measuring volume temperature mass distance

Measuring Volume Temperature Mass Distance

Reading the Meniscus

Always read volume from the bottom of the meniscus. The meniscus is the curved surface of a liquid in a narrow cylindrical container.

Try to avoid parallax errors.Parallax errors arise when a meniscus or needle is viewed from an angle rather than from straight-on at eye level.

Correct: Viewing the meniscus

at eye level

Incorrect: viewing the meniscus

from an angle

Graduated CylindersThe glass cylinder has etched marks to indicate volumes, a pouring lip, and quite often, a plastic bumper to prevent breakage.

Measuring Volume Determine the volume contained in a graduated cylinder by reading the bottom of the meniscus at eye level.

Read the volume using all certain digits and one uncertain digit.

Certain digits are determined from the calibration marks on the cylinder. The uncertain digit (the last digit of the reading) is estimated.

Use the graduations to find all certain digits

There are two unlabeled graduations below the meniscus, and each graduation represents 1 mL, so the certain digits of the reading are…

52 mL.

Estimate the uncertain digit and take a reading

The meniscus is about eight tenths of the way to the next graduation, so the final digit in the reading is . The volume in the graduated cylinder is

0.8 mL

52.8 mL.

10 mL GraduateWhat is the volume of liquid in the graduate?

_ . _ _ mL6 26

25mL graduated cylinder What is the volume of liquid in the graduate?

_ _ . _ mL1 1 5

100mL graduated cylinder What is the volume of liquid in the graduate?

_ _ . _ mL5 2 7

Self TestExamine the meniscus below and determine the volume of liquid contained in the graduated cylinder.

The cylinder contains:_ _ . _ mL7 6 0

The Thermometero Determine the temperature by reading the scale on the thermometer at eye level.

o Read the temperature by using all certain digits and one uncertain digit. o Certain digits are determined from the

calibration marks on the thermometer. o The uncertain digit (the last digit of the

reading) is estimated. o On most thermometers encountered in a

general chemistry lab, the tenths place is the uncertain digit.

Do not allow the tip to touch the walls or the bottom of the flask.

If the thermometer bulb touches the flask, the temperature of the glass will be measured instead of the temperature of the solution. Readings may be incorrect, particularly if the flask is on a hotplate or in an ice bath.

Reading the ThermometerDetermine the readings as shown below on Celsius thermometers:

_ _ . _ C _ _ . _ C8 7 4 3 5 0

Measuring Mass – Electronic balance

Our balances read to the hundredths place,the uncertain digit is the hundredths place ( _ _ _ . _ X)

Balance Rules In order to protect the balances and ensure accurate results, a number of rules should be followed:

Always check that the machine is zeroed before adding any substance on it. This is done by pushing the ZERO button before weighing any substance.

Never weigh directly on the balance pan. Always use a piece of weighing paper, weigh boat, or beaker to protect it.

Do not weigh hot or cold objects. Clean up any spills around the balance

immediately.

Determining Mass

1. Place container (weigh boat or beaker) on the pan

2. Zero the balance by hitting the zero bottom

3. Place object on pan

4. Read the digits to the last digit

Metric Ruler

• Rulers are used to measure distance• Can be used to measure the volume

of a cube (L x W x H) length x width x height

• Determine the temperature by reading the scale on the thermometer at eye level.

• Read the ruler by using all certain digits and one uncertain digit.

Reading a Ruler• Determine the readings as shown below

on centimeter ruler:

_ _ . _ cm

_ _ . _ cm

The SI System of Measurement

The Nature of Measurement

• Part 1 - number• Part 2 - scale (unit)

• Examples:• 20 grams

• 6.63 x 10-34 Joule·seconds

A Measurement is a quantitative observation consisting of TWO parts

The Fundamental SI Units (le Système International, SI)

Base Quantity Name of unit Symbol

Length Meter m

Mass Kilogram kg

Time Second s

Temperature Kelvin KAmount of Substance Mole mol

SI PrefixesCommon to Chemistry

Prefix Unit Abbr. ExponentKilo k 103

Deci d 10-1

Centi c 10-2

Milli m 10-3

Micro 10-6

Metric Acronym

• K King• H Henry• D Died• M Monday• D Drinking • C Chocolate • M Milk

Metric ConversionsgmL 10-1 10-2 10-3101102103

Baseunit

deci centi millidekahectokilo

Conversions in the metric system are merely a matter of moving a decimal point. The “base unit” means the you have a quantity (grams, meters, Liters, etc without a prefix.


Baseunit


Example #1: Convert 18 liters to milliliters

18 L1 2 3

18 liters = 18 000 milliliters


Baseunit


Example #2: Convert 450 milligrams to grams

123450 mg450 mg = 0.450 g

Metric Conversions Ladder Method

T. Trimpe 2008 http://sciencespot.net/

KILO1000Units

HECTO100

UnitsDEKA

10Units

DECI0.1

UnitCENTI

0.01Unit

MILLI0.001Unit

MetersLitersGrams

Ladder Method

How do you use the “ladder” method? 1st – Determine your starting point.

2nd – Count the “jumps” to your ending point.

3rd – Move the decimal the same number of jumps in the same direction.

4 km = _________ m

12

3

How many jumps does it take?

Starting Point Ending Point

4.1

__.2

__.3

__. = 4000 m

Try these conversions using the ladder method.

1000 mg = _______ g 1 L = _______ mL 160 cm = _______ mm

14 km = _______ m 109 g = _______ kg 250 m = _______ km

Conversion Practice

Compare using <, >, or =.

56 cm 6 m 7 g 698 mg

Write the correct abbreviation for each metric unit.

1) Kilogram _____ 4) Milliliter _____ 7) Kilometer _____

2) Meter _____ 5) Millimeter _____ 8) Centimeter _____

3) Gram _____ 6) Liter _____ 9) Milligram _____

Try these conversions, using the ladder method.

10) 2000 mg = _______ g 15) 5 L = _______ mL 20) 16 cm = _______ mm

11) 104 km = _______ m 16) 198 g = _______ kg 21) 2500 m = _______ km

12) 480 cm = _____ m 17) 75 mL = _____ L 22) 65 g = _____ mg

13) 5.6 kg = _____ g 18) 50 cm = _____ m 23) 6.3 cm = _____ mm

14) 8 mm = _____ cm 19) 5.6 m = _____ cm 24) 120 mg = _____ g

Metric Conversion Challenge

Compare using <, >, or =.

25) 63 cm 6 m 27) 5 g 508 mg 29) 1,500 mL 1.5 L

26) 536 cm 53.6 dm 28) 43 mg 5 g 30) 3.6 m 36 cm

Metric Dimensional Analysis Practice

Problem #1Convert 400 mL to Liters

400 mL= L

mLL

1 0001 .400

= 0.4 L= 4x10-1 L

Problem #2Convert 10 meters to mm

10 m= m

mmmm1

1 000 10 000

= 1x104 mm

Problem #3Convert 73 grams to kg

73 g= kg

gkg

1 0001 0.073

= 7.3x10-2 kg

Problem #4Convert 0.02 kilometers to m

0.02 km= m

kmm

11 000 20

= 2x101 m

Problem #5Convert 20 centimeters to m

20 cm= m

cmm

1001 0.20

= 2x10-1 m

Problem #6Convert 450 milliliters to dL

450 mL= dL

mLdL

1001 4.5

Problem #7Convert 10 kilograms to grams

10 kg = gkgg

11 000 10 000

= 1x104 g

Problem #8Convert 935 mg to cg

935 mg = cgmgcg

101 93.5

= 9.35x101 cg

Problem #9Convert 5.2 kg to mg

5.2 kg = mgkg

mg1

1 000 000

= 5 200 000 mg= 5.2x106 mg

Problem #10Convert 175 mL to cm3

175 mL = cm3

mLcm3

11 175

= 1.75x102 cm3

Uncertainty and Significant Figures

Cartoon courtesy of Lab-initio.com

Uncertainty in Measurement

• A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty.

Why Is there Uncertainty? Measurements are performed with

instruments No instrument can read to an infinite

number of decimal places

Which of these balances has the greatest uncertainty in measurement?

Precision and Accuracy• Accuracy refers to the agreement of a

particular value with the true value.

• Precision refers to the degree of agreement among several measurements made in the same manner.

Neither accurate

nor precise

Precise but not accurate

Precise AND accurate

Types of Error

• Random Error (Indeterminate Error) - measurement has an equal probability of being high or low.

• Systematic Error (Determinate Error) - Occurs in the same direction each time (high or low), often resulting from poor technique or incorrect calibration.

Rules for Counting Significant Figures - Details

• Nonzero integers always count as significant figures.

• 3456 has • 4 significant figures


• Zeros• Leading zeros do not count as significant

figures.

• 0.0486 has• 3 significant figures


• Zeros• Captive zeros always count as significant

figures.



• Zeros• Trailing zeros are significant only if

the number contains a decimal point.



• Exact numbers have an infinite number of significant figures.

• 1 inch = 2.54 cm, exactly

Sig Fig Practice #1How many significant figures in each of the following?

1.0070 m 5 sig figs

17.10 kg 4 sig figs

100,890 L 5 sig figs

3.29 x 103 s 3 sig figs

0.0054 cm 2 sig figs

3,200,000 2 sig figs

Rules for Significant Figures in Mathematical Operations

• Multiplication and Division: # sig figs in the result equals the number in the least precise measurement used in the calculation.

• 6.38 x 2.0 =• 12.76 13 (2 sig figs)

Sig Fig Practice #2

3.24 m x 7.0 m

Calculation Calculator says: Answer

22.68 m2 23 m2

100.0 g ÷ 23.7 cm3 4.219409283 g/cm3 4.22 g/cm3

0.02 cm x 2.371 cm 0.04742 cm2 0.05 cm2

710 m ÷ 3.0 s 236.6666667 m/s 240 m/s

1818.2 lb x 3.23 ft 5872.786 lb·ft 5870 lb·ft

1.030 g ÷ 2.87 mL 2.9561 g/mL 2.96 g/mL

Rules for Significant Figures in Mathematical Operations

• Addition and Subtraction: The number of decimal places in the result equals the number of decimal places in the least precise measurement.

• 6.8 + 11.934 =• 18.734 18.7 (3 sig figs)

Sig Fig Practice #3

3.24 m + 7.0 m

Calculation Calculator says: Answer

10.24 m 10.2 m

100.0 g - 23.73 g 76.27 g 76.3 g

0.02 cm + 2.371 cm 2.391 cm 2.39 cm

713.1 L - 3.872 L 709.228 L 709.2 L

1818.2 lb + 3.37 lb 1821.57 lb 1821.6 lb

2.030 mL - 1.870 mL 0.16 mL 0.160 mL

Scientific Notation

In science, we deal with some very LARGE numbers:

1 mole = 602000000000000000000000

In science, we deal with some very SMALL numbers:

Mass of an electron =0.000000000000000000000000000000091 kg

Scientific Notation

Imagine the difficulty of calculating the mass of 1 mole of electrons!

0.000000000000000000000000000000091 kg x 602000000000000000000000 ???????????????????????????????????

Scientific Notation:A method of representing very large

or very small numbers in the form:

M x 10n

M is a number between 1 and 9

n is an integer

2 500 000 000

Step #1: Insert an understood decimal point

.Step #2: Decide where the decimal must end up so that one number is to its leftStep #3: Count how many places you bounce the decimal point

123456789

Step #4: Re-write in the form M x 10n

2.5 x 109

The exponent is the number of places we moved the decimal.

0.0000579

Step #2: Decide where the decimal must end up so that one number is to its leftStep #3: Count how many places you bounce the decimal point

Step #4: Re-write in the form M x 10n

1 2 3 4 5

5.79 x 10-5

The exponent is negative because the number we started with was less than 1.

PERFORMING CALCULATIONS IN SCIENTIFIC

NOTATION

ADDITION AND SUBTRACTION

Review:Scientific notation expresses a number in the form: M x 10n

1 M 10n is an integer

4 x 106

+ 3 x 106IF the exponents are the same, we simply add or subtract the numbers in front and bring the exponent down unchanged.

7 x 106

4 x 106

- 3 x 106The same holds true for subtraction in scientific notation.

1 x 106

4 x 106

+ 3 x 105If the exponents are NOT the same, we must move a decimal to make them the same.

4.00 x 106

+ 3.00 x 105 + .30 x 106

4.30 x 106Move the decimal on the smaller number!

4.00 x 106

A Problem for you… 2.37 x 10-6

+ 3.48 x 10-4

+ 3.48 x 10-4

Solution…002.37 x 10-

6

+ 3.48 x 10-4

Solution…0.0237 x 10-4

3.5037 x 10-4

Density

Density Density is an important intensive

property, which can be used to help determine the identity of an unknown substance.

While the mass or the volume of a substance will vary from sample to sample, the density will remain the same at a given temperature.

As you know, the density of a substance is a measure of how much mass is present in a given unit of volume.

Density is the measure of the “compactness” of a material

How close the atoms or molecules are to each other

More than “heaviness” - density includes how much space an object takes up!!

All substances have density including liquids, solids, and gases

Density

D = m/vm = mass (g)Which is the amount of matter in a

substance DO NOT confuse mass with weightWeight – is the force of gravity

exerted on an objectv = volume (ml) or (cm3)Is measured by water displacement or

using a ruler

QuestionWhich weighs more?50 kilograms of iron

Or

50 kilograms of feathers

QuestionWhich has a greater density?

iron

Or

feathers

Water

• Waters density = 1.000 g/ml at 4oC • Ice floats on water

Densities of Common Substances

If water has a density of 1.00 g/cm3 , which of the following will float or sink?

Factors that Affect Density

Temperature• For most substances, as temperature

increases the volume increases and as a result the density decreases.

Pressure

Dissolved solids in Liquids

2 Ways to Measure the Volume of a Solid

1. Objects with regular sides calculate L x W x H using a ruler

2. Water displacement. Place irregularly shaped solid into a volumetric cylinder half filled with water and place the solid into the cylinder and subtract the volumes

Useful Conversion Factors

1 cm3 = 1 ml1 dm3 = 1 L

Manipulating the Density Formula

Sample Problem 1

A student determines that a piece of an unknown material has a mass of 5.854 g and a volume of 7.57 cm3. What is the density of the material, rounded to the correct number of significant digits?

Sample Problem 2

Iron has a known density of 7.87 g/cm3. What would be the mass of a 2.5 cm3 piece of iron?

Sample Problem 3

Mercury has a density of 13.5 g/cm3. How much space would 50.0 g of mercury occupy?

Percent Error

Percent Error

Students often assume that each measurement that they make in the laboratory is true and accurate. Likewise, they often assume that the values that they derive through experimentation are very accurate. However, sources of error often prevent students from being as accurate as they would like. Percent error calculations are used to determine how close to the true values, or how accurate, their experimental values really are.

The value that the student comes up with is usually called the observed value, or the experimental value. A value that can be found in reference tables is usually called the true value, or the accepted value. The percent error can be determined when the true value is compared to the observed value according to the equation below:

Percent Error

Sample Problem 1

A student measures the mass and volume of a piece of copper in the laboratory and uses his data to calculate the density o the metal. According to his results, the copper has a density of 8.37 g/cm3. Curious about the accuracy of his results, the student consults a reference table and finds that the accepted value for the density of copper is 8.92 g/cm3. What would be the student's percent error?

Sample Problem 2

A student experimentally determines the specific heat of water to be 4.29 J/g x Co. He then looks up the specific heat of water on a reference table and finds that is 4.18 J/g x Co. What is his percent error?

Sample Problem 3

A student takes an object with an accepted mass of 200.00 grams and masses it on his own balance. He records the mass of the object as 196.5 g. What is his percent error?

Classification of Matter

Composition of Matter Flowchart

MATTERCan it be physically

separated?

Homogeneous Mixture

(solution)

Heterogeneous Mixture Compound Element

MIXTURE PURE SUBSTANCE

yes no

Can it be chemically decomposed?

noyesIs the composition uniform?

noyes

Classifying Matter by Composition

Elements- simplest kind of matter, made of one type of atom

An atom is the smallest unit of an element that maintains the properties of that element.

Cannot be broken down into simpler substances by ordinary chemical means

Ex. gold, copper, oxygen (on the periodic table)

Element


Compounds – matter composed of the atoms of two or more elements chemically bonded

Compounds can be broken down by chemical methods

When they are broken down, the components have completely different properties than the compound.

Ex. Sugar, salt, water, carbon dioxide

Compound


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

A mixture is mixed together physically.

Variable composition, often expressed by a percent composition by mass or volume (Ex. 5% salt and 95% water)

Mixture


Homogeneous Mixture– matter with a uniform compositionHomogeneous mixtures are also called solutions.

Ex. Salt water and Kool –aid

Homogeneous Mixture


A heterogeneous mixture is not the same throughout (not uniform).

Examples: M & M’s, Chocolate chip cookie, gravel, soil, rocks such as granite, blood, milk, salad, ocean water, etc.

Heterogeneous Mixture

Classify Itcopper wire, aluminum foil

Classify It

» EX: table salt (NaCl)

Classify ItGranite Apple Juice

Classify It

Examples:» magnesium» Pizza» Calcium chloride» Orange juice» Club soda

Classify ItExamples:

» magnesium» pizza» Calcium

chloride» Orange juice» Club soda

elementhetero. mixturecompoundhetero. mixtureHomo. (solution)

States of MatterSolid- matter that can not flow and

has definite volume and shapeLiquid- definite volume but no

definite shape and can flowGas- a substance without definite

volume or shape and can flow.Plasma- a substance that is similar to

a gas, but loses electrons due to its high temperature

States of Matter

Gas

Definite Volume?

NO

Definite Shape?

NO

NO

Particle position and movement

Close together, can move past each other - flow

Far apart, move rapidly - flow

Solid

Liquid

YES

YES

YES Packed tightly, vibrate about fixed point

States of Matter

Separating MixturesMixtures are separated by their

physical properties. Primary methods of separatingmixtures are:filtrationdistillationcentrifugechromatography

Separating Mixtures

Filtration is a method used to separate the components of mixtures that contain an insoluble solid and a liquid. Example: sand and water

Filtration

Separating Mixtures

Distillation is a method of separating substances in a mixture by evaporation of a liquid and subsequent condensation of its vapor. Example: desalination of salt water

Separating MixturesCentrifuge Used to separate solid-liquid mixtures such as those in blood. The centrifuge spins rapidly and causes the solid to settle to the bottom.

Ex. Separating blood

Separating MixturesChromatography is a method of

separating mixtures that uses a stationary phase and a mobile phase. Paper chromatography can be used to separate pigments because they move at different rates on the paper.

Chromatography

Properties & Changes of Matter

Properties of Matter

Physical Property- a property that can be observed and measured without changing the identity of the substance.

Examples? Mass, Density, Melting and Boiling Points

Chemical Property-relates to a substance’s ability to undergo changes that transform it into different substances.

Examples? Reactivity, Toxicity, and Chemical stability

Properties of MatterChemists use properties to identify and

separate matter. More than one property must be used for identification.

Intensive Properties – do not depend on the amount of matter present

Ex. Melting pt., boiling pt, density, conduct electricity

Extensive Properties – depend on the amount of matter present

Ex. Volume, mass

Changes in Matter

A physical change does not change the composition or identity of the substance.• Examples? • Boiled water is still water.• All phase changes are physical

changes

Solid Liquid

Gas

Melt Evaporate

Condense

Freeze

Changes in Matter

Sublimation is a process in which a solid changes directly to a gas without going through the liquid phase.

Ex: dry ice CO2

Deposition is a process in which a gas changes directly to a liquid without going through the liquid phase

Ex: liquid vapor to ice (frost on windshields)

Changes in Matter

A chemical change occurs when one or more substances are changed into new substances.

Reactants- substances that reactProducts- substances that formProducts have NEW PROPERTIES

2H2 + O2 → 2H2O reactants product

Law of Conservation of mass – Also known as Conservation of Matter. Matter can be neither created nor destroyed, though it can be rearranged. Mass remains constant in an ordinary chemical change.

Indications of chemical change

1. Production of energy in the forms of heat, light, sound, or electricity

2. Production of a gas3. Formation of a

precipitate4. A change in color5. A change in odor

What is Energy?Energy - is the ability to or capacity to

do work or to produce changeConservation of Energy (1st Law of Thermodynamics) - Energy can be neither created nor destroyed; the energy of the universe is constant.Two types: Kinetic energy – energy of motionPotential energy – is the energy due to

position of object

Chemical Energy

Chemical energy - is a special kind of potential energy• Is the energy involved in chemical

reactions

Energy Changes

• Some changes in matter release energy.

• For example, the explosion that occurs when hydrogen and oxygen react to form water is a release of energy.

• Heat energy and light energy are released as the reaction takes place.

Energy Changes• A change in matter in which energy is absorbed

from the surroundings is an endothermic process (heat enters).•EXAMPLES: melting ice & boiling water •When barium hydroxide reacts ammonium nitrate are mixed the test-tube feels cold to touch because energy has been absorbed

Energy Changes

• A change in matter in which energy is released is an exothermic process (heat exits).

• Examples: freezing water & condensation

• Burning of paper gives off heat to the surroundings.

Phase Changes

Courtesy www.lab-initio.com

Water phase

changes

What do you notice about the temperatureduring a phase change?

Phase Diagram

Represents phases as a function of temperature and pressure. Critical temperature: temperature above which the vapor can not be liquefied. Critical pressure: pressure required to liquefy AT the critical temperature. Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).

Phase changes by Name

Carbon dioxidePhase

Diagram for

Carbondioxide

CarbonPhase Diagram

for Carbon

Intro to Thermochemist

ryHeat and

Temperature

Thermochemistry

Thermochemistry – is the study of the transfer of energy as heat that accompany chemical reactions and physical changes

Heat and Temperature

Heat - is energy transferred between a system and its surroundings

• Heat absorbed or released is measured by a calorimeter

Heat

• Energy that flows from something warm to something cooler

• A hotter substance gives KE to a cooler one

• When heat is transferred (lost or gained), there is a change in the energy within the substance

QuestionsA. When you touch ice, heat is transferred from

1) your hand to the ice2) the ice to your hand

B. When you drink a hot cup of

coffee, heat is transferred from

1) your mouth to the coffee2) the coffee to your mouth

Answer

A. When you touch ice, heat is transferred from

1) your hand to the ice

B. When you drink a hot cup of coffee, heat is transferred from

2) the coffee to your mouth

Question

C. When you heat 200 g of water for 1 minute, the water temperature rises from 10°C to 18°C.

If you heat 400 g of water at 10°C in the same pan with the same amount of heat for 1 minute, what would you expect the final temperature to be?1) 10 °C 2) 14°C 3) 18°C

200 g 400 g

Answer2)14°C Heating twice the mass of water

using the same amount of heat will raise the temperature only half as much.

200 g 400 g

TemperatureTemperature – is a measure of the average

kinetic energy of particles• Measured in Kelvin• Common Temperatures you need to

KNOWCelsius Scale (0C)

– 00C water freezes, 1000C water boilsFahrenheit Scale (0F)

– 320F water freezes, 2120F water boilsKelvin Scale (K)

– 273 K water freezes, 373 K water boils– ABSOLUTE Zero 0 K

Temperature

K oC

50357

32298

• Measured in Kelvin K = 273 + oC

Convert the following

Some Equalities for Heat

Heat is measured in calories or joules• 1 kcal = 1000 cal• 1 calorie = 4.18 J (Joules)• 1 kJ = 1000 J

Specific Heat

• Why do some foods stay hot longer than others?

• Why is the beach sand hot, but the water is cool on the same hot day?

Specific HeatDifferent substances have different capacities for storing energy

It may take 20 minutes to heat water to 75°C. However, the same mass of aluminum might require 5 minutes and the same amount of copper may take only 2 minutes to reach the same temperature.

Specific Heat Values Specific heat - is the amount of heat needed

to raise the temperature of 1 g of a substance by 1°Ccal/g°C J/g°Cwater 1.00 4.18 aluminum0.22 0.90

copper 0.093 0.39silver 0.057 0.24

gold 0.031 0.13

QuestionsA. A substance with a large specific heat 1) heats up quickly 2) heats up

slowlyB. When ocean water cools, the

surrounding air 1) cools 2) warms 3) stays the same

C. Sand in the desert is hot in the day, and cool at night. Sand must have a 1) high specific heat 2) low specific heat

AnswersA. A substance with a large specific heat 2) heats up slowlyB. When ocean water cools, the

surrounding air 2) warms

C. Sand in the desert is hot in the day, and cool at night. Sand must have a

2) low specific heat

Measuring Heat

Requires• Grams of substance • Temperature change T

T = Tf - Ti

• Specific heat of the substance

Energy lost or gained

Calculating Heat

= mass x temp. change x specific heat

q = m x ∆T x Cp

A few key ideas:• If a substance receives heat and

experiences an increase in temperature then Q is a positive number and ∆T is a positive number.

• If a substance loses heat and experiences a decrease in temperature then Q is a negative number and ∆T is a negative number.

A few key ideas:• Q, heat energy, can be measured in

either Joules or calories. Just make sure that your units for c are consistent with your units for Q.

• ∆T, change in temperature, can be measured in K, °C, or °F. Just make sure that your units for c are consistent with your units for ∆T.

• Always start a problem by listing the given information (with units) and writing down the specific heat capacity equation.

A few key ideas:• You must ALWAYS show all work and make

sure you have consistent units on your final answer.

• The First Law of Thermodynamics states that if two substances exchange heat, the quantity of heat gained by one substance is exactly equal and opposite to the quantity of heat lost by the other substance.

ProblemThe element hydrogen has the highest specific heat of all elements. At a temperature of 25°C, hydrogen’s specific heat capacity is 14300J/(kg ∙ K). If the temperature of a 0.34 kg sample of hydrogen is to be raised by 25 K, how much heat will have to be transferred to the hydrogen?

ProblemAt 25°C, radon’s specific heat capacity is 94J/(kg ∙K). If the temperature of a 0.34kg sample of radon is to be raised by 25K, how much heat will have to be transferred to the radon?

Problem A 0.59 kg brass candlestick has an initial temperature of 98.0°C. If 21,100J of heat is removed from the candlestick to lower its temperature to 6.8°C, what is the specific heat capacity of brass?

ProblemA 0.38kg drinking glass is filled with a hot liquid. The liquid transfers 7032J of heat to the glass. If the temperature of the glass increases by 22K, what is the specific heat capacity of the glass?

End Of Unit 1

measuring volume temperature mass distance

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