Chemistry Class on Significant Figures and Atomic Structure

Significant Figures (Sig Figs)

• Significant figures (sig figs) represent digits in a measurement that provide meaningful information on its precision.
• Non-zero digits are always significant.
• Zeros between non-zero digits are significant.
• Leading zeros are not significant.
• Trailing zeros in a decimal number are significant.
• Trailing zeros in a whole number (without a decimal) are not significant.
• Understanding significant figures is important to correctly round results in calculations.

Atomic Number and Atomic Mass

• Atomic Number (Z): The number of protons in an atom, defining the element.
• Example: Oxygen (atomic number 8) has 8 protons.
• Atomic Mass (A): The sum of protons and neutrons in an atom's nucleus (atomic weight).
• Example: Carbon (atomic mass ~12 amu) has 6 protons and 6 neutrons.

Isotopes and Radioactive Decay

• Isotopes: Different versions of the same element with varying neutron counts, changing atomic mass.
• Example: Carbon-12 (6 protons, 6 neutrons) and Carbon-14 (6 protons, 8 neutrons).
• Stable Isotopes: Do not change over time. Example: Carbon-12.
• Radioactive Isotopes: Decay over time, emitting radiation. Example: Carbon-14 decaying into Nitrogen-14.

Nuclear Reactions

• Nuclear reaction: A change in the atom's nucleus (not electrons), involving protons and neutrons.
• Alpha Decay: Nucleus emits an alpha particle (2 protons + 2 neutrons).
• Example: Uranium-238 → Thorium-234 + α particle.
• Beta Decay: A neutron transforms into a proton, releasing a beta particle (electron).
• Nuclear Fission: Splitting a heavy atom, releasing energy.
• Example: Uranium-235 fission.

Valence and Core Electrons

• Valence Electrons: Electrons in the outermost shell, dictating bonding behavior.
• Example: Oxygen (O) has 6 valence electrons.
• Core Electrons: Inner shell electrons not involved in bonding.
• Example: Oxygen has 2 core electrons in its first shell.

Ions and How They Form

• Ions: Charged atoms formed by gaining or losing electrons.
• Cations: Positive ions formed by losing electrons.
• Example: Sodium (Na) losing 1 electron to become Na⁺.
• Anions: Negative ions formed by gaining electrons.
• Example: Chlorine (Cl) gaining 1 electron to become Cl⁻.

Ionic Compounds

• Ionic compounds: Formed when metals transfer electrons to nonmetals, creating oppositely charged ions that attract.
• Example: Sodium (Na) to Chlorine (Cl) forming NaCl (table salt).

Writing Formulas for Ionic Compounds

• Balancing charges: Determine ion charges and balance positive and negative charges.
• Example 1: Mg²⁺ and Cl⁻ form MgCl₂
• Example 2: Na⁺ and O²⁻ form Na₂O

Polyatomic Ions

• Polyatomic ions: Ions with more than one atom, acting as a single unit with a charge.
• Example: ammonium (NH₄⁺), sulfate (SO₄²⁻), nitrate (NO₃⁻)
• Use parentheses for multiple instances: Ca(SO₄)₂

Transition Metals

• Transition metals: Groups 3-12, exhibiting multiple charges (oxidation states).
• Specify charge using Roman numerals:
• Example: Copper(I) chloride (CuCl), Copper(II) chloride (CuCl₂).

Physical vs. Chemical Changes

• Physical Change: Changes state or appearance, but substance remains the same.
• Examples: Ice melting, paper tearing.
• Chemical Change: Forms new substances.
• Examples: Burning wood, iron rusting.

Electron Configurations

• Electron configuration: Shows arrangement of electrons in atomic orbitals.
• Aufbau Principle: Fill lowest energy levels first.
• Pauli Exclusion Principle: Each orbital holds a maximum of 2 electrons.
• Hund's Rule: Fill each orbital individually before pairing electrons.
• Example: Carbon (C) electron configuration: 1s² 2s² 2p²

Binary Ionic Compounds

• Binary ionic compounds: Formed from 2 different elements (metal and nonmetal).
• Example 1: Mg²⁺ and Cl⁻ form MgCl₂
• Example 2: Na⁺ and S²⁻ form Na₂S

Rule of Zero Change

• Rule of Zero Change: Total positive charge from cations must equal total negative charge from anions for a neutral compound.
• Example: Mg²⁺ + 2Cl⁻ = MgCl₂ (zero charge).

This quiz covers essential concepts in chemistry including significant figures, atomic numbers, atomic mass, and isotopes. Understanding these principles is crucial for accurately interpreting measurement precision and elemental characteristics. Test your knowledge and solidify your grasp of foundational chemistry topics.