Chem 1151: Ch. 1

Chem 1151: Ch. 1

Matter, Measurements and Calculations

Chemistry: The Central ScienceChemistry: The Central Science

Chemistry

Biology

GeologyAstronomy

Physics

• Chemistry is how matter is organized/reorganized/changed at the molecular level

• Chemistry is often called the central science because it is an essential component of the natural and life sciences.

•Elements are created from the fusion reactions inside stars.•Hydrogen burning – 4 protons become alpha particle (helium nucleus).•Helium burning - 3 alpha particles form 12C then 16O.•Carbon and oxygen burning produce 28Si, 24Mg , 32S, and other elements.•Each of these requires more heat than the fusion reaction before it.•Uranium is last naturally-occuring element as a result of solar activity.

http://www.geophysics.rice.edu/department/faculty/sawyer/ESCI324/esci324_spr_06_lec2.ppt#866,19,Slide 19

Source of ElementsSource of Elements

Chemistry and AstronomyChemistry and Astronomy

• Elemental composition of stars can be determined by different wavelengths of visible light emitted.

• When starlight passes through a planets atmosphere, certain frequencies of light disappear because they are absorbed by compounds in the atmosphere.

http://cs.fit.edu/~wds/classes/cse5255/cse5255/davis/text.html

Chemistry and GeologyChemistry and Geology• Geochemistry: Study of the chemical composition of the earth• Chemical transformations in solids

– Ex. Polymorphism.– How limestone becomes marble.

http://geology.com/rocks/limestone.shtml; http://www.italartworld.com/

http://geology.com/rocks/limestone.shtml

PolymorphismPolymorphism

Graphite: Each Carbon is covalently bonded to 3 other carbons in ring

Mohs scale hardness: 1-2

Diamond: Each carbon is bonded to 4 other carbons

Mohs scale hardness: 10

T+P

Value: $0.10 Value: $1000.00

http://en.wikipedia.org/wiki/Image:Graphit_gitter.png

http://en.wikipedia.org/wiki/Image:Diamond_unit_cell.PNG

Chemistry and BiologyChemistry and Biology

• You reach a certain level in biology where processes can only be understood in terms of chemistry.

• Chemistry in biology explains:– Why your adrenaline levels increase when you are afraid or excited – Why a body fails to produce insulin (diabetes)– Why cells become cancerous– Neurotransmitter (e.g., dopamine, norepinephrine) imbalances that

can produce:• Euphoria when you have a few beers or fall in love• Depression

http://cs.fit.edu/~wds/classes/cse5255/cse5255/davis/text.html

Matter, Mass and WeightMatter, Mass and Weight• Matter is anything that has mass and occupies space.

• Mass is a measurement of the amount of matter in an object.• Mass is independent of the location of an object.• An object on the earth has the same mass as the same

object on the moon.• Weight is a measurement of the gravitational force acting on

an object.• Weight depends on the location of an object.• An object weighing 1.0 lb on earth weighs about 0.17 lb

on the moon.

Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

Properties of MatterProperties of Matter

Matter has Physical Properties:•These can be observed/measured without changing the composition of matter

– Appearance– Texture– Color– Odor– Shape– Melting Point (MP)– Boiling Point (BP)– Density (D)– Solubility

Physical Changes of Matter: – Smashing brick with a hammer– Freezing water or melting ice

Properties of MatterProperties of Matter

Matter has Chemical Properties:•These can be observed when you try to change matter from one form to another, i.e., you have different arrangements of atoms or molecules

– Reactivity– Flammability (whether something is ignitable, flash point < 100 °F)– Combustibility (whether something will burn, flash point > 100 ° F)

Chemical Changes of Matter:– Setting a piece of paper on fire– Reaction between an acid and a base

Particulate Model of MatterParticulate Model of Matter

• All matter is made up of tiny particles called molecules and atoms.

• MOLECULES• A molecule is the smallest particle

of a pure substance that is capable of a stable independent existence.

• ATOMS• Atoms are the particles that

make up molecules.


Molecule ClassificationsMolecule Classifications1. By number of atoms:

Diatomic molecules contain two atoms. Triatomic molecules contain three atoms. Polyatomic molecules contain more than three atoms.

2. By type of atoms: Homoatomic molecules contain same kind of atoms Heteroatomic molecules contain 2 or more kinds of atoms


Molecule Classifications (continued)Molecule Classifications (continued)3. By chemical and physical properties:


•Constant composition•Fixed physical and chemical properties•Formed by chemical interactions of atoms

•Composition can vary•Physical and chemical properties can vary•Compounds retain their chemical identities and can be separated

Mixtures

Pure Substances

Molecule Classifications (continued)Molecule Classifications (continued)3. By chemical and physical properties:


Matter

Mixtures Pure Substances

Heterogeneous Homogeneous (Solutions)

Elements Compounds

•More than 1 kind of atom (heteroatomic)•Can be divided into simpler compounds or elements (complex sugarsimple sugar atoms)•H2O, CO2, CO, NaCl

•Only 1 kind of atom (homoatomic)•Simplest pure substances•Can be divided but chemically same•O2, H2, Au

•Properties of a sample vary by location•Oil and water

•Properties of a sample same throughout•Sugar and water

Measurements and UnitsMeasurements and Units

• Scientific measurements consist of a number and a standard metric unit.

• Measurements are made using measuring devices (e.g. rulers, balances, graduated cylinders, etc.).


Measurements and UnitsMeasurements and Units

• Scientific measurements consist of a number and a standard metric unit.

Common Metric Prefixes Prefix Abbrev. Relationship to base unit Common Examples

mega- M 106 MB - Megabytes (computer)

kilo- k 103 kg – kilograms (weight/mass)

deci- d 10-1 or 1/101 dL – deciliters (volume)

centi- c 10-2 or 1/102 cm – centimeters (length)

milli- m 10-3 or 1/103 mL – milliliters (volume)

mm – millimeters (length) ms – millseconds (time)

micro- µ 10-6 or 1/106 µL – microliters (volume)

µm – micrometers or microns (length)

nano- n 10-9 or 1/109 nm – nanometers (length) ns – nanoseconds (time)

pico- p 10-12 or 1/1012 ps – picoseconds (time)

Temperature ScalesTemperature Scales

• The three most commonly-used temperature scales are the Fahrenheit, Celsius and Kelvin scales.

• The Celsius and Kelvin scales are used in scientific work.


Temperature ConversionsTemperature Conversions

• Readings on one temperature scale can be converted to the other scales by using mathematical equations.

• Converting Fahrenheit to Celsius.

• Converting Celsius to Fahrenheit.

• Converting Kelvin to Celsius.

• Converting Celsius to Kelvin.

32F9

5C

32C5

9F

273KC

273CK


Scientific NotationScientific Notation

• Scientific notation provides a convenient way to express very large or very small numbers.

• Numbers written in scientific notation consist of a product of two parts in the form M x 10n, where M is a number between 1 and 10 (but not equal to 10) and n is a positive or negative whole number.

• The number M is written with the decimal in the standard position.

Scientific Notation (continued)Scientific Notation (continued)

• STANDARD DECIMAL POSITION• The standard position for a decimal is to the right of the first

nonzero digit in the number M.

• SIGNIFICANCE OF THE EXPONENT n• A positive n value indicates the number of places to the right of

the standard position that the original decimal position is located.• A negative n value indicates

the number of places to the left of the standard position that the original decimal position is located.

Scientific ↔ Standard NotationScientific ↔ Standard Notation• Converting from scientific notation to standard numbers1.1 x 102 = 1.1 x 10 x 10 = 1.1 x 100 = 110 Decimal 1.1 x 10-2 = 1.1/ (10 x 10) = 1.1/100 = 0.011 Decimal

Converting ExponentsConverting ExponentsEx. 1 Ex. 2 0.67 x 10-5 0.067 x 10-4 0.0067 x 10-3 0.00067 x 10-2 0.000067 x 10-1 0.0000067 x 100 0.00000067 x 101 0.000000067 x 102

1.2 x 10-3 0.12 x 10-2 0.012 x 10-1 0.0012 x 100 0.00012 x 101

• When you move the decimal (l or r), the exponent will be

equal to number of places you moved the decimal.

Standard to Scientific NotationStandard to Scientific Notation

60023.5 345.233 -345.233 0.00345 0.10345 1.42

6.00235 × 104

3.45233 × 102

-3.45233 × 102

3.45 × 10-3

1.0345 × 10-1

1.42 × 100

Scientific to Standard NotationScientific to Standard Notation

7.932 × 105 4.01 × 104 3.220 × 102 5.673 X 101 3.142 X 10-3 7.6 X 10-6 4.5655 X 10-4

79320040100322.056.730.0031420.00000760.00045655

Math Operations with Scientific NotationMath Operations with Scientific Notation

Multiplication

Division

)10)(()10)(10( 2424 baba

)10()10(

)10( 242

4

b

a

b

a

Addition/Subtraction Convert numbers to the same exponents

(5.00 x 102) + (6.01 x 103) =(0.500 x 103) + (6.01 x 103) = (5.00 x 102) + (60.10 x 102) = (5.00 + 60.10) x 102 = (65.10 x 102) = 6510

(6.01 x 103) - (5.00 x 102) =(6.01 x 103) - (0.500 x 103) = (60.10 x 102) - (5.00 x 102) = (60.10 - 5.00) x 102 = (55.10 x 102) = 5510

Examples of Math OperationsExamples of Math OperationsMultiplicationa. (8.2 X 10-3)(1.1 X 10-2) = (8.2 X 1.1)(10(-3+(-2))) = 9.02 X 10-5 Not using SigFigs b. (2.7 X 102)(5.1 X 104) = (2.7 X 5.1)(102+4) = 13.77 X 106 Now change to Scientific Notation 1.377 X 107

Division a. 3.1 X 10-3 = (3.1/1.2)(10-3-2) = 2.6 X 10-5

1.2 X 102

b. 7.9 X 104 = (7.9/3.6)(104-2) = 2.2 X 102

3.6 X 102

Adding/Subtracting 3.05 X 103 + 2.95 X 103 = (3.05 + 2.95)(103) = 6.0 X 103

Significant FiguresSignificant Figures

• Significant figures are the numbers in a measurement that represent the certainty of the measurement, plus one number representing an estimate.

Q: When is a number NOT significant?A: Look at the zerosLeading zeros are NOT significant. 0.00123Confined zeros ARE significant. 0.00103Trailing zeros ARE significant, when decimal visible 0.0012300But NOT significant if no decimal 12300


Calculations with Significant FiguresCalculations with Significant Figures

Product (multiplication) or quotient (division) must have same number of sig figs as value with the fewest number of sig figs.


SF 2 SF 2SF 4

194625.19 5.4 325.4

SF 2 SF 2SF 4

96.0169.0 5.4 325.4

Ex. 01

Ex. 02

Calculations with Significant FiguresCalculations with Significant Figures

• Sum (addition) or difference (subtraction) must contain the same number of places to the right of the decimal (prd) as the quantity in the calculation with the fewest number of places to the right of the decimal.


prd 1 prd 1prd 3

8.10825.10 5.5 325.5

prd 1 prd 1prd 3

2.0175.0 5.5 325.5

Ex. 01

Ex. 02

Calculations with Significant FiguresCalculations with Significant FiguresRules for Rounding•If the first nonsignificant figure to drop from your answer is ≥ 5, all nonsignificant figures dropped, last significant figure increased by 1.•If the first nonsignificant figure to drop from your answer is < 5, all nonsignificant figures dropped, last significant figure stays the same.

Exact NumbersNumbers with no uncertainty, or are known values. Exact numbers do not change.

Ex. 1 foot is always = 12 inches. It will never be = 12.5 inches.

Not used to determine sig figs•Na = 1 mol = 6.02 x 1023

•π= 3.142•1m = 1000 mm

Dimensional Analysis (Factor-unit method)Dimensional Analysis (Factor-unit method)• The factors used in the factor-unit method are fractions derived from

fixed relationships between quantities• An example of a definition that provides factors is the relationship

between meters and centimeters: 1m = 100cm. This relationship yields two factors:

m 1

cm 100

cm 100

m 1

mm 1000

cm 100

mm 1000

m 1


Factor Unit Method ExamplesFactor Unit Method Examples• A length of rope is measured to be 1834 cm. How many meters is

this?• Solution: • Write down known quantity (1834 cm). • Set known quantity = units of the unknown quantity (meters). • Use factor (100 cm = 1 m), to cancel units of known quantity (cm)

and generate units of the unknown quantity (m). • Do the math.

m 34.18cm 100

m 1cm 1834

m cm 1834


Factor Unit Method ExampleFactor Unit Method Example

Q: If an arrow shot from a bow travels 30 yards in 1 second, many cm does it travel in 4 seconds?

•Time = 4 s•Rate = 30 yards/sec•1 yard = 3 feet•1 foot = 12 in.•1 in = 2.54 cm

cm ?in 1

cm 2,54

foot 1

in 12

yard 1

feet 3

s 1

yards 30 s 4


Percentage Percentage • The word percentage means per one hundred. It is the

number of items in a group of 100 such items.

• PERCENTAGE CALCULATIONS• Percentages are calculated using the equation:

• In this equation, part represents the number of specific items included in the total number of items.

100whole

part%


Percentage CalculationPercentage Calculation• A student counts the money she has left until pay day and

finds she has $36.48. Before payday, she has to pay an outstanding bill of $15.67. What percentage of her money must be used to pay the bill?

• Solution: Her total amount of money is $36.48, and the part is what she has to pay or $15.67. The percentage of her total is calculated as follows:

%96.4210048.36

67.15100

whole

part%

Density and Specific GravityDensity and Specific Gravity• Density is the ratio of the mass of a sample of matter divided by

the volume of the same sample.• Specific gravity is the ratio of the density of a compound relative

to the density of water (DHOH = 1.0 g/cm3)

volume

massdensity

V

mD

HOH

cmpd

D

DSG

Density CalculationDensity Calculation• A 20.00 mL sample of liquid is put into an empty beaker that

had a mass of 31.447 g. The beaker and contained liquid were weighed and had a mass of 55.891 g. Calculate the density of the liquid in g/mL.

• Solution: The mass of the liquid is the difference between the mass of the beaker with contained liquid, and the mass of the empty beaker or 55.891g -31.447 g = 24.444 g. The density of the liquid is calculated as follows:

mL

g222.1

mL 20.00

g 444.24

v

md

Energy CalculationsEnergy CalculationsQ: In order to lose 1 lb/week, you need to cut 500.0 Cal from your diet each day, or use an equivalent number of joules by working each day. How many equivalent joules would you have to spend on work to achieve this each day? How many joules would you have to expend to achieve this over 7 days?

To answer this, you need to first know the following:• 1Cal = 1 kcal = 1000 scientific calories or 1 nutritional calorie• 1 scientific calorie = 1 cal = 4.184 J

Useful Conversions/ConstantsUseful Conversions/Constants

• Na = 1 mol = 6.02 x 1023

• π= (Circumference of circle/diameter of circle) = 3.142

• 1 scientific calorie = 4.184 J• 1Cal = 1 kcal = 1000 scientific calories or 1

nutritional calorie• 1 N = (kg)(m)/s2

• g = 9.81 m/s2 (earth’s gravity)

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Chem 1151: Ch. 1

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