Download - Chemistry Review
Chemistry ReviewYou need to remember some basic things
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)
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”
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
Draw on your Whiteboard
•A neutral boron atom (for the nucleus you can just write B)
•A neutral fluorine atom
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
Draw
•A neutral magnesium atom
•A neutral phosphorus atom
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
Filling Valence Shells
•Generally chemical reactions occur that fill valence electron shells
•Either by gaining/losing electrons OR
•By sharing electrons with other atoms
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
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)
The column can be used to figure out how many bonds an atom will normally form
4 3 2 1 0
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)
6b.Polar vs. Non-Polar Covalent BondsNonpolar Polar
•Electrons shared equally
•Both atoms have similar electronegativity (affinity for electrons)
•Neither atom ends up with any charge
•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.
Non-Polar Polar
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
•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.
6c. Ionic Bonding
•Opposites attract!
•Significantly weaker than a covalent bond
•Can also occur between ionic molecules
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
Hydrogen Bonds
•Weak attraction between the partial charges of polar covalently bonded molecules
•In water, between O and H Means partial