Chemistry Chapter 1 Study Guide PDF

CHEMISTRY CHAPTER 1 STUDY GUIDE Matter, Atoms, and Molecules 1. Matter: ○ Anything that has mass and occupies space. ○ Exists in three primary states: solid, liquid, and gas. 2. Atoms: ○ The smallest unit of an element, made up of subatomic particles: Protons: Positive char...

CHEMISTRY CHAPTER 1 STUDY GUIDE Matter, Atoms, and Molecules 1. Matter: ○ Anything that has mass and occupies space. ○ Exists in three primary states: solid, liquid, and gas. 2. Atoms: ○ The smallest unit of an element, made up of subatomic particles: Protons: Positive charge, found in the nucleus. Neutrons: Neutral charge, found in the nucleus. Electrons: Negative charge, orbit around the nucleus. ○ The number of protons determines the element's identity (atomic number). 3. Molecules: ○ A group of two or more atoms bonded together, which can be the same (e.g., O ₂) or different (e.g., H₂O). ○ Molecules can be simple (like O₂) or complex (like proteins). ○ The bonds between atoms in molecules can be covalent (sharing electrons) or ionic (transferring electrons). Scientific Method The scientific method is a way to investigate and understand the world around us. It follows these steps: 1. Observe: Notice something interesting or a problem. 2. Ask a Question: Based on your observation, ask a clear question. 3. Make a Hypothesis: Guess an answer to your question that can be tested. 4. Test the Hypothesis: Perform an experiment to see if your guess is correct. 5. Collect Data: Measure and record what happens during the experiment. 6. Analyze: Look at the data to figure out what it means. 7. Conclude: Decide if your hypothesis is correct or if you need to change it. 8. Share Results: Tell others what you found. 9. Repeat: Do the experiment again or test a new idea based on what you learned. This method helps us get clear, reliable answers by following a step-by-step process. States of Matter States of Matter Phase Transitions 1. Solid ○ Tightly packed molecules 1. Melting: Solid → Liquid (gain heat) ○ Definite shape and volume 2. Freezing: Liquid → Solid (lose heat) 2. Liquid 3. Vaporization: Liquid → Gas (gain heat) ○ Molecules move past each other ○ Definite volume, takes shape of container 4. Condensation: Gas → Liquid (lose heat) 3. Gas 5. Sublimation: Solid → Gas (gain heat) ○ Molecules are spread out 6. Deposition: Gas → Solid (lose heat) ○ No definite shape or volume Types of Solids 4. Crystalline ○ Regular pattern of molecules ○ Definite shape and melting point 5. Amorphous ○ Random arrangement of molecules ○ No defined shape, melts over a range Elements, Compounds, and Mixtures Element Definition: A pure substance made up of only one type of atom. Example: Oxygen (O), Hydrogen (H) Compound Definition: A substance made of two or more different elements chemically bonded. Example: Water (H₂O), Carbon Dioxide (CO₂) Mixture Definition: A combination of two or more substances where each retains its properties. Types of Mixtures 1. Homogeneous Mixture ○ Uniform composition throughout. 2. Heterogeneous Mixture ○ Not uniform, different components can be seen. Properties Physical Properties Definition: Characteristics that can be observed or measured without changing the substance's chemical composition. Chemical Properties Definition: Characteristics that describe a substance's ability to undergo a chemical change. Intensive vs. Extensive Properties Intensive Properties Definition: Properties that do not depend on the amount of substance. Extensive Properties Definition: Properties that depend on the amount of substance present. Changes Physical Changes Definition: Changes in which the substance's composition remains the same, but its appearance or form may change. Examples: ○ Melting ice ○ Cutting paper Chemical Changes Definition: Changes where the substance undergoes a chemical reaction, resulting in a new substance with different properties. Examples: ○ Burning wood ○ Rusting of iron Key Differences Physical Change: No new substance formed. Chemical Change: New substance(s) formed, often with a change in color, temperature, or gas production. Energy Energy Types of Energy Definition: The ability to do work or cause change. Kinetic Energy (KE) Units: Joules (J) Definition: The energy of motion. Formula: Work KE=1/2mv2 Definition: Work is done when a force acts on an Potential Energy (PE) object and causes it to move. Formula: Definition: The stored energy of an object due to its Work = Force×Distance position or condition. Formula: Units: Joules (J) PE = mgh Key Concepts Thermal Energy Energy can transform: Kinetic energy can be Definition: The total energy of all the particles in an object converted to potential energy (and vice versa). due to their motion. Energy is conserved: In a closed system, energy is neither created nor destroyed, it only changes form (Law of Conservation of Energy). Units of Measurement SI Units (Metric System) Used worldwide in science and most countries. Length: meter (m) Mass: kilogram (kg) Time: second (s) Temperature: Kelvin (K) Amount of Substance: mole (mol) Derived SI Units Combinations of base units. Area: square meter (m²) Volume: cubic meter (m³) Speed/Velocity: meter per second (m/s) Force: Newton (N) = kg·m/s² Pressure: Pascal (Pa) = N/m² Energy/Work: Joule (J) = N·m Key Measurements and Units 1. Volume (V) 4. Density (ρ) Definition: The mass per unit volume of a Definition: The amount of space occupied by an substance. object or substance. Formula: Density(ρ) = Mass / Volume​ SI Unit: Cubic meter (m³), but commonly SI Unit: (kg/m³), (g/mL) (g/cm³). measured in liters (L). Conversions: 5. Time (t) ○ 1 m³ = 1,000 liters ○ 1 liter = 1,000 milliliters (mL) Definition: The duration in which events occur. SI Unit: Second (s) Other Units: 2. Mass (m) ○ Hour = 60 minutes = 3,600 seconds Definition: The amount of matter in an object. 6. Temperature (T) SI Unit: Kilogram (kg) Conversions: Definition: A measure of the average kinetic ○ 1 kilogram = 1,000 grams energy of particles in a substance. ○ 1 gram = 1,000 milligrams SI Unit: Kelvin (K) Common Scale: ○ Celsius (°C) = K - 273.15 or vise versa Metric Prefixes Tera (T): 1012 or 1,000,000,000,000 Giga (G): 109 or 1,000,000,000 Mega (M): 106 or 1,000,000 Kilo (k): 103 or 1,000 Hecto (h): 102 or 100 Deca (da): 101 or 10 Base (unit): 1 Deci (d): 10-1 or 0.1 Centi (c): 10-2 or 0.01 Milli (m): 10-3 or 0.001 Micro (µ): 10-6 or 0.000001 Nano (n): 10-9 or 0.000000001 Pico (p): 10-12 or 0.000000000001 Significant Figures (Sig Figs) Simplified What are Significant Figures? Significant figures show the precision of a measurement. They help us know how accurate a number is. Easy Rules to Remember: 1. Non-zero digits are always significant. ○ Example: 123 has 3 sig figs. 2. Zeros between nonzero digits are significant. ○ Example: 101 has 3 sig figs. 3. Leading zeros (zeros before the first non-zero digit) are NOT significant. ○ Example: 0.0045 has 2 sig figs. 4. Trailing zeros in a decimal are significant. ○ Example: 45.00 has 4 sig figs. 5. Trailing zeros in a whole number without a decimal are NOT significant. ○ Example: 1500 has 2 sig figs. Exact Numbers Simplified What Are Exact Numbers? Exact numbers are numbers that are counted or defined, and they have unlimited significant figures. These numbers are considered to have no uncertainty. Examples of Exact Numbers: 1. Counting objects ○ Example: 3 apples (exactly 3 apples, no approximation). 2. Defined quantities ○ Example: 1 dozen = 12 items (by definition). 3. Conversion factors ○ Example: 1 inch = 2.54 cm (exact by definition). 4. Pure constants in formulas ○ Example: Pi (π) is an exact value in calculations (though we typically use a rounded version like 3.14). Why They Don’t Affect Sig Figs: Exact numbers don’t limit the precision of your calculation. They are treated as having infinite significant figures, so they don’t change the number of sig figs in your result. Calculations with Significant Figures (Sig Figs) When performing calculations, it's important to follow rules for significant figures (sig figs) to ensure precision in your results. Here's how to handle sig figs in different types of calculations: 1. Addition and Subtraction Rule: The result should have the same number of decimal places as the least precise number in the calculation. 2. Multiplication and Division Rule: The result should have the same number of sig figs as the number with the least sig figs in the calculation. 3. Mixed Operations Rule: Do addition/subtraction first, rounding to the correct decimal places, then perform multiplication/division with the correct number of sig figs. Quick Summary: Addition/Subtraction: Round to the least number of decimal places. Multiplication/Division: Round to the least number of sig figs. Accuracy, Precision, and Experimental Error Accuracy Definition: How close your measurement is to the true or accepted value. Example: If the real temperature is 25°C, and your measurement is 24.8°C, your measurement is accurate. Precision Definition: How close your measurements are to each other, no matter if they are close to the true value. Example: If you measure the temperature five times and get 24.8°C, 24.7°C, 24.9°C, 24.8°C, and 24.7°C, your measurements are precise. Experimental Error Definition: The difference between the measured value and the true value. Types: ○ Systematic Error: Consistent, repeatable errors caused by a faulty instrument or method. ○ Random Error: Unpredictable variations that happen naturally and are difficult to control.. Dimensional Analysis What is Dimensional Analysis? Definition: A method used to convert between different units using conversion factors. Purpose: Ensure that equations are dimensionally consistent, and to convert measurements from one unit to another. Key Steps: 1. Identify Units: Know the starting and desired units. 2. Set Up Conversion Factor(s): Use ratios that express equivalent values (like 1 inch = 2.54 cm). 3. Multiply by Conversion Factor: Ensure that units cancel out properly. 4. Solve: perform the math, and make sure the units match what you need. Helpful Tip: Always double-check your conversion factors! The unit you want should be left over after cancellation. Why Use Dimensional Analysis? It's a powerful tool to avoid errors and ensure consistency in unit conversions!