Chapter 3 Atoms and Moles

Atomic Models

3.1 Matter Made of Atoms  Atomic Theory  Mikhail Lomonosov ( ) and Antoine Lavosier ( ): developed law of conservation of mass  states that mass of reactants equals mass of products

Law of Conservation of Mass

3.1 Matter Made of Atoms  Atomic Theory  Joseph Proust ( ): proposed law of definite proportions  states that two samples of a given compound are made of the same elements in exactly the same proportions by mass

3.1 Matter Made of Atoms  Atomic Theory  Claude-Louise Berthollet ( ): proposed law of multiple proportions  states that when two elements combine to form two or more compounds, the mass of one element that combines with a mass of the other is in the ratio of small whole #’s

3.1 Matter Made of Atoms  Dalton’s Atomic Theory  John Dalton ( ): developed a new atomic theory  1. all matter is made of atoms, which cannot be subdivided, created, or destroyed  2. atoms of a given element are identical in their chemical and physical properties

3.1 Matter Made of Atoms  Dalton’s Atomic Theory  3. atoms of different elements differ in their physical and chemical properties  4. atoms of different elements combine in simple, whole- number ratios to form compounds

3.1 Matter Made of Atoms  Dalton’s Atomic Theory  5. in chemical reactions, atoms are combined, separated, or rearranged but not created, destroyed, or changed

3.1 Matter Made of Atoms  Further Progress  Jons Berzelius ( )  studied proportions in which elements combine with one another (over 2000)  experimental underpinning of Dalton’s theory  made table of atomic weights  named ‘halogens’

3.1 Matter Made of Atoms  Further Progress  Jons Berzelius ( )  invented alphabetical nomenclature of elements  coined terms ‘organic chemistry’, ‘catalysis’, and ‘protein’

3.1 Matter Made of Atoms  Further Progress  valency and bonding described in 1850’s  Stanislao Cannizzaro (remember?): distinction between atoms and molecules  periodic table developed in 1860’s

3.2 Structure of Atoms  Subatomic Particles  Heinrich Geissler ( ): invented the vacuum tube (late 1850’s)  vacuum tube: hollow, glass tube in which the air has been removed; electrodes at either end  produces a glow when current flows between electrodes

3.2 Structure of Atoms  Subatomic Particles  Eugen Goldstein ( )  named glowing rays ‘cathode rays’ (1876)  showed that they were deflected by magnetic fields; could cast shadows  discovered rays coming from anode; called them ‘canal rays’ (1886)

3.2 Structure of Atoms  Subatomic Particles  William Crookes ( )  showed that cathode rays were made of particles, not light (1879)  convincing to the British, but not mainlanders

3.2 Structure of Atoms  Subatomic Particles  J. J. Thomson ( ): showed that rays were slower than light (1894)  Jean Perrin ( ): showed that metal plates hit by rays became negatively charged (1895)

Three Random Walks

3.2 Structure of Atoms  Subatomic Particles  J. J. Thomson (again)  measured mass/charge; found that particles were small or charge was large (1897)  measured electric charge itself; found electrons to be 1/2000 mass of a H atom (1899)  new atomic model

Deflections of Cathode Rays

Thomson’s Atomic Model

3.2 Structure of Atoms  Subatomic Particles  Ernest Rutherford ( )  discovered  and  radiation (1890’s)  discovered  radiation (1900)  discovered that  particles are a He nucleus (1908)

, ,  Radiation

 Radiation Image

3.2 Structure of Atoms  Subatomic Particles  Ernest Rutherford ( )  gold foil experiment (1909)   particles fired at gold foil  most went through, some deflected  conclusion: most of the mass and charge of an atom is in the nucleus; electrons in cloud

Gold

Gold Foil Experiment

Expectations versus Reality

Explanation

Rutherford’s Paper

3.2 Structure of Atoms  Subatomic Particles  Francis Aston ( ): showed that atoms come in different varieties (different weights) (1912)  called isotopes: atoms with the same number of protons but different numbers of neutrons  E.R. discovered proton (1918)

Evidence for Isotopes

3.2 Structure of Atoms  Subatomic Particles  James Chadwick ( ): discovered the neutron (sort of) (1932)

3.3 Electron Configuration  Electrons and Light  Light as a moving wave  c = f  c  speed of light = 3 x 10 8 m/s   wavelength (m)  distance between peak or troughs of a wave  f  frequency (1/s  1 hertz)  # of waves per second

Waves

Light Waves

Light

3.3 Electron Configuration  Electrons and Light  Albert Einstein ( )  atoms emit or absorb EM radiation in discrete (quantized) units (1905)  light has properties of waves and particles (1905)

3.3 Electron Configuration  Electrons and Light  Niels Bohr ( )  worked with Rutherford  new atomic model: electrons orbit nucleus at particular energy levels (1912)  electrons don’t give off energy (no spiraling allowed)  Why don’t electrons go straight to the nucleus???

3.3 Electron Configuration  Electrons and Light  Bohr’s model  electron in state of lowest possible energy is in ground state  if electron gains energy, it moves to an excited state(!)  if electron falls back to ground state, it releases energy as light

Excited State

Absorbance and Emission

Quantization

3.3 Electron Configuration  Electrons and Light  Bohr’s model, continued  Bohr predicted the wavelengths of light for hydrogen—he was right!  all light wavelengths together are called line-emission spectrum  each element has its own

Hydrogen Emission

H Absorbance and Emission

3.3 Electron Configuration  Electrons and Light  Louis de Broglie ( )  particles can be described as waves (1925)  therefore, electrons can only have certain frequencies (energy levels) and can’t fall toward nucleus  quantum atomic model

3.3 Electron Configuration  Quantum numbers  n  principal (main energy levels)  l  angular momentum (shape or type of sublevel)  l = 0  s orbital  l = 1  p orbital  l = 2  d orbital  l = 3  f orbital

Principal Quantum Number

Energy Level Transitions

3.3 Electron Configuration  Quantum numbers  m l  magnetic (subset of l quantum number)  m s  spin (orientation of magnetic field) +1/2 or -1/2

Quantum Numbers

Orbital Shapes

3.3 Electron Configuration  Electron Configurations  Pauli exclusion principle: each orbital can hold no more than two electrons  no two electrons can have the same four quantum numbers  Aufbau principle: electrons fill orbitals that have the lowest energy first  1s<2s<2p<3s<3p<4s<3d

Overlapping Orbital Energies

3.3 Electron Configuration  Electron Configurations  Hund’s rule: orbitals of the same n and l number are occupied by one electron before pairing occurs

Hund’s Rule