General Chemistry Study Notes: Measurement, Atoms, Molecules, and Mass Relationships

Chemical Tools: Experimentation & Measurement

Scientific Notation

Scientific notation is used to express very large or very small numbers in a concise form. It is essential for reporting measurements in chemistry.

• Large numbers (e.g., 50,000,000): Move decimal left; exponent is positive.

• Small numbers (e.g., 0.000045): Move decimal right; exponent is negative.

Example: 0.00032 = $3.2 \times 10^{-4}$

Significant Figures

Significant figures reflect the precision of a measurement. All non-zero digits are significant. Zeros between non-zero digits and trailing zeros after a decimal point are also significant.

• Leading zeros are not significant.

• Trailing zeros in a number with a decimal point are significant.

Example: 0.04050 has four significant figures.

Accuracy and Precision

• Accuracy: How close a measured value is to the true value.

• Precision: How close repeated measurements are to each other.

Metric Conversions and SI Units

Chemistry uses the International System of Units (SI) for measurements. Prefixes indicate powers of ten.

Prefix	Symbol	Meaning	Factor
kilo	k	thousand	$10^3$
centi	c	hundredth	$10^{-2}$
milli	m	thousandth	$10^{-3}$
micro	\mu	millionth	$10^{-6}$
nano	n	billionth	$10^{-9}$
pico	p	trillionth	$10^{-12}$

Example: 1 km = 1000 m; 1 cm = 0.01 m

Dimensional Analysis

Dimensional analysis (factor-label method) is used to convert between units and solve problems.

• Set up conversion factors so units cancel appropriately.

• Always include units in calculations.

Example: Convert 25 cm to meters: $25\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.25\ \text{m}$

Atoms, Molecules & Ions

Elements, Compounds, and Mixtures

• Element: Substance that cannot be broken down into simpler substances.

• Compound: Substance composed of two or more elements chemically combined in fixed ratios.

• Mixture: Combination of two or more substances physically blended, not chemically bonded.

Atomic Structure

• Proton (p+): Positive charge, located in nucleus.

• Neutron (n0): No charge, located in nucleus.

• Electron (e-): Negative charge, located outside nucleus.

Atomic Number (Z): Number of protons in the nucleus; defines the element.

Mass Number (A): Total number of protons and neutrons.

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

Periodic Table Organization

• Elements are arranged by increasing atomic number.

• Groups (columns) share similar chemical properties.

• Metals are on the left, nonmetals on the right, metalloids along the staircase line.

Metals, Nonmetals, and Metalloids

• Metals: Good conductors, malleable, ductile, shiny, usually solid at room temperature.

• Nonmetals: Poor conductors, brittle, dull, can be solid, liquid, or gas.

• Metalloids: Properties intermediate between metals and nonmetals.

Formation of Ions

Main group elements form ions to achieve noble gas configurations.

Group	Typical Ion	How to Determine
1 (alkali metals)	+1	Lose 1 electron
2 (alkaline earth metals)	+2	Lose 2 electrons
16 (chalcogens)	-2	Gain 2 electrons
17 (halogens)	-1	Gain 1 electron
18 (noble gases)	0	Stable, rarely form ions

Common Polyatomic Ions

Name	Formula	Charge
Ammonium	NH4+	+1
Nitrate	NO3-	-1
Sulfate	SO42-	-2
Phosphate	PO43-	-3
Hydroxide	OH-	-1
Cyanide	CN-	-1

Naming Compounds

• Ionic compounds: Name cation first, then anion. Use Roman numerals for transition metals with variable charge.

• Molecular compounds: Use prefixes to indicate number of atoms (mono-, di-, tri-, etc.).

Example: CO2 is carbon dioxide; FeCl3 is iron(III) chloride.

Mass Relationships in Chemical Reactions

The Mole Concept

The mole is the SI unit for amount of substance. One mole contains Avogadro's number ($6.022 \times 10^{23}$) of particles.

• Molar mass: Mass of one mole of a substance (g/mol).

Example: 1 mole of H2O = 18.02 g

Stoichiometry

Stoichiometry involves the calculation of reactants and products in chemical reactions using balanced equations.

• Coefficients in balanced equations represent the number of moles of each substance.

• Use molar mass to convert between grams and moles.

Example: $2H_2 + O_2 \rightarrow 2H_2O$

Percent Composition

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

• Calculate molar mass of compound.

• Divide mass of each element by total molar mass, multiply by 100%.

Example: In H2O, percent H = $(2.02/18.02) \times 100 = 11.2\%$

Empirical and Molecular Formulas

• Empirical formula: Simplest whole-number ratio of atoms in a compound.

• Molecular formula: Actual number of atoms of each element in a molecule.

Example: Empirical formula of C6H12O6 is CH2O.

Solution Concentration: Molarity

Molarity (M) is the number of moles of solute per liter of solution.

$M = \frac{\text{moles of solute}}{\text{liters of solution}}$

Example: 0.5 mol NaCl in 1 L solution = 0.5 M

Solution Dilution

To dilute a solution, use the equation:

$M_1V_1 = M_2V_2$

Where $M_1$ and $V_1$ are the initial molarity and volume, $M_2$ and $V_2$ are the final molarity and volume.

Stoichiometry with Solutions

• Use molarity and volume to find moles of solute.

• Apply stoichiometric coefficients from balanced equations to relate reactants and products.

Percent Yield

Percent yield compares the actual yield to the theoretical yield:

$\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100$

Types of Chemical Reactions

• Combination (synthesis): Two or more substances form one product.

• Decomposition: One substance breaks into two or more products.

• Single replacement: One element replaces another in a compound.

• Double replacement: Exchange of ions between two compounds.

• Combustion: Substance reacts with oxygen, producing heat and light.

Balancing Chemical Equations

• Write correct formulas for all reactants and products.

• Balance atoms one element at a time using coefficients.

• Check that all atoms are balanced.

Example: $2H_2 + O_2 \rightarrow 2H_2O$

Lab Techniques and Procedures

• Use dimensional analysis for unit conversions.

• Apply significant figures in calculations.

• Understand the use of laboratory glassware for measuring volume and mass.

Additional info: These notes cover foundational concepts from Chapters 1–3 and include bonus material on lab techniques and mathematical operations, as outlined in the General Chemistry curriculum.