Chapter 1 matter and measurements
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Chapter 1: Matter and Measurements. What is Chemistry?. Biology vs. Chemistry vs. Physics. 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?.

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Chapter 1: Matter and Measurements

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Chapter 1 matter and measurements

Chapter 1: Matter and Measurements


What is chemistry

What is Chemistry?

Biology vs.

Chemistry vs.

Physics


What is chemistry1

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 chemistry2

What is Chemistry?

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

(We formally define “stuff” as matter!)


What is chemistry3

What is Chemistry?

Think, pair, share:

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


Review scientific methods

Review: Scientific Methods

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.


Review scientific methods1

Review: Scientific Methods

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

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

    b. Scientific theories explain WHY something is observed.

    • Example of a scientific theory:


  • Review scientific methods2

    Review: Scientific Methods

    6. Steps of the Scientific Method—Review

    a.

    b.

    c.

    d.

    e.

    f.


    What is chemistry4

    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 chemistry5

    What is Chemistry?

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

    (We formally define “stuff” as matter!)


    Matter

    Matter


    Chapter 1 matter and measurements

    Matter


    Elements

    Elements

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


    Compounds

    Compounds

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


    Compounds1

    Compounds

    • Fixed compositions

      • A given compound contains the same elements in the same percent by mass


    Compounds2

    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


    Electrolysis

    Electrolysis


    Mixtures

    Mixtures

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

      • Homogeneous

      • Heterogeneous


    Homogeneous mixtures

    Homogeneous Mixtures

    • 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)


    Heterogeneous mixtures

    Heterogeneous Mixtures

    • Non-uniform; composition varies throughout the mixture


    Separating mixtures

    Separating Mixtures

    • Filtration


    Separating mixtures1

    Separating Mixtures

    • Distillation


    Separating mixtures2

    Separating Mixtures

    • Chromatography


    Scientific measurements

    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 system

    SI 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

    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 units

    INSTRUMENTS & UNITS

    Use SI units — based on the metric system

    Length

    Mass

    Time

    Temperature

    Meter, m

    Kilogram, kg

    Seconds, s

    Celsius degrees, ˚C

    Kelvins, K


    Length

    Length

    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 –


    Length1

    Length

    Base unit: METER

    Conversions:

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

    1 Gm = 106 m


    Units of length

    O—H distance =

    9.58 x 10-11 m

    9.58 x 10-9 cm

    0.0958 nm

    Units of Length

    • 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


    Volume

    Volume

    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!)


    Chapter 1 matter and measurements

    Mass

    • 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 scales

    Temperature Scales

    Fahrenheit

    Celsius

    Kelvin

    Anders Celsius

    1701-1744

    Lord Kelvin

    (William Thomson)

    1824-1907

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


    Temperature scales1

    Temperature Scales

    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


    Temperature scales2

    Temperature Scales

    100 oF

    38 oC

    311 K

    oF

    oC

    K


    Si system1

    SI System

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


    Calculations using temperature

    Calculations Using Temperature

    Fahrenheit/Celsius

    T (F) = 1.8 t (˚C) + 32


    Calculations using temperature1

    Calculations Using Temperature

    • 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


    Chapter 1 matter and measurements

    Problem

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


    Si system base units

    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)


    Chapter 1 matter and measurements

    Time

    Base unit: SECOND (sec)

    • Conversions:

      only non-decimal base unit

      60 sec = 1 min 60 min = 1 hr


    Precision and accuracy in measurements

    Precision and Accuracy in Measurements

    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)


    Precision and accuracy in measurements1

    Precision and Accuracy in Measurements

    Percent Error - a way to calculate accuracy in the lab

    Equation:

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

    Accepted Value


    Precision and accuracy in measurements2

    Precision and Accuracy in Measurements

    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 notation

    Scientific Notation

    Exponential (Scientific) Notation—See Worksheet


    Significant figures why are they important

    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 THROUGHOUT THIS COURSE!


    Graduated cylinder example

    Graduated Cylinder Example

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


    What are significant figures aka sig figs

    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.


    How rounding influences sig figs

    How Rounding Influences Sig Figs

    • 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


    Why this concept is important

    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!!


    How do we find the correct number of sig figs in an answer

    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.


    Rules for counting sig figs

    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!


    Rules for counting sig figs1

    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.


    Rules for counting sig figs2

    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.


    Rules for counting sig figs3

    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)


    Practice counting sig figs

    Practice Counting Sig Figs

    • 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


    Answers to practice

    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)


    General rounding rule

    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!


    Process for addition subtraction

    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 ofdecimal places from Step #1.


    Addition subtraction examples

    Addition/Subtraction Examples


    Process for multiplication division

    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.


    Multiplication division examples

    Multiplication/Division Examples


    Practice write the answers to the following computations using the correct number of sig figs

    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 =


    Important rounding rule

    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 factors

    Use of Conversion Factors

    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


    Using dimensional analysis

    Using Dimensional Analysis

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

    Conversion factors:


    Using dimensional analysis1

    Using Dimensional Analysis

    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 analysis2

    Using Dimensional Analysis

    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 analysis3

    Using Dimensional Analysis

    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 analysis4

    Using Dimensional Analysis

    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 analysis area conversions

    Using Dimensional Analysis: Area Conversions

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

    Conversion factors needed:


    Using dimensional analysis volume conversions

    Using Dimensional Analysis:Volume Conversions

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


    Properties of substances

    Properties of Substances

    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


    Properties of substances1

    Properties of Substances

    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 and physical properties

    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 changes

    Physical Changes

    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.


    Physical vs chemical change

    Physical vs. Chemical Change

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


    Example

    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


    Density

    Density

    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.


    Density review definition ratio of mass to volume for an object

    mass

    (

    g

    )

    =

    Density

    3

    volume

    (

    cm

    )

    Platinum

    Mercury

    Aluminum

    DENSITY : ReviewDefinition: ratio of mass to volume for an object

    2.7 g/cm3

    13.6 g/cm3

    21.5 g/cm3


    Chapter 1 matter and measurements

    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 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


    Chapter 1 matter and measurements

    SAMPLE PROBLEM: Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in grams? In pounds?

    Solve the problem using DENSITY AS A CONVERSION FACTOR.


    Chapter 1 matter and measurements

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


    Chapter 1 matter and measurements

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


    Chapter 1 matter and measurements

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


    Chapter 1 matter and measurements

    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.


    Chapter 1 matter and measurements

    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

    Density 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 units

    Derived 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)


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