unit 1- chemical foundations-1 gg

8/8/2019 Unit 1- Chemical Foundations-1 Gg

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Chapter 1- Matter and Change

Chapter 2- Measurements and Calculations



I. Matter- anything that takes up space and has

mass.

A. Properties and Changes in Matter

1. P

hysical property- an characteristicobserved w/out changing the identity of

the substance (melting pt, color, temp«)

a. Physical change- a change in a substancethat does not involve a change in the identity

(cutting, melting, boiling«)



2. Chemical property- relates to a

substance¶s ability to undergo changesthat transform it to a different substance.

a. Chemical change- a change in which one

or more substances are converted intodifferent substances.

Reactants products

Na+ + Cl- NaCl

yields



3. Extensive (Quantitative) Properties-

depend on the amount of matter .

a. Mass- the amount of matter an object contains.

b. Weight- the gravitational pull on an object.

c. Volume- the space something takes up.

4. Intensive (Qualitative) Properties- do not

depend on the amount of matter .

a. Density- the ratio of an objects

mass to its volume.

a. Melting/boiling point- the temp.

at which a substance melts/boils.



II. Classification of Matter

Matter

Can it be separated?

Mixtures Pure Substances

Is the composition

uniform?

Can it be decomposed by

ordinary chemical means?

Homogenous

Mixtures(air, sugar in water)

Heterogeneous

Mixtures(granite, blood)

Compounds(water, sugar)

Elements(oxygen, sodium)



II. Classification of Matter

A. Mixtures- consist of a physical blend of twoor more substances of variable

composition.

1. Types of Mixtures

a. Homogenous- the same throughout

(milk, saline solution«) b. Heterogenous- not the same throughout

(pizza, sand and water).



B. Pure Substances- consist of a fixed

composition.

1. Elements- the simplest form of matter

(Periodic Table).

a. Chemical symbols- shorthand for the

element on periodic table.



III. The Scientific Method- incorporates observations, hypotheses,

ex periments, theories and laws.

A. Stages in the Scientific Method

1. Observing/collecting data

2. Formulating hypotheses

3. Testing hypotheses

4. Collecting data

5. Theorizing



B. Collecting data

1. Observations- the noting and recording

of facts.

2. Inference- an

interpretation of an

observation.



IV. Units of Measurement

A. SI system- a standard system of scientificmeasurement.

1.

The 7 Fundamental SI

Units



2. Derived SI Units- produced by

multiplying or dividing SI units.

a. Volume- V = L x W x H

= cm x cm x cm

= cm3

1 cm3 = mL

b. Density- D = mass/ volume

= g/cm3



V. Using Scientific Measurements

A. Accuracy vs. Precision

1. Accuracy- how close a measurement

comes to the actual true value.

2. Precision- the reproducibility of the

measurement.



B. Percent Error- the accuracy of an average

ex perimental value compared to theaccepted value.

Percent Error = Valueex p.

± Valueacc.

x 100

Value acc.



VI. Developing Tools For Analysis

A. Making Tables and Graphs- graphs visuallyshow proportional relationships among

data.

1. Slope = y2 - y1

x2 - x1

= (y( x

= rise

run



B. Rules for Good Graphing

1. Give your graph a descriptive title.

2. Indent the axes from the edge of the

graph paper .

3. Label each axis and the units used.

4. Choose an appropriate scale.

5. Choose a convenient scale.

6. Locate points with a small circle around

them.

7. Draw a smooth curve or straight line to

represent the general tendency of the

data points.



VII. Significant Figures

A.Significant Figures- all thedigits in a measurement that can be

known accurately plus a last

digit that must be estimated.



1. Which digits are significant?

a. Every nonzero digit in a measurement.

b. Zeros appearing between nonzero digits.

c. Zeros at the end of a number and to the

right of a decimal point.

d. Zeros at the end of a measurement only if the number contains a decimal point.

e. All digits in scientific notation.

f . Zeros appearing in front of all nonzero

digits are NOT significant.



VIII. Significant Figures in Measurement

A.No measurement is e

xact!

1. It is important to state ex perimental

results with a number of significant

digits which give a reasonableimpression of the accuracy of the

measurement.

Consider an example of a measured density.

Suppose you measured the mass of a rock to be

9.3 grams and it¶s volume to be 3.4 cm3.

Density = mass = 9.3 g = 2.44736 g/cm3

volume 3.8 cm3

Mass measured

to 2 SF¶sVol. measured to

2 SFs

Not science but fiction!Maybe this doesn¶t look as impressive but it is a better answer

= 9.3 g = 2.4 g/cm3

3.8 cm3



B. Multiplying and dividing

1. The measurement with the fewest

significant figures determines the

number of significant figures in the

answer .

4.28 m x 9.2567 m = 3.96 m2

0.62 cm x 1.56 cm = 0.97 cm2

985.33 g / 65.2 mL = 15.1 g/mL

27.30 L2 / 2.73 L = 10.0 L or

1.00 x 101 L



C. Adding and subtracting

1. The measurement with the digit

having the lowest decimal value

determines the decimal value in the

answer .

12.1 g + 435.673 g = 447.8 g

62 m - 25.321m = 37 m

1.20 mL - .0021 mL = 1.20 mL



IX. Dimensional Analysis (Unit-Factor Method)

A. Dimensional Analysis- the converting of

unit into another without changing the

value of the original amount.



1. Conversion Factors - shows the

relationship between two measurements.

a. A ratio where the measurement on the

top is equal to the measurement on the

bottom.

100 cm = 1 & 1 m_ = 1

1 m 100 cm



2. Steps in Dimensional Analysis

(1)W

rite down what is given(2) Create an equality between what is

given and what you want to change to.

(3) Set up a conversion factor (the unit to

cancel goes on the bottom)(4) Do math. Multiply tops, divide by

bottoms.

HOW MANY CM IN 6.4 INCHES?

6.4 INCHES 2.54 CM =

1 INCH

Correct

for SFs!

16.3 CM

16.256 CM

unit 1- chemical foundations-1 gg

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