All matter is composed of atoms. ATOMIC STRUCTURE Understanding the structure of atoms is critical to understanding the properties of matter HISTORY OF THE ATOM DALTONS ATOMIC THEORY 1808 John Dalton suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them ATOMS 16 X + 8 Y 8 X 2 Y 1
HISTORY OF THE ATOM 1898 Joseph John Thompson found that atoms could sometimes eject a far smaller negative particle which he called an ELECTRON A = alpha B = gamma C = beta J.J. Thomson, measured mass/charge of e - (1906 Nobel Prize in Physics) CHARGE OF AN ELECTRON HISTORY OF THE ATOM gold foil 1910 Ernest Rutherford oversaw Geiger and Marsden carrying out his famous experiment. helium nuclei they fired Helium nuclei at a piece of gold foil which was only a few atoms thick. they found that although most of them Millikan oil drop experiment passed through. About 1 in 10,000 hit 2
Plum Pudding model of an atom. Rutherford s experiment. Results of foil experiment if Plum Pudding model had been correct. Actual Results. 3
A nuclear atom viewed in cross section. Rutherford s Model of the Atom atomic radius ~ 100 pm = 1 x 10-10 m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m Atoms are composed of -protons positively charged particles -neutrons neutral particles -electrons negatively charged particles Protons and neutrons are located in the nucleus. Electrons are found in orbitals surrounding the nucleus. Particle Subatomic Particles Mass (g) Charge (Coulombs) Charge (units) Electron (e - ) 9.1 x 10-28 -1.6 x 10-19 -1 Proton (p) 1.67 x 10-24 +1.6 x 10-19 +1 Neutron (n) 1.67 x 10-24 0 0 mass p = mass n = 1840 x mass e - 4
proton HELIUM ATOM Shell Every different atom has a characteristic number of protons in the nucleus. - + N N + - atomic number = number of protons electron neutron Atoms with the same atomic number have the same chemical properties and belong to the same element. Each proton and neutron has a mass of approximately 1 dalton. The sum of protons and neutrons is the atom s atomic mass. Isotopes atoms of the same element that have different atomic mass numbers due to different numbers of neutrons. ATOMIC NUMBER (Z) = number of protons in nucleus MASS NUMBER (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons ISOTOPS are atoms of the same element (X) with different numbers of neutrons in the nucleus Mass Number Atomic Number 1 H 2 A Z X 1 1 H (D) 1 H (T) 235 238 3 92 U 92 U Element Symbol 5
Two isotopes of sodium. HISTORY OF THE ATOM 1913 Niels Bohr studied under Rutherford at the Victoria University in Manchester. Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons. 6
Line Emission Spectrum of Hydrogen Atoms Every element has a unique emission spectrum LIGHT EMISSION OF SODIUM ATOM Line spectrum 7
The Bohr Model of the Atom Line spectrum of some elements The Bohr Model of the Atom: Ground and Excited States In the Bohr model of hydrogen, the lowest amount of energy hydrogen s one electron can have corresponds to being in the n = 1 orbit. We call this its ground state. When the atom gains energy, the electron leaps to a higher energy orbit. We call this an excited state. The atom is less stable in an excited state and so it will release the extra energy to return to the ground state. Either all at once or in several steps. The Bohr Model of the Atom: Hydrogen Spectrum Every hydrogen atom has identical orbits, so every hydrogen atom can undergo the same energy transitions. However, since the distances between the orbits in an atom are not all the same, no two leaps in an atom will have the same energy. The closer the orbits are in energy, the lower the energy of the photon emitted. Lower energy photon = longer wavelength. Therefore, we get an emission spectrum that has a lot of lines that are unique to hydrogen. 8
The Bohr Model of the Atom: Hydrogen Spectrum Neutral atoms have the same number of protons and electrons. Ions are charged atoms. -cations have more protons than electrons and are positively charged -anions have more electrons than protons and are negatively charged An ion is formed when an atom, or group of atoms, has a net positive or negative charge (why?). If a neutral atom looses one or more electrons it becomes a cation. ELECTRONS IN ATOMS Na 11 protons 11 electrons Na + 11 protons 10 electrons If a neutral atom gains one or more electrons it becomes an anion. Cl 17 protons 17 electrons Cl - 17 protons 18 electrons 9
ELECTRONS IN ORBIT ABOUT THE NUCLEUS ELECTRON DENSITY OF 1s ORBITAL Where 90% of the e - density is found for the 1s orbital Schrödinger Wave Equation Ψ = fn(n, l, m l, m s ) l for a given value of n, l = 0, 1, 2, 3, n-1 e - density (1s orbital) falls off rapidly as distance from nucleus increases n = 1, l = 0 n = 2, l = 0 or 1 n = 3, l = 0, 1, or 2 l = 0 l = 1 l = 2 l = 3 s orbital p orbital d orbital f orbital Shape of the volume of space that the e - occupies 10
l = 0 (s orbitals) l = 2 (d orbitals) l = 1 (p orbitals) Energy of orbitals in a single electron atom Energy depends only on principal quantum number n Energy of orbitals in a many-electron atom Energy depends on n and l n=3 n=3 l = 2 n=2 1 E n = -R H ( ) n 2 n=3 l = 0 n=2 l = 0 n=3 l = 1 n=2 l = 1 n=1 n=1 l = 0 11
Fill lowest energy orbitals first (Aufbau principle) Fill lowest energy orbitals first (Aufbau principle) He 2 electrons He 1s 2 Li 3 electrons Li 1s 2 2s 1 Hund s rule: The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins. Hund s rule: The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins. O 8 electrons O 1s 2 2s 2 2p 4 F 9 electrons F 1s 2 2s 2 2p 5 12
Hund s rule: The most stablearrangement of electrons in subshells is the one with the greatest number of parallel spins. Outermost subshell being filled with electrons Ne 10 electrons Ne 1s 2 2s 2 2p 6 Electrons determine all of the chemical properties and some of the physical properties of elements. 13