The Basics

The Basics Observations Measurement and Units Uncertainty and Significant Figures Unit Conversions Chemical and Physical Changes Density Simple Error Analysis

Be the Younger Smarter One

Observations All of chemistry revolves around detailed observation in lab Two Types of Observation Quantitative Having to do with quantity, numbers Measurements Qualitative Having to do with qualities, words and descriptions.

Measurements All measurements must contain two things: Some number Some unit Both are equally important and neither can be left out.

The Unit Units are equally important as the number A number without a unit is meaningless.

Use Units or Else Air Canada Flight Pilots used a conversion factor with no units. Unfortunately it was the wrong conversion factor Needed conversion between L to kg Accidentally given L to lbs Ended up with half the fuel they needed.

Or Just Thoroughly Embarrass Yourself… Cents vs Dollars.mp4

Standard Units Système International d'unités (SI) International System of Units Length – m – meter Mass – kg – kilogram Time – s – second Temperature – K – Kelvin Volume – L – liter Energy – J - Joule

SI Prefixes Name Symbol Scientific Notation giga- G 1x109 mega- M 1x106 kilo- k 1x103 Base unit - 1x100 deci- d 1x10-1 centi- c 1x10-2 milli- m 1x10-3 micro- μ 1x10-6 nano- n 1x10-9 pico- p 1x10-12

The Number How long is the black bar? Write down your answer.

Measurements The Rules: Use all digits known with certainty – the 2 and 6 Use one additional digit that is estimated. (3, 4, 5, 6)

Uncertainty The use of an estimated digit introduces some amount of uncertainty into the process. In any measurement the last digit is always assumed to have some uncertainty about it.

Measurements Meniscus – curved surface made in cylindrical containers Read the volume at the bottom of the meniscus Parallax error – error created by failing to read an instrument at eye level. Think the speedometer from the passenger seat

Measurement Activity

Results Balance #1 Balance #2 Ruler #1 Ruler #2 Buret #1 Buret #2 GC #1 GC #2

Back to Uncertainty The last digit in a measurement always has some uncertainty By convention, not all digits represent the uncertainty in a number Significant figures – digits used to represent the amount of certainty in a measured or calculated amount. Sig figs are based on the concept that placeholders only give information about the size of the measurement, not its certainty

Significant Figures The Rules: All non-zero digits are significant. Zeros that are sandwiched between significant figures are significant. Otherwise, zeros are significant only if they follow a decimal point AND a non-zero digit (the non-zero digit may be before OR after the decimal point).

Sig Fig Practice How many sig figs do the following numbers have:

Sig Figs in Calculations The least certain number in a calculation limits the amount of certainty in the answer. Round your answers to the lowest number of sig figs in the problem.

Sig Figs Practice What is the volume of a cube with 1.25cm sides? What is the volume of a rectangular prism that has side lengths of 11.0cm, 5.5cm, and 112cm? 22

Unit Conversions We will use the factor label method Multiply a value by conversion factors to convert to different units Must cancel the units you don’t want Must leave the units you do want The numerator and denominator must be equivalent to each other

Practice Problems Convert 240s to min Convert 590 inches to yards

Practice Problems Convert 450mg to g Always put the scientific notation you memorized with the base unit. Base unit – unit without a prefix

Practice Problems 85 cm to km Where is the base unit? Convert to the base unit first, then convert to the desired unit.

Practice Problems 101.325 kPa to μPa What is the base unit here?

Compound Units Convert 55m/s to km/hr Put the seconds on the bottom and then convert both units separately.

Compound Units Convert 0.789g/mL to kg/cm3 Write this down: 1mL = 1cm3 YOU NEED TO MEMORIZE THIS

Compound Units Convert 120 ft/min to m/s. (2.54cm = 1 inch) 31

Density Density – the ratio of the mass of a substance to the volume of a substance. Typical units are g/mL

Expectations for Algebra Work When working algebra problems, you must: Solve the equation you are using symbolically (and show this step) before you substitute in numbers. Example: When solving for mass in a density problem you must solve it first and show that m = DV Substitute the numbers in with units. Show the units cancelling and the resulting units in the answer.

Density What volume of ethanol would have a mass of 16.9g? What is the mass of 4.91mL of lead?

Error Analysis Ways to describe data: Accuracy – how close a value is to the correct value Precision – how close measured values are to each other

Challenge You are measuring an object that is exactly 2.61cm long. Come up with a set of 3 measurements that is: Accurate and precise Precise but not accurate Accurate but not precise Neither accurate nor precise

Chemical and Physical Changes Chemical change – a change where the chemical makeup of a substance is changed, a change where a substance’s identity is changed. Physical change – some change where no substance’s chemical structure or identity is changed. The appearance may change but the chemical composition does not

Indicators of a Chemical Change Color change Formation of a precipitate Precipitate – a solid produced by a chemical reaction in a solution. Formation of a liquid or gas

Indicators of a Chemical Change Absorption or release of heat Exothermic – the reaction releases heat You feel it getting hotter Endothermic – the reaction absorbs heat You feel it getting colder Light production

Physical Changes Dissolving and dilution. Phase changes are physical changes.