Chemistry 111 Fall, 2004

Elizabeth Martin: Room 311 Science Center, 953-5516, martine@cofc.edu

Week 1, Fall, 2004 Chemistry 111 Welcome to Chemistry 111

This course invites you to join the body of human beings who have sought and who are presently seeking to understand the observable universe and to build models to explain its behavior. It is said that Knowledge is Power. Since every THING is chemical, knowledge of chemistry is quite powerful. It is Power to the biologist who needs to understand how life “works”, Power to the doctor who needs to prevent and cure disease, Power to the entrepreneur who wants to make a new, useful, safe product.

In this course you will discuss and wrestle with a lot of ideas, concepts, models, mathematical relationships that form the basis of dynamic investigation of the world that is CHEMISTRY. Open your text and look through the first 11 chapters. What a gold mine of information! What an opportunity for you to build a powerful base for understanding the physical world and how it works!

How you master new knowledge and connect new ideas to ideas and experiences you already have as part of your knowledge-base is, therefore, one of the main keys to success.

A review checklist: From Chapter One of your textbook: Chemistry by Raymond Chang

* Define physical properties of matter and give some examples.

* Recognize the different states of matter (solids, liquids, gases); their characteristics.

* Understand the basic ideas of the kinetic molecular theory.

* Convert between temperatures on the Celsius and Kelvin scales.

* Use density as a way to connect the volume and mass of a substance. The density of a substance is defined as the ratio of its mass to its volume. Density = mass/ volume In chemistry we generally use density in units of grams per cubic centimeter.

* Identify the name or symbol for an element, given its symbol or name .

* Be able to use the terms atom, element, molecule, ion and compound correctly .

* Explain the difference between chemical and physical change.

* Recognize the difference between homogeneous and heterogeneous mixtures .

* Recognize and know how to use the prefixes that indicate the sizes of metric units .

* Begin using dimensional analysis to carry out unit conversions and other calculations.

* Know the difference between precision and accuracy .

* Understand the reason and the rules for using significant figures.

* Use the concept of percent in chemical problems.

Matter: “Any thing that has mass and takes up space”. Properties of Matter:

Observables:

qualitative properties of matter__________________________________________

quantitative properties of matter:____________________________________________

extensive properties:(depend on amount of sample)properties:______________________

intensive (do not depend on amt. of sample)properties____________________

physical properties_______________________________________________

chemical properties_______________________________________________

Matter: Our Working Theory

1. Elements: metals, non-metals, metalloids

2. Compounds:

* *

Molecular compounds ionic compounds

Composed of molecules composed of ions(ions are charged particles)

Cations have + charge

Anions have - charge

Classification of matter

Matter: Heterogeneous Homogeneous

Non-uniform Mixtures Solutions(homogeneous mixtures)

Pure Materials: Elements and Compounds

Phases of Matter:

Description(solids, liquids, gases) and Microscopic description:

·

· solid liquid gas

The Kinetic Molecular Theory: Particles are in constant motion. In solids the particles are close together and have limited motion. In a liquid some of the attraction between particles is overcome which allows the particles more freedom of movement. In a gas particles attraction between particles is minimized and the particles move freely throughout the container.

Two Properties related to the phase of a sample of matter:

1.Density D=mass/volume

2.Temperature: Temperature scales [Fahrenheit, Centigrade(Celsius), Kelvin]

Converting from one to the other:

^o F= 9/5 (^o C) + 32 ^o F

K(for Kelvin scale) = ^o C + 273.15 ^o C

The difference between temperature and heat

temperature= indication of the average kinetic energy.

heat= a form of energy(the ability to do work.)

How does adding heat affect particles? Three things that can happen:

1. Physical Change: The temperature rises. What happens to the particles?

2. Physical Change: A solid melts, a liquid evaporates. What is happening to the

particles?

3. Chemical Change: A chemical reaction may occur. What is happening to the particles?

Measurement- Metric/SI Units

A. Accuracy and precision [A similar diagram is on p. 26]

Accuracy tells us how close a measurement is to the true value of the quantity that was measured. The middle diagram represent accurate aim.

Precision refers to how closely two or more measurements of the same quantity agree with one another. The darts in the middle diagram show high precision. However, the darts in the last diagram also indicate

high precision but low accuracy!

*

B. Significant Figures: Recording and using measured values:

(Guidelines pp 24 & 25). The FIRST important thing is to realize that you have three tasks to learn:

How to significantly record data from measuring devices when YOU are the measurer

How to use measured values in calculations

How to recognize the number of significant digits in a recorded measurement

1. When you are the measurer: Read every digit from the scale on the measuring device AND one estimated place value.

2. When you use measured values. There are two major rules: [p.24&25]

a. When adding or subtracting: Watch the decimal place: Keep the answer to no more decimal places than the least specific value. In this example, the answer must be rounded to the tenths place.

Example : 89.332

+1.1___

________________(Check your answer on p. 24 of Chang)

b. When multiplying or dividing: Count the # of digits: Keep the answer to

the same number of digits as the value with the least number of significant digits. In this

example, the answer must be rounded off to have a total of two digits.

Example: 2.8 x 4.5039 = ________(Check your answer on p. 25 of Chang)

3. When you must recognize the number of significant digits in a value to be used in a calculation: Keep in mind that the person who recorded the measured value is following the same rules that you follow when recording a measurement and that significant digits are the

that tell you about the instrument on which the measurement was made.

a. In a measured value all non-zero digits are significant.

b. In a measured value the rules/examples for handling zeros are found on page 24.

A zero between two non-zero digits is significant.

Zeros to the left of the first non-zero digit are not significant. (Their purpose is to

indicate the decimal point.)

Zeros to the right of the decimal:

For a number greater than 1, zeros to the right of the decimal are

significant.

For a number less than 1, zeros following the first non-zero digit are

significant.

For a measured value with no decimal point, trailing zeros (those zeros that follow

the last non-zero digit) are a problem for the person reading and using the

value. The problem is that the user does not know whether the zeros there are to indicate that the measuring device read “zero” in those decimal places or whether the zeros there are just to indicate the size of the number. It is the responsibility of the person recording the number to use scientific notation to make the role of the trailing zeros clear. When the measurer has not done that, you must assume that trailing zeros are not significant.

c. In a counted number there are an infinite number of significant digits.

Check your laboratory manual Self-Tests for more practice on significant figures.

Solving Problems in Chemistry:

1. Read to understand the situation the problem is describing. It is helpful to draw a simple picture to help concentrate on translating the words into a description of a situation for which you will be asked to give an answer.
2. List and Label the information given in the problem.
3. List and Label what you are being asked to determine. Example: Mass in grams
4. List any relationships/formulas that you know that involve the given information and the quantity you are being asked to provide.
5. Solve for what you can with the information you have. If that does not solve the problem, review your notes, class materials and text for other relationships that do link what you know with what you need to determine.

In-Class Quiz #1: Due at the beginning of the second class period.

1. Give the symbol for each of the following elements:

(a) Lithium ______ (e) Silicon ________

(b) Titanium ______ (f) Cobalt ______

(c) Iron ______ (g ) Zinc _______

(d) Fluorine ______ (h) Bromine _______

2. Of the elements in question 1, which are:

metals? ________________

non-metals? ________________

metalloids? _______________

3. List the symbols and names for: 2 Alkali Metals __________, __________;

2 Noble Gases _____________, _____________; 2 Halogens _____________, ______________

4. Make a drawing, based on the kinetic-molecular theory and your ideas about atoms and molecules, of a submicroscopic model of the arrangement of particles in each of the cases listed below. Represent each atom as a circle and distinguish each different kind of atom by shading its circle. [See pages 13 and 14 for examples.]

(a) A sample of solid aluminum (b) A sample of liquid water

(consists of aluminum atoms) (consists of H₂O molecules)

5. A chemist needs 2.00 g of a liquid compound.

(a) What volume of the compound is necessary if the density of the liquid is 0.718 g/cm ³ ?

(b) If the compound costs $2.41 per milliliter, what is the cost of 2.00 g?

6. Many laboratories use 25 °C as a standard temperature. What is this temperature in kelvins?

7. Give the number of significant figures in each of the following measured values:

(a) 9.87 m _______ (c) 1050 km ______

(b) 0.00823 cm ______ (d) 1.607 g _______

8. A typical laboratory beaker has a volume of 800. mL. What is its volume in cubic centimeters? _______ In liters? _______

9. What is the average mass of three objects whose individual masses are 10.3 g, 9.234 g, and 9.35 g?

__________

10. Assume these are number s are all measured values. Solve the problem to the correct number of significant figures. 1.0000 (1.68)(7.847) = ___________

11. The platinum-containing cancer drug cisplatin is 65.0% platinum. If you have 1.53 g of the compound, how many grams of platinum can be recovered from this sample? ______________

12. Given:

Pressure (P) is equal to the force (F) an object exerts on a given area (A), i.e., the

relationship defining pressure is P= F/A.

The force an object exerts is the mass (m) of the object times its acceleration (a),

i.e., the relationship defining force is F=(m) x (a).

With this information, determine the mass in grams (g) of an object moving with an acceleration of 10 cm/sec² which is found to exert a pressure of 40 dynes/cm² on an area of 15 cm²? (A dyne is = 1 (g)(cm)/sec²). [Check problem-solving strategy discussed above.]

Atomic Theory

Democritus and Leucippus(Greek)

Lucretius(Roman)

Empirical “Laws” generally accepted by the end of the 1700”

Law of Conservation of Matter: Antoine Lavoisier

Matter is neither created nor destroyed in chemical or physical changes(Those

were the only types of changes they knew about at the time.)

Law of Constant Composition: Joseph Louis Proust

A Compound is always composed of the same elements and, when analyzed by mass, the ratio of the masses of the elements making up the samples is the same. For example: Analysis of a Compound X: If 2 grams of hydrogen reacted with 16 grams of oxygen, they would find that 1 gram of hydrogen reacted with 8 grams of oxygen.

Law of Multiple Proportions”: If the same two elements(A and B) formed two

different compounds, there is a whole number relationship between the masses

of B in the two different compounds for each gram of A in the compound. For

example: 1 gram of Element A reacts with 8 grams of Element B to make

compound X; whereas 1 gram of Element A reacts with 16 grams of B to make

compound Y. For the same amount of A in the two compounds, there was twice

as much B in compound Y compared to compound X.

John Dalton’s Atomic Theory: (1803) Proposed that all matter was composed of atoms of

different elements. Atoms were small, indestructible, indivisible, particles. He was able to propose values for relative atomic weighs of each element. His theory proposed:

All Matter is made of tiny particles called atoms.

Atoms are small, indestructible, indivisible particles.

All atoms of an given element are identical in atomic weight.

Compounds are formed by combination of two or more different kinds of

atoms.

Chemical chang is due to atoms recombining in different combinations

Early Success of AtomicTheory:

It could explain the “Law of Conservation of Matter”

It could explain the “Law of Constant Composition”

It could explain the “Law of Multiple Proportions”

A Look at What’s Happening in Chemistry in the years following Dalton’s Theory- in the Middle 1800s ?

Development of The Periodic Table: Demetri Mendeleev(Russia)

Lewis R. Gibbes(Charleston)

The Periodic Tab;e

groups or “families” = vertical columns

periods

metals, non-metal, metalloids

Families(Groups)

Names of some of the families

alkali metals

alkali earth metals

halogens

noble gases

transition metals

lanthanides and actinides

Some Discoveries that led to Modern Atomic Theory

Electricity: Ben Franklin (late 1700’s)

Electrolysis: Michael Faraday (1833)

Radioactivity: Henri Becquerel/the Curies (1896-98)

Electrons: J.J. Thomson- Charge to mass ratio for the electron

Results of the Cathode Ray Experiment: In the experiment both negative (Cathode rays) and positive(Canal rays)were formed. The negative Cathod ray was found to be particles that were the same no matter which element was in the tube. By passing these negatively charged particles(electrons) through an electric field(top diagram) and a magnetic field(bottom diagram), Thomson was able to find the ratio of the charge to mass of these sub-atomic particles.

*

Thomson’s Theory: The Raisin Bun/Chocolate Chip Theory of the Atom

Atoms were composed of smaller particles. He proposed that all atoms were composed of very small negative particles(electrons) that were imbedded in some kind of positive “stuff” like raisins in a raisin bun. His experiments gave a value for the ratio of charge to mass of the electron: Charge/mass = - 1.76 x 10¹¹ coulombs/kg of electrons

Electrons: Millikan- Oil-drop Experiment: Millikan was able to find the charge of the electron. When that is combined with Thomson’s ratio of charge to mass, it is then possible to calculate the mass of an electron as well.

*

Protons: Investigation of the positive(Canal Rays) generated in the “Cathode Ray “ tube led to the proposal for the existence of a positively charged particles(ions) which were different for each element. These particles differed in mass and charge depending on the element that was being tested. Later, it was proposed that they were the “rest of the atom” after some electrons had been lost. This picture is of the “canal rays” at the other end of the Cathode Ray Tube.

The Nucleus: Rutherford- The Gold Foil Experiment was to test Thomson’s theory of the atom. What happened to the alpha particles? Most went right through the foil but a number were scattered at angles and some essentially bounced back toward the source!

*

Rutherford’s Theory of the Atom: a very small, heavy, positive region(called the nucleus) with the positive particles(called protons) and some other neutrally charged material(later discovered and named neutrons) was surrounded by the small, negative electrons. Atoms were mostly empty space with a very small, dense nucleus.

Atomic Composition:

Protons: The number of protons/atom defines the atom and is called the atomic #.

Electrons

Neutrons

Isotopes

Discovering the existence of isotopes:

The Mass Spectrometer: In a mass spectrometer, charged particles are passed through a magnetic field. The path of the stream of particles will be bent by the magnetic field depending on the mass and charge of the particles. When naturally-occuring samples of most elements are charged and passed through a mass spectrometer, the spectrum indicates more than one form of the element. These are isotopes of the element: the number of protons is the same but the number of neutrons/atom differs for each isotope giving each isotope a slightly different mass.

Isotopes

Definition

Define symbolism to represent a particular atom of an isotope:

A charge A= Mass number = # protons + # neutrons

X

Z Z= atomic number

Examples:

Symbol Cu ______

# protons _______ 10

# neutrons _______ 10

# electrons _______ _____

Charge of particle _______ 0

Atomic number(Z) _______ ______

Mass number(A) _______ ______

Name of the element ___________ ____________