Chemistry: The Central Science – Chapter 1 Study Notes

1. What is Chemistry?

• Chemistry: The study of matter, its properties, and the changes it undergoes.

• It is fundamental to understanding many scientific fields.

2. Matter

• Matter: Anything that has mass and takes up space.

• Classified by:

  • State: solid, liquid, gas

  • Composition: element, compound, mixture

3. States of Matter

1. Solid: Definite shape and volume

2. Liquid: Definite volume, no definite shape

3. Gas: No definite shape or volume

Example: Ice (solid), liquid water, water vapor (gas)

4. Classification by Composition

Substances

• Pure, distinct properties, constant composition.

  • Element: Cannot be broken down into simpler substances (e.g., O₂, Fe).

  • Compound: Made of two or more elements (e.g., H₂O, CO₂).

Mixtures

• Blend of two or more substances.

  • Homogeneous: Uniform composition (e.g., saltwater, air).

  • Heterogeneous: Non-uniform composition (e.g., salad, granite).

5. Atoms and Molecules

• Atoms: Building blocks of matter.

• Molecules: Groups of atoms bonded together.

• Elements: Made of one type of atom.

• Compounds: Made of two or more different elements.

6. Properties of Matter

Physical Properties

• Can be observed without changing the substance.

• Examples: color, density, melting point, hardness.

Chemical Properties

• Observed only when the substance changes.

• Example: flammability, reactivity.

Intensive vs. Extensive Properties

• Intensive: Independent of amount (e.g., density, color).

• Extensive: Depends on amount (e.g., mass, volume).

7. Changes in Matter

Physical Changes

• Do not change composition.

• Examples: melting, boiling, cutting.

Chemical Changes

• Produce new substances.

• Examples: burning, rusting, decomposing.

8. Separating Mixtures

Based on physical properties:

• Filtration: Separates solids from liquids.

• Distillation: Separates based on boiling points.

• Chromatography: Separates based on adhesion.

9. Energy

• Energy: Capacity to do work or transfer heat.

• Work: Force causing displacement.

• Heat: Energy that increases temperature.

Forms of Energy

• Kinetic: Energy of motion → KE=12mv2KE=21​mv2

• Potential: Stored energy due to position.

10. Units of Measurement

SI Base Units

• Mass: kilogram (kg)

• Length: meter (m)

• Time: second (s)

• Temperature: Kelvin (K)

• Amount: mole (mol)

Metric Prefixes

• kilo (k) = 10³

• centi (c) = 10⁻²

• milli (m) = 10⁻³

• micro (μ) = 10⁻⁶

Volume

• Derived from length.

• Common units: Liter (L), milliliter (mL)

• 1 L = 1000 mL = 1 dm³

Temperature Scales

• Celsius (°C): Based on water (0°C freezing, 100°C boiling)

• Kelvin (K): SI unit, no negatives → K=°C+273.15K=°C+273.15

• Fahrenheit (°F): Not scientific

11. Density

• Density=MassVolumeDensity=VolumeMass​

• Units: g/cm³ or g/mL

• Intensive property

12. Numbers in Science

Exact vs. Inexact Numbers

• Exact: Counted or defined (e.g., 12 eggs in a dozen)

• Inexact: Measured (always some uncertainty)

Accuracy vs. Precision

• Accuracy: Closeness to true value.

• Precision: Closeness of repeated measurements.

13. Significant Figures

• All certain digits + one uncertain digit.

• Rules:

  • Non-zero digits are significant.

  • Zeros between non-zero digits are significant.

  • Leading zeros are NOT significant.

  • Trailing zeros ARE significant if there’s a decimal point.

Calculations with Sig Figs

• +/–: Round to least precise decimal.

• ×/÷: Round to least number of sig figs.

14. Dimensional Analysis

• Method to convert units using conversion factors.

• Example:
  2.54cm=1in2.54cm=1in
  → 2.54cm1in1in2.54cm​ or 1in2.54cm2.54cm1in​

Key Equations & Relationships

• K=°C+273.15K=°C+273.15

• °C=59(°F−32)°C=95​(°F−32)

• Density=mVDensity=Vm​

• KE=12mv2KE=21​mv2

• 1 cal = 4.184 J
  1 Cal (nutritional) = 1000 cal

KE= ½ mv²= unit of energy Joule

k= degree celsius + 273.15

most common units

g/mL or g/cm³