Download Atomic Theory notes.notebook

Atomic Theory notes.notebook
February 01, 2016
John Dalton: (1803) Atomic Theory
1. All matter made up of tiny indivisible
particles called atoms
2. All atoms of the same element are identical
3. Atoms of different elements are different
4. Atoms combine in simple whole number
ratios to form compounds
5. Atoms cannot be divided
John Dalton: (1810) Law of Multiple Proportions
– the same two elements can combine in
different ratios to form different compounds.
When comparing the masses between
compounds, they are still small whole number
ratios.
Feb 1­8:27 AM
Dalton came up with his ideas based solely on Large Scale experimentation (Macro World). He didn't have access to the kinds of equipment that we do today that allow us to see and do Small Scale experimentation (Micro World)
The scientists who do most of the Micro World experiments begin their work shortly after the first world war (1920's). But Just Before That...
Feb 1­8:32 AM
Atomic Theory notes.notebook
February 01, 2016
JJ Thompson, a British Chemist, (1897)
discovered the electron and calculated the charge­to­mass ratio of e­.
His Cathode ray tube experiment was instrumental in beginning the study of the SubAtomic Particles.
(e­/m = 1.759 x 108 coulomb/g)
Based on his findings, Dalton's Solid Sphere Atom is transformed into the Plum­pudding model of the atom (blueberry muffin/Chocolate Chip Cookie model)
Feb 1­8:34 AM
JJ Thompson becomes the head of Britain's Foremost research facility and attracts scientists from all over the world to come to Britain to perform experiments and study.
Robert Millikan: (1911) Oil‐Drop experiment established charge on e‐ (1.62 x 10‐19
coulomb)
1.
sprayed oil into the upper chamber, transferred e‐ to drops
2.
oil drops fall through chamber due to gravity
3.
charge on plates adjusted to offset gravity and suspend drop
charge on oil drops calculated **charges varied, but were all multiples of one
fundamental charge, the charge of an e‐**
Ernest Rutherford: (1911) Gold Foil Experiment – discovered the nucleus, led to discovery
and identification of protons
1.
2.
Bombarded gold foil with alpha particles
Most particles passed right through Some particles greatly deflected
James Chadwick: (1932) discovered the neutron; high energy, neutral particle emitted
from the nucleus
Feb 1­8:37 AM
Atomic Theory notes.notebook
February 01, 2016
Neils Bohr, A Danish Scientist, begins studying chemistry and physics, writes incredible papers that astound the scientific community, Pwns Einstein, and creates the following atomic Model:
Bohr’s Planetary Model of the Atom: (1913) 1. Electrons travel in definite energy levels without releasing energy
2. Electrons in each orbit have a certain amount of energy
3. Energy of orbits increases as distance from the nucleus increases
4. Electrons lose energy when they drop to lower energy levels
Feb 1­8:41 AM
All of which leads up to Heisenberg and Schrodinger ­ Einstein hates this, Bohr loves it.
Heisenberg’s Uncertainty Principle: (1927) It is impossible to know
the position and velocity of an e‐ at the same time.
Schrödinger’s Wave Equation: (1926) complex mathematical
function that treats the e‐ as a wave. The equation determines the
probability of finding the e‐ in any given place around the nucleus.
When the changing probabilities are plotted in 3‐D, the probability
area for finding the e‐ becomes a “cloud”. Schrödinger’s wave
equation used “quantum numbers”
Feb 1­8:43 AM
Atomic Theory notes.notebook
February 01, 2016
So what do you need to know?
Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U­235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Chm.1.1.2 • Analyze diagrams related to the Bohr model of the hydrogen atom in terms of allowed, discrete energy levels in the emission spectrum. • Describe the electron cloud of the atom in terms of a probability model. • Relate the electron configurations of atoms to the Bohr and electron cloud models. Chm.1.1.3 • Understand that energy exists in discrete units called quanta. 1. Describe the concepts of excited and ground state of electrons in the atom: 1. When an electron gains an amount of energy equivalent to the energy difference, it moves from its ground state to a higher energy level. 2. When the electron moves to a lower energy level, it releases an amount of energy equal to the energy difference in these levels as electromagnetic radiation (emissions spectrum). • Articulate that this electromagnetic radiation is given off as photons. • Understand the inverse relationship between wavelength and frequency, and the direct relationship between energy and frequency. • Use the “Bohr Model for Hydrogen Atom” and “Electromagnetic Spectrum” diagrams from the Reference Tables to relate color, frequency, and wavelength of the light emitted to the energy of the photon. • Explain that Niels Bohr produced a model of the hydrogen atom based on experimental observations. This model indicated that: 1. an electron circles the nucleus only in fixed energy ranges called orbits; 2. an electron can neither gain or lose energy inside this orbit, but could move up or down to another orbit; 3. that the lowest energy orbit is closest to the nucleus. • Describe the wave/particle duality of electrons. Feb 1­8:45 AM
We'll take these one at a time:
Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U­235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Particle Symbol location
relative relative charge mass
Proton
1
1
H1+, 11P1+
Nucleus
Plus 1
1 amu
Neutron
1
n0
Nucleus
Zero
1 amu
Electron
0
Cloud
Negative1
1/1857th amu
0
­1
e1­
Feb 1­8:46 AM
Atomic Theory notes.notebook
February 01, 2016
Next:
Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U­235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Atomic number (Z) – the numbers of protons in the nucleus,
identifies the element
Mass number (A) – The total number of Protons and
Neutrons in an element
Feb 1­9:02 AM
Next:
Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U­235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Isotopes – atoms of the same element with different numbers of neutrons.
Examples: protium (11H), deuterium (21H), tritium (31H)
Ion – an atom that has gained or lost electrons.
Lost electrons (cations) positive charge
Gained electrons (anions ) negative charge
Symbolic notation – shorthand method of describing a nuclide or ion
Example – A nuclide with A = 17 (8p+ + 9n), Z = 8 (8p+), and charge = ‐2 (8p+ ‐ 10e‐)
Feb 1­9:04 AM
Atomic Theory notes.notebook
February 01, 2016
Next:
Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U­235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Isotope Worksheet
Feb 1­8:48 AM
Next:
Chm.1.1.1 • Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U­235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding). Atomic mass
The atomic mass system is a relative system based on the mass of a carbon‐12 nuclide. The unit for
measuring atomic mass is the atomic mass unit (amu). 1 amu = 1/12 the mass of a carbon‐12 nuclide.
Atomic masses are specific to nuclides!
Average atomic mass
All naturally occurring elements exist as a combination of nuclides. For any element, the average atomic
mass is an average of all the masses of all the isotopes in a naturally occurring sample of the element.
Avg atomic mass = Σ (abundance in decimal form)(mass of nuclide)
Example: Find the average atomic mass for boron, given that 10B is 19.781% and 11B is 80.219%.
avg atomic mass = (0.19781)(10) + (0.80219)(11) = 10.812 amu
Beanium Activity
Feb 1­9:23 AM