Chemistry Review Flashcards

Chapter 1

• Observation vs. Hypothesis:
  • An observation is a qualitative or quantitative measurement of an event, a statement of fact without explanation.
  • A hypothesis is an educated guess explaining the observation; it can be true, untrue, or modified.
• Hypothesis to Theory or Law:
  • A hypothesis, when proven true on multiple levels, can become a theory or a law.
• Scientific Theory vs. Scientific Law:
  • A scientific theory explains why an observation occurs, based on principles and equations.
  • A scientific law explains what and how an observation occurs, based on observations and tangible means.
• Examples:
  • The leaves turned brown. (Observation)
  • The water boils because it got hot. (Hypothesis)
  • The boiling point of water is 212 °F. (Observation)
  • The Earth is flat because the map I have shows the Earth is flat. (Hypothesis)

Chapter 2

• Units of Measurement:
  • Liter (L) = volume
  • Meter (m) = length / distance
  • Gram (g) = mass
  • Second (s) = time
• Significant Figures:
  • 1.254 has 4 SF
  • 0.14580 has 5 SF
  • 0.008274 has 4 SF
  • 255000 has 3 SF
  • 5.24 \times 10^6 has 3 SF
• Calculations with Significant Figures and Units:
  • 12.0 - 1.587 + 150 = 160
  • (1087 \times 1.2) / 10.5 = 120
  • 1.254 \times 12.7 \times 0.0089 = 0.14
  • 124000 / (58.4 - 2.5) = 2200
• Unit Conversions:
  • 2500 cg to L (density = 2.00 g/mL) = 2500 \text{ cg} \times \frac{1 \text{ g}}{100 \text{ cg}} \times \frac{1 \text{ mL}}{2.00 \text{ g}} \times \frac{1 \text{ L}}{1000 \text{ mL}} = 0.013 \text{ L}
  • 100. mL to L = 100. \text{ mL} \times \frac{1 \text{ L}}{1000 \text{ mL}} = 0.100 \text{ L}
  • 0.0254 kg to mg = 0.0254 \text{ kg} \times \frac{1000 \text{ g}}{1 \text{ kg}} \times \frac{1000 \text{ mg}}{1 \text{ g}} = 25400 \text{ mg}
  • 35 cg to g = 35 \text{ cg} \times \frac{1 \text{ g}}{100 \text{ cg}} = 0.35 \text{ g}
  • 1. 05 g to mL (density = 2.15 g/mL) = 1.05 \text{ g} \times \frac{1 \text{ mL}}{2.15 \text{ g}} = 0.488 \text{ mL}
  • 80. mph to m/s = 80 \frac{\text{miles}}{\text{hr}} \times \frac{5280 \text{ ft}}{1 \text{ mile}} \times \frac{12 \text{ in}}{1 \text{ ft}} \times \frac{2.54 \text{ cm}}{1 \text{ in}} \times \frac{1 \text{ m}}{100 \text{ cm}} \times \frac{1 \text{ hr}}{60 \text{ min}} \times \frac{1 \text{ min}}{60 \text{ sec}} = 36 \text{ m/s}
• Sorting Numbers:
  • Ascending order: 2.0 \times 10^{-5} < 2.0 \times 10^{0} < 2.0 \times 10^{5}

Chapter 3

• Properties of Matter:
  • Matter has mass and takes up space.
• States of Matter:
  • Solid: Definite shape and definite volume.
  • Liquid: Definite volume and indefinite shape.
  • Gas: Indefinite volume and indefinite shape.
• Physical vs. Chemical Changes:
  • Melting butter: Physical
  • Boiling water: Physical
  • Burning wood: Chemical
  • Rusting steel: Chemical
  • Digesting food: Chemical
  • Tearing paper: Physical
• Element, Compound, or Mixture:
  • Water: Compound
  • Pure Air: Mixture
  • Iron bar: Element
  • Sugar: Compound
  • Hydrogen Gas: Element
  • Ocean Water: Mixture
• Calorie Calculation:
  • Calories needed to heat 1600 g of water from 25 °C to 100 °C:
  • q = m \times C \times \Delta T = (1600 \text{ g}) \times (1.00 \frac{\text{cal}}{\text{g} \cdot {}^{\circ}\text{C}}) \times (100 {}^{\circ}\text{C} - 25 {}^{\circ}\text{C}) = 120,000 \text{ cal} = 1.2 \times 10^5 \text{ cal}
• Calculating Mass of Copper Metal:
  • Grams of copper metal (C = 0.385 J/g°C) heated from 16 °C to 27 °C with 275 J applied:
  • m = \frac{q}{C \times \Delta T} = \frac{275 \text{ J}}{(0.385 \frac{\text{J}}{\text{g} \cdot {}^{\circ}\text{C}}) \times (11 {}^{\circ}\text{C})} = 66 \text{ g}
• Temperature Change Calculation:
  • Temperature change when 1.5 kg of aluminum is heated with 18 J (C = 0.899 J/g°C):
  • \Delta T = \frac{q}{m \times C} = \frac{1800 \text{ J}}{(1500 \text{ g}) \times (0.899 \frac{\text{J}}{\text{g} \cdot {}^{\circ}\text{C}})} = 1.3 {}^{\circ}\text{C}
• Temperature Conversions to Fahrenheit:
  • 0°C to F: F = (0 {}^{\circ}\text{C} \times 1.8) + 32 = 32 {}^{\circ}\text{F}
  • 0 K to F:
    • C = 0 \text{ K} - 273 = -273 {}^{\circ}\text{C}
    • F = (-273 {}^{\circ}\text{C} \times 1.8) + 32 = -460 {}^{\circ}\text{F}
  • 212 K to F:
    • C = 212 \text{ K} - 273 = -61 {}^{\circ}\text{C}
    • F = (-61 {}^{\circ}\text{C} \times 1.8) + 32 = -78 {}^{\circ}\text{F}
  • 23 °C to F: F = (23 {}^{\circ}\text{C} \times 1.8) + 32 = 73 {}^{\circ}\text{F}
• Phase Changes:
  • Melting: Endothermic
  • Evaporation: Endothermic
  • Freezing: Exothermic
  • Deposition: Exothermic
  • Sublimation: Endothermic
  • Condensation: Exothermic

Chapter 4

• Types of Elements:
  • H: Nonmetal
  • Mg: Metal
  • As: Metalloid
  • P: Nonmetal
  • Ge: Metalloid
  • Fe: Metal
• Subatomic Particles:
  • Proton: Has mass and positive charge.
  • Neutron: Has mass and no charge.
  • Electron: Has "no" mass and negative charge.
• Location of Subatomic Particles:
  • Protons and neutrons are in the nucleus.
  • Electrons orbit in the electron cloud.
• Isotopes:
  • ¹⁸O: p=8, n=10, e=8
  • ¹³C: p=6, n=7, e=6
  • ¹⁰⁰Mo: p=42, n=58, e=42
  • Selenium-76: p=34, n=42, e=34
  • Calcium-42: p=20, n=22, e=20
  • Iron-55: p=26, n=29, e=26
• Determining Isotopes:
  • 16 protons, 28 neutrons, 16 electrons: ^{44}S
  • 28 protons, 36 neutrons, 28 electrons: ^{64}Ni
  • 98 protons, 142 neutrons, 98 electrons: ^{240}Cf
  • 1 proton, 1 neutron, 1 electron: ^2H
  • 22 protons, 45 neutrons, 22 electrons: ^{67}Ti
  • 48 protons, 73 neutrons, 48 electrons: ^{121}Cd
• Identifying Elements:
  • (The content about identifying the element by its period and group number is not present in the reference text.)