CHAPTER 3Matter & Energy

MATTERMatter is defined as anything that has

mass and takes up space

We can describe matter usingVolume (may change under different

conditions)Weight (changes with gravity)Mass (not affected by temperature,

pressure, or gravity)

MATTER Matter can be classified into 3 groups:

1) Elements2) Compounds3) Mixtures

ELEMENTS

COMPOUNDSCompounds are made of 2 or more

elements that are chemically combinedThe elements in a compound are

combined in definite proportions by massH20 (2 hydrogens bond to 1 oxygen)NaCl (1 sodium bonds to 1 chlorine)

More common than pure elements (elements tend to be very chemically

reactive and combine with other elements to form compounds)

COMPOUNDS The chemical and physical properties of a

compound are different from the chemical and physical properties of the elements that make up the compound

Can be formed from simpler substances by chemical change, and they can be broken down into simpler substances by chemical change

***Elements and compounds are considered to be pure substances**

MIXTURESMixtures are made of 2 or more

substances that keep their own properties while combined

A mixture can form when:An element is mixed with one or more other

elementsA compound is mixed with one or more other

compoundsOne or more elements are mixed with one or

more compounds

MIXTURESMixtures are different than elements or

compounds

Mixtures have the properties of its constituents (the things that make it up), while elements/compounds have one set of properties

The composition of a mixture varies while the composition of an element/compound is fixed

MIXTURES Mixtures can be homogenous or

heterogeneous

Homogenous mixtures have uniform characteristics throughout i.e. a sample from one part of the mixture is the same as a

sample from another part Solutions Kool-Aid

Heterogeneous mixtures have different compositions i.e. a sample from one part of the mixture has a different composition than a sample from another part

Suspensions M&M’s, Munchies

http://blog.austinkids.org/2011/03/17/st-pattys-day-experiment/

Matter

Pure Substances

Mixtures

Elements

CompoundsHomogeneo

usHeterogeneo

us

PROPERTIES OF MATTER Properties are a set of characteristics that

identify a substanceExtensive properties- depends on the amount of

a substance (Volume, weight, mass) Intensive properties- depends on the nature of a

substance 1. Used to identify substances 2. Boiling point, melting point, malleability, crystalline

shape 3. Density- the mass of a substance that occupies one

unit of volume (the amount of “stuff” crammed into something) The ratio of mass to volume Density of a liquid or solid will change slightly with

changes in temperature and pressure (changes in volume are too slight to be noticeable)

Changes in the densities of gases can be very large

PROPERTIES OF MATTER Metals have interesting properties that

identify them as metals, but not necessarily a specific metal Ductile = the metal can be made into wires Malleable = the metal can be hammered into a

shape Luster = shine Conduct electricity

PHYSICAL PROPERTIES Extensive properties and intensive

properties are considered to be physical properties

Physical properties can be observed without the production of a new substance Examples: color, hardness, odor, taste, density,

melting & boiling points, electrical conductivity, and the other properties of metals discusses previously

A physical change is a change that occurs that DOES NOT alter the identity of the substances AND can be reversed without a chemical reaction

Evidence of a physical change: 1) phase changes (melting, boiling, etc)2) grinding3) changing its shape4) magnetizing a substance

CHEMICAL PROPERTIES A chemical property describes how a substance

reacts, or fails to react, with other substances to produce a new substance

A chemical change results in a new substance being formed Chemical changes are only reversible by another

chemical reaction (if they are reversible at all).

Evidence of a chemical change: 1) dramatic color change (not necessarily reliable) 2) release of gas or bubbles (effervesces!!!) 3) energy changes (heat or cold) 4) precipitate formation (an insoluble solid “falls out” of a

liquid)*** usually a chemical change will have 2 or more of these

occurring

Chemical Reactions Chemical reactions are written so that the

reactants are on the left of the arrow and products are to the right of the reaction arrow

6CO2 + 6H2O C6H12O6 + 6O2

In a chemical reaction, atoms are not created or destroyed, they are rearranged!! That is why you can balance chemical reactions

Reactants Products

CONSERVATION OF MASS The Law of Conservation of Mass states that…

Matter is neither created nor destroyedTHERE IS NO NEW MATTER

During both physical and chemical changes, the total amount of matter remains the same

(Nuclear reactions do not always follow this though)

Butane + Oxygen Carbon Dioxide + Water

58 g + 208 g 176 g + 90 g

266 g 266 g

Practice Problem

If a 35.0 gram sample of IRG is made up of 29.6 grams of KIX and 5.4 grams of GIRF, how many grams of KIX would there be in 245.3 grams of IRG?

Set this up as a factor-label method problem!!!

ENERGY Energy is defined as, “the capacity to do

work” The Law of Conservation of Energy states…

Energy can neither be created nor destroyed

-In chemistry we discuss Chemical EnergyChemical Energy is the energy that is involved with potential chemical changes (reactions) in chemical systems

CHEMICAL ENERGY Potential Energy – stored energy

Examples: gravitational PE, nuclear, chemical (located in the bonds between atoms)

Kinetic Energy - the energy of movement Examples: electromagnetic, mechanical (sound

& earthquakes), heat (internal motion of particles)

Microscopic particles contain potential and kinetic energy as heat energy, chemical energy, electrical energy (radiant energy) and sound energy

CHEMICAL ENERGY Every substance has a certain amount of

potential energy based on the number and strength of the chemical bonds

All chemical reactions have energy changes because they involve forming or breaking bonds between atoms!!!

Basic Units of Energy - joule

The SI unit of energy is the joule (J)

This unit was named after James P. Joule who determined that work (or energy) can be measured according to how much heat it could produce

Basic Units of Energy- calorie

A second unit of energy is the calorie

calorie (cal) – (heat calorie) the amount of heat required to raise the temperature of 1 gram of water 1°C. (notice the lower case c!!!!)

The nutritional or “food” Calorie (Cal) is 1000 calories(notice the capital C!!)

1 Cal = 1000 cal

The kilocalorie (kcal) equals 1000 calories

1 calorie (cal) = 4.184 J** a calorie is actually larger than a joule!!

Measuring Energy Changes

A calorimeter can be used to measure energy changes A calorimeter contains an inner container that

contains water (this is where the reaction takes place)

Energy changes from the reaction can be measured by the temperature change of the water (whether it absorbs or loses heat)

The heat required to raise the temperature of 1 gram of a substance 1 degree Celsius is the substance’s specific heat

The specific heat of water is 1 cal/g•C° (or 4.184 J/g•C°)

http://olc.spsd.sk.ca/de/physics20/heat/calorimeter.htm

Energy Reminder!!

Because of the Law of Conservation of Energy, whenever energy changes occur in a chemical reaction, the total amount of the energy remains the same.

In order for a reaction to get started, many chemical reactions require a specific amount of activation energy

Activation energy is the energy added to a reaction to get it started (can be heat or electrical)

Exothermic Reactions

If the reactants have more potential energy than the products:

http://www.kentchemistry.com/links/Kinetics/PEDiagrams.htm

The reaction is EXOTHERMIC

Energy is lost

(converted to heat energy and released)

What is an example of an exothermic reaction you’ve

experienced?

•The PE of the reactants is converted to heat (KE) and released (feels warm) which leaves the products with lower PE than the reactants

Endothermic Reactions If the products have more potential

energy than the reactants…

http://www.kentchemistry.com/links/Kinetics/PEDiagrams.htm

The reaction is

ENDOTHERMIC

The potential energy increases

Heat energy is absorbed and the reaction feels “cold”

What is an example of an endothermic reaction that

you’ve experienced?

•Heat energy (KE) is absorbed (feels cold) and becomes stored PE in the products leaving the products with a higher PE than the reactants

Temperature

Energy is usually measured by measuring temperature changes Heat and temperature are not the same thing!!!

Heat refers to the exchange of thermal energy Temperature is a measure of the thermal energy of

matter (atoms)

Temperature Scales

1)Fahrenheit (° F) – not used in Science2)Celsius (° C) – metric standard3)Kelvin (K) – used because negative Celsius temps can result in unrealistic answers (i.e. negative volume). 0 K = -273°C (absolute zero)

CONVERSIONS BETWEEN TEMP SCALES

°C = K – 273 or 5/9(°F - 32)

°F = (9/5 • °C) + 32

K = °C + 273°

SPECIFIC HEAT Specific heat – the amount of heat energy (J) needed

to raise the temperature of 1 g. of a substance by 1 °C

Many substances have different specific heat values called the specific heat constant (c) Units:

You can find these values in Table 3.4 on p.71

The greater the specific heat constant (c), the more heat is needed to change the temperature of 1g by 1 °C

J g • ° C

SPECIFIC HEAT FORMULA

q = heat absorbed or released (J)

m = mass of substance (g)

c = specific heat constant (p.71)

Δt = temperature change (° C)

The quantity of heat released or absorbed depends on 3 things: 1. mass or quantity of substance 2. what substance is (c) 3. how much the temp changes

q = m • c • Δt

J g • ° C

SPECIFIC HEAT CALCULATION

How much heat would have to be added to 5000.0 g of water to change its temp from 20.0 °C to 80.0 °C?

q = m • c • Δt

SPECIFIC HEAT CALCULATION

How much heat would have to be added to 5000.0 g of water to change its temp from 20.0 °C to 80.0 °C?

q = m • c • Δt

q = ??

m = 5000.0 g

C = 4.184

Δt = 80.0 °C – 20.0 °C = 60.0 °C

J g • ° C

q = (5000.0g)(4.184 )(60.0 °C )J g • ° C

q = 1,255,200 J

q = 1,260,000 J

Naming Phase Changes Review Solid Liquid = Melting Liquid Solid = Freezing

Liquid Gas = Boiling (or evaporation) Gas Liquid = Condensing Solid Gas = Sublimation Gas Solid = Deposition