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General Chemistry Review

Chapter 1 - Matter and Measurement

I. Types of Matter

a. Matter is anything that has mass and occupies space.

b. Matter exists commonly in three phases:

i. Solids – rigid shape and fixed volume.

ii. Liquids – fixed volume and not rigid in shape.

iii. Gases – has neither fixed volume or rigid shape

- gases take on the volume and shape of its

container.

c. Matter is classified into two categories:

i. Pure substances

a. fixed composition and unique set of properties

b. pure substances are either elements or compounds.

ii. Mixtures

a. composed of two or more substances

b. are classified as either homogeneous or

heterogeneous.

MATTER

Mixtures Pure Substances

Homogeneous (Solutions)

Heterogeneous (Coarse)

Elements Compounds

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iii. Elements

a) A type of matter that cannot be broken down into two

or more pure substances.

b) 112 known elements (115, if the three newly

discovered elements are confirmed1), of which 91

occur naturally.

iv. Compounds

a) A pure substance that contains more than one

element and has fixed composition.

b) Example: Water

-Water is a compound; composed from the

elements hydrogen and oxygen.

c) Example: Table Salt (Sodium Chloride)

1) The properties of compounds are very different

from the properties of the elements from which

they are made.

2) Sodium chloride is formed from the elements

sodium (Na), a highly reactive metal and chlorine

(Cl), a greenish-yellow poisonous gas.

3) Sodium chloride, a crystalline solid, exhibits its

own set of properties, unique from those of either

sodium metal or chlorine gas.

1 See Supplement 1 – New Elements Pop In

NaCl

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v. Mixtures

a. Contain two or more substances combined in such

a way that each substance retains its chemical

identity.

b. Example: Copper sulfate and sand.

a) When copper sulfate and sand are shaken

together, the two substances do not react with

one another.

b) Each substance retains their original properties.

c. Homogeneous Mixtures

a) Homogenous Mixtures are uniform, meaning

the composition is the same throughout.

b) Another name for a homogeneous mixture is

solution.

1) The solvent is usually the substance present

in the largest amount; solute is the substance

is lesser amount.

2) Solvents are usually liquids.

3) Solutes may be solids, liquids, or gases.

c) Examples of homogeneous solutions include

1) Soda water solution of carbon dioxide in

water.

2) Seawater, complex mixture of solids, one of

which is sodium chloride.

3) Metal alloys, such as

- brass (copper and zinc)

- sterling silver (92.5% silver and 7.5% other

metals, usually copper)

- steel (iron ~98%, ~2% carbon, or other

metal such as manganese or tungsten)

- rose gold (gold, copper)

- bronze (copper and usually tin)

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d. Heterogeneous Mixtures

a) Nonuniform mixtures; composition varies

throughout.

b) Examples

1) Most Rocks, such as granite (feldspar, quartz

and quartz)

2) Salad, trail mix, milk (before homogenized)

c) Heterogeneous mixtures can be separated by

physical means

1) Filtration – used to separate a solid-liquid

mixture.

2) Distillation – separates liquids (which are

vaporized, and re-condensed) from solids

that are left behind in the distillation flask.

3) Chromatography

- can be used to separate many types of

mixtures.

- uses the difference in solubility and/or

extent of adsorption on a solid surface.

- an analogy which is sometimes useful is to

suppose a mixture of bees and wasps

passing over a flower bed. The bees would

be more attracted to the flowers than the

wasps, and would become separated from

them.

4) Gas-liquid Chromatography

- emissions tests on auto exhaust are done

by by GLC

- pollutants show up as peaks on a graph.

- a computer can analyze the area under

the peaks, translating into concentrations

- Drug and alcohol testing is typically

done by GLC.

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II. Measurements

a. Scientific measurements are expressed in the metric

system.

Metric Prefixes

Meaning Prefix Abbrev. Factor Prefix Abbrev. 1,000,000 X’s larger mega- M 1000 X’s smaller milli- m

1,000 X’s larger kilo- k 1,000,000 X’s smaller micro- µ 10 X’s smaller deci- d

!

109 X’s smaller nano- n

100 X’s smaller centi- c

!

1012

X’s smaller pico- p

b. Uncertainty in measurement

i. Every time that you make a measurement, there is

some human error (uncertainty) inherent in the

measurement.

ii. The proper way to record a measurement is to

“estimate” one decimal place beyond the scale of the

instrument.

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iii. Temperature

a) Temperature Units

- Celsius (oC)

a. Uses the normal freezing point (0 oC) and

the normal boiling point (100 oC) to set

this temperature scale.

- Fahrenheit (oF)

a. Water freezes at 32 oF and boils at 100 oF

on this scale.

b. The following formula can be used to

convert from oF to oC.

!

ToF

=1.8ToC

+ 32o

- Kelvin (K)

a. Defined as 1/273.15 of the difference

between the lowest attainable

temperature (0K) and the triple point of

water (0.01 oC)

b. No degree sign is used for Kelvin.

c. The following formula can be used to

convert from oC to K.

!

TK

= ToC

+ 273.15

c) Example: Conversion of Units/Temperature

Express normal body temperature, 98.60 oF, in oC and K.

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III. Significant Figures a. The following rules determine how many significant figures are in a

number:

i. Non-zero digits are always significant.

- Hopefully, this rule seems rather obvious. If you measure something and the device you use (ruler, thermometer, triple-beam balance, etc.) returns a number to you, then you have made a measurement decision and that ACT of measuring gives significance to that particular numeral (or digit) in the overall value you obtain.

- How many significant figures are there in the following _____26.38 _____ 7.94 _____ 0.00980 _____ 28.09

ii. Any zeros between two significant digits are significant.

- Suppose you had a number like 406.

- By the first rule, the 4 and the 6 are significant.

- However, to make a measurement decision on the 4 (in the hundred's place) and the 6 (in the unit's place), you HAD to have made a decision on the ten's place.

iii. The leftmost nonzero digit is the first or most significant

figure. For example, in the number 0.02340, the first significant figure is the 2.

- Leading zeros (0.02340) are never significant.

iv. If there is a decimal point, the rightmost digit is the last significant figure.

- For example, in 0.02340 the first two zeros from the left are not significant but the zero after the 4 is significant.

v. If there is no decimal point explicitly shown, the rightmost non-

zero digit is the least significant figure.

- For example, in 3400 the 4 is the least significant figure since neither zero is significant in this case.

- A decimal point can be placed an the end of a whole number that ends in zero to override this rule.

THE 3 RULES OF ZEROS

1. Leading Zeros (0.0001)

NEVER significant

2. MIDDLE ZEROS (209)

ALWAYS significant

3. Trailing Zeros

SOMETIMES significant

+ No decimal point, not

significant

Ex. 250 (2 sig figs)

+ with a decimal point, is

significant

Ex. 250. (2 sig figs)

250.000 (6 sig figs)

0.0900 (3 sig figs)

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- The value 3400. , when written with the decimal point, indicates to the reader that the measurement is certain to 4 significant figures, rather than 2.

b. Managing Significant Figures in Calculations

i. MULTIPLICATION or DIVISION

1. Keep the same number of significant figures as the factor with the least number of significant figures.

2. Example: Report the following calculation to the proper

number of significant figures

1.2 x 4.56 = 5.472 ii. ADDITION or SUBTRACTION

1. Keep the same number of decimal places as the factor with the least amount.

2. Example: Report the following calculation to the proper

number of significant figures

1.234 + 5.67 = 6.904

iii. Combined Calculations 1. In calculations involving addition/subtraction and

multiplication/division, significant figure guidelines must be applied before and after each calculation involving addition or subtraction.

2. example: (3.2 x 4 x 0.035 / 7) + (12 x 0.5) = �0.06 + 6 = 6

c. Logarithms i. There are specific rules that are followed when determining the

number of significant figures resulting from logarithm calculations, which will be discussed later in this class.

d. Focus on these rules and learn them well. They will be used extensively throughout the remainder of this course.

i. Exam and homework questions are only considered to be completely correct if they are rounded to the proper number of sig figs.

Round at the end

of your calculation

unless there is a

need to change

rounding rules.

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IV. Converting Units - Dimensional Analysis

a. A technique for converting units in chemistry that involves

“cross canceling” like units.

b. Example: Convert 2.5 atm of pressure into mmHg.

c. Metric System Conversions

Examples

Convert 23 cm to kilometers.

Convert 0.00357 kg to grams.

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V. Properties of Substances

a. Every pure substance has its own unique set of properties

that distinguish it from all other substances.

b. Intensive/Extensive Properties

i. Intensive properties (color) are independent of amount.

ii. Extensive properties (mass) depend upon the amount.

Extensive Intensive

c. Chemical/Physical Properties

i. Chemical properties are observed when a substance

takes place in a chemical reaction.

Examples of chemical properties:

ii. Physical properties are observed without changing the

identity of a substance.

Examples of physical properties:

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d. Density

i. Density Formula

!

density =mass

volume or

!

D =m

V

ii. Example

To determine the density of ethyl alcohol, a student pipets a 5.00 mL

sample into an empty flask weighing 15.246 g. He finds that the mass

of the flask + ethyl alcohol = 19.171 g. Calculate the density of ethyl

alcohol.

iii. Density can be treated as a conversion factor to relate

mass and volume.

iv. Example

Mercury has a density of 13.6 g/mL. Calculate the mass of 2.6 mL of

mercury.

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e. Solubility

i. The process by which a solute dissolves in a solvent is

a physical process.

ii. The extent to which a substance dissolves is commonly

expressed as the number of grams of dissolved

substance per 100 g of water (H2O) at a given

temperature.

iii. At 20 oC, about 32 g of potassium nitrate dissolves in

100 g of water.

iv. Temperature is known to effect solubility. At greater

temperatures, more solute can be dissolved.

v. At 100 oC, about 246 g of potassium nitrate dissolves

in 100 g of water.

vi. Example

Taking the solubility of potassium nitrate, KNO3, to be 246 g per 100 g

of water at 100 oC per 100 g of H2O; and 32g per 100 g of water at

20 oC. Calculate

(a) the mass of water required to dissolve 100 g of KNO3 at 100 oC

(b) The amount of KNO3 that remains in solution when the mixture is

cooled to 20 oC.