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Energy & change
The quantum theory
Electrons in atoms
The periodic table
- Describe early milestones in the development of modern atomic theory.
- State and apply the
law of conservation of mass
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.
Lecture outlineElementary atomic theory is presented from an experimental perspective.
Development of the Atomic Theory
|Nothing exists but atoms and empty space; all else is opinion.
- 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
More about Dalton's atomic theory
- 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.
|CO, C2O2, C3O3, ...
|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 received the Nobel Prize in Physics in 1906 for his discovery of the electron.
- 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
Table: Hypothetical properties of the electron. How J. J. Thomson used properties of cathode rays to hypothesize properties of the electron.
|ray properties are independent of the cathode material
... cathode ray stuff is a component of all materials
|cathode rays bend near magnets
... magnets bend the paths of moving charged particles;
maybe cathode rays are streams of moving charged particles
|rays bend towards a positively charged plate.
rays impart a negative charge to objects they strike.
... cathode rays are streams of negative charges
|Cathode rays don't bend around small obstacles,
cast sharp shadows,
can turn paddlewheels placed in their path, and travel in straight lines
... cathode rays behave like streams of particles
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
|Rutherford received the 1908 Nobel Prize in Chemistry for his pioneering work in nuclear chemistry.
- 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
Table: Subatomic particles important in chemistry.
- 1 amu (aka 1 dalton) = exactly 1/12 the mass of a carbon-12 nucleus
- 1 dalton = 1.67 x 10-24 g
|All nuclei with more than 83 protons are unstable. But size isn't the only factor- the neutron/proton ratio must be just right for a nucleus to be stable. The optimal ratio is about 1:1 for light nuclei; it increases to about 1.5:1 for bismuth.
- 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
- large nuclei tend to be unstable (radioactive)
- 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
- number of electrons equals number of protons
- chief role of nucleus in chemistry: nucleus determines number of electrons
- 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: mass number varies for atoms of a given element
|Compounds containing different isotopes react at slightly different rates. For example, electrolysis of water is faster for water containing 1H than for water with 2H. The difference in reaction rate is used to isolate "heavy water".
|The first mass spectrograph was built in 1919 by F. W. Aston, who received the 1922 Nobel Prize for this accomplishment.
- 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
Table: Metal cations with more than one common charged form
- 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
- anion name = nonmetal root ending with "-ide"
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General Chemistry Online! Atoms & ions
Copyright © 1997-2005 by Fred Senese
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Last Revised 02/23/18.URL: http://antoine.frostburg.edu/chem/senese/101/atoms/index.shtml