Healthcare and Life Sciences

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  • Histamines allergic reaction



    An allergic reaction prompts the body to produce histamines.

    • Summarize the effects histamines can have on the body.
    • What is the purpose of an antihistamine?
    • Give at least 2 examples of an antihistamine.
    • What functional groups are present?
    • What are the benefits and dangers of using them?
    • There are several books and articles that have been published regarding “Biblical Healing”. Find one of these articles or books and summarize it. Be sure to cite it as a source in APA format.
    • Do you believe that an allergic reaction can be healed through faith alone? Does it mean you have weak faith if you do not?

    Post your answers below and respond to at least two classmates’ posts with substantive responses.

  • EXPERIMENT 3 Properties of Solutions- 2 day lab


    EXPERIMENT 3 Properties of Solutions

    Student Kit Materials:

    Lab Kit:

    Item Item # Price
    Electronic Pocket Scale, 100 x 0.01 g (already purchased)

    3 x Pyrex® Screw-Cap Bacteriological Culture Tubes, 13 x 100 mm, 9 mL 731501 1.70 x 3
    Graduated Cylinder, Polypropylene, 10 mL (already purchased)

    Safety Goggles 646706C 3.00
    Beaker, 100 mL, 250 mL, 500 mL (already purchased)


    Student Responsibility:

    Fine NaCl crystals
    Small Bowl (weighs less than 60 g)
    Coarse NaCl Crystals (Rock Salt)
    Coffee Mug
    Cooking Oil
    Iodine Tincture (Available at any drug store ~$4)- 1 oz.
    Isopropyl Alcohol (Rubbing Alcohol)- 16 oz.
    Mineral Oil- 16 oz bottle
    Table Sugar (Sucrose)


    Properties of Solutions- 2 day lab

    Materials: Solids: sodium chloride (NaCl) fine and coarse crystals, sucrose (table sugar, C12H22O11).  Solutions: saturated sodium chloride (NaCl) and iodine tincture. Liquids: Cooking oil, mineral oil, isopropyl alcohol (rubbing alcohol) (C3H7OH). 


    A solution is a homogeneous mixture in which one (or more) substance is dissolved in one or more other substances. For simplicity, we will investigate solutions consisting of one substance dissolved in one other substance. The solute is the substance that is dissolved. The solvent is the dissolving substance (also called the dissolving medium). The solvent is the substance present in the greater amount. The name of the solution is based on the name of the solute. In a “sodium chloride solution,” sodium chloride is the solute and water is the solvent.

    A solution has a variable composition (ratio of solute to solvent) as more or less of the solute may be dissolved in a given amount of a solvent. A solution is a homogeneous mixture because the solute remains uniformly dispersed throughout the solution after mixing. In other words, there is only one phase.

    Formation of a solution depends on the nature of the solute and the solvent.  In general water, composed of polar molecules, is a good solvent for many ionic compounds which are composed of ions. Nonpolar solvents like benzene and hexane, composed of nonpolar molecules, are good solvents for other nonpolar compounds. The phrase in chemistry that captures the principle expressed in the preceding sentences is “Like dissolves like.” In other words, substances will dissolve other substances of similar polarity. In this experiment, we apply this principle to solids dissolved or not dissolved in liquids and liquids which are miscible or immiscible in other liquids.

    The rate of dissolving a solute depends on:

    1. the particle size of the solute.
    2. whether the system is agitated (mixed) or not and how vigorously it is mixed.
    3. the temperature of the system.
    4. the concentration of the solute in the system.

    The following terms are important:

    Solubility is the amount of solute that will dissolve in a given amount of solvent at a given temperature. The terms used to describe solubility of a solid solute are: insoluble, slightly soluble, and soluble. When a liquid is added to another liquid, the term miscible is used to indicate that the liquid present in a lesser amount formed a homogeneous mixture with the liquid present in the greater. When the liquids form layers, the term immiscible is used.

    Concentration of a solution may be expressed in words or mathematically as a ratio. It expresses the ratio of solute to solvent in a solution. The following word expressions are common and important to know:

    Dilute solution:  A solution that contains a small amount of solute per unit volume of solution.

    Concentrated solution:  A solution that contains a large amount of solute per unit volume of solution.

    Saturated solution:  A solution that contains the maximum amount of dissolved solute possible at a given temperature. In the laboratory, saturated solutions appear as homogeneous liquids on top of undissolved solid particles. The liquid, called a “supernatant liquid” because it is over (above) the undissolved solid, is in dynamic equilibrium with the undissolved solid.

    Unsaturated solution:  A solution containing less solute per unit volume than the corresponding saturated solution

    Supersaturated solution:  A solution that contains more dissolved solute than is normally present in the corresponding saturated solution at a given temperature. The supersaturated solution is not stable and will crystallize if the solution is disturbed.

    The following mathematical expressions of concentration are important to know

    Mass percent  =   x   100%

    Molarity  =


    1. Concentration of a saturated solution (Plan ahead- this needs to set overnight)
    1. Preparation of the saturated NaCl solution: Weigh 36.6 g of fine NaCl crystals and place the crystals in a coffee mug.  Add 100.0 mL of water to the mug.  Place the mug in the microwave and heat the solution until it boils.  Stir until all (or the great majority) of the crystals dissolve.  Let the solution cool.
    2. Weigh a small bowl (needs to weigh no more than 50-60 g).  Record the mass.
    3. Transfer 6 mL of the saturated NaCl solution to the bowl.  Make sure to only transfer liquid- leave any solid residue behind.  Weigh the bowl with the solution. Record the mass.
    4. Allow the bowl to set out overnight or until all the liquid has evaporated.  Reweigh. Add water to the solid to dissolve it and dispose of it in the sink.

    B.  Relative solubility of a solute in two solvents

    1. Add 2mL of cooking oil and 2mL of water to a test tube, stopper it, and shake gently for 5 seconds. Allow the layers to separate and note which liquid has the greater density.
    2. Add 5mL of iodine tincture to the same test tube, note the color of each layer. Stopper and shake gently for 20 seconds. Allow the liquids to separate and note the color of each layer.
    3. Dispose of the mixture and clean the test tube.

    C.  Miscibility of liquids

    1. Take 3 dry test tubes and add liquids as follows:
    2. 1mL mineral oil and 1 mL isopropyl alcohol to test tube #1.
    3. 1mL mineral oil and 1 mL water to test tube #2.
    4. 1 mL water and 1 mL isopropyl alcohol to test tube #3.
    5. Stopper each test tube and shake for about 5 seconds. Note which pair is miscible.Dispose of waste in the designated container.

    D. Effect of particle size on rate of dissolving

    1. Fill a dry test tube to about 0.5cm deep with fine crystals of sodium chloride. Fill another test tube to about the same depth (0.5cm) with coarse sodium chloride crystals.
    2. Add 10mL of tap water to each test tube and shake each tube. Note the number of seconds required to dissolve the salt in each test tube.

    E. Effect of temperature on rate of dissolving

    1. Weigh two 0.5g samples of fine sodium chloride crystals.
    2. Add 50mL of cold tap water to a 100mL beaker.  Add 50mL of boiling water to a 250mL beaker.
    3. Add the 0.5g of sodium chloride to each beaker and slowly tilt the solutions in each beaker.  Observe how fast the sodium chloride dissolves in each solution. (Do not stir the solutions)

    F.  Solubility versus temperature; saturated and unsaturated solutions

    1. Label 4 weighing boats or papers as follows and weigh the stated amounts onto each one:

    a- 1.0g NaCl              b- 1.4g NaCl               c- 1.0g Sucrose            d- 1.4g Sucrose

    1. Place each 1.0g sample into a separate labeled test tube. Add 5mL of tap water to each, stopper, and shake until the solid in each test tube dissolves.
    2. Add 1.4g NaCl to the NaCl solution of part F.2. Add 1.4g Sucrose to the Sucrose solution of part F.2. Stopper and shake both test tubes. Note whether the crystals have dissolved in each.
    3. Place both unstoppered test tubes into a 500 mL Beaker with boiling water and stir frequently and vigorously, being careful not to break the bottom of the test tubes. Note results after 5 minutes. (You will need to continually add boiling water to the beaker so that the test tubes are consistently heated.)
    4. Cool the test tubes in running water for about a minute and let them stand for a few minutes. Record your observations.
    Section _______ Date _____________

    Report for Experiment 11          Instructor_______________________

    1. Concentration of a saturated solution
    2. a- Mass of evaporating dish __________________

    b- Mass of evaporating dish and saturated NaCl solution      __________________

    c- Mass of dish and NaCl after evaporation                           __________________

    1. Calculate: (Show work. For more space, use the reverse side.)

    a- Mass of saturated NaCl solution                                         __________________

    b- Mass of NaCl dissolved in the solution                              __________________

    c- Mass of water in the solution                                              __________________

    d- Mass % of NaCl in the solution                                         __________________

    e- Grams of NaCl per 100g of water                                      __________________


    1. Relative solubility of a solute in two different solvents


    1. Which liquid is more dense, water or cooking oil? __________________
    2. What evidence supports your answer? ____________________________________
    3. What is the color of iodine in water? __________________
    4. What is the color of iodine in cooking oil? __________________
    5. In which solvent was iodine more soluble? __________________
    6. What is the experimental evidence? _____________________________________



    1. Miscibility of liquids
    1. Which liquid pair(s) tested is miscible? __________________
    2. How do you classify the water/mineral oil mixture? __________________


    1. Rate of dissolving versus particle size
    2. Time (seconds) required for fine salt to dissolve __________________
    3. Time (seconds) required for coarse salt to dissolve             __________________
    4. What general conclusion can you draw?


    1. Rate of dissolving versus temperature
    2. Did salt dissolve faster under hot or cold conditions? __________________
    3. Length of time (seconds) for salt to dissolve in hot water __________________


    1. Solubility Versus Temperature
    2. Which of the 1.0g solutions is saturated? __________________
    3. Evidence? _________________________________________________________
    4. Which of the 2.4g solutions is saturated? __________________
    5. Evidence?_________________________________________________________
    6. Which solid is least soluble at the elevated temperature? __________________

    LAB EXERCISE 2                               Name____________________________

    Gas Laws

    This week we covered gas laws in chapter 6.  Please review this material and complete the following practice problems.  Show all calculation setups, including units, for all problems.

    1. A sample of methane gas, CH4, occupies 3.25 L at temperature of 19.0 oC.  If the pressure is held constant, what will be the temperature be if the volume expands to 10.00 L?
    2. A sample of oxygen gas occupies 1.9 L at pressure of 1156 torr.  What volume will it occupy when the pressure is changed to 912 torr and temperature remains constant?
    3. The pressure of hydrogen gas in a constant volume cylinder is 5.01 atm at 21.0 oC.  What will be the pressure if the temperature is raised to 70.0 oC?
    4. A 988 mL sample of air is at 852 mm Hg and 34.1 oC.  What will the temperature of this gas be, in Fahrenheit, at 955 mm Hg and a volume of 602 mL?
    5. A sample of a gas occupies 9850 mL at STP.  What volume will the gas occupy at 95 oC and 675 torr?

    LAB EXERCISE 2                        Name_____________________________

    1. A sample of nitrogen gas occupies 28.5 L at STP.  How many moles of nitrogen are present?
    2. A 795.0 mL volume of hydrogen gas is collected at 23 oC and 1055 torr.  What volume will it occupy at STP?
    3. What would the pressure be of 25.0 g of chlorine gas at –10.0 ºC in a 4.50 L container?
    4. Calculate the density of CH4 at STP.
    5. A volume of 495 mL of argon gas was collected at 21.0 oC and 779 torr.  What does this sample weigh?
  • Antiseptics and Oxidation



    In General, Organic, & Biological Chemistry read the Chapter 5.4 HealthLink titled, Antiseptics and Oxidation.  Research an antiseptic of your choice.

    • What antiseptic did you choose?
    • Is it an oxidizing agent?  How can you determine this?
    • What is the main purpose of your antiseptic?
    • What functional groups are present in the antiseptic?
    • Define the terms germicide and bacteriostat?  Classify your antiseptic in the appropriate category.

    Post your answers below and respond to at least two classmates’ posts with substantive responses.

  • EXPERIMENT 2 VSEPR- Structure and Shape


    Valence Shell Electron Pair Repulsion Theory Lab VSEPR

    Student Kit Materials:

    Lab Kit:


    Item #


    Molymod® Individual Organic Pack




    Student Responsibility:



     VSEPR- Structure and Shape 

    Materials and Equipment: Molecular Model Kit.


    This exercise provides procedures to determine the structure and shape of molecules.  This information is important because the properties of molecules are dependent upon their structure.  The first step in determining the structure (Lewis structure) of a molecule is to draw a structure accurately showing the location of all valence electrons.  From the Lewis structure, you can use a method called valence shell electron pair repulsion (VSEPR) theory to predict the shape of a molecule or ion.

    In order to use VSEPR, you need to be able to determine the number of electron groups bonded to the central atom and the number of atoms bonded to the central atom.  Lastly, after you have determined the molecule’s shape, you can determine whether electron density in the molecule is arranged symmetrically (a nonpolar molecule) or asymmetrically (a polar molecule).  In a polar molecule, one end of the molecule has a partial positive charge, one end has a partial negative charge. The polarity of a molecule has important implications for the properties of molecules.


    Theory of VSEPR

    In order to use VSEPR, it is necessary to have a completed Lewis structure for the molecule.  VSEPR is based on the principle that electron groups in a molecule tend to stay as far apart from each other as possible due to the repulsive forces that exist between like charges (the electrons).  An electron group could be a lone pair of electrons, a single bond, a double bond or a triple bond around the central atom.  The most probable arrangement of two, three, or four electron groups around a central atom are given in the table below.  This arrangement allows groups to spread out as far as possible.

    Table 1.  Electron Group Geometries


    # of Electron Groups Electron Group Geometry
    2 Linear
    3 trigonal planar
    4 Tetrahedral


    As an example, let’s consider methane, CH4.  The Lewis structure for methane is given below:





    In this case we can see that there are four electron groups (4 single bonds) surrounding the carbon atom, hence the geometric arrangement of the electrons about the carbon atom is tetrahedral.

    Ammonia, NH3, is a little more difficult.





    The Lewis structure for ammonia shows that there are four electron groups (3 single bonds and 1 lone pair of electrons) therefore the electron group geometry is also tetrahedral.  It should be noted however that CH4 has 4 atoms bonded to the central atom, while NH3 only has 3 atoms bonded to the central atom.  Ammonia therefore they will not have the same shape as CH4.  Molecular shape describes the arrangement of atoms about the central atom.  When determining the molecular shape, you must consider the electron group geometry and the number of atoms bonded to the central atom.  (Lone pairs are ignored at this point.)  The possible combinations of electron groups and bonded atoms are summarized below.


    Table 2.  Electron Group Geometries and Molecular Shapes


    # of Electron Groups # of Bonded Atoms Electron Group Geometry Molecular Shape
    2 2 linear linear
    3 2 trigonal planar bent  (120°)
    3 3 trigonal planar trigonal planar
    4 2 tetrahedral bent   (109.5°)
    4 3 tetrahedral trigonal pyramidal
    4 4 tetrahedral tetrahedral


    Using Table 2, we can predict that CH4 has a tetrahedral molecular shape while NH3 has a trigonal pyramidal molecular shape.


    After the geometries have been assigned to a molecule, we decide if there is more than one correct structure for it.  These correct structures are called resonance structures.  Lastly, we can use the molecular shape to determine if electron density is evenly distributed across the molecule.  If electron density is unevenly distributed across the molecule, the molecule is said to be polar.  A molecule with a uniform charge distribution is nonpolar.  But first you must learn how to draw Lewis dot structures…



    1. Drawing Lewis structures. This procedure will be illustrated using SO2 as an example.
    2. Determine the total number of valence electrons in the molecule. The number of valence electrons from an atom can be calculated by its location in the periodic table.  So, in this case, S and O are both in group VIA, so each atom contributes 6 electrons.  Hence the total number of valence electrons in SO2 is 18 (3 atoms ´ 6 valence electrons).

    For ions, it is necessary to add or subtract electrons depending on the charge of the ion.  For anions, the magnitude of the charge should be added as additional valence electrons.  For example, for OH, the total number of valence electrons is eight:  six from oxygen, one from hydrogen, and 1 for the negative charge  (6 + 1 + 1 = 8).  For cations, the magnitude of the charge should be subtracted from the number of valence electrons.  For NO+, the total number of electrons is 10 (5 for nitrogen plus 6 for oxygen, and subtract one for the charge).

    1. Determine which atom is the central atom and place a pair of electrons between it and the other atoms. Generally, look for the atom that there is only one of in the formula.  In SO2, there is only 1 sulfur atom (and 2 oxygen atoms) therefore sulfur is the central atom. Knowing this, we can construct the following crude sketch:



    1. Subtract the number of electrons used to connect the atoms from the total number of valence electrons. Remember each single bond is composed of 2 electrons.  For SO2, 18 electrons (total) – 4 electrons (from the two bonds in step 1) = 14 electrons left over.  Therefore we have 14 electrons left to place around the molecule.
    2. Add the appropriate number of electrons around each atom. Hydrogen requires 2 electrons, boron requires 6 electrons, and all other elements require 8 electrons.  Start by placing electrons on the outer atoms to give them a compete octet.  If more electrons are available, place them on the central atom.  If the central atom lacks an octet, form multiple bonds with outer atoms (see below).

    In our example, S and O both require 8 electrons.  So first we put 6 electrons around one oxygen (which gives it 8 including the two in the bond) and another 6 electrons around the other oxygen.  At this point our structure will look like this:



    Now we have only two electrons left to place around the S atom.  The question is, do we have enough electrons?  If we place the two electrons around the S atom, sulfur will have only 6 electrons (as shown in the following structure), and we need 8.



    We need to use the electrons more efficiently by making one of the lone pairs on an O atom a double bond.  If we move a lone pair to make a double bond, we get the following structure.



    This is the completed Lewis structure for SO2 because all of the atoms are surrounded by eight electrons (octet rule!).  Remember an octet for hydrogen (H) is only two.

    1. Procedure to Determine the Electron Group and Molecular Shapes of a Molecule. The electron group geometry can be determined by counting the number of groups of electrons (atoms + lone pairs) around the central atom and then looking up the appropriate geometry in Table 1.  In the case of SO2, we count three groups from the Lewis structure (2 atoms + 1 lone pair).  From Table 1, the electron group geometry is trigonal planar.

    The molecular shape can be determined by counting the number of atoms bonded to the central atom, and using the number of electron groups determined above to select the appropriate geometry from Table 2.  SO2 has two bonded atoms and three electron pairs, so Table 2 indicates that the molecular shape is 120° bent.


    1. Resonance Structures. Some molecules have more than one correct Lewis structure.  These are called resonance structures.  In order for a molecule to have resonance structures, it must have at least one multiple bond.  Molecules with only single bonds cannot have resonance structures.

    In the case of SO2, the molecule has been drawn above with the double bond to the oxygen to the right of the sulfur.  However, it could have also been drawn between the sulfur and the oxygen on the left, as shown below.  These are the resonance structures for SO2.




    1. Polarity of a Molecule. The last piece of information to be obtained about a molecule concerns the distribution of electron density and charges around the molecule.  A molecule with a uniform distribution of electron density is nonpolar; and one with an asymmetrical distribution is polar.  A molecule is nonpolar only if it has no lone pair electrons about the central atom and all groups attached to the central atom are identical (both conditions must be met to be nonpolar).  Another way to state this is if the electron group and molecular shapes are the same and the atoms attached to the central atom are identical, then the molecule is nonpolar.

    In the case of SO2, the Lewis structure shows us that the molecule is polar because the sulfur atom has a lone pair.




    Determine the Lewis structure, electron group geometry, molecular shape, presence or absence of resonance structures, and the polarity for a series of molecules given on the work­sheet.  Once you have filled in the worksheet, build a model of each compound using your model kit.  Compare the model you built with the responses you provided in the lab.  Carbon tetrachloride is worked out for you as an example.




    Names:                                                                      Date:                                                             


    CCl4 BF3 SO3 CO2 ClO2
    Crude Sketch 




    Calculations (# of valence electrons, # of bonds, etc. 





    1(4) + 4(7)=32


    32-4(2) = 24

    Lewis Structure 







    # electron groups,electron group geometry


    # of bonded atoms, molecular shape 


    Resonance structures (if any) 







    Polar or nonpolar 






    H2O SO42- NO2+ PO43- NO3
    Crude Sketch 




    Calculations (# of valence electrons, # of bonds, etc. 





    Lewis Structure 







    # electron groups,electron group geometry


    # of bonded atoms, molecular shape 


    Resonance structures (if any) 







    Polar or nonpolar 






    CO32- SO2 NO2 PF3 SiI4
    Crude Sketch 




    Calculations (# of valence electrons, # of bonds, etc. 





    Lewis Structure 







    # electron groups,electron group geometry


    # of bonded atoms, molecular shape 


    Resonance structures (if any) 







    Polar or nonpolar 





    NH3 H3O+ NH4+ SO32- CHCl3
    Crude Sketch 




    Calculations (# of valence electrons, # of bonds, etc. 





    Lewis Structure 







    # electron groups,electron group geometry


    # of bonded atoms, molecular shape 


    Resonance structures (if any) 







    Polar or nonpolar 







    LAB EXERCISE 1                             Name____________________________


    Formulas and Names


    1. Complete the table below by providing the formula and name of the compound formed in the rectangle where the cation (positive ion) and anion (negative ion) intersect.


    Cl                                                 O 2-                              P 3-






    Lithium Chloride






    Ca 2+


















    Mn 4+









    LAB EXERCISE 1                                   Name______________________________


    1. Complete the table below by providing the formula and name of the compound formed in the rectangle where the cation (positive ion) and anion (negative ion) intersect.


    NO3                                 CO3 2-                                     PO4 3-






    sodium nitrate















    Fe 3+


















    LAB EXERCISE 1 (continued)               Name______________________________


    1. Give the formulas for the following ionic compounds.


    a- manganese (III) chloride          _________________________________


    b- magnesium sulfide                   _________________________________


    c- barium bromide                        _________________________________


    d- potassium phosphide                _________________________________


    e- cadmium acetate                       _________________________________


    f- calcium permanganate              _________________________________


    g- cesium bromate                        _________________________________


    h- beryllium dichromate               _________________________________


    i- ammonium peroxide                 _________________________________


    j- iron (III) chloride                      _________________________________


    k- uranium (V) sulfide                  _________________________________


    l- nickel oxide                               _________________________________


    m- tin (IV) nitride                         _________________________________


    n- bismuth (III) hypochlorite        _________________________________


    o- mercury (I) thiocyanate            _________________________________


    LAB EXERCISE 1 (continued)               Name______________________________

    1. Name the following ionic compounds.


    a- BaI2                               _________________________________


    b- K3PO4                           _________________________________


    c- AlI3                               _________________________________


    d- Ag3N                            _________________________________


    e- Li2Cr2O7                       _________________________________


    f- MnCO3                          _________________________________


    g- CsHCO3                       _________________________________


    h- Sr(BrO2)2                      _________________________________


    i- Li2CrO4                         _________________________________


    j- CrF6                               _________________________________


    k- TiO2                              _________________________________


    l- Co3P2                             _________________________________


    m- MnS2                           _________________________________


    n- Zn(OH)2                       _________________________________


    o- Ga(ClO)3                      _________________________________


    LAB EXERCISE 1 (continued)               Name______________________________

    1. Name the following molecular compounds.


    a- NF3                               _________________________________


    b- P2S5                              _________________________________


    c- SBr2                              _________________________________


    d- CO                                _________________________________


    e- P4S3                               _________________________________


    f- SiI4                                   _________________________________


    g- SCl6                              _________________________________


    h- OCl4                              _________________________________


    i- SeO2                              _________________________________


    j- N2O3                              _________________________________


    k- B2H6                             _________________________________


    l- AsI3                                 _________________________________


    LAB EXERCISE 1 (continued)         Name______________________________

    1. Give the formulas for the following molecular compounds.


    a- nitrogen disulfide                     _________________________________


    b- phosphorus pentabromide        _________________________________


    c- carbon tetrachloride                  _________________________________


    d- sulfur hexabromide                   _________________________________


    e- sulfur trioxide                           _________________________________


    f- silicon disulfide                        _________________________________


    g- diarsenic pentasulfide               _________________________________


    h- boron trichloride                       _________________________________


    i- dinitrogen monoxide                 _________________________________


    j- carbon tetrahydride                   _________________________________


    LAB EXERCISE 1 (continued)         Name______________________________

    1. For each acid, name the anion bonded to hydrogen, then name the acid.


    Anion name                                                     Acid name


    a- HI                      _______________________                          _______________________


    b- HCl                   _______________________                          _______________________


    c- H2S                   _______________________                          _______________________


    d- H3PO4               _______________________                          _______________________


    e- H3PO3               _______________________                          _______________________


    f- H2SO4               _______________________                          _______________________


    g- H2SO3               _______________________                          _______________________


    h- HClO4               _______________________                          _______________________


    i- HClO                 _______________________                          _______________________












    LAB EXERCISE 1 (continued)          Name______________________________


    1. For each acid, give the formula of the anion bonded to hydrogen, then give the formula of the acid.


    Anion formula                                     Acid formula


    a- hydrobromic acid    ________________________            ________________________


    b- hydrofluoric acid    ________________________            ________________________


    c- hydrocyanic acid    ________________________            ________________________


    d- nitric acid                ________________________            ________________________


    e- sulfuric acid                        ________________________            ________________________


    f- periodic acid            ________________________            ________________________


    g- acetic acid               ________________________            ________________________


    h- sulfurous acid         ________________________            ________________________

    i- hypobromous acid   ________________________            ________________________


  • EXPERIMENT XX Solutions




    Materials and Equipment: Digital balance (0.01g precision); thermometer; 150mL and 250mL beakers; Teaspoon; 1 Cup Measuring Cup; Coffee Mug; 10mL and 100mL graduated cylinders; watch glass; marble.


    In this experiment we will become familiar with several instruments. Every instrument has some level of uncertainty, in other words, no measuring devise is perfect or exact. The smaller the increments of the measuring device, that is, the smaller the spaces between the actual markings (lines, tick marks) on the device, the closer the observer is able to determine the quantity being measured to its true value. The procedure used to report non-digital instrument readings is as follows: 1. Determine the size of the increments (spaces) on the measuring device (For example, 0.1cm, or 1mL). 2. Record your reading as a numerical value including one decimal place further to the right of the incremental value (For example, if the increment is 0.1cm, a one-tenth increment, you must record your reading as a numerical value that includes a digit or zero in the hundredths place). Examples of how to read non-digital volumetric, temperature, and length measuring instruments are provided below. For a digital measurement, record all the digits and zeros displayed.    The number of significant figures in your measurement is based off of the accuracy of the instrument.  The more significant figures, the more accurate the measurement is.

    Mass (weight) measurements in this lab are recorded to the 0.01g (hundredth of a gram) or 0.001g (thousandth of a gram), depending on the precision of the digital balance used.

    Volumetric measurements in this lab are made in graduated cylinders.  Most of liquids we use are water (aqueous) solutions. The curved surface (curve “pointing” downward) that you observe when reading the volume contained in a graduated cylinder is called the meniscus.  Observe the meniscus at eye level to read the volume correctly.  In the 10mL graduated cylinder, the increments (spaces) are 0.1mL. In the 100mL graduated cylinder, the increments are 1mL. An example reading technique is shown in Figure 2-1.

    Figure 2-1

    Temperature measurements in this lab are recorded using a thermometer with increments of 1oC. Example temperature readings are shown in Figure 2-2.

    Figure 2-2

    Length measurements in this lab are recorded using printed metric rulers (below) with different size increments. On Ruler A, (Figure 2-3) the increments are 1cm.  On Ruler B, (Figure 2-3) the increments are 0.1cm.  On Ruler A, an example length might be estimated as 12.5cm. However, on Ruler B, the same example length might be estimated as 12.55cm because Ruler B has smaller increments.

    Figure 2-3



    One distinguishing physical property of matter is density. Each pure substance has its own density. Therefore, density can be used to help identify the substance. In words, density is the ratio of a substance’s mass to the volume occupied by that mass. Mathematically, with density, d, mass, m, and volume, V,


    d = _m_




    In this lab, you will measure mass in grams, g, and volume in milliliters, mL. Therefore, the units of density will be, in words, grams per milliliter, or in symbols, g/mL or _g_ .


    Liquid volumes are easily measured by means of graduated cylinders. However, solids, unless precisely machined, for example, have volumes that are difficult to determine by direct measurement with a ruler or other similar instrument. The volume of irregularly shaped solids that do not dissolve (for example, rocks, pieces of metal) may be measured by the volume of a liquid that the solid displaces when the solid is completely submerged. If a graduated cylinder is partially filled with water and an insoluble solid is carefully submerged in the water, the water level will rise. The volume of the solid is the difference between the water levels before and after the solid is added.


    In this lab, you will determine the density of water and the density of a metal slug.




    Part A: Measuring Length

    1. Measure the diameter of a watch glass using Ruler A (Figure 2-3).
    2. Record the measurement on your report.
    3. Measure the diameter of the same watch glass using
    4. Ruler B (Figure 2-3). Record the measurement on your report.






    Part B: Measuring Mass


    Digital balances are sensitive. Proceed carefully. Be sure the balance is energized. The display should read “0.00 g.” If it does not, press the “TARE” key (the one with T/0) and wait a few seconds.  Place the item to be weighed on the balance pan and read the digital display.  The digital display represents the mass of the item in grams.  Remember that all the decimal places shown on the display are to be recorded.


    1. Use a digital balance to determine the mass of a 100mL beaker.
    2. Record the mass.
    3. Using a 100mL graduated cylinder, pour about 20mL of tap water into the same beaker. (See NOTE at the end of this paragraph.)
    4. Determine the combined mass of the beaker and water.
    5. Record the combined mass.
    6. Calculate the mass of the water. Record the result.


    **(NOTE: When you read an instruction about a volume that states “about” or “approximately,” there is no need to go through the careful procedure described in the introduction to this lab. Unless your instructor tells you otherwise, use the “plus or minus 10%” rule. For example, for “about 20mL of tap water,” obtain between 18mL and 22mL of tap water.)

    Part C: Measuring Volume

    1. Pour 1 Teaspoon of water in the 10 mL graduated cylinder.
    2. Record the volume of the water in the 10 mL cylinder.
    3. Pour 1/3 cup of water in the 100 mL graduated cylinder.
    4. Record the volume of the water in the 100 mL cylinder.


    Part D: Measuring Temperature

    1. Fill a 250mL beaker about halfway with tap water.
    2. Use a thermometer to record the temperature.
    3. Using a microwave proof glass (such as a coffee mug), heat 1 cup of water to boiling.
    4. Carefully remove the water from the microwave and use a thermometer to record the temperature of the water.


    Part E: Measuring the Density of a Liquid

    1. Dry a 10mL graduated cylinder, weigh it, and record its mass.
    2. Put about 9mL of deionized water into the cylinder and carefully record the volume of the water. (Refer to the introduction.)
    3. Weigh the cylinder with the water and record the combined mass.
    4. Calculate the mass of the water.
    5. Calculate the density of the water.

    Part F: Measuring the Density of a Solid

    1. Weigh the marble and record its mass.
    2. Fill a 100mL graduated cylinder about halfway with tap water.
    3. Read and record this initial volume of water following the rules.
    4. Tilting the graduated cylinder at an angle without spilling the water, carefully slide the marble down the cylinder without splashing water up along the sides.
    5. Read and record this final volume of water.
    6. Calculate the volume of the marble and its density and record both values on your report.




    Report for Experiment 2                  Instructor_______________________




    Part A: Measuring Length


    Diameter of watch glass using Ruler A                      __________________________


    Diameter of watch glass using Ruler B                      __________________________



    Part B: Measuring Mass


    Mass of 100mL beaker                                               __________________________


    Mass of 100mL beaker with water                             __________________________


    Mass of water                                                             __________________________

    (show calculation)

    Part C: Measuring Volume


    Volume of water in 10mL cylinder                            __________________________


    Volume of water in 100mL cylinder                          __________________________




    Part D: Measuring Temperature


    Temperature of cold water                                         __________________________


    Temperature of boiling water                                     __________________________


    Part E: Measuring the Density of a Liquid


    Mass of cylinder                                                        __________________________


    Mass of cylinder with water                                       __________________________


    Mass of water                                                             __________________________

    (show calculation)


    Volume of water                                                         __________________________


    Density of water                                                         __________________________

    (show calculation)



    Part F: Measuring the Density of a Solid


    Mass of marble                                                           _________________________


    Initial volume of water                                               __________________________


    Final volume of water                                                 __________________________


    Volume of marble                                                       __________________________

    (show calculation)


    Density of marble                                                       __________________________

    (show calculation)



    1. Determine the number of significant figures in each of the following measurements:


    1. 35s ___________ 1.05cm __________


    1. 89g ___________ f. 10.5mL __________


    1. 012g ___________ g. 10.00mL __________


    1. 0cm ___________ h.  -40.0oC __________


    1. Perform the following calculations; apply sig. fig. and round off rules:


    1. 6g + 50.05g + 50.432g  =


    1. 77mL – 23.4mL  =


    1. (41.5cm)(0.4cm)  =


    1. 8 mm3 / 25.4 mm  =


    1. Convert 37.6oC to Fahrenheit. (Remember: In the formula for converting between Fahrenheit and Celsius, the “32” and “1.8” are exact numbers.)
    2. Convert 112oF to Celsius.
    3. What is the mass of a liquid that has a volume of 50.0mL and a density of 2.40g/mL?
    4. Calculate the volume, in milliliters, of a solid that has a density of 7.142g/cm3 and a mass of 232.51g. (HINT: 1cm3 = 1mL exactly)
  • Genes of Genesis Solution


    Week 1 | Discussion – The Origin of Life

    Respond And Discuss

    One of the theories of how life originated is that organic molecules came together in the ocean and life was created.  There is a fascinating article regarding this on the website in a new window).  Please download Genes of Genesis [DOWNLOAD] and read from “Molecules that Begin Life” through “Enzymes and Nucleic Acid Formation” and respond to the following prompts:

    1. Explore this topic in more detail and summarize how specifically a living molecule was supposed to have formed from an organic substance.
    2. According to the Bible, life started when God spoke it into existence. What flaws do you see with the theory that life came from organic compounds?
    3. The article cites several different types of organic molecules that were to have formed in early earth. Pick one molecule and describe how scientists theorize it formed. What are the inaccuracies with their method?
    4. With respect to the article, what evidence is there to support creationism? Find an outside source other than Wikipedia that will support this evidence.

    Post your answers below and respond to at least two classmates’ posts with substantive responses.

  • Cabbage Juice pH Indicator Solutions


    Acids and Bases: Cabbage Juice pH Indicator


    Liquids all around us have either acidic or basic (alkaline) properties. For example, acids taste sour; while, bases taste bitter and feel slippery. However, both strong acids and strong bases can be very dangerous and burn your skin, so it is important to be very careful when using such chemicals. In order to measure how acidic or basic a liquid is one must use the pH scale as illustrated below:…/ ph-scale.gif

    In this lab, students will use the juice from red cabbage as a pH indicator to test common household liquids and determine their pH levels. You will mix cabbage juice with different household liquids and see a color change produced by a pigment called flavin (an anthocyanin) in red cabbage. Through this color change, you will be able to successfully identify the approximate pH of common household liquids using the table below:

    Strength increases at extremes of this scale.

    Cutting the cabbage:

    1. Cut ¼ of a cabbage into thin strips like shown below:




    1. Place the cut cabbage into a Ziploc bag. Note: Keep the red cabbage refrigerated.



    • Pre-cut cabbage
    • Blender
    • Strainer
    • Large container
    • 1liter beaker
    • 7 plastic cups
    • 7 plastic spoons
    • Lemon soda
    • White vinegar
    • Apple juice
    • Baking soda
    • Shampoo (preferably clear)
    • Conditioner (preferably clear)
    • Hand sanitizer

    Pre-Lab Questions

    Look at each of the liquids being tested. Predict whether each of the substances is acidic, neutral or basic. Circle one. (Think about the properties of acids and bases.)

    Part 1: Preparing the Cabbage Juice

    Preparing the Cabbage Juice:

    1. Put the red cabbage leaves into the blender with 800mL of water.
    2. Close the top and let it blend at high power for 30 seconds.
    3. Once it is blended, filter out the leaves inside the mixture with the strainer and pour the mixture into a large container.

    *This should provide you with 600-800 ml of cabbage juice.

    Part 2: Mixing the Cabbage Juice

    1. Label each cup with each of the liquids. (Example: vinegar, apple juice, etc.)
    2. Pour 100 ml of each individual liquid into its respective cup (except for baking soda).
    3. For baking soda, add 3 tablespoons of baking soda into 100 ml water.


    1. Pour 50 ml of cabbage juice into each of the cups. Do this one at a time and record the color change below:

    Now look up the actual pH of each of the substances and see how accurate the cabbage juice indicator was!

    How did your reasoning for your predictions change after seeing the approximate pH level?

    Does the color intensity of the liquid change? If so, how and why do you think this is?

    Concept Questions:

    1. Does the addition of water (baking soda + water) alter the pH of weak acids/bases? How does it change the pH of strong acids/bases? Why or why not?
    2. How does a difference in 1 pH unit change in terms of H+ concentration? Example: How does a pH of 3 differ from pH of 4? Which one is stronger or weaker? Why?
    3. Look at the ingredients for each liquid you tested. Which ingredients contribute to each of the liquid’s pH level?

    Real Life Applications:

    1. Neutralization: Whenever you mix an acid with a base, they neutralize each other. If this is the case, why is Alka-Seltzer used to treat stomach aches? (Note: excess stomach acids cause stomach aches)
    2. What is acid rain and how does it negatively impact oceans, rivers, lakes, and other natural environments?
  • Nursing and the budget


    Assignment 1: Discussion Assignment

    All health care organizations have complicated budgeting policies and procedures. The more the nurse understands the process, the more effectively they can participate in the process. The budget process usually starts with an interdisciplinary approach.

    Describe the potential members of the interdisciplinary team for budget development and the role of each individual.
    What are the specific responsibilities of nursing in the development of the budget?
    Are the responsibilities for budget specific only to the leadership of the nursing department or are they found throughout the organization?

  • Most health problems


    Assignment 1: Discussion Assignment

    The discussion assignment provides a forum for discussing relevant topics for this week based on the course competencies covered.

    Health Problems
    Which area, rural or urban, has the most health problems? Why? Justify your response based on the readings or articles from the South University Online Library.

    To support your work, use your course and text readings and also use the South University Online Library. As in all assignments, cite your sources in your work and provide references for the citations in APA format.
    Start reviewing and responding to the postings of your classmates as early in the week as possible. Respond to at least two of your classmates. Participate in the discussion by asking a question, providing a statement of clarification, providing a point of view with a rationale, challenging an aspect of the discussion, or indicating a relationship between two or more lines of reasoning in the discussion. Cite sources in your responses to other classmates.

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