Chapter 2 Matter and Energy. 2.1 Classification of Matter Matter is anything that has mass and...

Chapter 2Chapter 2

Matter and Energy

2.1 Classification of 2.1 Classification of MatterMatterMatter is anything that has mass

and occupies space. Classification of matters are◦Pure substance◦Mixture

Pure substancePure substance

Has fixed or definite composition◦An element, the simplest type of pure

substance and is composed of atoms E.g. H, Na, O, C etc…

◦A compound consists of atoms of two or more elements Chemically combined in proportion and

held together by a bond E.g. H2O, NaCl, CO2 etc…

MixturesMixtures

Two or more substances are physically mixed, not chemically combined.◦Homogeneous mixture or solution

Has a uniform composition E.g. air contains oxygen and nitrogen

◦Heterogeneous mixture Does not have a uniform composition E.g. raisins in cookie

2.2 States and Properties of 2.2 States and Properties of MattersMatters

Solids have

• a definite shape.

• a definite volume.

• particles that are close together in a fixed arrangement.

• particles that move very slowly.

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2.2 States and Properties of 2.2 States and Properties of MattersMatters

Liquids have

• an indefinite shape, but a definite volume.

• the same shape as their container.

• particles that are close together, but mobile.

• particles that move slowly.

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2.2 States and Properties of 2.2 States and Properties of MattersMatters

Gases have

• an indefinite shape.

• an indefinite volume.

• the same shape and volume as their container.

• particles that are far apart.

• particles that move very fast.

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Summary of the States of Summary of the States of MatterMatter

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Table 2.1

2.2 States and Properties of 2.2 States and Properties of MattersMatters

Physical Properties: Characteristics that do not involve a change in a sample’s chemical makeup.

Chemical Properties: Characteristics that do involve a change in a sample’s chemical makeup.

2.2 States and Properties 2.2 States and Properties of Mattersof MattersPhysical change occurs when

matter changes its appearance but the composition stay the same

Chemical change takes place when the original substance is converted into one or more new substances◦Have different physical and chemical

properties

ExamplesExamples

Identify each as: 1) solid, 2) liquid, or 3) gas.

___ A. It has a definite volume, but takes the shape of

the container.__ B. Its particles are moving rapidly.__ C. It fills the volume of a container.__ D. It has particles in a fixed arrangement. __ E. It has particles close together that are

mobile.

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ExampleExample

What type of change, physical or chemical, takes place in the each of the following◦Water vapor condenses to form rain◦Cesium metal reacts explosively with

water◦Gold melts at 1064 oC◦Food is digested

EnergyEnergy

Energy

• makes objects move.

• makes things stop.

• is needed to “do work.”

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WorkWork

Work is done when

• you climb.

• you lift a bag of groceries.

• you ride a bicycle.

• you breathe.

• your heart pumps blood.

• water goes over a dam.

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Potential EnergyPotential Energy

Potential energy is• stored energy.

Examples are

• water behind a dam.

• a compressed spring.

• chemical bonds in gasoline, coal, or food.

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Kinetic EnergyKinetic Energy

Kinetic energy is the

• energy of motion.

Examples are

• swimming.

• water flowing over a dam.

• working out.

• burning gasoline.

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Units for Measuring Energy or Units for Measuring Energy or HeatHeat

Heat is measured in joules or calories.

4.184 Joules (J) = 1 calorie (cal)

1 kJ = 1000 J

1 kilocalorie (kcal) = 1000 calories

(cal)

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Examples of Energy In JoulesExamples of Energy In Joules

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ExamplesExamples

Identify the energy as potential or kinetic.

A. RollerbladingB. a peanut butter and jelly sandwichC. mowing the lawnD. gasoline in the gas tank

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ExampleExampleThe energy needed to keep a 75-watt light

bulb burning for 1.0h is 270KJ. Calculate the energy required to keep the light bulb burning for 3.0 h in each of the following energy units:a. Joules b.

kilocalories

2.32.3 Temperature ConversionTemperature Conversion• Temperature is a measure of how hot or cold

an object is compared to another object.• indicates that heat flows from the object with a

higher temperature to the object with a lower temperature.

• is measured using a thermometer.

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Temperature ScalesTemperature Scales

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• are Fahrenheit, Celsius, and Kelvin.

• have reference points for the boiling and freezing points of water.

ExamplesExamples

A. What is the temperature of freezing water? 1) 0 °F 2) 0 °C 3) 0 K

B. What is the temperature of boiling water? 1) 100 °F 2) 32 °F 3) 373 K

C. How many Celsius units are between the boiling and freezing points of water?

1) 100 2) 180 3) 273

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Fahrenheit FormulaFahrenheit Formula

• On the Fahrenheit scale, there are 180 °F between the freezing and boiling points; on the Celsius scale there are 100 °C.

180 °F = 9 °F = 1.8 °F 100 °C 5 °C 1 °C

• In the formula for the Fahrenheit temperature, adding 32 ° adjusts the zero point of water from 0 °C to 32 °F.

TF = 9/5 TC + 32

orTF = 1.8 TC + 32

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Celsius FormulaCelsius Formula

• TC is obtained by rearranging the equation for TF.

TF = 1.8TC + 32 °

• Subtract 32 ° from both sides.

TF - 32 ° = 1.8 TC ( + 32 ° – 32 °)

TF - 32 ° = 1.8 TC

• Divide by 1.8 = °F - 32 ° = 1.8 TC

1.8 1.8

TF - 32 ° = TC

1.8

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Solving A Temperature Solving A Temperature ProblemProblem

A person with hypothermia has abody temperature of 34.8 °C. What is that temperature in °F?

TF = 1.8 TC + 32

TF = 1.8 (34.8 °C) + 32 ° exact 3 SFs exact

= 62.6 + 32 ° (addition)

= 94.6 °F tenth’s

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Kelvin Temperature ScaleKelvin Temperature Scale

The Kelvin temperature scale• has 100 units between the freezing and boiling

points of water. 100 K = 100 °C or 1 K = 1 °C

• is obtained by adding 273 to the Celsius temperature.

TK = TC + 273

• contains the lowest possible temperature, absolute zero (0 K).

0 K = –273 °C

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Learning CheckLearning Check

The normal body temperature of a chickadee is 105.8 °F. What is that temperature on the Celsius scale?

1) 73.8 °C 2) 58.8 °C3) 41.0 °C

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ExamplesExamples

A pepperoni pizza is baked at 455 °F. What temperature is needed on the Celsius scale?

1) 423 °C2) 235 °C3) 221 °C

and in the Kelvin scale?

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2.5 Specific Heat2.5 Specific Heat

Specific heat

• is different for different substances.

• is the amount of heat that raises the temperature of 1 g of a substance by 1 °C.

• in the SI system has units of J/g °C.

• in the metric system has units of cal/g °C.

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Examples of Specific HeatsExamples of Specific Heats

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TABLE 2.7

ExamplesExamples

A. When ocean water cools, the surrounding air 1) cools. 2) warms. 3) stays the same.

B. 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.

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Heat = mass x specific heat x ΔT

Heat EquationHeat Equation

The amount of heat lost or gained by a substance iscalculated from the

•mass of substance (g).•temperature change (ΔT).•specific heat of the substance (J/g °C).

This is expressed as the heat equation.

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ExampleExample

How many kJ are needed to raise the temperature of 325 g of water from 15.0 °C to 77.0 °C?

1) 20.4 kJ2) 77.7 kJ3) 84.3 kJ

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ExampleExampleCalculate the Calories of one stalk of celery

that produces energy to heat 505 g of water from 25.2oC to 35.7oC

ExampleExample

What is the specific heat if 24.8 g of a metal

absorbs 275 J of energy and the temperature rises

from 20.2 °C to 24.5 °C?

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ExamplesExamples

How many kcal are absorbed by ocean water if 3 x 1018 L of water in the Caribbean has an increase of 1 °C. Assume the specific heat of ocean water is the same as water. Assume the density of ocean water is 1.0 g/mL.1) 3 x 1015 kcal2) 3 x 1018 kcal3) 3 x 1021 kcal

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2.6 Energy and Nutrition2.6 Energy and NutritionEnergy Values in

Nutrition◦1 Cal = 1 kcal =

1000 cal◦1 Cal = 4.184 kJ

= 4184 JThe number of

Calories in a food is determined by using an apparatus called a calorimeter

Energy Value for FoodsEnergy Value for FoodsThe energy (caloric) value of food are the

kilocalories or kilojoules obtained from burning 1g of carbohydrate, fat or protein.

Table 2.8 Typical Energy (caloric) values of the three food types

ExamplesExamplesAt a fast-food restaurant, a hamburger

contains 37g of carbohydrate, 19g of fat, and 24g of protein. What is the total energy content in kilocalories? Round off the answer to the tenth place

ExampleExampleUsing the energy values for food (Table 2.8),

determine the grams of fat in one avocado that has 405 kcal, 13g of carbohydrate and 5g of protein. Round the answer to the tenth place

2.7 Changes of State2.7 Changes of State

A substance

• is melting while it changes from a solid to a liquid.

• is freezing while it changes from a liquid to a solid.

• such as water has a freezing (melting) point of 0 °C.

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Calculations Using Heat of FusionCalculations Using Heat of Fusion

The heat of fusion

• is the amount of heat released when 1 gram of liquid freezes (at its freezing point).

• is the amount of heat needed to melt 1 gram of a solid (at its melting point).

• for water (at 0 °C) is 80. cal 1 g water

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Calculation Using Heat of Calculation Using Heat of FusionFusion

The heat needed to freeze (or melt) a specific mass of water (or ice) is calculated using the heat of fusion.

Heat = g water (ice) x 80. cal 1 g water (ice)

Example: How much heat in cal is needed to melt 15. g of ice?

15. g ice x 80. cal = 1200 cal 1 g ice

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ExamplesExamples

A. How many calories are needed to melt 5.0 g of ice at 0 °C?1) 80. cal 2) 4.0 x 102 cal 3) 0 cal

B. How many calories are released when 25 g of water at 0 °C freezes?

1) 80. cal 2) 0 cal 3) 2.0 x 103 cal

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SublimationSublimation

Sublimation• occurs when particles

change directly from solid to a gas.

• is typical of dry ice, which sublimes at -78 C.

• takes place in frost-free refrigerators.

• is used to prepare freeze-dried foods for long-term storage.

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Copyright © 2009 by Pearson Education, Inc.

Evaporation and Evaporation and CondensationCondensationWater

• evaporates when molecules on the surface gain sufficient energy to form a gas.

• condenses when gas molecules lose energy and form a liquid.

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BoilingBoiling

At boiling,

• all the water molecules acquire enough energy to form a gas.

• bubbles appear throughout the liquid.

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Heat of VaporizationHeat of Vaporization

The heat of vaporization is the amount of heat

• absorbed to vaporize 1 g of a liquid to gas at the boiling point.

• released when 1 g of a gas condenses to liquid at the boiling point.

Boiling Point of Water = 100 °C

Heat of Vaporization or condense (water) = 540 cal

1 g water

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ExamplesExamples

How many kilocalories (kcal) are released when 50.0 g of steam from a volcano condenses at 100 °C?

1) 27 kcal2) 540 kcal 3) 2700 kcal

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Summary of Changes of StateSummary of Changes of State

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Heating CurveHeating Curve

A heating curve • illustrates the

changes of state as a solid is heated.

• uses sloped lines to show an increase in temperature.

• uses plateaus (flat lines) to indicate a change of state.

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ExamplesExamples

A. A flat line on a heating curve represents 1) a temperature change. 2) a constant temperature. 3) a change of state.

B. A sloped line on a heating curve represents 1) a temperature change. 2) a constant temperature. 3) a change of state.

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Cooling CurveCooling Curve

A cooling curve • illustrates the

changes of state as a gas is cooled.

• uses sloped lines to indicate a decrease in temperature.

• uses plateaus (flat lines) to indicate a change of state.

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ExamplesExamples

Use the cooling curve for water to answer each.

A. Water condenses at a temperature of 1) 0 °C. 2) 50 °C. 3) 100 °C.

B. At a temperature of 0 °C, liquid water 1) freezes.2) melts. 3) changes to a gas.

C. At 40 °C, water is a 1) solid. 2) liquid. 3) gas.

D. When water freezes, heat is 1) removed. 2) added.

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Combined Heat CalculationsCombined Heat Calculations

To reduce a fever, an infant is packed in 250. g of ice. If the ice (at 0 °C) melts and warms to body temperature (37.0 °C), how many calories are removed from the body?

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ExampleExampleHow many kilojoules of heat are released

when 75g of steam at 100. oC is converted to ice at 0oC?