Unit 2- Atoms

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Atomic Structure:

• Atoms- The small particles which make up everything and anything.
• Subatomic particles- particles inside the atom. (Smaller than atoms)
• 2 main Parts:
  • Nucleus- In the center of the atom. Contains 2 subatomic particles.
  • Protons- Positively charged (+1 charge), subatomic particles. Weighs 1 amu
    • Protons = atomic number
  • Neutrons- Neutrally charged (0 charge), subatomic particles. Weighs 1 amu
    • amu = atomic mass unit
  • Clouds/Orbitals- circles on the outside of the atom. Contains 1 subatomic particle.
  • Electrons- negatively charged (-1 charge), subatomic particles. Has a negligible mass.
    • negligible- so small we pretend it dosent exist
• Atomic number- number of protons in an atom. Defining characteristic of all elements.
• Atomic Mass- Number of protons and neutrons in an atom (The 2 subatomic particles with weight)
  • (Is more shown as the average number of all the isotopes of an element)
  • Atomic Mass is the same as a mass number.
• Isotopes- Atoms that have the same atomic number,but different atomic masses. (The same element but different masses)
• Avogadro’s number- a mole of a substance contains 6.022x10^23 of that molecule or particle.

Theories:

Going in chronological order: (They all expanded off each other)

• Dalton’s theory- Realized that matter had to be made of particles, that could not be broken apart or destroyed.
  • The particle model (Just a sphere representing a particle)
• Thomson’s Atomic Theory- Used the plum pudding model (A machine where an electrode was shot through it) in there he found that particles had a negative charge. This proved the existence of electrons.
  • Found the existence of electrons. (Just put them in a particle model)
• Rutherford’s Atomic Theory- He shot particles at gold foil thinking they would just bounce back. Some ended going through, others deflected, and some did bounce back. This meant there was space inside atoms.
  • The nuclear model (model with nucleus in the center and electrons outside)
  • Found the nucleus and protons.
• Bohr’s Atomic Theory- Using some advanced physics dealing with electromagnetic energy, he figured out the orbitals/ clouds which held the electrons. Basically they were held together by electromagnetic forces (Like gravity).
  • The planetary model (Most common model of an atom showing a nucleus and circles that look like planet rings, which were the orbitals)
• Schrodinger’s model- A lot of upper level math made a atomic model which had the orbitals looking like balloon, basically an atom in 3D. (Isn’t used in the class very often)

Attraction:

• Protons and electrons work like magnets. They attract because they are positively and negatively charged respectively.
• The bigger the distance between them, the less attractive force there is
• The more Protons, the greater the attractive force.

Calculating Isotopic Mass (Atomic Mass):

It’s on the periodic table (which well have for the regents), but we might need to show how they got the mass number in the first place. Calculation steps + examples: (Not on flashcards)

They will give us everything to be able to do this on the test.

Ions:

Elements or atoms with the wrong number of electrons. (Has a charge)

• Total charge = Protons - Electrons
• Normally protons = electrons. But when they don’t it’s an ion.
• add “ide” to the end of ions
  • Ex: Oxygen → Oxide Ion
• When the electrons are taken out, the atoms becomes more positively charged.
  • When electrons are added in, the atom becomes more negatively charged.
  • This can be found in the upper right corner of the periodic table for each element.

Bright Line spectrum:

• Bohr’s model shows that as electrons gain energy they jump up orbitals
  • This means they also jump up energy levels
• They will always return back though
  • As they do they release energy and light.
  • The light will be colored, and each element has its own stripes of colors
    • Like Gel Electrophoresis (D.N.A Fingerprinting), but for atoms
• Wave lengths:
  • Shorter wavelengths give off colors like: Violet and blue. They have more energy
  • Longer wavelengths give off colors like: Orange and Red. They have less energy
• Excited state- When the electrons have more energy and jump up
• Ground state- When the electrons are back to their original orbitals
• Photon- Particle of light.

Valence Electrons:

Electrons on the outside of the atom (farthest orbital/cloud)

• Groups/ columns on the periodic table can help you figure out how many each element has.
  • Group 1= Has 1 valence electron
  • Group 2= Has 2 electrons on the outside (Valence electron)
  • Middle groups (3-12) don’t work like this they are random
  • Group 13= Has 3 valence electrons
  • Group 14= 4 valence electrons
  • Group 15= 5 valence electrons
  • Group 16= 6 valence electrons
  • Group 17= 7 valence electrons
  • Group 18= 8 valence electrons
  • After 12 just subtract ten from the group number ^

Electron notation:

• Orbitals:
  • S- can hold up to 2 electrons
  • P- can hold up to 6 electrons
  • d- can hold up to 10 electrons
  • f- can hold up to 14 electrons
• Energy levels are assigned numbers:
  • Level 1- Has orbital S only.
  • Level 2- Has orbitals S and P only.
  • Level 3- Has orbitals S, P, and d only.
  • Level 4-Has all orbitals S, P, d, and f. (All levels above 4 as well)
• Pattern examples:
  • Hydrogen= 1s^1
  • Helium= 1s^2
  • Lithium= 1s^2, 2s^1
  • Oxygen= 1s^2, 2s^2, 2p^2
  • Eventually you would get something like:
    • 1s^2 2s^2 2p^6 3s^2 3p^6 3d^10, and so on.
  • Once you get to 3p^6 it gets a bit weird because you follow it with 4s^2, and then 3d^10. This is because 4s^2 contains less energy than 3d^10
    • Ex: 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^2, 3d^10.

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Next Unit: Unit 3- Periodic Table

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