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Chapter 1: Matter and Measurements

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Biology vs.

Chemistry vs.

Physics

Biology

Physics

Chemistry

The study of living organisms

The study of forces & motion

The study of matter and its reactions and properties

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

(We formally define “stuff” as matter!)

Think, pair, share:

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

1. Hypothesis

- Suggested solution to a problem
2. Experiment

- A controlled method of testing a hypothesis
3. Data

- Organized observations
a. Data is always reproducible.

4. Scientific Law

- Statement which summarizes results of many observations and experiment
a. Scientific laws explain WHAT is observed.

- Example of a scientific law:
5. Scientific Theory

- Example of a scientific law:

b. Scientific theories explain WHY something is observed.

- Example of a scientific theory:

6. Steps of the Scientific Method—Review

a.

b.

c.

d.

e.

f.

Biology

Physics

Chemistry

The study of living organisms

The study of forces & motion

The study of matter and its reactions and properties

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

(We formally define “stuff” as matter!)

Matter

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

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

- Fixed compositions
- A given compound contains the same elements in the same percent by mass

- 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

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

- Composition is the same throughout the mixture
- Examples: salt water, soda water, brass

- A.k.a. a solution
- Solute in a solvent (salt dissolved in water)

- Non-uniform; composition varies throughout the mixture

- Filtration

- Distillation

- Chromatography

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.

- 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)

- 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:

Use SI units — based on the metric system

Length

Mass

Time

Temperature

Meter, m

Kilogram, kg

Seconds, s

Celsius degrees, ˚C

Kelvins, K

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 –

Base unit: METER

Conversions:

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

1 Gm = 106 m

O—H distance =

9.58 x 10-11 m

9.58 x 10-9 cm

0.0958 nm

- 1 kilometer (km) = 1000 meters (m)
- 10-2 meter (m) = 1 centimeter (cm)
- 102 meter (m) = 1 hectometer (Hm)
- 1 nanometer (nm) = 1.0 x 10-9 meter

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!)

- 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

Fahrenheit

Celsius

Kelvin

Anders Celsius

1701-1744

Lord Kelvin

(William Thomson)

1824-1907

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

212 ˚F

100 ˚C

373 K

100 K

180˚F

100˚C

32 ˚F

0 ˚C

273 K

Fahrenheit

Celsius

Kelvin

Boiling point of water

Freezing point of water

Notice that 1 kelvin degree = 1 degree Celsius

100 oF

38 oC

311 K

oF

oC

K

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

Fahrenheit/Celsius

T (F) = 1.8 t (˚C) + 32

- Some calculations are in kelvins (especially important for Ch 5!!)
- T (K) = t (˚C) + 273.15 (273)
- Body temp = 37 ˚C + 273 = 310 K
- Liquid nitrogen = -196 ˚C + 273 = 77 K

Problem

Example 1L.1 A baby has a temperature of 39.8oC. Express this temperature in oF and K.

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)

Base unit: SECOND (sec)

- Conversions:
only non-decimal base unit

60 sec = 1 min 60 min = 1 hr

Precision vs. Accuracy

Definitions:

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

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

Percent Error - a way to calculate accuracy in the lab

Equation:

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

Accepted Value

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

Exponential (Scientific) Notation—See Worksheet

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 THROUGHOUT THIS COURSE!

http://learningchemistryeasily.blogspot.com/2013/07/precision-of-measurement-and.html

- 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.

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

- 1.024 x 1.2 = 1Too few numerals(sig figs)
Not precise enough

- 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!!

- 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.

- 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!

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

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

- 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)

- How many sig figs in each of the following?
- 1.2304 mm
- 1.23400 cm
- 1.200 x 105 mL
- 0.0230 m
- 0.02 cm
- 8 ounces = 1 cup
- 30 cars in the parking lot

- 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)

- 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!

- 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 ofdecimal places from Step #1.

- 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.

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 =

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.

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

Dimensional analysis/ Use of conversion factors

Definition: technique to change one unit to another using a conversion factor

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

original unit

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

Conversion factors:

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?

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.

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

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.

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

Conversion factors needed:

Example 1L.15 Convert 17.5 quarts to cubic meters. (1 L = 1.057 qt, 1 ft3 = 28.32 L)

1. Every pure substance has its own unique set of propertiesthat 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. density

b. mass

c. melting point

d. volume

Brick

Styrofoam

- 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.

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 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.

- Another name for a Chemical change is achemical reaction — change that results in a new substance.

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

Brick

Styrofoam

- 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.

mass

(

g

)

=

Density

3

volume

(

cm

)

Platinum

Mercury

Aluminum

2.7 g/cm3

13.6 g/cm3

21.5 g/cm3

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 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

Solve the problem using DENSITY AS A CONVERSION FACTOR.

Ex1L.9 What is the density of Hg if 164.56 g occupy a volume of 12.1cm3?

Ex1L.10 What is the mass of 2.15 cm3 of Hg?

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

Example 1L.12: Given the following densities: chloroform 1.48 g/cm3 and mercury 13.6 g/cm3 and copper 8.94 g/cm3. 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 and a 10.0 g chunk of copper.

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

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)

- 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)