LECTURE_NOTES_SEPT_10_2024.pdf

CHEM 1100, Section A04 Introductory Chemistry 1: Atomic and Molecular Structure and Energetics Lecturer: Dr. Frank Schweizer ([email protected]) Lectures: Tuesdays | Thursdays 8:30 am - 9:45 am Location: Education 290  Assignment Due Dates...

CHEM 1100, Section A04 Introductory Chemistry 1: Atomic and Molecular Structure and Energetics Lecturer: Dr. Frank Schweizer ([email protected]) Lectures: Tuesdays | Thursdays 8:30 am - 9:45 am Location: Education 290  Assignment Due Dates/Term test dates Assignment Due Dates (by 5:00 pm) Term tests #1 Thursday, Sept 24 Test #1: Oct 08 #2 Tuesday, Oct 03 Test #2: Nov 07 #3 Thursday, Oct 24 #4 Tuesday, Nov 05 #5 Tuesday, Dec 03 Corrected !  Assignment Due Dates/Term test dates Assignment Due Dates (by 5:00 pm) Term tests #1 Thursday, Sept 24 Test #1: Oct 08 #2 Tuesday, Oct 03 Test #2: Nov 07 #3 Thursday, Oct 24 #4 Tuesday, Nov 05 #5 Tuesday, Dec 03 Corrected ! Assignment 1 is now available on UMLearn  SCHEDULE DATE CHAPTER NOTES Sept 05 INTRO/Chapter 3 Sept 10 Chapter 3 Sept 12 Chapter 3 Sept 17 Chapter 3 Sept 19 Chapter 4 Sept 24 Chapter 4 Assignment-1 chapter 3 due, 5 pm Sept 26 Chapter 4 Oct 01 Chapter 4 Oct 03 Chapter 5 Assignment-2 chapter 4 due, 5 pm Oct 08 Term test 1 Oct 10 Chapter 5 Oct 15 Chapter 5 Oct 17 Chapter 5 Oct 22 Chapter 6 Oct 24 Chapter 6 Assignment-3 chapter 5 due, 5 pm Oct 29 Chapter 6 Oct 31 Chapter 6 Nov 05 Chapter 7 Assignment-4 chapter 6 due, 5 pm Nov 07 Term test 2 Nov 12 No class Nov 14 No class Nov 19 Chapter 7 Nov 21 Chapter 7 Nov 26 Chapter 7 Nov 28 Chapter 14 Dec 03 Chapter 14 Assignment-5 chapter 7 due, 5 pm Dec 05 Chapter 14  Supplemental Instruction (SI) What are SI study groups ? Supplemental Instruction sessions are free weekly review sessions that are available to students who want to improve their understanding of certain courses. These voluntary sessions offer students an opportunity to interact on an informal basis so that they can: § ask each other questions about the course § compare notes § solve practice problems together § develop new study strategies § discuss course content  When do SI sessions start ? CHEMISTRY 1100 Section Leader Time Location A01 Aliyhia Bushie Monday (2:30-3:30 p.m.) 386 University College A01 Aliyhia Bushie Tuesday (7:00-8:00 p.m.) Online A02 Kanwar Lubana Tuesday (1:00-2:00 p.m.) 302 Tier A02 Kanwar Lubana Thursday (1:00-2:00 p.m.) 215 Tier A03 Kaiya Biluan Wednesday (11:30 a.m.-12:30 p.m.) 308 Tier A03 Kaiya Biluan Friday (11:30 a.m.-12:30 p.m.) 308 Tier A04 Jodh Ghuman Monday (11:30 a.m.-12:30 p.m.) 308 Tier A04 Jodh Ghuman Thursday (10:00-11:00 a.m.) E2-125 Engineering A05 Olumide Ayeni Monday (1:30-2:30 p.m.) 319 Allen A05 Olumide Ayeni Friday (2:30-3:30 p.m.) 111 Armes A06 Marcus Navasca Monday (4:00-5:00 p.m.) 215 Tier A06 Marcus Navasca Wednesday (4:00-5:00 p.m.) 215 Tier A07 Sara Akram Wednesday (1:30-2:30 p.m.) 319 Allen A07 Sara Akram Thursday (7:00-8:00 p.m.) Online  SI Leader for A04 session For the SA04 session the SI leader is Jodh Ghuman [email protected]. While you are encouraged to attend Jodh’s SI sessions, you are welcome to attend the SI sessions for any section. SUMMARY OF LAST CLASS 3.1 Types of Energy: Objective Key Learning Objective: Recognize the types of energy of interest to chemists energy | (definition) the ability to do work (displacement of an object) Work: displacement of an object against an opposing force 3.1 Types of Energy 1. Kinetic Energy 2. Potential Energy: a. Electrical b. Chemical c. Mass 3. Thermal Energy 4. Radiant Energy Kinetic Energy Every moving object 🏃 has kinetic energy (Ekinetic) ↪ includes objects, people, atoms… ↪ Ekinetic depends on the velocity (v) and mass (m) of the moving object 1 Ekinetic = mv2 2 ↪ measures in Joules (J), the SI unit of all forms energy m2 1 Joule = 1 J = kg s2 3.1 Types of Energy 1. Kinetic Energy 2. Potential Energy: a. Electrical b. Chemical c. Mass 3. Thermal Energy 4. Radiant Energy Electrical Potential Energy ↪ energy from charged (positive / negative) ions held a small distance apart q1 / q2 are the charges of the ions q1q2 Eelectrical = k r = distance between ions (in pm) r k = 2.31 x 10−16 J pm ← r→ ← r→ ←r → q1 > 0 q2 < 0 q1 > 0 q2 > 0 q1 < 0 q2 < 0 ↪ Eelectrical < 0 ↪ Eelectrical > 0 ↪ Eelectrical > 0 ↪ attraction! ↪ repulsion! ↪ repulsion! Chemical Potential Energy ↪ atoms bind together to form molecules through the influence of electrical forces ↪ chemical energy results from the sum of attraction of (negatively charged) electrons and (positively charged) nuclei in molecules ↪ overall, electrons and nuclei in a molecule will arrange so as to minimize (make most negative) electrical potential energy ↪ we call this “bond energy” or “chemical The formation of chemical energy” more generally bonds makes molecules more stable than a collection of separated atoms… Mass as Potential Energy As formulated by Einstein, mass represents a form of potential energy This equivalence is represented by the famous equation: E = mc2 While not relevant in standard chemical transformations, nuclear energy results from the transformation of large amounts of nuclear mass into useable energy (heat) https://en.wikipedia.org/ 3.1 Types of Energy 1. Kinetic Energy 2. Potential Energy: a. Electrical b. Chemical c. Mass 3. Thermal Energy 4. Radiant Energy Thermal Energy (= Kinetic Energy) The total energy of random movements of molecules in a sample is defined as its thermal energy ↪ because its energy associated with movement, the thermal energy of a sample is actually a measure of a sample’s average molecular kinetic energy (Ēkinetic) R = ideal gas constant (8.131451 J K-1 mol-1) 3RT Ēkinetic = NA = Avogadro’s number (6.02213 x 1023 mol-1) 2NA T = temperature (K) ↪ so as the temperature 🌡 (T) of a sample increases, so does its average molecular kinetic energy (molecules “move” faster) ↪ higher T = greater thermal E while lower T = lesser thermal E ↪ note: changing thermal E is the cause; change in T is the effect Thermal Energy (= Kinetic Energy) Molecules contribute to thermal energy in three main ways: Translational kinetic energy (movement in space) Rotational kinetic energy → Vibrational energy (by internal movement of atoms) → In a gas, atoms and molecules can free translate/rotate/vibrate In a liquid, the nearby presence of other atoms/molecules restricts this freedom of motion In a solid, atoms/molecules can no longer translate freely but then can still vibrate back and forth 3.1 Types of Energy 1. Kinetic Energy 2. Potential Energy: a. Electrical b. Chemical c. Mass 3. Thermal Energy 4. Radiant Energy Radiant Energy Radiant energy describes the energy contained in electromagnetic radiation (aka... LIGHT!) And not just visible light – radiation is classified according to its wavelength/frequency which is associated with its energy: Matter will interact with radiation differently depending on the energy of the radiation; for example, microwaves cause water molecules to rotate faster, which increases the thermal energy of a sample and heats it up! NEW MATERIAL Chapter 3: Energy and its Conservation, sections 3.1 – 3.6 Learning Objectives Recognize the types of energy of interest to chemists. ✔ Understand the first law of thermodynamics and the concepts of heat and work. Understand the origins of energy changes in chemical reactions. Apply the principles of calorimetry to determine energy changes in a chemical reaction. Understand and calculate enthalpy and internal energy. Energy Transfers and Transformations Energy can be transferred from one type to another (🍔 → 🏃 ; 🧗 → 🙀) and/or from one object to another (☕→🧊) ↪ importantly, the total energy never changes though it can seem that way sometimes if energy is “lost” as heat 🔥 Energy transformations accompany chemical reactions: ↪ If a reaction releases energy, chemical energy is converted into other forms of energy ↪ A reaction that releases energy almost always produces thermal energy, but chemical energy can be transformed into a variety of other forms, depending on the conditions under which the reaction occurs 3.2 Thermodynamic Terms ther mo dy nam ics | ˌTHərmōˌdīˈnamiks | plural noun [treated as singular] the branch of physical science that deals with the relations between heat and other forms of energy (such as mechanical, electrical, or chemical energy), and, by extension, of the relationships between all forms of energy ↪ study of energy transfers and transformations ↪ “how much energy goes where” 3.2 Thermodynamic Terms ↪ study of energy transfers and transformations ↪ “how much energy goes where” Key terms: 👆 System: whatever we want to describe and study by itself Surroundings: everything else but the system; defined as “not the system” Boundary: what separates the system from its surroundings ↪ matter/energy can move across! ⚠ Learning Objective Alert! ⚠ Understand the first law of thermodynamics and the concepts of heat and work Sometimes referred to as the Law of Conservation of Energy and it is generally postulated as: Energy can neither be created nor destroyed but only changed from one form to another Joule or The energy of the universe is constant or The energy of a system which is isolated wikipedia.org from its surroundings is constant 3.2 Thermodynamic Terms: Heat 🔥 A system can exchange thermal energy with its surroundings ↪ we refer to the energy transferred in this way as heat (= q) which we measure in joules (J = kg m2 s-2; SI unit alert!) ↪ remember 1 J is pretty small, so we usually discuss “kJ mol-1” ↪ chemical reactions typically result in heat flow in one direction or the other 3.2 Thermodynamic Terms: Heat 🔥 Flow and Temperature Some heat flows result only in changes in temperature ↪ in these cases, the amount of heat can be determined from the change in temperature (ΔT = Tfinal - Tinitial) 3.2 Thermodynamic Terms: Heat 🔥 Flow and Temperature Such changes in temperature depend on four factors: 1. ΔT depends on q, the amount of heat transferred 2. ΔT depends on the direction of heat flow: If a substance absorbs heat → ΔT > 0 If a substance releases heat → ΔT < 0 3. ΔT depends inversely on the amount of material; that is, for a given amount of heat transferred (q), the change in temperature will be smaller if there’s more ”stuff” present ↪ For example, ΔT for the transfer of 50 J of heat to 1 mole of a substance is twice as large as ΔT for the transfer of 50 J of heat to 2 moles of the same substance 4. ΔT depends on the identity of the material; different substances absorb heat differently! 3.2 Molar Heat Capacity Molar heat capacity (Cm) is defined as the amount of heat (q) needed to raise the temperature of one mol of substance by one degree Celcius (℃) e.g., Cm of gaseous oxygen (O2) at 295 K is 29.355 J mol-1 ℃-1 while for carbon dioxide (CO2) Cm = 37.11 J mol-1 ℃-1 ↪ heat capacity can also be expressed in terms of mass in which case it has units of J g-1 ℃-1 and is called the specific heat of a substance 3.2 Thermodynamic Terms: Heat 🔥 Flow and Temperature The following equation describes the temperature change resulting from heat flow: q ΔT = nCm q is the amount of heat transferred n is the number of moles of material Cm is the molar heat capacity of the substance in J mol−1 ℃−1 3.2 Thermodynamic Terms: Heat 🔥 Flow and Temperature We can rearrange this equation to calculate the heat transferred from temperature measurements: q = nCm ΔT Remember: energy transfer is directional, so we must keep track of the signs associated with heat flow ↪ for example, putting an ice cube into a cup of hot soup, we know that heat will flow from the soup to the ice cube ↪ that is, q is negative for the soup & positive for the ice cube ↪ but the amount of heat flow is the same! -qsoup = qice cube 3.2 Thermodynamic Terms: Heat 🔥 Flow and Temperature In more general terms, when heat is transferred from the system to the surroundings, we can say: qsurr = -qsys Don’t forget, however, that a change in temperature does NOT equate directly with heat flow! q = nCm ΔT ↪ we need to know the identity and amount of the material absorbing or releasing heat Also, note: ΔT = final temperature - initial temperature Be sure to use the correct sign! If Tf < Ti, then the ΔT will be negative Let’s work through an example! Calculate the temperature change that results from adding 250 J of thermal energy to 0.75 mol of Hg (= mercury) To calculate the temperature change, we’ll need to use this form of the equation: q ΔT = nCm From the information we’ve been given, we know q and n… what about the molar heat capacity of mercury? Cm we can look up in our textbook – Table 3.1 Let’s work through an example! Calculate the temperature change that results from adding 250 J of thermal energy to 0.75 mol of Hg (= mercury) Let’s work through an example! Calculate the temperature change that results from adding 250 J of thermal energy to 0.75 mol of Hg (= mercury) So Cm of Hg(l) is 27.983 J mol-1 ℃-1 We can now insert all this into our equation and solve for ΔT! q ΔT = nCm 250 J ΔT = (0.75 mol)(27.983 J mol−1 ℃−1) = 11.9 ℃ = 12 ℃ (significant figures!)  q Let’s work through an example! ΔT = nCm What about the temperature change that results from adding 250 J of thermal energy to 0.35 mol of Hg? With less mercury, the temperature change is now bigger: 250 J ΔT = (0.35 mol)(27.983 J mol−1 ℃−1) = 25.53 ℃ = 26 ℃ A smaller quantity of Hg experiences a larger temperature change upon absorbing the same amount of thermal energy  q Let’s work through an example! ΔT = nCm What about the temperature change that results from adding 250 J of thermal energy to 0.35 mol of H2O 💧? The molar heat capacity of liquid water is bigger than liquid mercury (Table 3.1): 250 J ΔT = (0.35 mol)(75.291 J mol−1 ℃−1) = 9.5 ℃ 👆 So, liquid water experiences a smaller temperature change compared to an equimolar amount of liquid mercury upon absorbing same amount of thermal energy Let’s work through another example! Hooray! 🎉 Lunch time! Who feels like freshly fried fries? 🍟 An aluminum frying pan that weighs 745 g is heated on a stove from 25 ℃ to 205 ℃ What is q for the frying pan? How much heat is transferred? q So now we’ll need this form of the equation: (ΔT = ) nCm q = nCm ΔT What do we know? ΔT = Tfinal – Tinitial = 205 - 25 ℃ = 180 ℃ Let’s work through another example! Hooray! 🎉 Lunch time! Who feels like freshly fried fries? 🍟 An aluminum frying pan that weighs 745 g is heated on a stove from 25 ℃ to 205 ℃ What is q for the frying pan? How much heat is transferred? From Table 3.1 we can get the molar heat capacity (Cm) for solid aluminum: Let’s work through another example! Hooray! 🎉 Lunch time! Who feels like freshly fried fries? 🍟 An aluminum frying pan that weighs 745 g is heated on a stove from 25 ℃ to 205 ℃ What is q for the frying pan? How much heat is transferred? Plugging these values into our equation… wait – n? q = nCm ΔT We need to convert 745 g of solid aluminum into mols! mass (g) 745 (g) n (mols) = = = 27.61 mol molar mass (g mol−1) 26.98 (g mol−1) Let’s work through another example! Hooray! 🎉 Lunch time! Who feels like freshly fried fries? 🍟 An aluminum frying pan that weighs 745 g is heated on a stove from 25 ℃ to 205 ℃ What is q for the frying pan? How much heat is transferred? Plugging these values into our equation… wait – n? q = nCm ΔT You can get this from the periodic We need to convert 745 g of solid tableinto aluminum on the inside mols! cover of your textbook…. mass (g) 745 (g) n (mols) = = = 27.61 mol molar mass (g mol−1) 26.98 (g mol−1) Let’s work through another example! Hooray! 🎉 Lunch time! Who feels like freshly fried fries? 🍟 An aluminum frying pan that weighs 745 g is heated on a stove from 25 ℃ to 205 ℃ What is q for the frying pan? How much heat is transferred? Plugging these values into our equation… q = (27.61 mol)(24.35 J mol−1 ℃−1)(180 ℃) = 121,027.98 J = 1.21 x 105 J = 121 kJ Let’s work through another example! Hooray! 🎉 Lunch time! Who feels like freshly fried fries? 🍟 An aluminum frying pan that weighs 745 g is heated on a stove from 25 ℃ to 205 ℃ What is q for the frying pan? How much heat is transferred? Plugging these values into our equation… q = (27.61 mol)(24.35 J mol−1 ℃−1)(180 ℃) = 121,027.98 J Let’s check our units! = 1.21 x 105 J = 121 kJ 3.2 Thermodynamic Terms: Work Just like heat, work is a flow of energy between objects or between a chemical system and its surroundings We’ll define work (w) as “the energy used to move an object against an opposing force” ↪ like climbing a mountain, converting chemical energy into gravitational potential energy! ↪ in these terms work is the product of force (F) and displacement (d): w = Fd and the amount of work depends on the magnitude of the opposing force 3.2 Thermodynamic Terms: Work Just as with heat flow, it’s important to keep track of the signs associated with work In chemistry, we’re usually viewing things from the perspective of the system ↪ when a system does work on the surroundings, the sign of w is negative because the system loses energy ↪ when the surroundings do work on the system, the sign of w is positive because the system gains energy But the amount of work is the same for both! That is: wsurroundings = – wsystem Visualizing work "flows" When a system does work on its surroundings... w ↪ wsys < 0, wsurr > 0 reaction surroundings When the surroundings do work on the system... w ↪ wsys > 0, wsurr < 0 reaction surroundings 3.2 Thermodynamic Terms: 1st Law of Thermodynamics Energy can therefore be exchanged between the system and the surroundings in only two ways: as heat (q) or as work (w) system surroundings We can link work and heat to the internal energy change (ΔU) of a system using an equation: ΔU = q + w → this is simply a restatement of the law of conservation of energy → “U” is for internal energy but you may see “E” sometimes… 3.2 State vs Path Functions State Functions: properties that depend only on the conditions that describe the system (“as the crow flies…”) Path Functions: properties that depend on how the change occurs (“how you get there matters”) ↪ distance travelled is a path function route 1 "as the crow flies" A B route 2 3.2 State vs Path Functions energy is a state function – for example, in the combustion of methane (CH4; see below), the overall energy change (ΔU) for the process just depends on the strengths of the bonds being broken or formed – not on how they are broken/formed we can therefore use tables of values for bond energies, for example, to analyze the thermodynamics of a reaction This is a particularly important concept for thinking about the thermodynamics of catalytic ΔU processes (including enzymatic processes!) reactants products CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l) 3.2 State vs Path Functions So, whether this combustion process happens in a car fueled by natural gas or in a furnace, the combustion reaction has the same overall energy change ↪ ΔU depends on the net change in the conditions of the system, not on how that change occurs! 3.3 Energy Changes in Chemical Reactions ⚠ Learning Objective Alert! ⚠ In this section, we want to understand the origins of energy changes in chemical reactions Chemical reactions involve rearrangements of atoms: some bonds break, some bonds form ↪ breaking bonds requires an input of energy (= endothermic) ↪ forming bonds releases energy (= exothermic) * The sum of these inputs / outputs determines the net energy change for the reaction 3.3 Energy Changes in Chemical Reactions When the bonds formed in the products of a chemical reaction are stronger (i.e., more stable) than the bonds broken in the reactants, the system will release energy to the surroundings ↪ we call this a “downhill” reaction When the bonds formed in the products of a chemical reaction are weaker (i.e., less stable) than the bonds broken in the reactants, the surroundings must transfer energy to the system (energy is absorbed by the system) for the reaction to occur (e.g., photosynthesis which requires radiant energy to occur) ↪ we call this an “uphill” reaction 3.3 Energy Changes in Chemical Reactions Notation alert! Chemists sometimes show depict stability implicitly based on the vertical position on a page For example, in this diagram → E or U The products are more stable and so are drawn as lower down on the page compared with the reactants which are less stable and higher up on the page ↪ think of it like there’s an invisible vertical axis marked “Energy” Example Time! The Haber reaction for the formation of ammonia releases energy: N2 + 3 H2 ® 2 NH3 ΔU = -40.9 kJ How much energy is released in the production of 1.00 kg of ammonia? The ΔU (= -40.9 kJ) here is actually a molar value (= kJ mol-1) However, note that in the above reaction 2 equivalents of NH3 are produced! So, ΔUmolar = -40.9 kJ / 2 mol = -20.45 kJ / mol of NH3 … but how many mols of NH3 are in 1.00 kg? Example Time! The Haber reaction for the formation of ammonia releases energy: N2 + 3 H2 ® 2 NH3 ΔU = -40.9 kJ How much energy is released in the production of 1.00 kg of ammonia? The molecular weight (= molar mass) of NH3 is 17.03 g mol-1 m [g] 1000 g n [mols] = = = 58.72 mols MW [g mol−1] 17.03 g mol−1 We now just multiply ΔUmolar by the # of mols of NH3! Energy released = (-20.45 kJ / mol)(58.72 mols) = -1200.8 kJ Example Time! The Haber reaction for the formation of ammonia releases energy: N2 + 3 H2 ® 2 NH3 ΔU = -40.9 kJ How much energy is released in the production of 1.00 kg of ammonia? The molecular weight (= molar mass) of NH3 is 17.03 g mol-1 m [g] 1000 g n [mols] = = = 58.72 mols MW [g mol−1] 17.03 g mol−1 We now just multiply ΔUmolar by the # of mols of NH3! Energy released = (-20.45 kJ / mol)(58.72 mols) = -1200.8 kJ 3.3 Energy Changes: Path “Independence” A change in any state function is independent of path (doesn’t matter how you get there, just the difference between where you start and where you finish!) ↪ a good analogy is altitude 🗻 🏂 Δaltitude 🏂 3.3 Energy Changes: Path “Independence” Energy is a state function: so, the overall energy change in a chemical reaction is independent of how the reaction takes place Consider our old friend, the Path 2 Path 1 combustion of methane: CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l) This figure shows three different ways this reaction could take place Path 3 If we want to determine the overall energy change for the reaction, we can do this using any path we like!  End of class Read up on sections 3.3 and 3.4 for next class !