Unit One Review

Accuracy and Precision The important things to remember about

accuracy and precision:

You want measurements that are both accurate and precise!

Accurate: You hit the bulls-eye! (to also be precise you must hit it every time)

Precise: you hit the same spot every time (you get the same measurement every time) You can be precise without being accurate

Accuracy vs. Precision Accuracy - how close a measurement is

to the accepted value

Precision - how close a series of measurements are to each other

ACCURATE = CORRECT

PRECISE = CONSISTENT

Accuracy and Precision

Accuracy and Precision If you are measuring the temperature of

boiling water:

1st measurement: 980C; 2nd measurement: 980C; 3rd measurement: 970C

1st measurement: 1010C; 2nd 1020C; 3rd 970C

1st measurement: 1000C; 2nd 1000C; 3rd 1000C

Measurements

When we measure something, we can (and do) always estimate between the smallest marks.

21 3 4 5

Measurements

The better marks the better we can estimate.

Scientist always understand that the last number measured is actually an estimate.

21 3 4 5

Significant Figures Indicate precision of a measurement.

Recording significant figures Significant figures in a

measurement include the known digits plus a final estimated digit

2.35 cm

Significant Figures

Counting Sig Figs Count all numbers EXCEPT:

Leading zeros -- 0.0025

Trailing zeros without a decimal point -- 2,500

4. 0.080

3. 5,280

2. 402

1. 23.50

Significant FiguresCounting Sig Fig Examples

1. 23.50

2. 402

3. 5,280

4. 0.080

4 sig figs

3 sig figs

3 sig figs

2 sig figs

Significant Figures

How many sig figs in the following measurements?

458 g

4085 g

4850 g

0.0485 g

0.004085 g

40.004085 g

Significant Figures Calculating

Multiply/Divide - The # with the fewest sig figs determines the # of sig figs in the answer.(13.91g/cm3)(23.3cm3) = 324.103g

324 g

4 SF 3 SF3 SF

Significant Figures Calculating

Add/Subtract - The # with the lowest decimal value determines the place of the last sig fig in the answer.

3.75 mL

+ 4.1 mL

7.85 mL

224 g

+ 130 g

354 g 7.9 mL 350 g

3.75 mL

+ 4.1 mL

7.85 mL

224 g

+ 130 g

354 g

Scientific Notation

Converting into Sci. Notation: Move decimal until there’s 1 digit to its

left. Places moved = exponent.

Large # (>1) positive exponentSmall # (<1) negative exponent

Only include sig figs.

65,000 kg 6.5 × 104 kg

Graphs Circle graphs – compare

parts in a whole

Bar graphs – compare quantities

Line graphs – compare sets of data, show change and patterns over time.

Base Units

Length - meter - m

Mass - grams or kilogram - g

Time - second - s

Temperature – Kelvin - K

Energy - Joules- J

Volume - Liter - L

Amount of substance - mole - mol

Prefixes

kilo k 1000 times

deci d 1/10

Centi c 1/100

milli m 1/1000

kilometer - about 0.6 miles

centimeter - less than half an inch

millimeter - the width of a paper clip wire

Conversion PracticeKilo- 1000 units

Hecto-

100 units

Deka- 10 units Base Unit Deci-

0.1units

Centi-

0.01units

Milli-0.001units

To convert to a smaller unit, move the

decimal point to the right or multiply.

To convert to a larger unit, move the decimal point to the left or

divide.

Converting

k h D d c m how far you have to move on this chart, tells you how far,

and which direction to move the decimal place.

The box is the base unit, meters, Liters, grams, etc.

Conversions

Change 5.6 m to millimeters

k h D d c m

starts at the base unit and move three to the right.move the decimal point three to the right

56 00

Density

How heavy something is for its size.

The ratio of mass to volume for a substance.

D = M / V

Independent of how much of it you have

gold - high density

air low density.

Volume and DensityRelationship Between Volume and Density for Identical Masses of Common Substances

Cube of substance Mass Volume Density Substance (face shown actual size) (g) (cm3) (g.cm3)

Lithium

Water

Aluminum

Lead

10 19 0.53

10 10 1.0

10 3.7 2.7

10 0.58 11.4

Celsius & Kelvin Temperature Scales

Temperature Scales

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 136

Measuring Temperature

Kelvin starts at absolute zero (-273 º C)

degrees are the same size

C = K -273

K = C + 273

Kelvin is always bigger.

Kelvin can never be negative.

273 K

Specific Heat

States of Matter

3 states

Gas No definite volume or shape

Liquid Definite volume, but not shape

Solid Definite shape and volume

Controlled Experiments

1. Manipulated Variable: is one thing (factor) you change or allow to be different in an experiment.

2. Responding Variable: is one thing (factor) that you are measuring or observing (the outcome of the experiment).

3. Controlled Variables: are factors or things that are kept the same at the beginning of the experiment.

Tulip Experiment

Materials:

Tulip plant

Water

Sun Minerals Fertilizer

CONTROL (Pot 2)

Tulip plant

Water

Sun

Fertilizer

EXPERIMENTAL (Pot 1)

Tulip plant Water Sun Minerals Fertilizer

What are the variables in this experiment?

CV: Tulip plant, Water, Sun, Fertilizer

MV: Minerals

RV: Color of the tulip flowers

(ECC: no minerals)