Chem 101 Exam 1 Study Guide: Chapters 1–5.3 (Atoms, Matter, Measurement, Nuclear Chemistry, and Ionic Compounds)

Chapter 1: Introduction to Chemistry and Matter

Classification of Matter

• Heterogeneous Mixtures: Mixtures with visibly different components or phases (e.g., salad, oil and water).

• Homogeneous Mixtures (Solutions): Mixtures with uniform composition throughout (e.g., saltwater, air).

• Elements: Pure substances consisting of only one type of atom (e.g., O2, Fe).

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

• Compounds: Substances composed of two or more elements chemically combined in fixed ratios (e.g., H2O).

Structure of the Atom

• Subatomic Particles:

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

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

  • Electron: Charge = –1; Relative mass ≈ 1/1836 amu; Located in electron cloud outside nucleus.

States of Matter

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

• Liquids: Definite volume, no definite shape; particles close but can flow past each other.

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

Physical vs. Chemical Changes

• Physical Change: Change in state or appearance without altering composition (e.g., melting ice, dissolving sugar).

• Chemical Change: Change that produces new substances (e.g., rusting iron, burning wood).

Elements and the Periodic Table

• Know symbols and names for the first 20 elements and Fe, Ag, Au, Cu.

• Groups (Columns): Elements with similar properties; numbered 1A–8A.

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

• Families:

  • Group 1A: Alkali metals

  • Group 2A: Alkaline earth metals

  • Group 7A: Halogens

  • Group 8A: Noble gases

• Metals: Left and center of table; conduct electricity, malleable, ductile.

• Non-metals: Right side; poor conductors, brittle in solid form.

Chapter 2: Measurement in Science and Medicine

Units and Prefixes

• SI Base Units: Meter (m), kilogram (kg), second (s), mole (mol), kelvin (K), ampere (A), candela (cd).

• Metric Prefixes:

  • kilo- (k): $10^3$

  • deca- (da): $10^1$

  • centi- (c): $10^{-2}$

  • milli- (m): $10^{-3}$

  • micro- (μ): $10^{-6}$

  • nano- (n): $10^{-9}$

Scientific Notation

• Express numbers as $a \times 10^n$ (e.g., $3.2 \times 10^4$).

• Convert between standard and scientific notation as needed.

Significant Figures

• Indicate the precision of a measurement.

• Rules for counting significant figures:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

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

• In calculations:

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

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

Dimensional Analysis and Conversions

• Use conversion factors to change units.

• Set up calculations so units cancel appropriately.

Volume, Mass, and Density

• Volume: Amount of space occupied; units: L, mL, cm3.

• Mass: Amount of matter; units: g, kg.

• Density: Mass per unit volume.

  • Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Can solve for mass or volume if density and one other variable are known.

Temperature Conversions

• Fahrenheit to Celsius: $T_{\text{C}} = \frac{5}{9}(T_{\text{F}} - 32)$

• Celsius to Fahrenheit: $T_{\text{F}} = \frac{9}{5}T_{\text{C}} + 32$

Chapter 3: Atoms: The Building Blocks of Chemistry

Atomic Structure and Notation

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

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

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

• Atomic symbol notation: $^{A}_{Z}\text{X}$ (e.g., $^{14}_{6}\text{C}$)

• Number of electrons in a neutral atom equals the atomic number.

Isotopes

• Atoms of the same element with different numbers of neutrons (different mass numbers).

• Isotopes have identical chemical properties but different physical properties (e.g., $^{12}\text{C}$, $^{13}\text{C}$, $^{14}\text{C}$).

Electron Arrangement

• Electrons occupy energy levels (shells), sublevels (s, p, d, f), and orbitals.

• Box (orbital) diagrams show electron arrangement using arrows for electrons.

• Electron configuration: e.g., $1s^2 2s^2 2p^6$

• Abbreviated configuration uses noble gas core (e.g., [Ne] $3s^2 3p^4$).

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

• Electron dot (Lewis) structures show valence electrons as dots around the element symbol.

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 atomic number by 2 and mass number by 4.

• Beta Decay (β–): Emission of an electron; neutron converts to proton.

• Positron Emission (β+): Emission of a positron; proton converts to neutron.

• Gamma Decay (γ): Emission of high-energy photons; no change in atomic or mass number.

Balancing Nuclear Equations

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

• Identify new isotopes formed after decay.

Radiation and Measurement

• Relative Energies: Gamma > Beta > Alpha.

• Curie (Ci): Measures radioactivity (disintegrations per second).

• Absorbed Dose: Measured in rads and grays (Gy).

Half-Life

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

• Half-life formula: $N = N_0 \left(\frac{1}{2}\right)^{t/t_{1/2}}$

• Used to calculate remaining quantity after a given time.

Medical Applications

• Most used radioisotope: Technetium-99m ($^{99m}\text{Tc}$).

• PET scans use $^{18}\text{F}$ (fluorine-18).

Nuclear Fission and Fusion

• Fission: Splitting of heavy nuclei (e.g., $^{235}\text{U}$); initiated by neutron absorption; produces chain reactions; used in nuclear reactors.

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

Chapter 5 (through 5.3): Ionic Compounds

Valence Electrons and Bonding

• Number of valence electrons determines how many electrons are gained or lost during bonding.

• Bonding: Atoms combine to achieve stable electron configurations (octet rule).

• Types of Bonds:

  • Ionic Bonds: Formed by transfer of electrons from metals to non-metals.

  • Covalent Bonds: Formed by sharing electrons (not covered in detail here).

Ions and Ionic Compounds

• Metals lose electrons to form cations (positive ions).

• Non-metals gain electrons to form anions (negative ions).

• Charge on an ion: $\text{Charge} = \text{Number of protons} - \text{Number of electrons}$

• Ionic compounds are electrically neutral (total positive charge = total negative charge).

Formulas and Charges

• Formulas show the simplest ratio of ions that results in neutrality (e.g., NaCl, MgCl2).

• Polyvalent elements (transition metals) can have multiple charges; specified with Roman numerals (e.g., Fe(II), Fe(III)).

• To determine the charge of a metal in a compound, use the known charge of the anion and the formula (e.g., in CuCl2, Cu is +2).

Ion	Symbol	Charge
Sodium	Na+	+1
Chloride	Cl–	–1
Calcium	Ca2+	+2
Oxide	O2–	–2
Copper (II)	Cu2+	+2
Iron (III)	Fe3+	+3
Silver	Ag+	+1
Gold (III)	Au3+	+3
Additional info: Table includes common ions relevant to ionic compounds.

Example:

• What is the charge of copper in CuCl2? Chloride is –1, so two Cl– ions = –2; copper must be +2 to balance.