Chemistry Final Review

2nd Semester, 2014-2015

Unit 1: Basics

• How many sig figs are in the following numbers?

• 2300• 2• 314• 3• 2.06• 3

• 0.0025• 2• 9.001• 4• 6.02x1023

• 3• 587.0• 4

Unit 1: BasicsDiscuss 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 = 6.02x1023 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• 2.71x104

• 27100• 6.4x10-3

• 0.0064• 2687• 2.687x103

• 0.012654• 1.2654x10-2

Unit 1: Basics

• What is the equation for density?• Density = mass volume• What are the common units for density?• g/mL or g/cm3

Unit 1: Basics

• Give an example of an element and an example of a compound.

• Element: He, C, Mn• Compound: CO2, 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: Atoms & Periodic Table

• Complete the following table:Element/ion

Atomic number

Atomic mass

Protons Neutrons electrons

Fe

Cl-

K+

26 55.85 26 2630

17 17 181835.45

19 1819 2039.10

Unit 2: 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 2: Atoms & Periodic Table

• The halogens make a charge of ____ when they become ions.

• -1• 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 2: Atoms & Periodic Table

• The halogens make a charge of ____ when they become ions.

• -1• 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 2: 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 2: Atoms & Periodic Table

• Complete the following:• Type of Decay• ________ 99m

43Tc 9943Tc + ______

• ________ 24795Am 0

-1e + _____

• ________ 17593Np 4

2He + ____

00γ

24796Cm

17191Pa

Gamma

Beta

Alpha

Unit 2: Atoms & Periodic Table

• Draw a Bohr model for Beryllium.

• Draw a Bohr model for Silicon.

Unit 2: Atoms & Periodic Table

• There will be no electron configuration on the final (s, p, d, f).

Unit 3: 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 3: Compounds & Bonding

Name the following compounds:

• MgO• Magnesium oxide• AlF3

• Aluminum fluoride• NiSO4

• Nickel (II) sulfate• FeCl2

• Iron (II) chloride

• N2O5

• Dinitrogen pentoxide• SF4

• Sulfur tetrafluoride

Unit 3: 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 3: Compounds & Bonding

• There will be no molecular shapes on the final.

Unit 3: 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 4: Energy Transfer

• Draw a heating curve for water.

Unit 4: Energy Transfer

• Label the Q-equations for each section.

Q=mcΔT

Q=mHv

Q=mHf

Unit 4: Energy Transfer

• Find the value for Hf, c, and Hv in your data book.

• Hf = 79.72 cal/g• C = 1.0 cal/g°C• Hv = 539.4 cal/g

Unit 7: SolutionsUse 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: SolutionsUse 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 (NH3) 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

Unit 8: EquilibriumDescribe a system at

equilibrium for each case:liquid and gas phases of the same substance in a closed

container reactants and products in a

chemical systema saturated solution

33A.liquid + gas in a closed container with equal # of molecules condensing as are evaporating; rate of condensing = rate of evaporatingrate of forward reaction is equal to rate of reverse reactionrate of dissolving is equal to rate of precipitation or crystal growth/formation

Unit 8: EquilibriumUse LeChatelier’s Principle and

the following reaction:

PCl3(g)+3NH3(g)P(NH2)3(g)+3HCl(g)

If [PCl3] increases then [HCl]…If pressure increases then [NH3]…

If [P(NH2)3] increases then [HCl]…

35A.

If [PCl3] is increased then [HCl]… increases.

If pressure is increased then [NH3]…is constant.

If [P(NH2)3] is increased then [HCl]…decreases.

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.

0.062 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 1024 atoms

47. How much space does 64 g of oxygen

occupy?

47A.

45 L O2

48. What is the mass of 3.75 x 1022

molecules of CO2?

48A.

2.74 g CO2

49. In ammonia production, nitrogen and hydrogen are synthesized into ammonia (NH3). 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.

N2 (g) + 3H2 (g) 2NH3(g)

 1.8 kg NH3

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.

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 CaSO4

measured in the experiment was 3.99 g, what is the percent error?

f) What is the percent yield?

50A.a) CuSO4 (aq) + Ca(NO3)2 (aq)CaSO4 (s) + Cu(NO3)2 (aq)

b)Cu2+ +SO42-+ Ca2+ + 2NO3

- CaSO4 (s) +Cu2+ + 2NO3-

c)SO42- + Ca2+ CaSO4(s)

d)4.3 g CaSO4

e)-7.2%f)93%