WELCOME TO CHEMISTRY!

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Meet the Elements!

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WHAT IS CHEMISTRY?

• Chemistry is the study of matter and the changes that it undergoes.

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WHAT IS MATTER?

• Matter is anything that has mass and volume.

• Mass refers to the measure of the amount of material in an object ( measures the resistance to an object being moved).

• Volume refers to the amount of space an object occupies.

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MATTER HAS 4 STATES

• Solid

• Liquid

• Gas

• Plasma

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3 STATES OF MATTER

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SOLIDS

• State of matter in which the particles are closely packed.

• Solids have definite shape and definite volume

• Solids essentially cannot be compressed.

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LIQUIDS

• Particles are arranged so they can slide past one another.

• Liquids take the shape of their container and have definite volume.

• Liquids essentially cannot be compressed.

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GASES

• Particles are spread out widely.

• Gases have neither definite shape nor definite volume.

• Are very compressible.

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PLASMA

• An electrically neutral gas of ions (charged particles) and electrons (negatively charged particles)…VERY HIGH ENERGY!! Thought to be found in stars.

• Is present when nuclear fusion occurs..requires a temp of 100,000,000 0C.

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The smallest unit of matter is the atom

• are the smallest unit of matter capable of existing by themselves.

• 2 or more atoms chemically bonded together are referred to as a molecule.

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MODELS OF THE ATOM

orbits

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Atoms are composed of the following subatomic particles.• Protons- Positively charged particles located in

the nucleus of the atom.

• Neutrons- Non-charged nuclear particles located in the nucleus of the atom.

• Electrons- Negatively charged particles that move around the nucleus located in the electron cloud

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Matter with only one type of atom are called elements.

• Each element has a unique number of protons in the nucleus of its atoms. This is what defines that particular element.

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ELEMENTS (CONT).

• Chemists use elemental symbols as a shorthand way of representing elements.

• These symbols consist of 1, 2, or 3 letters (only the 1st is uppercase).

• Many elemental symbols show the Latin origin of the element’s name.

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Periodic Table of Elements

• 114 elements • Vertical column =

groups• Horizontal rows =

periods• Traits are organized

by similar

properties

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ALL MATTER CAN BE CLASSIFIED AS EITHER A

SUBSTANCE OR A MIXTURE• A substance is matter that has a fixed

composition and distinct properties.

• A mixture is a combination of 2 or more substances, each of which retains it’s properties.

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Substances can be either elements or compounds.

• An element is a substance in which all atoms have the same number of protons.

• Elements cannot be decomposed into simpler substances.

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Compounds

• Compounds are substances in which 2 or more elements chemically combine to form a new substance with new properties.

• Compounds can be decomposed into simpler substances.

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COMPOUNDS CAN BE CHEMICALLY BROKEN DOWN

• This can be accomplished through the use of heat energy (thermal decomposition) or electric current (electrolysis).

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EVERY SUBSTANCE HAS A UNIQUE SET OF PROPERTIES

• Properties are characteristics that allow us to recognize and distinguish a substance from other substances.

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PHYSICAL PROPERTIES

• Physical properties are properties that can be observed or measured without changing the identity and composition of the matter.

• Physical properties include color, odor, density, melting pt., boiling pt., malleability, ductility, hardness, etc.

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CHEMICAL PROPERTIES

• Chemical properties describe the way a substance may change or react to form other substances.

• To observe this property you must carry out a chemical change.

• An example would be burning in the presence of oxygen.

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MATTER UNDERGOES CHANGE

• This may be physical change or chemical change.

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PHYSICAL CHANGES

• In physical changes, the substance changes its physical appearance but not its composition.

• An example would be changing states.– Evaporation of water into water vapor

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CHEMICAL CHANGES

• In a chemical change, a substance is changed into 1 or more new substances with new properties.

• Chemical change involves rearrangement of atoms.– Aluminum rusting, fireworks exploding

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PHYSICAL CHANGE

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INTENSIVE PROPERTIES

• Are independent of sample size.

• Examples include density, temperature, melting point, etc. (1 gal of water or 1 cup of water: both have same boil point.)

• Can be used to identify a substance.

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EXTENSIVE PROPERTIES

• Are properties that depend on sample size.

• Examples include mass, volume, and length

• 1 gal of water will have a greater mass than 1 cup of water.

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Mixtures can be either homogeneous or heterogeneous.

• Homogeneous mixtures are uniform (have the same composition) throughout. Homogeneous mixtures are also called solutions.

Ex: salt, air, water• Heterogeneous mixtures are non-

uniform throughout.Ex: sand, wood, rock

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MIXTURES CAN BE PHYSICALLY SEPARATED

• Filtration

• Distillation

• Chromatography

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FILTRATION

• In filtration, the mixture is separated based on differences in particle size.

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DISTILLATION

• Distillation separates components based on differences in boiling point.

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DISTILLATION APPARATUS

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CHROMATOGRAPHY

• Separates mixture components based on differences in attraction to a surface.

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Name Those Molecules

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CHEMICAL CHANGE

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Measurement

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SCIENTIFIC DATA FALLS INTO 2 CATEGORIES

• Qualitative data

• Quantitative data

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Qualitative

• Qualitative data consists of descriptive terms. There is no use of numbers.

• Determines the presence or absence of a particular substance in a mix.

– Ex: water is a liquid

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Quantitative

• Quantitative data consists of measurements (numbers)

• Determines the amount of substance that is in a sample. – Water has a boiling point of 100degrees

Celsius

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Scientific Method

• Hypothesis : possible explanation of an observation.

• Theory: set of tested hypothesis that give an explanation of a natural phenomena

• Law: concise statement or explanation

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Apples, Volkswagens, Pigeons or Metric?

UNITS ARE EVERYTHING!

Chemistry is in the details. The slightest change or mistake can mean a world of

difference. This IS going to drive you nuts this year, and you will hate chemistry, me,

and the universe, but we go’tta do it!

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MEASUREMENT

• At one time measurement was inconsistent and therefore, unreliable.

• Scientists need to be able to repeat each other’s experiments to verify/modify scientific knowledge.

• To assist in this process, all scientist use the SI system of measurement. (AKA your new best friend!)

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Base Units

mass gram g length meter m time seconds s electric current ampere A temperature Kelvin K light intensity candela cd amt. Subst. mole mol l

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MEASUREMENT (CONT)

• Scientist often use prefixes in conjunction with the basic unit of measure.

• The indicate decimal fractions or multiples of various units. – Ex: milli = 10-3

1 milligram = 1 mg = 10 -3 grams

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Like what is a centimeter again?

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MEASUREMENT (CONT)Prefix Symbol Exp. Represent.Giga G 109

mega- M 106

kilo- k 103

hecto- h 102

deka- da 101

deci- d 10-1

centi- c 10-2

milli- m 10-3

micro- µ 10-6

nano- n 10-9

pico- p 10-12

MEMORIZE THIS TABLE !!!! ROYO pg 17 table 1.3

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Question

• What is the name for:

• 10-9 gram =

• 10-6 seconds =

• 10-3 meter =

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USEFUL STUFF

• Back cover of your text “useful conversions”

• Memorize: 454 grams =1 lb 2.54 cm = 1 in 1.6 km = 1 mi 1 m = 3.33 ft 1 hr = 3600 sec

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TEMPERATURE

• Think of temperature as a measure of the “hotness” of an object.

• Scientist use both the Kelvin (SI) and Celsius scale to express temperature.

• Kelvin scale is an absolute scale.

Idea: All molecular motion would stop at

0 K.

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TEMPERATURE (CONT)

• The Celsius scale is based on the freezing point and boiling point of water.

• 1 Kelvin unit of measure is equal to 1 unit of measure on the Celsius scale…the 2 scales simply have different “starting points”

• There is no Degrees symbol in Kelvin

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TEMPERATURE (CONT)

K = 0C + 273

MEMORIZE THESE CONVERSIONS!!!

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Scientific notation

• Break VERY LARGE or very small numbers into manageable bits

• The number 123,000,000,000 in scientific notation is written as :

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To write a number in scientific notation:

Put the decimal after the first digit and drop the zeroes.

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To find the exponent

• count the number of places from the decimal to the end of the number.

In 123,000,000,000 there are 11 places.

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UNCERTAINTY IN MEASUREMENT

• In scientific work, numbers can be exact or inexact.

• This will be a huge part of your research paper and how you collect your data!!!!

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UNCERTAINTY IN MEASUREMENT (CONT)

• Inexact numbers – numbers obtained by measurement. They can be inexact due to

human error and or the limitations of the equipment that you use.

Exact numbers – known values or quantities Ex: 1 dozen eggs = 12 eggs

All conversion factors are exact numbers!

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• THERE ARE ALWAYS LIMITATIONS IN THE EQUIPMENT USED TO MEASURE QUANTITIES (EQUIPMENT ERRORS) AND THERE ARE DIFFERENCES IN HOW DIFFERENT PEOPLE MAKE THE SAME MEASUREMENT (HUMAN ERROR).

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• UNCERTAINTIES ALWAYS EXIST IN MEASURED QUANTITIES!!!

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To Describe Our Measurements , We Use Precision and Accuracy

• Precision refers to how close a series of measurements are to one another.

• Accuracy refers to how closely individual measurements are to a “true” or accepted value.

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Precision, Accuracy, Both ?

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In general, the more precise a measurement, the more accurate it

will be• This is assuming that all measuring

devices are calibrated and the person doing the measuring is well trained.

• BECAUSE OF THIS, SCIENTISTS PERFORM MULTIPLE TRIALS OF THE SAME EXPERIMENT.

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Which would you use?

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IDEA: EVERY MEASURING DEVICE HAS ITS LIMITATIONS

• Because of this, when recording data, scientists always include an “uncertain” number (the rightmost digit in the number).

• Tied in with this is the concept of significant figures.

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SIGNIFICANT FIGURES (CONT)

• Significant figures are all the certain digits of a measured quantity plus one uncertain digit (the last).

• “sig figs” indicate how precise a number is

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SIGNIFICANT FIGURE RULES

1. All numbers are significant except zeros at beginning of the number.

9.12 0.00912 0.0912 0.912 all have three SF

Leading Zeros ARE NOT significant

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Rule 2

Zeros between non- zeros are significant

1005 = 4 SF

1.03 = 3 SF

Captive Zeros ARE significant

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Rule 3

• Zeros that are at the end of a number and after a decimal point are significant.

9.00 9.10 90.0

all have three SF

Trailing Zeros ARE significant

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SIG FIGS RULES (CONT)

4. When a number ends in zeros but has no decimal point, the zeros may or may not be significant. We use scientific notation to clarify.

e.g 100 has 1 known sig fig, but 1.00 x 102 has 3 sig figs.

NOTE!!!!!!!

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Question

• How many Sig figs?

4.0036.023 x 1023

50000.00134

Write 5000 with three sig figs.

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SIG. FIG CALCULATION RULES

• When multiplying or dividing, the answer has the same total number of sig figs as the measurement with the least number of sig figs.

Example:

1.06 m x 13.400 m = 14.204 m2 14.2 m2 3SF 5SF 3 SF

Round when the answer contains too many digits. You must have correct SF and UNITS for credit!

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More Rules!

• In addition or subtraction, the answer has as many digits after the decimal point as the measurement with the fewest number of digits after the decimal point measurement).– Example:

• 38.5200 cm - 11.4 cm = 27.12 cm 27.1 cm 6 SF 3 SF 3 SF

4 digits 1 digit 1 digit

• When doing multi-step calculations, round at the end.

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Question

• Find the volume of a box with the following measurements:

L = 27.3 cm W = 15.5 cm H = 5.4 cm

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Answer

L = 27.3 cm W = 15.5 cm H = 5.4 cm

3 SF 3 SF 2 SF

2285.01 cm3 -> 2.3 x 103 cm3

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P E M D A S

Please Excuse My Dear Aunt Sally

( ) Zy X ÷ + -

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Multiple function equations

8.925 - 8.904 x 100 = 0.23529

8.925

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(9.04 – 8.23 + 21.954 + 81.0) /3.1416

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9.2 x 100.65

8.321 + 4.026

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• 0.1654 + 2.07 – 2.114 = 0.12

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• Derived units , uncertainty in measurement, dimensional analysis

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RECALL

• SI units tell us our base units

• Mass = grams g

• Time = seconds s

• Temperature = Kelvin K

• Length = meter m

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DERIVED SI UNITS (CONT)

• In the lab, scientists use graduated cylinders, burets, pipets, and syringes to measure volume.

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PROBLEM

• Sometimes scientist have to convert from one unit of measure to another (e.g. it would be inappropriate to describe the distance from the Earth to the moon in µm).

• To do so, scientist use a tool called dimensional analysis or factor-label method.

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DIMENSIONAL ANALYSIS (CONT)

• Dimensional analysis is a problem-solving technique where a quantity is expressed equivalently using a different type of unit through the use of 1 or more conversion factors.

• This insures that we will calculate the covert numerical value as well as the correct units.

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• A conversion factor is a fraction whose numerator and denominator are the same quantity expressed in different units.

• E.g. 1 ft/12 inch or 365 days/year

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• What volume, in L, does a 37.3 mL sample occupy?

• 37.3 mL x __10-3_L____ = 0.0373 L

1 mL

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• How many cm exist in a 8.50 in sample?

• 8.50 in x 2.54 cm = 21.6 cm

1 in

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General Rule of Thumb

1. Begin the conversion by looking at the units given and desired.

2. Identify one/ many conversion factors that will take you from the given to the desired units

Given unit x Desired Unit = Desired Unit

Given Unit

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Convert 8.00 m to in

8.00 m x 1 cm x 1 in = 315 in 10-2 m 2.54 cm

OR

8.00 m x 102 cm x 1 in = 315 in 1 m 2.54 cm

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Question

1.) 3.85m to mm

2.) How many cm are in 6.5 miles (hint start with what you know)

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Answer3.85 m x 1mm = 3850 m 10-3m

6.5mi x 1600m x 102cm = 1.0 x 106 cm 1 mi 1m

6.5mi x 5280 ft x 12in x 2.54 cm = 1.0 x 106cm 1 mi 1 ft 1 in

6.5 mi x 1 km x 105cm = 1.0 x 106cm 0.62137 mi 1 km

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DERIVED SI UNITS

• Derived SI units are used to measure quantities expressed as combinations of SI units.

• Area = length in m x width in m = m2

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COMMON DERIVED UNITSWant to find Formula Derived Unit

Area length x width m2

Volume length x width x height

m3

Density mass/volume g/L

Molar mass mass/amt of substance

g/mol

Concentration amt. of substance/volume

Mol/L

Energy force x length N-m

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DERIVED SI UNITS (CONT)

• Volume = l x w x h = m x m x m = m3

The volume of liquids is often measured in mL.

1 cm3 = 1 mL 1 dm3 = 1 L

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Cubing conversation factors

• 13 cm3 to dm3

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Density

• Mass in a unit volume of a substance

Density = Mass / Volume

Units = g/cm3 or g/ml

MEMORIZE!!!!

Density of Water = 1.00 g/ml

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Density Cont.

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Question

1. Find the density of 1.00 x 102 g of mercury that occupies a volume of 7.36 cm3 .

2. Find the volume of 65.0 grams of methanol with a density of 0.791 g/ml

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Challenge question

• What is the mass in grams of a cube of gold (d = 19.32 g/cm3). The length of the cube of gold is 2.00 cm

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Answer

• Find the mass from the volume of the cube and its density. Volume of a cube is given by its length.

Mass = Volume x Density

8.00 cm3 (19.32 g) = 154 g

1 cm3

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• Flash Example

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WELCOME TO CHEMISTRY!