Atomic Theory and Structure

1
Atomic Theory and Structure

• Chapters 4-5

2
Atomic Theories

• Democritus 400 BC
• believed that atoms were indivisible and
  indestructible
• Dalton 1800s
• Developed through experiments
• First Atomic Model

3
Daltons Atomic Model

• All elements are composed of tiny indivisible
  particles called atoms
• Atoms of the same element are identical. The
  atoms of any one element are different from those
  of any other element.

4
Daltons Atomic Model (cont)

• Atoms of different elements can physically mix
  together or can chemically combine in simple
  whole-number ratios to form compounds.
• Chemical reactions occur when atoms are
  separated, joined, or rearranged. Atoms of one
  element, however, are never changed into atoms of
  another element as a result of a chemical
  reaction.

5
Discovery of Electron

• 1897 JJ Thomson, using cathode ray tube,
  discovered negatively charged particles called
  electrons
• 1909 Robert Millikan - Oil Drop Experiment
• Determined charge on an electron.

6
Plum Pudding Model

• Uniform positive sphere with negatively charged
  electrons embedded within.

7
Radiation

• Late 1800s discovery of radiation
• Three Types
• Alpha
• Beta
• Gamma

8
Rutherford Gold Foil Experiment - 1909

• Shot alpha particles at gold foil
• Most went through foil with little or no
  deflection.
• Some were deflected at large angle and some
  straight back.
• A.K.A. Geiger Marsden Experiment

9
Rutherford Gold Foil Experiment - 1909
10
Rutherford Model

• Conclusions from Gold Foil Experiment
• Atom is Mostly Empty Space
• Dense positive nucleus
• Electrons moving randomly around nucleus

11
Subatomic Particles

• Electron
• Discovered in 1897 by JJ Thomson
• Negative charge (-1)
• Mass 9.10938910-28g
• Approx mass 0
• Found outside of nucleus

12
Subatomic Particles

• Proton
• Discovered in 1919 by Rutherford
• Positive charge (1)
• Mass 1.67262310-24g
• Approx mass 1 atomic mass unit (u)
• Found inside nucleus

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Subatomic Particles

• Neutron
• Discovered in 1932 by James Chadwick
• No charge (0)
• Mass 1.674928610-24g
• Approx mass 1 atomic mass unit (u)
• Just slightly larger than a proton
• Found inside nucleus

14
Atomic Structure

• Atoms have no net charge
• of electrons of protons
• of electrons around nucleus of protons in
  nucleus

15
Atomic Structure

• Atomic Number
• Number of protons in an element
• All atoms of the same element have the same
  number of protons
• Mass Number
• Number of protons and neutrons in an atom

16
Atomic Structure

• of Neutrons Mass Number Atomic Number
• Atoms of the same elements can have different
  numbers of neutrons
• Isotope atoms of the same element with
  different number of neutrons

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Chemical Symbols
Mass Number

• Cl-35
• Chlorine-35

Atomic Number
18
Ion

• Atom or group of atoms that have gained or lost
  one or more electrons
• Have a charge
• Example
• H, Ca2, Cl-, OH-

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Ions

• H 1 proton 0 electrons
• Ca2 20 protons 18 electrons
• Cl- 17 protons 18 electrons
• OH- 9 protons 10 electrons

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Atomic Theories

• Rutherfords model could not explain the
  chemical properties of elements
• Niels Bohr believed Rutherfords model needed to
  be improved
• Bohr proposed that electrons are found only in
  circular paths around the nucleus

22
Bohr Model

• Dense positive nucleus
• Electrons in specified circular paths, called
  energy levels
• These energy levels gave results in agreement
  with experiments for the hydrogen atom.

23
Bohr Model
24
Bohr Model

• Each energy level can only hold up to a certain
  number of electrons
• Level 1 ? 2 electrons
• Level 2 ? 8 electrons
• Level 3 ? 18 electrons
• Level 4 ? 32 electrons

25
Electron Configuration

• The way in which electrons are arranged in the
  atom
• Example Na 2-8-1
• Valence Electrons
• Electrons in the outermost energy level

26
Energy Level Transitions

• Electron energy is quantized
• Electrons can move between energy levels with
  gains or losses of specific amounts of energy.

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Energy Level Transitions

• Gaining energy will move an electron outward to a
  higher energy level (Absorption)
• When an electron falls inward to a lower energy
  level, it releases a certain amount of energy as
  light (Emission)

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Energy Level Transitions
29
Ground State vs. Excited State

• Ground State
• When the electrons are in the lowest available
  energy level
• Ex Na 2-8-1
• Excited State
• When one or more electrons are not in the lowest
  available energy level
• Ex Na 2-7-2 or 2-8-0-1 or 2-6-1-1-1

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Line Spectra

• Emission Spectra
• Shows only the light that is emitted from an
  electron transition
• Absorption Spectra
• Shows a continuous color with certain wavelengths
  of light missing (absorbed)

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Energy Level Transitions
32
Energy Level Transitions
33
Wave Mechanical Model

• More detailed view of the Bohr Model
• Schrödinger Wave Equation and Heisenberg
  Uncertainty provides region of high probability
  where electron COULD be.
• Orbital
• Modern Model
• AKA Quantum Mechanical Model, Electron Cloud Model

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Wave Mechanical Model

• Orbital
• Regions of space where there is a high
  probability of finding an electron

35
Wave Mechanical Model

• Each energy level is divided into sublevels
• 1st Energy level has 1 sublevel, s
• 2nd Energy level has 2 sublevels, s and p
• 3rd Energy level has 3 sublevels, s, p, and d
• 4th Energy level has 4 sublevels, s, p, d, and f
• These sublevels start to overlap as you move away
  from the nucleus

36
Wave Mechanical Model

• Sublevels are divided into orbitals
• s sublevel has 1 orbital
• p sublevel has 3 orbitals
• d sublevel has 5 orbitals
• f sublevel has 7 orbitals
• Each orbital can hold up to 2 electrons

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Atomic Orbitals
38
Electron Orbital Configuration

• Sublevel order
• 1s
• 2s 2p
• 3s 3p 3d
• 4s 4p 4d 4f
• 5s 5p 5d 5f 5g
• 6s 6p 6d 6f 6g 6h
• 7s 7p 7d 7f 7g 7h

39
Electron Orbital Configuration

• One sublevel must be full before you can move to
  the next sublevel
• For sublevels with multiple orbitals
• Each orbital must have one electron before you
  can double up

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Electron Orbital Configuration

• H ____ 1s1
• He ____ 1s2
• Li ____ ____ 1s2 2s1

1s
1s
1s
2s
41
Electron Orbital Configuration

• C ____ ____ ____ ____ ____
• C 1s2 2s22p2
• N ____ ____ ____ ____ ____
• N 1s2 2s22p3

1s
2s
2p
1s
2s
2p
42
Electron Orbital Configuration

• O ____ ____ ____ ____ ____
• O 1s2 2s22p4
• F ____ ____ ____ ____ ____
• F 1s2 2s22p5

1s
2s
2p
1s
2s
2p
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MMs Demo

• What colors are found in a regular MMs bag?
• Green
• Yellow
• Orange
• Blue
• Red
• Brown

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MMs Demo

• Do you get an equal amount of each color in each
  bag?
• If we opened up all the regular MM bags in the
  world would we get an equal number of each color?
• Are you supposed to?

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MMs Demo
Color 1 bag World
Blue 24
Green 16
Yellow 14
Orange 20
Red 13
Brown 13
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MMs Demo

• MMs come in certain abundances (percentages)
• So do isotopes of each element
• Relative Abundance
• Percent of each naturally occurring isotope found
  in nature

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Average Atomic Mass

• Atomic Mass
• Weighted average based on the relative abundance
  and mass number for all naturally occurring
  isotopes
• Example
• C-12 98.9 12.011u
• C-13 1.1

49
Atomic Mass

• C-12 98.9
• C-13 1.1
• Carbon 0.98912 0.01113 12.011u