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Atomic Structure

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Atomic Structure

Atomic Theory

• Democritus (460-370 B.C.)– Greek philosopher – Democritus proposed that

the world is made up of empty space and tiny particles called atomos (or atoms)

Development of Atomic Theory

John Dalton (1766-1844) – developed Dalton’s Atomic Theory

– All elements are made up of atoms– Atoms of the same element are identical to each other but are

different than the atoms of other elements– Atoms of different elements can combine to form compounds– Atoms are indestructible and cannot be divided into smaller

particles.

Which one is NOT TRUE???

• J.J. Thompson

– Discovered electrons

– Cathode Ray Experiment

– Plum Pudding Model: electrons stuck into a lump of positive charge, like raisins stuck in dough (or chocolate chips stuck in cookie dough)

• Ernest Rutherford– Proposed that the atom is mostly empty space

with all of the mass and positive charge concentrated within the nucleus

– Conducted the “Gold Foil Experiment”

– Concluded that the atom is mostly empty space but contains a small, dense, positively charged central core called an atom.

– Nuclear Model

Atomic StructureParticle Location Charge Relative

Mass

Proton

(p+)

Nucleus + 1

Relatively large

Neutron (n0)

Nucleus 0 1

Same size as protons

Electron

(e-)

Shells - 0

Very small

The Atom

Nucleus – center core of an atom,

very dense, contains protons

and neutrons, overall positive

charge

Electron shell/orbit – surrounds the nucleus, where electrons are located

Atomic Number

• Tells the number of protons in the nucleus

• Because an atom is electrically neutral, the atomic number also indicates the number of electrons (#p+ = #e-)

6 12.011

C

Carbon

Mass Number

• Tells the number of protons PLUS neutrons

• Mass number - atomic number number of neutrons

6 12.011

C

Carbon

• Protons – 6• Electrons – 6• Neutrons –

12 - 6 = 6

6 12.011

C

Carbon

6p+

6n0

2e- 4e-

Isotopes

• Atoms that have the same number of protons but a different number of neutrons

• The atomic number stays the same in isotopes (the number of protons identifies the element)

• The mass number changes

Isotope Examples

6p+

6n0

2e-

4e-

6p+

7n0

2e-

2e-

4e-

C-12:

6p+

8n0

4e-

C-13:

C-14:

Isotope Examples

H-1 (Protium): 1p+

0n0

1e-

H-2 (Deuterium):1p+

1n0

1e-

H-3 (Tritium):1p+

2n0

1e-

Ions

• Cations – an atom with a positive charge because electrons have been lost

• Anion – an atom with a negative charge because electrons have been gained

A

Negative

I

O

N

Average Atomic Mass

• Weighted average of all naturally occurring isotopes of an element

• Explains why the mass number on the periodic table is a decimal number

Chlorine-35 has a percent abundance of 75.77 and an amu of 34.969. Chlorine-37 has a percent abundance of 24.23 and an amu of 36.966. Calculate the atomic mass of chlorine.

Cl-35: 34.969 X 0.7577 = 26.496

Cl-37: 36.966 X 0.2423 = 8.957

35.453

% abundance must be converted to relative abundance by dividing it by 100

The number after the symbol represent the atomic mass of the isotope

Electrons in Atoms

• Electrons move very quickly around the nucleus.

• The can also move up or down energy levels.

• Electrons move up an energy level when they are excited or given energy.

• Electrons move down an energy level when they lose energy.

Valence Electrons

• Electrons located in the outermost energy level (the last shell)

• Number of valence electrons = group number

Lewis Dot Diagrams• Use dots to represent the valence

electrons

• Steps to drawing dot diagrams:– Write the chemical symbol– Determine the number of

valence electrons– Draw out the dots in the

following configuration

Cl

The Bohr Model• Electrons are found in

specific paths (orbits) located around the nucleus.

• An electron must be ON an energy level.

• A “quantum” of energy will allow an electron to move from one energy level to another.

• Higher energy levels are located farther away from the nucleus

• Drawing Bohr models:

HIGHER ENERGY LEVEL

LOWER ENERGY LEVEL

Nucleus

6p+

6n0

2e-

4e-

1st energy level – 2 electrons

2nd energy level – 8 electrons

3rd energy level – 18 electrons

4th energy level – 32 electrons

Quantum Mechanical Model

• Developed by Erwin Schrodinger

• Describes the probability of finding an electron at various locations around the nucleus

Quantum Numbers

• Describes the location of the outermost electrons

• The electrons “zip code”

• Each element on the periodic table has a unique four digit quantum number

• Pauli Exclusion Principle – no two elements can have the same set of quantum numbers

Practice Problems

• K4, 0, 0, -1/2

• W5, 2, 1, -1/2

• Cu3, 2, 1, 1/2

• 5, 0, 0, ½Sr

• 5, 3, -2, -1/2Th

• 4, 2, 2, ½Cd

Practice Problems• Nitrogen

1s22s22p3

• Silicon

1s22s22p63s23p2

• Helium

1s2

• Chromium

1s22s22p63s23p64s13d5

• Gold

1s22s22p63s23p64s23d104p65s24d105p66s14f145d10

Short-Hand Electron Configuration

• Write the symbol for the noble gas (column 18) from the row before the element in brackets

• Write out the electron configuration for the entire row the element is located in

• Silicon– [Ne]3s23p2

• Chromium– [Ar]4s13d5

• Copper– [Ar]4s13d10

• Bismuth– [Xe]6s24f145d106p3

Practice Problems

Orbital Diagrams

• Series of lines and arrows used to represent the order in which the electrons fill the orbitals

• Write out the electron configuration

• Draw lines to represent the orbitals (s gets 1, p gets 3, d gets 5, and f gets 7)

• Draw arrows to represent the electrons (the superscript)

• Atomic Spectra– Atoms that have absorbed energy have

electrons that move to a higher energy level– When these electrons lose energy, they emit

light as they drop back to a lower energy level (their ground state)

– The light that is emitted by these atoms have very specific frequencies that appear as discrete lines when viewed through a diffraction grate

– The atomic spectrum of each element is unique because each element has a unique electron configuration

– The light emitted by the electron is directly proportional to the energy change of the electrons; the greater the energy change, the greater the frequency of light emitted