chem 100 fall 2014 la tech instructor: dr. upali siriwardane e-mail: [email protected] office: cth...

CHEM 100 Fall 2014 LA Tech

Instructor: Dr. Upali Siriwardanee-mail: [email protected]: CTH 311 Phone 257-4941Office Hours: M,W, 8:00-9:30 & 11:30-12:30 a.m Tu,Th,F 8:00 - 10:00 a.m. Or by appointmentTest Dates:

Chemistry 100(02) Fall 2014

September 29, 2014 (Test 1): Chapter 1 & 2 October 20, 2014 (Test 2): Chapter 3 & 4 November 12, 2014 (Test 3) Chapter 5 & 6November 13, 2014 (Make-up test) comprehensive: Chapters 1-6 9:30-10:45:15 AM, CTH 328

REQUIRED :Textbook: Principles of Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro - Pearson Prentice Hall and also purchase the Mastering Chemistry Group Homework, Slides and Exam review guides and sample exam questions are available online: http://moodle.latech.edu/ and follow the course information links.OPTIONAL : Study Guide: Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro 2nd EditionStudent Solutions Manual: Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro 2nd

Text Book & Resources

http://moodle.latech.edu/


Chapter 1. Matter, Measurement, and Problem Solving

1. 1 Atoms and Molecules………………………………….. 11 .2 The Scientific Approach to Knowledge…………….. 31 .3 The Classification of Matter…………………………… 51 .4 Physical and Chemical Changes and Physical and Chemical Properties…………………………………….. 91 .5 Energy: A Fundamental Part of Physical and Chemical Change…………………………………………………….. 121 .6 The Units of Measurement……………………………... 131 .7 The Reliability of a Measurement……………………… 201 .8 Solving Chemical Problems……………………………. 27


Chapter 1. KEY CONCEPTS

• What is chemistry?• Scientific Method.• Properties of the three states of

matter• Physical changes and

properties.• Chemical change and

properties.• Categories of matter. • Elements and Compounds• Atomic symbols• Chemical Elements and

properties• Chemical Symbolism• Separating Mixtures.• Scientific Measurement

• Prefixes of SI units• Macro, micro and nano-scales • Conversion factors.• Factor label method.• Uncertainty and significant

figures• Temperature Conversions.• Density Calculations.• Three chemical Laws• Dalton's atomic theory • Interpreting chemical formulas

and chemical reaction.• Concept of mole• Gram to mole conversion


Units


Length meter (m)

SI UNITS OF MEASUREMENT

Mass kilogram (kg)

Time second (s)

Temperature kelvin (K)

Amount mole (mol)

6.02 x 1023 units

Five Basic Units


The Units of Measurement

1) Give the name and abbreviation of the SI Unit for:

a) Length b) Mass c) Time

d) Amount of substance e) Temperature


Prefix SymbolDecimal

EquivalentPower of 10

mega- M 1,000,000 Base x 106

kilo- k 1,000 Base x 103

deci- d 0.1 Base x 10−1

centi- c 0.01 Base x 10−2

milli- m 0.001 Base x 10−3

micro- m or mc 0.000 001 Base x 10−6

nano- n 0.000 000 001 Base x 10−9

pico p 0.000 000 000 001 Base x 10−12

Metric SystemHow to change measurements to reasonable numbers


Macroscale, Microscale, and Nanoscale


2) Give the abbreviation for the following units and describe what they are used to measure:

a) cubic centimeter b) micrometer

c) nanoseconds d) millimole


3) Give the name and the abbreviation (without looking in the book) of the SI or metric prefix for:

a) 10-12 b) 106 c) 10-9 d) 10-2

e) 10-3 f) 109 g) 103 i) 10-6

CHEM 100 Fall 2014 LA Tech12

Uncertainty in a Measurement

• The last digit is an estimate.

• .Measurement ≈ 26.13 cm

types of errors, random and systematic.

Careless measurementsLow resolution instrumentsCalibration errors


Uncertainty and Significant Figures

All measurements involve some uncertainty.

Scientists write down all the digits that have no uncertainty plus one additional uncertain digit.

If an object is reported to have a mass = 6.3492 g, the last digit (“2”) is uncertain ( it is probably close to 2, but may be 4, 1 … etc).

There are five significant figures in this number. All the digits are meaningful.



To find the number of significant figures:

Read a number from left to right and count all digits, starting with the first non-zero digit.

All digits are significant except those zeros that are used to position a decimal point (“placeholders”).

0.00034050

5 sig. figs.

Scientific Notation (3.4050 x 10-4)

placeholders

significant

significant


• All non-zero digits are significant.– Example: 536 has three significant figures.

• Zeros between non-zero digits are also significant.– Example: 6703 has four significant figures.

• Place holder zeros:– Zeros to the left of a non-zero digit are NOT significant.

• Example: 0.0043 has two significant figures.• Example: 0.0600 has three significant figures.

– Zeros to the right or after a non-zero digit to the decimal point are NOT significant.• Example: 7000 has only one significant figure.• Example: 32040 has four significant figures.• Example: 50.0 has three significant figures; all the

zeros are significant.

Significant Figure Primer



Number Sig. figs. Comment

2.12 3

4.500 4 The zeros are not placeholders. They are significant.

0.002541 4 The zeros are placeholders (not significant).

0.00100 3 Only the last two zeros are significant.

500 1, 2, 3 ? Ambiguous. If a number lacks a decimal point the zeros may be placeholders or may be significant.

500. 3 Adding a decimal point is one way to show that the zeros are significant.

5.0 x 102 2 No ambiguity.

Examples


The Reliability of a Measurement

4) Express the following numbers in scientific notation with the appropriate number of significant figures:

a) 10980000000 b) 414100

c) 0.000095162 d) 746.5 x 107


Consider rounding 37.663147 to 3 significant figures.

RoundingLook at the 1st non-significant digit (the digit after the last one retained). If it:

is > 5, round the last retained digit up by 1.

is < 5, make no change.

equals 5, and the 2nd non-significant digit is:absent, round the last retained digit up by 1.

odd, round the last retained digit up by 1.

even, make no change.

last retained digit

1st non-significant digit

It rounds up to 37.7

2nd non-significant digit


RoundingExamplesRound the following numbers to 3 significant figures:

1st non-sig. 2nd non-sig.Rounded

Number digit digit Number

2.123 2.123 - 2.12

51.372 51.372 51.372 51.4

131.5 131.5 - 132.

24.752 24.752 24.752 24.7

24.751 24.751 24.751 24.8

0.06744 0.06744 - 0.0674


Can you define accuracy vs. precision?

Precision = the degree of exactness of a measurement that is repeatedly recorded.Accuracy = the extent to which a measured value agrees with a standard value• Which set is more precise?

18.2 , 18.4 , 18.3517.9 , 18.3 , 18.8516.8 , 17.2 , 19.44


Three targets with three arrows each to shoot.

Can you hit the bull's-eye?

Both accurate and precise

Precise but not accurate

Neither accurate nor precise

How do they compare?


Speed of light is 3.00 x 108 m s-1 . Convert the speed of light to miles per year (1 mile = 1.61 km).

Unit Conversion Calculation


Using Conversion Factors

5) Convert 78.01 inches into:

a) feet b) meters

6) Convert 15.42 meters into:

a) kilometers b) micrometers


Example: 301 in scientific notation is 3.01 x 102.NOTE: The decimal point was moved two

places to the left.

Example: 0.0301 in scientific notation is 3.01 x 10-2. NOTE: The decimal point was moved two

places to the right.

Both of these value indicate THREE significant figures.

Significant Figures and Scientific Notation


adp = 4adp = 3

Significant Figures in Add/Sub

The answer you report in a problem should only include significant digits.Addition and subtraction

Find the number of digits after the decimal point (adp) in each number.

answer adp = smallest input adp.

Example

Add: 17.245 + 0.1001 17.3451 Rounds to: 17.345 (adp = 3)


Examples: 2 3 .4 6 7 in + 3 1 3 .2 1 in 3 3 6 .6 7 7 in but you would report

336.68 in

4 5 7 cm - 0 . 6 8 cm

4 5 6 . 3 2 cm but you would report 456 cm

Addition and Subtractions Examples:


Significant Figures Add/Sub

adp = 2adp = 4

Example

Subtract 6.72 x 10-1 from 5.00 x 101

Write the numbers down with the same power of 10:

5.00 x 101

– 0.0672 x 101

4.9328 x 101

Rounds to: 4.93 x 101 adp = 2


sig. fig. = 4sig. fig. = 5

Significant Figures Mul/Div

Multiplication and DivisionFind the number of significant figures (sig. fig.) in each number.

Answer has sig. fig = smallest input sig. fig.

Example

Multiply 17.425 and 0.100117.245

x 0.1001 1.7262245

Rounds to: 1.726 sig. fig. = 4 ExampleMultiply 2.346 x 12.1 x 500.99

Rounds to: 1.42 x 104 (3 sig. fig.)

= 14,221.402734


7) Perform the following calculations and give the answer in the correct number of significant figures. a) 30.84 + 9.74 Answer:

Scientific notation:

b) 30.84 + 9.74486 Answer: Scientific notation:

c) 145 + 1.54 x 10-6 Answer: Scientific notation:


7) Perform the following calculations and give the answer in the correct number of significant figures.

2 adp 2 adp 2 adp

a) 30.84 + 9.74 Answer: 40.58 Scientific notation: 4.058 x 101

2 adp 5 adp 2 adp

b) 30.84 + 9.74486 Answer: 40.58486 Scientific notation: 4.058 x 101

0 adp 8 adp 0 adp

c) 145 + 1.54 x 10-6 Answer: 145.00000154 Scientific notation: 1.45 x 102


7) Perform the following calculations and give the answer in the correct number of significant figures. d) 40.79 - 1.18432 Answer: Scientific notation:

e) 1.43 x 0.848 Answer: Scientific notation:

f) (7.601x107) x (8.09x10-4) Answer

Scientific notation:


7) Perform the following calculations and give the answer in the correct number of significant figures.

2 adp 5 adp 2 adpd) 40.79 - 1.18432 Answer: 41.97432 Scientific notation: 4.197 x 101

Multiplication

3 sfg 3 sfg 3 sfg

e) 1.43 x 0.848 Answer: 1.21264 Scientific notation: : 1.21 x 100 = 1.21

4 sfg 3 sfg 3 sfg

f) (7.601x107) x (8.09x10-4) Answer 61492.09 =615

Scientific notation: 6.15 x 102


Temperature Scales: Comparison


Temperature Conversions

Human body temperature is 98.6 oF. Convert this temperature to oC and K scale

oC = 5/9 (98.6 - 32) = 5/9 (66.6) = 37.0

oC--> K = 37.0 oC +273.15 = 310.2 K

Shift the scale to zero

Convert the scale 100/180

Shift the scale to zero K


Temperature ConvesionsoF -- > oC ; C = 5/9 (F - 32)oC -- > oF ; F =9/5 C + 32oC -- > K ; K = C + 273.15

8) Convert 98.6 °F into:

a) °C b) K.


Problem Solving by Factor Label Method

– State question in mathematical form– Set equal to piece of data specific to the problem– Use conversion factors to convert units of data

specific to problem to units sought in answer– Other names used Unit Conversion Method

or dimensional (Unit) Analysis


Common Conversion Factors

Length 1 kilometer = 1000 m = 0.62137 mile1 inch = 2.54 cm (exactly)1 angstrom (Å) = 1 x 10-10 m

Volume 1 liter (L) = 1 x 10-3 m3

= 1000 cm3 = 1000 mL= 1.056710 quarts

1 gallon = 4 quarts = 8 pints

Mass1 amu = 1.6606 x 10-24 g1 pound = 453.59237 g = 16 ounces1 ton (metric) = 1000 kg1 ton (US) = 2000 pounds


Solving Chemical Problems 9) Convert 75 miles per hour into: m s-1.


Solving Chemical Problems 10) Convert 100 m2 into cm2

11) Convert 1 m3 into cm3

.


Mathematical operations dictate the reporting of significant figures in an answer.

1. Multiplication and Division

The least precise measured value determines the number of significant figures in the reported answers.

2. Addition and Subtraction

The value with the smallest decimal measurement determines the answer’s significant figure.

Significant Figures andMathematical Operations


Solving Chemical Problems

13) Perform the following mathematical operations and give the answer with the correct number of significant figures

a) (2.481 x 12.74) + 0.27=

2.69

b) (4.73 x 10-4) - (72.85) =

(872.3) - (0.305)

:


Density = mass (g)

volume (cm3) An INTENSIVE physical property

The physical property does not depend on amount of substance.

The physical properties of mass and volume that determine a substance’s density are EXTENSIVE.Extensive physical properties are dependent on amount. Densities of liquid and gases are affected by temperature.

Density


PROBLEM:

Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in grams? In pounds?

Strategy:1. Convert mL to cm3.2. Solve for mass (in grams) using density relationship.3. Convert grams to pounds.

Density Calculations


PROBLEM: Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in grams? In pounds?

Density = mass (g)volume (cm3)

Step 1: 95 mL x (1cm3/1mL) = 95 cm3

Step 2: 13.6 g/cm3 = x / 95 cm3

x = 1.29 x 103 g, but report 1.3 x 103 g

Step 3: 1.3 x 103 g x (1 lb/454 g) = 2.9 lb

A Density Calculation


Solving Chemical Problems 12) Aluminum block weighs 14.2 g and has a density of

2.70 g cm-3. Calculate the volume of the block.


Problem:

The density of octane (C8H18) is 7.00 lb/gal.a) What is density in mg/cm3?b) What is the mass in grams of 1.25 liters of octane?

Strategy:1. Convert 7.00 lbs to mg.2. Change gallons to cm3.3. Determine the density of octane in mg/cm3.4. Convert 1.25 L to mL.5. Determine the mass of octane in 1.25 L using density.

chem 100 fall 2014 la tech instructor: dr. upali siriwardane e-mail: [email protected] office: cth...

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