Fall Semester 2024 Exam Review

Honors Chemistry Fall Semester 2024 Exam Review 1. Describe the four laws of thermodynamics. Give a real world example of each. Zeroth law (“thermal equilibrium”): Two systems in equilibrium with a third system are in thermal equilibrium with each other; example: if you leave a glass of cold water in a room, the water will absorb heat from the room until it is at the same temperature as the rest of the room First Law (“law of conservation”): Energy can change forms, but is neither created nor destroyed; When you rub your hands together, the amount that your hands heat up equals the mechanical work done to rub your hands together Second Law: Entropy of an isolated system always increases; You would never see a scattered pile of cards leap off the table to form an ordered house of cards, but a house of cards will easily fall and scatter (a system can become more ordered, but it requires work from outside the system) Third Law: Entropy of a system approaches a constant value as its temperature approaches absolute zero; example: the entropy of a perfect crystal is zero 2. A football is left outside on a hot sunny day in Texas, raising the temperature of the gases in the ball. a. What will happen to the pressure of the gas in the ball? Temperature and pressure have a direct relationship - the pressure inside of the ball will increase b. Pick one of the four laws of thermodynamics and describe how it applies to this scenario. Zeroth Law - the air inside of the ball will reach thermal equilibrium with its surroundings First law - the energy from the sun is absorbed by the ball, causing the air molecules to increase in kinetic energy; the energy from the sun is converted into kinetic energy by the air molecules Second law - the energy from the sun causes the air molecules inside of the ball to increase in kinetic energy, therefore increasing the entropy of the system 3. A soccer ball hits the ground, causing the volume of the gas inside to decrease. What happens to the pressure of the gas at that moment? Pressure and volume have an indirect relationship - the pressure inside of the ball will increase. Unit 2: Atomic Structure Learning Targets 2.1 I can illustrate the structure of an atom and identify the location of protons, neutrons, and electrons 2.2 I can explain the functions of the subatomic particles 2.3 I can create/illustrate the development of the modern atomic theory over time using the discoveries of Dalton, Thomson, Rutherford, Bohr, and Heisenberg 4. Complete the following table: Subatomic Particle Proton Neutron Electron Charge (+) positive Neutral (-) negative Location In the nucleus In the nucleus In the electron cloud surrounding the nucleus Function Determines the identity of an atom Provides stability in the nucleus with all the protons Determines the chemical properties of an atom Describe the experiments/discoveries and their model of the atom of the following scientists: ● Democritus: Proposed the idea that you couldn’t divide something in half for forever and you must eventually reach the smallest indivisible part. ○ Model - Atomos ● Dalton: Stated that atoms are hard spheres (indivisible), Atoms of the same element are identical, Combinations of different atoms make compounds, Began thinking about chemical reactions ○ Model - Billiard ball ● Thomson: Stated that atoms are made from electrons scattered through a sea of positive charge like plums in a pudding; experiment showed that cathode rays were deflected toward a positively charged plate and away from a negatively charged plate ○ Model: Plum pudding model ● Rutherford: The atom is mostly empty space with a positive middle; discovered by firing alpha particles at gold foil, some deflected but most passed through (Gold foil experiment) ○ Model: Nuclear model ● Bohr: Stated that electrons move around the nucleus in orbits of specific energy and size, Electrons fit into specific energy levels; Studied line spectra of atoms ○ Model: Bohr model ● Schrödinger: Theory stated that there are clouds called orbits where electrons exist; It is impossible to know both the speed and the position of a subatomic particle at the same time ○ Model: Quantum Mechanical Model (QMM) 5. Explain how the behavior/movement of electrons differ in Bohr’s model and the Quantum Mechanical model of the atom. In the Bohr model of the atom, electrons move in well-defined circular orbits around the nucleus. Each orbit represents a different energy level that electrons can occupy. In the Quantum Mechanical model of the atom, electrons occupy loosely-defined regions called orbitals. The shape of the orbitals differ depending on the energy level and sublevel they are found in. 6. John Dalton is credited as the Father of Modern Atomic Theory. Discuss a portion of his Theory that has since been disproved. Dalton stated that atoms are indivisible, meaning that they cannot be broken down into smaller pieces or portions. This was later disproved by the discovery of the subatomic particles (protons, neutrons, and electrons). Learning Targets 2.4 I can explain and calculate how wavelength, frequency, and energy relate to one another and the electromagnetic spectrum 2.5 I can relate element emission data to the electromagnetic spectrum 7. A typical X-ray machine emits radiation with a frequency of 4.68 x 1018 Hz. Calculate the amount of energy each x-ray is carrying. E=h*v E = (6.63x10-34)*(4.68x1018) E = 3.10 x 10-15 J 8. A chemist observes green light emitted from an element when heated over a Bunsen burner. She estimates the wavelength of this light to be 4.52 x 10-7 m. Calculate the frequency of the light. C = ƛ*v (v) = 3.00x108 / (4.52x10-7) v = 6.64 x 1014 Hz 9. Use Planck’s constant, 6.626 x 10-34, to solve the following: What is the amount of energy carried by a photon that has a frequency of 5.71 x 1014 Hz? E = hv E = (6.626x10-34)*(5.71x1014) E = 3.78 x 10-19 J 10. Using the terms quantum, ground state, excited state and photon, explain the process an electron goes through to produce light. Illustrate your explanation below. An electron in its ground state can absorb a quantum of energy and move to an excited state in a higher energy level around the nucleus. It will then release that energy as a photon of light and move back to its ground state. Unit 3: Nuclear Chemistry Learning Targets 3.1 I can write nuclear equations for alpha, beta, and gamma radiation 3.2 I can distinguish between fission chain reactions and fusion reactions 3.3 I can give examples of applications of nuclear chemistry 3.4 I can do calculations with the half-life of a radioisotope. 🔢 11. Write out the following nuclear equations: ● The beta decay of 5326 Fe 53 26 Fe → 0 -1 e + 5327 Co ● The alpha decay of 3819 K 38 19 K→ 4 2 He + 3417 Cl 12. The half-life of calcium-47 is about 5 days. Starting with 80 g of this isotope, what would be the amount remaining after 20 days? There will be 4 half-life cycles (20/5 = 4) Time 0 5 days 10 days 15 days 20 days Amount 80 g 40 g 20 g 10 g 5g There will be 5 grams left after 20 days. 13. The radioisotope technetium-99 is often used as a radiotracer to detect disorders of the body. It has a half-life of 6.0 hours. If a patient received a 30.0-mg dose of this isotope during a medical procedure, how much would remain 12.0 hours after the dose was given? There will be 2 half-life cycles (12 hours / 6 hours = 2 HL) Time 0 6 hours 12 hours Amount 30 mg 15 mg 7.5 mg There will be 7.5 mg left after 12 hours. 14. Explain the difference between fusion and fission reactions. Which process is used in power plants to produce electricity? Fusion = two smaller nuclei combining into a single larger nucleus Fission = a single large nucleus splitting into smaller nuclei; used in nuclear power plants Fusion Fission Unit 4: Periodic Table Learning Targets 4.1 I can explain the chemical and physical properties in the historical development of the periodic table 4.2 I can determine the number of protons, neutrons, and electrons in an atom using various chemical notations 4.3 I can calculate the atomic mass of an element using its isotopic information 4.4 I can identify and predict the properties of chemical families (alkali metals, alkaline earth metals, halogens, noble gasses, and transition metals) on the periodic table 4.5 I can identify periodic trends based on elemental data (atomic radius, electronegativity, ionization energy) 4.6 I can determine the electron configuration of an element Alkali Metals Location on the periodic table (Group #) Properties Group 1 or 1A Most Reactive Metals Loses Electrons Forms + IONS Cation 1 Valence Electron Alkali Earth Metals Group 2 or 2A Reactive Metals Loses Electrons Forms + IONS Cations 2 Valence Electrons Transition Metals Groups 3-12 Properties of Metals Can form 2 different ions Ductile, malleable, good conductors of heat and electricity Strong, hard and less reactive Halogens Group 17 or 7A Noble Gases Group 18 or 8A Most Reactive Nonmetals (some are gases) Inert (unreactive) Gases Forms - IONS Anions 8 valence electrons 7 valence electrons Valence Electrons = s + d Valence Electrons 1 2 Most have 2 (s shell) since they are in D block 7 8 +1 +2 Varies - depends on loss of e- = they form positive ions -1 n/a Ionic Charge 15. Which scientist rearranged the periodic table in order of increasing atomic number? Dmitri Mendeleev 16. What term describes vertical columns across the periodic table? Families or groups 17. What term describes horizontal or left-to-right rows across the periodic table? Periods 18. What element(s) would be chemically similar to Neon? Helium, Argon, Krypton, Xenon - the noble gases 19. Define the following terms and identify the trends that each follow on the periodic table: ● Electronegativity: Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons ● Ionization energy: Amount of energy required to take an electron away from an atom ● Atomic Radius: Distance from the center of the nucleus to the outermost ring 20. How many neutrons does an isotope of Calcium have with a mass number of 42 amu? 22 neutrons (Mass number - Atomic Number = Neutrons) 21. Naturally occurring chlorine is 75.53% chlorine-35, which has an atomic mass of 34.969 amu, and 24.47% chlorine-37, which has an atomic mass of 36.966 amu. Calculate the average atomic mass of chlorine. Isotope Name: Abundance Mass Abundance*Mass Chlorine-35 75.53%/100 = 0.7553 34.969 amu 0.7553 * 34.969 = 26.412 Chlorine - 37 24.47%/100 = 0.2447 36.966 amu 0.2447 * 36.966 = 9.046 Average: 35.458 amu Average atomic mass = 35.458 amu Unit 5: Chemical Bonding Learning Targets 5.2 I can express the number of valence electrons using electron dot diagrams 5.3 I can represent how different types of covalent bonds are formed by drawing electron dot structures 5.4 I can use the VSEPR theory to predict the shape of a molecule. 22. What is the correct lewis dot structure for nitrogen trihydride (NH₃) and what is its molecular shape and bond angles based on VSEPR? Molecular shape = trigonal pyramidal Bond angles = 109.5 23. Draw the Lewis dot structure for carbon tetrafluoride (CF₄). What is its molecular geometry and bond angles? Molecular shape = tetrahedral Bond angles = 109.5 Learning Targets 5.1 I can determine the bond character based on electronegativity differences 5.6 I can explain how intramolecular & intermolecular attractions determine physical properties of molecular, ionic, and metallic substances 24. Which best represents a molecular compound? Explain your answer. the molecules of most molecular compounds are free to move around relative to each other and are not held in place by any bonds like in an ionic crystal lattice. 25. Explain why the structure of ionic compounds is described as a “crystal lattice”. The ions of an ionic compound are bonded together in an ordered, repeating arrangement of alternating positive and negative charges that form a crystalline structure. 26. Define molecule. How does the arrangement of molecules differ when the compound is a solid, liquid, and gas? Molecule - a group of two or more atoms covalently bonded together. As a solid, molecules are very close together and are not free to move. The molecules of a liquid are slightly farther apart and able to flow around each other. The molecules of a gas are extremely far apart and move independently. 27. Identify three physical properties of ionic compounds that are associated with ionic bonds and relate them to bond strength. ● Solid at room temperature - strong ionic bonds hold ions in place, not allowing them to flow or easily separate ● Brittle - applying pressure to ionic solids will line up ions of similar charge, causing them to repel and break the crystal lattice ● High boiling point - strong ionic bonds require a large amount of energy to break and separate ions 28. List three properties of molecular compounds in the solid phase. Explain these properties in terms of attractions between the molecules. ● Liquids or gases at room temperature - weak attractions between molecules allow them easily flow or vaporize ● Soft and malleable - applying pressure to molecular solids will move molecules around without breaking any covalent bonds ● Low boiling point - weak intermolecular attractions are easily overcome by adding a small amount of energy to spread out the individual molecules Use the electronegativity chart below as a reference for the following questions. 29. Predict whether each molecule is polar or nonpolar based on the information given. Explain each of your predictions. Include details about the molecular geometry and bond polarity of the molecule. a. A molecule has 3 bonds and 1 lone pair of electrons on the central atom. All 3 bonds are between nitrogen (N) and hydrogen (H). Nitrogen is the central atom. The molecular geometry of the molecule is trigonal pyramidal. The electronegativity difference between Nitrogen and Hydrogen is .84 making all the bonds polar (greater than .5 difference). The molecule is polar due to having one lone pair of electrons and polar bonds being on the same side. (asymmetrical) b. The molecular geometry of the molecule is trigonal planar. The electronegativity difference between Boron and Chlorine is 1.12 making all the bonds polar (greater than .5 difference). The molecule is nonpolar due to having polar bonds being on opposite sides. (symmetrical) 30. Rank the intermolecular attractions in order of decreasing strength (IMF’s: hydrogen bonding, london dispersion, dipole-dipole) Hydrogen bonding, dipole-dipole, london dispersion Use the electronegativity chart below as a reference for the following questions. 31. Which of these pure samples would have the highest boiling point? Compound 1 Compound 2 Explain your answer choice in terms of the types and strengths of the intermolecular forces that each molecule experiences. Compound 2 would have the highest boiling point because it experiences dipole-dipole forces (stronger) and compound 1 only experiences London dispersion forces (weaker). Stronger forces require more energy, like higher boiling points, to separate molecules. 32. Describe how metal cations and their valence electrons are arranged in a metallic lattice. The metal cations are arranged in a crystal lattice structure similar to ionic compounds, while their valence electrons are freely moving throughout the lattice in what is called the “sea” of electrons. 33. Define alloy: a mixture of two or more elements with metallic properties 34. Why are alloys so much more useful than pure metals? Alloys are created as they are stronger than each of the metals in their pure form. The presence of multiple atom types makes it more difficult for the layers to slide and rearrange, making the alloy strong. Alloys also have different purposes based on which metals are combined