Download Chemistry Review Study Guide 1. Atomic number (Z) = number of P

Survey
yes no Was this document useful for you?
   Thank you for your participation!

* Your assessment is very important for improving the work of artificial intelligence, which forms the content of this project

Document related concepts
no text concepts found
Transcript
Chemistry Review Study Guide
1. Atomic number (Z) = number of P or E
2. Atomic mass (A) = number of P + N
3. Atomic mass is same as molar mass
and is in grams
4. The reason why mass is not a whole
number is because it is an avg. of the mass of
the isotopes.
5. Order of orbitals (1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2)
6. valence electrons = outer electrons (group 1 = 1 valence electron, etc…, ignore the transition metals)
7. An element with an electron configuration of 1s2 2s2 2p4 has six valence electrons because there are six electrons in
the second energy level.
8. non-metal – non-metal = covalent (mono, di, tri, tetra, penta, etc…)
9. metal – non-metal = ionic (name of metal + ide (monatomic ions) or name of polyatomic ion (on chart)
10. covalent – gases, liquids, solids, made of molecules. Low melting/boiling points. Poor electrical conductors (brittle)
Ionic - crystalline solids made of ions, high melting/boiling points. Good electrical conductors in solution
(An ion is an atom that has gained or lost an electron)
Metallic - generally high melting/boiling points, conducts electricity as solids, conducts heat, ductile, malleable.
11. Ranking of bond strength from strongest to weakest
1.Covalent bonding (bonding that happends by the sharing of electrons).
2.Ionic-bonding (most commonly found in salts, it forms by the losing and gaining electrons ** electrons are not shared in this type of bonding).
3.Metallic bonding (between metals only).
4.H-bonding (this type of bonding tends to form with Nitrogen, Oxigen and FLuorine)
5.Dipole-Dipole(forc e that exist because of the interaction of dipoles on polar molecules in close contact).
6.London Dispersion (simultaneously Dipole-dipole moments)
12. Bond length and bond strength. C-C (longer bond, weaker bond strength) C=C (shorter bond length, stronger bond)
13. Periodic table trends
1. Electronegativity – over and up (increases)
2. First ionization energy – over and up (increases)
3. atomic radius (size)- across table (decreases), down table (increases)
4. Ionic radius – loses electrons (metals)- gets smaller , gains electrons (non-metals) – gets bigger
5. Reactivity of metals (increases as you go down a column), non-metals (increases as you go up a column)
6. alkali metals and halogens are most reactive (most likely to gain or lose electrons)
7. Same group = similar properties (i.e. alkali metals behave in a similar way)
14. Balancing equations – do an atom inventory, add coefficients (numbers at front of formula), don’t forget to multiply
numbers in parentheses (i.e. Mg(OH)2 = 2 O and 2 H atoms)
15. Don’t forget to criss-cross charges to make a formula (i.e. Iron (III) Oxide. Fe3+
O2- = Fe2O3
16. Molar mass – Mass is in grams (find it on periodic table) Add elements together. (i.e. NaCl = 23 + 35.45 = 58.45g)
17. Mole conversions.
1. start with what you know.
2. Cannot go from mass to molecules or atoms, put moles in the middle
18. Percent composition = mass of element/mass of compound
19. Empirical formula 1. cross out percent and put grams
2. convert grams to moles
3. Divide by the smallest number
4. Multiply by 2 if you get a half number
20. Molecular formula 1. Divide mass of unknown compound by mass of empirical formula
2. Use this number to multiply the subscripts
21. Chemical change – cannot go back
22. Kinetic energy
Solids – low kinetic energy, low movement, dense
Liquids – higher kinetic energy, more movement, less densely packed
Gases – highest kinetic energy, most movement, least densely packed
23. Electrolytes - NaCl  Na+ + Cl- (positive and negative ions in solution are electrolytes)- come from ionic compounds
and can conduct electricity
24. Arrhenious acid/base (acid = H+) (base = OH-)
25. Acid + Base  salt and water (neutralization)
26. Things that speed up chemical reactions
1. heat
2. a catalyst (enzymes are an example) – work by decreasing activation energy
3. increasing surface area (little pieces vs. big chunk)
27. Electromagnetic Radiation
1. shorter wavelength = higher frequency and more energy, longer wavelength = lower frequency/energy
2. gamma rays (no mass, penetrates most), alpha (2 protons, 2 neutrons, penetrates only air), beta (electron,
penetrates paper)
28. Heat of fusion – use when dealing with freezing or melting (find on reference table)
Heat of vaporization – use when dealing with vaporization or condensation (on reference table)
-
When you add energy to liquid-vapor system = vapor increases
When you take away energy from a solid-liquid system = solid increases (freezing)
29.
30.
31. Moles of a gas = 22.4L per mole (on reference table!)
32. Gas laws (look at units given) are on the reference table!
33. Potential energy diagram
34. Equilibrium
1. The mass action expression consists of the product of the products, each raised to the power given by
the coefficient in the balanced chemical equation, over the product of the reactants, each raised to the
power given by the coefficient in the balanced chemical equation. This mass action expression is set
equal to the equilibrium constant, Keq:
2. COCl2 (g)  CO (g) + Cl2 (g)
If you add more COCl2, reaction will produce more products, if you remove products, it will produce more
product, if you remove COCl2, it will go in the reverse and make more COCl2. Think about balancing!
35. pH = -log [H+], pOH –log [OH-]
pH + pOH = 14 (find pH by subtracting pOH from 14)
36.