PhyChem LU 1.1

LEARNING UNIT 1: COMPONENTS OF MATTER

1.1 INTRODUCTION TO MATTER

  • Definition of Matter:

    • Anything that occupies space, possesses mass, and has volume.

  • Classification of Matter:

    • Matter is classified into three states: solids, liquids, and gases.

1.2 STATES OF MATTER

  • Three States of Matter:

    • Solids:

    • Have fixed shape and volume.

    • Particles are closely packed in a relatively fixed position.

    • Liquids:

    • Have fixed volume, but not a fixed shape.

    • Particles are in close proximity but can move freely.

    • Gases:

    • Have neither fixed volume nor shape.

    • Particles move randomly and freely in three dimensions.

1.3 PHYSICAL PROPERTIES OF MATTER

  • Characteristics:

    • Can be observed or measured without altering the composition of matter.

  • Examples:

    • Odor, color, density, specific gravity, melting point, boiling point.

  • Physical Changes:

    • Changes that affect the appearance but not the composition.

    • Example of Physical Change: Ice melting.

1.4 CHEMICAL PROPERTIES OF MATTER

  • Definition:

    • Related to the capacity or tendency of a substance to undergo chemical reactions.

  • Examples:

    • Rusting of iron (reaction with oxygen)

    • Formation of water from the reaction between hydrogen and oxygen.

1.5 COMPOSITION OF MATTER

  • Molecules: Can consist of more than one type of element.

    • Element:

    • Molecules with only one type of atom.

    • Compound:

    • Molecules formed from more than one type of atom that are chemically bonded.

    • Mixture:

    • Contains more than one atom, element, or compound mixed physically, not chemically bonded.

1.6 TYPES OF MIXTURES

  • Heterogeneous Mixtures:

    • Composition is not uniform throughout; parts are visible and vary in texture and appearance.

    • Examples: Cement, rocks, oil-water mixture.

  • Homogeneous Mixtures:

    • Composition is uniform throughout; parts are not visible.

    • Examples: Soft drinks, milk, salt water, air.

SUMMARY: COMPOSITION OF MATTER

  • If homogeneous matter can be separated by physical means, it is a mixture.

  • If it cannot, it is a pure substance.

  • If a pure substance can be decomposed, it is a compound.

  • Methods of Separation:

    • Physical Separation Methods: Evaporation, filtration, etc.

    • Chemical Separation Methods: Isolation, electrolysis, etc.

1.2 ELEMENT AND COMPOUND

  • Element:

    • A fundamental (pure) substance that cannot be chemically transformed into simpler substances.

    • Current Knowable Elements: 118 elements as known in the periodic table.

    • Chemical Symbol: Composed of one or two letters (first letter capitalized).

    • Examples include H (Hydrogen), He (Helium).

PERIODIC TABLE OF ELEMENTS

  • Represents organization of elements by increasing atomic number.

  • Contains various groups such as noble gases, nonmetals, transition metals, and alkali metals.

COMPONDS

  • Definition:

    • Compounds consist of two or more elements chemically bonded in defined proportions and can be broken down into simpler substances.

  • Examples:

    • Water (H₂O), Glucose (C₆H₁₂O₆).

  • Processes involving Compounds:

    • Example A (Combustion): Hydrogen gas burns in oxygen producing water.

    • Example B (Electrolysis): Water can be decomposed into hydrogen and oxygen by electrical current.

1.3 UNIT OF MEASUREMENT

  • SI Units (Systeme Internationale):

    • The modern metric system accepted globally.

SI BASE UNITS

  • Defined measurements of various quantities in terms of standard units.

TEMPERATURE SCALES

  • Kelvin & Celsius Relation Formula:

    • K=C+273.15K = C + 273.15

  • Fahrenheit & Celsius Relation Formula:

    • F=(Cimesrac95)+32F = (C imes rac{9}{5}) + 32

PREFIXES IN SI UNITS

  • SI base units can be adjusted using prefixes to convey larger or smaller measurements.

DERIVED SI UNITS

  • Examples:

    • Volume:

      • V=LimesBimesHV = L imes B imes H

      • SI unit for volume is m³ (cubic meters).

      • Relation: 1 mL = 1 cm³ = 1 × 10⁻⁶ m³.

      • 1 L = 1000 mL = 1 dm³.

DENSITY

  • Definition:

    • Measurement of mass per unit volume.

  • Derived SI Units: g/cm³ or kg/m³

1.4 UNCERTAINTY IN MEASUREMENT

  • Numbers in scientific work fall into two categories:

    • Exact Numbers:

    • Values known precisely (e.g., 12 eggs in a dozen).

    • Has zero uncertainty.

    • Inexact Numbers:

    • Values with some uncertainty (e.g., measured values subject to error).

ACCURACY AND PRECISION

  • Accuracy:

    • Measurements close to the true value.

  • Precision:

    • Measurements that are close to each other.

SIGNIFICANT FIGURES

  • Guidelines for Determining Significant Figures:

    • Rule 1: Count all digits from left to right starting with the first non-zero digit.

    • Rule 2: For addition/subtraction, the number of decimal places in the result should match the least in the inputs.

    • Rule 3: For multiplication/division, the number of significant figures in the output should match that of the input with the fewest.

    • Rule 4: Rounding occurs such that the last digit is retained and incremented only if the next number is 5 or greater.

1.5 STOICHIOMETRIC

  • Dalton’s Atomic Theory:

    • All elements are made of extremely small particles called atoms.

    • All atoms of a given element are identical (same size, mass, chemical properties).

    • Atoms of one element cannot transform into another by chemical reactions, neither created nor destroyed.

    • Compounds form from combinations of more than one element.

    • The relative number of atoms in compounds remains consistent.

LAWS CONNECTED TO CHEMICAL REACTIONS

  • Law of Constant Composition:

    • In a compound, relative numbers and kinds of atoms are consistent; samples must have the same mass composition.

    • Exception: Isotopes have different mass numbers (e.g., Na⁻³⁵Cl₁⁷, Na⁻³⁷Cl₁⁷).

  • Law of Conservation of Mass:

    • Total mass pre- and post-reaction remains constant. Matter cannot be created or destroyed.

  • Law of Multiple Proportions:

    • For compounds, the mass of one element that combines with a fixed mass of another element is in a ratio of small whole numbers.

    • Example:

    • 8.0 g of oxygen combines with 1.0 g of hydrogen for H₂O.

    • 16.0 g of oxygen combines with 1.0 g of hydrogen for H₂O₂ (Ratio: 2:1).

MODERN VIEW OF ATOMIC STRUCTURE

  • Structure of Atom:

    • Consists of protons (positively charged), electrons (negatively charged), and neutrons (neutral).

    • Protons and neutrons exist in the nucleus, which holds most of the atom's mass.

    • Electrons orbit the nucleus, defining the atom's volume.

SUBATOMIC PARTICLES OF ATOMS

  • Atomic Definitions:

    • Atomic Number (Z):

    • Number of protons in the nucleus.

    • Mass Number (A):

    • Number of protons + number of neutrons,

    • (A=Z+N)(A = Z + N)

  • Isotopes:

    • Atoms of an element with different neutron counts.

PERCENT ISOTOPE ABUNDANCE & AVERAGE ATOMIC MASS

  • Calculation:

    • ext{Atomic Mass} = rac{ ext{No. of atoms of isotope}}{ ext{Total no. of atoms of all isotopes}} imes 100 ext{%}

    • ext{Average Atomic Mass} = rac{ ext{(% Isotope 1)}}{100} imes ext{mass of isotope 1} + rac{ ext{(% Isotope 2)}}{100} imes ext{mass of isotope 2} + …

EXAMPLES OF AVERAGE ATOMIC MASS CALCULATIONS

  • Example:

    • Boron Isotopes: 10B (19.91%), 11B (80.09%).

    • extAverageAtomicMass=(0.1991)imes10.0+(0.8009)imes11.0=10.81extamuext{Average Atomic Mass} = (0.1991) imes 10.0 + (0.8009) imes 11.0 = 10.81 ext{ amu}

    • Bromine Isotopes: 78.9Br (50.69%), 80.9Br (49.31%).

    • extAverageAtomicMass=(0.5069)imes78.9+(0.4931)imes80.9=79.70extamuext{Average Atomic Mass} = (0.5069) imes 78.9 + (0.4931) imes 80.9 = 79.70 ext{ amu}