barium sulfate,

• e. calcium carbonate.

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States of Matter and Phase Changes: ⚗️

Understanding characteristics of solids, liquids, and gases

• Solid substances have fixed shape and volume, liquid substances have fixed volume but not shape, and gaseous substances have neither fixed shape nor volume. Understanding this involves understanding the molecular behavior and kinetic energy in different states.

Differentiating between evaporation and boiling

Evaporation is a surface phenomenon and occurs at any temperature, while boiling involves the whole liquid and occurs at a specific temperature. Understanding this requires knowledge about how temperature affects phase transitions.

Understanding latent heat of fusion and vaporization

These are the amounts of energy required to change a substance from solid to liquid (fusion) and from liquid to gas (vaporization), respectively, without any temperature change. This requires understanding the role of energy in phase transitions.

. a. A liquid is heated at its boiling point. Although energy is used to heat the liquid, its temperature does not rise. Explain.

Pure Substances & Mixtures: 🧪

Identifying elements and compounds as pure substances

Elements are substances composed of identical atoms (like Iron, Fe), and compounds are substances composed of two or more different atoms in a fixed ratio (like Water, H₂O). Identifying them requires understanding atomic structure and chemical bonding.

Distinguishing between homogeneous and heterogeneous mixtures

Homogeneous mixtures are uniform in composition (like salt water), while heterogeneous mixtures are not (like oil and water). Distinguishing them involves understanding the nature of substances and their miscibility.

Understanding colloids and suspensions

Colloids are mixtures where tiny particles are dispersed in another substance but do not settle (like milk), whereas suspensions are mixtures where particles will settle over time (like sandy water). Understanding them requires knowledge of particle sizes and dispersion stability.

Separation Techniques: 🔍

Applying filtration to separate solids from liquids

Understanding and performing simple distillation

Applying sublimation to separate specific substances

• Performing Filtration 🧑‍🔬: Filtering mixtures to separate a solid from a liquid (like coffee grounds from coffee) relies on understanding particle sizes and the principle of filtration.

• Executing Distillation 🧪: Separating a mixture of liquids with different boiling points, such as in the production of alcoholic beverages, requires knowledge of boiling points and the distillation process.

• Applying Sublimation ☁️➡️🧊: Understanding how some substances can transition directly from solid to gas (and vice versa) without a liquid phase, such as dry ice (solid CO₂), requires understanding of phase transitions and the special conditions under which sublimation occurs.

Concentration & Molarity: 🧮

Calculating molarity given moles of solute and volume of solution

Performing dilution calculations

Converting between different concentration units (e.g., molarity, molality, % mass)

• Calculating Molarity 🧠💡: This involves using the formula M = n/V, where M is molarity, n is the number of moles of solute, and V is the volume of solution. Mastery requires understanding of moles, volume units, and the molarity concept.

• Performing Dilution Calculations 📐📏: Using the formula M1V1 = M2V2, where M and V are the molarity and volume of the solution before and after dilution, respectively. Requires understanding of molarity, volume, and the mathematical relation between initial and final states in a dilution.

• Interpreting Solubility Curves 📈: Understanding how solubility changes with temperature, and being able to use solubility curves to find the solubility of a substance at a given temperature. This requires knowledge of solubility, temperature units, and how to read and interpret graphs.

Reactions in Solution: 💥

Understanding solvation/dissolution process

Predicting precipitates using solubility rules

Writing balanced molecular, total ionic, and net ionic equations

• Balancing Chemical Equations ⚖️: Understanding the law of conservation of mass and ensuring that the number of atoms of each element is the same on both sides of the reaction. This requires knowledge of atomic masses and stoichiometry.

• Predicting Precipitation Reactions ☔️: Determining which combinations of aqueous ionic compounds will produce a precipitate. Requires understanding of solubility rules and ion interaction in solution.

• Writing Net Ionic Equations ✍️: Writing only the species directly involved in the reaction, excluding spectator ions. This requires understanding of complete ionic equations, spectator ions, and the principles of ionic reactions.

Stoichiometry: 🎛️

Balancing chemical equations

Performing stoichiometric calculations with moles

Performing stoichiometric calculations with mass

• Mole-to-Mole Conversions 🔄: Understanding and using the stoichiometric coefficients in balanced chemical equations to calculate how many moles of one substance react with or produce another.

• Mole-to-Mass and Mass-to-Mass Conversions ⚖️: Using molar mass along with stoichiometric coefficients to calculate the mass of reactants or products.

• Limiting Reactant Calculations 📏: Identifying which reactant is completely consumed in a reaction and therefore determines the maximum amount of product that can be formed.

Chemical Formulas & Naming Compounds: ✍️

Writing formulas for ionic compounds

Writing formulas for covalent compounds

Naming ionic and covalent compounds using IUPAC nomenclature

• Writing Chemical Formulas 📝: Being able to correctly write the formula of a compound given its name, which requires understanding of ionic and molecular compounds and their nomenclature rules.

• Naming Compounds 🗣️: Being able to correctly name a compound given its chemical formula, which requires understanding the rules of nomenclature for ionic and molecular compounds.

• Recognizing Polyatomic Ions 👁️: Identifying common polyatomic ions by their charge and composition, such as sulfate (SO4^2-), nitrate (NO3^-), and ammonium (NH4^+).

Ionic Compounds: ⚛️

Understanding the crystal lattice structure of ionic compounds

Understanding the concept of formula units

Predicting properties of ionic compounds based on ionic bonds

• Understanding Ionic Structures 🏗️: Recognizing how positive and negative ions arrange themselves in a crystal lattice, leading to the unique properties of ionic compounds such as high melting and boiling points.

• Predicting Ionic Formulas 👩‍🔬: Using the charges on ions to predict the formula of the compound they will form, such as knowing that Na^+ and Cl^- form NaCl.

• Exploring Properties of Ionic Compounds 🧪: Examining the characteristics that result from their ionic nature, such as electrical conductivity when dissolved in water or melted.

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Safety Practices: 🥽

Understanding the use of personal protective equipment (PPE)

Knowing how to handle and dispose of chemicals safely

Understanding safety protocols for handling flammable and hazardous substances

• Recognizing Hazards 👀: Identifying potentially dangerous substances or procedures in a lab setting, such as toxic or corrosive chemicals, and understanding appropriate responses to these hazards.

• Using Protective Equipment 🧤🥽: Understanding the importance of using goggles, gloves, and lab coats, when they should be used, and how to use them properly.

• Handling Chemical Spills 🧹: Knowing the procedures for safely cleaning up different types of chemical spills, including using absorbent materials and neutralizing agents as necessary.

189. Condensed phases of matter are solid and liquid.

190. Gaseous elements (under room conditions) are found at the top right hand side of the Periodic Table.

191. One gram of steam, H2O (g) causes more severe burns than one gram of water, H2O(l) at 100oC. At the same temperature, both have the same average kinetic energy but steam has a higher potential energy than water.

192. A volatile liquid is a liquid that evaporates at room temperature. A liquid with a low boiling point is easy to vaporize.

193. Vapor pressure of a liquid: is the pressure of the gas above the liquid with which it is at equilibrium (Both liquid and gas exist indefinitely).

194. Vapor pressure of a liquid in a sealed container depends on temperature of the flask. As the temperature increases the vapor pressure of a liquid increases.

195. At the boiling point, the temperature of a pure substance stays constant as the liquid is being heated until all the liquid changes into gas. The heat given to the liquid causes more liquid to change into gas.

196. Molar heat of vaporization is the minimum energy required to change one mole of a substance from liquid to gas at the same temperature.

197. General equation for Molar heat of vaporization: X (l) + heat ⇌ X (g)

198. General equation for Molar heat of condensation: X (g) ⇌ X (l) + heat

199. In general, a substance that has a higher boiling point is expected to have a higher molar heat of vaporization.

200. Minimum conditions for liquid molecules to vaporize: 1) Molecules are supposed to be on the surface. 2) Molecules are supposed to have an average kinetic energy greater than the energy keeping the molecules in the liquid state.

201. Boiling point: is the temperature at which the liquid vaporizes anywhere in the solution.

202. At the boiling point: a. Vapor pressure is equal to the surrounding pressure. b. Bubbles of vapor can form anywhere within the liquid. c. Molecules escape from the surface of the liquid to enter the gas phase as vapor (this also happens at room temperature). d. With increasing altitude, atmospheric pressure decreases and so does boiling point.

203. Normal boiling point: is the temperature at which the vapor pressure is exactly 1 atm or 760 mmHg.

204. Molar heat of fusion: is the energy required to change one mole of a substance from solid to liquid at the same temperature and constant pressure.

225. If we collect the crystals from a freezing aqueous solution, melt it and freeze it again it will freeze at 0⁰C.

226. Diagrammatic representations of elements compounds in the 3 states of matter.

227. Demonstration: filtration

228. Selective solubility

229. Alcohol is flammable, therefore it cannot be heated directly. To heat alcohol, we should use a steam bath or an electric heater.

230. If you need to collect sugar from sugar alcohol solution heat the solution using an electric heater to crystallization point. Leave the solution to cool and crystals to form.