Introduction: Matter &

Measurement

CHEMISTRY The Central Science

Cpt. 1 and 2

What is Chemistry?

The central science

The study of the matter, its composition, properties, and the changes it undergoes.

The Molecular Perspective of Chemistry

• Matter is the physical material of the universe. It

has mass and occupies space.

• On the microscopic level, matter consists of atoms

and molecules.

• Atoms combine to form molecules.

• Molecules may consist of the same type of atoms

or different types of atoms.

The Molecular Perspective of Chemistry

Why Study Chemistry?

• Chemistry is central to our understanding of other

sciences.

• Chemistry is also encountered in everyday life.

States of Matter

• Matter can be a gas, a liquid, or a solid.

• Gases have no fixed shape or volume.

• Gases can be compressed to form liquids.

• Liquids have no shape, but they do have a defined volume. Liquids are fluid.

• Solids are rigid and have a definite shape and volume.

The Three States of Matter

• Elements and compounds are referred to as

Pure Substances. They contain one type of

particle.

• Mixtures are not considered pure

substances. They contain more than one

type of particle.

Pure

Substances

and Mixtures

Elements

• If a pure substance cannot be decomposed into

something else, then the substance is an element.

• There are 118 elements known.

• Each element is given a unique chemical symbol

(one or two letters).

• Elements are building blocks of matter.

Elements

Elements

• Chemical symbols with one letter have that letter

capitalized (e.g., H, B, C, N, etc.)

• Chemical symbols with two letters have only the

first letter capitalized (e.g., He, Be).

C Cu Na U

Compounds

• Most elements interact to form compounds.

• Law of Constant Composition (or Law of

Definite Proportions):

–The composition of a pure compound is always

the same.

Substances

• If water is decomposed, then there will always be

twice as much hydrogen gas formed as oxygen

gas.

• Pure substances that can be decomposed are

compounds.

Mixtures

• Heterogeneous mixtures are not uniform

throughout.

• Homogeneous mixtures are uniform throughout.

• Homogeneous mixtures are called solutions.

Physical and Chemical Properties

• Intensive physical properties do not depend on

how much of the substance is present. – Examples: density, temperature, color, melting point.

• Extensive physical properties depend on the

amount of substance present. – Examples: mass, volume, pressure.

Chemical Properties describe the chemical

behavior of a substance

-Example: Flammable, corrodes metal, basic

Physical and Chemical Changes

• When a substance undergoes a physical change, its

physical appearance changes. – Ice melts: a solid is converted into a liquid.

• Physical changes do not result in a change of

composition.

• When a substance undergoes a chemical change, it

changes its composition: – When pure hydrogen and pure oxygen react completely, they

form pure water.

Chemical Changes

Separation of Mixtures

• Mixtures can be separated if their physical

properties are different.

• Solids can be separated from liquids by means of

filtration.

• Homogeneous liquid mixtures can be separated by

distillation.

Distillation

Separation of Mixtures

• Chromatography can be used to separate mixtures

that have different polarity.

• Chromatography can be used to separate the

different colors of inks in a pen.

Gas Chromatography

The Periodic Table

Periodic Table Basics

• Most elements are metals – left of zig-zag

• Non metals are on the right of zig-zag

• Borderline elements called metalloids

• Horizontal rows are called periods

• Vertical columns are called groups or

families. (similar properties – more later)

Metals and Nonmetals and

Their Ions

• Metals

– Good conductors of heat and electricity.

– Malleable and ductile.

– Moderate to high melting points.

• Nonmetals

– Nonconductors of heat and electricity.

– Brittle solids.

– Some are gases at room temperature.

C-2 Measures and metrics

SI Units

• There are two types of units: – fundamental (or base) units;

– derived units.

• There are 7 base units in the SI system.

• Powers of ten are used for convenience with

smaller or larger units in the SI system.

SI Units

SI Units

• Note the SI unit for length is the meter (m)

whereas the SI unit for mass is the kilogram (kg). – 1 kg weighs 2.2046 lb.

Temperature

There are three temperature scales:

• 1. Kelvin Scale – Used in science.

– Same temperature increment as Celsius scale.

– Lowest temperature possible (absolute zero) is zero Kelvin.

– Absolute zero: 0 K = -273 oC.

Temperature

• 2. Celsius Scale – Also used in science.

– Water freezes at 0 oC and boils at 100 oC.

– To convert: K = oC + 273.

• 3. Fahrenheit Scale – Not generally used in science.

– Water freezes at 32 oF and boils at 212 oF.

– To convert:

32-F

9

5C 32C

5

9F

Temperature

Scientific Notation

•Numbers written in scientific notation include a

numeral with one digit before the decimal point,

multiplied by some power of ten (6.022 x 1023)

•In scientific notation, all digits are significant.

•You should be able to convert from non-scientific

notation to scientific and vice-versa (on calc. as

well).

Examples

1. Convert to scientific notation:

a. 450 000 000

b. 0.000 000 047

2. Convert to standard notation:

a. 7.09 x 10-6

b. 3.39 x 105

Derived Units

• Derived units are obtained from the 7 base SI

units.

• Example:

m/s

seconds

meters

timeof units

distance of units velocityof Units

Volume

• The units for volume are

given by (units of

length)3. – SI unit for volume is 1 m3.

• Common: 1 mL=1 cm3.

• Other volume units: 1 L = 1 dm3 = 1000 cm3 = 1000mL.

Volume

Density

• Used to characterize substances.

• Defined as mass divided by volume:

• Units: g/cm3 or g/mL

• Originally based on mass (density was defined as

the mass/vol. of 1.00 g of pure water).

volume

massDensity

Examples:

1. An object with a mass of 17.95 g occupies a

volume of 11.8 mL. What is its density?

2. A sample with a density of 3.75 g/cm3 occupies a

volume of 10.44 cm3. What is the mass of the

sample?

3. A graduated cylinder is filled with 15.0 cm3 of

water. An object with a mass of 29.66 g causes

the total volume to increase to 23.4 mL. What is

the density of the sample?

1. 1.52g/mL 39.2 g 3.53g/ml

Dimensional Analysis

• Method of calculation utilizing a knowledge of

units. Conversion factors are simple ratios

(fractions):

unitgiven

unit desired factor Conversion aswritten

Using a Conversion Factor

• Example: convert length in meters to length in

centimeters:

• 3.25 meters x 100 cm = 325 cm

1 m

Using Two or More Conversion Factors

• Example to convert length in meters to length in

inches:

• 3.00 meters x 100 cm x 1 inch = 118 in

1 m 2.54 cm

Problem solving

• In dimensional analysis always ask three

questions:

• What data are we given?

• What quantity do we need?

• What conversion factors are available to take us

from what we are given to what we need?

Uncertainty in Measurement

• All scientific measures are subject to error.

• These errors are reflected in the number of figures

reported for the measurement.

Precision and Accuracy

• Measurements that are close to the “correct” value

are accurate.

• Measurements that are close to each other are

precise.

Precision and Accuracy

Significant Figures

• The number of digits reported in a measurement

reflect the accuracy of the measurement and the

precision of the measuring device.

• All the figures known with certainty plus one extra

figure (estimated digit) are called significant

figures.

• In any calculation, the results are reported to the

fewest significant figures (for multiplication and

division) or fewest decimal places (addition and

subtraction).

Significant Figures Rules:

1. Non-zero numbers are always significant.

2. Zeros between digits are always significant.

3. Initial zeros are not significant. (Example:

0.0003 has only 1 sf.)

4. Zeros at the end of the number are significant IF

there is a decimal point.(24.0 cm has 3 sf)

Examples: How many significant figures are in

each of the following?

1. 52300 m

2. 0.000487 kg

3. 29.0400 s

4. 507 people

5. 230,050 cm

6. 45.600 A

Reporting Uncertainties and Errors

in Measurements

• All instruments have limitations. When

recording a measurement, you should report

the measurement to the limits of the

instrument.

• The uncertainty (or tolerance) of an

instrument may be found on the instrument

( given as ±) , or you may be told what it is.

• Percent Error – used to compare data

(observed values) to accepted values in a

meaningful way

• │ Obs – Acc│ x 100 = % error

Acc

0

2

4

6

8

10

1 2 3 4 5 6

Series2

Direct Inverse

Proportions Proportions

0

5

10

15

20

25

30

1 2 3 4 5 6 7

Series2