Atomic structure
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Atomic Structure. What is a theory?. a well-substantiated explanation of some aspect of the natural world; an organized system of accepted knowledge that applies in a variety of circumstances to explain a specific set of phenomena;

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Atomic Structure

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Atomic Structure


What is a theory?

  • a well-substantiated explanation of some aspect of the natural world;

  • an organized system of accepted knowledge that applies in a variety of circumstances to explain a specific set of phenomena;

  • "theories can incorporate facts and laws and tested hypotheses"


"If I have seen further it is by standing on the shoulders of Giants." --Isaac Newton


Early Theories

  • Democritus: 4 B.C.: “atom”

    • Believed there were 4 elements:

    • Fire, Air, Water, Earth


Dalton: 1766-1844

>All elements composed of tiny particles

called atoms

>Atoms of same element are identical;

atoms of different elements are different

>Atoms of different elements can physically

mix together or chemically combine to form

compounds

>Chemical reactions cannot change atoms

of one type of element to another


Thomson: 1856-1940

  • >discovered electrons in 1897

  • >used a cathode ray tube

  • >the ray produced was deflected by an electrical field (showed that atoms had particles with (-) charge)


Cathode Ray Tubes

  • A cathode ray tube or CRT is a specialized vacuum tube in which images are produced when an electron beam strikes a phosphorescent surface.

  • TVs, PCs, ATMs, video games, video cameras, and monitors all contain cathode-ray tubes.

  • Displays millions of colors.


Rutherford: 1871-1937

>Gold Foil Experiment

>Discovered the nucleus


Rutherford’s Gold Foil Experiment

Experiment

Shot positively charged alpha particles at gold foil

  • Results

  • Most particles passed through the foil

  • A few were deflected


Demonstration of Gold Foil Experiment


Rutherford’s Gold Foil Experiment

Conclusions

  • small, dense, positively charged core (nucleus)

  • the rest of the atom is empty space


Modern Theories

Bohr  planetary model

  • electrons arranged in concentric circular patterns

  • paths or orbits around nucleus (energy level)

    Wave-Mechanical Model  Electron Cloud Model

  • based on the ideas that orbitals are the area of highest probability where an electron will be found.

  • Orbitals have a variety of shapes and names (s, p, d, f)


Example: Wave Mechanical Model

  • Ψ2 (psi2) is a calculation that can predict the probability of finding an electron in a given area.


Summary- Atomic Models

Dalton’s Cannonball


Thomson’s Plum Pudding


Rutherford’s Nuclear Model


Bohr’s Planetary Model


Wave-Mechanical Model


1 amu = 1/12th mass of a carbon-12 atom

Subatomic Particles

**Note: amu = atomic mass unit


Atomic Number

  • Equal to the number of protons

  • Every element has its own atomic number

  • See Periodic Table

C

6


Mass Number

  • Equal to the sum of the protons and the neutrons (whole number)

  • Can be written as carbon-12

C

12


To find:

# of protons 

look up atomic number on Periodic Table


To find:

# of electrons 

in a neutral atom, it is equal to the number of protons


To find:

# of neutrons 

if protons + neutrons = mass then,

# of neutrons = mass # - # protons


Practice

20

40

20

20

20

12

24

12

12

12

11

23

11

12

11

2

4

2

2

2


Ions

  • Defined as “charged particles”

  • Ions are formed when the number of electrons changes.

  • If a (+) ion is formed, electrons are lost (called cations).

  • If a (-) ion is formed, electrons are gained (called anions).


Examples

  • Ca2+

    A Ca atom has 20 protons and 20 electrons.

    A Ca2+ ion has lost two electrons to have 18.


Examples

  • Cl-

    A Cl atom has 17 protons and 17 electrons.

    A Cl-ion has gained one electron to have 18.


Practice

30

30

65

30

35

26

56

26

30

23

9

19

9

10

9

127

53

74

54

53

3

4

3

7

2


Isotopes

  • Definition: elements that have the same atomic number but different mass (different # of neutrons)


Isotopic Symbols

  • Must write isotopic symbol to show mass

  • Same atomic #, different mass #

    X

Mass #

Atomic #


Write the isotopic symbol for:

  • Carbon-14

    C

14

6


Write the isotopic symbol for:

  • Oxygen-17

    O

17

8


Write the isotopic symbol for:

  • Chlorine-37

    Cl

37

17


Common Isotopes of Hydrogen

1

1

2

1

3

1


Why is atomic mass not a whole number?

  • The atomic mass on the periodic table is a weighted average of the isotopes of the elements.

  • The weighted atomic mass takes into account the relative abundances of all the naturally occurring isotopes.


Example of a general weighted average

  • Your grade in chemistry

  • 60% exams 85

  • 15% quizzes 100

  • 15% labs 95

  • 10% HW/CW 80

88.25

(0.60)85 + (0.15)100 + (0.15)95 + (0.10)80 =


Example 1:

  • Determine weighted atomic mass

  • Boron-10 19.78% 10.013 amu

  • Boron-11 80.22% 11.009 amu

10.812 amu

(0.1978) 10.013 + (.8022) 11.009 =


Example 2

  • Determine weighted atomic mass

  • Potassium-39 93.12% 38.964 amu

  • Potassium-41 6.88% 40.962 amu

39.101 amu

(0.9312) 38.964 + (0.0688) 40.962 =


Bohr models

How do electrons “orbit” the nucleus?

Each principal energy level …

  • is a fixed distance from the nucleus

  • can hold a specific number of electrons

  • has a definite amount of energy


  • The greater the distance from the nucleus…the greater the energy of the electrons in it.

  • The orbits are called principal energy levels or shells.


Energy levels or shells

energy levelnumber of e-

1 2

2 8

3 18

4 32

Increasing distance from nucleus

Increasing energy


Bohr models: examples

-energy levels

and total number

of electrons

8 e-

P+

n0

nucleus---

# protons

And neutrons

2 e-

Electron configuration: element’s symbol and number of

electrons in each orbit; LOOK UNDER ATOMIC NUMBER

ON PT of E


TRY THESE

12 p+

12 n0

Mg

2 e- 8 e- 2 e-

Electron configuration (bottom left corner on PT): Mg 2-8-2

HNaFC


answers

HNa

H 1Na 2-8-1

FC

F 2-7C 2-4

11 p+

12 n0

1p+

1 e-

2 e- 8e-1e-

9 p+

10 n0

6 p+

6 n0

2 e- 7 e-

2 e- 4 e-


Lewis Dot Diagrams

Valence shell: outer most shell of an atom that contains electrons

Valence electrons: electrons that occupy the valence shell (last number in electron configuration)

Electron dot diagrams or Lewis dot diagrams:

show only the valence shell of the atom

Ex: Lewis dot for nitrogen: N


TRY THESE

OFC

NeIK


Ions

For ions: remember that ions have gained or lost electrons.

(+)  indicate charge

(-)  use brackets and charge


Ca  Ca+2

Cl  [ Cl ]-1


Ground State vs. Excited State

  • When all electrons in an atom occupy the lowest available orbitals, it is said to be in the ground state.

  • When electron(s) absorb energy, they have the ability to jump to higher energy levels.

  • The excited state is when electrons have absorbed energy and no longer occupy the lowest available energy levels.


Possible Excited States

2-6-3

2-5-4


Absorption

  • When an electron “jumps” to a higher energy level it absorbs energy.

  • The excited state is a temporary state.

Excited State

(i.e. energy level 2)

e-

Ground State

(i.e. Energy level 1)


Emission

  • The electron then falls back down to the ground state, emitting energy.

  • The energy is in the form of light.


  • This radiant energy has a characteristic color and wavelength that can be determined.

  • Every electron transition produces a specific wavelength of light and all transitions for an element blend together.

  • This light can be separated through a prism into its various wavelength components.

  • Every element has its own unique bright line spectrum that can be used to help identify the presence of that element.

  • Ex: elements in a star, forensic analysis, flame tests, spectroscopy


Light and Atomic Spectra (bright line spectra)

Electromagnetic spectrum consists of light that exists as waves.


Sunlight and prisms

  • Sunlight produces a continuous range of wavelengths and frequencies that can be separated into all the colors of the rainbow.

  • R O Y G B I V.


  • Atomic emission spectra produce narrow lines of color called bright line spectra.

  • Each line corresponds to an exact wavelength.


Experiments – Flame Tests

  • Flame Tests – demonstrates the emission spectrum of a substance.

  • Completed by heating elements to high temperatures so they may enter excited state.

  • Characteristic color will be emitted as excited electrons return to ground state.

  • Used to determine metal ion presence in unknown substance.


Experiments – Spectroscopy

  • Spectroscopy – used to view the bright line spectra for given gases.

  • Completed by viewing a gas tube through which an electric current is passed.

  • Use an instrument called a spectroscope, contains a prism to separate emitted light into line spectra.


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