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Welcome to Chemistry 116 Course

• Instructor: Prof. Khotseng

• Location: New Chemical Sciences Building, 3rd Floor, Room 3.23

• Contact: lkhotseng@uwc.ac.za

• Extension: 4031

• Consultation Times: Monday - Thursday 12:00 - 1:00 PM, or by appointment (email ahead). Casual drop-ins possible but may result in waiting due to busy schedule.

Course Topics

• Chapter 1: Basic concepts of Chemistry

• Chapter 2: Atoms, molecules, and Ions

• Chapter 3: Chemical Reactions

• Chapter 4: Stoichiometry

• Chapter 5: Thermochemistry

• Chapter 6: Electronic structure of atoms

• Chapter 7: The structure of atoms and periodic trends

• Chapter 8: Chemical bonding and Molecular geometry

• Chapter 10: Gases and their properties

• Chapter 13: Solutions and Their Behavior

• Chapter 19: Oxidation-Reduction Reactions

• A total of 11 chapters covered

Chapter 1: Basic Concepts of Chemistry

Understanding Chemistry

• Definition: Chemistry is the science of change, focusing on:

  • Study of matter: Composition, properties, and changes.

  • Qualitative information: Non-numerical data (e.g., color, physical appearance).

  • Quantitative information: Numerical data (e.g., mass, temperature).

• Key Principle: Law of conservation of matter – Mass is conserved in chemical reactions.

Matter

Definition of Matter

• Matter: Anything that occupies space and has mass.

• Mass: Measure of the amount of matter an object contains.

States of Matter

• Solids: Rigid shape, fixed volume.

• Liquids: Fixed volume, no fixed shape; do not fill container completely.

• Gases: No fixed volume or shape; expand to fill their container.

Kinetic Molecular Theory of Matter

• All matter consists of tiny particles (atoms, molecules, or ions) in constant motion:

  • Solids: Particles closely packed, lowest kinetic energy.

  • Liquids: Randomly arranged particles, intermediate kinetic energy.

  • Gases: Far apart particles, highest kinetic energy.

Chemistry & Matter

• Macroscale vs. Particulate:

  • Explore the macroscopic world (observable) to understand the particulate world (invisible).

  • Symbols: Written to describe aspects of matter.

A Chemist's View of Matter

• Chemical equation example:

  • 2 H2(g) + O2(g) --> 2 H2O(g)

Classification of Matter

Is it uniform throughout?

• Heterogeneous mixture: Not uniform throughout.

• Homogeneous mixture: Uniform composition.

Variable Composition

• Homogeneous Mixture/Solution: Uniform.

• Pure substance: Consistent composition.

Separation into Simpler Substances

• Element: Cannot be separated.

• Compound: Can be separated into elements.

Chemical Elements

Definition

• Pure substances that cannot be broken down into other substances by ordinary means.

• Examples: Sodium (Na), Bromine (Br), Aluminum (Al).

Chemical Symbols

• Abbreviations: Simple one or two-letter symbols to represent elements, capitalized (e.g., C for carbon, Ne for neon).

• Symbols based on Latin/Greek names for some elements (e.g., Sb for Antimony, Pb for Lead).

Periodic Table

• Contains names and symbols for all elements (118 known as of 2012).

Atoms and Molecules

Definition of an Atom

• The smallest particle of an element that has chemical properties.

Definition of a Molecule

• Particles consisting of more than one atom held together by chemical bonds.

Chemical Compounds

• Definition: Substances composed of two or more different elements bonded together chemically (e.g., H2O).

Properties of Matter

Chemical Properties

• Ability of matter to undergo a change in composition under certain conditions (e.g., rusting of iron).

Physical Properties

• Observed or measured without altering the composition (e.g., color, odor, melting point).

Classification of Properties

• Intensive properties: Not dependent on the amount of substance (e.g., density).

• Extensive properties: Depend on the amount of substance (e.g., mass, volume).

Density

• Definition: Relation of mass to volume; a useful physical property.

• Intensive property: Density independent of sample size.

• Example values: Mercury (13.6 g/cm3), Aluminum (2.7 g/cm3).

Changes in Matter

Physical Changes

• Examples include evaporation, melting, and dissolving solids.

Chemical Changes

• Transformation of substances into new compositions (e.g., iron + oxygen = rust).

Observations and Measurements

• Qualitative observations: Non-numerical (color, state).

• Quantitative measurements: Based on numerical data using SI units.

Units of Measurement

• Common SI Units:

  • Length: meter (m)

  • Time: seconds (s)

  • Mass: kilograms (kg)

  • Temperature: kelvin (K)

  • Amount of substance: moles (mol)

Temperature Scales

• Celsius and Kelvin: Boiling point of water at 100°C (373.15 K) and freezing point at 0°C (273.15 K).

• Formula for conversion: T(K) = t(°C) + 273.15

Significant Figures

Rules for Counting Significant Figures

1. All non-zero digits are significant.

2. All zeros between non-zero digits are significant.

3. Leading zeros are not significant.

4. Trailing zeros count if there’s a decimal.

Calculations with Significant Figures

• Use the number of significant figures in the least precise measurement for final result accuracy.

Dimensional Analysis

Steps

1. Write down the value to convert.

2. Use parentheses to divide units and relate to each other.

3. Assign numbers to units and multiply for final result, ensuring correct units and significant figures.

Example Problem

• Calculating Density: mass/volume.

Conclusion

What You Should Learn

• Define key terms (matter, element, atom, etc.).

• Memorize metric prefixes.

• Use dimensional analysis and significant figures properly.

• Be adept at converting measurements and calculating density.

Questions

• Open for student queries!