Chemistry Final Review 2 nd Semester,

Unit 1: Basics How many sig figs are in the following numbers? x

Unit 1: Basics Discuss the following in terms of accuracy and precision

Unit 1: Basics List the three mole conversions you learned at the beginning of the year. 1 mol = 6.02x10 23 particles (atoms, molecules, ions, etc) 1 mol = molar mass in grams from the periodic table 1 mol = 22.4 L of GAS at STP  see how that makes more sense now ?

Unit 1: Basics Write the following in standard or scientific notation 2.71x x x x10 -2

Unit 1: Basics What is the equation for density? Density = mass volume What are the common units for density? g/mL or g/cm 3

Unit 1: Basics Give an example of an element and an example of a compound. Element: He, C, Mn Compound: CO 2, NaCl, etc (anything with more than one element)

Unit 1: Basics Write the equation for percent yield. % yield = actual x 100 expected Write the equation for percent error. % error = (actual-expected) x 100 expected

Unit 2: Energy Transfer Draw a heating curve for water.

Unit 2: Energy Transfer Label the Q-equations for each section. Q=mcΔT Q=mHv Q=mHf

Unit 2: Energy Transfer Find the value for Hf, c, and Hv in your data book. Hf = cal/g C = 1.0 cal/g°C Hv = cal/g

Unit 3: Atoms & Periodic Table Complete the following table: Element/ ion Atomic number Atomic mass ProtonsNeutronselectrons Fe Cl - K+K

Unit 3: Atoms & Periodic Table Positive ions form when: Atoms lose electrons (usually metals, on left of table) Negative ions form when: Atoms gain electrons (usually nonmetals, on right of table) Why do atoms form ions? To become more stable, get the configuration of a noble gas.

Unit 3: Atoms & Periodic Table The halogens make a charge of ____ when they become ions. The alkali metals make a charge of ___ when they become ions. +1 The alkali earth metals make a charge of ___ when they become ions. +2

Unit 3: Atoms & Periodic Table The halogens make a charge of ____ when they become ions. The alkali metals make a charge of ___ when they become ions. +1 The alkali earth metals make a charge of ___ when they become ions. +2

Unit 3: Atoms & Periodic Table What are the three types of nuclear decay? Alpha, Beta, Gamma What type of particle does each emit? Alpha = helium nucleus (2 protons, mass of 4) Beta = electron (no mass, -1 proton) Gamma = high energy (no mass, no proton change, but product is more stable)

Unit 3: Atoms & Periodic Table Complete the following: Type of Decay ________ 99m 43 Tc  Tc + ______ ________ Am  0 -1 e + _____ ________ Np  4 2 He + ____ 00γ00γ Cm Pa Gamma Beta Alpha

Unit 4: Atoms & Periodic Table Draw a Bohr model for Beryllium. Draw a Bohr model for Silicon.

Unit 4: Compounds & Bonding Which types of elements participate in ionic bonding? Metals (+) and non-metals (-) (also polyatomic ions) Which types of elements participate in covalent bonding? Non-metals and non-metals (they share electrons instead of charges sticking together)

Unit 4: Compounds & Bonding Name the following compounds: MgO Magnesium oxide AlF 3 Aluminum fluoride NiSO 4 Nickel (II) sulfate FeCl 2 Iron (II) chloride N 2 O 5 Dinitrogen pentoxide SF 4 Sulfur tetrafluoride

Unit 4: Compounds & Bonding Describe a polar bond. Covalent bond (non-metal and non-metal) in which electrons are shared UNEVENLY (one atom is more electronegative than the other). Draw a water molecule and show its polarity.

Unit 4: Compounds & Bonding List and describe the 4 types of Intermolecular forces (IMF). Dispersion forces (weakest, between nonpolar molecules) Dipole-dipole interaction (stronger, between polar molecules) Hydrogen-bonding (STRONG, between molecules with N, O, or F bonded to H, explains water’s high boiling point) Ion-molecule interaction (strongest, how water dissolves salt – pulls apart the ions)

Unit 5: Chemical Reactions Write the generic equation for all 6 types of chemical reactions Synthesis: A + B  AB Decomposition: AB  A + B Single Replacement: A + CB  AB + C (remember, there are 3 sub-types of Single Replacement) Double Replacement: AB + CD  AD + CB Combustion: Fuel + __O 2  __CO 2 + __H 2 O Acid/Base: HD + COH  CD + H 2 O

Unit 5: Chemical Reactions What table do you use to decide if a single replacement reaction happens? What do you look for? Table N Strong metals are on the bottom…if the LONE metal is closer to the bottom it will replace the metal in a compound. Strong halogens are at the top…if the LONE halogen is closer to the top, it will replace the halogen in a compound. Hydrogen follows the metal rule.

Unit 5: Chemical Reactions List the 7 diatomic elements. H, O, N, Cl, Br, I, F

Unit 5: Chemical Reactions What table do you use to decide if a double replacement reaction happens? What do you look for? Table E (solubility) Both reactants must be aqueous (soluble), AT LEAST one of the products must be NOT AQUEOUS (solid, liquid, or gas)

Unit 5: Chemical Reactions Single and Double Replacement – predict the products and write the balance equation 1. silver nitrate + nickel  2AgNO 3 (aq) + Ni (s)  Ni(NO 3 ) 2 (aq) + 2Ag (s) 2. lead + zinc acetate  Pb + Zn(C 2 H 3 O 2 ) 2  N.R. 3. NaOH + CaBr 2  2 NaOH (aq) + CaBr 2 (aq)  2 NaBr (aq) + Ca(OH) 2 (s) 4. Pb(NO 3 ) 2 + HCl  Pb(NO 3 ) 2 (aq) + 2 HCl (aq)  2 HNO 3 (aq) + PbCl 2 (s)

Unit 5: Chemical Reactions 1. Consider this molecular (normal) equation: H 3 PO 4 (aq) + 3KOH (aq)  K 3 PO 4 (aq) + 3H 2 O (l) 2. Write the complete ionic equation: 3H + (aq) + PO 4 3- (aq) + 3K + (aq) + 3OH - (aq)  3K + (aq) + PO 4 3- (aq) + 3H 2 O(l) 3. Write the net ionic equation: 3H + (aq) + 3OH - (aq)  3H 2 O(l) 4. What are the spectator ions? PO 4 3- (aq), 3K + (aq)

Unit 6: Stoichiometry 49. In ammonia production, nitrogen and hydrogen are synthesized into ammonia (NH 3 ). What mass of ammonia will be produced from 1.5 kg of nitrogen assuming that hydrogen is in excess? (Hint: first write the complete, balanced equation)

49A. N 2 (g) + 3H 2 (g)  2NH 3 (g) 1.8 kg NH 3

Unit 6: Stoichiometry 50. A solution made from 5.0 g of copper (II) sulfate is mixed with a solution containing excess calcium nitrate. A precipitate of calcium sulfate is formed. a) Write the complete, balanced equation for the reaction. b) Write the complete ionic equation for the reaction. c) Write the net ionic equation for the reaction.

Unit 6: Stoichiometry 50. A solution made from 5.0 g of copper (II) sulfate is mixed with a solution containing excess calcium nitrate. A precipitate of calcium sulfate is formed. Continued: d) How much (mass) calcium sulfate is expected? e) If the amount of CaSO 4 measured in the experiment was 3.99 g, what is the percent error? f) What is the percent yield?

50A. a) CuSO 4 (aq) + Ca(NO 3 ) 2 (aq)  CaSO 4 (s) + Cu(NO 3 ) 2 (aq) b)Cu 2+ +SO Ca NO 3 -  CaSO 4 (s) +Cu NO 3 - c)SO Ca 2+  CaSO 4 (s) d)4.3 g CaSO 4 e)-7.2% f)93%

Unit 7: Solutions Use Table E to determine the solubility of each substance: ammonium chloride barium carbonate silver iodide mercury (II) bromide

ammonium chloride soluble barium carbonate nearly insoluble silver iodide nearly insoluble mercury (II) bromide slightly soluble

Unit 7: Solutions Use Table D: 1. How many grams of sodium nitrate will dissolve in 100 g of water at 25 C? 2. How many grams of ammonia (NH 3 ) will dissolve in 100 g of water at 100C? 3. If 140 g of KI is dissolved in 100 g of water at 30 C, is the solution saturated, supersaturated, or unsaturated?

92 g 7 g unsaturated

The last seven questions will give you practice doing basic chemistry conversions and doing stoichiometry.

44. How many moles are in 3.6 g of sodium chloride?

44A mol

45. What is the volume of 2.3 moles of oxygen? (at STP)

45A. 52 L

46. How many atoms of mercury are in 5.0 moles of mercury?

46A. 3.0 x atoms

47. How much space does 64 g of oxygen occupy?

47A. 45 L O 2

48. What is the mass of 3.75 x molecules of CO 2 ?

48A g CO 2