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Midterm Review Chemistry Level 2

Midterm Review Chemistry Level 2. Chapter 1. What is chemistry? The field of science that studies the composition and structure of matter What is matter ? Anything that has mass and occupies space

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Midterm Review Chemistry Level 2

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  1. Midterm ReviewChemistry Level 2

  2. Chapter 1 • What is chemistry? • The field of science that studies the composition and structure of matter • What is matter? • Anything that has mass and occupies space • Be able to describe each step of the scientific method. Provide a couple examples where you used the scientific method to solve a problem in your life and label each part. • State a Problem, Hypothesis, Experiment, Reach a Conclusion

  3. Chapter 1 - Continued • Theory v. Scientific Law --- What is the difference? Think of examples for both and identify the limitations. Can they be proved/disproved supported/falsified? Which one is more powerful than the other? • Law summarizes observations but theories provide an explanation • Limitations: • A scientific Theory CAN NEVER BE PROVED

  4. Chapter 1 - Continued • What are variables? What step of the scientific method involves manipulating variables? • Variables are factors that can change • Experimenting • In large, why has science advanced as much as it has over the years? • Advancements in science has taken part largely because of advancements in technology

  5. Chapter 2 • Name some examples of physical properties. • Mass, volume, density, boiling point, melting point, solubility, phase (solid, liquid or gas), color, texture, hardness, etc…

  6. Chapter 2 (continued) • Define the particle arrangement of a solid, liquid and gas. What can you say about their shapes and the volume they occupy? • Solids have an orderly arrangement of particles with a definite shape and a definite volume • Liquids have a less orderly arrangement of particles with no definite shape and definite volume • Gases have a random arrangement of particles with no definite shape and no definite volume.

  7. Chapter 2 (continued) • Identify physical properties from chemical properties. For example zinc metal is hard, silver in color, is easily shaped and reacts vigorously with hydrochloric acid. What are some physical and chemical properties of zinc? • Physical • Metal, hard, silver, malleable • Chemical • Reacts with hydrochloric acid

  8. Chapter 2 (continued) • What is the difference between a physical and chemical change? How do you know when one occurred? Think of examples. • A physical change is one in which the substance still retains its identity • Evaporating, breaking, bending, cutting, etc… • A chemical change is one in which its chemical identify has changed • Rusting, burning, explosion, decaying, etc…

  9. Chapter 2 (continued) • What are symbols? What are formulas? When should you capitalize letters? What are subscripts? • Symbols represent elements • Formulas represent compounds • The first figure in a properly written chemical symbol always is capitalized.

  10. Chapter 2 (continued) • An object has a volume of 12mL and a mass of 3.5g. What is the object’s density? • 0.29 g/mL • An object has a mass of 357g and volume of 75mL. What is the object’s density? • 4.8 g/mL • An object has a mass of 4.2g and a volume of 19.3mL. What is the object’s density? • 0.22 g/mL

  11. Chapter 3 • What is an advantage of the metric system compared to the English system? • Based on a power of 10 • What is the SI unit of mass? What is the SI unit for temperature? • kilogram – mass • Kelvin - temperature • Define weight. Does weight change based on location? • the pull on a given mass by gravity • the weight of an object depends on its location • Which temperature scale has no negative numbers? • Kelvin

  12. Chapter 3 (continued) Convert the following: 50oC = ___323_______K -25oC = ___248___K Count the number of significant figures in these problems 50,400 mg = _____3_______ 0.00046 kg = _____2_______ 923.110 mg = _____6______ 1.0045 L = _______5_______

  13. Chapter 3 (continued) Round problems 1-4 to two significant figures and write the answer in scientific notation. 50,400 mg = ____5.0 X 104 mg________ 0.00046 kg = ____4.6 X 10-4 kg_______ 923.110 mg = ____9.2 X 102 mg______ 1,045 L = _______1.0 X 103 L______

  14. Chapter 3 (continued) Review rules regarding adding/subtracting and multiplying/dividing • What is the result of adding 2.01g +4.5g = 6.5 g • What is result of subtracting 4.356m – 3.6m = 0.8 m • What is the result of multiplying 5.2 X 102 by 1.367 X 10-4m = 7.1 X 10-2 m • What is the result of dividing 4.3 X102 by 2.0 X 104g = 2.2 X 10-2 g

  15. Chapter 3 (continued) • Define density. • Ratio of an object’s mass to its volume • Density is found by dividing mass by volume • volume = 90.0 mL density = 0.70g/mL mass = 63 g • mass = 55 g density = 12.3 g/mL volume = 4.5 mL • mass = 7.0 g volume = 9.1 cm3 density = 0.77 g/cm3 • What is the density of an object that has a mass of 5.6g and a volume of 25cm3? • 0.22 g/cm3

  16. Chapter 4 Fill in the table below:

  17. Chapter 4 (continued) • What does the atomic number of an element indicate? • # of protons • What does the mass number of an element indicate? • number of protons + neutrons • If an atom has an atomic number of 36 and a mass number of 84 how many protons, neutrons and electrons are present in this atom? • 36 Protons, 36 electrons, and 48 neutrons

  18. Chapter 4 (continued) • What does the number 13 represent in carbon-13? • Mass Number (number of protons and neutrons) • How do you calculate the number of neutrons in an atom? Mass Number - # of Protons = # of Neutrons • Who was Democritus? • One of the first to suggest the idea of atoms (460 -370 B.C.) • What is the unit that is used to measure a weighted average atomic mass? • Amu (atomic mass unit)

  19. Chapter 4 (continued) • What is an isotope? • Isotopes of the same element have different numbers of neutrons • What is an ion? • Ions are atoms that have number of protons not equal to electrons • How many protons, neutrons and electrons are present in this neutral atom? 3919K • 19 protons, 19 electrons, and 20 neutrons

  20. Chapter 5 • Describe the quantum mechanical model of an atom? • Based on the probability of finding an electron • Which scientist developed the quantum mechanical model of an atom? • Erwin Schrödinger • What are the shapes of an s and p orbitals? • s is a spherical shape • p is a dumbbell shape

  21. Chapter 5 (continued) • What is a principal energy level, sublevel and atomic orbital? • Principal energy level contains one or more types of sublevels • Sublevels are atomic orbital shape represented by the following letter s, p, d and f • Atomic orbitals are the regions of space where electrons are located in the different sublevels

  22. Atomic Orbitals P sublevel S sublevel 1s 2s 2p Atomic Orbitals Principal Energy Level 2 Principal Energy Level 1

  23. Chapter 5 (continued) • What is the maximum number in each s, p, d and f orbitals? • S (1 orbital), p (3 orbitals), d (5 orbitals), f (7 orbitals) • What types of atomic orbitals are in the 1st, 2nd and 3rd principal energy levels? • 1st (s atomic orbitals) • 2nd (s and p) • 3rd (s, p, and d)

  24. Chapter 5 (continued) • If the spin of one electron is clockwise in an orbital the spin on the second electron must be counterclockwise? • Using the Aufbau diagram what orbital would come after 3p? • 4s • What is the number of electrons in the outermost energy level of sulfur? • 6

  25. Chapter 5 (continued) • What happens when an electron moves from a lower to higher energy level? • absorb a quantum of energy • When an electron moves from a higher to lower energy level what happens? • lose a quantum of energy • When does emission of light occur? • Drops from a higher to lower energy level

  26. Chapter 5 (continued) • Which of the following quantum leaps would be associated with the greatest energy emitted? • n = 4 to n=2 • n = 5 to n= 1

  27. Chapter 5 (continued) • How are frequency and wavelength of light related? • They are inversely proportional to each other • Which variable is directly proportional to energy? • frequency • Define a photon. • Quanta of light

  28. Chapter 5 (continued) • What is the wavelength of a radiation with a frequency of 5 X 1018 Hz (s-1)? (c = 3.00 X 108 m/s) 6 X 10-11 m • What is the frequency of light having a wavelength of 4.2 X 10-8m? 7.1 X 1015 m • What is the energy of a photon with a frequency of 2.5 X 1015 Hz? [E = h v] (h = 6.63 X 10-34 J s) 1.7 X 10-18 J

  29. Chapter 5 (continued) • What are the electron configurations for the following elements: • Na: 1s2 2s2 2p6 3s1 • Ar: 1s2 2s2 2p6 3s2 3p6 • Zn: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 • As: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3 • Br: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 • Sr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2

  30. Chapter 6 • Another name for the representative elements is Group A elements. Where are these elements located on the periodic table? • In the s and p sublevels • Who was the first scientist to arrange the elements according to similar chemical and physical properties in order of increasing atomic mass? • Dmitri Mendeleev

  31. Chapter 6 (continued) • What is characteristic of the electron configurations of noble gases? • The highest occupied s and p sublevels are completely filled • Which subatomic particle plays the greatest role in determining the properties of an element? • electron

  32. Chapter 6 (continued) • What is the periodic trend for: • Atomic radius • Increase Top to Bottom and Right to Left • Ionic radius • Increase Top to Bottom and Right to Left for both Cations and Anions • Ionization energy • Increase Bottom to Top and Left to Right • Electronegativity • Increase Bottom to Top and Left to Right

  33. Chapter 6 (continued) • What is the energy required to remove an electron from an atom in the gaseous state called? • Ionization Energy • Which element would have the lowest first ionization energy? • Na, Mg, Al, Si, S, or Cl • Which element has the largest atomic radius? H, Li, Na, K, Rb, Cs or Fr Li, Be, B, C, N, O, or F

  34. Chapter 6 (continued) • Which element has the highest electronegativity value on the periodic table? • Fluorine • Define ion, cation, and anion. • Ion: charged atom • Cation: positively charged atom • Anion: negatively charged atom

  35. Chapter 6 (continued) • What are the ionic charges for the representative elements? +1, +2, +3, +/- 4, -3, -2, -1 • Write the charges for each one of these ions as a superscript Li +1 Ca+2 O-2 Al+3 S-2 Br-1 Na+1 Rb+1

  36. Chapter 7 • How many valence electrons would these atoms have? • Li: 1 valence electron • N: 5 valence electrons • O: 6 valence electrons • Cl: 7 valence electrons

  37. Chapter 7 (continued) • What is an ionic compound? • A compound composed of cations and anions • Held together by ionic bonds • What are some characteristics of ionic compounds? • Can conduct electric current when dissolved or in molten state • Has a high melting point • Crystalline solid • Brittle

  38. Chapter 7 (continued) • What is the electron configurations of the following ions: • N3- 1s2 2s2 2p6 • Al3+ 1s2 2s2 2p6 • Cl1- 1s2 2s2 2p6 3s2 3p6 • Ca2+ 1s2 2s2 2p6 3s2 3p6

  39. Chapter 7 (continued) • What is a metallic bond? • The attraction of metal ions to mobile electrons • What are some properties of metals? • Conductive (heat and electricity) • Malleable • Ductile • Luster • Why do metals have these properties • They have mobile valence electrons

  40. Chapter 8 • In a single covalent bond, how would two atoms still achieve a stable noble-gas electron configuration? • Two atoms share two electrons • How many electrons are required to create a single, double and triple covalent bond? • 2 shared electrons makes a single covalent bond • 4 shared electrons makes a double covalent bond • 6 shared electrons makes a triple covalent bond

  41. Chapter 8 (continued) • What is an unshared pair of electrons • Pair of valence electrons that does not participate in bonding • According to VSEPR theory, why do molecules adjust their shapes? • to keep pairs of valence electrons as far apart as possible

  42. Chapter 8 (continued) • What is a polar covalent bond? What is a non-polar covalent bond? How do you determine which type of bond would form? • Polar Covalent bonds are the result of uneven sharing of electrons • Based on the range of electronegativity values for the two atoms • Those atoms closest to Fluorine will acquire the most polar charge • Non-polar covalent bonds are the result of an even sharing of electrons

  43. Chapter 8 ( continued) • Which type of solid has the strongest bond structure and therefore the highest melting point? • Ionic Solids • Metallic Solids • Network Solids

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