CHM11-3Lecture Unit #1

General Chemistry 1 LECTURE UNIT No. 1

Introduction to Chemistry

Engr. Edgie Estopace

School of Chemical Engineering and Chemistry

2

Outline

1. The Study Matter

2. States of Matter

3. Chemical and Physical Properties

4. Chemical and Physical Changes

5. Classification of Matter

6. Measurements

3

Matter - Vocabulary

Chemistry Science that describes matter its properties, the

changes it undergoes, and the energy changes that accompany those processes

Matter Anything that has mass and occupies space.

4

Natural Laws

Law of Conservation of Mass

Law of Conservation of Energy

Law of Conservation of Mass-Energy

Einsteins Relativity

E=mc2

5

States of Matter

Solids

6

States of Matter

Solids

Liquids

7

States of Matter

Solids

Liquids

Gases

Solid: particles maintain a regular ordered structure; maintains size and shape.

Liquid: particles remain close but no longer ordered; takes shape of container.

Gas: particles are widely separated and move independently of one another; fills available volume of container.

States of Matter

9

States of Matter

Change States

heating

cooling

10

States of Matter

Illustration of changes in state

requires energy

11

Chemical and Physical Properties

Chemical Properties - chemical changes rusting or oxidation

chemical reactions

Physical Properties - physical changes changes of state

density, color, solubility

Extensive Properties - depend on quantity

Intensive Properties - do not depend on quantity

Physical Change

During a physical change, chemical composition does not change.

Heating liquid water to make gaseous water (steam)

During a chemical change, a chemical reaction occurs that changes the chemical composition of the matter involved.

Using electricity to convert water into oxygen and hydrogen molecules

Chemical Change

14

Density and Specific Gravity

density = mass/volume

What is density?

Why does ice float in liquid water?

15


density = mass/volume

What is density?

Why does ice float in liquid water?

H2O(l) H2O(s)

H

C

HH

H

H

C

HH

H

16


SP1: Calculate the density of a substance if 742 grams of it occupies 97.3 cm3.

Vm density

mL 3.97cm 97.3 mL 1 cm 1 33

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SP2: Suppose you need 125 g of a corrosive liquid for a reaction. What volume do you need?

liquids density = 1.32 g/mL

You do it!

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SP2 Suppose you need 125 g of a corrosive liquid for a reaction. What volume do you need?


density

mV

V

mdensity

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SP2 Suppose you need 125 g of a corrosive liquid for a reaction. What volume do you need?


mL 94.7 1.32

g 125V

density

mV

V

mdensity

mLg

20


Waters density is essentially 1.00 at room T.

Thus the specific gravity of a substance is very nearly equal to its density.

Specific gravity has no units.

)water(density

)substance(densityGravity Specific

Density

mass Volume

Density at 20C Substance d (g/mL)

ethanol 0.789

water 0.998

magnesium 1.74

aluminum 2.70

titanium 4.50

copper 8.93

lead 11.34

mercury 13.55

gold 19.32

m V

d =

Water, copper and mercury

A piece of metal has mass = 215.8 g. It is placed into a measuring cylinder and it displaces 19.1 mL of water. Identify the metal. Density at 20C

Substance d

(g/mL)

magnesium 1.74

aluminum 2.70

titanium 4.50

copper 8.93

lead 11.34

mercury 13.55

gold 19.32

Short Quiz 1:

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Heat and Temperature

Heat and Temperature are not the same thing T is a measure of the intensity of heat in a body

3 common temperature scales - all use water as a reference

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Heat and Temperature are not the same thing

T is a measure of the intensity of heat in a body

3 common temperature scales - all use water as a reference

25


MP water BP water

Fahrenheit 32 oF 212 oF

Celsius 0.0 oC 100 cC

Kelvin 273 K 373 K

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Relationships of the Three Temperature Scales

273KC

or

273 C K

ipsRelationsh Centigrade andKelvin

o

o

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Relationships of the Three Temperature Scales

1.8

32FC

or

32C 1.8F

1.85

9

10

18

100

180

ipsRelationsh Centigrade and Fahrenheit

oo

oo

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SP3: Convert 211oF to degrees Celsius.

1.8

32112C

1.8

32FC

o

oo

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Classification of Matter

Substance matter in which all samples have identical

composition and properties

Elements substances that cannot be decomposed into

simpler substances via chemical reactions

Elemental symbols found on periodic chart

30


31

Name Symbol Name Symbol Name Symbol

Aluminum Al Fluorine F Oxygen O

Arsenic As Gold Au Phosphorus P

Argon Ar Germanium Ge Palladium Pd

Barium Ba Hydrogen H Platinum Pt

Bromine Br Iodine I Potassium K

Calcium Ca Iron Fe Silicon Si

Carbon C Lead Pb Silver Ag

Chlorine Cl Magnesium Mg Sodium Na

Chromium Cr Mercury Hg Sulfur S

Cobalt Co Nickel Ni Tin Sn

Copper Cu Nitrogen N Zinc Zn


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Compounds

substances composed of two or more elements in a definite ratio by mass

can be decomposed into the constituent elements

Water is a compound that can be decomposed into simpler substances hydrogen and oxygen


33


34

Mixtures

composed of two or more substances

homogeneous mixtures

heterogeneous mixtures


Mixtures are either:

homogeneous

two or more substances in the same phase.

completely uniform.

heterogeneous

properties vary from point to point.

may need a microscope to see variation.

Classifying Matter: Substances & Mixtures

Classify each of the following as an element or a compound:

a. Sodium chloride c. Alcohol

b. Helium d. Platinum

Mixtures can be separated by physical methods.

e.g. magnetic separation of iron filings from sulfur powder.

Separation and Purification


Matter (may be solid, liquid, or gas): anything that occupies space and has mass

Homogeneous matter: uniform composition throughout

Heterogeneous matter: nonuniform composition

Substances: fixed composition; cannot be further purified

Solutions: homogeneous mixtures; uniform compositions that may vary widely

Elements: cannot be subdivided by chemical or physical changes

Compounds: elements united in fixed ratios

Physically

separable into

Physically

separable into

Chemically

separable into

Combine chemically

to form

2008 Brooks/Cole

Measurements

Atoms are very small.

1 tsp of water contains 3x as many atoms as there are tsp of water in the Atlantic Ocean!

Impractical to use pounds and inches...

Need a universal unit system

The metric system.

The SI system (Systeme International) - derived from the metric system.

2008 Brooks/Cole

Metric Units

Prefix Factor Example

mega M 106 1 megaton = 1 x 106 tons

kilo k 103 1 kilometer (km) = 1 x 103 meter (m)

deci d 10-1 1 deciliter (dL) = 1 x 10-1 liter (L)

centi c 10-2 1 centimeter (cm) = 1 x 10-2 m

milli m 10-3 1 milligram (mg) = 1 x 10-3 gram (g)

micro 10-6 1 micrometer (m) = 1 x 10-6 m

nano n 10-9 1 nanogram (ng) = 1 x 10-9 g

pico p 10-12 1 picometer (pm) = 1 x 10-12 m

femto f 10-15 1 femtogram (fg) = 1 x 10-15 g

A decimal system.

Prefixes multiply or divide a unit by multiples of ten.

2008 Brooks/Cole

1 pm = 1 x 10-12 m ; 1 cm = 1 x 10-2 m

How many copper atoms lie across the diameter of a penny? A penny has a diameter of 1.90 cm, and a copper atom has a diameter of 256 pm.

x 1 x 10-2 m

1 cm

= 7.42 x 107 Cu atoms

1 pm 1 x 10-12 m

x 1.90 cm = 1.90 x 1010 pm

Number of atoms across the diameter:

1.90 x 1010 pm x 1 Cu atom 256 pm

Metric Units

2008 Brooks/Cole

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

Volume 1 liter (L) = 1000 cm3 = 1000 mL = 1.056710 quarts 1 gallon = 4 quarts = 8 pints

Mass 1 amu = 1.66054 x 10-24 g 1 pound = 453.59237 g = 16 ounces 1 ton (metric) = 1000 kg 1 ton (US) = 2000 pounds

Some Common Unit Equalities

2008 Brooks/Cole

5.0 lb

Report the mass of a 5.0 lb bag of sugar in kilograms.

1 lb = 453. g

= 2265 g x 453. g 1 lb

= 2.3 x 103 g

= 2.3 kg


2008 Brooks/Cole

165 mg dL

A patients blood cholesterol level measured 165 mg/dL. Express this value in g/L

1 mg = 1 x 10-3 g ; 1 dL = 1 x 10-1 L

x 1 x10-3 g

1 mg = 1.65 g/L x 1 dL

1 x10-1 L


2008 Brooks/Cole

All measurements involve some uncertainty.

Reported numbers include one uncertain digit.

Uncertainty and Significant Figures

Consider a reported mass of 6.3492 g

Last digit (2) is uncertain

Close to 2, but may be 4, 1, 0

Five significant figures in this number.

46 2008 Brooks/Cole

Read numbers from left to right.

Count all digits, starting with the 1st non-zero digit.

All digits are significant except zeros used to

position a decimal point (placeholders).

0.00024030

5 sig. figs.

(2.4030 x 10-4)

placeholders significant

significant


47 2008 Brooks/Cole

Number Sig. figs. Comment on Zeros

2.12 3

4.500 4 Not placeholders. Significant.

0.002541 4 Placeholders (not significant).

0.00100 3 Only the last two are significant.

500 1, 2, 3 ? Ambiguous. May be placeholders or may be significant.

500. 3 Add a decimal point to show they are significant.

5.0 x 102 2 No ambiguity.


48 2008 Brooks/Cole

dp = 4

dp = 3

Addition and subtraction

Find the decimal places (dp) in each number.

answer dp = smallest input dp.

Add: 17.245 + 0.1001

17.3451

Rounds to: 17.345 (dp = 3)

Significant Figures in Calculations

49 2008 Brooks/Cole

dp = 2 dp = 4

Subtract 6.72 x 10-1 from 5.00 x 101

Use equal powers of 10:

5.00 x 101 0.0672 x 101

4.9328 x 101

Rounds to: 4.93 x 101 dp = 2


50 2008 Brooks/Cole

sig. fig. = 4

sig. fig. = 5

Multiplication and Division

Answer sig. fig = smallest input sig. fig.

17.245

x 0.1001

1.7262245

Rounds to: 1.726 sig. fig. = 4

Multiply 2.346, 12.1 and 500.99

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

= 14,221.402734


51 2008 Brooks/Cole

Round 37.663147 to 3 significant figures.

Examine the 1st non-significant digit. If it:

> 5, round up.

< 5, round down.

= 5, check the 2nd non-significant digit.

round up if absent or odd; round down if even.

last retained

digit

1st non-

significant digit

Rounds up to 37.7 2nd non-

significant

digit

Rules for Rounding

52 2008 Brooks/Cole

Round the following numbers to 3 sig. figs.

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

Rules for Rounding

53 2008 Brooks/Cole

Significant figures?

99.12444 6.321 27.5256

= 92.80344

27.5256 = 3.37153195571

= 3.3715 (5 sig. figs.)

dp = 5 dp = 3

Answer dp = 3.

92.803 is the significant result.

(5 sig. figs).

6 sig. figs.

Rules for Rounding

54 2008 Brooks/Cole

To avoid rounding errors

Carry additional digits through a calculation.

Use the correct number of places in the final answer.

Note

Exact conversion factors:

(100 cm / 1 m) or (2H / 1 H2O)

Have an infinite number of sig. figs.

Rules for Rounding

55

End of Lecture Unit No. 1

56

Outline

1. The Study Matter

2. States of Matter

3. Chemical and Physical Properties

4. Chemical and Physical Changes

5. Classification of Matter

6. Measurements

General Chemistry 1 LECTURE UNIT No. 1

Introduction to Chemistry

Engr. Edgie Estopace

School of Chemical Engineering and Chemistry

CHM11-3Lecture Unit #1

Documents

specific gravity density

liquids density

vm density

states of matter change

specific gravity sp2

chemical composition

processes matter

classification of matter