Chemistry 105 ~ Skills Summary for Laboratory Final · PDF fileChemistry 105 ~ Skills Summary for Laboratory Final Exam ... to the appropriate number of significant figures/decimal

C. Graham Brittain Page 1 of 12 11/18/2008

Chemistry 105 ~ Skills Summary for Laboratory Final Exam

The Final Exam questions will have a variety of formats, including short answer, multiple-choice, fill-in-the-blank, and

essay/problem-solving.

Each student must work individually, and come prepared with a calculator. A Periodic Table and key mathematical formulas

(% Error, % Difference, % Yield, etc.) will be provided.

Good study materials include:

• Your Pre-Lab Assignments and Pre-Lab Quizzes (though you should realize that these questions were intended to

test for a BASIC level of understanding BEFORE conducting the experiment; this means you should know a bit more

now that you’ve completed the experiment)

• The “Concepts you need to know to be prepared” discussions in the lab handouts

• The post-lab calculations on the Report Sheets

• The problem-solving and discussion questions on the Report Sheets

For ALL calculations performed on the lab final exam, realize that your lab instructor can’t possibly grade work done in your

head or on your calculator, ONLY the work that you present neatly and legibly on your test paper!

So to receive maximum credit on calculations, you MUST show them neatly, with correct units, and express your answers,

with correct units, to the appropriate number of significant figures/decimal places.

If you use one of the formulas provided on the exam (such as PV = nRT), you should show the formula FIRST, and THEN

substitute the values with their units.

Realize that it’s quite likely that you’ll be given sample experimental data, and asked to work the same sorts of calculations

that you did in completing your Report Sheets.

For example: Consider the Mole Relationship (“Tracking Zn+2”

) experiment. You might be given the balanced chemical

equation and an initial mass of ZnCO3, and asked to predict the mass of ZnO (or CO2) that will be produced by the reaction (a

stoichiometry calculation). If you’re given the mass of product that was experimentally produced, you could then be asked to

calculate the % Yield.

If you feel that you’ll need assistance in reviewing any of the skills listed below, please make arrangements to meet with your

lab instructor (or other chemistry teaching assistant) in the Pastore 215 Help Office.

Realize that with the exception of a few skills specific to the laboratory, these are the SAME skills needed for success in your

CHM 103 lecture course.

To demonstrate mastery of CHM 105 Skills, students should be able to:

From the Laboratory Safety Activit:y:

Interpret the chemical hazard information (color and number) indicated on National Fire Protection Association (NFPA)

labels.

From Lab 1: Laboratory Measurements

Given a measurement, recognize the number of “significant figures” it contains (and the digit that contains the uncertainty).

When making a measurement, identify the smallest increment on the measuring device, and then estimate the digit that

contains the uncertainty. Record the measurement with units and the appropriate number of significant figures. Use rulers,

graduated glassware and thermometers to make measurements.

Use significant figures correctly when working ALL calculations (multiplying /dividing or adding/subtracting), and round

appropriately to express the uncertainty in the result.

Convert numbers between standard and scientific (exponential) notation.

Use Dimensional Analysis (aka the Factor Unit Method or Factor Label Method) to solve “unit conversion” problems:

• Use “unit relationship” EQUALITIES as “conversion factors.”

• Set up calculations so that units CANCEL correctly and the result has the desired units.


Explain the difference between accuracy and precision:

• Accuracy is the degree of agreement between a measured value and the true value. A measurement is said to be

“accurate” if it is very close to the true value, and “inaccurate” if it is not close to the true value.

• Precision refers to the agreement of replicate measurements of the same quantity.

Given an “accepted” value and an “experimental” value, calculate the Percent Error.

State the base metric unit for mass, length, volume, and temperature:

• The base metric unit of mass is the gram (g)

• The unit of length is the meter (m)

• The unit of volume is the liter (L)

• The unit of temperature is the Celsius degree (oC)

Define the most commonly used metric prefixes (mega-, kilo-, deci-, centi-, milli-, and micro-). These are related to the base

unit by powers of ten.

Use the metric prefixes to convert between larger to smaller units (so as to use a unit of the appropriate size for what is being

measured).

Explain how volume is a unit that is derived

from length.

State equivalencies between the more commonly

used cubic length and laboratory volume units.


Explain the relationships between the Celsius and Fahrenheit temperature scales.

Given the appropriate formulas, convert a temperature from one temperature

scale to another.

℉ = �1.8 ℉1 ℃ ℃ + 32

℃ = �℉ − 32� � 1 ℃1.8 ℉

Distinguish between extensive and intensive properties:

• Extensive properties depend on the quantity of a substance. Mass and

volume are examples of extensive properties.

• Intensive properties are those that are inherent to (and characteristic of) a

particular substance. Thus they are independent of the amount of substance

that is being measured, and they can be used to help identify an unknown

substance. Density, freezing/melting point, and boiling point are examples

of intensive properties.

Explain density: the mass per unit volume of a material.

Given any two of the three quantities (mass, volume, density), calculate the third.

Use the density unit relationship between mass and volume as a conversion factor in working dimensional analysis problems.

From Lab 2: Densities of Solids and Liquids

Given mass and volume measurements of different samples of the same substance, prepare a plot of the data from which the

density of the substance can be determined.

From a plot of mass versus volume, graphically determine the density of a substance. (Draw the best fit straight line, and

determine the slope of that line.)

From Lab 3: Separation of a Mixture into Pure Substances ~ Paper Chromatography of Metal Cations

Fully describe the Classification System for Matter:

• Any material is either a pure substance or a mixture of pure substances.

• A pure substance cannot be separated into other kinds of matter by a physical change.

• The two types of pure substances are elements and compounds.

• Elements are the basic building blocks of matter; they cannot be decomposed by a chemical reaction.

• Compounds are made from two or more elements by a chemical reaction; thus they are a chemical combination of

elements. Compounds have constant composition; for a particular compound, the elements always combine in a fixed

proportion.

• Mixtures are a physical combination of pure substances (elements and compounds). They can be separated by a physical

change, by leveraging the differences in the physical properties of the components. They also have variable

composition: elements and/or compounds can be combined in any proportion to make a mixture.

• Heterogeneous mixtures have multiple phases.

• Homogeneous mixtures have just a single phase (they are consistent throughout). Homogeneous mixtures are also called

solutions.


The Classification System for Matter:

Explain how substances can be separated from one another by paper chromatography, on the basis of differences in their

affinities for the mobile and stationary phase.

Given a paper chromatogram, determine the Retention Factor (Rf) value of the separated substances.

Explain what a larger Rf value (moves high up the paper with the mobile phase) indicates about a substance.

Explain what a lower Rf value (leaves the mobile phase and adheres to the stationary phase) indicates about a substance.

From Lab 4: Structure, Geometry, and Polarity of Molecules

Explain the electron sharing (covalent bonding) that occurs between nonmetals in molecular (covalently-bonded)

compounds.

Explain the “covalence” of the nonmetal atoms. (Covalence is the number of electrons needed to achieve a noble gas

valence electron configuration, and is thus the number of covalent bonds an atom will typically form.)

Given the chemical formula of a simple molecule (one to three central atoms), draw a Lewis structure for the molecule.

Explain the information contained in the chemical formula and in the Lewis structure for molecules (covalent compounds):

• The chemical formula expresses the composition of the molecule (exactly how many of each type of atom), but doesn’t

indicate how the atoms are connected.

• Lewis structures indicate the connectivity of the atoms, but do NOT clearly represent the molecular geometry about the

central atoms.


Use electron group repulsion theory to the molecular geometry of the central atoms in a covalently-bonded compound:

• Count the groups of electrons around a central atom and determine the “electron-group geometry.”

• From the number of bonding and nonbonding electron groups, determine the “molecular geometry” of central atoms.

Explain what it means for a

covalent bond to be polar.

State the trends in

electronegativity (Table 4.4).

Use electronegativity values

to evaluate the polarity of

covalent bonds.

Use the molecular geometry

of a central atom to determine

whether bond dipoles are

canceled or enhanced by the

geometry.

From Lab 5: Determining the Chemical Formula of an Ionic Compound

State Avogadro’s number: 1 mole of any “substance” = 6.02 x 1023

“particles” (where “particles” refers to atoms, ions, or

molecules, depending on the nature of the substance).

Write an “Avogadro’s statement” for any element or compound:

• 1 mole of an element contains 6.02 x 1023

atoms of that element, and has a mass in grams equal to the atomic weight of

the element.

• 1 mole of a molecular element (such as O2, H2, etc.) contains 6.02 x 1023

molecules of that element, and has a mass in

grams equal to the formula weight (molecular weight) of the molecule.

• 1 mole of an ionic compound contains 6.02 x 1023

“formula units” of that compound, and has a mass in grams equal to

the formula weight of the compound.

• 1 mole of a covalent (molecular) compound contains 6.02 x 1023

molecules of that compound, and has a mass in grams

equal to the formula weight (molecular weight) of the compound.

NOTE: The mass in grams of 1 mole of ANY substance is commonly referred to as the “Molar Mass.”

Use the Avogadro’s unit relationships (expressed in the “Avogadro’s Statement”) to convert between numbers of “items,”

grams, and moles for any element or compound.

Given sample data (the experimental masses of each of the elements in a compound):

• Use the molar masses of each element to convert from grams to moles.

• Determine the mole ratio of the elements in the compound.

• Determine the chemical formula of the compound.

• Determine the experimental mass percentage of each element in the compound.

Given the chemical formula of a compound, calculate the theoretical mass percentage of each element.


Explain that for an ionic compound, the chemical formula expresses the fixed, relative proportions of the elements in the

compound.

Remember that ionic compounds should NEVER be thought of as just a

PAIR of oppositely-charged ions, as the chemical formula seems to

imply.

If you look closely at the illustration of a sodium chloride (NaCl) salt

crystal to the right, you’ll see that it is definitely NOT just pairs of Na+

and Cl- ions, but an extensive, orderly, three-dimensional network of

charged particles. Each Na+ ion is attracted to the Cl

- ions immediately

above, below, to the right, to the left, in front, and behind it. And each

Cl- ion is attracted to the Na

+ ions surrounding it.

So the chemical formula for any salt does NOT describe a small

grouping of ions, it describes the fixed proportion in which vast

numbers of the oppositely-charged ions must combine in order to

appropriately balance out the positive and negative charges.

From Lab 6: A Six-Bottle Study of Ionic Compounds

Given a chemical equation describing a reaction, BALANCE the chemical equation by placing numerical coefficients in front

of the chemical formulas of the participants in the reaction. Realize that you CANNOT balance an equation by changing the

chemical formulas (by changing the fixed proportion of the elements in a compound).

Given a sentence describing a chemical reaction, translate it into a balanced chemical equation. (Realize that the first step is

to write the correct chemical formula for every element or compound participating in the reaction.)

Given a table of the solubilities of common ionic compounds, use it to identify the solid or gaseous products of a chemical

reaction. Then indicate the physical states of all participants in the balanced chemical equation using the appropriate

subscripts: (g), (l), (s), (aq).


From Lab 7: The Mole Relationship in Chemical Reactions

Use the mole relationships in a balanced chemical equation to work “stoichiometry” problems. That is, for a given amount

(grams, moles) of one participant in a chemical reaction, predict the amounts of other participants in the reaction as dictated

by the stoichiometry of the balanced chemical equation.

Use the percentage yield calculation to compare an experimental (actual) yield to a predicted (theoretical) yield.

From Lab 8: Determining the Ideal Gas Constant

Use the Kinetic Molecular Model to explain the properties of gases, liquids and solids:

• In the solid state: Particle motion is restricted to vibrations. Intermolecular attractive forces are the strongest.

• In the liquid state: Particles move freely around one another, but in close proximity. Intermolecular attractive forces are

weaker.

• In the gas state: Particles move freely and are widely separated from one another. They experience elastic collisions,

and intermolecular attractive forces are at a minimum.

A Kinetic Molecular view of solids, liquids, and gases:


Use the Gas Law relationships to predict the behavior of a fixed number of moles of a gas when it is compressed, expanded,

heated or cooled.

• Boyle’s Law: For a constant number of moles of gas at a constant temperature, the VOLUME of the gas is INVERSELY

proportional to its pressure. This means that as the pressure on the gas is increased at constant temperature, the volume

of the gas must decrease. Or as the pressure on the gas is released, the gas can expand (increase its volume).

• Charles’ Law: For a constant number of moles of gas at a constant pressure, the VOLUME of the gas is DIRECTLY

proportional to its TEMPERATURE. This means that as the gas is heated at constant pressure (its temperature is

increased, the gas will expand (its volume increases).

• For a constant number of moles of gas at a constant volume, the PRESSURE of the gas is DIRECTLY proportional to its

TEMPERATURE. This means that as the gas is heated (its temperature is increased, the pressure of the gas must

increase if volume is held constant.

• Avogadro’s Law: If two samples of gas have exactly the same volume, and exist at exactly the same temperature and

pressure, they must contain the same number of moles of gas.


Use the Ideal Gas Law for calculations regarding the behavior of a sample of gas (n moles) at specified conditions of

pressure, volume, and temperature (P, V, T).

�� = ��

In using the Ideal Gas Law:

P = pressure of the gas in atmospheres (atm)

V = volume of the gas in liters (L)

T = temperature of the gas in Kelvin (K)

The Ideal Gas Constant, R = 0.0821 L atm/mol

From Lab 9: Solubilities of Ionic and Molecular Substances

Explain what a solution is (a homogeneous mixture, with the components present as atoms, ions, or molecules). Give

specific examples of various types of solutions (solid in liquid, gas in liquid, etc.)

Explain the “like dissolves like” rule for solubility. Be able to predict the solubility of given solutes in specified solvents,

based on an assessment of their chemical structure. Consider the thought process required to make this prediction:

• Is the solute ionic or molecular?

• IF the solute is molecular, are the solute molecules primarily polar or nonpolar? (What is the degree of polarity? Mostly

hydrophilic – water loving? Or mostly hydrophobic – water fearing? Are there any hydrogen-bonding groups? How

many? And how large is the nonpolar – “hydrophobic” – portion of the molecule?)

• Are the solvent molecules primarily polar or nonpolar? (Again, what is the degree of polarity?)

• What type(s) of intermolecular attractive forces can the solute and solvent molecules use in interacting with one another?

To the right: How an IONIC compound dissolves in water: the

ions “dissociate” and are “hydrated” by the water molecules.

To the left: How a POLAR molecule dissolves in water: the

molecules “go for a swim” among the water molecules and

interact with them via dipole-dipole interactions or hydrogen

bonding.


From Lab 10: Factors Affecting the Rate of a Chemical Reaction

Explain how Reaction Rate is defined, as well as how it can be measured.

Reaction Rate can be defined as:

• The change in concentration of the REACTANT as a function of time (the rate of DISAPPEARANCE of the REACTANT)

• The change in concentration of the PRODUCT as a function of time (the rate of FORMATION of the PRODUCT)

Concentration is always expressed in MOLARITY units (M = moles/liter).

So a reaction rate typically has units of: moles per liter per unit time: moles/L sec, OR moles/L min, OR moles/L hour, etc.

�� = ∆�� !∆ "#$% = &∆�'%(� () !

∆ "#$% = $�*%+,# %� +%�

Given sample data (change in reactant concentration, change in time), calculate the reaction rate.

Describe the “Collision Theory” that explains reaction rates:

• In order for reactant “particles” (atoms, ions, or molecules) to react, they must collide.

• At the instant of collision, the kinetic energy of particle motion is converted to collision energy.

• The collision energy must be greater than or equal to the Activation Energy (Ea) in order to have an effective (successful)

collision – one that has sufficient energy to break key bonds and rearrange the atoms.

• The Activation Energy (Ea) is the minimum energy that must be supplied by the collision in order for the reaction to

occur by a particular pathway. It is the difference between the energy level of the reactants and the energy level of the

“transition state” (the intermediate that exists at the moment of collision, when the kinetic energy, converted into

collision energy, is used to break bonds and rearrange the atoms).

• A certain fraction of all molecules in a sample will have the necessary Activation Energy to react, and that fraction will

increase with increasing temperature.

• The proper spatial orientation of the colliding molecules is also essential to an effective collision.

Explain (in terms of the Collision Theory) the four factors that influence the rate of a chemical reaction:

• The ionic or molecular nature of the reactants

• The concentration of the reactants

• The temperature of the reactants

• The presence of a catalyst

From Lab 11: An Introduction to Acids, Bases, pH, and Buffers

Explain the Bronsted-Lowry definitions of an acid, a base, and an acid-base reaction:

• An acid is a compound which can donate a proton (hydrogen ion, H+)

• A base is a compound that can accept a proton.

• An acid-base reaction is then simply the transfer of a proton from the acid to the base.

Write chemical reactions which illustrate an acid

donating a proton to produce its conjugate base,

and a base accepting a proton to produce its

conjugate acid.

Note that the brackets [ ] mean “concentration of,”

and that the concentration is ALWAYS expressed in

MOLARITY units (moles per liter = moles/L = M).


Explain what the terms “strong” and “weak” mean when applied to acids (and bases):

• Strong acids dissociate completely in water (ionize 100%).

• Weak acids dissociate to an equilibrium extent, which can be described by the value of the Acid Dissociation Equilibrium

Constant, Ka.

Given the molarity of a strong acid or strong base solution, calculate the hydronium ion concentration, and then the pH of the

solution.

Realize that the dissociation of ANY weak acid (HA) in water can be represented as:

HA (aq) + H2O (l) � H3O+

(aq) + A -

(aq)

And that the equilibrium constant expression for the dissociation of ANY weak acid can be written as:

-( = �./01!�2&! �.2!

As this equilibrium constant describes the extent of dissociation of a weak acid in water, it’s known as the Acid Dissociation

Constant, Ka.

And realize that the ionization of ANY weak base (B) in water can be represented as:

B (aq) + H2O (l) � OH-

(aq) + HB +

(aq)

And that the equilibrium constant expression for the dissociation of ANY weak base can be written as:

-3 = �0.&!�.41! �4!

As this equilibrium constant describes the extent of ionization of a weak base in water, it’s known as the Base Ionization

Constant, Kb.

Use equilibrium constants for the dissociation of weak acids (Ka, acid dissociation constant) to compare the proton-donating

ability of weak acids. That is, use the Ka values to compare the relative amounts of hydronium ion, H3O+, present in each

weak acid solution at equilibrium.


Use equilibrium constants for the ionization of weak bases (Kb, base ionization constants) to compare the proton-accepting

ability of weak bases. That is, use the Kb values to compare the relative amounts of hydroxide ion, OH-, present in each weak

base solution at equilibrium.

Explain how weak acids and their conjugate bases (or weak bases and their conjugate acids) can be used to “buffer” the pH of

a solution.

Show/describe the amphiprotic nature of water

(the self-ionization of water) that occurs in EVERY

aqueous solution.

Use the Ion Product Constant of Water to convert

between hydroxide and hydronium ion

concentrations in ANY aqueous solution.

-5 = �./01!�0.&! = 1.0 × 10&89

Explain the pH method of expressing the hydronium ion concentration of an aqueous solution. pH = – log [H3O+]

Given a hydronium ion concentration, calculate the pH.

Given a pH, calculate the hydronium ion concentration. That is, if pH = – log [H3O+], then [H3O

+] = 10

– pH

Explain the pH (and relative

hydronium and hydroxide ion

concentrations) of pure water, a

neutral solution, an acidic

solution, and a basic solution.

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Chemistry 105 ~ Skills Summary for Laboratory Final · PDF fileChemistry 105 ~ Skills Summary for Laboratory Final Exam ... to the appropriate number of significant figures/decimal