AS

Chemistry Chapter 1 Reading Notes (copy)

1.1 Chemistry in Context

  • Chemistry: The study of matter and its properties, how matter changes, and how these changes are accompanied by energy changes.

  • Matter: Anything that has mass and occupies space.

    • Our universe is composed of matter and energy.

  • Scientific Method: A systematic approach to research.

    • Observation: Gathering qualitative (non-numerical) and quantitative (numerical) data.

    • Hypothesis: A tentative explanation for observations, testable through experimentation.

    • Experiment: Procedures designed to test the hypothesis; must be controlled and reproducible.

    • Law: A concise statement or mathematical equation that summarizes a vast number of experimental observations and describes a natural phenomenon (e.g., Law of Conservation of Mass).

    • Theory: A well-substantiated, comprehensive, and testable explanation of a particular aspect of nature (e.g., Atomic Theory).

1.2 Phases and Classification of Matter

  • States of Matter:

    • Solid: Definite shape and volume; particles closely packed in fixed positions, vibrating.

    • Liquid: Definite volume, indefinite shape; particles close but can move past each other.

    • Gas: Indefinite shape and volume; particles far apart and move randomly.

    • Plasma: Gaseous state containing significant numbers of electrically charged particles (ionized atoms or molecules).

  • Classification of Matter:

    • Pure Substances: Matter with a fixed composition and distinct properties.

    • Elements: Cannot be broken down into simpler substances by chemical means (e.g., Oxygen, Gold). Composed of only one type of atom.

    • Compounds: Two or more different elements chemically bonded together in fixed proportions (e.g., Water ( ext{H}2 ext{O}), Carbon Dioxide ( ext{CO}2)). Can be broken down into elements by chemical means.

    • Mixtures: Two or more substances physically combined, each retaining its own chemical identity.

    • Homogeneous Mixture (Solution): Uniform composition and properties throughout (e.g., salt water, air, alloys).

    • Heterogeneous Mixture: Non-uniform composition; components are visibly distinct (e.g., sand and water, oil and vinegar).

  • Atoms: The smallest particle of an element that retains the chemical identity of the element.

  • Molecules: Two or more atoms joined together by chemical bonds.

1.3 Physical and Chemical Properties

  • Physical Properties: Characteristics of matter that can be observed or measured without changing the substance's chemical identity.

    • Examples: Density, color, hardness, melting point, boiling point, electrical conductivity, state of matter.

    • Physical Change: A change in the state or properties of matter without altering its chemical composition (e.g., melting ice, boiling water, dissolving sugar).

  • Chemical Properties: Characteristics that describe a substance's ability to undergo a specific chemical change or reaction, forming new substances.

    • Examples: Flammability, reactivity with acids, toxicity, corrosion resistance.

    • Chemical Change (Chemical Reaction): A process that produces one or more new substances that differ from the original substances (e.g., burning wood, rusting iron, food digestion).

    • Evidence of chemical change: Color change, temperature change, gas production (bubbles), precipitate formation, emission of light.

  • Extensive Properties: Depend on the amount of matter present (e.g., mass, volume, heat).

  • Intensive Properties: Do not depend on the amount of matter present (e.g., temperature, density, melting point, boiling point).

1.4 Measurements

  • SI Units (International System of Units): The standard system of units used globally in science.

    • Base Units:

    • Length: meter (m)

    • Mass: kilogram (kg)

    • Time: second (s)

    • Temperature: Kelvin (K)

    • Amount of Substance: mole (mol)

    • Electric Current: ampere (A)

    • Luminous Intensity: candela (cd)

    • Prefixes: Used to denote multiples or fractions of base units (e.g., kilo ( imes 10^3), milli ( imes 10^{-3}), micro ( imes 10^{-6}), nano ( imes 10^{-9})).

  • Derived Units: Units obtained from combinations of base units (e.g., density in ext{kg/m}^3, volume in ext{m}^3).

    • Volume: Common derived unit. 1 ext{ L} = 1000 ext{ mL} = 1000 ext{ cm}^3 = 1 ext{ dm}^3.

    • Density (d): A physical property defined as mass per unit volume (d = m/V). Common units are ext{g/cm}^3 for solids/liquids and ext{g/L} for gases.

  • Temperature Scales:

    • Celsius (^ ext{ extdegree} ext{C}): Water freezes at 0^ ext{ extdegree} ext{C} and boils at 100^ ext{ extdegree} ext{C}.

    • Kelvin (K): Absolute temperature scale where 0 ext{ K} = -273.15^ ext{ extdegree} ext{C} (absolute zero).

    • Conversion: T{ ext{K}} = T{ ext{ extdegree} ext{C}} + 273.15

    • Fahrenheit (^ ext{ extdegree} ext{F}): Primarily used in the United States for weather, body temperature, etc.

    • Conversion: T{ ext{ extdegree} ext{C}} = (T{ ext{ extdegree} ext{F}} - 32^ ext{ extdegree}) / 1.8 or T{ ext{ extdegree} ext{F}} = 1.8 imes T{ ext{ extdegree} ext{C}} + 32^ ext{ extdegree}.

1.5 Measurement Uncertainty, Accuracy, and Precision

  • Uncertainty in Measurement: All measurements are subject to some degree of uncertainty due to limitations of measuring instruments and human error.

    • Recorded measurements should include all certain digits and one estimated uncertain digit.

  • Significant Figures (Sig Figs): The number of meaningful digits in a measured quantity.

    • Rules for Counting Significant Figures:

    1. Non-zero digits are always significant (e.g., 28.7 has 3 sig figs).

    2. Zeros between non-zero digits are significant (e.g., 1005 has 4 sig figs).

    3. Leading zeros (at the beginning of a number before the first non-zero digit) are not significant (e.g., 0.0025 has 2 sig figs).

    4. Trailing zeros (at the end of a number) are significant only if the number contains a decimal point (e.g., 12.00 has 4 sig figs; 1200 has 2 sig figs, unless specified by scientific notation or a decimal point at the end like 1200. which would have 4).

    • Exact Numbers: Numbers from definitions or direct counts have infinite significant figures.

  • Accuracy: How close a measurement is to the true or accepted value.

  • Precision: How closely multiple measurements agree with one another (reproducibility).

1.6 Mathematical Treatment of Measurement Results

  • Significant Figures in Calculations:

    • Addition and Subtraction: The result should have the same number of decimal places as the measurement with the fewest decimal places.

    • Multiplication and Division: The result should have the same number of significant figures as the measurement with the fewest significant figures.

  • Rounding Rules:

    1. If the first non-significant digit (the one to be dropped) is less than 5, simply drop it and all subsequent digits.

    2. If the first non-significant digit is 5 or greater, round up the last significant digit by adding one to it.

  • Dimensional Analysis (Factor-Label Method): A problem-solving technique that uses conversion factors to convert units.

    • Conversion Factor: A ratio of equivalent measurements that expresses the same quantity in different units (e.g., 1 ext{ m} = 100 ext{ cm}). The ratio (1 ext{ m} / 100 ext{ cm}) or (100 ext{ cm} / 1 ext{ m}) can be used.

    • Units are treated like algebraic quantities and can be cancelled out to ensure the final answer has the desired units.