MATTER

What is Matter?

• Must have mass

• Occupies space (has volume)

Density

• Because all matter has mass and volume, all matter has density.

D = M

V

Law of Conservation of Matter

• Matter cannot be created or destroyed

• When 400 grams of wood are burned only about 30 grams of ash remain.

• What happened to the missing matter?

• During a chemical reaction matter is conserved

States of Matter

• Matter can exist in three basic states:– Solid (s)– Liquid (l)– Gas (g)

A Fourth State of Matter?

• Plasma– Occurs when the

atoms of matter have been stripped of their electrons.

Basic Types of Matter

Pure Substances

Elements

Elements

• Made up of atoms of one specific type

• Cannot be broken down further by a chemical reaction

• Have specific physical and chemical properties.

The Periodic Table of Elements

• Each element has its own atomic number

• Atomic number tells us the number of protons in that elements nucleus.

Examples of Elements

• What elements can you find in the classroom now?

• What elements are important in the food we eat?

Where Do We Get Elements?

• Obtained in nature in their pure state.

• Some elements are combined together in compounds and are separated out through chemical reactions.

• Some elements don’t exist naturally, and are man-made through nuclear bombardment reactions

Elements and Their Symbols

• Elements have either one or two letters when they are written as a symbol.

• If two letters, the second letter is written lower case. (Ex: Write Cu for copper not CU)

• Usually it is pretty easy to match an element to its symbol (Ex: Neon = Ne)

• Sometimes the symbol is derived from a Latin name for the element.– Ex: Au = gold (comes from Latin word aurum)

Tricky Elements

• Here are some of the harder elements to match their name to their symbol.

– Na = sodium– K = potassium– Fe = iron– Cu = copper – Sb = antimony– Sn = tin– Pb = lead– Hg = mercury– W = tungstun– Au = gold– Ag = silver

Phases and Elements• Under standard conditions,

most elements exist as solids.

• The Gases:

– H, He, Ne, Ar, Kr, Xe, Rn, O, F, Cl

• The Liquids:

– Hg, Br

Diatomic Elements

• Two atoms of the same element joined

• Bromine, Iodine, Nitrogen, Chlorine, Hydrogen, Oxygen, Fluorine

• BrINClHOF

Compounds

Compounds

• Most elements in nature do not exist in their pure state, but in compounds.

Compound = Two or more different elements

bonded together chemically

Ex: C6H12O6, MgCl2, CH4, NO2, CO2, NaCl

Law of Definite Proportions

• Elements in compounds have definite fixed

proportions by mass.

Getting New Properties

• Compounds have different chemical and physical properties than the elements that make them up.

- Sodium (Na)

Toxic, reactive metal that will explode in water

- Chlorine (Cl)

Brown toxic gas that was used as a chemical weapon to kill people in WWI

Sodium Chloride (NaCl)

White edible crystal, and salt for my french fries!!

Focus Question

• What are the differences between elements and compounds?

Mixtures

Mixtures

• Two or more pure substances physically mixed together.

• Can be two or more elements, compounds or both.

• Mixtures can be solid, liquid or gaseous.

Composition Not Fixed

• The composition of a mixture is NOT FIXED.

• It can vary depending on how much of each component is added.

Retaining Their Properties

• The substances in mixtures retain their own individual properties.

• Ex: Iron filings (Fe) and sand (SiO2)

• Even if mixed the iron retains its magnetic properties

Types of Mixtures

Homogeneous Mixtures

• Components are distributed uniformly at the molecular level.

• All true solutions are homogeneous.

• Aqueous = (aq)

Mean the substance is dissolved in water.

• Ex: NaCl (aq) means a salt water solution

Heterogeneous Mixtures

• Components are not uniformly distributed at the molecular level.

• Colloids and suspensions are included in this category.

How to Tell if a True Solution

• Tyndall Effect– When you have a suspension or a colloid the larger

particles will scatter a beam of light. In a true solution the light will not be scattered

Techniques to Separate Mixtures

Separating Mixtures

• Components of a mixture retain their own original properties, you can use these properties to separate the components.

• Ex: density, particle size differences, solubility differences, boiling point temperature, magnetic properties

Filtration

• Ex: Separate water and sand

• Technique cannot be used to separate components of solutions

Separatory Funnel

• Ex: Separate oil and water

Evaporation

• Ex: Separate sugar from water

Chromatography

• Separates different compounds dissolved in the same liquid.

• Use Chromatography Paper• Liquid runs up the paper

and components separate out along the length of the paper.

• Ex: Separate different types of chlorophyll

Distillation

• Separates two or more liquids, using differences in boiling point temp.

• Use distillation apparatus

• Ex: Separate alcohol and water

Focus Questions

• Describe the differences between a pure substance and a mixture.

• What is the difference between a homogenous mixture and a heterogeneous mixture?

• Describe several different methods to separate mixtures.

Physical vs. Chemical

Physical vs.. Chemical Properties

• Physical Properties– Can be observed without changing the

identity of the substance.– Ex: color, hardness, density, odor, malleability, ductility,

melting point temp.

• Chemical Properties– Relates to a chemical change that type of

matter will undergo. – Ex: flammable, explosive, corrosive, rusts, decomposes

Chemical vs. Physical Changes

• Physical Change– Does not alter the chemical composition,

substance still has the same chemical formula.– Ex: bend, cut, grind, dissolve, undergo a phase change

• Chemical Change– The chemical composition changes as a

reaction occurs. New products are formed.– Ex: cook food, sour milk, burn wood, explode dynamite

What Type of Change?

What Type of Change?

• H2O(s) → H2O(l)

• 2H2O (l) → 2H2(g) + O2(g)

• NaCl(s) → NaCl(aq)

• CH4(g) + O2(g) → CO2(g) + H2O(g)

Is it Chemical or Physical?

• Sometimes it is possible to determine if a chemical change has occurred using the naked eye.

• Things to look for:– Gas bubbles produced, color changes,

new odor produced, precipitate forming

– Precipitate: an insoluble solid that forms when two solutions are mixed.

Lab Write Up Format (Typed)

• Title (Caps)

• Objective: What did we set out to do?

• Materials Used (list all)

• Procedure Followed: (simple, numbered steps)

• Data Tables and Graphs: Make sure that all are labeled clearly and neatly with units indicated

• Conclusion: What did you conclude based upon your data?

• Discussion: What does this conclusion mean? Why did you arrive at it? Discuss the basic scientific concepts involved. Sources of Error: (list at least 3 things that could have affected your data)

Measuring Density (D = M/V)

• Mass– Measured on a

scale, often you need to use a weighing dish or a beaker.

• Volume– Can be measured

in different ways.– For solids the

method of water displacement is often used.

Phases of Matter and Phase Changes

Phase

• Depends on the strength of the forces of attraction between particles.

.

Solids

• Definite shape and volume. • Regular crystalline lattice structure. • Denser than other phases (exception is water!).• Difficult to compress.• Highest forces of attraction between particles.• Atoms vibrate in fixed positions

• Note: Amorphous solids include

glass, plastic, wax, and silly putty

Liquids

• Definite volume, but no definite shape

• Takes shape of container

• Hard to compress

• Particles can slide past each other

• Forces of attraction between particles still pretty high

Gases

• No definite shape or volume• Expands to completely fill

container• Lowest density• Density depends on pressure• Particles have very little

attraction to each other

• “Vapor” = term to describe a gaseous state of something that is normally liquid (Ex: water vapor)

Phases Applet

• http://www.harcourtschool.com/activity/states_of_matter/

• http://mutuslab.cs.uwindsor.ca/schurko/animations/phasescontainers/phasescontainer.html

Changes in Phase

GasCondensation Vaporization

(Boiling or Evaporating)

LiquidSolidification Melting (fusion)

Solid

Let’s Skip a Phase

• Sublimation– When a substance goes

directly from the solid phase to the gas phase.

– Happens with substances with weak intermolecular forces (they separate easily!)

– Ex: CO2(s) dry ice, Iodine

CO2(s) → CO2 (g)

Energy

• Energy = the capacity to do work or produce heat. It can be anything that causes matter to move or change direction.

• Ex: electrical, radiant, atomic, mechanical, magnetic, sound, chemical

Law of Conservation of Energy

• In any process energy isn’t created or destroyed, just transferred from one form to another

PE vs. KE

• Potential Energy = stored energy– Energy can be stored in

bonds between atoms and released during chemical rxns.

• Kinetic Energy = energy of motion– All atoms are moving and

vibrating unless at absolute zero

Heat Energy

• A form of energy that increases the random motion of particles

• Measured in Joules or calories.

Heat Flow

• Heat energy travels from an object of higher temp. to one of lower temp. until both reach the same temp.

Temperature

• A measure of the average kinetic energy of all the particles in a sample.

• It is not a form of energy

• But if you add heat energy or take it away, it causes the particles to move faster or slower and thus changes the temp.

Temperature Scales and

Conversions

K = ˚C + 273

Heat vs. Temperature

• Teacup vs. Bathtub

• Both at 25˚C

• Which one contains more heat energy?

• Which one has the greater average KE?

Exothermic vs. Endothermic

• All changes in matter are accompanied by changes in energy.

• Exothermic Change: A + B → C + D + energy– Energy is released– Energy “ex”its

• Endothermic Change: A + B + energy → C + D– Energy is absorbed– Energy “en”ters

Energy and Phase Changes

• Convert from s→l→g– You must add energy to overcome the

attractive forces between particles– Add energy = endothermic

• Convert from g→l→s– As the particles come closer together energy

is released– Release energy = exothermic

Phase Change Applet

• http://mutuslab.cs.uwindsor.ca/schurko/animations/waterphases/status_water.htm

Heating & Cooling Curves

• Graphically represents temp. changes as heat energy is added or taken away.

Interpreting the Graph• The slanted portions =

temp is changing– Single phase is heating

up or cooling down– KE is changing

• The flat portions = temp not changing– Substance undergoing

a phase change– PE is changing

Applet

• http://images.google.com/imgres?imgurl=http://www.kentchemistry.com/images/links/matter/aim10.14.jpg&imgrefurl=http://www.kentchemistry.com/links/Matter/HeatingCurve.htm&h=290&w=473&sz=24&hl=en&start=38&um=1&usg=__LbCYC0DjOmKIRT180A3leRnBPuU=&tbnid=pC9q8ROC-XslJM:&tbnh=79&tbnw=129&prev=/images%3Fq%3Dchemistry%2Bpotential%2Benergy%2Bkinetic%2Benergy%26start%3D20%26ndsp%3D20%26um%3D1%26hl%3Den%26safe%3Dactive%26sa%3DN&safe=on

Label This Graph

Heating Curve for Water

What is Melting Pt? Boiling Pt?

Heat Equations

• Calculates the energy involved when a substance changes in temperature or undergoes a phase change.

• When Temp. of Substance Changes • Q = mcΔT

• When Undergoing Phase Change • Q = mHf or Q = mHv

Specific Heat Capacity

• Every substance has a different ability to absorb heat when energy is applied.

– Some objects rise in temp. quickly when heated – Some objects need more heat to raise the same number

of degrees

• Specific Heat = Joules of heat needed to raise 1 gram of substance 1°C.

• Ex: Metal pot vs. water in it. Which gets hotter faster?

Change in Phase (Temp. Constant)

• Q = mHf Use when changing from solid to liquid (melting) or

liquid to solid (freezing)

Hf = heat of fusion (for water = 334J/g)

• Q = mHv Use when changing from liquid to gas (vaporization)

or gas to liquid (condensing)

Hv = heat of vaporization (for water = 2260J/g)

Calorimeters

• Instrument to determine the amount of heat lost or gained in a reaction by measuring changes in the temp. of water surrounding the system.

Q = mcΔT

Multi-step Heat Problems

• Need to use more than one of the heat equations and add up the total heat.

• Ex: Calculate the heat energy to raise 10 grams of water at -25°C to 80°C.– Draw a heating curve. Figure out # of steps.– 1.) Heat ice from -25° to 0° q = mcΔT– 2.) Melt ice to liquid at 0° q = mHf– 3.) Heat liquid water from 0° to 80° q = mcΔT

Heat Lost = Heat Gained

• When two objects of different temperatures are placed together in a closed system, heat will flow from the hotter to the colder object until they reach the same temperature.

• The total heat lost = total heat gained

mcΔT = mcΔT

• http://www.bbc.co.uk/schools/ks3bitesize/science/chemistry/physical_changes_4.shtml