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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.

Next Unit: Unit 3- Periodic Table