GOB Chemistry Exam 1 Study Guide

Classifying Matter

Pure Substances and Mixtures

Matter can be classified based on its composition and uniformity. Understanding these classifications is fundamental in chemistry.

• Pure Substance: A form of matter with a constant composition and distinct chemical properties. Examples include elements (e.g., O2, Fe) and compounds (e.g., H2O, NaCl).

• Mixture: A physical blend of two or more substances, each retaining its own properties. Mixtures can be separated by physical means.

• Homogeneous Mixture: Uniform composition throughout (e.g., saltwater).

• Heterogeneous Mixture: Non-uniform composition (e.g., salad, sand in water).

Example: Air is a homogeneous mixture; salad is a heterogeneous mixture.

Elements, Compounds, and the Periodic Table

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

• Element: A pure substance consisting of only one type of atom.

• Compound: A pure substance composed of two or more elements chemically combined in fixed ratios.

• Groups: Vertical columns; elements in a group have similar chemical properties.

• Periods: Horizontal rows; elements in a period have the same number of electron shells.

• Metals: Located on the left and center; typically shiny, malleable, and good conductors.

• Nonmetals: Located on the right; often gases or brittle solids, poor conductors.

Example: Sodium (Na) is a metal in Group 1; chlorine (Cl) is a nonmetal in Group 17.

How Matter Changes

Physical and Chemical Changes

Matter can undergo changes that are either physical or chemical in nature.

• Physical Change: Alters the form but not the composition (e.g., melting ice).

• Chemical Change (Reaction): Alters the composition, forming new substances (e.g., burning wood).

• States of Matter: Solid, liquid, gas.

• Chemical Equation: Represents a chemical reaction using symbols and formulas.

Example: $2H_2 + O_2 \rightarrow 2H_2O$ (formation of water from hydrogen and oxygen).

Math in Chemistry

Units and Measurements

Accurate measurement and unit conversion are essential in chemistry.

• Base Units: Meter (distance), gram (mass), liter (volume), Kelvin/Celsius (temperature).

• Metric Prefixes: kilo- (103), centi- (10-2), milli- (10-3).

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

• Scientific Notation: Expresses numbers as $a \times 10^n$.

Example: 0.00045 = $4.5 \times 10^{-4}$

The "Stuff" of Chemistry

Mass, Volume, and Density

Understanding the properties of matter involves measuring mass, volume, and density.

• Mass: Amount of matter in an object (measured in grams).

• Volume: Space occupied by matter (measured in liters or milliliters).

• Density: Mass per unit volume. $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Specific Gravity: Ratio of the density of a substance to the density of water.

Example: If a block has a mass of 10 g and a volume of 2 mL, its density is $5\,\text{g/mL}$.

Temperature Scales

• Celsius (°C), Kelvin (K), Fahrenheit (°F): Common temperature scales.

• Conversions:

  • $K = ^\circ C + 273.15$

  • $^\circ F = (1.8 \times ^\circ C) + 32$

Kinetic and Potential Energy

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position or composition.

Example: A rolling ball has kinetic energy; a stretched bow has potential energy.

States of Matter

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes shape of container.

• Gas: No definite shape or volume.

Measuring Matter

Accuracy and Precision

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

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

Unit Conversions

• Convert between SI (metric) and U.S. customary units using provided conversion factors.

• Apply conversion factors, drop units, and convert percentages in health-related measurements.

Atoms and Their Components

Subatomic Particles

• Proton: Positively charged particle in the nucleus.

• Neutron: Neutral particle in the nucleus.

• Electron: Negatively charged particle outside the nucleus.

Example: Carbon-12 has 6 protons, 6 neutrons, and 6 electrons.

Atomic Number and Mass Number

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

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

Example: $A = Z + N$ (where N = number of neutrons)

Symbolic Notation

• Written as $^{A}_{Z}\text{X}$, where X is the element symbol.

Example: $^{14}_{6}\text{C}$ for carbon-14.

Isotopes and Atomic Mass

Isotopes

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

• Mass Number: Number of protons + neutrons in a specific isotope.

• Atomic Mass: Weighted average mass of all isotopes of an element.

Example: $^{12}\text{C}$ and $^{14}\text{C}$ are isotopes of carbon.

Radioactivity and Radioisotopes

Radioactivity

• Radioactivity: The spontaneous emission of particles or energy from unstable atomic nuclei.

• Ionizing Radiation: Includes alpha, beta, and gamma radiation.

• Penetrating Power: Gamma > Beta > Alpha.

Example: Alpha particles can be stopped by paper; gamma rays require lead shielding.

Radiation Units and Half-Lives

• Activity Units: Curie (Ci), Becquerel (Bq).

• Half-Life: Time required for half the atoms in a sample to decay.

Half-life formula:

$N = N_0 \left(\frac{1}{2}\right)^{t/t_{1/2}}$

where $N$ = remaining amount, $N_0$ = initial amount, $t$ = elapsed time, $t_{1/2}$ = half-life.

Medical Applications

• Radioisotopes are used in diagnosis (e.g., PET scans) and treatment (e.g., cancer radiotherapy).

Electron Arrangements and the Octet Rule

Electron Shells and Valence Electrons

• Electrons are arranged in shells (energy levels) around the nucleus.

• Valence Electrons: Electrons in the outermost shell; determine chemical reactivity.

• Main-group elements in the first four periods follow predictable patterns for valence electrons.

The Octet Rule

• Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (octet), leading to stability.

• Three ways to reach stability:

  • Give electrons (form cations)

  • Take electrons (form anions)

  • Share electrons (form covalent bonds)

Example: Sodium gives up one electron to achieve an octet; chlorine takes one electron.

Ion Formation and Ionic Compounds

Ion Formation

• Application of the octet rule to form ions.

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

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

• Predict ionic charge using the periodic table (e.g., Group 1: +1, Group 17: -1).

• Write ion symbols: e.g., $\text{Na}^+$, $\text{Cl}^-$.

• Common polyatomic ions: phosphate ($\text{PO}_4^{3-}$), carbonate ($\text{CO}_3^{2-}$), hydroxide ($\text{OH}^-$), hydronium ($\text{H}_3\text{O}^+$).

Ionic Compounds

• Formed by the combination of cations and anions.

• Formula reflects the ratio of ions needed for charge neutrality.

• Transition metals may have variable charges, indicated by Roman numerals (e.g., Fe2+ is iron(II)).

Example: $\text{Na}^+ + \text{Cl}^- \rightarrow \text{NaCl}$

Covalent Compounds and Bonding

Covalent Bond Formation

• Atoms share electrons to achieve an octet.

• Nonmetals typically form covalent bonds.

• Number of bonds relates to the number of valence electrons (e.g., oxygen forms two bonds).

• Lewis structures represent bonding and lone pairs.

Example: $\text{H}_2\text{O}$: oxygen shares electrons with two hydrogens.

Ionic vs. Covalent Compounds

• Ionic: Metal + nonmetal, transfer of electrons, high melting points, conduct electricity when dissolved.

• Covalent: Nonmetal + nonmetal, sharing of electrons, lower melting points, do not conduct electricity.

The Mole: Counting Atoms and Compounds

• Mole (mol): Amount of substance containing $6.022 \times 10^{23}$ particles (Avogadro's number).

• Relates mass, number of particles, and volume (for gases).

Example: 1 mol of H2O contains $6.022 \times 10^{23}$ molecules.

Shapes and Polarity of Covalent Compounds

Molecular Shape (VSEPR Theory)

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

• Common shapes: linear, bent, trigonal planar, tetrahedral.

• Wedge-and-dash notation shows 3D arrangement.

Example: Methane (CH4) is tetrahedral.

Electronegativity and Molecular Polarity

• Electronegativity: Tendency of an atom to attract electrons in a bond.

• Bond Polarity: Difference in electronegativity leads to polar bonds.

• Molecular Polarity: Determined by bond polarities and molecular shape.

Example: Water (H2O) is polar due to its bent shape and polar O-H bonds.

Summary Table: Classification of Matter

Type	Definition	Example
Element	Pure substance, one type of atom	O2, Fe
Compound	Pure substance, two or more elements chemically combined	H2O, NaCl
Homogeneous Mixture	Uniform composition	Saltwater, air
Heterogeneous Mixture	Non-uniform composition	Salad, sand in water