General Chemistry 1007: Essential Concepts, Conversions, and Equations

Unit Conversions

Temperature Conversions

Temperature can be measured in Fahrenheit (°F), Celsius (°C), and Kelvin (K). Converting between these units is essential in chemistry calculations.

• Fahrenheit to Celsius: $\text{temp in } ^\circ\mathrm{F} = \frac{9}{5} \times (\text{temp in } ^\circ\mathrm{C}) + 32^\circ\mathrm{F}$

• Celsius to Fahrenheit: $\text{temp in } ^\circ\mathrm{C} = (\text{temp in } ^\circ\mathrm{F} - 32^\circ\mathrm{F}) \times \frac{5}{9}$

• Celsius to Kelvin: $K = ^\circ\mathrm{C} + 273.15$

Example: Convert 25°C to Fahrenheit: $\text{temp in } ^\circ\mathrm{F} = \frac{9}{5} \times 25 + 32 = 77^\circ\mathrm{F}$

Specific Heat of Water

Specific heat (SH) is the amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• For H2O(l): $\text{SH} = \frac{1.00\ \text{cal}}{\text{g}^\circ\mathrm{C}} = \frac{4.184\ \text{J}}{\text{g}^\circ\mathrm{C}}$

Useful Conversion Factors

Conversion factors are used to change units in calculations. Memorizing common factors is helpful for problem-solving.

Quantity	Conversion
1 lb	453.6 g
1 in	2.54 cm (exactly)
1 mi	1.609 km
1 km	0.621 mi
1 L	1.0567 qt
1 calorie	4.184 J

Equations

Density

Density is a physical property defined as mass per unit volume.

• Formula: $D = \frac{\text{grams}}{\text{mL}}$

• Example: If a sample has a mass of 10 g and a volume of 2 mL, $D = \frac{10}{2} = 5\ \text{g/mL}$

Molecular Weight (MW)

Molecular weight is the mass of one mole of a substance, expressed in grams per mole.

• Formula: $MW = \frac{\text{grams}}{\text{mole}}$

• Application: Used to convert between mass and moles in stoichiometric calculations.

Summation of Masses

The total mass of a system is the sum of the masses of its components.

• Formula: $m_t = a_1m_1 + a_2m_2 + \ldots + a_n m_n$

• Example: For a mixture with 2 components, $m_t = a_1m_1 + a_2m_2$

Constants

Avogadro's Number

Avogadro's number is the number of particles (atoms, molecules, ions) in one mole of a substance.

• Value: $N_A = 6.022 \times 10^{23}$

• Application: Used to convert between moles and number of particles.

Core Concepts in General Chemistry

What is a Chemical?

A chemical is a substance with a definite composition and properties, which can be an element or a compound. Chemistry studies the properties, composition, and changes of matter.

• Examples: Water (H2O), sodium chloride (NaCl), oxygen (O2).

Significant Figures and Scientific Notation

Significant figures reflect the precision of a measurement. Scientific notation expresses numbers as a product of a coefficient and a power of ten.

• Example: 0.00456 has 3 significant figures; written in scientific notation: $4.56 \times 10^{-3}$

Measurements and Units

Chemistry uses the metric system for measurements. Common units include grams (g), liters (L), meters (m), and moles (mol).

• Mass: Measured in grams (g)

• Volume: Measured in liters (L)

• Length: Measured in meters (m)

Physical vs. Chemical Changes

Physical changes alter the form of a substance but not its identity. Chemical changes result in the formation of new substances.

• Physical Change Example: Melting ice

• Chemical Change Example: Rusting of iron

Classification of Matter

Matter can be classified as elements, compounds, or mixtures.

• Element: Pure substance made of one type of atom (e.g., O2)

• Compound: Substance made of two or more elements chemically combined (e.g., H2O)

• Mixture: Physical blend of two or more substances (e.g., air)

Periodic Table Organization

The periodic table arranges elements by increasing atomic number. Groups (columns) share similar properties; periods (rows) indicate energy levels.

• Group 1: Alkali metals

• Group 2: Alkaline earth metals

• Group 17: Halogens

• Group 18: Noble gases

Atomic Structure

Atoms consist of protons, neutrons, and electrons. The atomic number equals the number of protons; the mass number is the sum of protons and neutrons.

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

Formulas and Nomenclature

Chemical formulas represent the composition of compounds. Nomenclature is the system for naming chemical substances.

• Example: NaCl is sodium chloride.

Percent Composition

Percent composition expresses the mass percentage of each element in a compound.

• Formula: $\%\ \text{composition} = \frac{\text{mass of element}}{\text{total mass of compound}} \times 100$

Empirical and Molecular Formulas

The empirical formula shows the simplest whole-number ratio of elements; the molecular formula shows the actual number of atoms in a molecule.

• Example: Glucose: Empirical formula CH2O, Molecular formula C6H12O6

Accuracy and Precision

Accuracy refers to how close a measurement is to the true value; precision refers to how reproducible measurements are.

• Example: Weighing a sample multiple times and getting similar results shows precision.

Homogeneous vs. Heterogeneous Mixtures

Homogeneous mixtures have uniform composition; heterogeneous mixtures have visibly different parts.

• Homogeneous Example: Saltwater

• Heterogeneous Example: Salad

Metric Conversions and Dimensional Analysis

Dimensional analysis uses conversion factors to solve problems involving unit changes.

• Example: Convert 5.0 inches to centimeters: $5.0\ \text{in} \times \frac{2.54\ \text{cm}}{1\ \text{in}} = 12.7\ \text{cm}$

Significant Figures in Calculations

When performing calculations, the result should reflect the correct number of significant figures based on the input values.

• Multiplication/Division: Use the least number of significant figures.

• Addition/Subtraction: Use the least number of decimal places.

Physical States and Phase Changes

Matter exists as solids, liquids, or gases. Phase changes include melting, freezing, boiling, and condensation.

• Example: Water boiling at 100°C changes from liquid to gas.

Atomic Mass and Moles

The mole is a counting unit in chemistry, relating mass to number of particles via Avogadro's number.

• Formula: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

Physical and Chemical Properties

Physical properties can be observed without changing the substance's identity; chemical properties describe how a substance reacts.

• Physical Property Example: Density

• Chemical Property Example: Reactivity with acid

Accuracy in Experimental Practice

Proper measurement techniques and understanding of error sources are crucial for reliable results in chemistry.

• Example: Calibrating balances and using proper glassware.

Summary Table: Key Conversion Factors

Unit	Equivalent
1 lb	453.6 g
1 in	2.54 cm
1 mi	1.609 km
1 km	0.621 mi
1 L	1.0567 qt
1 calorie	4.184 J

Skills You Should Master

• Metric conversions in units

• Translating chemical formulas and equations

• Classification of elements, compounds, and mixtures

• Reading and using the periodic table

• Dimensional analysis (unit conversions)

• Significant figures and scientific notation

• Identifying physical and chemical changes

• Calculating percent composition and empirical/molecular formulas

• Understanding physical states and phase changes

Additional info: Some context and examples have been expanded for clarity and completeness based on standard General Chemistry curriculum.