CHAPTER 1

Chemistry: The Study of Change

 

CHEMISTRY

UNITS

Table 1.2 SI Base Units

Base Quantity Name of Unit Symbol

Length meter m

Mass kilogram kg

Time second s

Temperature kelvin K

Amount of substance mole mol

Prefix Symbol Meaning

Tera- T 1012

Giga- G 109

Mega- M 106

Kilo- k 103

Hecto- h 102

Deca- da 101

Deci- d 10-1

Centi- c 10-2

Milli- m 10-3

Micro- 10-6

Nano- n 10-9

Pico- p 10-12

PREFIXES

The number of atoms in 12 g of carbon:

602,200,000,000,000,000,000,000

6.022 x 1023

The mass of a single carbon atom in grams:

0.0000000000000000000000199

1.99 x 10-23

N is a number between 1 and 10

n is a positive or negative integer

SCIENTIFIC NOTATION

568.762

n > 0

568.762 = __________

move decimal left

0.00000772

n < 0

0.00000772 = ___________

move decimal right

SCIENTIFIC NOTATION

SIGNIFICANT FIGURES

Once you start the counting don’t stop!

SIGNIFICANT FIGURES

Rule 1: Every nonzero digit in a measurement is significant.Examples: 24.7 0.22 569Rule 2:Zeros appearing between nonzero digits are significant.Examples: 7003 60.8 0.502Rule 3A ZERO is NOT significant when it is a placeholder. A placeholder is

used to show the location of the decimal point.Examples: .00099 5280 700

SIG FIGS CONT.

Rule 4:Zeros at the end of a number and to the right of a decimal

point are always significant.Examples: 86.0 46.00 1.010Rule 5:When a number is counted or defined within a system of

measurement, there is an infinite amount of significant digits.

Examples: 11 students 100 cm = 1 m

24 mL

3001 g

0.0320 m3

6.4 x 104 molecules

560 kg

COUNT THE SIG FIGS:

The answer cannot have more digits to the right of the decimalpoint than any of the original numbers.

89.3321.1+

90.432 round off to ________

one significant figure after decimal point

3.70-2.91330.7867

two significant figures after decimal point

round off to ________

SIG FIGS – Addition and Subtraction

The number of significant figures in the result is set by the original number that has the smallest number of significant figures

4.51 x 3.6666 = 16.536366 = _______

3 sig figs round to3 sig figs

6.8 ÷ 112.04 = 0.0606926

2 sig figs round to2 sig figs

= _______

SIG FIGS – Multiplication and Division

1. Determine which unit conversion factor(s) are needed

2. Carry units through calculation

3. If all units cancel except for the desired unit(s), then the problem was solved correctly.

1 L = 1000 mL

How many mL are in 1.63 L?

1L

1000 mL1.63 L x = 1630 mL

1L1000 mL

1.63 L x = 0.001630L2

mL

FACTOR LABEL METHOD

The speed of sound in air is about 343 m/s. What is this speed in miles per hour?

1 mi = 1609 m

1 min = 60 s 1 hour = 60 min

meters to miles

seconds to hours

Important things to consider when solving problems and performing

experiments….

Volume – SI derived unit for volume is cubic meter (m3)

1 L = 1000 mL = 1000 cm3 = 1 dm3

Matter - anything that occupies space and has mass.

mass – measure of the quantity of matter

SI unit of mass is the kilogram (kg)

weight – force that gravity exerts on an object

weight = c x mass

on earth, c = 1.0

on moon, c ~ 0.1

A 1 kg bar will weigh

1 kg on earth

0.1 kg on moon

Density – SI derived unit for density is kg/m3

1 g/cm3 = 1 g/mL = 1000 kg/m3

density = mass

volume D = mV

A piece of platinum metal with a density of 21.5 g/cm3 has a volume of 4.49 cm3. What is its mass?

For Water: 1g/mL

K = 0C + 273.15

0F = x 0C + 3295

273 K = 0 0C 373 K = 100 0C

32 0F = 0 0C 212 0F = 100 0C

Convert 172.9 0F to degrees Celsius.

0F = x 0C + 3295

0F – 32 = x 0C95

x (0F – 32) = 0C95

0C = x (0F – 32)95

0C = x (172.9 – 32) = 78.395

Graphing

Distance vs. Time

Dep

ende

nt V

aria

ble

Independent Variable

Line of Best Fit

Accuracy – how close a measurement is to the true value

Precision – how close a set of measurements are to each other

Percent Error:

Percents and Percent Error

Measured Value - Accepted Value

Accepted ValueX 100

Example: The mass of a compound measured in a lab was 25.0 grams. The accepted value for this compound is 24.5 grams. Calculate the percent error.

25.0 g- 24.5 g24.5 g

X 100 = 2.04 %

Scientific Method

SCIENTIFIC METHOD

logical approach to solving problems

ObservationProblemHypothesisExperiment

Data Analysis

Conclusion

Observations

Qualitative: quality, non-numeric terms

Quantitative: quantity, numerical description

MATTER: ANYTHING THAT OCCUPIES SPACE AND HAS MASS

MATTER

State of Matter

Volume Shape Density Compressibility Motion of Molecules

Gas

Liquid

Solid

Plasma

Three States of Matter

solid liquid gas

Phase Changes

phase diagram: summarizes the conditions at which a substance exists as a solid, liquid, or gas.

Phase Diagram of Water

Phase Diagram Points

________________: above this point a substance becomes a supercritical fluid critical temperature (Tc): temperature above which the gas

cannot be made to liquefy, no matter how great the applied pressure.

critical pressure (Pc): minimum pressure that must be applied to bring about liquefaction at the critical temperature.

________________: point at which all three phases coexist

1. Pure Substance: form of matter that has a definite composition and distinct properties

2. Mixture: combination of two or more substances in which the substances retain their own identities.

water, ammonia, sucrose, gold, oxygen

Classification of Matter

• mixed together physically• can usually be separated

1. Homogenous mixture – composition of the mixture is the same throughout.

2. Heterogeneous mixture – composition is not uniform throughout.

Examples:

1.4

Examples:

Types of Mixtures

Mixture Pictures

•solution: mixture that remains uniformly mixed•solute: •solvent:

•suspension: mixture where visible particles settle

•colloid: mixture where particles are unevenly distributed but do not separate, positive Tyndall Effect

Types of Mixtures

Hom/Het? Soln/Susp/Coll?

Fog ____________ ___________

Paint ____________ ___________

Syrup ____________ ___________

Physical means can be used to separate a mixture into its pure components.

magnet

1.4

distillation

Methods of Separation

Strainer

Filtration

Physical

Evaporation

Centrifuge

Distillation

PASTA/WATER SAND/IRON FILINGS

SALT /WATER

BLOOD FOOD COLORING/WATER

SAND/WATER

Element:

•all atoms are the same

•cannot be broken down by physical or chemical means

• 114 elements named on the Periodic Table• 83 elements occur naturally on Earth

gold, aluminum, lead, oxygen, carbon• many elements have been created by scientists

technetium, americium, seaborgium

Compound:

•2 or more elements combined

• Cannot be broken down by physical means

•Can be broken down by chemical means

•Appears different from original elements

•Fixed ratios in definite proportionsWater (H2O) Glucose (C6H12O6)

Ammonia (NH3)

PHYSICAL AND CHEMICAL CHANGES

Physical Properties and Physical Changes

physical property: characteristic that can be observed or measured without changing the identity of the substance.

• physical change: change in a substance that

does not involve a change in the identity of the substance.

•

Physical Properties

Intensive: INDEPENDENT of amount of matter present (sample size) Example:

Extensive: DEPENDENT on the amount of matter present (sample size) Example:

Chemical Properties and Chemical Changes

chemical property: a substance’s ability to undergo changes that transform it into different substances

Example:

chemical change: change in which one or more substances are converted into different substances

Example:

Evidence of a Chemical Change

1.

2.

3.

4.

5.

physical change does not alter the composition or identity of a substance.

chemical change alters the composition or identity of the substance(s) involved.

ice meltingsugar dissolving

in water

hydrogen gas burns in oxygen gas to form water

Physical or Chemical?

Remember: