CHEM210 X-rays old-time remedies reading Please read the following article:

http://www.bbc.co.uk/news/science-environment-2206…

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3: Find 3 concepts from within the article and relate them to 3 concepts within CHEM 210 we have discussed in class and cite 3 textbook references using the chapter and page number.

2: Find 2 concepts from within the article that you want to know more about (i.e. muddy points, have questions about, did not quite understand).

1: Write an exam question with the answer about 1 concept discussed from within the article. The exam question must be well thought out and appropriate to the subject matter. Chapter 1. Chemistry:

Matter on the Atomic Scale

1.1

1.2

1.3

1.4

What is Chemistry?

Classification of Matter

Units and Measurement

Unit Conversions

Chapter 1 introduces the fundamental components of matter, the

different types of structures they can make when they join

together, and the types of changes they undergo.

Interactive Figure 1.1.1 – Understand the Scale of Science

• Macroscopic scale

– Matter that can be seen with the naked eye and can be held

• Atomic scale

– Nanoscale and molecular scale

– Processes cannot be seen

2

1.2

The Properties of Matter

All properties of matter are either extensive or intensive.

The measured value of an extensive property depends on the

amount of matter.

Mass is an extensive property.

The value of an intensive property are independent on the

amount of matter.

Density and temperature are intensive properties.

3

Some Chemical Properties of the Elements

Physical Properties: Characteristics that do not

involve a change in a sample’s chemical makeup.

Chemical Properties: Characteristics that do involve a

change in a sample’s chemical makeup.

Temperature Conversions

• Fahrenheit to Celsius

9

T(°F) = T(°C) + 32

5

9

T(°F) = (82.63 °C) + 32 = 180.73 F

5

• Celsius to Kelvin

1K

× T°C + 273.15°C

1°C

1K

= 298.2 K

25.0°C + 273.15°C ×

1°C

T(K) =

5

Density Units

• Density: Physical property that relates the

mass of a substance to its volume

mass

Density =

volume

• Densities of solids and liquids are reported:

– As grams per milliliter (g/mL) or grams per cubic

centimeter (g/cm3) 1 mL = 1 cm3

– At a standard temperature, 25°C, close to room

temperature

6

Density

Example

Determine the mass of 3274 mL of mercury, with a

d = 13.55 g/mL.

m = V x d = 3274 mL x 13.55 g = 4.436 x 104 g

1 mL

Interactive Figure 1.2.5 – Classify Matter

Mixture combo of two or

more substances retain their

distinct identities

Substance definite composition

& distinct properties

Element cannot be separated

into simpler substances by

chemical means

Compound atoms of two or

more elements in fixed ratio.

Homogeneous

uniform throughout

Heterogeneous not

uniform throughout

8

1.3

1.3 Scientific Measurement

SI Base Units

Measured property

Unit name

Abbreviation

Mass

Kilogram

kg

Length

Meter

m

Time

Second

s

Temperature

Kelvin

K

Amount

Mole

mol

Electric current

Ampere

A

Luminous intensity

Candela

cd

9

Units of Measure

Metric prefixes are combined with SI units when reporting physical quantities

Prefix

Abbreviation

Meaning

Example

Giga

G

109 (billion)

1 gigahertz = 1 X 109 Hz

Mega

M

106 (million)

1megaton = 1 X 106 ton

Kilo

k

103 (thousand)

1 kilometer (km) = 1 X 103 m

Deci

d

10-1 (tenth)

1 decimeter (dc) = 1 X 10-1 m

Centi

c

10-2 (onehundredth)

1 centimeter (cm) = 1 X 10-2 m

Milli

m

10-3 (one

thousandth)

1 millimeter (MM)= 1 X 10-3 m

Micro

m

10-6 (one

millionth)

1 micrometer (mm)= 1 X 10-6 m

Nano

n

10-9 (one

billionth)

1 nanometer (nm) 1 X 10-9 m

Pico

p

10-12

1 picometer (pm) =1 X 10-12 m

Femto

f

10-15

1 femtometer (fm)= 1 X 10-15 m

10

Scientific Notation

• Numbers are expressed in a format that

conveys the order of magnitude

– General form: N × 10x

• Converting standard notation to scientific

notation

– Count the number of times the decimal point is

moved to the right or left

3285 ft = 3.285 103 ft

0.00215kg = 2.15 103 kg

1

Precision and Accuracy

• Precision: How close the values in a set of

measurements are to one another

• Accuracy: How close a measurement or a set of

measurements is to a real value

12

1.4

Uncertainty in Measurement

Chemistry makes use of two types of numbers: exact and

inexact.

Exact numbers include numbers with defined values with

infinite significant figures, such as

2.54 in the definition 1 inch (in) = 2.54 cm

1000 in the definition 1 kg = 1000 g

12 in the definition 1 dozen = 12 objects.

Numbers measured by any method other than counting are

inexact.

13

1.4

Uncertainty in Measurement

An inexact number must be

reported in such a way as to

indicate the uncertainty in its value.

Significant figures are the

meaningful digits in a reported

number.

14

Counting Significant Figures

Rule

Example

Number of significant

figures

1256

4

All zeros between non-zero

numbers are significant-captive

zeros

1056007

7

Leading zeros are NEVER

significant (zeros to the left of your

first non-zero number)

0.000345

3

0.00046909

5

1780

3

770.0

4

0.08040

4

All non-zero numbers are

significant

Trailing zeros are significant ONLY

if a decimal point is part of the

number (zeros to the right of your

last non-zero number)

1

1.3

Uncertainty in Measurement

Significant Figures

To avoid ambiguity in such cases, it is best to express such

numbers using scientific notation [Appendix 1].

1.3 × 102 two significant figures

1.30 × 102 three significant figures

50

Rounding Numbers

• Find the last digit that is to be kept

• Check the number immediately to the right

– If that number is less than 5 leave the last digit

alone

2.543 round down to 2.54

– If that number is 5 or greater increase the

previous digit by one

2.546 round up to 2.55

1

Round to Two Significant Figures

Find the

last digit

that is to be

kept

Check

number

to the

right

Is digit to

the right

less than

5?

Is digit to

right 5 or

greater?

Rounded

number

1056007

1056007

No

Yes

1100000

0.000345

0.000345

No

Yes

0.00035

1740

1740

Yes

No

1700

1

1.3

Uncertainty in Measurement

Calculations with Measured Numbers

Addition and Subtraction, the answer cannot have more

digits to the right of the decimal point than the original

number with the smallest number of digits to the right of

the decimal point

If the leftmost digit to be dropped is less than 5, round down.

If the leftmost digit to be dropped is equal to or greater than

5, round up.

55

1.3

Uncertainty in Measurement

Calculations with Measured Numbers

Multiplication and Division, the number of significant figures

in the final product or quotient is determined by the original

number that has the smallest number of significant figures.

2 SF

4 SF

2 SF

56

1.3

Uncertainty in Measurement

Addition and Subtraction with Multiplication and Division,

the number of significant figures in the final answer is based

on the Order of Operations and the corresponding rule

needed.

2.54 cm x (147.9 inch – 145.900 inch) = ??

inch

Do () math 1st:

147.9 inch

– 145.900 inch 2 SF

2.000 inch ~ 2.0 inch

3 SF

2 SF

3 SF

2 SF

Then multiply: 2.54 cm x (2.0 inch) = 5.08 cm ~ 5.1 cm

inch

56

1.3

Using Units and Solving Problems

A conversion factor is a fraction in which the same quantity is

expressed one way in the numerator and another way in the

denominator.

Both forms of this conversion factor are equal to 1, we can

multiply a quantity by either form without changing the value

of that quantity.

72

Dimensional Analysis

Useful US units to Metric Unit Conversions

• 1 in = 2.54 cm

• 1 lb = 453.6 g

• 1 oz (mass) = 28.3459231 g (mass)

• 1 oz (fluid) = 29.5735 mL (fluid)

FOLLOW THE UNITS!

2

Example of Conversion Factors

• Conversion factors – Used to make the

conversion between units

– The resulting quantity is equivalent to the original

quantity, it differs only by the units

Unit (1) ×conversion factor = unit (2)

– Example – Metric to metric equalities

1 m = 100 cm

1m

102 cm

or

2

10 cm

1m

• Use of exact conversion factors will not affect the

significant figures in a calculation

24

1.6

Using Units and Solving Problems

The use of conversion factors in problem solving is called

dimensional analysis or the factor-label method.

A proportionality (or conversion) factor was used.

known units x desired units = desired units

known units

FOLLOW THE UNITS!

Make sure UNITS Cancel Out before doing the math!

74

1.6

Using Units and Solving Problems

The speed of light in a vacuum is 2.998 x 108 m/s, what is

the speed in cm/min?

2.998 x 108 m x 100 cm x 60 s = 1.799 x 1012 cm

s

1m

1 min

min

FOLLOW THE UNITS!

Make sure UNITS Cancel Out before

doing the math!

74

Exercise: How Many Picometers Are in 25.4 nm?

Given: 25.4 nm

Find: pm

Roadmap:

Factors:

Solve:

nm m pm

1 m 1012 pm

;

9

10 nm 1 m

1m

1012 pm

25.4nm × 9

×

= 2.54 ×10 4 pm

10 nm

1m

FOLLOW THE UNITS!

Make sure UNITS Cancel Out before doing

the math!

Check: We would expect a greater number of pm and that is what we have

27

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