Foundations of Chemistry: Matter, Measurement, and Chemical Compounds

Chapter 1: Chemistry in Our Lives

Definition of Chemistry and Chemicals

Chemistry is the scientific study of matter, its properties, composition, and the changes it undergoes. Chemicals are substances with a definite composition, found both naturally and synthetically.

• Chemistry explores how substances interact, combine, and change to form new substances.

• Chemicals include everything from water (H2O) to table salt (NaCl).

• Example: The rusting of iron is a chemical process involving iron and oxygen.

The Scientific Method

The scientific method is a systematic approach to research and discovery in science.

• Steps include: Observation, Hypothesis, Experiment, Conclusion, and Theory.

• Ensures that scientific knowledge is based on evidence and repeatable results.

• Example: Testing whether salt affects the boiling point of water by conducting controlled experiments.

Strategies and Study Plan for Chemistry

• Practice problem-solving regularly.

• Review key concepts and definitions.

• Use visual aids like diagrams and tables.

• Form study groups for discussion and clarification.

Key Math Skills for Chemistry

• Place Values: Understanding the value of digits in numbers.

• Calculations: Performing arithmetic operations accurately.

• Percentages: Calculating parts per hundred, often used in concentration and yield.

• Algebraic Equations: Solving for unknowns in chemical formulas and reactions.

• Scientific Notation: Expressing very large or small numbers, e.g., $6.02 \times 10^{23}$.

• Graph Interpretation: Reading and analyzing data from graphs.

Absolute Uncertainty/Precision

• Graduated devices (like rulers or burettes) have inherent measurement uncertainty.

• Precision refers to how close repeated measurements are to each other.

• Example: Measuring 25.0 mL with a graduated cylinder may have an uncertainty of ±0.1 mL.

Chapter 2: Chemistry and Measurements

2.1 Units of Measurement

Standard units are essential for consistency in scientific communication.

• Length: meter (m)

• Volume: liter (L)

• Mass: gram (g)

• Temperature: Celsius (°C), Kelvin (K)

• Time: second (s)

2.2 Measured Numbers & Significant Figures

• Significant Figures (Sig Figs): Digits in a measurement that are known with certainty plus one estimated digit.

• Measured Numbers: Obtained by measurement; have uncertainty.

• Exact Numbers: Defined values or counted objects; have no uncertainty.

• Example: 2.50 g has three significant figures.

2.3 Significant Figures in Calculations

• Multiplication/Division: Result has as many sig figs as the value with the fewest sig figs.

• Addition/Subtraction: Result has as many decimal places as the value with the fewest decimal places.

• Mixed Operations: Apply rules stepwise, following order of operations.

• Rounding: Adjust final answer to correct sig figs.

2.4–2.6 Problem Solving Using Equalities and Conversion Factors

• Metric System: Uses base units and prefixes (e.g., kilo-, centi-, milli-).

• Conversion Factors: Ratios used to convert from one unit to another.

• Unit Line Equations: Set up conversions so units cancel appropriately.

• Example: To convert 5.0 cm to meters: $5.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.050\ \text{m}$

• Medical Dosage, Percent, and Density: Use conversion factors to solve practical problems.

2.7 Density

• Density: The mass of a substance per unit volume.

• Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Units: g/mL (liquids), g/cm3 (solids)

• Floating/Sinking: An object floats if its density is less than the liquid's density.

• Calculating by Displacement: Volume of solid = final volume - initial volume (when submerged in liquid).

Chapter 3: Matter and Energy

3.1 Classification of Matter

• Matter: Anything that has mass and occupies space.

• Pure Substances: Elements (single type of atom) and compounds (two or more elements chemically combined).

• Mixtures: Physical combinations of substances; can be homogeneous (uniform, e.g., saltwater) or heterogeneous (non-uniform, e.g., salad).

3.2 States and Properties of Matter

• States: Solid (definite shape/volume), Liquid (definite volume, variable shape), Gas (variable shape/volume).

• Particle Motion: Increases from solid to gas.

• Physical Properties: Observed without changing composition (e.g., melting point).

• Chemical Properties: Describe ability to undergo chemical change (e.g., flammability).

• Physical Change: No new substance formed (e.g., melting ice).

• Chemical Change: New substance formed (e.g., burning wood).

3.3 Temperature

• Scales: Fahrenheit (°F), Celsius (°C), Kelvin (K).

• Kelvin: Absolute zero (0 K) is the lowest possible temperature; no negative values.

• Conversions: Use provided equations to convert between scales.

3.4 Energy

• Potential Energy: Stored energy due to position or composition.

• Kinetic Energy: Energy of motion.

• Heat vs. Temperature: Heat is energy transfer due to temperature difference; temperature measures average kinetic energy.

• Energy Units: 1 calorie (cal) = 4.184 joules (J).

3.5 Energy & Nutrition

• Nutritional Calorie (Cal): 1 Cal = 1000 cal = 1 kcal = 4184 J.

• Used to quantify energy content in food.

3.6 Specific Heat

• Specific Heat (SH): Amount of heat needed to raise 1 g of a substance by 1°C.

• Heat Equation: $\text{Heat} = m \times \Delta T \times SH$

• Used to calculate energy changes in heating/cooling.

3.7 Changes of State

• Melting/Freezing Point: Temperature at which solid and liquid phases are in equilibrium.

• Vaporization/Condensation Point: Temperature at which liquid and gas phases are in equilibrium.

• Heating/Cooling Curves: Graphs showing temperature changes during phase transitions.

Chapter 4: Atoms, the Periodic Table, and Isotopes

4.1–4.5: Classification within the Periodic Table

• Metals: Left and center; conduct electricity, malleable.

• Nonmetals: Right side; poor conductors, brittle.

• Metalloids: Border between metals and nonmetals; properties of both.

• Groups/Families: Vertical columns; similar properties.

• Periods: Horizontal rows; properties change progressively.

• Special Groups: Alkali metals (Group 1), Alkaline earth metals (Group 2), Halogens (Group 17), Noble gases (Group 18), Transition metals (center block).

Subatomic Particles

Particle	Charge	Location	Relative Mass
Proton	+1	Nucleus	1
Neutron	0	Nucleus	1
Electron	-1	Outside nucleus	~0

• Atomic Number (Z): Number of protons; defines the element.

• Mass Number (A): Protons + neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Isotope Notation: $^{A}_{Z}\text{X}$, where X is the element symbol.

• Example: $^{14}_{6}\text{C}$ is carbon-14 (6 protons, 8 neutrons).

4.6 Electron Energy Levels

• Electromagnetic Radiation (EMR): Includes radio, microwave, infrared, visible, ultraviolet, X-ray, gamma (increasing energy/frequency).

• Atomic Spectra: Unique patterns of light emitted by elements due to electron transitions.

• Electron Energy Levels: Electrons occupy specific energy levels; each level holds a maximum number of electrons (2, 8, 18, ...).

• Valence Electrons: Outermost electrons; determine chemical properties.

4.7 Periodic Trends

• Electron-Dot (Lewis) Symbols: Dots represent valence electrons around element symbol.

• Valence Electrons: Equal to group number for main group elements.

• Trends:

  • Atomic Size: Increases down a group, decreases across a period.

  • Ionization Energy: Energy to remove an electron; increases across a period, decreases down a group.

  • Metallic Character: Increases down a group, decreases across a period.

Chapter 6 (Part 1): Ionic & Covalent Compounds

Ionic Compounds

• Formation: Metals lose electrons (form cations), nonmetals gain electrons (form anions).

• Representative Elements: Groups 1, 2, 13 (Al), 15, 16, 17 form predictable charges.

• Transition Metals: Can form multiple cations; charge indicated by Roman numeral (e.g., iron(III) chloride).

• Formula: Compounds must be charge neutral.

• Polyatomic Ions to Memorize:

  • Ammonium, NH4+

  • Nitrate, NO3-

  • Sulfate, SO42-

  • Phosphate, PO43-

  • Carbonate, CO32-

  • Bicarbonate, HCO3-

  • Cyanide, CN-

  • Hydroxide, OH-

• Naming: No prefixes; use element names and ion charges.

• Example: NaCl (sodium chloride), CuCl2 (copper(II) chloride).

Covalent/Molecular Compounds

• Covalent Bonding: Nonmetals share electrons to form molecules.

• Diatomic Elements: N2, H2, F2, O2, I2, Cl2, Br2.

• Identifying Compounds: Ionic = metal + nonmetal; covalent = nonmetals only.

• Naming: Use prefixes (mono-, di-, tri-, tetra-, etc.) for number of atoms.

• Valence Electrons: Equal to last digit of group number.

• Lewis Diagrams: Show bonding and lone pairs in molecules.

• Bond Prediction: Number of covalent bonds = number of unpaired valence electrons.

Electronegativity and Polarity of Molecular Compounds (6.7–6.9)

• Bond Types:

  • Nonpolar Covalent: Equal sharing of electrons (similar EN values).

  • Polar Covalent: Unequal sharing (moderate EN difference).

  • Ionic: Electron transfer (large EN difference).

• VSEPR Theory: Predicts 3D molecular shape based on electron pairs around central atom.

• Determining Polarity: Draw Lewis diagram, determine shape, and assess dipole moments.

• Intermolecular Forces:

  • Dispersion (London): Weakest; present in all molecules, dominant in nonpolar.

  • Dipole-Dipole: Stronger; present in polar molecules.

  • Hydrogen Bonding: Strongest; occurs when H is bonded to N, O, or F.

• Relative Strengths: Dispersion < Dipole-Dipole < Hydrogen Bonding.

• Identifying Hydrogen Bonds: Look for N-H, O-H, or F-H bonds.