Chemistry 1032 Final Exam Spring 2020Dr. Santra

Supplemental Instruction Leader: JadaApril 18th, 2020

This is not a full extent of the Exam. Please continue to review all previous materials, worksheets and lecture notes

Chapter 1: The Chemical World

1. Define chemistry:

2. What is the importance of the Scientific Method:

3. What is the difference between theories and laws: (give an example of each)

4. What is the law of conservation of mass and who was it developed by?

5. The atomic theory was developed by John Dalton, explain how this theory explained the law of conservation of mass?

Chapter 2: Measurement and Problem Solving

1. Uncertainty is indicated by what digit?

2. The greater the precision of the measurement, the (lower/ greater) the number of significant figures

3. List the rules on reporting significant figures:

4. How can we tell that we are working with an exact number?

5. Report the number of significant figures: 1.230

12,000

1 mole = 6.022 x 10^23

2371

0.0047

6. Define the multiplication/division rule for reporting significant figures:

7. Add 5.892 to 6.10

8. List the Standards for time, mass, length, and temperature

9. What is the difference between mass and weight?

Chapter 3: Matter and Energy

1. Discuss/write information and properties for the states of matter:Solids Liquids Gases

Matter

2. Complete the image

3. State the difference and possible examples of chemical and physical properties:

4. How are chemical and physical properties different from changes?

5. State the law of conservation of energy:

6. Draw the graphs for endothermic and exothermic

7. A student must use 225mL of hot water in a lab procedure. Calculate the amount of heat required to raise the temperature of 225mL of water from 20.0 degrees Celsius to 100 degrees Celsius.

9. A 40.0 g sample of ethanol releases 2952 J as it cools from 50.0 degrees celisus. Calculate the final temperature of the ethanol.

Chapter 4: Atoms and Elements

1. What are the three parts of Dalton’s atomic theory:

2. What are important take aways from these two experimentsThomson Rutherford

3. Write and state the meaning of chemical symbol with its number in the chemical way using letters A, x, z

4. What is the importance of Dmitri Mendeleev5. Recurring patterns tend to be shown in (columns/rows)

6. What do we use on the periodic table to read ionic charges?

7. T/F isotopes have the same number of neutrons and different number of protons

8. By looking at an isotope how can we find the number of neutrons?

9. Rubidium is a soft silvery metal and is known to have two isotopes: Rb 85 and Rb 87. If the abundance of Rb 85 is 72.2% and the abundance of Rb 87 is 27.8%, how can we find the average?

Chapter 5: Molecules and Compounds

1. What law states that all samples of a given compound have the same proportions of their constituent elements.

2. Write the ionic formulas for the following chemical names:Beryllium oxide

Manganese (IV) nitride

Vanadium (III) selenide

Ammonium nitrate

Copper (II) chloride

3. Write any helpful information on thisStructural Formula Empirical Formula Molecular formula

4. What is the difference between molecular/ atomic elements

5. Give an example of a molecular compound and an ionic compound

6. Naming (ionic/ molecular) compounds means that we use prefixes

7. Caffeine has the following percent composition: carbon 49.48%, hydrogen 5.19%, oxygen 16.48% and nitrogen 28.85%. It is known to have a mass of 194.19. What is the empirical formula?

8. Benzoic acid contains 68.8% carbon, 4.95% hydrogen and 26.25% oxygen. Find the empirical formula.

Chapter 7: Measurement and Problem Solving

1. ___ AgI + ___ Na2S ___Ag2S + ___ NaI

2. __Ba3N2 + ___ H2O ___ Ba(OH)2 + ___ NH3

3. ____ CaCl2 + ___ Na3PO4. ____ Ca3(PO4)2 + ___NaCl

4. Predict the precipitate reaction:

Hydrochloric acid + silver (I) nitrate

Sulfuric Acid + barium nitrate

5. From the two above chemical reactions determine the net ionic formula for both.

6. Classifying Chemical Reactions; place any helpful information

Precipitation Acid-Base Gas Evolution Oxidation-reduction

Chapter 8: Quantities in Chemical Reactions

1. Identify the limiting and excess reactants when 25 mL of nitrogen reacts with 25mL of hydrogen at STP. How many liters of ammonia gas are formed in this reaction?

___N2 + ___H2 ___ NH3

2. 50.0 mL of 2.00 M H2SO4 react with 75.0 mL of 2.00 M NaOH. Identify limiting and reactant. How many grams of Na2SO4 will form?

____H2SO4 + ___ NaOH ___ Na2SO4 + ____H2O

3.If you were to attempt to make 45.0g of methane from CO2 and H2O, (with O2 also being produced). How much heat could be absorbed during the reaction?

___ CO2 + ___ H2O __ CH4 + ___ O2. Delta H= +890.4 kJ

4. H2O (g) has an enthalpy of -242 kJ/mol, O3(g) has an enthalpy of +143 kJ/mol and H2(g) has an enthalpy of 0 kJ/mol. With this information given and the chemical equation find the enthalpy of reaction.

3H2 (g) + O3 (g) 3H2O(g)

5. How many grams of O2 will be formed from 3.76 grams of KClO3?

__KClO 3 ___ KCl + ___ O2

6. If enough H2O is reacted to produce 3.40 moles of H2, how many moles of O2 must have been made?

__H2O __ H2 + ___ O2

Chapter 10:

1. Draw Lewis structure for:

C3H8

CH3OH

CH3Cl

BeF2

2. Fill in Missing Information regarding shape and compoundsCompound Geometry Shape Atoms on Central atom

SF6

BrF5 Square pyramid Y

CO2 N

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Chapter 11

1. Fill in the chart with informationBoyles Law Charles Law Combined Gas

LawAvogadro’s Law Ideal Gas Law

CHAPTERS 12-14 WILL BE FUTHER DISCUSSED IN SESSION AND STUDY UNION