Matter, Energy, Kinetics, Equilibrium

1. Matter, Energy, Kinetics, Equilibrium

2. Brief Review of States of Matter • Solids: definite shape and definite volume. • Particles are held in fixed positions. • Liquids: no definite shape, but definite volume. • Particles slide past each other. • Liquids take the shape of their container. • Gases: no definite shape, and no definite volume. • Gases completely fill any container, regardless of volume.

3. Solid Liquid Gas

4. Heating Curve

5. Vocabulary of Heating Curves • Fusion is the melting of a solid into a liquid • The amount of heat needed to cause melting is the heat of fusion • Vaporization is the boiling of a liquid into a gas • The amount of heat needed to cause boiling is the heat of vaporization • Sublimation occurs when a solid changes directly into a gas • Fusion, vaporization, and sublimation are endothermic reactions • Endothermic = absorbs heat

6. A cooling curve

7. Condensation is the change of a gas into a liquid Freezing is the change of a liquid into a solid Deposition occurs when a gas changes directly into a solid Condensation, freezing, and deposition are all exothermic reactions Exothermic = releases heat During phase changes, the extra heat is being used to overcome the forces in the solid or liquid, not to change the temperature

8. During the diagonal portions of the heating curve, the kinetic energy of the substance increases. • The substance is being heated. • Only one phase is present. • During the straight portions of the heating curve, the potential energy of the substance increases. • The substance is undergoing a phase change. • Both phases are present. • The extra heat is being used to cause the phase change, not to raise the temperature.

9. Heating of liquid 1 phase present KE increases Boiling of liquid 2 phases present PE increases KE constant Heating of gas 1 phase present KE increases Heating of solid 1 phase present KE increases Melting of solid 2 phases present PE increases KE constant

10. Temperature • The temperature of a substance is a measure of its average kinetic energy • The more kinetic energy, the greater the temperature • The temperature shows which way the heat is flowing between two objects • Heat always flows from the warmer object to the cooler object

11. THINK ABOUT IT: When you place ice cubes into a glass of warm soda, what happens? Do the ice cubes cool the soda, or does the soda warm the ice cubes? • The soda warms the ice cubes.

12. Celsius and Kelvin • In order to convert Celsius into Kelvin: K = °C + 273

13. Ex: What Kelvin temperature is equal to 40oC? • Known: Celsius temp. = 40o • Formula: K = oC + 273 (**Table T**) • Solve: K = 40oC + 273 K = 313K

14. Ex: What Celsius temperature is equal to 100oK? • Known: Kelvin temp. = 100 • Formula: K = oC + 273 • Solve: 100K = oC + 273 oC= 100K - 273 Answer =-173oC

15. Temperature Scales

16. Energy • Energy is the ability to do work • There are many forms of energy: light, heat, atomic, kinetic, potential, etc. • In exothermic reactions, heat is released • In endothermic reactions, heat is absorbed

17. Measuring Energy • The energy that is involved in chemical reactions is measured in calories • 1 calorie is the amount of heat needed to raise the temperature of 1g of water by 1°C • Heat energy can also be measured in joules (J)

18. Measuring Energy q = mCΔT • q is the amount of heat in joules • m is the mass of the substance • C is the specific heat of the substance (given) • ΔT is the difference in temperatures

19. How many joules are absorbed when 41.0g of water are heated from 20.1oC to 48.5oC ? Step 1: List the “knowns” and formula to be used & “unknowns” m = 41.0g Known: Specific heat of water (C) = 4.18 J/g•oC (This is from Table B) ∆T = 48.5oC - 20.1oC = 28.4oC Formula from table Tq = mC∆T Step 2: Solve for q (how many joules) q = (41.0g)(4.18 J/g•oC)(28.4oC) q = 4867.192 = 4867 J= 4.87 x 103J

20. q = m × heat of fusion q = m × heat of vaporization Energy can also be measured using the heat of fusion and the heat of vaporization Key word for heat of fusion is “melt” Key words for heat of vaporization is “vaporize” or “evaporate”

21. How many joules are required to melt 410g of ice at 0oC ? Step 1: List the “knowns” and formula to be used & “unknowns” m = 410g Heat of Fusion = 334 J/g (This is from Table B) Formula from table Tq = mHf Step 2: Solve for q (how many joules) q = (410g)(334 J/g) q = mHf q = 136,940 J = 136.9kJ (you can divide by 1,000 to get kJ; kilo = 1,000)

22. Examples: 2. How many joules of energy are required to vaporize 520g of water at 100oC and 1 atm ? Step 1: List the “knowns” and formula to be used & “unknowns” m = 520g Heat of vaporization = 2260 J/g (This is from Table B) Formula from table Tq = mHv Step 2: Solve for q (how many joules) q = mHv q = (520g)(2260 J/g) q = 1,175,200 J = 1175.2kJ = 1175 kJ (you can divide by 1,000 to get kJ; kilo = 1,000)

23. Kinetic Molecular Theory • 1. All gases are made up of particles that are in constant, random, straight-line motion • 2. When gas particles bump into each other, a transfer of energy occurs, but the total amount of energy in the system stays the same • 3. The volume of the individual gas particles is extremely small compared to the volume of the space they move in

24. 4. No attractive forces exist between gas particles • The KMT is only entirely true for ideal (perfect) gases • For real gases: • 1. There is a small volume to gas particles • 2. There are attractive forces between gas particles • Real gases behave most like ideal gases at high temperatures and low pressure • The behavior of real gases can be seen best at low temperatures and high pressure

25. Gas Relationships • Volume and Pressure are inversely related (opposites). • If the volume of a gas decreases, pressure increases. • If the volume of a gas increases, pressure decreases. • Temperature and Volume are directly related. • If the temperature of a gas decreases, volume decreases. • If the temperature of a gas increases, volume increases.

26. Temperature and Velocity are directly related. • The higher the temperature, the greater the velocity. • Temperature and Pressure are directly related. • If the temperature of a gas decreases, pressure decreases. • If the temperature of a gas increases, pressure increases. • When temperature, pressure, & velocity increase, the KINETIC ENERGY increases.

27. Avogadro’s Hypothesis When the volume, temperature, and pressure of two gases are the same, they contain the same number of molecules. 22.4 liters of any gas at STP contain 1 mole of the gas. One mole of any substance contains 6.02 × 1023 molecules.

28. STP (Standard Temperature & Pressure) • Standard temperature = 0°C • Standard pressure = 1 atm

29. Combined Gas Law P1V1 P2V2 ____ ____ = T1 T2 - must use Kelvin, not Celsius - if any condition remains constant, cross it out - P can be atm, kPa, torr, mm Hg - V can be L or mL

30. Kinetics • Kinetics deals with how quickly chemical reactions occur • In order for a reaction to occur, particles must bump (collide) into each other • Collision theory • Remember: the particles at the start of a reaction are reactants, and the resulting particles are products

31. Factors That Affect Rate • Nature of the reactants: • Substances with covalent bonds react slower than substances with ionic bonds • Covalent substances have more bonds to break • Covalent particles must have more energy when they collide • Concentration: • The higher the concentration of particles, the more collisions occur, the higher the rate

32. Surface area: • The more surface area a substance has, the faster the rate • There is more material to react with • Powders react faster than solid pieces • Pressure: • For gases, an increase in pressure causes an increase in concentration, which increases the rate • No effect on solids and liquids

33. Catalysts: • A catalyst is a substance that increases the rate of a reaction (enzymes) • Catalysts are not changed or used up in reactions • Temperature: • Since temperature is kinetic energy, the higher the temperature, the more energy particles have • High temperature causes more frequent and more effective collisions

34. Potential Energy Diagrams • A potential energy diagram shows the change in potential energy (PE) that occurs during a reaction • PE is shown on the vertical axis • The reaction coordinate (“time”) is on the horizontal axis

35. As the reactants get closer together, their kinetic energy is converted to potential energy If the reactants collide with each other in the proper positions, they make an activated complex An activated complex is a temporary product that only exists for a little while The activated complex eventually breaks apart to form new products

36. Measuring PE • The amount of PE of the reactants is the distance between the reactants line and the horizontal axis • The amount of PE of the products is the distance between the products line and the horizontal axis • The amount of PE of the activated complex is the distance between the peak of the curve and the horizontal axis • The bigger the distance, the bigger the PE

37. You can also compare the relative PEs of the reaction • The amount of PE required to move from reactants to activated complex is the activation energy • The activation energy is the “push” the reaction needs to get going • The difference between the PE of the reactants and the PE of the products is called the heat of reaction (ΔH)

38. activated complex Activation energy reactants PE ΔH products PE of reactants PE of activated complex PE of products REACTION COORDINATE

39. Endo- and Exothermic • When the PE of the reactants is higher than the PE of the products, the reaction is exothermic • Heat is released (−ΔH) • When the PE of the reactants is lower than the PE of the products, the reaction is endothermic • Heat is absorbed (+ΔH)

41. Catalysts • Catalysts make reactions go faster by lowering the amount of activation energy needed • Catalysts provide a new, easier pathway for the reaction • ΔH, the energy of the reactants, and the energy of the products all remain the same • Only the activation energy gets lowered • The “hill” gets shorter

43. Equilibrium • Reactions can go both forwards and backwards • Can make the products or the reactants • When both the forward and reverse reactions occur at the same rate, the reaction is at equilibrium • Equilibrium is represented by⇋instead of an arrow • The concentrations of the reactants and products are constant, but not equal

44. Physical Equilibrium H2O (ℓ) ⇋ H2O (g) H2O (s) ⇋ H2O (ℓ) At the melting point and boiling point of a substance, the phases are in equilibrium At 0°C, both ice and liquid water exist at the same time The amounts of each phase do not need to be equal

45. Solution Equilibrium C6H12O6 (s) ⇋ C6H12O6 (aq) CO2 (g) ⇋ CO2 (aq) Saturated solutions are at equilibrium When solids are placed in liquids, they tend to dissolve at the same rate they recrystallize In liquids like soda, there is a balance between the free gas and the dissolved gas

46. Le Châtelier’s Principle • Le Châtelier’s principle explains how a system at equilibrium responds to stress • Stress makes equilibriums shift • Shifting usually occurs to move away from the stress • This will restore the balance of equilibrium • Use + and − signs above the products and reactants to see which way the equilibrium shifts

47. CH4 (g) + H2O (g) ⇋ 3 H2 (g) + CO (g) Which way will the equilibrium shift if the amount of CH4 (g) is increased? Step1: Mark the increase in CH4 by placing a plus sign in a circle above it. + CH4 (g) + H2O (g) ⇋ 3 H2 (g) + CO (g) Step 2: Any substance on the same side as the + will decrease. The extra CH4 uses them up. ‒ + CH4 (g) + H2O (g) ⇋ 3 H2 (g) + CO (g)

48. Step 3: Any substance across the arrow from the + will increase. The extra CH4 makes more of them. ‒ + + + CH4 (g) + H2O (g) ⇋ 3 H2 (g) + CO (g) So, if we increase CH4, the amount of H2O will decrease, and the amounts of H2 and CO will increase. The equilibrium shifts towards the products.

49. CH4 (g) + H2O (g) ⇋ 3 H2 (g) + CO (g) Which way will the equilibrium shift if the amount of CH4 (g) is decreased? Step1: Mark the decrease in CH4 by placing a minus sign in a circle above it. ‒ CH4 (g) + H2O (g) ⇋ 3 H2 (g) + CO (g) Step 2: Equations hate decreases. The equilibrium will try to replace what is being lost. The only way to make more CH4 is to use up the H2 and CO. ‒ ‒ ‒ CH4 (g) + H2O (g) ⇋ 3 H2 (g) + CO (g)

50. Step 3: Making more CH4 also makes more H2O since they’re on the same side of the arrow. ‒ ‒ ‒ + CH4 (g) + H2O (g) ⇋ 3 H2 (g) + CO (g) So, if we decrease CH4, the amount of H2O will increase, and the amounts of H2 and CO will decrease. The equilibrium shifts towards the reactants.