CHEM 1001 Midterm Study Guide: Chapters 1–4 (GOB Chemistry)

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Chapter 1: Chemistry Basics – Matter and Measurement

Classifying Matter

Understanding the classification of matter is fundamental in chemistry. Matter can be categorized as pure substances (elements and compounds) or mixtures (homogeneous and heterogeneous).

Pure Substances: Consist of a single type of particle; elements contain only one kind of atom, while compounds contain two or more elements chemically bonded.
Mixtures: Physical combinations of substances; homogeneous mixtures have uniform composition, while heterogeneous mixtures do not.
Example: Salt water is a homogeneous mixture; sand and iron filings form a heterogeneous mixture.

The Periodic Table

The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties.

Periods: Horizontal rows; indicate energy levels.
Groups: Vertical columns; elements in the same group have similar properties.
Metals, Nonmetals, Metalloids: Classified based on physical and chemical properties.
Example: Sodium (Na) is a metal in Group 1; chlorine (Cl) is a nonmetal in Group 17.

Significant Figures and Scientific Notation

Accurate measurement and reporting in chemistry require understanding significant figures and scientific notation.

Significant Figures: Digits that carry meaning in a measurement.
Scientific Notation: Expresses numbers as a product of a coefficient and a power of ten.
Example: 0.00456 has three significant figures; written as in scientific notation.

Physical and Chemical Properties

Properties of matter are classified as physical (observed without changing composition) or chemical (describe ability to undergo chemical change).

Physical Properties: Color, state, melting point, boiling point.
Chemical Properties: Reactivity, flammability, oxidation.

Chapter 2: Atoms and Radioactivity

Atomic Structure

Atoms are the basic units of matter, composed of protons, neutrons, and electrons.

Atomic Number (Z): Number of protons in the nucleus; defines the element.
Mass Number (A): Sum of protons and neutrons.
Isotopes: Atoms of the same element with different numbers of neutrons.
Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Atomic Symbols and Notation

Atomic symbols represent elements and their isotopes.

Notation: , where X is the element symbol, A is mass number, Z is atomic number.
Example: for Carbon-14.

Radioactivity

Some atoms are unstable and emit radiation to become more stable.

Types of Radiation: Alpha (), beta (), and gamma () radiation.
Half-life: Time required for half of a radioactive sample to decay.
Example: Uranium-238 undergoes alpha decay.

Periodic Trends

Elements exhibit trends in properties across periods and groups.

Atomic Radius: Increases down a group, decreases across a period.
Ionization Energy: Energy required to remove an electron; increases across a period.
Electronegativity: Tendency to attract electrons; increases across a period.

Chapter 3: Compounds – How Elements Combine

Electron Arrangement and the Octet Rule

Atoms combine to achieve stable electron configurations, often following the octet rule (eight electrons in the valence shell).

Electron Shells: Energy levels where electrons reside.
Valence Electrons: Electrons in the outermost shell; involved in bonding.
Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

Ionic and Covalent Bonding

Compounds form through ionic or covalent bonding.

Ionic Bonds: Transfer of electrons from metals to nonmetals; forms ions.
Covalent Bonds: Sharing of electrons between nonmetals.
Example: Sodium chloride (NaCl) is ionic; water (H2O) is covalent.

Lewis Structures and Molecular Geometry

Lewis structures depict the arrangement of electrons in molecules. Molecular geometry describes the 3D shape of molecules.

Lewis Structure: Shows bonding and lone pairs.
VSEPR Theory: Predicts molecular shapes based on electron pair repulsion.
Example: Methane (CH4) is tetrahedral.

Polarity and Intermolecular Forces

Polarity arises from differences in electronegativity; affects physical properties and interactions.

Polar Molecules: Unequal sharing of electrons; have dipole moments.
Intermolecular Forces: Include hydrogen bonding, dipole-dipole, and London dispersion forces.

Homogeneity and Molecular Polarity

Homogeneity refers to uniform composition; molecular polarity affects solubility and reactivity.

Example: Water is polar and forms hydrogen bonds; oil is nonpolar.

Chapter 4: Introduction to Organic Compounds

Representing Organic Compounds

Organic compounds can be represented in various ways: Lewis, condensed, and skeletal structures.

Lewis Structure: Shows all atoms and bonds.
Condensed Formula: Groups atoms together (e.g., CH3CH2OH).
Skeletal Structure: Lines represent bonds; vertices represent carbon atoms.

Hydrocarbons: Saturated and Unsaturated

Hydrocarbons are compounds of carbon and hydrogen; classified as saturated (alkanes) or unsaturated (alkenes, alkynes).

Saturated Hydrocarbons: Only single bonds (alkanes).
Unsaturated Hydrocarbons: Double or triple bonds (alkenes, alkynes).
Example: Ethane (C2H6) is saturated; ethene (C2H4) is unsaturated.

Alkanes and Cycloalkanes

Alkanes are straight-chain or branched hydrocarbons; cycloalkanes are ring-shaped.

General Formula for Alkanes:
General Formula for Cycloalkanes:
Example: Cyclohexane (C6H12).

Functional Groups

Functional groups are specific groups of atoms that impart characteristic properties to organic molecules.

Common Functional Groups: Alcohols (-OH), carboxylic acids (-COOH), amines (-NH2).
Example: Ethanol contains an alcohol group.

Drawing and Naming Organic Compounds

Organic compounds are named using IUPAC rules; structures can be drawn in various formats.

Structural Isomers: Compounds with the same molecular formula but different structures.
Example: Butane and isobutane (C4H10).