Chapter 1: Matter and Chapter 1: Matter and MeasurementsMeasurements

What is Chemistry?

Biology vs.

Chemistry vs.

Physics

3

What is Chemistry?

Biology

Physics

Chemistry

The study of living organisms

The study of forces & motion

The study of matter and its reactions and properties

What is Chemistry?

Chemistry is the study of CHANGES in the “stuff” around us.

(We formally define “stuff” as matter!)

What is Chemistry?

Think, pair, share:

What are all the chemicals you use in your daily life?

Review: Scientific Methods

1. Hypothesis

–Suggested solution to a problemSuggested solution to a problem

2. Experiment

–A controlled method of testing a hypothesisA controlled method of testing a hypothesis

3. Data

–Organized observationsOrganized observations

a. Data is always reproduciblereproducible..

Review: Scientific Methods

4. Scientific Law

• Statement which summarizes results of many Statement which summarizes results of many observations and experimentobservations and experiment

a. Scientific laws explain WHAT is observed.

» Example of a scientific law:

5. Scientific Theory

• Explanation that supports a hypothesis and which Explanation that supports a hypothesis and which has been supported with repeated testinghas been supported with repeated testing

b. Scientific theories explain WHY something is observed.

» Example of a scientific theory:

Review: Scientific Methods

6. Steps of the Scientific Method—Review

a.

b.

c.

d.

e.

f.

9

What is Chemistry?

Biology

Physics

Chemistry

The study of living organisms

The study of forces & motion

The study of matter and its reactions and properties

What is Chemistry?

Chemistry is the study of CHANGES in the “stuff” around us.

(We formally define “stuff” as matter!)

11

Matter

12

Matter

13

Elements

Type of matter that cannot be broken down into simpler, stable substances and is made of only one type of atom

14

Compounds

A pure substance that contains two or more elements whose atoms are chemically bonded

15

Compounds

Fixed compositions A given compound contains the same elements

in the same percent by mass

16

Compounds

The properties of a compound are VERY DIFFERENT from the properties of the elements they contain

Ex.) Sodium Chloride (NaCl) vs. Sodium & Chlorine

Sodium: http://www.youtube.com/watch?v=RAFcZo8dTcU

http://www.youtube.com/watch?v=92Mfric7JUc

17

Electrolysis

18

Mixtures

A blend of two or more kinds of matter, each of which retains its own identity and properties

Homogeneous

Heterogeneous

19

Homogeneous MixturesComposition is the same throughout the

mixtureExamples: salt water, soda water, brass

A.k.a. a solutionSolute in a solvent (salt dissolved in water)

20

Heterogeneous Mixtures

Non-uniform; composition varies throughout the mixture

21

Separating Mixtures

Filtration

22

Separating MixturesDistillation

23

Separating Mixtures

Chromatography

Scientific Measurements

Chemistry is a quantitative science. –This means that experiments and

calculations almost always involve measured values.

Scientific measurements are expressed in the SI (metric) system.–This is a decimal-based system in which all

of the units of a particular quantity are related to each other by factors of ten.

SI SystemSI System• Definition: modernized version of metric

system; uses decimals

• All units derived from base units; larger and smaller quantities use prefixes with base unit

• Must memorize prefixes from nano- (10-9) to tera- (1012)

Prefixes (see handout & Ebook)

• You will need to memorize all of the prefixes (factors, names and abbreviations from

109 (giga-) to 10-9 (nano)!

• One example of a memory device:

 

INSTRUMENTS & UNITSINSTRUMENTS & UNITS

Use Use SI unitsSI units — based on the — based on the metric systemmetric system

Length Length

MassMass

TimeTime

TemperatureTemperature

Meter, mMeter, m

Kilogram, kgKilogram, kg

Seconds, sSeconds, s

Celsius degrees, ˚CCelsius degrees, ˚CKelvins, KKelvins, K

LengthLength

The standard unit of length in the metric system is the METER

which is a little larger than a YARD.

USING THE PREFIXES WITH LENGTH:

cm – often used in lab

km –

Gm –

LengthLengthBase unit: METERMETER

Conversions:

1 km=1000 m 1 cm = 10 -2 m

1 Gm = 106 m

Units of Units of LengthLength

• 1 kilometer (km) = 1000 meters (m)1 kilometer (km) = 1000 meters (m)

• 1010-2-2 meter (m) = 1 centimeter (cm) meter (m) = 1 centimeter (cm)

• 101022 meter (m) = 1 hectometer (Hm) meter (m) = 1 hectometer (Hm)

• 1 nanometer (nm) = 1.0 x 101 nanometer (nm) = 1.0 x 10-9-9 meter meter

O—H distance =O—H distance =9.58 x 109.58 x 10-11 -11 mm9.58 x 109.58 x 10-9 -9 cmcm0.0958 nm0.0958 nm

O—H distance =O—H distance =9.58 x 109.58 x 10-11 -11 mm9.58 x 109.58 x 10-9 -9 cmcm0.0958 nm0.0958 nm

VolumeVolume

• THE COMMON UNITS OF VOLUME IN CHEMISTRY ARE:

the liter (milliliter) and cubic centimeter (cm3)

• THE COMMON INSTRUMENTS FOR MEASURING VOLUME IN CHEMISTRY ARE: graduated cylinder & buret

• Note that 1 cm3 = 1 mL (We will use this exact conversion factor throughout the year, so you will need to memorize it!)

MassMass

• THE COMMON UNIT OF MASS IN CHEMISTRY IS :

the gram (g) — often used in lab

• Mass IS A MEASURE OF THE AMOUNT OF MATTER IN AN OBJECT;

• Weight IS A MEASURE OF THE GRAVITATIONAL FORCE ACTING ON THE OBJECT. CHEMISTS OFTEN USE THESE TERMS INTERCHANGEABLY.

1000 g= 1 kg 1 Mg = 10 6 g

Temperature ScalesTemperature Scales• FahrenheitFahrenheit

• CelsiusCelsius

• KelvinKelvin

Anders Celsius1701-1744

Lord Kelvin(William Thomson)1824-1907

TEMPERATURE IS THE FACTOR THAT DETERMINES the direction of heat flow.

Temperature ScalesTemperature Scales

Notice that 1 kelvin degree = 1 degree Celsius1 kelvin degree = 1 degree Celsius

Boiling point Boiling point of waterof water

Freezing point Freezing point of waterof water

CelsiusCelsius

100 ˚C100 ˚C

0 ˚C0 ˚C

100˚C100˚C

KelvinKelvin

373 K373 K

273 K273 K

100 K100 K

FahrenheitFahrenheit

32 ˚F32 ˚F

212 ˚F212 ˚F

180˚F180˚F

TemperatuTemperature Scalesre Scales

100 100 ooFF38 38 ooCC311 K311 K

oF oC K

SI SystemEnglish Units (inches, feet, degrees F, etc.) are NEVER used to take measurements in the lab!

Calculations Calculations Using Using TemperatureTemperature

Fahrenheit/Celsius Fahrenheit/Celsius

T (F) = 1.8 t (˚C) + 32T (F) = 1.8 t (˚C) + 32

Calculations Calculations Using Using TemperatureTemperature

Some calculations are in kelvins Some calculations are in kelvins

(especially important for Ch 5!!)(especially important for Ch 5!!)

T (K) = t (˚C) + 273.15 (273)T (K) = t (˚C) + 273.15 (273)

•Body temp = 37 ˚C + 273 = 310 KBody temp = 37 ˚C + 273 = 310 K

•Liquid nitrogen = -196 ˚C + 273 = 77 KLiquid nitrogen = -196 ˚C + 273 = 77 K

ProblemProblem

Example 1L.1 A baby has a temperature of Example 1L.1 A baby has a temperature of 39.839.8ooC. Express this temperature in C. Express this temperature in ooF and K.F and K.

SI System: Base Units

ESTABLISHMENT OF THE INTERNATIONAL SYSTEM OF UNITS (SI)—

SI UNITS AS ESTABLISHED BY THE SI:

– LENGTH – meter (m)

– VOLUME – cubic meter (m3)

– MASS – kilogram (kg)

– TEMPERATURE – Kelvin (K)

TimeTime

Base unit: SECOND (sec)

• Conversions:

only non-decimal base unit

60 sec = 1 min 60 min = 1 hr

Precision and Accuracy in Precision and Accuracy in MeasurementsMeasurements

Precision vs. AccuracyPrecision vs. AccuracyDefinitions:

• Precision—how close answers are to each other (reproducibility)

• Accuracy—how close answer is to accepted (true) value (agreement to accepted value)

Precision and Accuracy in Precision and Accuracy in MeasurementsMeasurements

Percent Error - a way to calculate accuracy in the lab

Equation:

% Error = | Accepted Value – Exp. Value | x 100

Accepted Value

Precision and Accuracy in Precision and Accuracy in MeasurementsMeasurements

Ex1.9 A student reports the density of a pure substance to be 2.83 g/mL. The accepted value is 2.70 g/mL. What is the percent error for the student’s results?

Equation:% Error = | Accepted Value – Exp. Value | x 100

Accepted Value

Scientific NotationScientific Notation

Exponential (Scientific) Notation—See Worksheet

Significant Figures: Why are they Important?

Numbers in math: no units, abstract, no context, can read calculator output exactly for answer.

vs.

Numbers in chemistry: measurements – include units.

SIG FIGS WILL BE IMPORTANT SIG FIGS WILL BE IMPORTANT THROUGHOUT THIS COURSETHROUGHOUT THIS COURSE!!

Graduated Cylinder Example

http://learningchemistryeasily.blogspot.com/2013/07/

precision-of-measurement-and.html

49

What are significant figures?(aka sig figs)

Significant figures are all the digits in a measurement that are known with certainty plus a last digit that must be estimated.

With experimental values your answer can have too few or too many sig figs, depending on how you round.

50

How Rounding Influences Sig Figs

1.024 x 1.2 = 1.2288Too many numerals(sig figs)

Too precise 1.024 x 1.2 = 1

Too few numerals(sig figs)

Not precise enough

51

Why This Concept is Important

We will be adding, subtracting, multiplying and dividing numbers throughout this course.

You MUST learn how many sig figs to report each answer in or the answer is meaningless.

You must report answers on lab reports & tests/quizzes with the correct number of sig figs (+/- 1) or else you will lose points!!

52

How Do We Find the Correct Number of Sig Figs In an Answer?

First, we will learn to count number of sig figs in a number. You must learn 4 rules and how to apply them.

Second, we will learn the process for rounding when we add/subtract or multiply/divide. We will then apply this process in calculations.

53

Rules for Counting Sig Figs

Rule #1: Read the number from left to right and count all digits, starting with the first digit that is not zero. Do NOT count final zero’s unless there is a decimal point in the number!

3 sig figs

4 sig figs

5 sig figs

23.4

234

0.234

2340

203

345.6

3.456

0.03456

34560

3405

678.90

6789.0

0.0067890

67008

60708

54

Rules for Counting Sig Figs

Rule #2: A final zero or zero’s will be designated as significant if a decimal point is added after the final zero.

3 sig figs

4 sig figs

5 sig figs

2340

23400

234000

2340000

2340.

2000.

20000.

23400.

55

Rules for Counting Sig Figs

Rule #3: If a number is expressed in standard scientific notation, assume all the digits in the scientific notation are significant.

2 sig figs

3 sig figs 4 sig figs

2.3 x 102

2.0 x 103

2.30 x 102

2.00 x 103

2.300 x 102

2.000 x 103

56

Rules for Counting Sig Figs

Rule #4: Any number which represents a numerical count or is an exact definition has an infinite number of sig figs and is NOT counted in the calculations.

Examples: 12 inches = 1 foot (exact definition)1000 mm = 1 m (exact definition)25 students = 1 class (count)

57

Practice Counting Sig Figs How many sig figs in each of the following?

1.2304 mm1.23400 cm1.200 x 105 mL0.0230 m0.02 cm8 ounces = 1 cup30 cars in the parking lot

58

Answers to Practice

How many sig figs in each of the following? 1.2304 mm (5) 1.23400 cm (6) 1.200 x 105 mL (4) 0.0230 m (3) 0.02 cm (1) 8 ounces = 1 cup (infinite, exact def.) 30 cars in the parking lot (infinite,

count)

59

General Rounding Rule When a number is rounded off, the last digit to

be retained is increased by one only if the following digit is 5 or greater.

EXAMPLE: 5.3546 rounds to 5 (ones place)5.35 (hundredths place)5.355 (thousandths place)5.4 (tenths place)

You will lose points for rounding incorrectly!

60

Process for Addition/Subtraction

Step #1: Determine the number of decimal places in each number to be added/subtracted.

Step #2: Calculate the answer, and then round the final number to the least number of decimal places from Step #1.

61

Addition/Subtraction ExamplesExample #1:Round to tenths place.

Example #2:Round to hundredths place.

Example #3:Round to ones place.

23.456+ 1.2+ 0.05-------------- 24.706Rounds to: 24.7

3.56- 0.14- 1.3501--------------- 2.0699Rounds to: 2.07

14+ 0.735+ 12.0-------------- 26.735Rounds to: 27

62

Process for Multiplication/Division

Step #1: Determine the number of sig figs in each number to be multiplied/divided.

Step #2: Calculate the answer, and then round the final number to the least number of sig figs from Step #1.

63

Multiplication/Division ExamplesExample #1:Round to 1 sig fig.

Example #2:Round to2 sig figs.

Example #3:Round to 3 sig figs.

23.456x 1.2x 0.05-------------- 1.40736Rounds to: 1

3.56x 0.14x 1.3501---------------0.67288984Rounds to: 0.67

14.0/ 11.73

--------------1.193520887Rounds to: 1.19

64

PracticeWrite the answers to the following computations using the correct number of sig figs

a. 129.0 g + 53.21 g + 1.4365 g =

b. 10.00 m - 0.0448 m =

c. 23.456 × 4.20 × 0.010 =

d. 17 ÷ 22.73 =

65

Important Rounding Rule

When you are doing several calculations, carry out all the calculations to at LEAST

one more sig fig than you need (I carry all digits in my calculator memory) and

only round off the FINAL result.

Use of Conversion FactorsUse of Conversion Factors

• Also known as dimensional analysis or factor-label method (or unit conversions)

Dimensional analysis/ Use of conversion factorsDefinition: technique to change one unit to another using a conversion factor

Ex.) # in original unit x new unit = New # in new unit original unit

Using Dimensional Using Dimensional AnalysisAnalysis

• Express the quantity 1.00 ft in different dimensions (inches, meters).

• Conversion factors:

Using Dimensional Using Dimensional AnalysisAnalysis

Example 1L.5 Calculate the following single step conversions: a. How many Joules are equivalent to 25.5 calories if 1 cal = 4.184 joules? b. How many liters gasoline can be contained in a 22.0 gallon gas tank if 3.785 L = 1 gal?

Using Dimensional Using Dimensional AnalysisAnalysis

Example 1L.6 The following multiple step conversions can be solved, knowing that 1 in = 2.54 cm. Convert the length of 5.50 ft to millimeters.

Using Dimensional Using Dimensional AnalysisAnalysis

Example 1L.7 The average velocity of hydrogen molecules at 0oC is 1.69 x 105 cm/s. Convert this to miles per hour.

Using Dimensional Using Dimensional AnalysisAnalysis

Example 1L.8 A piece of iron with a volume of 2.56 gal weighs 168.04 lbs. Convert this density to scruples per drachm with the following conversion factors:

1.00 L = 0.264 gal, 1.000 kg = 2.205 lb, 1.000 scruple = 1.296 g, 1.000 mL = 0.2816 drachm.

Using Dimensional Using Dimensional Analysis: Area Analysis: Area ConversionsConversions

Example 1L.14 Express the area of a 27.0 sq yd carpet in square meters.

Conversion factors needed:

Using Dimensional Using Dimensional Analysis:Analysis:

Volume ConversionsVolume ConversionsExample 1L.15 Convert 17.5 quarts to cubic meters.

(1 L = 1.057 qt, 1 ft3 = 28.32 L)

Properties of Substances1. Every pure substance has its own unique set of

properties that serve to distinguish it from all other substances.

2. Properties used to identify a substance must be intensive; that is, they must be independent of amount.

– Extensive properties depend on the amount.Classify the following as either intensive (I) or extensive (E):a. densityb. massc. melting pointd. volume

Properties of Properties of SubstancesSubstances

Properties of Properties of SubstancesSubstances• Density is an INTENSIVE

property of matter, which does NOT depend on quantity of matter.

• Contrast with EXTENSIVE properties of matter, which do depend on quantity of matter.

–Examples of extensive properties: mass and volume.

StyrofoamStyrofoam BrickBrick

Chemical and Physical Properties

Chemical property – Observed when the substance changes to a new one.

Example of a chemical property:Copper reacts with air to form copper (II) oxide.

Physical property – Observed without changing the substance to a new one.

Example of a physical property:Water boils at 100oC.

Physical ChangesPhysical ChangesPhysical changes do not

result in a new substance:

• boiling of a liquid

• melting of a solid

• dissolving a solid in a liquid to give a SOLUTION.

Physical vs. Chemical Physical vs. Chemical ChangeChange

• Another name for a Chemical change is a chemical reaction — change that results in a new substance.

Example:

Classify the following as either physical (P) or chemical (C) changes:

a. ice melting

b. gasoline burning

c. food spoiling

d. log of wood sawed in half

DensityDensityDensityDensity• Density is an INTENSIVE

property of matter, which does NOT depend on quantity of matter.

• Contrast with EXTENSIVE properties of matter, which do depend on quantity of matter.

–Examples of extensive properties: mass and volume.

StyrofoamStyrofoam BrickBrick

DENSITY : Review

Definition: ratio of mass to volume for an object

Mercury

13.6 g/cm13.6 g/cm33 21.5 g/cm21.5 g/cm33

Aluminum

2.7 g/cm2.7 g/cm33

Platinum

Density mass (g)

volume (cm 3 )

82

Sample Problem A piece of copper has a mass of 57.54 g. It is 9.36 cm long, 7.23 cm wide, and 0.95 mm thick. Calculate density (g/cm3).

Density as a Conversion Factor

Density is a “bridge” between mass and volume, or vice versa

Volume (cm3) x density g = mass (g) cm3

Mass (g) density cm3 = Volume (cm3) g

84

SAMPLE PROBLEM: Mercury (Hg) has a density of SAMPLE PROBLEM: Mercury (Hg) has a density of 13.6 g/cm13.6 g/cm33. What is the mass of 95 mL of Hg in . What is the mass of 95 mL of Hg in grams? In pounds?grams? In pounds?

SAMPLE PROBLEM: Mercury (Hg) has a density of SAMPLE PROBLEM: Mercury (Hg) has a density of 13.6 g/cm13.6 g/cm33. What is the mass of 95 mL of Hg in . What is the mass of 95 mL of Hg in grams? In pounds?grams? In pounds?

Solve the problem using Solve the problem using DENSITY AS A DENSITY AS A CONVERSION FACTOR.CONVERSION FACTOR.

Ex1L.9 Ex1L.9 What is the density of Hg if 164.56 g occupy a What is the density of Hg if 164.56 g occupy a volume of 12.1cmvolume of 12.1cm33??

Ex1L.10 Ex1L.10 What is the mass of 2.15 cmWhat is the mass of 2.15 cm33 of Hg? of Hg?

Ex1l.11 Ex1l.11 What is the volume of 94.2 g of Hg?What is the volume of 94.2 g of Hg?

Example 1L.12: Example 1L.12: Given the following densities: chloroform Given the following densities: chloroform 1.48 g/cm1.48 g/cm33 and mercury 13.6 g/cm and mercury 13.6 g/cm33 and copper 8.94 g/cm and copper 8.94 g/cm33. . Calculate if a 50.0 mL container will be large enough to Calculate if a 50.0 mL container will be large enough to hold a mixture of 50.0 g of mercury, 50.0 g of chloroform hold a mixture of 50.0 g of mercury, 50.0 g of chloroform and a 10.0 g chunk of copper.and a 10.0 g chunk of copper.

Example 1L.13 Example 1L.13 How many kilograms of methanol How many kilograms of methanol (d = 0.791 g/mL) does it take to fill the 15.5-gal fuel (d = 0.791 g/mL) does it take to fill the 15.5-gal fuel tank of an automobile modified to run on methanol?tank of an automobile modified to run on methanol?

Density of WaterDensity of Water

Density of water changes with temperature

(As water temperature changes, volume changes)

Maximum density of water is at

4oC = 0.999973 g/cm3

(often rounded to 1.00 g/cm3)

Derived UnitsDerived Units• Definition: derived from base units

Example: m/sec (unit of speed)

Divide meters by seconds

• Volume examples

m3 (m x m x m) or cm3 (cm x cm x cm)