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Atomic Structure
Atomic Structure
All matter is composed of atoms.
Understanding the structure of atoms is
critical to understanding the properties of matter.
HISTORY OF THE ATOM
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
DALTONS ATOMIC THEORY
16X + 8Y 8X2Y
HISTORY OF THE ATOM
1898 Joseph John Thompson
found that atoms could sometimes eject a
far smaller negative particle which he
called an ELECTRON.
J.J. Thomson, measured mass/charge of e-
1906 Nobel Prize in Physics
CHARGE OF AN ELECTRON
1909 Millikan
oil drop experiment
charges were all multiplesof a certain base value, which was found to be 1.6 ×10−19 C
HISTORY OF THE ATOM
1910 Ernest Rutherford
oversaw Geiger and Marsden carrying out
his famous experiment.
they fired Helium nuclei at a piece of gold
foil which was only a few atoms thick.
they found that although most of them
passed through. About 1 in 10,000 hit
Plum Pudding model of an atom
Rutherford’s experiment
Results of foil experiment: if Plum Pudding model had been correct.
Actual Results
Rutherford’s Model of the Atom
atomic radius ~ 100 pm = 1 x 10-10 mnuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
A nuclear atom viewed in cross section
Atomic Structure
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
Subatomic Particles
Particle 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
Atomic Structure
Every different atom has a characteristic number of protons in the nucleus.
atomic number = number of protons
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 atoms of the same element (X) with different numbers of neutrons
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.
The Bohr Model of the Atom
Shell
The Bohr Model of the Atom
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.
Ground and Excited States
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 orbits are in energy, the lower energy of the photon emitted.
– Lower energy photon = longer wavelength.
• Therefore, emission spectrum has a lot of lines that are unique to hydrogen
Emission spectrum of Hydrogen
Every element has a unique emission spectrum
The Bohr Model of the Atom:
Hydrogen Spectrum
ELECTRONS IN ORBIT ABOUT THE NUCLEUS
ELECTRON DENSITY OF 1s ORBITAL
Ψ = fn(n, l, ml , ms ) n=principal quantum number
for a given value of n, l = 0, 1, 2, 3, … n-1
n = 1, l = 0
n = 2, l = 0 or 1
n = 3, l = 0, 1, or 2
Shape of the “volume” of space that the e- occupies
l = 0 s orbital
l = 1 p orbital
l = 2 d orbital
l = 3 f orbital
Schrödinger Wave Equation
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
He 2 electronsHe 1s2
Fill lowest energy orbitals first (Aufbau principle)
↓↑
O 8 electronsO 1s2 2s2 2p4
Hund’s rule: The most stable arrangement of electrons in subshells is one with the greatest number of parallel spins.
↓↑
↓↑
↓↑ ↑ ↑
8 octet
2
8
18
Outermost subshell being filled with electrons
Atomic Structure
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 .
If a neutral atom looses one or more electrons it becomes a cation.
If a neutral atom gains one or more electrons it becomes an anion.
Electrons determine all of the chemical properties
and some of the physical properties of elements.
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