Atoms, Molecules, and Ions - Key Terms (VOCABULARY Flashcards)
2.1 Early Ideas in Atomic Theory
Greek origin of the idea: Leucippus and Democritus in the 5th century BC proposed that matter is composed of tiny indivisible particles called atoms.
Term origin: atomos, from Greek for “indivisible.”
Later, Aristotle/post-Aristotelian thinking favored the four classical elements: fire, earth, air, and water, challenging the atomist view.
Dalton’s contribution (1807): English schoolteacher John Dalton proposed a modern atomic theory.
Dalton’s five postulates (summary):
Matter is composed of exceedingly small particles called atoms. An atom is the smallest unit of an element that can participate in a chemical change.
An element consists of only one type of atom, which has a mass that is characteristic of the element and is the same for all atoms of that element.
Atoms of one element differ in properties from atoms of all other elements.
A compound consists of atoms of two or more elements combined in a small, whole-number ratio. In a given compound, the number of atoms of each of its elements are always present in the same ratio.
Atoms are neither created nor destroyed during a chemical change, but rearrange to yield a different type(s) of matter.
Microscopic/macro connection: Dalton’s theory provides a microscopic basis for macroscopic properties.
Dalton-era visual aids:
Figure 2.2: A pre-1982 copper penny contains approximately $3 imes 10^{22}$ copper atoms; all share the same chemical properties. (illustrative)
Figure 2.3: Copper(II) oxide in a 1:1 Cu:O ratio (CuO) shows how different elements combine in fixed ratios.
Law of Conservation of Matter (mass):
If atoms are neither created nor destroyed in a chemical change, total mass remains constant: m{ ext{initial}} = m{ ext{final}}.
Dalton’s atomic theory and conservation of matter: microscopic atoms rearrange rather than disappear or appear.
Example 2.1 (Dalton’s theory test – redistribution vs destruction):
Scenario: Starting with two green atoms and two purple atoms; products show one green and one purple atom—this would imply destruction if atoms disappear, violating Dalton’s postulates.
Solution: Atoms are neither created nor destroyed; redistribution in whole-number ratios is required. If the product count reduces atoms, the change violates the postulates. (Key idea: mass and atom count must be conserved.)
Example 2.1 (redistribution in stable counts):
Starting with four green and two purple atoms; products also have four green and two purple atoms.
This does not violate Dalton’s postulates: Atoms are not created/destroyed but redistributed in small, whole-number ratios.
Law of Definite Proportions (Law of Constant Composition):
A pure compound contains the same elements in the same proportion by mass regardless of sample size.
Historical origin: Joseph Proust.
Table 2.1: Constant composition of isooctane (illustrative data)
Sample A: Carbon = 14.82 g, Hydrogen = 2.78 g
Sample B: Carbon = 22.33 g, Hydrogen = 4.19 g
Sample C: Carbon = 19.40 g, Hydrogen = 3.64 g
The mass ratio C:H is constant across samples, illustrating the law of definite proportions.
Law of Multiple Proportions: when two elements form more than one compound, a fixed mass of one element combines with small whole-number ratios of the other element.
Example: chlorine (Cl) and copper (Cu) compounds.
Green solid: 0.558 g Cl per 1 g Cu; Brown solid: 1.116 g Cl per 1 g Cu.
2.2 Evolution of Atomic Theory – Learning objectives preview
Outline milestones in modern atomic theory.
Summarize and interpret Thomson, Millikan, Rutherford experiments.
Describe subatomic particles; define isotopes and give examples.
2.2 Evolution of Atomic Theory
What are atoms composed of? Is there something smaller than an atom? We discuss key developments.
J. J. Thomson – Discovery of the Electron:
Experiment: cathode ray tubes.
Cathode ray tube (CRT): sealed glass tube with most air removed, two metal electrodes; a visible beam (cathode ray) appears between electrodes when high voltage is applied.
The beam deflection:
The beam always deflected toward the positive charge and away from the negative charge, regardless of the metal electrodes used.
Thomson used deflections in electric and magnetic fields to determine the charge-to-mass ratio of the cathode ray particles: rac{e}{m} ext{ (charge-to-mass ratio)} = 1.759 imes 10^{11} rac{ ext{C}}{ ext{kg}}.
Conclusion: cathode rays are composed of negatively charged particles (electrons) with a characteristic e/m ratio.
Thomson’s Plum Pudding Model (Fig. 2.8a): atoms as a positively charged sphere with embedded negative electrons (to explain overall neutrality).
Robert A. Millikan – Oil Drop Experiment (1909):
Method: microscopic oil droplets are charged and influenced by an electric field; the field is adjusted to suspend or retard drops.
Purpose: measure the electron’s charge by balancing gravitational and electric forces on each drop.
Result: the charge on an individual oil drop is a multiple of a fundamental charge, $e = 1.6 imes 10^{-19}$ C.
Millikan concluded that the elementary charge is $1.6 imes 10^{-19}$ C, i.e., the charge of a single electron.
Correlation with Thomson’s e/m: Thomson had already established the ratio; Millikan provided the elementary charge value.
Rutherford – Discovery of the Nucleus (Gold Foil Scattering):
Experiment: α particles (positively charged) directed at a thin gold foil; a fluorescent screen detects deflections.
Observations: most α particles passed through; a few were deflected; very few experienced large deflections.
Rutherford’s interpretation (Fig. 2.10):
The atom is mostly empty space.
A very small, dense, positively charged nucleus at the center contains most of the atom’s mass.
Negatively charged electrons surround the nucleus.
Nuclear model of the atom (Fig. 2.10, 2.9, 2.10): nucleus at center; electrons occupy a large volume around it.
Subsequent discoveries related to the nucleus:
Isotopes: atoms of the same element with different masses.
Neutrons: uncharged subatomic particles with mass ≈ proton mass; discovered by James Chadwick (1932) and located in the nucleus.
Isotopes and neutrons together explain mass differences and isotopic abundances.
Isotopes and symbol notation (transition to 2.3):
Isotopes are described by mass number A and atomic number Z: the isotope notation often places A as a superscript to the left of the element symbol and Z as a subscript to the left (when shown).
Example: Magnesium isotopes ${}^{24}{12} ext{Mg}$, ${}^{25}{12} ext{Mg}$, ${}^{26}_{12} ext{Mg}$ share 12 protons but differ in neutrons (12, 13, 14 neutrons respectively).
2.3 Atomic Structure and Symbolism
Atomic structure essentials:
The nucleus contains most of the atom’s mass; protons and neutrons (nucleons) are much heavier than electrons.
Electrons occupy nearly all of the atom’s volume; the diameter of an atom ~ 10^{-10} ext{ m}, while the nucleus is ~10^{-15} ext{ m}.
The Nuclear Model (illustrated): nucleus at center with electrons surrounding.
Subatomic particle properties (approximate values):
Proton: mass ≈ 1.0073 ext{ amu}, charge +1
Neutron: mass ≈ 1.0087 ext{ amu}, charge 0
Electron: mass ≈ 0.00055 ext{ amu}, charge -1
Atomic Number (Z):
Z = number of protons in the nucleus; defines the element (identity).
Example: any atom with 6 protons is carbon, Z = 6.
Neutral atoms:
In a neutral atom, number of protons equals number of electrons. Therefore, the atomic number Z also equals the number of electrons in a neutral atom.
Mass Number (A):
A = total number of protons and neutrons in the nucleus.
Neutrons count: N = A - Z.
Ions:
If protons and electrons are not equal, the atom is charged (ion).
Charge of an ion = number of protons − number of electrons.
Cations and Anions:
Cation: positive charge (loss of electrons). Example: Na (Z = 11) losing one electron becomes Na⁺ with charge +1.
Anion: negative charge (gain of electrons). Example: O (Z = 8) gaining two electrons becomes O^{2-} with charge −2.
Chemical Symbols:
A chemical symbol abbreviates the element; examples: Hg for mercury.
Some symbols derive from Latin names (e.g., Fe from ferrum, Pb from plumbum).
Most symbols have one or two letters; some elements with Z > 112 have three-letter symbols used historically; only the first letter is capitalized.
Isotopes and notation:
The isotope notation shows A as superscript and Z as subscript (when shown): e.g., ${}^{A}_{Z} ext{X}$.
All isotopes of an element have the same Z, but different A (and thus different neutrons and masses).
Common elements and symbols (partial table):
aluminum: Al; iron: Fe; bromine: Br; lead: Pb; calcium: Ca; magnesium: Mg; carbon: C; chlorine: Cl; nitrogen: N; chromium: Cr; oxygen: O; cobalt: Co; potassium: K; copper: Cu; silicon: Si; fluorine: F; silver: Ag; gold: Au; sodium: Na; helium: He; sulfur: S; hydrogen: H; iodine: I; zinc: Zn.
Isotopes in very light elements (Table 2.4, partial for hydrogen):
Hydrogen (Z = 1): protium (A = 1) mass ~ 1.0078 ext{ amu}, abundance ~ 99.989%; deuterium (A = 2) mass ~ 2.0141 ext{ amu}, abundance ~ 0.0115%; tritium (A = 3) mass ~ 3.01605 ext{ amu}, trace.
Worked Examples (2.1) – protons, neutrons, electrons from isotope notation:
Strategy: superscript = A (mass number), subscript = Z (atomic number). If no subscript, Z inferred from symbol.
Atoms are neutral: electrons = protons = Z.
(a) Z = 17 (Cl); A = 35 → protons = 17; neutrons = A − Z = 18; electrons = 17.
(b) Z = 17; A = 37 → protons = 17; neutrons = 20; electrons = 17.
(c) Potassium (K) with Z = 19; A = 41 → protons = 19; neutrons = 22; electrons = 19.
(d) Carbon (C) with Z = 6; A = 14 → protons = 6; neutrons = 8; electrons = 6.
The Nuclear Model implications:
Nucleus contains most of the mass; electrons occupy a large volume around the nucleus.
Nuclear model explains why most of the atom is empty space and why α particles mostly pass through in Rutherford’s experiment.
2.4 Chemical Bonds, Formulas, and the Mole
2.4 Chemical Bonds
Elements: substances composed of a single type of atom.
Compounds: substances composed of two or more elements held together by chemical bonds, in fixed definite proportions.
Example:
Elements: O, H, etc.
Compounds: Al, Fe, C, O2, P4 (illustrative).
Chemical bond: the attractive force that holds atoms together in compounds.
2.4 Representing Compounds: Chemical Formulas and Molecular Models
Chemical formula indicates the elements present and the relative number of atoms/ions of each.
Components:
Chemical symbols for each element.
Subscript indicating the number of atoms (omit 1).
Examples:
Water: ext{H}_2 ext{O}
Carbon dioxide: ext{CO}_2
Sodium chloride: ext{NaCl}
2.4 Types of Chemical Formulas
Empirical formula: simplest whole-number ratio of atoms in a compound.
Molecular formula: actual number of atoms of each element in a molecule.
Examples:
Hydrogen peroxide: Molecular formula ext{H}2 ext{O}2; Empirical formula ext{HO}.
Butene: Molecular formula ext{C}4 ext{H}8; Empirical formula ext{CH}_2.
Example: glucose
Molecular formula: ext{C}6 ext{H}{12} ext{O}_6.
Empirical formula: ext{CH}_2 ext{O}.
Practice 3: How many of the following formulas are empirical formulas? ext{C}6 ext{H}6, ext{H}2 ext{SO}4, ext{P}4 ext{O}{10}, ext{C}2 ext{H}4 ext{O}
Answer: 1 (namely $ ext{H}2 ext{SO}4$ is already in simplest whole-number ratio; others can be reduced: $ ext{C}6 ext{H}6 o ext{CH}$, $ ext{P}4 ext{O}{10} o ext{P}2 ext{O}5$, $ ext{C}2 ext{H}4 ext{O} o ext{CH}_2 ext{O}$).
Elements that exist as molecules
Diatomic molecules: ext{H}2, ext{N}2, ext{O}2, ext{F}2, ext{Cl}2, ext{Br}2, ext{I}_2
The most common form of elemental sulfur is ext{S}_8.
2.4 The Mole – The foundation concepts
Define the mole:
1 dozen = 12; 1 mole = 6.022 imes 10^{23} particles.
Avogadro’s number: N_A = 6.022 imes 10^{23} ext{ mol}^{-1}.
A mole is the amount of substance containing N_A particles.
2.4 Formula Mass
Formula mass is the sum of the atomic masses of all atoms in a molecule (units: amu).
Example: Chloroform, ext{CHCl}_3
Calculation: 1 imes 12.01 + 1 imes 1.008 + 3 imes 35.45 = 119.37 ext{ amu}.
2.4 The Mole – Molar mass
Molar mass is the mass of 1 mole of a substance (units: g/mol).
It is numerically equal to the formula mass: e.g., for chloroform, molar mass =119.37 ext{ g/mol}.
2.4 Converting between Mass and Number of Particles
Key relationships:
Mass (g) $ o$ moles: n = rac{m}{M} where $M$ is the molar mass (g/mol).
Moles $ o$ number of particles: N = n imes N_A.
Mass conversion example: 5.60 mol Cu contains N = 5.60 imes N_A = 3.37 imes 10^{24} atoms.
Small-mass sample example: 40.0 mg saccharin (molar mass 183.18 g/mol):
Mass in grams: 0.0400 ext{ g}
Moles: n = rac{0.0400}{183.18} ext{ mol} \
2.1 Early Ideas in Atomic Theory
Greek Origin: Leucippus and Democritus (5th century BC) proposed matter consists of indivisible particles called atomos (Greek for “indivisible”). Aristotle later favored four classical elements.
Dalton’s Atomic Theory (1807): John Dalton proposed:
Matter is composed of atoms, the smallest unit of an element in chemical change.
An element consists of one type of atom with characteristic mass.
Atoms of different elements have different properties.
Compounds combine atoms of two or more elements in small, whole-number ratios.
Atoms are rearranged, not created or destroyed, in chemical changes.
Laws Based on Dalton's Theory:
Law of Conservation of Matter (Mass): Total mass remains constant in a chemical change (m
_\text{initial} = m
_\text{final}), as atoms only rearrange.Law of Definite Proportions (Constant Composition): A pure compound always contains the same elements in the same proportion by mass.
Law of Multiple Proportions: When two elements form multiple compounds, a fixed mass of one element combines with whole-number ratios of the other.
2.2 Evolution of Atomic Theory
J. J. Thomson (Cathode Ray Tube): Discovered electrons (negatively charged particles) and determined their charge-to-mass ratio (\frac{e}{m} = 1.759 \times 10^{11} \frac{\text{C}}{\text{kg}}). Proposed the "Plum Pudding Model."
Robert A. Millikan (Oil Drop Experiment, 1909): Measured the elementary charge of an electron, e = 1.6 \times 10^{-19} \text{ C}.
Ernest Rutherford (Gold Foil Experiment): Alpha particles fired at gold foil showed most passed through, some deflected, few greatly deflected.
Interpretation: Atom is mostly empty space with a tiny, dense, positively charged nucleus containing most mass, surrounded by electrons. Led to the nuclear model.
Post-Rutherford Discoveries:
Isotopes: Atoms of the same element with different masses (different number of neutrons).
Neutrons: Uncharged subatomic particles (mass \$\approx\$ proton) found in the nucleus, discovered by James Chadwick (1932).
Isotope notation: {}^{A}
_{Z} \text{X}, where A = mass number, Z = atomic number.
2.3 Atomic Structure and Symbolism
Atomic Structure: Nucleus (protons, neutrons) holds most mass; electrons occupy most volume. Atom diameter \$\sim\$ 10^{-10} \text{ m}, nucleus \$\sim\$ 10^{-15} \text{ m}.
Subatomic Particles:
Proton: mass \$\approx\$ 1 \text{ amu}, charge +1
Neutron: mass \$\approx\$ 1 \text{ amu}, charge 0
Electron: mass \$\approx\$ 0 \text{ amu}, charge -1
Atomic Number (Z): Number of protons; defines the element. In neutral atoms, electrons = protons = Z.
Mass Number (A): Protons + Neutrons. Neutrons = A - Z.
Ions: Charged atoms (protons \neq electrons).
Cation: Positive charge (loses electrons).
Anion: Negative charge (gains electrons).
Chemical Symbols: Abbreviate elements (e.g., Hg).
Isotopes: Atoms of the same element (same Z) with different mass numbers (A, due to different neutrons). Noted as {}^{A}
_{Z} \text{X}.Determining Protons, Neutrons, Electrons:
Protons = Z. Electrons = Z (for neutral atoms). Neutrons = A - Z.
2.4 Chemical Bonds, Formulas, and the Mole
Chemical Bonds: Hold atoms together in compounds, forming substances with fixed proportions.
Elements: One type of atom (e.g., O, H).
Compounds: Two or more elements bonded (e.g., \text{H}_2\text{O}).
Chemical Formulas: Indicate elements and their relative atom counts (e.g., \text{H}_2\text{O}).
Empirical Formula: Simplest whole-number ratio of atoms.
Molecular Formula: Actual number of atoms in a molecule.
Molecules of Elements: Some elements exist as diatomic molecules (\text{H}_2, \text{N}_2, \text{O}_2, \text{F}_2, \text{Cl}_2, \text{Br}_2, \text{I}_2) or larger (e.g., \text{S}_8).
The Mole:
Amount of substance containing Avogadro’s number (N
_A = 6.022 \times 10^{23}) of particles.
Formula Mass: Sum of atomic masses in a molecule (amu).
Molar Mass: Mass of 1 mole of a substance (g/mol), numerically equal to formula mass.
Conversions:
Mass (g) \leftrightarrow Moles (n = \frac{m}{M})
Moles \leftrightarrow Number of Particles (N = n \times N
_A$$)