Chemistry Notes - Module 1

Chemistry Introduction

Chemistry Definition

• Chemistry is the study of matter and the changes it undergoes.
• Matter exists in three states: gas, liquid, and solid.

Matter

• Matter is defined as anything that has mass and takes up space.
• Examples include: oxygen, water, carbon dioxide, ethanol, ethylene glycol, and aspirin.
• Matter can exist as elements, molecules of an element, molecules of a compound, or a mixture of elements and a compound.

Atoms

• Atoms are the building blocks of matter.
• Each element is made of the same kind of atom.
• A compound is made of two or more different kinds of elements.

States of Matter

• Gas: Total disorder, much empty space, particles have complete freedom of motion, particles far apart.
• Liquid: Disorder, particles or clusters of particles are free to move relative to each other, particles close together.
• Crystalline Solid: Ordered arrangement, particles are essentially in fixed positions, particles close together.
• Changes in state occur with heating/cooling or increasing/reducing pressure.

Properties of Matter

Physical Properties

• Can be observed without changing a substance into another substance.
• Examples: boiling point, density, mass, volume.

Chemical Properties

• Can only be observed when a substance is changed into another substance.
• Examples: flammability, corrosiveness, reactivity with acid.

Intensive Properties

• Independent of the amount of the substance present.
• Examples: density, boiling point, color.

Extensive Properties

• Depend upon the amount of the substance present.
• Examples: mass, volume, energy.

Types of Changes

Physical Changes

• Changes in matter that do not change the composition of a substance.
• Examples: changes of state, temperature, volume.

Chemical Changes

• Result in new substances.
• Examples: combustion, oxidation, decomposition.

Chemical Reactions

• Reacting substances are converted to new substances.
• Example: Hydrogen and oxygen gases burn to form water.
• $2 H<em>2O \rightarrow 2H</em>2 + O_2$

Classification of Matter

• Matter
  • Does it have constant properties and composition?
    • Yes: Pure substance
      • Can it be simplified chemically?
        • No: Element
        • Yes: Compound
    • No: Mixture
      • Is it uniform throughout?
        • Yes: Homogeneous
        • No: Heterogeneous

Separation of Mixtures

Distillation

• Uses differences in boiling points to separate a homogeneous mixture into its components.

Filtration

• Solid substances are separated from liquids and solutions to separate heterogeneous mixtures.

Units of Measurement

SI Units (Système International d'Unités)

• Mass: Kilogram (kg)
• Length: Meter (m)
• Time: Second (s)
• Temperature: Kelvin (K)
• Amount of substance: Mole (mol)

Metric System Prefixes

• Prefixes convert the base units into units that are appropriate for the item being measured.
  • Giga (G): $10^9$ (e.g., 1 Gm = $1 \times 10^9$ m)
  • Mega (M): $10^6$ (e.g., 1 Mm = $1 \times 10^6$ m)
  • Kilo (k): $10^3$ (e.g., 1 km = $1 \times 10^3$ m)
  • Deci (d): $10^{-1}$ (e.g., 1 dm = 0.1 m)
  • Centi (c): $10^{-2}$ (e.g., 1 cm = 0.01 m)
  • Milli (m): $10^{-3}$ (e.g., 1 mm = 0.001 m)
  • Micro ($\mu$): $10^{-6}$ (e.g., 1 $\mu$m = $1 \times 10^{-6}$ m)
  • Nano (n): $10^{-9}$ (e.g., 1 nm = $1 \times 10^{-9}$ m)
  • Pico (p): $10^{-12}$ (e.g., 1 pm = $1 \times 10^{-12}$ m)
  • Femto (f): $10^{-15}$ (e.g., 1 fm = $1 \times 10^{-15}$ m)

Volume

• Commonly used metric units for volume are the liter (L) and the milliliter (mL).
  • 1 L is a cube 1 dm long on each side.
  • 1 mL is a cube 1 cm long on each side.
  • $1 mL = 1 cm^3$

Uncertainty in Measurements

• Different measuring devices have different uses and different degrees of accuracy.
• Examples of measuring devices: Graduated cylinder, Syringe, Buret, Pipet, Volumetric flask

Temperature

• Temperature is a measure of the average kinetic energy of the particles in a sample.
• Scales:
  • Kelvin (K)
  • Celsius (°C)
  • Fahrenheit (°F)

Temperature Scales

• Celsius scale is based on the properties of water.
  • 0°C is the freezing point of water.
  • 100°C is the boiling point of water.
• Kelvin is the SI unit of temperature.
  • Based on the properties of gases.
  • No negative Kelvin temperatures.
  • $K = °C + 273.15$
• Fahrenheit scale is not used in scientific measurements.
  • $°F = \frac{9}{5}(°C) + 32$
  • $°C = \frac{5}{9}(°F - 32)$

Temperature Conversion Examples

• Convert 172.9 °F to degrees Celsius:
  • $°C = \frac{5}{9} \times (172.9 - 32) = 78.3$

Density

• Density is a physical property of a substance.
• $d = \frac{m}{V}$
  • d = density
  • m = mass
  • V = volume

Density Example Calculations

• A 525 mL sample of an unknown liquid has a mass of 375 g. Determine the density of this liquid in g/mL.
  • $Density = \frac{375 g}{525 mL} = 0.714 g/mL$
• A 425 g piece of an unknown metal has dimensions of 8.0 cm x 4.0 cm x 2.0 cm. Determine the density of this metal in g/cm3.
  • $Volume = 8.0 cm \times 4.0 cm \times 2.0 cm = 64 cm^3$
  • $Density = \frac{425 g}{64 cm^3} = 6.6 g/cm^3$

Significant Figures

• Significant figures refer to digits that were measured.
• When rounding calculated numbers, pay attention to significant figures so we do not overstate the accuracy of our answers.

Rules for Significant Figures

• All nonzero digits are significant.
• Zeroes between two significant figures are themselves significant.
• Zeroes at the beginning of a number are never significant.
• Zeroes at the end of a number are significant as long as the number has a decimal point.

Examples of Significant Figures

• 1.234 kg: 4 significant figures
• 606 m: 3 significant figures
• 0.08 L: 1 significant figure
• 2.0 mg: 2 significant figures
• 0.00420 g: 3 significant figures
• 2040. mL: 4 significant figures

Significant Figures in Measurements

• 24 mL: 2 significant figures
• 3001 g: 4 significant figures
• 0.0320 m3: 3 significant figures
• $6.4 \times 10^4$ molecules: 2 significant figures
• 560 kg: 2 significant figures
• 560. kg: 3 significant figures

Significant Figures in Calculations

• Addition or Subtraction: Answers are rounded to the least significant decimal place.
• Multiplication or Division: Answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation.

Rounding Off

• Go one digit to the right of where you want to cut off the number.
• If this value is larger than 5, then increase the value of the cutoff number by 1.
• If this value is smaller than 5, the value of the cutoff number stays the same.

Rounding Off With 5: The Odd/Even Rule

• If this value to the right of the cutoff number is 5 and no other digits follow, then we use the odd/even rule.
• When the cutoff number is odd, this cutoff number must be rounded up.
• When the cutoff number is even, the cutoff number stays the same.

Examples of Rounding

• Round 16.35 to 3 significant figures: 16.4
• Round 16.65 to 3 significant figures: 16.6

Examples of Rounding to 3 Digits

• 68.34 mL round to 68.3 mL
• 3528 g round to 3530 g
• 0.32952 cm3 round to 0.330 cm3
• 6465 miles round to 6460 miles
• 563.5 kg round to 564 kg

Significant Figures: Addition and Subtraction

• The answer cannot have more digits to the right of the decimal point than any of the original numbers.
  • Example: 89.332 + 1.1 = 90.432 round off to 90.4 (one significant figure after decimal point)
  • Example: 3.70 - 2.9133 = 0.7867 round off to 0.79 (two significant figures after decimal point)

Significant Figures: Multiplication and Division

• The number of significant figures in the result is set by the original number that has the smallest number of significant figures.
  • Example: $4.51 \times 3.6666 = 16.536366$ round to 16.5 (3 sig figs)
  • Example: $\frac{6.8}{112.04} = 0.0606926$ round to 0.061 (2 sig figs)

Significant Figures: Exact Numbers

• Numbers from definitions or numbers of objects are considered to have an infinite number of significant figures.
• Example: 1 in. = 2.54 cm is the exact conversion between inches and centimeters and does not affect the number of digits we report.

Accuracy versus Precision

• Accuracy: Refers to the proximity of a measurement to the true value of a quantity.
• Precision: Refers to the proximity of several measurements to each other.

Dimensional Analysis

• We use dimensional analysis to convert one quantity to another.
• Most commonly dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm)
• $\frac{1 in.}{2.54 cm} or \frac{2.54 cm}{1 in.}$
• $Given\ unit \times \frac{Desired\ unit}{Given\ unit} = Desired\ unit$
• Example to convert 8.00 m to inches:
  • $8.00 m \times \frac{100 cm}{1 m} \times \frac{1 in.}{2.54 cm} = 315 in.$

Dimensional Analysis Example

• How many mL are in 1.63 L?
  • $1.63 L \times \frac{1000 mL}{1 L} = 1630 mL$

Speed of Sound Conversion Example

• The speed of sound in air is about 343 meters per second. What is this speed in miles per hour?
  • $343 \frac{m}{sec} \times \frac{1 mi}{1609 m} \times \frac{60 sec}{1 min} \times \frac{60 min}{1 hour} = 767 \frac{mi}{hour}$

Working with Squared and Cubed Units

• To convert squared units, we must square the conversion factors.
  • Example: Convert 1.53 m2 into in2
    • $1.53 m^2 \times (\frac{100 cm}{1 m})^2 \times (\frac{1 in.}{2.54 cm})^2 = 2370 in.^2$
• To convert cubed units, we must cube the conversion factors.
  • Example: Convert 3.31 m3 into in3
    • $3.31 m^3 \times (\frac{100 cm}{1 m})^3 \times (\frac{1 in.}{2.54 cm})^3 = 2.02 \times 10^5 in.^3$