Chapter 1: Chemistry Basics – Matter and Measurement (Study Guide)

Chapter 1: Chemistry Basics – Matter and Measurement

1.2 The Periodic Table: Groups, Periods, Metals, and Nonmetals

The periodic table is a fundamental tool in chemistry, organizing elements by their properties and atomic structure. Understanding its layout is essential for identifying element types and interpreting chemical formulas.

• Groups: Vertical columns on the periodic table. Elements in the same group have similar chemical properties due to the same number of valence electrons.

• Periods: Horizontal rows on the periodic table. Elements in the same period have the same number of electron shells.

• Metals: Located on the left and center of the periodic table. They are typically shiny, good conductors of heat and electricity, malleable, and ductile.

• Nonmetals: Found on the right side of the periodic table. They are often dull, poor conductors, and brittle in solid form.

• Chemical Formulas: Indicate the types and numbers of atoms in a compound. For example, H2O contains 2 hydrogen atoms and 1 oxygen atom.

Example: In NaCl, there is 1 sodium (Na) atom and 1 chlorine (Cl) atom per formula unit.

1.3 Physical and Chemical Changes

Understanding the difference between physical and chemical changes is crucial for classifying processes in chemistry.

• Physical Change: A change that affects the form of a substance, but not its chemical identity. Examples include melting, freezing, and dissolving.

• Chemical Change: A process where one or more substances are converted into new substances with different properties. Examples include burning, rusting, and digestion.

Example: Ice melting to water is a physical change; iron rusting is a chemical change.

1.4 Measurement and Significant Figures

Accurate measurement and correct reporting of significant figures are essential in scientific work. Significant figures reflect the precision of a measurement.

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

• Certain and Uncertain Digits: All digits except the last are certain; the last digit is uncertain (estimated).

• Rules for Counting Significant Figures:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros are significant only if there is a decimal point.

• Reporting Measurements: Use the correct number of significant figures based on the measuring device.

• Rounding: Round to the correct number of significant figures after calculations.

• Rules for Calculations:

  • Addition/Subtraction: The result should have the same number of decimal places as the measurement with the fewest decimal places.

  • Multiplication/Division: The result should have the same number of significant figures as the measurement with the fewest significant figures.

Example: 12.34 + 1.2 = 13.5 (rounded to one decimal place)

Example: 2.5 × 3.42 = 8.6 (rounded to two significant figures)

1.4 Scientific Notation and Unit Conversions

Scientific notation and unit conversions are essential for expressing and manipulating measurements in chemistry.

• Scientific Notation: A way to express very large or small numbers using powers of ten.

  • Example: 0.00056 = $5.6 \times 10^{-4}$

  • Example: 3,200,000 = $3.2 \times 10^{6}$

• Metric Prefixes: Common prefixes include:

  Prefix	Symbol	Value
  micro-	μ	$1 \times 10^{-6}$
  milli-	m	$1 \times 10^{-3}$
  centi-	c	$1 \times 10^{-2}$
  deci-	d	$1 \times 10^{-1}$
  kilo-	k	$1 \times 10^{3}$

• Conversion Factors: Ratios used to convert from one unit to another. Set up so that units cancel appropriately.

• 1-Step Unit Conversions: Use a single conversion factor to change units.

  • Example: Convert 5.0 cm to mm: $5.0\ \text{cm} \times \frac{10\ \text{mm}}{1\ \text{cm}} = 50\ \text{mm}$

1.5 Volume and Density Calculations

Volume and density are key physical properties measured in chemistry. Understanding how to determine these values is fundamental for laboratory work.

• Volume by Water Displacement: The volume of a solid can be found by measuring the change in water level when the object is submerged.

• Density: The mass per unit volume of a substance. Calculated using the formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Units: Commonly expressed in g/mL or g/cm3 for solids and liquids.

Example: If a 10.0 g object displaces 2.5 mL of water, its density is $\frac{10.0\ \text{g}}{2.5\ \text{mL}} = 4.0\ \text{g/mL}$

1.6 Dosage Calculations

Dosage calculations are important in health sciences and pharmacy, requiring the application of unit conversions and ratios.

• 1-Step Dosage Problems: Use a single conversion factor to determine the correct dose based on patient weight, concentration, or other factors.

• Formula Example: $\text{Dose} = \text{Patient Weight} \times \frac{\text{Dosage}}{\text{Weight}}$

Example: If a medication is prescribed at 5 mg/kg and the patient weighs 70 kg, the dose is $70\ \text{kg} \times \frac{5\ \text{mg}}{1\ \text{kg}} = 350\ \text{mg}$

Additional info: This study guide covers foundational measurement and calculation skills essential for success in GOB Chemistry and related laboratory work.