Chapter 4 Atomic Structure HONORS CHEMISTRY Chapter 4 Atomic Structure

History of the Atomic Theory

DEMOCRITUS (400 BC) 1st atomic theory “World is made of empty space & tiny particles called ‘atoms’.” Atomos - Greek for indivisible Smallest poss. particles of matter “Diff. types of atoms for every type of matter” General & not supported by experiment Not accepted - Contradicted Aristotle

ARISTOTLE “Matter is continuous” - not made of smaller particles “Hyle” Accepted until 17th Century

Isaac Newton & Robt. Boyle Published articles on belief in atomic nature of elements No Proof Attempted explanations, no predictions

John Dalton Logical hypothesis on existence of atoms Studied & explained work of other scientists Lavoisier - “In a closed syst., the mass of the reactants = the mass of the products” LAW OF CONSERVATION OF MASS Proust - “Specific substs. always contain elems. in the same ratio by mass” LAW OF DEFINITE PROPORTIONS

Dalton’s Atomic Theory Basis of modern atomic theory 1st atomic theory based on experimental evidence

Dalton’s Atomic Theory Four important statements: 1. All matter is composed of indivisible atoms. 2. All atoms of the same elem. are identical. 3. Atoms of diff. elems. are not alike. 4. Atoms unite in simple ratios to form compounds.

Dalton’s Atomic Theory Explains Law of Cons. of Mass atoms are rearranged in a chem. rxn. Explains Law of Definite Proportions Not exactly correct

DALTON ALSO STATED: Law of Multiple Proportions The ratio of masses of one element that combines w/ a constant mass of another elem. can be expressed in small whole numbers.

Other Scientists Gay Lussac - “Under constant conditions, the volumes of reacting gases & gaseous products are in the ratio of small whole numbers.” Avogadro explained this - “Equal volumes of gases, under the same conditions, contain the same # of molecs.”

Cathode Ray Tube Tube w/ charged metal electrodes in ea. end Anode - Positive electrode Cathode - Neg. electrode Rays in tube seemed to travel from cathode to anode Cathode Rays

J. J. Thomson Discovered electrons using cathode ray tube Determined charge to mass ratio of e-

Robt. Millikan Oil Drop Experiment Measured the charge on an e- std. unit of neg. charge (-1) e- mass is 1/1837 mass of a H atom

J. J. Thomson Discovered electrons using cathode ray tube Determined charge to mass ratio of e- Discovered the proton using a modified cathode ray tube same amt. of chg. as e- but positive std. unit of (+) chg. = +1 Calculated mass of p+ (1836 X mass of e-)

Lord Rutherford 1920 - predicted 3rd particle

James Chadwick Discovered the neutron high energy particle w/ no chg. & approx. same mass as p+

Dalton’s Theory was revised. Subatomic particles had been discovered.

J. J. Thomson Discovered ISOTOPES atoms of the same elem. that differ in mass have same # of p+’s, but diff. # of no’s

Henry Mosely 1913 - using x-rays, found the number of p+’s in the nucleus of an atom is always the same for a given element Atomic Number (Z) - # of p+’s in the nucleus # p+’s = # e-’s in a neutral atom

The number of p+’s determines the identity of the elem The number of p+’s determines the identity of the elem. and the # of no’s determines the particular isotope of the elem.

Dalton’s Theory revised again Not all atoms of the same element are exactly alike. Atoms are NOT indivisible!

Nucleons - particles that make up the atomic nucleus Nuclide - a particular type of atom containing a definite # of p+’s & no’s Nucleons - particles that make up the atomic nucleus p+’s & no’s Mass Number (A) - total # of nucleons in an atom Number of no’s = A - Z (mass # - atomic #)

Rutherford’s Gold Foil Experiment 1912-1913 led by Lord Rutherford, assisted by a team of physicists (Niels Bohr, Hans Geiger, & Ernest Marsden) Procedure: shot (+) charged subatomic particles @ very thin sheet of gold foil.

Rutherford’s Gold Foil Experiment Observations 1. Most particles passed straight thru foil. 2. Few particles were deflected @ large angles. 3. Very few (1 in 8000) bounced almost straight back. Conclusions: 1. Most of the atom is empty space. 2. + particles came close to “core” of atom which must have a + charge. 3. + particles almost hit core straight on.

Rutherford’s Gold Foil Experiment Overall Conclusion Atoms consist of (+) charged nucleus surrounded by e-’s

Diameter of an atom ~ 100-500 pm Radii of nuclei of atoms vary between 1.2x10-3 and 7.5 x 10-3 pm Nucleus is ~ 1 trillionth the vol. of the atom.

Henri Becquerel 1896 w/ Marie & Pierre Curie discovered Radioactive Substs. When brought near charged electroscope, leaves become discharged

Radioactivity Phenomenon of rays being produced spontaneously by unstable atomic nuclei mixture of particles & energy given off by nuclei during spontaneous nuclear decay amt. of energy very large - E = mc2 Half-Life - length of time needed for 1/2 an amt. of a radioactive nuclide to disintegrate.

Nuclear Force - force which holds p+’s and No’s together in nucleus effective over very short distance

Scientists agree on: 1. Nucleons have a prop. that corresponds to spinning on an axis. 2. e-’s don’t exist in nucleus, but can be emitted from nucleus.

Star Trek Science

Subatomic particles - particles composing atoms 2 broad classes Leptons - (light particles) - truly elementary best known: electrons Hadrons - appear to be made of smaller particles best known: neutrons & protons

For every particle, a mirror image particle called an Antiparticle is believed to exist antielectron is a positron When particle & its antiparticle collide, both are destroyed & energy is produced.

Several Leptons electrons neutrinos - essentially massless Muon - much more massive than e- Tau - much more massive than e-

Hadrons divided into 2 groups Mesons Baryons p+’s and no’s are baryons Both made of Quarks 6 kinds of quarks up, down, charmed, strange, top, bottom ea. quark comes in 3 different “colors” - red, blue, green ea. quark has antimatter counterpart - antiquark

If structure of nucleus is unstable, ejects particle or energy to become stable Some nuclei naturally unstable, some artificially unstable

3 forms of radiation from naturally radioactive nuclei 2 are particles Alpha particle - 42He - helium nucleus a Beta particle - 0-1e - an e- b 1 Form is energy Gamma Rays- g - very high energy x-rays

Short hand to represent particles Upper rt. “corner” - charge on ion Lower rt. - # of atoms in formula unit Upper left - mass # Lower left - charge on nucleus or particle

Examples 3216S - Sulfur nucleus or atom 0-1e - electron 42He - alpha particle (helium nucleus)

Scientists create radioactive nuclides by bombarding stable nuclei w/ accelerated particles or w/ neutrons in nuclear reactor Decay by emitting natural radiation & other methods.

Planetary Atomic Model Proposed by Rutherford and Bohr e-’s “orbit” around nucleus H atom similar to solar syst. w/ 1 planet

Bohr exposed atoms to radiant energy atoms absorb some energy Excited Atoms Excited atoms & molecs. produce energy changes unique & can be used to identify particle absorb and emit radiant energy

SPECTROSCOPY Method of studying substs. exposed to exciting energy

SPECTRUM Pattern of radiant energy studied in spectroscopy

ELECTROMAGNETIC (RADIANT) ENERGY Visible light, radio, ultraviolet, infrared, etc. Travels in waves variations in elect. & magnetic fields taking place in regular repeating fashion Frenquency - n- # of wave peaks that occur in a unit of time meas. in hertz (Hz) = 1 peak or cycle per sec.

ELECTROMAGNETIC (RADIANT) ENERGY Travels @ speed of light (c) 3.00 x 108 m/s in vacuum Wavelength - l - physical dist. betw. peaks Related by c = ln Amplitude - maximum displacement from zero

Excited atoms lose energy Energy emitted by gaseous atoms can be spread into a spectrum. Emission Spectrum - shows l’s of light given off by excited atoms Absorption Spectrum - have lines missing from continuous spectrum showing which l’s of light have been absorbed

Lines missing in absorption spectrum are the same as lines shown in emission spectrum unique to ea. elem. used to identify elems.

Electromagnetic Spectrum Radio waves - longest l’s Gamma waves - shortest l’s Visible light????

Planetary Model of Atom Developed by Bohr to explain H spectrum Used Quantum Theory - theory of energy emission stated by Max Planck

Quantum Theory Planck assumed energy was emitted in packets or Quanta - not continuously Quanta of radiant energy - Photons Amt. of energy given off is directly related to frequency of light emitted E = hn h = Planck’s constant (6.63 x 10-34 J/Hz) E = energy in a quanta

Planck’s Hypothesis Energy is given off in quanta instead of continuously

Bohr “Absorption of light by H @ definite l’s corresp. to definite changes in energy of e-.

Reasoned: 1. Orbits of e- around nucleus must have definite diameter. 2. e-’s can occupy only certain orbits. 3. Only orbits allowed - those w/ diff. in energy = energy absorbed when atom was excited \ e-’s can absorb quantum of energy & move to larger orbit Since quantum represents certain amt. of energy, next orbit must be definite dist. from 1st orbit.

When e- drops to lower orbit, energy is emitted (light). Orbit represents definite energy \ def. amt. of energy is given off. Ground State of e- is its smallest (lowest) orbit

Today’s atomic model differs from Bohr’s Major diff. - e-’s do not move around nucleus like planets orbit sun Idea of energy levels still basis of modern atomic theory.

Average Atomic Mass Mass of single atom too sm. to work with. use mass of large group of atoms Chemists meas. mass of single atoms in Atomic Mass Units.(amu or u) C-12 nuclide chosen at std. - all other atoms compared to it one C-12 atom defined as having a mass of 12 amu

Average Atomic Mass An Atomic Mass Unit - 1/12 the mass of a C-12 nuclide e- = 9.10953 x 10-28g = 0.000549 u p+ = 1.67265 x 10-24g = 1.0073u no = 1.67495 x 10-24g = 1.0087u Number in Periodic table based on “average atom” of the elem. Ave. atomic mass used for calculations

Average Atomic Mass 2 ways of determining masses for atoms of elem: 1. Experimentation & calculation 2. Mass Spectrometer - meas. masses & relative amts. of nuclides for all isotopes of an elem.

Average Atomic Mass If masses of isotopes & relative amts. are known, ave. atomic mass can be calculated Atomic Mass of the elem. Must use weighted average to find ave. atomic mass.