Honors Chemistry Summer WorkWakefield Memorial High School, Mrs. Brazile

Names of phase changesScientific notationBasic algebraFractions

Welcome to Honors Chemistry! Before coming to the first class in the fall, you must present this completed packet on the first day of school. You are responsible for learning the information included here and completing all of the PRACTICE problems. Though some sample problems are provided to help you, you may also use the Internet for assistance, including the following videos:

https://www.youtube.com/watch?v=KRvsPh4pU90. https://www.youtube.com/watch?v=hQpQ0hxVNTg

You will be tested on the material within the first week of school, so please review it carefully and complete the practice problems with detail.

Prerequisite Skills

Names of phase changes

Scientific notation

Basic algebra

Fractions

1.1 Atomic and Molecular Perspective

MATTER: physical material; has mass and occupies space

PROPERTY: characteristic that allows you to recognize a type of matter and distinguish it from others

ELEMENT: most basic form of matter

ATOM: smallest unit of an element; building block of matter

COMPOUND: chemical combination of atoms of different elements

MOLECULE: smallest unit of a compound

Atoms of elements combine to form molecules of compounds.

1.2 Classifications of Matter

Classify matter 2 ways

1. Physical state (phase)

2. Composition

States of Matter GAS (vapor)

No fixed volume or shape; conforms to size and shape of container

Can be compressed into smaller volume or can expand to larger one

Molecules are spread apart

Molecules collide with each other and walls of container

LIQUID Fixed volume but no fixed shape; conforms to shape of container

Cannot be compressed

Molecules closer than in gas

Molecules move rapidly but slide past one another allowing us to pour easily

SOLID Fixed shape and volume

Cannot be compressed

Molecules held tightly together, usually in specific arrangement

Molecules can only wiggle slightly in otherwise fixed positions

States of Matter (cont)

Three States of Matter of H2O

Pure Substances

Most things we encounter are not pure

PURE SUBSTANCE: matter that has distinct properties and composition that doesn’t vary from sample to sample

Elements

Compounds

Elements

90% of Earth’s crust is 5 elements; 90% of human body is 3 elements

Organized on the periodic table by atomic number

Represented by chemical symbols

Compounds

Most elements interact with other elements to form compounds which have a totally new set of properties than original elements

Represented by chemical formulas (symbols of elements that make them up)

Ex: water (H2O) is made of hydrogen (H) and oxygen (O)

Some elements exist as diatomic molecules (7)

Compounds (cont)

Joseph Louis Proust (1754-1826)

LAW OF CONSTANT COMPOSITION (definite proportions): the elemental composition of a pure compound is always the same

Pure compounds have same composition and properties regardless of how much, where it came from, etc

Mixtures

Not pure; composition varies

Substances making up mixture (components) are pure substances physically combined (can be separated)

2 types

HOMOGENEOUS (solutions): uniform throughout and appears as one

Ex: salt dissolved in water, air

HETEROGENEOUS: varied appearance and properties

Ex: many types of rocks, trail mix

Separation of Mixtures

Components in mixture retain their own properties which we utilize to separate them

1. FILTRATION: solids and liquids separated through porous medium like filter paper

Ex: pulp from orange juice

Separation of Mixtures (cont)

2. DISTILLATION: separates substances based on different boiling points

Ex: salt from water

Separation of Mixtures (cont)

3. CHROMATOGRAPHY: utilizes ability of a substance to adhere to a surface like paper

Ex: dyes in ink

Classifications of Matter Summary

PRACTICE #1

1. Classify the substances as element, compound, homogeneous mixture, or heterogeneous mixture.

a. Chocolate chip cookie dough

b. Gatorade

c. Helium inside a balloon

d. Baking soda (NaHCO3)

e. Orange juice with pulp

f. Dirt

g. Purified air

h. Table salt (NaCl)

i. Copper pipe

1.3 Properties of Matter

PHYSICAL PROPERTIES: can be measured without changing identity and composition of substance

Ex: color, odor, density, melting/boiling point

CHEMICAL PROPERTIES: describe how a substance changes (reacts) to form other substances

Ex: flammability, rusting

INTENSIVE PROPERTIES: do not depend on the amount of substance being examined

Ex: melting/boiling point, density, temperature

EXTENSIVE PROPERTIES: depend on quantity of sample

Ex: mass, volume

Physical and Chemical Changes

PHYSICAL CHANGE: substance changes in appearance but not composition

Ex: phase changes, breaking glass

CHEMICAL CHANGE (reaction): substance transforms into a chemically different substance

Ex: burning, rusting

PRACTICE #2

2. Classify as a physical property, physical change, chemical property, or chemical change.

a. Melting chocolate in the sun

b. The iron screws are rusting.

c. The density of water is 1.00 g/mL.

d. Paper is flammable.

e. The boiling temperature of water is 100 Celsius.

f. Paper is burning.

1.4 Units of Measurement

SI UNITS: international set of units used in scientific measurements

A measurement means NOTHING without a unit

Metric System

Prefixes used to indicate decimal fractions or multiples of various units

Memorizing Prefixes

Come up with sentence to help remember order

T G M k h d o d c m μ n p

Measurements in Lab

LENGTH: distance (m, cm)

MASS: amount of matter in an object (g, kg)

TEMPERATURE: measure of hotness or coldness of an object; determined direction of heat flow (K, °C)

K = °C + 273.15

Derived Units

Use units of other quantities to generate

VOLUME: amount of space an object occupies

Volume of a cube = Length3 = m × m × m = m3

DENSITY: amount of mass in a unit volume of a substance

Density = Mass ÷ Volume = g/cm3

PRACTICE #3

3. Answer the questions regarding units.

a. How many centimeters are in 1 meter?

b. How many bytes are in 1 megabyte?

c. How many millimeters make up a kilometer?

1.5 Uncertainty in Measurement

2 kinds of numbers

1. EXACT

Exactly 12 eggs in 1 dozen

2. INEXACT (contain uncertainty)

All measurements due to limitations in equipment and human error

Use precision and accuracy to discuss uncertainties in measured values

Precision and Accuracy

PRECISION: how closely individual measurements agree with one another

ACCURACY: how closely individual measurements agree with correct/true value

Significant Figures

Last digit of measured quantities is uncertain

SIGNIFICANT FIGURES: all digits in measured quantity, including uncertain one

Which has more certainty, 2.2 g or 2.2405 g?

Greater number of significant figures = greater certainty

Relates to precision of a measurement, not accuracy

Significant Figures (cont)

Rules for counting SigFigs:1. All nonzero digits are significant.

EX: 546 cm = 3 SF

2. Zeros between nonzero digits are significant.

EX: 4,005 g = 4 SF

3. Zeros at the beginning of a number are not significant.

EX: 0.003 m = 1 SF

4. Zeros at the end of a number are significant if the number contains a decimal anywhere in the number. EX: 450 = 2 SF

EX: 4.50 = 3 SF

Scientific notation must be used in certain cases to express a measurement with the proper number of significant figures

PRACTICE #4

4. Determine the number of SigFigs in each measurement below.

EXAMPLE: 3,454 cm = 4 SF

a. 4, 002,341 g

b. 0.984 m

c. 0.802 oz

d. 25,000 km

e. 598.00 g

f. 0.00908300 sec

PRACTICE #5

5. Determine how to express the measurements below in 2 SigFigs.

EXAMPLE: 3,434 cm = 3,400

a. 0.984 m

b. 0.805 oz

c. 25,800 km

d. 598.00 g

e. 0.009 sec

f. 4, 002,341 g

How do you know how many digits to write down when you’re taking a measurement?

Based on instrument…look at the tick marks!

RULE FOR SigFigs IN MEASURING: Record values with digits including smallest interval (tick mark) PLUS one estimated digit.

Significant Figures in the Lab

Number Places

PRACTICE #6

SMALLEST INTERVAL

Ex: 1.44 cm

PRACTICE #6 (cont)

SMALLEST INTERVAL

SMALLEST INTERVAL

ANSWER ANSWER

Significant Figures in Calculations

When measured quantities are used in calculations, the least certain measurement limits the SigFigs in the answer.

1. MULTIPLICATION & DIVISION: answer contains same number of SigFigs as measurement with fewest SigFigs.

Area = (6.221 cm)(5.2 cm) = 32.3492 cm2 = 32 cm2

2. ADDITION & SUBTRACTION: answer contains the same number of decimal places as the measurement with the fewest decimal places.

20.43 cm + 1.322 cm + 83.1 cm = 104.842 cm = 104.8 cm

PRACTICE #7

7. Practice using rules of SigFigs when performing the following calculations.

a. 837.6 g – 836.22 g =

b. 1.4 g ÷ 1.05×105 cm3

1.6 Dimensional Analysis

Mathematical strategy using conversion factors to convert units

CONVERSION FACTOR: fraction whose numerator and denominator are the same quantity expressed in different units

Select which conversion factor to use depending on what you are converting to/from

2.54 cm

1 in and

1 in

2.54 cm

Dimensional Analysis (cont)

Given unit ´ desired unit

given unit = desired unit

150.6 cm ´ 1 inch

2.54 cm = 59.3 in

72 in ´ 2.54 cm

1 inch = 180 cm

Utilize unit cancellation (numerator and denominator)

PRACTICE #8

8. Perform the conversions below using dimensional analysis.

a. 34 dozen = _________ eggs

b. 8.9 in = _________ ft

c. 0.234 ft = _________ cm

PRACTICE #9

9. Perform the conversions below using dimensional analysis.

a. 45 kg = ____________ g

b. 160 nm = _____________ km

Dimensional Analysis (cont)

Conversions involving volume (3-dimensional unit)

2.00 in3 = __________ cm3

2.00 in3 ´

2.54 cm

1 in

æ

èç

ö

ø÷

3

=

2.00 in3 ´ (2.54)

3 cm

3

(1)3 in

3=

2.00 in3 ´

16.387 cm3

1 in3

= 32.774 cm3 = 32.8 cm

3

PRACTICE #10

10. Perform the following conversion using dimensional analysis.

a. 13,450 cm3 = __________ ft3

PRACTICE #10 (cont)

10. What if the units you want to convert are in the denominator?

b. 15 mi/hr = ______________ mi/min

PRACTICE #10 (cont)

10. Perform the following conversions using dimensional analysis.

c. 15 mi/hr = ______________ ft/sec

PRACTICE #10 Challenge

The normal lead content of human blood is about 0.40 parts per million (that is, 0.40 g of lead per one million grams of blood). A value of 0.80 ppm is considered to be medically dangerous. The average adult has 6.0 kg of blood. Determine how many grams of lead are contained in the average adult if the lead content is 0.62 ppm.

ADDITIONAL PRACTICE

11. Convert 35C to Kelvin.

12. Convert 375 K to degrees Celsius.

ADDITIONAL PRACTICE (cont)

13. If the density of a substance is 48 g/L and a sample of that substance occupies a volume of 22.4 L, what is the mass of the substance within the sample? Predict what phase of matter this is most likely to be.

14. If you perform an experiment with three trials to determine the boiling point of water and your data is T1: 98.7C, T2: 99.4C, T3: 100.2C, was your data accurate, precise, both, or neither? Explain.