Chapter 1: Chemistry Basics - Matter and Measurement

Elements vs. Compounds

Understanding the distinction between elements and compounds is fundamental in chemistry.

• Element: A pure substance consisting of only one type of atom. Examples: Oxygen (O), Iron (Fe).

• Compound: A substance formed when two or more elements chemically combine in fixed ratios. Example: Water (H2O).

Characteristics of Elements in the Periodic Table

The Periodic Table organizes elements by increasing atomic number and groups them by similar properties.

• Groups (columns): Elements with similar chemical properties.

• Periods (rows): Elements with increasing atomic number.

• Metals, nonmetals, metalloids: Classified based on physical and chemical properties.

Periodic Table Organization

The arrangement reveals information about atomic structure and reactivity.

• Atomic number: Number of protons in the nucleus.

• Valence electrons: Determine chemical bonding behavior.

Conversions and Dimensional Analysis

Dimensional analysis is used to convert between units.

• Metric to metric: Example:

• English to metric: Example:

• Simple dosing calculations: Used in medical and laboratory settings.

Common Conversion Factors

Significant Figures and Scientific Notation

These concepts ensure precision and clarity in measurements.

• Significant figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Scientific notation: Expresses numbers as for easier handling of very large or small values.

• Operations: Rules for addition, subtraction, multiplication, and division with significant figures and scientific notation.

Accuracy vs. Precision

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

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

Percent Calculations

Percentages are used to express ratios and concentrations.

• Formula:

Density Calculations

Density is a physical property relating mass and volume.

• Formula:

• Units: Commonly g/mL or g/cm3

Physical vs. Chemical Changes

• Physical change: Alters form or appearance, not composition (e.g., melting ice).

• Chemical change: Produces new substances (e.g., burning wood).

Writing & Balancing Chemical Reactions

Chemical equations represent reactants and products; balancing ensures conservation of mass.

• Steps: Write formulas, count atoms, adjust coefficients.

• Example:

Chapter 2: Atoms and Radioactivity

Atomic Structure

Atoms consist of protons, neutrons, and electrons.

• Protons (p+): Positive charge, found in nucleus.

• Neutrons (n0): Neutral, found in nucleus.

• Electrons (e-): Negative charge, orbit nucleus.

Determining Number of Protons, Neutrons, Electrons

• Protons: Equal to atomic number.

• Neutrons:

• Electrons: Equal to protons in neutral atom; adjust for ions.

Allotropes

Allotropes are different structural forms of the same element.

• Example: Carbon as diamond, graphite, and graphene.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Characteristics: Same atomic number, different mass number.

• Example: and

Radioisotopes

Radioisotopes are unstable isotopes that emit radiation.

• Difference: Radioisotopes undergo radioactive decay; stable isotopes do not.

Half-life Problem Solving

Half-life is the time required for half of a radioactive sample to decay.

• Formula:

• n: Number of half-lives elapsed.

Chapter 3: Compounds - How Elements Combine

Valence Electrons and the Periodic Table

Valence electrons are the outermost electrons involved in bonding.

• Determined by: Group number for main group elements.

Why Atoms Bond

Atoms bond to achieve stable electron configurations, often following the octet rule.

• Octet rule: Atoms tend to have eight electrons in their valence shell.

• Duet rule: Applies to hydrogen and helium (two electrons).

Determining Ions: Cations vs. Anions

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Predicting charges: Based on group number and electron configuration.

Isoelectronic Species

Atoms or ions with the same number of electrons.

• Example: Na+ and Ne both have 10 electrons.

Recognizing Ionic Bonds & Compounds

Ionic bonds form between metals and nonmetals via electron transfer.

• Example: NaCl (sodium chloride)

Polyatomic Ions

Polyatomic ions are charged groups of covalently bonded atoms.

Naming & Writing Formulas of Ionic Compounds

• Combine cation and anion names. Example: Na+ + Cl- = NaCl (sodium chloride).

• Polyatomic ions: Use ion names and balance charges.

Recognizing Covalent Bonds & Molecules

Covalent bonds form between nonmetals via electron sharing.

• Example: H2O (water), CO2 (carbon dioxide)

Naming & Writing Formulas of Covalent Molecules

• Prefixes: mono-, di-, tri-, tetra-, etc. Example: CO2 is carbon dioxide.

Electron Dot Symbols and Lewis Structures

Electron dot symbols show valence electrons; Lewis structures depict bonding and lone pairs.

• Example: Lewis structure for H2O shows two bonds and two lone pairs on oxygen.

Electronegativity & Trends in the Periodic Table

Electronegativity is the tendency of an atom to attract electrons in a bond.

• Trend: Increases across a period, decreases down a group.

Determining Bond Polarity

Bond polarity depends on the difference in electronegativity between atoms.

• Nonpolar: Small or no difference.

• Polar: Significant difference.

Drawing Molecular Shape (VSEPR Theory)

VSEPR theory predicts molecular geometry based on electron pair repulsion.

• Example: H2O is bent; CO2 is linear.

Determining Overall Molecular Polarity

Overall polarity depends on bond polarity and molecular shape.

• Example: H2O is polar; CO2 is nonpolar.

Additional info: Academic context and examples were added to clarify brief points and ensure completeness.