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▫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 is filled first

•Electrons are distributed to each orbital in a shell before filling each orbital

•The outermost shell is called the valence shell

Electron Shells Con.•The first shell has only 1 orbital, so it can hold only 2 electrons

The 2nd/3rd Shell•Consists of 4

orbitals, so each shell can hold 8 electrons

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 D block (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

Draw the ions on your Whiteboard•Na+

•P3-

•Si2-

•H+

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)

Guidelines for Drawing Lewis Structures•Carbons make up core•Add Hs last (they can’t connect anything)

•Remember how many bonds each atom will make (using the periodic table

On your Whiteboard Draw or make:C3H8 CH3OH CF2O2H2

Draw : H3CCH2OH H2NCH2OH

Double Bonds•Use double bonds or triple bonds

when there aren’t enough atoms to form the proper number of bonds

•i.e. oxygen O=O, if it was a single bond O would not have the correct number of bonds

Draw and or make:N2 HCOOH HCN

Drawing ionic molecules (think about total valence electrons present)C should bring 4 valence electrons, N should bring 5. If neutral there would be 9.BUT if it’s – charged there should be 10 total

Notice atoms won’t form the correct number of bonds

NH4+ H3O+ OH -

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 are not shared equally

• 1 atom is more electronegative (O, F, N, Cl are the usual culprits)

• Electronegative atom ends up with a partial – charge since they have the electron more often

• Other atom ends up with a partial + charge as they are deprived of the electron

Non-Polar Polar

10. Ion Formation•Some atoms more

easily give up electrons (1st and 2nd columns) to end up with a full valence shell

•These electrons can be stripped by atoms in the 7th column (need 1 e)

•Forms ions

6c. Ionic Bonding•Opposites attract!

•Significantly weaker than a covalent bond

•Can also occur between ionic molecules

Hydrogen Bonds•Weak attraction between the partial

charges of polar covalently bonded molecules

•In water, between O and H Means partial