• Describe early milestones in the development of modern atomic theory.
• State and apply the law of conservation of mass and the law of definite proportions .
• State the premises of Dalton's atomic theory.
• Describe J. J. Thomson's experimental evidence for the existence of electrons.
• Describe Rutherford's scattering experiments and show how the results of the experiments imply the existence of atomic nuclei.
• List the three most important particles that all atoms are composed of, and describe their charges and relative masses.
• Understand the concept of atomic weight.
• Describe how isotopic masses and isotopic abundances are measured experimentally using mass spectrometry. Use a mass spectrum to compute an average atomic mass. Given a table of isotopic masses and abundances, sketch a mass spectrum.
• Predict the most common ion formed by a main group element by consulting a periodic table.
• Name and write the formulas for common transition metal ions.

Development of the Atomic Theory

• Problem: why do different materials have different properties?

  Early models for salt, water, & iron "atoms".

• Early atomists (Leukippos, Demokritos, and Epikouros)
  • samples can't be subdivided without limit
  • tiny, discrete, indestructible units of matter are atoms
  • atoms in constant motion through empty space
  • sizes and shapes of atoms determine all material properties
  • atoms have mass

• John Dalton, ca. 1803
  • Elements are composed of atoms with characteristic masses.
  • Compounds are composed of atoms combined in small, whole number ratios.

  Relationship between elemental composition and atom ratios in molecules. If C and O atoms weigh 12 and 16 units, respectively, the atom ratios in molecules convert to the element mass percentages in compounds.

  compound	% C by mass	% O by mass	possible molecules
  carbon monoxide	42.8%	57.2%	CO, C2O2, C3O3, ...
  carbon dioxide	30.0%	70.0%	CO2, C2O4, C3O6, ...

  • Chemical reactions rearrange connections between atoms.

Discovery of the Electron

• Michael Faraday, 1807
  • observation: in electrolysis of compounds, there is a definite relationship between the amount of element freed and the quantity of electricity passed through the compound
  • hypothesis: charge is somehow involved in binding elements together to form compounds
• George Stoney, 1891
  • hypothesis: atoms contain balanced positive and negative charges; the negative charges within atoms are "electrons"

• J. J. Thomson's cathode ray experiment
  • "cathode rays" pass from negative electrode towards positive electrode in an evacuated tube
    
  • hypothesis: cathode rays are streams of electrons
  • calculated mass to charge ratio for electrons by observing bending of cathode rays in electric and magnetic fields
  • proposed the plum pudding model of the atom

• R. A. Millikan's oil drop experiment
  • observation: charges on tiny droplets of liquid are always whole number multiples of a certain value
  • hypothesis: this value was the charge on a single electron
  • Nobel Prize, 1923
• electrons play a central role in chemistry
  • number and energies of electrons in an atom determine chemical properties of an element
  • electrons bind atoms into molecules

Discovery of the Nucleus

• Radioactivity
  • heavy elements are radioactive
  • electric field resolves radiation into 3 components: alpha, beta, and gamma

    Table: hypothetical description of alpha particles based on properties of alpha radiation

    observation	hypothesis
    alpha rays don't diffract	... alpha radiation is a stream of particles
    alpha rays deflect towards a negatively charged plate and away from a positively charged plate	... alpha particles have a positive charge
    alpha rays are deflected only slightly by an electric field; a cathode ray passing through the same field is deflected strongly	... alpha particles either have much lower charge or much greater mass than electrons

• Ernest Rutherford's scattering experiment
  • hypothesis: If the plum pudding model of the atom is correct, atoms have no concentration of mass or charge (atoms are 'soft' targets)
  • experiment to test hypothesis:
    • fire massive alpha particles at the atoms in thin metal foil
    • alpha particles should pass like bullets straight through soft plum pudding atoms
  • observation: a few alpha particles ricocheted!
  • new hypotheses:
    • all of the positive charge and nearly all of the mass of the atom is concentrated in a tiny, incredibly dense 'nucleus', about 10^-14 m in diameter
    • electrons roam empty space about 10^-10 m across, around the nucleus
• Composition of the Nucleus
  • nuclei are composed of "nucleons": protons and neutrons
  • atomic mass units
    • 1 amu (aka 1 dalton) = exactly 1/12 the mass of a carbon-12 nucleus
    • 1 dalton = 1.67 x 10^-24 g
    Table: Subatomic particles important in chemistry.

    particle	symbol	charge	mass, kg	mass, daltons
    electron	e-	-1	9.10953×10-31	0.000548
    proton	p+	+1	1.67265×10-27	1.007276
    neutron	n	0	1.67495×10-27	1.008665

• nuclear tug-of-war
  • electrostatic repulsion pushes protons in nuclei apart
  • strong nuclear force holds nucleons in nuclei together
  • energy required to break up a nucleus is millions of times the energy required to break up a chemical bond
    • that makes nuclei inert in chemical reactions.
  • range of strong nuclear force is about 10^-14 m; stable nuclei are small.
    • large nuclei tend to be unstable (radioactive)

Counting particles

• counting nucleons
  • atomic number
    • number of protons in the nucleus
    • Z is unique for each element
  • mass number (M)
    • number of protons and neutrons in the nucleus
  • symbols for nuclei:
    • nuclide symbols are 'top heavy': mass number is always the left superscript, and is always bigger than or equal to Z, the left subscript
    • sometimes Z is omitted
    • names are sometimes written out: element name dash mass number , e. g. uranium-235
    • nuclide names use mass not atomic number (carbon-12 is correct; carbon-6 is not)
• counting electrons
  • atoms
    • number of electrons equals number of protons
    • chief role of nucleus in chemistry: nucleus determines number of electrons
  • ions
    • atom (or molecule) with missing or extra electrons
    • charge = #protons - #electrons
    • charge given as a trailing superscript in nuclide symbols
    • positive ions are cations; negative ions are anions

Isotopes

• isotopes: mass number varies for atoms of a given element

• isotopes have different physical properties, but are chemically almost identical
• some isotopes are common; others are unstable and are very rare

  Table: Isotopes of carbon

  isotope	natural isotopic abundance	isotopic mass (daltons)	lifetime
  carbon-12	98.89%	12 exactly (by definition)	stable
  carbon-13	1.11%	13.003354	stable
  carbon-14	trace		5730 years
  carbon-15	nil		2.4 seconds
  carbon-16	nil		0.74 seconds

• isotopic abundance
  • abundances are "fingerprints" that can be used to trace source of samples
    • natural vs. synthetic compounds
    • black market fissionable materials
    • objects of extraterrestrial origin
    • age of objects
• isotopic mass
  • is NOT exactly equal to the total mass of nucleons

    total mass of nucleons = isotopic mass + mass defect
    

  • consequence: you can't calculate isotopic mass precisely from M, Z, and the nucleon masses
  • isotopic masses must be determined experimentally
    

Weighing atoms

• mass spectrometry is used to experimentally determine isotopic masses and abundances
• interpreting mass spectra
• average atomic weights
  • computed from isotopic masses and abundances
  • significant figures of tabulated atomic weights gives some idea of natural variation in isotopic abundances

Predicting ion charges

• rule for metal ion charges: metals lose electrons to get the same number of electrons as the nearest noble gas
  
  • rule only works if predicted charge is +3 or less;
    more than one common cation usually exists in this case
• rule for nonmetal ion charges:
  nonmetals gain electrons to get the same number of electrons as the nearest noble gas
  
  • rule only works if predicted charge is -3 or less;
    nonmetals that break this rule usually form covalent, not ionic, compounds
• when the rules don't help, get charges from a formula for a compound of the ion

Naming monatomic ions

• metal cations
  • cation name = metal name followed by "ion"
  • if more than one cation is possible, the charge must be specified in the name:
    • two naming styles
      • systematic name: metal name followed by charge on metal atom, written as a Roman numeral, in parentheses
      • common name: latin root followed by -ous for low charge form, -ic for high charge form
      • examples
  Table: Metal cations with more than one common charged form

  cation formula	systematic name	common name
  Fe2+	iron(II) ion	ferrous ion
  Fe3+	iron(III) ion	ferric ion
  Cu+	copper(I) ion	cuprous ion
  Cu2+	copper(II) ion	cupric ion
  Hg22+	mercury(I) ion	mercurous ion
  Hg2+	mercury(II) ion	mercuric ion
  Pb2+	lead(II) ion	plumbous ion
  Pb4+	lead(IV) ion	plumbic ion
  Sn2+	tin(II) ion	stannous ion
  Sn4+	tin(IV) ion	stannic ion

  • anion name = nonmetal root ending with "-ide"

  H-	hydride ion	O2-	oxide ion
  F-	fluoride ion	S2-	sulfide ion
  Cl-	chloride ion
  Br-	bromide ion	N3-	nitride ion
  I-	iodide ion

Copyright © 1997-2005 by Fred Senese
Comments & questions to fsenese@frostburg.edu
Last Revised 02/23/18.URL: http://antoine.frostburg.edu/chem/senese/101/atoms/index.shtml