chemistry: the science of matter. unit objectives classify matter according to its composition...

Chemistry: The Science of Matter

Unit Objectives

• Classify matter according to its composition• Distinguish among elements, compounds,

homogenous mixtures, and heterogeneous mixtures

• Relate the properties of matter to structure

Proper Lab Safety Procedures

Appropriate Dress for Lab

Goggles must be worn at all times

Long hair must be tied back

Close-toed shoes must be worn

Safety Equipment

Fire Blanket

Fire Extinguisher

Shower

Eye Wash Station

At your lab table…

• You may not have food or drink

• You must notify a teacher of any chemical spill…don’t touch it!

• You may not have bring any backpacks or other personal items other that your lab guide, a periodic table, a writing utensil, and a calculator

• Keep the area free from clutter…when your done with it, return it or dispose it!

• Spray and wipe down your lab tables.

• Don’t run, horse play, and always be aware of your surroundings.

Lab Measurement

Using the Metric System - Length• Using Rulers – always use the “cm” or “mm” side.• You can estimate out one decimal place beyond the

actual marking. • For each of the following, draw a downward arrow

pointing to the correct value and label the arrow. 1. Where is 4.6 cm?2. Where is 46 mm?3. Where is 6.85 cm?4. *Where is 35.5 mm?

Common Metric Units of Length

• Millimeter (mm): 1000 mm = 1 m• Centimeter (cm): 100 cm = 1 m• Meter• Kilometer (km) : 1 km = 1000 m

Using the Metric System - Mass

• Weighing a substance– All digital scales are in

grams and read out to two decimal places.

– ALWAYS put chemicals onto a weighing dish and NEVER directly on the scale.

– Find the mass of the weighing dish first, then press the “zero” button

Common Metric Units of Mass

• Milligram (mg): 1,000 mg = 1 g• Gram (g)• Kilogram (kg): 1 kg = 1000 g

Using the Metric System - Volume• Measured using

graduated cylinder, Erlenmeyer flasks, and beakers.

• Just like a ruler, You can estimate out one decimal place beyond the actual marking.

• Volume in our lab will be mostly measured in mL (milliliters)

• Always read the bottom of the meniscus.

Beaker Erlenmeyer Flask

Graduated Cylinders

Reading Volume

Common Metric Units of Volume

• Milliliter (mL) : 1,000 mL = 1 L• Centimeters cubed : 1 cm3 = 1 mL• Liter (L)

Using the Metric System - Temperature

• We ALWAYS use the unit Celcius (never Fahrenheit!!)

• Hint: Degrees Celcius will seem much smaller than Fahrenheit.

• Always make sure your digital thermometer is on the “°C” setting

Assignment 2 Part B – Complete in Groups

Accuracy and Precision

Accuracy• How close an experimental

measurement is to the actual/correct value

• If the correct measurement is 5.65 g, which actual value is the most accurate: 5.55 g, 5.63 g, 5.7

Precision• How close a series of

measurements are to one another

• Which series of measurements are the most precise?– 5.15 g, 5.55 g., 5.59 g.– 5.15 g., 5.99 g., 6.86 g.

% Error

Gives you a mathematic “gauge” as to how close (accurate) your results are to the accepted value

%Error actual experimental

actual100

Observation

Quantitative• An observation made that

involves a number.• Examples:

– The mass of the rock is 5 g.– The length of the block is 4

cm.– The chemical changed color 5

times.– It took 2 minutes for the

reaction to happen. – The temperature changed

from 20 °C to 30°C

Qualitative• An observation made that

does not involve a number.• Examples:

– The chemical reaction produced a gas.

– The color of the substance changed from clear to blue.

– The reaction got warmer.– Block A feels heavier than

Block B.

What is matter?

• Matter is anything that takes up space and has mass.– Where is matter? Everywhere! Can you name

some things around you that have matter? That don’t have matter?

• Mass is the measure of the amount of matter that an object contains– What has mass? Doesn’t have mass?

Properties of Matter

• Properties of matter – describe the characteristics and behavior of matter, including the changes that matter undergoes.

• Example (Figure 1.3, p. 6) – Iron is…

Melted at high temps

MagneticMalleable (bendable)Ductile (stretchable)

Able to rust

Macroscopic vs. Microscopic

Macroscopic• Matter that is large enough

to be seen• Examples?

Microscopic• Matter that can not be seen

without a microscope• Examples:

– Bacteria– Cells– Atoms– Subatomic particles - Protons,

Neutrons, Electrons

Pure substance vs. mixtures – The two categories of matter

Pure Substance– Substance – matter with the same fixed

composition and properties. – Any sample of pure matter is a substance.

• A pure substance is either a:– Compound– Element

Pure Substances

• Pure Substance that cannot be broken down into any other substances by chemical or physical means

Gold - element Manganese Dioxide - compound

Pure Substance

• Element– composed of identical atoms– Are found on the periodic table– EX: copper wire, aluminum foil

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Pure Substances

• Compound

– composed of 2 or more elements in a fixed ratio

– properties differ from those of individual elements

– Chemical bonds hold the elements together

– EX: table salt (NaCl)


Pure Substances - FYILaw of Definite CompositionLaw of Definite Composition

– A given compound always contains the same, fixed ratio of elements.


Two different compounds, each has a definite composition

Mixtures

• Mixture – a combination of two or more substances in which the basic identity of each substance is not changed

• Most of the matter you encounter every day is a mixture.

• Mixtures can either be:– Homogeneous – Heterogeneous

Mixtures Variable combination of two or more pure

substances. Each keep individual properties

Homogeneous- Evenly Mixed cannot see different parts. (Same)

Heterogeneous – Can see different parts (different)


Tyndall Effect• The scattering of light by particlesin a mixture

• http://www.youtube.com/watch?v=gheuYqQ6phE&feature=related

Homogeneous Mixtures

• A mixture that is evenly mixed in such a way that you cannot see its individual parts.

• A homogenous mixture is also known as a solution.

• Example: Apple Juice

Mixtures

Solution– homogeneous– very small particles– no Tyndall effect– particles don’t settle

– EX: – rubbing alcohol (ethyl alcohol and

water)– Air (nitrogen and oxygen)


Heterogeous Mixture

• A mixture in which you can distinguish the various components

• Types of homogenous mixtures:– Colloids– Suspensions

MixturesColloid

– heterogeneous– medium-sized particles– Tyndall effect– particles don’t settle– Particles scatter light– EX:

• Milk• Clouds• Smoke• mayo


Mixtures

Suspension– heterogeneous– large particles– Tyndall effect– particles settle– EX: • fresh-squeezed

lemonade• Sand in water


Classification of Matter

MATTER(gas. Liquid,

solid, plasma)

PURESUBSTANCES MIXTURES

HETEROGENEOUSMIXTURE

HOMOGENEOUSMIXTURESELEMENTSCOMPOUNDS

Separated by

physical means into

Separated by

chemical means into

Kotz & Treichel, Chemistry & Chemical Reactivity, 3rd Edition , 1996, page 31

MatterMatter

SubstanceDefinite composition

(homogeneous)

SubstanceDefinite composition

(homogeneous)

Element(Examples: iron, sulfur,

carbon, hydrogen,oxygen, silver)

Element(Examples: iron, sulfur,

carbon, hydrogen,oxygen, silver)

Mixture ofSubstances

Variable composition

Mixture ofSubstances

Variable composition

Compound(Examples: water.

iron (II) sulfide, methane,Aluminum silicate)

Compound(Examples: water.

iron (II) sulfide, methane,Aluminum silicate)

Homogeneous mixtureUniform throughout,also called a solution

(Examples: air, tap water,gold alloy)

Homogeneous mixtureUniform throughout,also called a solution

(Examples: air, tap water,gold alloy)

Heterogeneous mixtureNonuniform

distinct phases(Examples: soup, concrete, granite)

Heterogeneous mixtureNonuniform

distinct phases(Examples: soup, concrete, granite)

Chemicallyseparable

Physicallyseparable


Elements

only one kindof atom; atomsare bonded itthe element

is diatomic orpolyatomic

Compounds

two ormore kinds

of atomsthat arebonded

substancewith

definitemakeup

andproperties

Mixtures

two or moresubstances

that arephysically

mixed

two ormore

kinds ofand

Both elements and compounds have a definite makeup and definite properties.

Packard, Jacobs, Marshall, Chemistry Pearson AGS Globe, page (Figure 2.4.1)


uniformproperties?

fixedcomposition?

chemicallydecomposable?

no

no

no

yes

hetero-geneousmixture

solution

element

compound

http://antoine.frostburg.edu/chem/senese/101/matter/slides/sld003.htm

Matter Flowchart

Examples:

– graphite

– pepper

– sugar (sucrose)

– paint

– soda


element

hetero. mixture

compound

solution homo. mixture

hetero. mixture

The Composition of Air

AirAir

NitrogenNitrogen

OxygenOxygenHeliumHelium

Watervapor

WatervaporNeonNeon

Carbondioxide

Carbondioxide ArgonArgon

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 34

Top Ten Elements in the Universe

Top Ten Elements in the Universe

Percent Element (by atoms) 1. Hydrogen 73.92. Helium 24.03. Oxygen 1.14. Carbon 0.465. Neon 0.136. Iron 0.117. Nitrogen 0.0978. Silicon 0.0659. Magnesium 0.05810. Sulfur 0.044

A typical spiral galaxy(Milky Way is a spiral galaxy)

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 26

Physical Properties & Changes

• A physical property is a characteristic of a sample of matter that can be observed or measured without any change in its identity.– Examples/key words – solubility, melting point, boiling

point, color, density, electrical conductivity, and physical state (solid, liquid, gas) (they are all generally adjectives)

• A physical change is a change in matter that does not involve a change in the identity of the substance– Examples/key words – boiling, freezing, melting,

evaporating, dissolving, separating, and crystallizing (they are all generally verbs)

Chemical Properties & Changes

• A chemical property is the ability of a substance to undergo a change such as reacting with other substances or decomposing. (generally adjectives)

• A chemical change is the change of one or more substances into other substances. This is also referred to as a chemical reaction.– Examples & key words: decompose, explode, rust,

oxidize, corrode, tarnish, ferment, burn, react, changes color, bubbles/fizzes, or rot (generally verbs)

Physical or Chemical Property? -Identify each property below as either a chemical or physical property of a

substance:

Questions: 1. Aluminum bends easily2. Copper sulfate dissolves in

water3. Magnesium burns in air4. Gold jewelry is unaffected

by perspiration5. Basking soda is a white

powder6. Fluorine is a highly reactive

element

Answers: • 1.• 2.• 3.• 4.

• 5.

• 6.

Launch Lab – Why is the mass different? – p. 3

• Record your observations.• Answer the “Analysis” questions on p. 3• Class Discussion: How can three objects with

the same volume have different masses?

Density

• Density is the amount of matter (mass) contained in a unit of volume.

• Examples:– A foam cup has less

density than a stone because it has more mass in a identical-sized sample.

Density formula:D = m

v

Density Pyramid:

Density Practice Set #1 – Solving for density

• 1. Calculate the density of an object that has a mass of 2.53 g and a volume of 4.54 mL

• 2. Calculate the density of an object that has a mass of 16.0 g and a volume of 25.3 mL.

• *3. Calculate the density of an object that has a mass of 3.01 g and a volume of 5.08 cm3

• Which substance above has the greatest density?• …Least density?

Density Practice Set #2 – solving for mass or volume

• 1. What is the mass of a sample that has a density of 2.0 g/mL and a volume of 4.6 mL?

• 2. What is the volume of a sample that has a density of 0.23 g/mL and a mass of 2.5 g?

Fill in the missing cells on the chart below: Mass (m) Volume (v) Density (d)

1.4 g 7.6 mL

5.1 mL 2.1 g/mL

24.2 g 7.5 g/mL

States of Matter

• The states of matter on Earth are solid, liquid gas – these are physical properties of a substance

• A change of state is the temperature at which a substance changes from one state to another. – Water freezes/melts at 0°C– Table salt (sodium chloride) freezes/melts at 804°C– Oxygen freezes/melts at -218°C

• A change in state is a physical change

Changing States of Matter

Law of Conservation of Mass

• While atoms of substances do and can change, they never disappear or appear from no where.

• The law of conservation of mass (or matter) states that in a chemical change, matter is neither created nor destroyed.

Before = After … ALWAYS!!

Endothermic vs. Exothermic Reactions

• All chemical changes also involve some sort of energy change.

• Energy is either absorbed (endothermic reaction) or released (exothermic reaction) as a chemical reaction is taking place.

C Endothermic Reaction

chemistry: the science of matter. unit objectives classify matter according to its composition...

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