Chapter 3

Atoms and MolesAtoms and Moles

Atomic Models

3.1 Matter Made of Atoms Atomic Theory

Mikhail Lomonosov (1711-1795) and Antoine Lavosier (1743-1794): developed law of conservation of massstates that mass of reactants equals mass of products

Law of Conservation of Mass

Law of Conservation of Mass

3.1 Matter Made of Atoms Atomic Theory

Joseph Proust (1754-1826): proposed law of definite proportionsstates 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 (1748-1822): proposed law of multiple proportionsstates 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 (1766-1844): developed a new atomic theory1. 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 (1779-1848)studied proportions in which elements combine with one another (over 2000)experimental underpinning of Dalton’s theory

made table of atomic weightsnamed ‘halogens’

3.1 Matter Made of Atoms Further Progress

Jons Berzelius (1779-1848)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 (1814-1879): 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 (1850-1930)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 (1832-1919)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 (1856-1940): showed that rays were slower than light (1894)

Jean Perrin (1870-1942): 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 (1871-1937)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 (1871-1937)gold foil experiment (1909)

particles fired at gold foilmost 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

Explanation

Rutherford’s Paper

3.2 Structure of Atoms Subatomic Particles

Francis Aston (1877-1945): 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 (1891-1974): discovered the neutron (sort of) (1932)

3.3 Electron Configuration Electrons and Light

Light as a moving wavec = f

c speed of light = 3 x 108 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 (1879-1955)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 (1885-1962)worked with Rutherfordnew 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 modelelectron 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

Absorbance and Emission

Quantization

3.3 Electron Configuration Electrons and Light

Bohr’s model, continuedBohr predicted the wavelengths of light for hydrogen—he was right!

all light wavelengths together are called line-emission spectrumeach element has its own

Hydrogen Emission

H Absorbance and Emission

3.3 Electron Configuration Electrons and Light

Louis de Broglie (1892-1987)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 orbitall = 1 p orbitall = 2 d orbitall = 3 f orbital

Principal Quantum Number

Energy Level Transitions

3.3 Electron Configuration Quantum numbers

ml magnetic (subset of l quantum number)

ms 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 electronsno two electrons can have the same four quantum numbers

Aufbau principle: electrons fill orbitals that have the lowest energy first1s<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