Unit 1: Matter, Atoms, Elements, and Particles - Vocabulary Flashcards

Learning Objective 1.1: Explain and represent the structure and properties of matter, including macro- and micro-particulate representations

  • Chemistry definition: the study of the properties of matter and the changes that matter undergoes; matter is anything that has mass and occupies space.

  • Matter can be described at two levels:

    • Macro view: what we can observe directly (bulk properties).

    • Micro view: the underlying particles and interactions that explain those observations.

  • Matter is classified by its states and by its composition.

  • Key terms:

    • Mass: a measure of the amount of matter in a sample.

    • Volume: space that matter occupies.

    • Phase of matter: solid, liquid, or gas.

    • Condensed phases: solids and liquids (collectively).

  • Visual representations often use macro-particulate representations to describe substances (e.g., macroscopic observation with underlying particle ideas).

  • Quick concepts to remember:

    • All substances can, in principle, exist as a solid, liquid, or gas.

    • Phase changes interconvert states without changing chemical composition: extsolid⇌liquid⇌gasext{solid} \rightleftharpoons \text{liquid} \rightleftharpoons \text{gas}

  • Real-world relevance: understanding material behavior (melting, boiling, sublimation) in engineering, environment, and health.

Learning Objective 1.2: Identify experiments and conclusions leading to the Modern Atomic Theory

  • Historical progression from macro to micro: scientists used observable changes to infer atomic-level behavior.

  • Core experiments and concepts (as covered):

    • Rutherford’s gold foil experiment (nucleus concept) and the cathode ray tube experiments (discovery of electrons).

    • Millikan’s oil-drop experiment (charge of the electron).

    • Thomson’s cathode-ray experiments leading to electron discovery and the Plum Pudding model.

  • Key laws underpinning atomic theory:

    • Law of Conservation of Mass (1789, Lavoisier): in a chemical reaction, matter is neither created nor destroyed. m<em>extbefore=m</em>extafterm<em>{ ext{before}} = m</em>{ ext{after}}

    • Law of Definite Proportions (Joseph Proust): a given compound always has the same elemental mass ratio. For water, mass ratio O:H ≈ racm<em>Om</em>H=16.02.0=8:1.rac{m<em>O}{m</em>H} = \frac{16.0}{2.0} = 8:1. (atom ratio: O: H = 1:2)

    • Law of Multiple Proportions: when two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in small whole-number ratios (e.g., CO vs CO₂).

  • Dalton’s Atomic Theory (four postulates):

    1. Elements are composed of tiny, indivisible particles called atoms.

    2. All atoms of a given element have the same mass and properties that distinguish them from atoms of other elements.

    3. Atoms combine in simple, whole-number ratios to form compounds.

    4. Atoms cannot be changed into atoms of another element in chemical reactions.

  • Modern Atomic Theory refinements (summary):

    • All matter is composed of atoms, which contain subatomic particles (protons, neutrons, electrons).

    • Atoms of one element cannot be converted into atoms of another element in chemical reactions; conversion occurs in nuclear reactions.

    • Atoms of an element have the same number of protons and electrons in neutral atoms, determining chemical behavior.

    • Compounds form by chemical combinations of specific elements in fixed ratios and involve changes in electron structure during bonding.

  • Isotopes: atoms of the same element (same Z) with different numbers of neutrons (A varies). Neutrons contribute to mass but carry no charge. Isotopes may be represented as:

    • ZAX^{A}_{Z}X or AX^{A}X (mass number A, atomic number Z, symbol X).

  • Nucleus and electron cloud: protons and neutrons in the nucleus; electrons occupy a diffuse region around the nucleus.

  • Subatomic particle properties (conveniently in amu and elementary charge units):

    • Proton: mass ≈ 1.00727,extcharge=+11.00727, ext{ charge} = +1, mass ≈ 1.67262×10−27 extkg1.67262 \times 10^{-27}\ ext{kg}, charge +1.60218×10−19  extC+1.60218 \times 10^{-19}\ \, ext{C}

    • Neutron: mass ≈ 1.00866,extcharge=01.00866, ext{ charge} = 0

    • Electron: mass ≈ 0.00055,extcharge=−10.00055, ext{ charge} = -1, mass ≈ 9.1×10−31 extkg9.1 \times 10^{-31}\ ext{kg}, charge −1.60218×10−19  extC-1.60218 \times 10^{-19}\ \, ext{C}

  • Atomic mass unit (amu): defined by exactly 12 amu for the mass of the 12C isotope. 1 extamu=1.66054×10−24 g1\ ext{amu} = 1.66054 \times 10^{-24}\ \text{g}

  • Why amu? Because atomic masses are tiny; amu provides a practical scale for atomic-scale masses.

Learning Objective 1.3: Characterize subatomic particles, isotopes, and ions

  • Location of subatomic particles within an atom:

    • Protons and neutrons reside in the nucleus.

    • Electrons reside in the electron cloud surrounding the nucleus.

    • In a neutral atom: number of protons equals number of electrons (
      Z=e−extcount=extelectronsZ = e^- ext{ count} = ext{electrons}
      ).

  • Notation and masses:

    • Atomic number: ZZ = number of protons (identity of the element).

    • Mass number: A=Z+NA = Z + N, where NN is the number of neutrons.

    • Neutron number: N=A−ZN = A - Z.

    • Isotopes of an element have the same ZZ but different NN and thus different AA.

  • Isotopes can be represented and identified by:

    • Symbol with mass number: ZAX^{A}_{Z}X or a compact form like AX^{A}X.

  • Ions: charged atoms or groups of atoms (polyatomic ions).

    • Cations: positively charged (loss of electrons, e.g., extNa+,extAl3+ext{Na}^+, ext{ Al}^{3+}).

    • Anions: negatively charged (gain of electrons, e.g., extCl−,extO2−ext{Cl}^- , ext{O}^{2-}).

    • For many transition metals, ionic charge is shown with a Roman numeral: e.g., extTi3+,extCr2+ext{Ti}^{3+}, ext{Cr}^{2+}, etc.

    • Naming:

    • Nonmetals: root name + \