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Chemistry Test Study Guide: All Flashcards
Here are all 87 flashcards, numbered in one continuous list, with each question followed by its answer. You can copy these into Quizlet or use them to study directly.
Section 1: Matter, atoms, molecules, and chemical formulas (1–13)
1. What is matter? Anything that has mass and takes up space.
2. What is an atom? The smallest unit of an element that retains its chemical identity.
3. What is an element? A pure substance made of only one type of atom.
4. What is a molecule? Two or more atoms chemically bonded together.
5. What is a chemical bond? A force of attraction that holds atoms together.
6. What is a chemical symbol? A one- or two-letter abbreviation for an element. Examples: O = oxygen, Na = sodium.
7. What is a chemical formula? A combination of element symbols and subscripts showing which elements and how many atoms are present.
8. What does a subscript tell you in a chemical formula? The number of atoms of the element immediately before it. Example: H₂O contains 2 hydrogen atoms.
9. How do you recognize a formula containing only one element? It contains only one type of element symbol. Example: O₂ contains only oxygen.
10. How do you recognize a formula containing different elements? It contains two or more different element symbols. Example: H₂O contains hydrogen and oxygen.
11. How many atoms are in CO₂? 3 total atoms: 1 carbon and 2 oxygen.
12. How many atoms are in 2H₂O? 6 total atoms: 4 hydrogen and 2 oxygen. The coefficient 2 multiplies the entire formula.
13. What is the difference between a coefficient and a subscript? A coefficient multiplies the entire formula; a subscript counts atoms of the element before it.
Section 2: Solids, liquids, gases, and ice (14–27)
14. How are particles arranged in a solid? Very close together in an organized arrangement.
15. How do particles move in a solid? They vibrate in place.
16. What are the attractive forces between particles in a solid? They are generally strong enough to hold particles close together in fixed positions.
17. How are particles arranged in a liquid? They are close together but not in fixed positions.
18. How do particles move in a liquid? They slide past one another.
19. How are particles arranged in a gas? They are far apart, with lots of empty space between them.
20. How do particles move in a gas? They move freely and rapidly in random directions.
21. How strong are attractive forces between gas particles? Usually weak compared with the particles' motion, allowing them to spread apart.
22. Which state of matter has a definite shape and definite volume? A solid.
23. Which state has a definite volume but takes the shape of its container? A liquid.
24. Which state takes both the shape and volume of its container? A gas.
25. Why does ice have an unusual particle arrangement? When water freezes, hydrogen bonds hold water molecules in an open, organized structure.
26. Why is ice less dense than liquid water? Its open structure spaces the molecules farther apart, so the same mass occupies more volume.
27. What happens to the particle arrangement when ice melts? Some of the open structure collapses, allowing molecules to pack closer together.
Section 3: Physical properties and gas particles (28–39)
28. What is an observable property? A property you can notice through your senses or observation, such as color or appearance.
29. What is a measurable property? A property that can be measured numerically using a tool, such as density or melting point.
30. What is appearance? How a substance looks, including its color, shape, and texture.
31. What is conductivity? The ability of a material to conduct heat or electricity.
32. What is bendability? The ability of a material to bend without breaking.
33. What is density? The amount of mass in a given volume. Formula: density = mass ÷ volume.
34. What is melting point? The temperature at which a solid changes into a liquid.
35. What is boiling point? The temperature at which a liquid boils and changes into a gas.
36. What is a monatomic gas? A gas whose particles are individual atoms. Example: helium (He).
37. What is a diatomic gas? A gas whose particles are molecules containing two bonded atoms. Examples: H₂, N₂, and O₂.
38. What is the difference between monatomic and diatomic gases? Monatomic gas particles contain one atom each, while diatomic gas particles contain two bonded atoms each.
39. Why can a gas be compressed? Gas particles have large spaces between them that can be reduced.
Section 4: Atomic structure and the periodic table (40–52)
40. Where are protons located, and what is their charge? In the nucleus; positive charge (+1).
41. Where are neutrons located, and what is their charge? In the nucleus; neutral charge (0).
42. Where are electrons located, and what is their charge? In the electron cloud outside the nucleus; negative charge (−1).
43. What is the relative mass of a proton? Approximately 1 atomic mass unit (amu).
44. What is the relative mass of a neutron? Approximately 1 amu.
45. What is the relative mass of an electron? Approximately 1/1836 of a proton's mass, often treated as 0 amu in basic calculations.
46. What is the atomic number? The number of protons in an atom's nucleus.
47. What is the mass number? The total number of protons and neutrons in one atom.
48. How do you calculate the number of neutrons? Subtract the atomic number from the mass number.
Formula: Neutrons = mass number − atomic number.
49. How many electrons does a neutral atom have? The same number as its protons.
50. An atom has 11 protons and a mass number of 23. How many neutrons does it have? 23 − 11 = 12 neutrons.
51. A neutral atom has an atomic number of 8. How many protons and electrons does it have? 8 protons and 8 electrons.
52. What determines which element an atom is? Its number of protons. Changing the number of protons changes the element.
Section 5: Scientists and the atomic model (53–60)
53. What did John Dalton contribute to the atomic model? He proposed that matter is made of atoms and pictured atoms as solid spheres. He also proposed that different elements have different types of atoms.
54. What did J. J. Thomson discover? He discovered the electron through cathode ray experiments.
55. What was Thomson's atomic model? The plum pudding model, which pictured negatively charged electrons embedded in a positively charged sphere.
56. What did Ernest Rutherford discover? His gold foil experiment showed that atoms have a small, dense, positively charged nucleus and are mostly empty space.
57. What happened in Rutherford's gold foil experiment? Most alpha particles passed straight through the foil, while a few were deflected sharply or bounced back.
58. What did most particles passing straight through the gold foil demonstrate? Atoms are mostly empty space.
59. What did the few particles that deflected sharply demonstrate? The atom's positive charge and most of its mass are concentrated in a small, dense nucleus.
60. How did the atomic model change from Dalton to Thomson to Rutherford?
Dalton: atoms are solid spheres.
Thomson: electrons are embedded in positive material.
Rutherford: atoms have a small, dense nucleus and are mostly empty space.
Section 6: Isotopes and average atomic mass (61–73)
61. What are isotopes? Atoms of the same element that have different numbers of neutrons.
62. How are isotopes of the same element alike? They have the same number of protons and the same atomic number.
63. How do isotopes of the same element differ? They have different numbers of neutrons and different mass numbers.
64. Why is the atomic mass on the periodic table usually not a whole number? It is a weighted average of the masses of the element's naturally occurring isotopes.
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65. What is fractional abundance? The fraction of all atoms in a sample that are a particular isotope.
66. How do you calculate fractional abundance from counts? Divide the number of atoms of that isotope by the total number of atoms.
Formula: Fractional abundance = isotope count ÷ total count.
67. How do you convert a percentage to a decimal? Divide by 100. Example: 25% = 0.25.
68. How do you convert a decimal to a percentage? Multiply by 100. Example: 0.25 = 25%.
69. How do you calculate weighted average atomic mass? Multiply each isotope's mass by its decimal abundance, then add all the contributions.
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70. An isotope has 20 atoms out of 100 total atoms. What is its fractional abundance? 20 ÷ 100 = 0.20, or 20%.
71. An isotope has a mass of 10 amu and an abundance of 0.20. What is its contribution to average atomic mass? 10 × 0.20 = 2.0 amu.
72. Isotope A has a mass of 10 amu and 20% abundance. Isotope B has a mass of 11 amu and 80% abundance. What is the average atomic mass? (10 × 0.20) + (11 × 0.80) = 2.0 + 8.8 = 10.8 amu.
73. What should all isotope decimal abundances add up to? 1.00, or 100% when expressed as percentages.
Section 7: Kinetic energy, potential energy, and thermal energy (74–87)
74. What is kinetic energy? Energy an object has because of its motion.
75. What is potential energy? Stored energy due to an object's position, arrangement, or condition.
76. What is chemical potential energy? Energy stored in chemical substances because of their composition and arrangement. It can be released or absorbed during chemical reactions.
77. What is the difference between kinetic energy and potential energy? Kinetic energy is energy of motion, while potential energy is stored energy.
78. What is thermal energy? Energy associated with the random motion and interactions of particles in a substance.
79. How does temperature relate to particle motion? Higher temperature generally means particles have greater average kinetic energy.
80. What happens to particle motion when a substance is heated? Its particles generally move faster or vibrate more.
81. What is a system? The object or group of objects being studied.
82. What are the surroundings? Everything outside the system that can interact with it.
83. In which direction does thermal energy transfer between objects at different temperatures? From the warmer object to the cooler object.
84. What happens when two objects reach thermal equilibrium? They reach the same temperature, so there is no net thermal energy transfer between them.
85. A hot spoon is placed in cold water. Which way does thermal energy transfer? From the hot spoon to the colder water.
86. A cold ice cube is placed in warm juice. Which way does thermal energy transfer? From the warmer juice to the colder ice cube.
87. Does an object always contain more thermal energy just because it has a higher temperature? No. Total thermal energy also depends on the amount of substance and the particles and material involved