CHAPTER 1 Chemistry: The Study of Change

CHEMISTRY

UNITS Table 1.2 SI Base Units Base QuantityName of UnitSymbol Lengthmeterm Masskilogramkg Timeseconds TemperaturekelvinK Amount of substancemolemol

PrefixSymbolMeaning Tera-T10 12 Giga-G10 9 Mega-M10 6 Kilo-k10 3 Hecto-h10 2 Deca-da10 1 Deci-d10 -1 Centi-c10 -2 Milli-m10 -3 Micro-  Nano-n10 -9 Pico-p PREFIXES

The number of atoms in 12 g of carbon: 602,200,000,000,000,000,000, x The mass of a single carbon atom in grams: x N is a number between 1 and 10 n is a positive or negative integer SCIENTIFIC NOTATION

n > = __________ move decimal left n < = ___________ move decimal right SCIENTIFIC NOTATION

SIGNIFICANT FIGURES

Once you start the counting don’t stop!

SIGNIFICANT FIGURES Rule 1: Every nonzero digit in a measurement is significant. Examples: Rule 2: Zeros appearing between nonzero digits are significant. Examples: Rule 3 A ZERO is NOT significant when it is a placeholder. A placeholder is used to show the location of the decimal point. Examples:

SIG FIGS CONT. Rule 4: Zeros at the end of a number and to the right of a decimal point are always significant. Examples: Rule 5: When a number is counted or defined within a system of measurement, there is an infinite amount of significant digits. Examples: 11 students100 cm = 1 m

24 mL 3001 g m x 10 4 molecules 560 kg COUNT THE SIG FIGS:

The answer cannot have more digits to the right of the decimal point than any of the original numbers round off to ________ one significant figure after decimal point two significant figures after decimal point round off to ________ SIG FIGS – Addition and Subtraction

The number of significant figures in the result is set by the original number that has the smallest number of significant figures 4.51 x = = _______ 3 sig figsround to 3 sig figs 6.8 ÷ = sig figsround to 2 sig figs = _______ SIG FIGS – Multiplication and Division

1.Determine which unit conversion factor(s) are needed 2.Carry units through calculation 3.If all units cancel except for the desired unit(s), then the problem was solved correctly. 1 L = 1000 mL How many mL are in 1.63 L? 1L 1000 mL 1.63 L x = 1630 mL 1L 1000 mL 1.63 L x = L2L2 mL FACTOR LABEL METHOD

The speed of sound in air is about 343 m/s. What is this speed in miles per hour? 1 mi = 1609 m 1 min = 60 s1 hour = 60 min meters to miles seconds to hours

Important things to consider when solving problems and performing experiments….

Volume – SI derived unit for volume is cubic meter (m 3 ) 1 L = 1000 mL = 1000 cm 3 = 1 dm 3

Matter - anything that occupies space and has mass. mass – measure of the quantity of matter SI unit of mass is the kilogram (kg) weight – force that gravity exerts on an object weight = c x mass on earth, c = 1.0 on moon, c ~ 0.1 A 1 kg bar will weigh 1 kg on earth 0.1 kg on moon

Density – SI derived unit for density is kg/m 3 1 g/cm 3 = 1 g/mL = 1000 kg/m 3 density = mass volume D = m V A piece of platinum metal with a density of 21.5 g/cm 3 has a volume of 4.49 cm 3. What is its mass? For Water: 1g/mL

K = 0 C F = x 0 C K = 0 0 C 373 K = C 32 0 F = 0 0 C F = C

Convert F to degrees Celsius. 0 F = x 0 C F – 32 = x 0 C 9 5 x ( 0 F – 32) = 0 C C = x ( 0 F – 32) C = x (172.9 – 32) =

Graphing Distance vs. Time Dependent Variable Independent Variable Line of Best Fit

Accuracy – how close a measurement is to the true value Precision – how close a set of measurements are to each other

Percent Error: Percents and Percent Error Measured Value- Accepted Value Accepted Value X 100 Example: The mass of a compound measured in a lab was 25.0 grams. The accepted value for this compound is 24.5 grams. Calculate the percent error g g 24.5 g X 100 = 2.04 %

Scientific Method

SCIENTIFIC METHOD logical approach to solving problems Observation Problem Hypothesis Experiment  Data  Analysis Conclusion

Observations Qualitative: quality, non-numeric terms Quantitative: quantity, numerical description

MATTER: ANYTHING THAT OCCUPIES SPACE AND HAS MASS MATTER

State of Matter VolumeShapeDensity Compressibility Motion of Molecules Gas Liquid Solid

Plasma Three States of Matter

solid liquidgas Phase Changes

phase diagram: summarizes the conditions at which a substance exists as a solid, liquid, or gas. Phase Diagram of Water

Phase Diagram Points ________________:  above this point a substance becomes a supercritical fluid  critical temperature (T c ): temperature above which the gas cannot be made to liquefy, no matter how great the applied pressure.  critical pressure (P c ): minimum pressure that must be applied to bring about liquefaction at the critical temperature. ________________:  point at which all three phases coexist

1. Pure Substance: form of matter that has a definite composition and distinct properties 2. Mixture : combination of two or more substances in which the substances retain their own identities. water, ammonia, sucrose, gold, oxygen Classification of Matter mixed together physically can usually be separated

1. Homogenous mixture – composition of the mixture is the same throughout. 2. Heterogeneous mixture – composition is not uniform throughout. Examples: 1.4 Examples: Types of Mixtures

Mixture Pictures

solution: mixture that remains uniformly mixed solute: solvent: suspension: mixture where visible particles settle colloid: mixture where particles are unevenly distributed but do not separate, positive Tyndall Effect Types of Mixtures

Hom/Het?Soln/Susp/Coll? Fog_______________________ Paint _______________________ Syrup _______________________

Physical means can be used to separate a mixture into its pure components. magnet 1.4 distillation

Methods of Separation Strainer Filtration Physical Evaporation Centrifuge Distillation PASTA/WATER SAND/IRON FILINGS SALT /WATER BLOOD FOOD COLORING/WATER SAND/WATER

Element: all atoms are the same cannot be broken down by physical or chemical means 114 elements named on the Periodic Table 83 elements occur naturally on Earth gold, aluminum, lead, oxygen, carbon many elements have been created by scientists technetium, americium, seaborgium

Compound: 2 or more elements combined Cannot be broken down by physical means Can be broken down by chemical means Appears different from original elements Fixed ratios in definite proportions Water (H 2 O)Glucose (C 6 H 12 O 6 ) Ammonia (NH 3 )

PHYSICAL AND CHEMICAL CHANGES

Physical Properties and Physical Changes physical property: characteristic that can be observed or measured without changing the identity of the substance. physical change: change in a substance that does not involve a change in the identity of the substance.

Physical Properties Intensive: INDEPENDENT of amount of matter present (sample size)  Example: Extensive: DEPENDENT on the amount of matter present (sample size)  Example:

Chemical Properties and Chemical Changes chemical property: a substance’s ability to undergo changes that transform it into different substances Example: chemical change: change in which one or more substances are converted into different substances Example:

Evidence of a Chemical Change

physical change does not alter the composition or identity of a substance. chemical change alters the composition or identity of the substance(s) involved. ice melting sugar dissolving in water hydrogen gas burns in oxygen gas to form water Physical or Chemical? Remember: