Chemistry Midterm Examination Review Sheet

with answers

 

General Information:

 

àThe Mid-term Examination will cover Chapters 1-4 and 7-10 plus selected topics from

 chapters 13 (phases of matter)  and 14 (molarity and solution stoichiometry).

àYou will be provided with a periodic table, solubility rules, and the activity series.

àYou will NOT be allowed to use a calculator.

à Multiple-choice questions on Scantron (75 questions = 75 points) will be followed by a

 section of free-response questions and  problems (20 questions = 95 points).

 

Chapter 1:  Introduction to Chemistry

 

TERMS:

Chemistry:  the branch of science that studies the composition and properties of matter.

Scientific Method:     Observationà Hypothesis à Experiment à Theory

 

Chapter 2:  Scientific Measurements

 

TERMS:

Qualitative vs. Quantitative measurements.

 

Scientific Notation

 

Significant Figures:    1)  All non-zero digits are significant.

2)      Zeros are significant only when they are between two non-zero digits

      or if they occur after a decimal point and after a non-zero digit.

3)      In multiplication or division, you cannot end up with more significant figures than you started with.

4)      In addition or subtraction, you cannot have more digits to the right of the decimal than you started with.

 

Percent           =          one quantity      x   100

                                    total sample

 

EXAMPLES:

 

1.                  Which of these measurements is NOT uncertain?

 

67 g                       4.5 mL                         457 K              22 people        6.7 x 1011 moles

 

2.                  Convert the following into standard notation:

 

            4.89 x 10-5 nm              _____0.0000489 nm_________________

 

                        7.9 x 102 mm                _____790 mm_________________

 

 

3.                  Convert the following into scientific notation:

 

0.00056 mg                  ______5.6 x 10-4 mg________________

 

78,341 kg                    ______7.8431 x 104 kg________________

 

 

4.                  Give the number of significant figures in each of these measurements:

 

            46 crucibles      __infinite_____

 

                        607 g               ___3____

 

                        0.00900 mL     ___3____

 

                        4.90 x 10-5 s     ___3____

 

5.                  Subtract 45.6 from 67.89 and round off the answer.

22.3

 

 

6.                  Multiply 240 by 2.10 and round off the answer.

5.0 x 102

 

 

7.                  An alloy contains 12 parts platinum and 4 parts gold.  Find the percents of platinum and gold in this alloy.

75% Pt

                                                      25% Au

 

 

 

                                   

Chapter 3:  The Metric System

 

TERMS:                                            

Metric              SI

Units:                          Temperature                 (°C)                 (K)                   K = °C  +  273

Length                          (m)                   (m)

Mass                            (g)                    (kg)

                                    Volume                        (L)                   (L)

                                    Time                             (s)                    (s)

                                    Amount                        (mol)                (mol)

                                    Energy                          (cal)                 (J)

 

Metric Unit Conversions

Prefix

Symbol

Conversion factor

kilo-

k

1 km = 1000 m

centi-

c

1 cm = 0.01 m

milli-

m

1 mm = 0.001 m

micro-

m

1 mm = 0.000001 m

nano-

n

1 nm = 0.000000001 m

 

Density           =           mass

                                    volume

 

Heat:  a measure of the total energy in a system.

 

Specific Heat              =          __heat__

                                               mass x ºC

 

 

8.                  Make the following unit conversions:

 

5 g into kg                                      ____.005 kg______

 

 

40 mL into L                                  ____0.04 L______

 

 

319 nm into cm                               ___0.0000319 cm_______

 

 

316 K into Celsius                          ___43°C_______

 

 

9.   A student is trying to determine the identity of an unknown metal. 

He finds that its mass is 20.0 g and its volume is 1.0 mL. 

The density of the metal is:                    ____2.0 x 101 g/mL_________

 

The metal that the student has is:            _____Gold___________

 

Known Densities of Metals

Gold                 19.3 g/cm3

Mercury           13.6 g/cm3

Lead                11.4 g/cm3

Silver                10.5 g/cm3

Copper                          8.9 g/cm3

Aluminum           2.7 g/cm3

 

 

 

 

      10. The specific heat of copper is 0.092 cal/g x °C.  If 378 calories are applied to 50 g of copper,

 how much will its temperature increase?

 

                                                            82°C

 

 

11.  The specific heat of iron is 0.10 cal/g x °C.  How many calories are required to raise the temperature of 40 g of iron from 298 K to 312 K?

 

 

56 cal

 

 

Chapter 4:  Matter and Energy

 

TERMS:

Element:  a pure substance that cannot be broken down by a chemical reaction.          

Compound:  a pure substance that can be broken down into two or more substances by a chemical reaction.   

Mixtures:  Homogeneous (the same properties throughout) vs. Heterogeneous (different properties throughout).

 

Physical States:  Solid (s), Liquid (l), Gas (g), Aqueous (aq).

 

Chemical vs. Physical properties and changes.

 

Metals:

Non-Metals:

Semi-metals or Metalloids:

 

Law of Conservation of Mass:  Matter can neither be created nor destroyed.

Law of Conservation of Energy:  Energy can neither be created nor destroyed.

 

Potential vs. Kinetic Energy

 

EXAMPLES:

 

  1. Give examples of the following:

 

Element                                    __________hydrogen_____________________

 

Compound                               ___________water____________________

 

Homogeneous mixture  __________air_____________________

 

Heterogeneous mixture ___________salad____________________

 

 

  1. How many atoms are in a glucose molecule, C6H12O6?

24 atoms

 

  1. Name the following phase changes:

 

Solid à  Liquid            _____________melting____________________

 

Liquid à  Gas              ____________vaporization_____________________

 

Gas à  Solid               _____________deposition____________________

 

Solid  à  Gas              ______________sublimation___________________

 

Gas à  Liquid              _____________condensation____________________

 

Liquid à  Solid            ______________freezing___________________

 

  1. Consider a glass of water.  What are three of its physical properties?  What are three of its chemical properties?

Physical:  clear, liquid, density of 1g/mL

 

      Chemical:  reacts with active metals to produce hydrogen gas, can be electrolyzed to produce hydrogen and oxygen, is non-flammable, reacts with carbon dioxide in photosynthesis.

 

 

16.  What are three properties of a gas?  Three properties of a liquid?  Three properties of a solid?

 

Gases:  indefinite shape, indefinite volume, low density, compressible, mix uniformly

      Liquids:  indefinite shape, definite volume, higher density than gases, incompressible

      Solids:  definite shape, definite volume, usually higher density than liquids, incompressible, do not mix uniformly.

 

Chapter 7:  Nomenclature

 

TERMS:

Ions:    Cation = a positively charged ion.

            Anion = a negatively charged ion.

            Ions are formed by gaining or losing electrons.

 

àYou must know the name, formula and charge for any given monatomic or polyatomic ion (see attached list or your flashcards).ß

 

Making Neutral Compounds: You will be asked to make ionic compounds using cations and anions.  You must use as many cations and anions as is necessary to balance the charge.  Example:  Magnesium is a 2+ cation.  Nitrate is a 1- anion.  Two nitrates are required to balance the charge on one magnesium, therefore the chemical formula is Mg(NO3)2.

 

Nomenclature:

I.                   Binary Ionic Compounds (metal + non-metal)

Name the metal cation first.

Name the non-metal anion second, change the suffix to –ide.

Example:  CaCl2 is calcium chloride.

 

II.        Binary Ionic Compounds with a metal of variable charge (metal + non-metal)

Name the metal cation first.

Give the charge on the metal using Roman Numerals (I), (II), (III), (IV), (V).

Name the non-metal anion second, change the suffix to –ide.

                        Example:  Fe2O3 is iron (III) oxide.

 

III.       Ternary Ionic Compounds (containing polyatomic ions.)

Name the cation first, using Roman Numerals if necessary.

Name the anion second, do not change the suffix if it is a polyatomic ion.

                        Examples:  Ni3(PO4)2  is nickel (II) phosphate;  NH4I is ammonium iodide.

 

IV.       Binary Molecular Compounds (non-metal + non-metal)

Name the most metallic element first (the one furthest to the left table).

Name the least metallic element second, change the suffix to –ide.

Use prefixes to indicated the number and type of each atom (mono-, di-, tri-, tetra- penta-, hexa-, hepta-, octa-, nona-).

                        Example:  SF6 is sulfur hexafluoride (or monosulfur hexafluoride).

 

V.        Acids

            Acids can often be recognized by hydrogen beginning a formula.

            You should be able to name these five common acids: 

            HCl                  HNO3              H2SO4              H3PO4              HC2H3O2

 

VI.       Bases

            Bases can often be recognized by hydrogen ending a formula.

            Bases are named just like other ternary ionic compounds (see III);

            the exception is NH3, which is ammonia.

 

EXAMPLES:

 

17.       Name the following compounds:

 

            CCl4                 ______carbon tetrachloride_________________________

 

            MgCr2O7         _______magnesium dichromate________________________

 

            NiSO4              _______nickel (II) sulfate________________________

 

            NaOH              ________sodium hydroxide_______________________

 

            H2SO4                  _________sulfuric acid______________________

 

18.       Write formulae for the following compounds:

 

Zinc sulfite                                _____ZnSO3______________

 

Silicon dioxide              _____SiO2______________

 

Iron (II) phosphide                   ______Fe3P2_____________

 

Potassium dichromate               _____K2Cr2O7______________

 

Ammonium nitrate                     _____NH4NO3______________

 

 

Chapter 8:  Chemical Reactions

 

TERMS:

Balancing Equations:You should be able to predict and balance equations:  make sure there

                        are the same number of atoms of each type on each side by using coefficients.

 

Types of Reactions:

            A.        Combination Reactions

 

                        Two or more reactants à  One product

                                    2Mg(s) + N2(g)  à  2MgN(s)

 

B.                 Decomposition Reactions

 

One reactant à  Two or more products

                            2NaHCO3(s)  à  Na2CO3(s)  +  H2O(l)  +  CO2(g)

    Na2CO3(s)  à  Na2O(s)  +  CO2(g)

    2NaO(s)  à  2Na(s)  +  O2(g)

 

C.                 Single-replacement Reactions

 

One element replaces another element in a compound.

                        2Na(s)  +  ZnCl2(aq)  à  2NaCl(aq)  +  Zn(s)

                        Predictable:  Consult the activity series (attached to this sheet). 

More active metals can replace less active ones.

 

D.                Double-replacement Reactions

 

Two ionic compounds exchange partner ions, forming two new compounds.

2AgNO3(aq)  +  CuCl2(aq)  à  2AgCl(s)  +  Cu(NO3)2(aq)

Predictable:  Swap the ions to make two neutral compounds.

Consult the solubility rules (attached to this sheet) to determine if the products are solid or aqueous.

 

 

EXAMPLES:

 

19.       Predict the products of the following reactions and then balance the chemical equations.

Indicate the type of chemical reaction for each.  Indicate physical states for double-replacement reactions (consult the Solubility Rules).

 

a.                   4Na(s)  +  O2(g)  à  2Na2O

 

 

 

Type:  ___combination______

                                                                                                           

b.                  Ag(s)  +  H2O(l)  à        NR

 

 

 

Type:  ______single replacement______

 

c.                   MgBr2(aq)  +  K2SO4(aq)  à       2KBr(aq)  +  MgSO4

 

 

 

 

Type:  ____double replacement______

 

d.                  2Al(s)  +  3Pb(NO3)2(aq)  à       2Al(NO3)3  +  3Pb

 

 

 

 

Type:  ______single replacement_____

 

 

 

e.                   2MgO(s)  à     2Mg  +  O2

 

 

 

 

Type:  ______decomposition_______

 

f.          H2O(aq)  +  KMnO4(aq)  à

 

 

 

 

                                                                        Type:  reactions of water will not be on the exam

 

Chapter 9:  The Mole Concept

 

TERMS:

   6.02 x 1023    atoms or molecules       (Avogadro’s number)

                                    mole

  

22.4