Chemistry Review

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Chemistry Review

• You need to remember some basic things

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The Atom

• Smallest possible unit that maintains properties
  of the element
• Made of
• Protons positively charged particles, define
  the element, atomic number
• Neutrons- neutral particles
• Together form the atomic nucleus
• Electrons- negatively charged particles
• Fly around the nucleus
• Each element has a unique number of protons
  (atomic number)

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Electron Orbitals/Shells

• Electrons are found in characteristic areas
  around the nucleus, called an orbital
• Each one represents a different energy level
• Simplifying things, orbitals are grouped into
  shells

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Electron Shells

• The innermost shell of orbitals is filled first
• Electrons are distributed to each orbital in a
  shell before filling each orbital
• The outermost shell is called the valence shell

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Draw on your Whiteboard

• A neutral boron atom (for the nucleus you can
  just write B)
• A neutral fluorine atom

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Using the Periodic Table

• Ignore the metals
• The row tells you the of shells the atom should
  have
• The column tells you the of valence electrons a
  neutral atom should have in its valence shell

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Draw

• A neutral magnesium atom
• A neutral phosphorus atom

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Ions

• Aka charged atoms
• ions occur when there are more protons than
  electrons
• - ions occur when there are more electrons than
  protons
• Atoms can gain and lose electrons

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Filling Valence Shells

• Generally chemical reactions occur that fill
  valence electron shells
• Either by gaining/losing electrons
• OR
• By sharing electrons with other atoms

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6a. Covalent Bond

• Sharing of electrons between two atoms
• A single bond consists of 2 shared electrons,
  which occupy the valence shell of both atoms
• Double bond 4 electrons
• Triple bond 6 electrons

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Guidelines of Bonding

• Atoms almost always will end up with 8 electrons
  in their valence shell (may be lone pairs or
  shared electrons)
• So an atom that normally has 6 valence electrons
  needs to get 2 more from bonding

(only showing the valence electrons)
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The column can be used to figure out how many
bonds an atom will normally form
4 3 2 1 0
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Lewis Structures

• A line represents 2 electrons, usually shared in
  a covalent bond
• Dots represent electrons that are held by only
  one atom (lone pairs)
• Only valence electrons are shown
• Each atom should have a total of 8 electrons
  (except H and He which hold 2)

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6b.Polar vs. Non-Polar Covalent Bonds

• Nonpolar

• Polar

• Electrons not shared equally
• 1 atom is more electronegative (typically O, F,
  N, Cl )
• Electronegative atom ends up with a partial
  charge since they often hog the electron
• Other atom ends up with a partial charge as
  they have the electron less.

• Electrons shared equally
• Both atoms have similar electronegativity
  (affinity for electrons)
• Neither atom ends up with any charge

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• Non-Polar

• Polar

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10. Ion Formation

• Some atoms more easily give up e- (1st and 2nd
  columns) to get a full valence shell
• They commonly form bonds with atoms in the 6th
  7th column (respectively) since they need 1 or 2
  e-
• This is 1 way to form ions

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• There other 2 ways to turn an atom into an ion.
• Light, e.g. photoelectric effect where the
  energy of the incident photon kicks the electron
  out of its orbit. EX PHOTOSYNTHESIS
• Heat where the kinetic energy of atom and
  electron vibrations is so large that the electron
  vibrates away from the atom and does not return.

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6c. Ionic Bonding

• Opposites attract!
• Significantly weaker than a covalent bond
• Can also occur between ionic molecules

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11. Intermolecular Bonds

• Between 2 different molecules (think interstate
  highway is between 2 different states)
• I.e. hydrogen bonds in water
• Much weaker than intramolecular bonds
• aka intermolecular forces, attractions

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Hydrogen Bonds

• Weak attraction between the partial charges of
  polar covalently bonded molecules
• In water, between O and H

Means partial
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