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. 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!

•  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 α – Beta particle - 0-1e - an e- β

•  1 Form is energy – Gamma Rays- γ - 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 - ν- # 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 - λ - physical dist. betw. peaks – Related by c = λν

•  Amplitude - maximum displacement from zero

Excited atoms lose energy

•  Energy emitted by gaseous atoms can be spread into a spectrum. – Emission Spectrum - shows λ’s of light given

off by excited atoms – Absorption Spectrum - have lines missing

from continuous spectrum showing which λ’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 λ’s •  Gamma waves - shortest λ’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 = hν

–  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 λ’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.