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Chapter X! I. Energy A. Potential Energy! Rock on a cliff! Stretched rubber band! Wood burning! Meteor showers! Molecular Motion! B. Kinetic Energy! Bullet from a gun! Page 1! C. Temperature and Kinetic Energy! How would you describe the molecular motion:! ! a. Of the black solid before it is put in the water compared to after?! T=100.0 oC T=25.0 oC T=32.5 oC ! b. Of the water (liquid) before the metal is put in compared to after?! Page 2! D. Law of Conservation of Energy! 1. Visible changes in energy! Describe the changes in energy that occur during a lightning storm:! Sky before lightning! Sky during lightning! Page 3! 2. Exothermic Chemical Reactions! Match burning! 3. Endothermic Chemical Reactions! NH4Cl NH4+ Cl- Page 4! II. Measuring Changes in Energy A. Units of Energy! 1. calorie:! ! ! ! 2. Calorie:! ! ! 3. Joule:! Water temperature increase is measured Water absorbs the heat B. The Meaning of the Sign (+ or -)! Food is burned here How food calories are determined! Warming hands by a fire! Page 5! C. Equation and Sample Problem! q = m c !T How much energy is required to heat 1 cup of water (mass = 237 g) from 20.5 oC to 95.6 oC? The specific heat capacity of water is 4.184 J/goC.! Page 6! D. Specific Heat Capacities! Substance! Specific Heat Capacity (J/goC)! How much energy is required to raise 1 g Al by 1.0 oC?! water! rubbing alcohol! aluminum! iron! How much energy is required to raise 1 g water by 1.0 oC?! copper! gold! Average temp in Hilo, Hawaii:! ! January:! July:! Average temp in Fargo, North Dakota:! ! January:! July:! Page 7! E. More Calculations and Problems! 1. How much heat is given off when a 0.0327 kg block of aluminum (c = 0.900 J/goC) at 75.0 oC cools down to 36.2 oC?! ! ! ! ! ! ! ! ! ! ! 2. What is the specific heat capacity of a 542 g piece of metal if it takes 12,500 J of energy to raise the temperature of the metal from 15.0 oC to 75.0 oC?! ! Page 8! 3. Two blocks of copper and aluminum (the same mass, at the same initial temperature) sit in the sun in the same location for 20 minutes. Which metal gets hotter? (c of copper = 0.385 J/goC and c of aluminum = 0.900 J/goC )! ! ! ! ! ! ! ! ! 4. What is the final temperature of a 500.0 g block of iron (c = 0.473 J/goC), originally at 22.0 oC, if 359 calories of energy are added?! Page 9! Blank Page Chapter 9! I. Electromagnetic Radiation A. Types / Applications of Electromagnetic Radiation! X-ray! Ultraviolet (sunburns)! Visible light! Microwaves (ovens and cell phones)! B. “Light” (Radiation) has both Wave and Particle Nature:! 1. Wave Nature:! 2. Particle Nature:! Proposing the particle nature of light got Einstein his Nobel prize in 1921! Page 1! C. Properties of Waves! 1. Wavelength (!):! 2. Frequency (v):! 0 sec.! Radio stations broadcast in Megahertz! 1 sec.! Page 2! D. Electromagnetic Spectrum! Page 3! Page 4! E. Wavelength and Frequency Relationship to each other! 1. Speed of light:! Earth’s circumference is 40,075,000 meters: light would go around Earth 7.5 times in 1 second if it travelled in a curve.! ! Sunlight takes 8 minutes 17 seconds to reach the Earth (distance of 92,000,000 miles)! 2. A red light has a wavelength of 680 nanometers (680 nm). What is its frequency? ! (Given: 1 meter = 1 x 109 nm)! c = ! v! Page 5! F. Wavelength and Frequency Relationship to Energy! 1. Equation:! E = h v! 2. Calculate the energy of one photon of red light that has a ! = 680 nm (6.8 x 10-7 m).! Page 6! G. Proportionalities! 1. Wavelength and Frequency:! c = ! v! 2. Energy and frequency! E = h v! 3. Energy and wavelength! E = h c! !" Page 7! H. Conceptual Problems! A certain cellphone transmits a 9.60 x 108 Hz signal and a light bulb emits a 5.50 x 1014 Hz signal.! 1. Which has a higher frequency? ! !Cellphone or Light bulb! 2. Which has a longer wavelength? ! !Cellphone or Light bulb! 3. Which type of radiation has a higher energy? !Cellphone or Light bulb! 4. Which wave travels faster? ! !Cellphone or Light bulb! Gamma Rays! X- Rays! Ultraviolet! ! Visible! Infrared! Microwaves! Radio! Page 8! II. Atomic Emission Spectrum A. Emission Spectrum! Flaming salts! Fireworks! Page 9! B. Where light comes from! C. The size of the transition (jump) relates to the color of light! E! E! Page 10! D. Why only certain colors are emitted! 1. The Bohr Model! 2. Quantized Energy Levels! purple blue green red ! Page 11! Page 12! III. Modern View of the Atom A. Wave Model of the Atom! Moths have random movement around a porch light, but have a high probability to be found near it.! probability picture! “orbital”! B. Principle Energy Level (n)! Page 13! C. Orbital Shapes! D. Orbital Orientations! E. Electron Spin! Page 14! F. Orbital “Order” and the Periodic Table! Page 15! IV. Electronic Structure A. Orbital Diagrams and Electronic Configurations! Atom! # e"! Orbital Diagram! Electron Configuration! H! He! Li! B! C! O! Page 16! Atom! # e"! Mg! Orbital Diagram! Electron Configuration! ! ! ! ! ! ! ! Ti! Which atom has more unpaired electrons, F or N?! Page 17! B. Ions and Electron Configuration! 1. Oxygen2"! 2. Silicon2+! Page 18! noble gases! 2 C. Noble Gases, Valence Electrons and Condensed Electron Configurations! 1. Noble Gases are all very unreactive.! He 4.00 10 Ne 20.18 18 Ar 39.95 36 Kr 83.80 54 Xe 2. Valence Electrons + Condensed e- Configurations! Lithium! 131.3 Sodium! Orbital Diagram! Electron configuration! Condensed Electron configuration! # valence electrons! Atom / Ion! Page 19! Condensed Electron Configuration! # Valence e"! Ti! P! Br! Ni+! Page 20! D. Noble Gas Electron Configuration is very common! 1. Common Ion Forms! Atom! Condensed e" configuration! Most Common Ion! Ion e"Configuration! K! Mg! Cl! 2. Use electron configurations to explain why aluminum is commonly 3+ in an ionic compound.! Page 21! Blank Page Chapter 10! I. Bonding A. Bonding Affects Properties! Diamond! Substance! Graphite (coal, pencil lead)! Buckminsterfullerene! Molecular Structure! B. What is a “Bond”?! No! Yes! Interactions within H2O! Interactions within LiF! Page 1! C. Ionic Bonds (in ionic compounds)! 1. Bonding in KCl! Potassium! Chlorine! Potassium chloride! e− configuration! dot structure! 2. “Octet Rule”! 3. Bonding in CaBr2! Page 2! D. Covalent Bonds (mostly in molecular substances)! 1. Bonding in F2! One Fluorine! Other Fluorine! Molecule F2! e− configuration! dot structure! 2. “Octet Rule”! 3. Bonding in O2! Page 3! E. Polar Covalent Bonds! 1. Covalent vs. polar covalent vs. ionic! 2. Electronegativity! Page 4! 3. Bonding in HF! 4. Classify each substance and draw a representation of each.! Substance! Na2O! HCCH! NO! Ionic (I), Covalent (C) or Polar Covalent (PC) ?! Draw a representation showing charge or partial charge if present! Page 5! Substance! N2! MgBr2! HCl! Ionic (I), Covalent (C) or Polar Covalent (PC) ?! Draw a representation showing charge or partial charge if present! 5. Which bond is more “polar”?! H H C H O H H H C N H H H Then decide which bond is more polar and show partial charges.! Cl---Se vs. Br---Se! Page 6! II. Lewis Structures (for covalent bonding) A. Uses for Lewis Structures! Explaining shape! CO2 is straight! NO2 is bent! PCl3 has Cl’s on one side! Explaining bond strengths! The C-C bond in acetylene (C2H2) is 1.3 times stronger than the C-C in ethylene (C2H4)! B. Method! 1. Sum all the valence electrons. If ionic, add an electron for each negative charge; subtract an electron for each positive charge.! ! 2. Draw the skeleton (often the least electronegative atom is in the middle). Connect each atom to the central atom with single bonds.! ! 3. Arrange the valence electrons in pairs to give each atom an octet (8 e− ) and hydrogen atoms a duet (2 e−).! ! 4. If there are too few electrons to give all atoms octets, form multiple bonds.! Page 7! C. Examples! Substance! H2O! NCl3! CH2FBr! Valence e− ! Lewis Structure! Page 8! Substance! HCN! NH4+! CO2! Valence e− ! Lewis Structure! Page 9! D. Resonance Structures! Carbonate:! E. Bonding in Polyatomic Ions! Sodium carbonate:! Page 10! III. Molecular Shape A. Valence Shell Electron Pair Repulsion Theory (VSEPR)! B. Molecular Shape and Bond Angles! # groups !! # lone pairs "! 2! 3! 4! 0! 1! 2! Page 11! C. What Counts as a Group?! D. Practice! Give the shape and bond angles for each substance.! Shape! Bond angles! Page 12! Draw the Lewis Structure for each, then give the shape of the substance.! Formula! H2OBr +! SiH4! CS2! COF2! ClO3−! SCl2! Lewis Structure! Shape! Formula! Lewis Structure! ! Shape! Page 13! E. Molecular Polarity! 1. Review of bond polarity! 2. Molecular polarity: is there a “net” dipole?! Formula! CO2! H2O! CCl4! CHCl3! Lewis Structure! Polarity! Page 14! F. Problems! Draw the Lewis structure, including resonance structures if possible.! Determine the shape of the substance, bond angles and polarity.! Formula! SO2! PF3! CHOCl! Lewis Structure! Shape! Bond angles! Polarity! Page 15!
