General Chemistry 101: Comprehensive Study Guide

Temperature Conversions

Understanding Temperature Scales

Temperature is a fundamental measurement in chemistry, often requiring conversion between Fahrenheit (°F), Celsius (°C), and Kelvin (K). Accurate conversions are essential for laboratory work and problem solving.

• Fahrenheit to Celsius: $\mathrm{\degree C = (\degree F - 32) \times \frac{5}{9}}$

• Celsius to Fahrenheit: $\mathrm{\degree F = 1.8 \times (\degree C) + 32}$

• Celsius to Kelvin: $\mathrm{K = \degree C + 273}$

• Kelvin to Celsius: $\mathrm{\degree C = K - 273}$

Example: Convert 100 °F to °C: $\mathrm{(100 - 32) \times \frac{5}{9} = 37.8\degree C}$

Metric Conversions & Scientific Notation

Unit Conversions and Scientific Notation

Chemistry uses the metric system and scientific notation to express measurements and quantities efficiently. Understanding how to convert units and compare numbers in scientific notation is crucial.

• Metric Prefixes: Examples include giga (G, $10^9$), kilo (k, $10^3$), milli (m, $10^{-3}$), micro (μ, $10^{-6}$).

• Example: 1 GL = $10^9$ L

• Scientific Notation: Move the decimal point to adjust the power of 10. For example, $4.5 \times 10^3$ is 4500.

• Comparing Sizes: Larger exponents indicate larger numbers.

Significant Figures

Rules for Counting Significant Figures

Significant figures reflect the precision of a measurement. Properly identifying and using significant figures is essential for calculations and reporting results.

• Leading zeros: Not significant (e.g., 0.003 has 1 sig fig).

• Captive zeros: Zeros between digits are significant (e.g., 205 has 3 sig figs).

• Trailing zeros: Significant only if a decimal is present (e.g., 50.0 has 3 sig figs; 400.0 has 4 sig figs).

Example: 30 (1 sig fig), 50.0 (3 sig figs), 400.0 (4 sig figs)

Laws of Chemistry

Fundamental Chemical Laws

Chemistry is governed by several foundational laws that describe the behavior of matter and reactions.

• Law of Conservation of Mass: Mass is neither created nor destroyed during a chemical reaction.

• Law of Multiple Proportions: Elements combine in small whole-number ratios to form compounds.

• Dalton’s Atomic Theory:

  • Elements are made of tiny particles called atoms.

  • Atoms of the same element are identical.

  • Atoms combine in simple whole-number ratios to form compounds.

  • Atoms cannot be created or destroyed in chemical reactions.

Density & Layering

Density and Floating Behavior

Density determines whether a substance will float or sink when placed in another substance. Lower density materials float on higher density materials.

• Order of Density (top to bottom):

  1. Wood (0.54 g/mL)

  2. Acetone (0.79 g/mL)

  3. Water (1.00 g/mL)

  4. Neoprene Rubber (1.33 g/mL)

  5. Osmium (22.4 g/mL)

Example: Wood floats on water because its density is lower.

Periodic Table Concepts

Groups, Diatomic Elements, and Isotopes

The periodic table organizes elements by properties and groups. Understanding these groups and special types of elements is essential.

• Groups:

  • Alkali metals: Group 1

  • Alkaline earth metals: Group 2

  • Halogens: Group 17

  • Noble gases: Group 18

• Diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2

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

Ions

Counting Subatomic Particles in Ions

Ions are atoms or molecules with a net charge due to loss or gain of electrons. Calculating the number of protons, neutrons, and electrons is a key skill.

• Protons: Equal to atomic number.

• Neutrons: Mass number minus atomic number.

• Electrons: Protons ± charge (add electrons for negative charge, subtract for positive).

• Example: N3− has 3 extra electrons compared to neutral nitrogen.

Moles and Molar Mass

Key Relationships and Calculations

The mole is a fundamental unit in chemistry, relating mass, number of particles, and molar mass. Calculations involving moles are central to chemical quantification.

• Key Equations:

  • $\mathrm{moles = \frac{mass}{molar\ mass}}$

  • $\mathrm{mass = moles \times molar\ mass}$

  • $\mathrm{particles = moles \times 6.022 \times 10^{23}}$

• Molar Mass: Sum of atomic masses from the periodic table.

Example: 18 g of water ($\mathrm{H_2O}$) is 1 mole.

Naming Ionic & Covalent Compounds

Rules for Naming Compounds

Compounds are named according to systematic rules. Ionic compounds consist of metals and nonmetals, while covalent compounds consist of nonmetals only.

• Ionic Compounds:

  • Metal name stays the same.

  • Nonmetal ends in "-ide".

  • Transition metals require Roman numerals for charge, except fixed-charge ions (Group 1, Group 2, Ag+, Zn2+, Al3+).

• Examples:

  • K2SO4 = potassium sulfate

  • Fe3P2 = iron (II) phosphide

• Polyatomic Ions: Common ions include:

  • CO32− carbonate

  • HCO3− hydrogen carbonate

  • NH4+ ammonium

  • ClO− hypochlorite

  • H2PO4− dihydrogen phosphate

Hydrates

Hydrate Compounds and Naming

Hydrates are compounds that include water molecules in their structure. The number of water molecules is indicated by a prefix and the formula.

• General Form: compound • nH2O

• Example: Copper (II) sulfate pentahydrate: CuSO4 • 5H2O

Chemical Properties of Metals

Physical and Chemical Properties

Metals exhibit characteristic properties that distinguish them from other elements.

• Malleable: Can be hammered into thin sheets.

• Ductile: Can be drawn into wires.

• Conduct heat: Efficient transfer of thermal energy.

• Conduct electricity: Free movement of electrons.

• Shiny/lustrous: Reflect light.

Example: Copper is both ductile and conducts electricity.

Empirical Formula Steps

Determining Empirical Formulas

The empirical formula represents the simplest whole-number ratio of elements in a compound. The process involves several steps.

1. Convert percent composition to grams (assume 100 g sample).

2. Convert grams to moles using molar mass.

3. Divide each mole value by the smallest number of moles.

4. Multiply to obtain whole numbers if necessary.

Example: Sucrose composition can be used to find the C:H:O ratio.

Types of Compounds

Ionic vs. Covalent Compounds

Compounds are classified based on the types of elements involved.

• Ionic: Metal + nonmetal (e.g., Ag3PO4, Be(NO3)2, MnO2)

• Covalent: Nonmetal + nonmetal (e.g., C8H16, N2, NF5)

Practice Problems to Prepare

Essential Problem Types

Mastery of chemistry requires practice with key problem types.

• Unit conversions

• Density and floating problems

• Percent composition calculations

• Moles ↔ grams ↔ molecules conversions

• Significant figures identification

• Naming ionic compounds and hydrates

• Writing formulas from compound names

Polyatomic Ion Table

Common Polyatomic Ions

Polyatomic ions are groups of atoms with a charge, commonly found in ionic compounds.

Ion	Name
CO32−	Carbonate
HCO3−	Hydrogen carbonate
NH4+	Ammonium
ClO−	Hypochlorite
H2PO4−	Dihydrogen phosphate

Additional info: More polyatomic ions can be found in standard tables, such as nitrate (NO3−), sulfate (SO42−), and phosphate (PO43−).