Electron configuration ch 11
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Electron Configuration Ch 11. review. What do we know about electrons? How many electrons does a neutral atom contain? To learn: How are these electrons arranged around the nucleus?. Development of Atomic Theory. Elements (Boyle – 1600’s)

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Electron Configuration Ch 11

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Electron configuration ch 11

Electron ConfigurationCh 11


Review

review

  • What do we know about electrons?

  • How many electrons does a neutral atom contain?

  • To learn:

  • How are these electrons arranged around the nucleus?


Development of atomic theory

Development of Atomic Theory

  • Elements (Boyle – 1600’s)

  • Elements are made up of atoms, which cannot be further broken down. All atoms of elements are exactly alike. (Dalton early 1800’s)

  • Atoms with positive & negative charges mixed together (JJ Thompson – mid 1800’s)

  • Atoms with solid nucleus & electrons on outside (Rutherford – later 1800’s)

  • Electrons in concentric orbits, like planets, around nucleus (Bohr – early 1900’s)


Electron locations bohr model

Electron locations – Bohr model

  • First electrons closest to nucleus

  • Only room for two – why?

  • Draw the Bohr model for hydrogen, helium

  • Next “orbit” holds 8 electrons

  • Draw the Bohr model for nitrogen, neon

  • 3rd “orbit” holds 8 electrons

  • Draw Bohr model for magnesium, sulfur

  • What is last atom with three levels?


Development of atomic theory cont d

Development of Atomic Theory – cont’d

  • Wave-mechanical model (deBroglie – 1920’s)

  • Firefly analogy – orbitals are not orbits

  • Orbitals = 3-dimensional region in which there is a high probability of finding an electron in an atom


Electron configuration rules see ch 11 section 4

Electron Configuration – rulesSee Ch 11, section 4

  • Rules of filling:

  • a maximum of 2 electrons per orbital

  • Electrons in any one orbital spin in opposite directions.

  • Electrons “fill in” from the lowest energy level to the highest.

  • Periods correspond to energy levels for s & p blocks. The d block energy level = period -1.

    Format: 3d4, where the 3 = energy level,

    • d = orbital shape,

    • 4 = number of electrons in the orbital


Orbital configuration

Orbital configuration

  • Much like electron configuration, but also:

  • Shows spin

  • Shows individual orbitals within blocks

  • Shows how electrons fill in

  • 3 p orbitals in each energy level: x, y & z

  • First: px1 py1 pz1

  • Then: px2 py2 pz2


Valence electrons

Valence electrons

  • Valence electrons = electrons at the highest energy level

  • Why aren’t d orbitals ever valence electrons?

  • b/c they always fill in at one lower energy level than the valence s orbital for that element (period – 1)

  • What is the maximum number of valence electrons an element can ever have?

  • 8 (two from s + 6 from p)


Lewis dot structures

Lewis Dot Structures

  • Named after G.N Lewis (1902)

  • Shorthand way of showing valence electrons

  • Dots around element symbol

  • s orbital first (two dots on top)

  • p orbitals next, one at a time going around symbol.

    HW: draw the lewis dot structures for all the elements 11-18 & 19 - 36


Electrons and energy

Electrons and Energy

  • Energy as light – waves or particles?

  • Both or either

  • Called wave-particle duality

  • Photon = “particle” of light (electromagnetic radiation)

  • When electrons absorb energy, they get “excited” and jump to a higher energy state

  • When the electrons go to lower energy state:

  • Emit photons of light


Electron configuration ch 11

Electromagnetic Spectrum

  • Different elements emit different visible colors

  • Why?

  • Because the different energy levels correspond to different wavelengths


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