General and Inorganic Chemistry: Introduction to Chemistry

Foundations of Chemistry and Scientific Method

Definition of Chemistry:
- Chemistry is the comprehensive study of chemicals, focusing on how physical and chemical properties depend on composition and structural organization.
- It investigates what substances are and how they undergo fundamental transformations.
- The term originates from El Khemid, meaning "the transformation," underscoring the discipline's primary focus on converting initial substances into distinct new materials.
- Historical applications evolved from early metal processing to the production of cosmetics, pharmaceuticals, ceramics, and glassmaking.
Early Chemical Investigations:
- Initial chemical practices focused on identifying, classifying, and describing distinct physical substances.
- As empirical knowledge accumulated, systematic curiosity led chemists to seek underlying patterns and theories that explain chemical behavior, rather than attempting to memorise every individual substance.
The Scientific Method:
- Observation: An objective gathering of data. An observation accompanied by a statement of certainty is established as a fact.
- Laws: Generalized statements or mathematical relationships that summarize a broad range of consistent observations.
- Hypothesis: A tentative explanation or predictive proposition formulated to explain observed phenomena.
- Experiment: A controlled procedure performed specifically to test the validity of a hypothesis.
- Theory: A thoroughly tested, refined, and validated hypothesis that explains why natural phenomena occur.
- Model: An integrated combination of multiple theories forming a comprehensive explanation of a wide variety of phenomena.
Nature of Scientific Knowledge:
- Science differs from art in that scientific knowledge requires consensus—first regarding empirical facts, and subsequently regarding theoretical explanations.
- Chemistry operates under a delicate balance between established empirical facts and theoretical interpretations.
- Scientific progress relies on provisional truth: hypotheses and theories are accepted as best-available explanations while maintaining rigorous scientific skepticism.
Serendipity in Science:
- Scientific results frequently occur by accident, captured by the adage: "Chance favors the prepared mind."
- When Wilhelm Conrad Röntgen observed the outline of his hand bones on a fluorescent screen near a cathode ray tube (leading to the discovery of X-rays), his response regarding what he thought was: "I did not think. I investigated."
Role as a Central Science:
- Chemistry serves as the central science interconnecting physics, biology, geology, environmental science, medicine, and engineering.

Classification and States of Matter

Definition and Physical States of Matter:
- Matter is defined as anything that occupies physical space and possesses mass.
- Mass represents the fundamental quantity of matter in an object, whereas weight is the variable gravitational force exerted on that mass.
- Matter exists in three principal physical states:
- Solid: Characterized by particles packed tightly together in a rigid, highly ordered structural arrangement.
- Liquid: Characterized by loosely packed particles that move past one another, maintaining a fixed volume while taking the shape of their container.
- Gas: Characterized by randomly packed particles separated by large relative distances that expand dynamically to fill the container.
Kinetic Molecular Theory (KMT):
- All matter is composed of extremely small particles that remain in continuous motion.
- Thermal energy directly correlates with particle movement: an increase in temperature increases the kinetic energy and speed of the constituent particles.
Levels of Observation in Chemistry:
- Macroscopic Level: Physical phenomena that can be directly observed and measured with the unassisted eye.
- Microscopic Level: Structural features observable with optical magnification instruments.
- Submicroscopic Level: Atomic, ionic, and molecular interactions that provide the foundational explanation for chemical reactions and physical processes.
Hierarchy and Organization of Matter:

Matter is subdivided into Substances:
- Pure Substances: Samples where every portion exhibits identical physical and chemical properties.
  - Elements: Fundamental chemical substances that cannot be decomposed into simpler substances by ordinary chemical means.
  - Currently, $108$ elements are identified.
  - A group of $40$ elements accounts for $99.9\%$ of all known matter.
  - A core group of $10$ elements constitutes $99\%$ of the Earth's crust.
  - Subclassified into Metals and Non-metals.
  - Compounds: Substances composed of two or more different elements chemically bonded in fixed proportions. They can be chemically broken down into their constituent elements.
  - Subclassified into Molecular compounds and Ionic compounds.
- Mixtures: Combinations of two or more pure substances that retain their individual chemical identities and can be separated by physical techniques.
  - Homogeneous Mixtures: Uniform in composition throughout down to the molecular level.
  - Subclassified into Solutions and Colloids.
  - Heterogeneous Mixtures: Non-uniform compositions possessing distinct phases or visual boundaries.
- Methods for Physical Separation of Mixtures:
Mixtures can be resolved into pure components via physical separation techniques including:
- Filtration: Separation based on particle size differences using a porous medium.
- Mechanical separation: Manual or physical sorting of components.
- Distillation: Separation leveraging differences in boiling points.
- Dissolving: Selective extraction utilizing solubility differences in specific solvents.
- Chromatography: Separation based on differential partitioning between stationary and mobile phases.
- Properties of Metals and Non-Metals:
Metals:
- Malleable (can be hammered into thin sheets).
- Ductile (can be drawn into wires).
- Possess metallic luster.
- Excellent conductors of heat and electricity.
- Elemental symbols differing from English names:
  - Lead: $\text{Pb}$
  - Potassium: $\text{K}$
  - Iron: $\text{Fe}$
  - Sodium: $\text{Na}$
  - Tungsten: $\text{W}$
  - Copper: $\text{Cu}$
  - Mercury: $\text{Hg}$
Non-metals:
- Solid non-metals at room temperature: Arsenic ( $\text{As}$ ), Phosphorus ( $\text{P}$ ), Sulfur ( $\text{S}$ ), Iodine ( $\text{I}$ ), Boron ( $\text{B}$ ), Selenium ( $\text{Se}$ ), and Carbon ( $\text{C}$ ).
- Liquid non-metal at room temperature: Bromine ( $\text{Br}$ ).
- All remaining non-metals exist naturally as gases at room temperature.

Physical and Chemical Properties of Matter

Distinguishing Properties:
- Physical Properties: Characteristics observed or measured without altering the underlying chemical composition or identity of the substance (e.g., color, odor, melting point, boiling point, density, index of refraction).
- Chemical Properties: Characteristics describing the reactivity of a substance as it transforms into entirely new chemical substances (e.g., reactivity with acids to liberate carbon dioxide gas, $\text{CO}_2$ ).
Physical Changes vs. Chemical Changes:
- Physical Changes: Processes altering the state or appearance of matter without changing its chemical identity (e.g., boiling a liquid, melting a solid, dissolving a solute in a solvent to form a homogeneous solution).
- Chemical Changes (Chemical Reactions): Processes involving the rearrangement of atoms and molecules to form chemically distinct substances (e.g., burning hydrogen gas, $\text{H}_2$ , in oxygen gas, $\text{O}_2$ , to yield water, $\text{H}_2\text{O}$ ).
Density as an Intensive Property:
- Intensive Properties: Physical properties independent of the total quantity of matter present (e.g., density, temperature, melting point).
- Extensive Properties: Physical properties directly proportional to the quantity of matter present (e.g., mass, volume).

Styrofoam and brick balance demonstration

Density Formula: $\text{Density} = \frac{\text{mass (g)}}{\text{volume (cm}^3\text{)}}$
Specific reference densities:
- Mercury ( $\text{Hg}$ ): $13.6\,\text{g/cm}^3$
- Gold ( $\text{Au}$ ): $19.3\,\text{g/cm}^3$
- Worked Example 1: Density Calculation of Copper:
Problem: A piece of copper has a mass of $57.54\,\text{g}$ . Its dimensions are length = $9.36\,\text{cm}$ , width = $7.23\,\text{cm}$ , and thickness = $0.95\,\text{mm}$ . Calculate its density in $\text{g/cm}^3$ .
Step 1: Convert all dimensional measurements into common units ( $\text{cm}$ ): $0.95\,\text{mm} \times \frac{1\,\text{cm}}{10\,\text{mm}} = 0.095\,\text{cm}$
Step 2: Calculate volume in cubic centimeters ( $\text{cm}^3$ ): $\text{Volume} = (9.36\,\text{cm}) \times (7.23\,\text{cm}) \times (0.095\,\text{cm}) = 6.4\,\text{cm}^3$
Step 3: Calculate density: $\text{Density} = \frac{57.54\,\text{g}}{6.4\,\text{cm}^3} = 9.0\,\text{g/cm}^3$
- Worked Example 2: Dimensional Analysis with Mercury:
Problem: Mercury ( $\text{Hg}$ ) has a density of $13.6\,\text{g/cm}^3$ . Determine the mass of $95\,\text{mL}$ of mercury in grams and in pounds ( $1\,\text{lb} = 454\,\text{g}$ ).
Step 1: Note volumetric equivalency: $1\,\text{cm}^3 = 1\,\text{mL}$
Step 2: Calculate mass in grams using dimensional analysis: $95\,\text{cm}^3 \times \frac{13.6\,\text{g}}{\text{cm}^3} = 1.3 \times 10^3\,\text{g}$
Step 3: Convert mass to pounds: $1.3 \times 10^3\,\text{g} \times \frac{1\,\text{lb}}{454\,\text{g}} = 2.8\,\text{lb}$

Measurement, Units, and Temperature Scales

Quantitative vs. Qualitative Observations:
- Qualitative Observations: Non-numerical descriptions such as color changes, gas evolution, or phase transitions.
- Quantitative Measurements: Numerical determinations accompanied by specific physical units.
International System of Units (SI Units):
- Base physical quantities:
- Length: meter ( $\text{m}$ )
- Mass: kilogram ( $\text{kg}$ ) or gram ( $\text{g}$ )
- Time: second ( $\text{s}$ )
Metric Prefixes and Powers of Ten:
- Mega: $10^6$
- Kilo: $10^3$
- Deci: $10^{-1}$
- Centi: $10^{-2}$
- Milli: $10^{-3}$
- Micro: $10^{-6}$
- Nano: $10^{-9}$
- Femto: $10^{-12}$
Length Metric Conversions and Molecular Scale:
- $1\,\text{km} = 10^3\,\text{m}$
- $1\,\text{m} = 100\,\text{cm}$
- $1\,\text{cm} = 10\,\text{mm}$
- $1\,\text{nm} = 1.0 \times 10^{-9}\,\text{m}$
- Molecular example: The $\text{O-H}$ bond distance in a water molecule ( $\text{H}_2\text{O}$ ) measures: $9.4 \times 10^{-11}\,\text{m} = 9.4 \times 10^{-9}\,\text{cm} = 0.094\,\text{nm}$
Temperature Scale Comparisons:

Comparison of Fahrenheit, Celsius, and Kelvin scales

Key Historical Figures:
- Anders Celsius ( $1701\text{--}1744$ )
- Lord Kelvin / William Thomson ( $1824\text{--}1907$ )
Reference Thermal Points:
- Boiling Point of Water: $212^\circ\text{F} = 100^\circ\text{C} = 373.15\,\text{K}$
- Freezing Point of Water: $32^\circ\text{F} = 0^\circ\text{C} = 273.15\,\text{K}$
- Scale Intervals: The interval between freezing and boiling is $180^\circ$ on the Fahrenheit scale, $100^\circ$ on the Celsius scale, and $100\,\text{K}$ on the Kelvin scale.
- Degree Equivalence: An increment of $1\,\text{K}$ is identical in magnitude to an increment of $1^\circ\text{C}$ .
Temperature Conversions: $T(\text{K}) = T(^\circ\text{C}) + 273.15$
- Human Body Temperature: $37^\circ\text{C} + 273 = 310\,\text{K}$
- Liquid Nitrogen Temperature: $-196^\circ\text{C} + 273 = 77\,\text{K}$
- Ambient Warm Temperature Comparison: $100^\circ\text{F} \approx 38^\circ\text{C} \approx 311\,\text{K}$

Data Quality: Precision, Accuracy, and Error

Accuracy vs. Precision:
- Accuracy: The closeness of a measured value to the true or accepted standard value.
- Precision: The reproducibility or agreement between replicate measurements of the same quantity under identical conditions.
Target Graphical Analogy:

Target diagrams illustrating precision and accuracy

Poor Precision and Poor Accuracy: Darts are widely scattered across the target away from the center.
Good Precision and Poor Accuracy: Darts are tightly grouped together in a cluster far off-center.
Good Precision and Good Accuracy: Darts are tightly grouped together directly in the central bullseye.
- Quantifying Experimental Error: $\text{Error} = \text{experimental value} - \text{accepted value}$ $\text{Percent Error} = \frac{\text{Error}}{\text{accepted value}} \times 100\%$

Significant Figures and Mathematical Rules

Standard Rules for Determining Significant Figures:
1. Non-zero digits and captured zeros: All non-zero digits and zeros situated between non-zero digits are significant.
- Example: $6023$ contains $4$ significant figures.
1. Leading zeros: Zeros located to the left of the first non-zero digit serve solely as decimal place locators and are not significant.
- Example: $0.0006023$ contains $4$ significant figures.
1. Trailing zeros with decimal point: Zeros appearing at the end of a number to the right of a decimal point are significant.
- Example: $2.200$ contains $4$ significant figures.
1. Trailing zeros without decimal point: Zeros at the end of a whole number without an explicitly written decimal point are not significant.
- Example: $600000$ contains only $1$ significant figure.
Arithmetic Rules for Calculations:
- Multiplication and Division: The final calculated result cannot contain more significant digits than the measurement possessing the fewest total significant figures.
- Calculated Example: $\frac{104.1 \times 1.2780}{7.3} = 18.22463013699$
  - Since $7.3$ contains only $2$ significant figures, the result must be rounded to $2$ significant figures: 18$.\n * **Addition and Subtraction**: The final calculated result cannot contain more decimal places (digits to the right of the decimal point) than the measurement possessing the fewest decimal places.\n * *Calculated Example*:\n 123.456 + 34.432 + 1.28 + 100.02 = 259.188\n * Since 1.28 $and$ 100.02 $contain$ 2 $decimal places, the result must be rounded to$ 2 $decimal places:$ 259.19$.

Fundamental Laws of Chemical Combination

Law of Constant Composition:
- The relative mass proportions of each constituent element in a given chemical compound are fixed, definite, and invariant regardless of the source or sample size.
Mass Percentage (Percent Composition): $\text{Mass Percentage} = \frac{\text{Mass of Element}}{\text{Mass of Compound}} \times 100\%$
Worked Calculation Example:
- A sample of iron sulfide compound weighing $2.46\,\text{g}$ contains $1.56\,\text{g}$ of iron ( $\text{Fe}$ ) and $0.9007\,\text{g}$ of sulfur ( $\text{S}$ ). Calculate the mass percentages of iron and sulfur.
- Mass Percentage of Iron ( $\text{Fe}$ ): $\%\,\text{mass}_{\text{Fe}} = \frac{1.56\,\text{g}}{2.46\,\text{g}} \times 100\% = 63.4146\% = 63.4\%$
- Mass Percentage of Sulfur ( $\text{S}$ ): $\%\,\text{mass}_{\text{S}} = \frac{0.9007\,\text{g}}{2.46\,\text{g}} \times 100\% = 36.61\% = 36.6\%$