Chemistry Review

1. Chemistry Review You need to remember some basic things

2. 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)

4. 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”

5. 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

7. Draw on your Whiteboard • A neutral boron atom (for the nucleus you can just write B) • A neutral fluorine atom

8. 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

9. Draw • A neutral magnesium atom • A neutral phosphorus atom

10. 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

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

12. 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

13. 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)

14. The column can be used to figure out how many bonds an atom will normally form 4 3 2 1 0

15. 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)

16. 6b.Polar vs. Non-Polar Covalent Bonds Nonpolar 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.

17. Non-Polar Polar

18. 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

19. 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.

20. 6c. Ionic Bonding • Opposites attract! • Significantly weaker than a covalent bond • Can also occur between ionic molecules

22. 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

23. Hydrogen Bonds • Weak attraction between the partial charges of polar covalently bonded molecules • In water, between O and H Means partial