Chemistry Review You 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

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 Bonds NonpolarPolar 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-PolarPolar

10. Ion Formation Some atoms more easily give up e- (1 st and 2 nd columns) to get a full valence shell They commonly form bonds with atoms in the 6 th & 7 th 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