Chem 101 Exam 1 Study Guide: Chapters 1–5 (GOB Chemistry)

Chapter 1: Introduction to Chemistry

Classification of Matter

Understanding the types of matter is fundamental in chemistry. Matter can be classified based on its composition and uniformity.

• Heterogeneous Mixture: A mixture with non-uniform composition; different parts can be distinguished. Example: Salad, granite.

• Homogeneous Mixture: A mixture with uniform composition throughout; also called a solution. Example: Saltwater, air.

Elements, Atoms, and Compounds

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

• Atom: The smallest unit of an element that retains its chemical properties.

• Compound: A substance formed from two or more elements chemically bonded in fixed proportions. Example: Water (H2O).

Subatomic Particles

• Proton: Charge = +1; Relative mass ≈ 1 amu; Located in nucleus.

• Neutron: Charge = 0; Relative mass ≈ 1 amu; Located in nucleus.

• Electron: Charge = -1; Relative mass ≈ 0.0005 amu; Located in electron cloud around nucleus.

States of Matter

• Solid: Definite shape and volume; particles closely packed; vibrate in place.

• Liquid: Definite volume, no definite shape; particles less tightly packed; flow freely.

• Gas: No definite shape or volume; particles far apart; move freely.

Comparison Table: Properties of States of Matter

State	Shape	Volume	Particle Arrangement	Flow
Solid	Definite	Definite	Close, orderly	No
Liquid	Indefinite	Definite	Close, random	Yes
Gas	Indefinite	Indefinite	Far apart	Yes

Physical vs. Chemical Changes

• Physical Change: Change in state or appearance without altering composition. Examples: Melting ice, dissolving sugar.

• Chemical Change: Change that produces new substances. Examples: Rusting iron, burning wood.

Periodic Table Basics

• First 20 Elements: Know names and symbols (H, He, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca) plus Fe, Ag, Au, Cu.

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

• Periods: Horizontal rows; elements arranged by increasing atomic number.

• Families: Group 1A (Alkali metals), 2A (Alkaline earth metals), 7A (Halogens), 8A (Noble gases).

• Metals: Located on the left and center; Non-metals: Located on the right.

Chapter 2: Measurement in Science and Medicine

Metric and SI Units

Scientific measurements use standardized units for consistency.

• Base Units: Length (meter, m), Mass (kilogram, kg), Time (second, s), Temperature (kelvin, K), Amount (mole, mol).

• Prefixes: Used to express large/small quantities. Examples: nano (n, 10-9), milli (m, 10-3), centi (c, 10-2), deca (da, 101), kilo (k, 103).

Scientific Notation

• Expresses numbers as a product of a coefficient and a power of ten. Example: 0.00045 = $4.5 \times 10^{-4}$

• Convert between standard and scientific notation as needed.

Significant Figures

• Significant Figures (Sig Figs): Digits that reflect the precision of a measurement.

• Rules for counting sig figs:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros are significant only if there is a decimal point.

• In calculations:

  • Multiplication/Division: Answer has same number of sig figs as the least precise measurement.

  • Addition/Subtraction: Answer has same number of decimal places as the least precise measurement.

Dimensional Analysis

• Method for converting between units using conversion factors.

• Example: To convert 5.0 cm to meters: $5.0 \text{ cm} \times \frac{1 \text{ m}}{100 \text{ cm}} = 0.050 \text{ m}$

Volume, Mass, and Density

• Volume: Amount of space occupied; SI unit is liter (L) or cubic meter (m3).

• Mass: Amount of matter; SI unit is kilogram (kg).

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

Temperature Conversions

• Convert between Fahrenheit and Celsius:

  • $\text{C} = \frac{5}{9}(\text{F} - 32)$

  • $\text{F} = \frac{9}{5}\text{C} + 32$

Chapter 3: Atoms: The Building Blocks of Chemistry

Atomic Structure

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

• Mass Number (A): Total number of protons and neutrons; $A = Z + \text{number of neutrons}$

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

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

Electron Structure

• Electrons are arranged in energy levels, sublevels, and orbitals.

• Box Orbital Diagram: Visual representation of electron arrangement in orbitals.

• Electron Configuration: Notation showing distribution of electrons among orbitals. Example: Oxygen: 1s2 2s2 2p4

• Abbreviated Electron Configuration: Uses noble gas core. Example: Sodium: [Ne] 3s1

Valence Electrons and Electron Dot Diagrams

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

• Can be found from periodic table group number or electron configuration.

• Electron Dot Configuration: Dots around element symbol represent valence electrons.

Periodic Trends

• Atomic Size: Increases down a group, decreases across a period.

• Ionization Energy: Energy required to remove an electron; decreases down a group, increases across a period.

Chapter 4: Nuclear Chemistry

Radioactive Decay Processes

• Alpha Decay (α): Emission of an alpha particle ($^4_2\text{He}$); decreases mass number by 4, atomic number by 2.

• Beta Decay (β): Emission of a beta particle (electron, $^0_{-1}\text{e}$); increases atomic number by 1.

• Positron Emission (β+): Emission of a positron ($^0_{+1}\text{e}$); decreases atomic number by 1.

• Gamma Decay (γ): Emission of gamma radiation; no change in mass or atomic number.

Balancing Nuclear Equations

• Sum of mass numbers and atomic numbers must be equal on both sides.

• New isotopes are formed depending on the decay process.

Radiation Dose and Units

• Curie (Ci): Measures radioactivity.

• Rad and Gray: Measure absorbed dose (energy deposited per mass).

Half-Life

• Time required for half the nuclei in a sample to decay.

• Half-life calculation: $N = N_0 \times (\frac{1}{2})^{t/T_{1/2}}$

Medical Radioisotopes

• Most used: Technetium-99m (Tc-99m).

• PET scans: Use Fluorine-18 (F-18).

Nuclear Fission and Fusion

• Fission: Splitting of heavy nuclei (e.g., uranium-235); initiated by neutron absorption; produces chain reactions; used in nuclear power.

• Fusion: Combining of light nuclei (e.g., hydrogen isotopes); powers the sun; requires high temperature and pressure.

Chapter 5: Ionic Compounds (through 5.3)

Valence Electrons and Bonding

• Number of valence electrons determines how atoms bond and whether they gain or lose electrons.

• Bonding: Atoms combine to achieve stable electron configurations.

• Types of Bonds:

  • Ionic Bond: Formed by transfer of electrons from metal to non-metal.

  • Covalent Bond: Formed by sharing electrons between non-metals. (Covered in later chapters)

Ion Formation and Charges

• Atoms gain or lose electrons to form ions with full outer shells (octet rule).

• Metals: Lose electrons, form positive ions (cations).

• Non-metals: Gain electrons, form negative ions (anions).

• Charge of Ion: $\text{Charge} = \text{Number of protons} - \text{Number of electrons}$

Ionic Compounds

• Formed from cations and anions; overall charge is zero (neutral).

• Formula: Ratio of ions balances charges.

• Polyvalent Elements: Some metals (e.g., Fe, Cu) can form ions with multiple charges; specified using Roman numerals (e.g., Fe2+ is iron(II)).

• Determining Ion Charges from Formulas: For CuCl2, Cl is -1, so Cu must be +2.

Table: Common Polyvalent Metals and Their Charges

Element	Possible Charges	Notation
Iron (Fe)	+2, +3	Iron(II), Iron(III)
Copper (Cu)	+1, +2	Copper(I), Copper(II)
Lead (Pb)	+2, +4	Lead(II), Lead(IV)
Tin (Sn)	+2, +4	Tin(II), Tin(IV)

Additional info: Some explanations and tables have been expanded for clarity and completeness, including the table of states of matter and common polyvalent metals.