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Chemistry Review

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 Neutrons- neutral particles Together form the atomic nucleus Electrons- negatively charged particles

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Chemistry Review

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  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 • 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)

  3. 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”

  4. 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

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

  6. The 2nd/3rd Shell • Consists of 4 orbitals, so each shell can hold 8 electrons

  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 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

  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. Draw the ions on your Whiteboard • Na+ • P3- • Si2- • H+

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

  13. 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

  14. 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)

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

  16. 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)

  17. 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

  18. On your Whiteboard Draw or make:C3H8 CH3OH CF2O2H2

  19. Draw : H3CCH2OH H2NCH2OH

  20. 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

  21. Draw and or make:N2 HCOOH HCN

  22. Drawing ionic molecules (think about total valence electrons present) NH4+ H3O+ OH - 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

  23. 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 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

  24. Non-Polar Polar

  25. 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

  26. 6c. Ionic Bonding • Opposites attract! • Significantly weaker than a covalent bond • Can also occur between ionic molecules

  27. 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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