Lecture 1 - Nassau Community College

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Lecture 1Professor Hicks

General Chemistry (CHE131)

Scientific notation

scientists describe things very large/small

diameter earth = 12000000 meters

diameter atom = 0.00000000011 meters

more conveniently expressed in scientific notation

(without so many zeros) as

diameter earth = 1.2 x 107 meters

diameter atom = 1.1 x 10-10 meters

6

Scientific notation

diameter earth = 1.2 x 107 meters

number between

1 and 10

decimal places

abbreviated as

powers of 10

diameter earth = 12000000 meters

7 decimal

places

numbers larger than 10

have positive powers of 10

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Scientific notation

= 1.1 x 10-10 meters

number between

1 and 10

decimal places abbreviated

as powers of 10

diameter atom = 0.00000000011 meters

10 decimal

places

numbers smaller 1 have

negative exponents

1.21 Express these numbers in scientific notation:

(a) 0.000000027, (b) 356, (c) 0.096.

1.23 Convert these to nonscientific notation:

(a) 1.52 × 104, (b) 7.78 × 10−8.

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Qualitative vs Quantitative Data

• Data can be Qualitative- meaning it is descriptive such as

- “The solution turned green”

- “The test tube got hot”

or

• Data can be Quantitative- meaning it involves numbers

• The process of collecting quantitative data is called

measurement

- Quantitative data is often analyzed using graphs

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Accuracy vs Precision

• Precision is the consistency of a measurement made in different trials

• Accuracy is the agreement of a measure value with an accepted value

accurate

and precise

not accurate

but precise

not accurate

not precise

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13

Precision

small radius

good precision XX

XXX

more precise

precise

no matter how small the

circle gets (precise) if it

is not near center it is

not accurate

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Accuracy vs Precision

in measurements

• Making systematic errors is when you repeat a

mistake without realizing it

• Systematic errors can lead to results that are very

precise, but not accurate

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Systematic Errors

• Forgetting to zero the scale

• Forgetting to subtract the mass of a container

• Using a ruler with a worn end

• Repeatedly miscalculating

• Using the wrong chemical

• Misspelling a word over and over

if results are precise, but far from expected

values a systematic error is probably the cause

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Random Errors

• Errors that are not caused by choices made by

the scientist - like the luck of the draw

• Variation when devices are read by eye

• Random fluctuations such as the draft in the

room changing mass reading on a digital scale

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accurate

but not

precise

precise

but not

accurate

accurate

and

precise

minimal

random &

systematic

error

Experiment

minimal

random

error

minimal

systematic

error

What do the accuracy and

precision tell us about the

sorts off errors that occurred?

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Controlling errors

• Best way to control random

errors is to repeat trials and

average – this improves the

precision

• Best way to control systematic

errors is comparing your result

to accepted values

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19

(16)

scale marked in units of 1 cm – read to 0.1 cm

Significant figures• All measurements have a precision that describes how

much uncertainty is in the measurement

• Rules to describe precision called rules of significant figures

• Rule 1 – when making a measurement record values to 1

decimal place more than the scale is marked

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(16)

• Estimate decimal place

between 15 and 16

•15.?

Significant figures• Rule 1 – record values to 1 decimal place more than the

scale is marked

I say it looks like 15.4

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Significant figuresI say it looks like 15.4

some people may read it as 15.3 and some 15.5

the value you record is assumed to be within

one unit in last decimal place

15.4 means value is most likely from 15.3 to 15.5

or 15.4 + 0.1

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15.4

15.3 15.5

Rules of significant figures assume everyone’s judgment

can achieve a precision of 1 unit in last decimal place

radius of circle = 0.1 cm in this case

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Reading liquid levels

magnify

curved edge of liquid

called meniscus

5 ml

6 ml

I read 5.2 milliters at

bottom edge of meniscus

position eye at liquid level

to read liquid level

it can face up (water in plastic)

or down (water in glass)

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Number of significant figures

• the value 5.2 ml has 2 significant figures

• this is the number of digits that are considered reliable

• most likely in range 5.1-5.3 or 5.2 0.1

rules for counting number of significant figures

1) all non-zero digits are significant

2) zeros between other digits are significant

3) zeros to the left of all other digits are not significant (they are just placeholders)

4) zeros to the right of all other digits and the decimal place are significant

5) zeros to the right of other digits, but left of the decimal place are ambiguous

(they should be rewritten in scientific notation)

calculators know nothing about significant figures

these rules apply to numbers that were assigned by a

scientist based on the methods used to arrive at the value

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Rules for counting number of significant figures

1) all non-zero digits are significant

2) zeros between other digits are significant

3) zeros to the left of all other digits are not

significant

4) zeros to the right of all other digits and the

decimal place are significant

5) zeros to the right of other digits, but left of

the decimal place are ambiguous

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How many significant figures?

• 1614.1

• 22.0000

• 1001

• 0.000136

• 100

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4 zeros between other digits are significant

zeros to left of other digits NOT significant

ambiguous zeros right of other digits, but left of

decimal place are ambiguous

must be present as placeholders so we are

not sure if they are significant

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same number in

scientific notation

1.0 x 102 2

all the digits in scientific

notation are significant

zeros to the right of decimal

place are significant

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Multiplication/division

and significant figures

1.369 * 2.5 = 3.4225 this is what your calculator reads

the result does NOT have 5 s.f.

Rule for multiplication/division

final answer must be rounded to the least number of

significant figures that either of the factors had.

4 s.f. 2 s.f.

round to 2 s.f.

3.4

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100.00 cm

9 cm

+ /- 1 cm

8 cm

10 cm

Understanding the multiplication rule

+ /- 0.01 cm

100.01 cm99.99 cm

largest area 100.01 10 = 1000.1 cm2

smallest area 99.99 8 = 799.92 cm2

most likely area 100.00 9 = 900 cm2 + /- about 100 cm

which is the prediction of the

rules of significant figures

each side was measured to a different precision

largest possible area

smallest possible area

possible range is

about 800-1000 cm2

900 cm2

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Addition/subtraction

and significant figures

12.344 – 11.2 = 1.144 This is what your calculator reads

the result does NOT have 4 s.f.

Rule for addition/subtraction

Final answer must be rounded to the larger of the decimal places

that were significant in either number.

0.001 orthousandths

decimal

0.1 ortenths

decimal

round to 0.1 or tenths decimal place

1.1 finally count significant figures

2 s.f.

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K2 is the second

highest mountain

in world 28,252 ft

Understanding the

addition/subtraction rule

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Understanding the

addition/subtraction rule

K2 is the second

highest mountain

in world 28,251 + 1 ft

penny 0.00523 feet thick

max height K2 + penny = 28252 + 0.00524

min height K2 + penny = 28250 + 0.00522

range = 28250.00522 28252.00524

about the same as the

uncertainty in height K2

2825028252 ft

No! Rules of significant figures say the range is + 1

same as less precise quantity 28251 ft

Is it taller now (by rules of sig fig rules)?

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Addition/subtraction

and multiplication/division

7.65*(12.344 – 11.2)0.001 or

thousands

decimal

0.1 or

tenths

decimal

1.144 2 s.f

= 7.65 *(1.144)

do not round!!!!

keep track of # s.f

3 s.f. 2 s.f.

= 8.7516

round to 2 s.f.

8.8

mult/div ask how many SF

add/subt ask what decimal place

do each step and express your answer

in terms the next step requires

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Numbers with unlimited

significant figures

• Defined quantities

12 inches = 1 foot

60 seconds = 1 minute

• Quantities you can

count

7 donuts

6 pencils

#sig figs = infinity!

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1.27 What is the number of significant figures in each of these

measured quantities? (a) 4867 miles, (b) 56 mL, (c) 60,104

tons, (d) 2900 g.

1.29 Carry out these operations as if they were calculations of

experimental results, and express each answer in the correct

units and with the correct number of significant figures:

5.6792 m + 0.6 m + 4.33 m

3.70 g − 2.9133 g

4.51 cm × 3.6666 cm

(3 × 104 g + 6.827 g)/(0.043 cm3 − 0.021 cm3)

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1.35 Three students (X, Y, and Z) are assigned the task of

determining the mass of a sample of iron. Each student makes

three determinations with a balance. The results in grams are

X (61.5, 61.6, 61.4); Y (62.8, 62.2, 62.7); Z (61.9, 62.2, 62.1).

The actual mass of the iron is 62.0 g. Which student is the

least precise? Which student is the most accurate?

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SI units used in science

Some common units

are not SI units

• Centimeters

• Celsius degree

• Grams – convenient unit in

student lab also popular with

street level drug dealers

• SI unit of kilogram used in

chemical industry and

popular with druglords

metric but

not SI units

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Prefix multipliers

• Words used instead of 10something

Example: kilo means 103 so

5.7 kilometers = 5.7 × 103 meters

On the exam it will be given to you like this

Example: Express 250 kilometers in scientific

notation without prefixes.

250 kilometers

1) move 2 decimal places left 102

2.5 x 102 kilometers

2) replace prefix with number

2.5 x 102 x103 meters

3) simplify exponents

2.5 x 105 meters

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• Example: Express 0.0000537 seconds in

microseconds.

0.0000537 seconds

need to express answer as something x 10-6 seconds

1) insert the factor 106 × 10-6 = 1

0.0000537 106 10-6 seconds

2) combine these

decimal moves

right 6 places

3) rewrite 10-6 as micro

53.7 microseconds

here’s the trick!

Units “Math”

• Units are included in calculations you can do the same kind of operations on units as you can with numbers

cm × cm = cm2

cm + cm = cm

cm ÷ cm = 1

• Using units as a guide to problem solving is called Dimensional Analysis

Conversion factors and units

• Converting one unit into another often involves ratios called Conversion Factors

• Conversion factors come from Equivalence Statements

1 inch = 2.54 cm can give two factors

divide both sides by 1 inch or divide both sides by 2.54 cm

in1

cm54.21 = 1

cm54.2

in0.1= multiplying by either factor is

equivalent to multiplying by 1

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Using conversion factors

select conversion factors so that the old unit

cancels and is replaced by the new desired unit

unit newunit old

unit newunit old =

conversion factor

equivalent to multiplying by 1

Example: Convert 0.299 pounds to grams

lb 1.0

g 453.59

1 lb = 453.59 grams

g 136lb 1.0

g 453.59lb 0.299 =

look at equations you have

involving pounds and grams

quantity in

old unitconversion factor

cancels old unitquantity in

new unit

g 453.59

lb 1.0

gives 2 conversion factors

pick the conversion factor that

will cancel the old unit and has

new unit on top

Example: Convert 1.76 yd to centimeters

1 yd = 0.9141 m

1 m = 100 cm

cm 161m 1

cm 001

yd 1

m 0.9141yd .761 =

yd m cm

look at equations you have involving yards and cm

yards can be converted to meters then

meters converted to centimeters

quantity in

old unitconversion factors

quantity in

new unit

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Example: Convert 125 decimeters into meters.

1 meter = 1 meter

1 decimeter = 1 x 10-1 meter

gives 2 possible conversion factors

1 decimeter

1 x 10-1 meters

1 x 10-1 meters

1 decimeter

125 decimeters

old unit

x1 x 10-1 meters

1 decimeter

conversion factor

= 12.5 meters

new units

you will have this table for the exam

Example: Convert 235 nanometers into

micrometers.

1 meter = 1 meter

1 nanometer = 1 x 10-9 meter

1 meter = 1 meter

1 micrometer = 1 x 10-6 meter

gives 2 possible conversion factors gives 2 possible conversion factors

1 nanometer

1 x 10-9 meters

1 x 10-9 meters

1 nanometer

1 micrometer

1 x 10-6 meters

1 x 10-6 meters

1 micrometer

235 nanometers

old unit

x1 x 10-9 meters

1 nanometer

1 micrometer

1 x 10-6 metersx

conversion factors

= 0.235 micrometers

tip 2: When converting between units with prefixes

use two conversion factors: one to go to the un-prefixed

unit and one to go to the new prefixed unit.

e.g. in this case nanometersmetersmicrometers

Derived Units

• Units built up from base units are called derived units

• Can be multiplied or divided

- “per” means division of units

3) Pressure unit “pounds per square inch”pounds

inch2

1) All formulas for area involve two length dimensions multiplied meter * meter

meter2 or m2

2) Units of velocity “miles per hour” miles

hour

area rectangle =l*w area circle = r2 area triangle = ½b*h

Derived unit

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Example: Convert 2.11 yard3 to meters3

2.11 yard3

x1 meter

1.0936 yard

3

2.11 yard3x

13 meter3

1.09363 yard3

2.11 yard3x

1 meter3

1.3079 yard3= 1.61 meter3

or 1.61 m3

tip 3: conversion factors for units of area

or volume can be derived by writing down

the conversion factor for the base unit of

length and squaring or cubing it

Understanding conversion factors for area/volume

1 meter

1 meter

area = 1 decimeter squared

or 1 dm2

1 decimeter

1 decimeter

even though a decimeter is

1/10 the length of a meter

it would require 100 square

decimeters to cover 1 m2

1 m2 = 100 dm2

1 decimeter = 10-1 meter

1 dm2 = 10-2 m2

or 100 dm2 = 1 m2

area = 1 meter squared

or 1 m2

(deci = 10-1)

How could we have figured

that out without a diagram?

Converting base units within a derived unit

the SI unit of energy is the derived unit called the Joule.

1 Joule = 1kg*m2

sec2

• any part of a derived unit can be converted

as if it was a base unit alone

x1000 g

1 kg=

kg*m2

sec2

0.251 251 g*m2

sec2

units of

Joules

conversion

factor for

kg to g

new derived

units has grams

instead of kg

Example: Convert 0.251 joules into units of g*m2/sec2 .

it is as if we just

converted

kg into grams

kg g

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Carry out these conversions:

(a) 12.6 decimeters to m

(b) (b) 252.4 mg to kilograms.


(a) 142 lb to milligrams

(b) 18.3 nm3 to cubic meters.


(a) A 5.0-ft person weighs 136 lb. Express this person's height

in meters and weight in kilograms. (1 lb = 453.6 g; 1 m =

3.28 ft.)

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(b) The current speed limit in some states in the United States

is 55 miles per hour. What is the speed limit in kilometers

per minute?


(c) The speed of light is 3.0 × 108 m/s. How many miles does

light travel in 1 minute?


(d) Lead is a toxic substance. The “normal” lead content in

human blood is about 0.40 part per million (that is, 0.40 g

of lead per million grams of blood). A value of 0.80 part

per million (ppm) is considered to be dangerous. How

many grams of lead are contained in 6.0 × 103 g of blood

(the amount in an average adult) if the lead content is 0.62

ppm?

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The density of ammonia gas under certain conditions is 0.625

g/L. Calculate its density in g/µm3.

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Temperature Scales

• Temperature reflects the random motion of matter at the microscopic level

• At higher temperatures motion is faster

• Most matter expands as it gets warmer and shrinks as it cools

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• Based upon the expansion of matter as it is warmed

• Calibrated using reference points like boiling water, or ice water as fixed points

Thermometers

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4

add alcohol

food coloring

ice bath

mark as 0 degrees

boiling water bath

alcohol expands

as it warms

mark as 100 degrees

make 100 uniform marks

between 0 and 100 degrees

each is 1 degree on

the Celsius scale

empty

glass

tube

Celsius Temperature Scale

Anders Celsius

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Other temperature scales

• Fahrenheit scale

similar to Celsius scale 3 points used

1) ice bath (32 degrees Fahrenheit)

2) ice bath with a compound added (0 degrees Fahrenheit)

3) Daniel Fahrenheit's armpit (98.6 degrees Fahrenheit)

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Kelvin Scale

• Based on similar principles to Celsius /Fahrenheit using gases not liquids

• Step sizes same as Celsius scale

• 0 degrees Kelvin was originally defined as the temperature at which gases would shrink to zero volume

William Thompson(Lord Kelvin)

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Converting between temperature scales

F = 9/5C + 32

• Converts a temperature from Celsius to Fahrenheit.

• Example: Convert 37 C to the Fahrenheit temperature.

F = 9/5*37 +32= 66.6 + 32= 99 F

this equation is on your units conversion page

37 C is about human body temperature

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Converting Fahrenheit to Celsius

• Rearrange the equation used to convert Celsius to Fahrenheit

F = 9/5C + 32

F - 32 = 1.8 C

(F – 32)/1.8 = C

C = (F-32)/1.8

this equation is on your

units conversion page

hedgehog

Example: Convert the body temperature of a hibernating hedgehog 26.8 F to degrees Celsius.

C = (26.8 -32)/1.8 = -5.2/1.8 = -2.88 C

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Converting between temperature scales

K = C + 273.15

• Converts a temperature from Celsius to Kelvin.

• Example: Convert 25 C (room temperature) to the Kelvin scale.

K = C + 273.15= 25 + 273.15= 298.15 K= 298 K

this equation is on your units conversion page

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Intensive vs Extensive Properties

• Extensive properties of matter depend on the amount of matter considered.

• Intensive properties do not depend on the amount of matter considered

Extensive Intensivecost of a bag candy cost per pound candy? temperaturemass density

(mass per 1 unit volume)

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Density

• Intensive property of matter (can be measured on any sample size)

• D = mass/volume

• Mass is its related extensive property

• Determines if an object will sink or float

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Measuring Density

Density = mass/volume

• Mass and volume must both be measured

• Any sample size OK because density is an intensive property

- Both mass and volume must be measured on the same sample to determine density

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Measuring Density

• Volume

- Liquids can be directly measured in glassware

- Solids with geometric shapes can have their individual length(s) measured and volume calculated

• Irregular shaped solids can be measured by water displacement

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Measuring Density

• What about irregular solids that will dissolve in water, like a chunk of salt?

• Sand could be used instead of a liquid in a liquid displacement-like experiment

• Instead of water a different liquid could be used that the substance would not dissolve in, like oil.

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A gold sphere has a mass of 1.36 × 102 g, and its

volume is 7.039 cm3. Calculate the density of gold.

Mercury is the only metal that is a liquid at room

temperature. Its density is 13.6 g/mL. How many grams

of mercury will occupy a volume of 16.8 mL?

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This procedure was carried out to determine the volume

of a flask. The flask was weighed dry and then filled with

water. If the masses of the empty flask and the filled

flask were 96.12 g and 197.18 g, respectively, and the

density of water is 0.9976 g/cm3, calculate the volume of

the flask in cubic centimeters.

An object weighing 116.3 g is placed in a graduated

cylinder containing 236.01 mL of water. The volume

of water now reads 260.56 mL. From these data

calculate the density of the object.

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• Elements – building blocks of all matter

each box on the periodic table is one element

• Compounds – built up of

elements in definite proportion

• Mixtures – elements or compounds combined

in any proportion

Classifications of Matter

metals

non-metals

elements in same groups (columns) undergo similar chemical reactions

memorize the organization of metals and non-metals on the PT

this group (column)

called the alkali metals

this group (column)

called the noble gases

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in any proportion


Compounds

ionic compounds molecular compounds

metal + non-metal 2 non-metals

Properties: high melting points

conduct electricity when melted

or dissolved.

Properties: do not conductor electricity

low melting points.

Hydrogen – metal or non-metal?

• Appears in group I Li, Na, K, etc all metals

• Also is grouped with group VII F, Cl, Br etc all non-metals

• Some compounds with H are ionic

• Some compounds with H are molecular

• H acts like a nonmetal when it reacts with metals andnon-metals

• H + metal ionic compound

• H + non-metal molecular compound

• If a compound contains H we treat H as a non-metal and classify the compound accordingly

3

Chemical Formula

• Inventory of how many atoms of each element are present in the substance

NO2

element symbol

for nitrogen

element symbol

for oxygen

number of nitrogen

atoms = 1

number of oxygen

atoms = 2

element more to left on PT is written first

the atom is the smallest amount of an element

Molecular compounds

form molecules

• Molecules are individual particles with definite size and number of atoms

• Chemical formulas of molecular compounds do not always have smallest whole number ratios of atoms

NO2

N2O4

different compounds

both are molecules

N and O = non-metals

form molecular compounds

Molecular elements

• Some elements exist as molecules

• They are called molecular elements

S8

element symbol

for sulfur

number of sulfur

atoms = 8

molecular

sulfur

They are molecular because they have a definite number of atoms

They are elements because theya are made up of only one type of atom

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Many molecular elements are diatomic molecules

(two atom)

F2 Cl2 Br2 I2 At2

F F

O O H HN N

Cl Cl

hydrogen

H2

nitrogen

N2

oxygen

O2

Br Br I I At At

molecular

fluorine

molecular

chlorineetc.

DiatomicsF2 Cl2 Br2 I2 At2

F FF FF F

O OO OO O H HH HH HN NN NN N

Cl ClCl ClCl Cl

hydrogen

H2

nitrogen

N2

oxygen

O2

Br BrBr BrBr Br I II II I At AtAt AtAt At

halogens

hydrogen sometimes

grouped with F, Cl, Br etc

elements found as

diatomic molecules

He Ne Ar Kr Xe Rn

Noble Gases are Atomic Elements

• Elements found in nature as atoms

• They do not form compounds

• All are gases at room temperature

They are atomic because they are found as single atoms

They are elements only one type of atom is present

5

Ionic compounds form lattices

1.8 nm

1.0 mm 2.0 cm

• Geometric arrangement of ions

• No definite size

sodium chloride = NaCl

Waters of hydration

• Ionic compounds form

crystal lattices

• Water molecules fit in

spaces in the lattice

heat

CuSO4 5H2O CuSO4 + 5H2O

(dot means it is packed into the crystal)

hydrated form anhydrous form






in any proportion


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homogenous heterogeneous

Mixtures

Every sample of a

homogenous mixture

has same composition

every sample of a heterogeneous mixture

does not have the same composition

mixtures differ from compounds because

they do not have to be made up with a

definite ratio of their components

(can have variable composition)

Serial dilution

• Homogenous mixtures uniform throughout

1.00 gram

of blue compound

in 1000 mL of water

1 mL

+ 999 mL

water

1 mL

+ 999 mL

water

1 mL

contains

0.000000000100 grams

(1.00 nanogram)

of blue compound

0.00100 grams

of blue compound

in 1000 mL

compared to a scale

that is limited to

0.001 g ( 1 mg)

0.00000100 grams

of blue compound

in 1000 mL

Physical properties/changes

• Physical property = a property that can be checked without

changing the chemical identity

- Example matter in solid, liquid, or gaseous state

• Physical change = a change that occurs without changing

the chemical identity

- Example heating to cause solid liquid gas

heroin solid

heat heat

heroin gas

heroin liquid

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Solid, Liquid, Gas

• Solids - atoms/molecules in contact with

each other in fixed positions

• Liquid - atoms/molecules in contact but can

slide past each other

• Gas - atoms/molecules widely separated

Physical vs Chemical change

Physical changes

• identity of substance(s) does not change!

– solid liquid gas

– grind, dissolve, spread out,

concentrate, make a mixture, or

separate a mixture

Chemical changes

• identity of substance does change!

• A Chemical Reaction occurs

Chemical reactions convert:

– elements to compounds

– compounds to elements

– compounds to different compounds

Chemical Reaction

reactants products

Solutions and Like Dissolves Like

• Homogenous mixtures that are liquid are called Solutions

• Substances are classified based on whether they dissolve to form solutions in oil or water

• Water and substances that dissolve in it are said to be Polar

• Oil and substances that dissolve in it are said to be Non-Polar

• Solubility can often be estimated based on the “Like Dissolves Like Rule”- polar substances dissolve in polar liquids

- non-polar substances dissolve in non-polar liquids

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Separation of Mixtures by Evaporation

wet sand

heterogenous mixture

water + sand

heatclouds = water as a gas

dry sand

sun

• Based on differences in physical properties

• Water evaporates lower temperature than sand

Separation of

Mixtures by

Filtration

• Based on solubility

• Soluble materials form homogenous

mixture (solution) and pass through filter

• Insoluble materials cannot pass through

filter

caffeine

soluble in

hot water

Separation of Mixtures by Distillation

• Based on different physical properties

• Substance with lower boiling point evaporates

more

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Separation of Mixtures by Paper

Chromatography

• Method used to separate mixtures

• Used for mixtures with many components

Steps

1) mixture added to a solid material = stationary phase

2) “washed out” or eluted with a liquid solvent = mobile phase

• components elute (move up paper) at different rates

Paper chromatography

• Stationery phase = paper

• Mixture spotted toward bottom of paper

• Solvent drawn into the paper by capillary

action

• As solvent moves up paper mixture

separates

• Components more soluble in mobile

phase elute faster

Classification of matter

more than one

type of atom?

every sample has the

same composition?

Can the composition

be varied?

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1.8 Does each of these describe a physical change

or a chemical change?

(a) The helium gas inside a balloon tends to leak out

after a few hours.

(b) A flash-light beam slowly gets dimmer and finally

goes out.

(c) Frozen orange juice is reconstituted by adding

water to it.

(d) The growth of plants depends on the sun's

energy in a process called photosynthesis.

(e) A spoonful of table salt dissolves in a bowl of

soup.

1.9 Which of these properties are intensive and

which are extensive?

(a) length, (b) volume, (c) temperature, (d) mass.

1.10 Which of these properties are intensive and

which are extensive?

(a) area (b) color (c) density.

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1.11 Classify each of these substances as an

element or a compound:

(a) Hydrogen

(b) Water

(c) Gold

(d) sugar

1.12 Classify each of these as an element or a

compound:

(a) Sodium chloride (table salt)

(b) Helium

(c) Alcohol

(d) Platinum

1.41 Which of these describe physical and which

describe chemical properties?

(a) Iron has a tendency to rust.

(b) Rainwater in industrialized regions tends to be

acidic.

(c) Hemoglobin molecules have a red color.

(d) When a glass of water is left out in the sun, the

water gradually disappears.

(e) Carbon dioxide in air is converted to more

complex molecules by plants during

photosynthesis.

1



Coulomb’s Law

• Some types of matter can acquire a

property called charge when rubbed

together

• There are two types

positive (+) negative (-)

+-

Coulomb’s Law

• Oppositely charged objects are attracted

• Like charges are repelled

+ -

+

+ -

-

2

Cathode ray tube

+-

screen

e- e-

+ electrons attracted

to positive platehigh voltage

©

Discovery of the electron

• Electrons can be made to flow out of all kinds of matter

• They are therefore building blocks of all matter

• They are negatively charged and matter is not charged

there must be a positive part also

©

Rutherford’s Gold Foil Experiment• particles are positive particles emitted from radioactive substances

• They are also seen flowing towards the negative plate when helium is

placed in the cathode ray tube before evacuating

• Most of the particles go straight through!

The atom must have lots of empty space

Those that are deflected must have hit something so most of the matter of

the atom is compacted into small regions we call the nucleus

3

Positive part of matter

• Positive part of matter called protons

• Protons forms lumps called the nucleus

(not easily pulled apart)

• Nucleus + electrons = atom

the smallest amount of any element

+ = proton

Organization of the nucleus

nucleus

+ +

+

+

+

N

N

N

N

N

N

+ = proton

N = neutron

~10-15 m• Composed of protons and

particles called neutrons

• Protons and electrons have the

same charge but opposite sign

Charge proton = +1.6 x10-19 C

Charge electron = -1.6 x10-19 C

• Neutrons uncharged

• Nuclear radii range 2-15 fm

Organization of the atom

• Typical size 10-10 m

• Most of volume

occupied by electrons

• eEectrically neutral

#protons = # electrons

electrons

nucleus

10-15 m

10-10 m+ charge = - charge

4

Each element has a different

number of protons

# protons

# protons

#protons increases

Periodic table is a list of elements with

increasing number of protons

Atomic mass units (amu)

+ = proton = 1.0 amu

N = neutron = 1.0 amu

N

NN

NN

N N

N N

N

N

NN

NN

N N

N N

N

N

NN

NN

N N

N N

N

N

NN

NN

N N

N N

N

N

NN

NN

N N

N N

N

N

NN

NN

N N

N N

N

N

NN

NN

N N

N N

NN

NN

NN

N N

N N

N

N

NN

NN

N N

N N

N

N

NN

NN

N N

N N

NN

NN

NN

N N

N N

N

N

NNN

N

N N

N N

N

6.02 x 1023 amu

= 1.0 g

• Based on mass protons, neutrons

• Proton and neutron each have

about the same mass = 1.0 amu

• 6.02 x1023 amu = 1.0 gram

• 6.02 x 1023 called

Avogadro's number (NA)

e- = electron = 1/1836 amu

(5.4 x 10 -4 amu)

5

• Same # protons

different # neutron

• Isotopes are versions of an

element with different # neutrons

Isotopes

boron-10

+ +

+

+

+

N

N

N

N

N

boron-11

+ +

+

+

+

N

N

N

N

N

Nboth are boron (5 protons)

different # neutrons

different isotopes of boron

= an isotope

Isotope notations

C13

6 element symbol

nuclear charge (Z)

# protons

mass number (A)

= # protons

+ # neutrons

pronounced “carbon 13”

How many neutrons?

mass number - # protons= # neutrons

= A-Z

=13 – 6 = 7 neutrons

can also be written “carbon-13”

or “C-13”

More on isotopes

two isotopes of boron

boron-10

+ +

+

+

+

N

N

N

N

N

boron-11

+ +

+

+

+

N

N

N

N

N

N

• Elements in nature are mixtures

of different isotopes

• Isotopes have different masses

• Nuclei of some isotopes are unstable

- break down in nuclear reactions

7

Weighted averages

• a way to calculate different weight for different types

of contribution

Example: Grade is calculated as weighted averages

grade = 15% Quizzes + 60% Exams + 25% Lab

To write it as a math equation convert % /100

grade = 15 x Quizzes + 60 x Exams + 25 x Lab

100 100 100

percentages are the weights

8

Weighted averages

How does it compare to a regular average?


100 100 100

grade = Quizzes + Exams + Lab

3


100 100 100

which you can rewrite as

regular average = all same weight

Atomic mass and periodic table

he atomic mass is

weighted average

of isotope masses

# protons

natural Sn is a mixture of:

Sn120

50

Sn116

50

Sn118

50

Sn122

50

Sn119

50Sn117

50

Sn115

50Sn114

50

Sn124

50

Sn115

50Sn114

50

Isotopes and weighted averages

• Each isotope has a different mass in amu

• The percentage of it found in nature is the weight

atomic mass = % isotope x mass isotope

+ % isotope x mass isotope

+ % isotope x mass isotope

… etc

weighted average

of isotope masses

(aka atomic weight)

9

Example. Carbon found in nature is mostly two

isotopes, carbon-12, and carbon-13. They have

abundances of 98.9% and 1.1%. What is the

atomic mass of natural carbon in amu?

atomic mass = 98.9 x 12.00 amu

100

+ 1.1 x 13.00 amu

100

carbon-12 has mass of 12.00 amu

carbon-13 has mass of 13.00 amu

12.01 amu

C12

6

C13

6

The atomic masses of B-10 and B-11

are 10.0129 amu and 11.0093 amu,

respectively. Calculate the natural

abundances of these two isotopes. The

average atomic mass of boron is 10.81

amu.

1


General Chemistry (CHE 131)

Main Group Elements

Transition Metals

Lanthanides/Actinides

alkali

metals

alkaline

earth metalshalogens

noble

gases

Periodic Properties

(elements in same group are similar)

react with water

form hydrogen gas

react with acids to

form hydrogen gas

Cl, Br, I used

disinfectants

F is too dangerous

At is radioactive

do not form

compounds

memorize the names

of these groups

2

Ions

• Atoms are electrically neutral

positive charge = negative charge

#protons = # electrons

• Atoms can gain or lose electrons to form

monatomic ions

• Gaining electrons forms anions (negative ions )

• Losing electrons forms cations (positive ions)

I-K+

lost 2

electrons

Ca2+

N3-

gained 3

electrons

Notating chargesNa+

lost 1

electron K+

lost 1

electron

Al3+

lost 3

electrons

F-

gained 1

electron

gained 2

electrons

O2-

charge written as number of electrons lost (+) or

gained (-) in superscript

Examples

Na+ K+ Ca2+ Al3+ F- O2- N3-

ions are visualized

as spheres

Hydrogen atom has 1 proton in nucleus

Hydrogen atom has 1 electron

charge hydrogen atom = 0

If 1 electron is removed hydrogen ion is formed

charge hydrogen ion = +1 (+1.6 x 10-19 C)

written as H+

Example of an ion H+

nucleus

+

e-

Hydrogen atom

+

electrically

neutral

3

Trend in type I (fixed-charge) ions

(a main group periodic property)

• Same columns (groups) = same charge

• Metals form positive ions

• Non-metals form negative ions

metals

non-metals

metals in this area can have

more than one ion

called variable charge ions

+1

+2+3 -3

-1-2

memorize these trends

also a few of these

metals can also form

variable charge ions

Ionic compounds

• Electrically neutral

• Charge positive ions = charge negative ions

NaCl = 1 Na+ and 1 Cl-

Li2S = 2 Li+ and 1 S2-

CaF2 = 1 Ca2+ and 2 F-

-1+1

-2+1 +1

2+ -1-1

only one compound for

each combination elements

Na2Cl NaCl2

4

Molecular compounds are more

numerous than ionic compounds• Binary compounds are one with only two

elements

• Knowing what elements are in a molecular compound is not enough to determine its chemical formula

N2O4

NO Viagra

activates NO

Rocket fuel

NO2 Smog

N2O

Laughing gas

aka whippets, hippie crack

How do we tell if a compound is ionic or molecular

before we know its chemical formula?

-anion

+cation

• Most molecular compounds are

non-electrolytes

They do not separate into ions

when dissolved or melted

do not conduct electricity

• Ionic compounds conduct electricity when

dissolved in water or melted

separate into ions when

dissolved in water or melted

• Mobile ions conduct electricity

• Ionic compounds are said to be

strong electrolytes

water is a molecular compound and a non-electrolyte

(remember H is an exception it acts like a non-metal)

tap water only conducts b/c

ions are dissolved in it

Formula Unit• Chemical formula of ionic compound called Formula Unit

• Smallest whole number ratio of ions that will be electrically neutral

SrO = 1 Sr2+ and 1 O2-

64 Na+ and 64 Cl-

1019 Na+ and 1019 Cl-

1022 Na+ and 1022 Cl-

CaCl2 = 1 Ca2+ and 2 Cl- Li2O = 2 Li+ and 1 O2-

NaCl = 1 Na+ and 1 Cl-

Na64Cl64

Na Cl1019 1019

Na Cl1022 1022

smallest whole number ratio is same for different size lattices

5

Crossing over rule(how to figure out the formula unit of an ionic compound)

Ca Ca2+ N3- N

3-

2+

2+

2+

3-

total +

charge

3 x 2

total -

charge

2 x 3

=

3 2

=

3 2=

periodic

table

periodic

table

What is the formula unit of the ionic compound

made from calcium and nitrogen?

why?b/c ionic compounds

electrically neutral

Ca3N2

Crossing over rule

Mg Mg2+ O2- O2 2

periodic

table

periodic

table


made from magnesium and oxygen?

why not

Mg2O2?

b/c the formula unit has smallest whole number ratio

of ions that will be electrically neutral (1 to 1 smaller 2 to 2)

MgO

Polyatomic ions

• Groups of atoms bonded

together that have a charge

Hg2 2+

PO43-NH4

+CN-

OH-

SO42-

• Acts as a single ion

H COO O

atoms

+ e-

H

CO O

O

-1

HCO3-

HCO3-

6

Polyatomic ions form ionic

compounds

Hg2 2+ PO4

3-

NH4+

CN-

OH-

SO42-HCO3

-

• Positive polyatomic ions can substitute a metal ion

• Negative polyatomic ions can substitute a non-metal ion

you will be

given

this chart

( )2

Crossing over rule (polyatomic ions)

Ca Ca2+ PO43-

phosphate3

periodic

table

table of

polyatomic

ions


made from calcium and phosphate ions?

why use

( ) ? Ca3(PO4)2

(a polyatomic ion)

PO43- PO4

3-Ca2+ Ca2+ Ca2+

2 PO43- ions

not PO423-

Counting atoms in chemical formula

Ca3(PO4)2

PO43- PO4

3-Ca2+ Ca2+ Ca2+

# Ca 2+ ions = 3

# PO43- ions = 2

1 P + 4 O 1 P + 4 O

total = 2 P and 8 O

3 Ca

7

2.28 Give an example of each of the following: (a) a

monatomic cation, (b) a monatomic anion, (c) a

polyatomic cation, (d) a polyatomic anion.

Lecture 1 - Nassau Community College

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