chemistry – chapter 1 & 2

CHEMISTRY – Chapter 1 & 2

Matter, Measurements, and Calculations

Chapter 1 – Section 1

Objectives:1. Define chemistry2. List examples of branches of

chemistry3. Compare and contrast basic research,

applied research, and technological development

What objects in this room are related to chemistry? Plastics Fabrics Clothes Cooking oil Motor oil Make-up Radio Batteries Computers

Chemistry in our daily lives.

Antibiotics Food Transportation Sports Farming Military Industry

Chemistry

Study of the composition and properties of matter and the changes that matter undergoes- What something is made of- What is the internal arrangement

Chemical

Any substance that has a definite composition

6 Main Branches of Chemistry

1. Organic – substances containing C2. Inorganic – substances other than organic3. Biochemistry – living things4. Physical chemistry – changes of matter5. Analytical chemistry – id components of

materials6. Theoretical chemistry – use math and

computers to understand chemical behavior

All branches involve some type of research.

Basic research – to increase knowledge- how and why

Applied research – to solve problems

Technological development – production and use of products- lags behind discoveries

- application of knowledge

Review and Assignment

1. Define chemistry2. List examples of branches of

chemistry3. Compare and contrast basic research,

applied research, and technological development

Assignment: WS 1-1

Quiz

1. Name two branches of chemistry.2. List two ways that chemistry affects

our daily lives.3. Definition of chemistry.

Chapter 1 - Matter


Objectives:

1. Distinguish between a mixture and a pure substance.

2. Define what matter is.

Matter

- anything that has mass and occupies space

- includes almost everything- exceptions are light, heat, and sound- properties are used to measure matter

ex. mass

Mass – measure of quantity of matter- not affected by temp, location, or any other factor

Demo. Mass vs. matter What caused the change in mass? Is air matter?

Matter (cont.)

Classified into 2 groups:1. pure substances2. mixtures

Pure substance – matter that has the same properties throughoutex. element or compound

Pure SubstancesElement – substance that cannot be broken down by

ordinary chemical change- only 1 type of atom- symbols abbreviated w/1 or 2 letters- can be an allotrope

allotrope – one of a number of different molecular forms of an element in the same

state

Compound – substance made up of 2 or more elements chemically combined- can be broken down by chemical change- more than 1 type of atom

Compounds1. Elements that make up a compound are

combined in definite proportion by massex. 100 g water has 11.2 g H and 88.8 g of O

2. Chemical and physical properties of compound differ from those of its partsex. water is liquid, H and O are gases

3. Compounds can be formed from simpler substances by chem change and can be broken down into simpler substances

example

100 of water has 11.2 g H and 88.8 g OHow many g of H is in a 120g sample of

water?

120 g water | 11.2 g H = 13.4 g H| 100 g water

Mixtures

- contain 2 or more substances that have different properties- vary in composition and properties from sample to sampleex. rock, wood, salt water

- Not chemically combined- Can be separated by simple physical means

- ie. filtration, evaporation, distillation

Formation of Mixtures

A mixture can be formed 3 ways:1. Element mixed w/1 or more other elements

ex. carbon w/sulfur2. Compound mixed w/ 1 or more other

compoundsex. salt w/sugar

3. 1 or more elements mixed w/1 or more compoundsex. sulfur w/sugar

Characteristics of Mixtures

- retain properties of each of its partsex. iron and sulfur

- iron remains magnetic

- composition can vary widely

- can be homogeneous or heterogeneous

Types of mixtures

Homogeneous – uniform composition throughout- called solutionsex. alloys, pop, air, coffee

Heterogeneous – not uniform throughoutex. concrete, soil, dry soup, spaghetti and meat balls

Matter

Pure substance Mixture

Element Compound Homogeneous Heterogeneous


1. Distinguish between a mixture and a pure substance.

2. Define what matter is.

Assignment: WS


Objectives:

1. Distinguish between the physical properties and chemical properties of matter.

2. Classify changes of matter as physical or chemical.

3. Explain the gas, liquid, and solid states in terms of particles.

Properties of Matter

- allow us to distinguish btwn substances

- characteristics of a substance- what can be observed- way that a substance behaves

ex. color, taste, odor, gas, liquid, solid

Properties (cont.)

- can be extensive or intensive

Extensive – d/o amount of matterex. volume, weight, mass, and E

Intensive – does not d/o amount of matterex. melting point, boiling point, density, and conductivity

Demonstration Properties

- water and glycerinHow do they compare?

- look, feel, weight, flow

- water and salt waterHow do they compare?

- conductivity

Physical Properties

Can be observed or measured w/out changing the substance Can describe the substance Odor, taste, hardness, density, melting

point, and boiling point Metals – ductile (pulled into wire),

malleable (hammered into sheets), luster (shine), good conductors

Chemical Properties

A transformation of a substance into a different one rusting, flammability, tarnishing, new

substance formed

Physical Change

No new substance is formed CHANGE IN PHASE, pounding, grinding,

cutting Changes of phase

When a substance changes phase there is no change in composition

Physically different, chemically the same Solid, liquid, or gas are the three states

of matter

States of Matter

Solid – definite volume and shape Particles are in fixed positions Held w/strong attractive forces

Liquid – definite volume and no definite shape Takes shape of container Particles can move past each other

States of Matter (cont.)

Gas – neither definite volume nor definite shape Particles move easily and are very far

apart

Plasma – high temperature state in which atoms lose their electrons

Chemical Change

One or more substance is changed to something new Rusting, burning, gas formed, digestion,

heat or light added, explosion, color change, odor change, water formed


1. Distinguish between the physical properties and chemical properties of matter.

2. Classify changes of matter as physical or chemical.

3. Explain the gas, liquid, and solid states in terms of particles.

Assignment: p. 18 and WS

CHEMISTRY – Chapter 1 – Section 3

Objectives:1. Perform density calculations.2. Describe conservation of mass.

Properties of Matter

- E is always involved in both physical and chemical changes- Physical are not at noticable- Chemical are more noticable

- Heat and light are given off

Density is a physical property is always the same for

a solid substance in gases and some

liquids a change in temperature will change the density

increase in temperature will decrease density

D = m/V

Density problem

Use the 5 steps in problem solving to solve the following problem.

Lead has a mass of 22.7 g and its volume is

2.00 cm3. What is its density?

m = 22.7 g V = 2.00 cm3

D = m/V = 22.7 g/2.00 cm3 = 11.4 g/ cm3

Examples

Conservation of Mass

In reactions matter cannot be created or destroyed by a chemical change- mass stays the same, it may just change form

Density Lab Results

Group 1 –

Group 2 –

Group 3 –

Group 4 –

Group 5 -


1. Perform density calculations.2. Describe conservation of mass.

Assignment: WS and Density lab

Chapter 2 - Sec.1

Objectives:1. Describe the purpose of the scientific

method.2. Distinguish between qualitative and

quantitative observations.3. Describe the steps to making a graph.4. Distinguish between inversely and

directly proportional relationships.

Scientific Method- a logical approach to solving problems1. Make observations

- observe your surroundings

2. State the problem- stated as a question

3. Collect data4. Form hypothesis

- testable statement

5. Test hypothesis6. Conclusion7. Modify hypothesis and retest

Observing

Involves making measurements and collecting data Data can be qualitative or quantitative

Qualitative – non-numerical information- descriptive (the sky is blue)

Quantitative – numerical information- the mass is 25.7 grams

Conclusion

Can be explained by using models

Model – explanation of how phenomena occur or how things are related

- visual- verbal- mathmatical

Theory

- models may become part a theory

Theory – broad generalization that explains facts or phenomena

- must be able to predict resultsex. kinetic-molecular theory

collision theory

Controlled Experiments

Use manipulated variable (independent) Use responding variable (dependent) One variable manipulated at a time Measurements are called data

Making a Graph Shows results of an experiment in a

meaningful pattern Dependent variable is on the vertical axis1. Always include a title2. Determine variables3. Set up scale4. Plot points5. Draw best-fit line

Oxygen obtained from electrolysis of water

Electrolysis of water

3

10

1823

28

05

10152025

0 5 10 15 20 25

water (g)

oxyg

en (g

)

Oxygen Water2.7 38.9 1016 18

20.4 2324.9 28

Relationships in graphs Directly proportional – if dividing one by the

other gives you a constant value If one increases so does the other If started at point (0,0)

Inversely proportional – if their product is constant If one increases the other decreases Produce a curve


1. Describe the purpose of the scientific method.

2. Distinguish between qualitative and quantitative observations.

3. Describe the steps to making a graph.4. Distinguish between inversely and

directly proportional relationships.

Assignment: graphing WS

Quiz

1. List three steps of the scientific method.

2. List two steps in making a graph.

Chapter 2 Sec.2

Objectives:

1. Distinguish between a quantity, a unit, and a measurement standard.

2. Name SI units for length, mass, time, volume, and density.

3. Distinguish between mass and weight.

Measurements

Basic part of science Make observations more meaningful Needs to be more than just a number

or quantity Need a common system of units

For consistency Measure your desk w/anything you have

available

SI System

- The International System of Units- Used in all science- A standard- Based on 10

- Makes it easier to convert from one unit to another

SI System (continued)- 7 base units

1. Length – meter (m)2. Mass – kilogram (kg)3. Time – second (s)4. Amount – mole (mol)5. Temperature – Kelvin (K)6. Electric Current – ampere (amp)7. Luminous intensity – candela (cd)

Weight vs. massMass – quantity of matter

- how much space it takes up- measured w/a balance- unit kg

Weight – F gravity pulls on matter with- measured w/spring scale- unit Newton

On the moon will our weight or mass stay the same?

SI Prefixes

You must know these.

Kilo- 1000Deca – 10Base unit (m, s, L) Centi – 1/100 or 0.01Milli – 1/1000 0r 0.001

Derived Units

- combination of base units

Examples- Area = m2

- Volume = m3

- Density = kg/m3

- Newton = m٠kg/s2

Derived Units (cont.)

Area – determined by multiplying 2 lengths

Volume – determined by multiplying 3 lengths for a solid- for liquids unit is cm3 or mL

** 1 mL = 1 cm3


1. Distinguish between a quantity, a unit, and a measurement standard.

2. Name SI units for length, mass, time, volume, and density.

3. Distinguish between mass and weight.

Assignment: WS 2-2 and p. 42 ~1-3

Quiz

1. What is the base SI unit for mass?2. Kilo = ______3. Centi = _____4. What is a derived unit?5. 1 cm3 = _____ mL

Chapter 2 - Sec.3

Objectives:1. Distinguish between accuracy and

precision.2. Determine the number of significant

figures in measurements.3. Perform mathematical operations involving

significant figures.

Accuracy and Precision

Accuracy – closeness of a measurement to correct value

Precision – closeness of a set of measurements to each other Consistency Do not have to be correct d/o measuring instrument

Bullseyes

Significant Figures

- digits in a measurement that are know with certainty and one digit that is estimated

- CALCULATORS DO NOT KEEP TRACK OF SIGNIFICANT FIGURES

Significant Figure Rules1. Digits other than zero are ALWAYS significant

ex. 61.4 3 sig. fig.2. All zeros at the end of a number and to the right of the

decimal with a # preceding the decimal are ALWAYS sigex. 4.7200 km 5 sig. fig.

3. Zeros used only for spacing are NOT significantex. 7000 1 sig. fig.

20 1 sig. fig. 100.0 4 sig. fig.

4. Zeros between sig. fig are significant5. Zeros in front of a non-zero are NOT sig.

- don’t count until you get to 1st non-zero from lf to rt0.004 1 sig. fig.0.0009 1 sig. fig.

Significant Figures

1,000 = _____ sig figs

100.0 = _____ sig figs

0.00012340 = _____ sig fig

10.0340 = _____ sig fig

Calculating w/Significant Figures

Addition and Subtraction- use same # of decimal places as the measurement w/the least decimal placesex. 2.098 3 DECIMAL places

+6.2 1 DECIMAL place8.298 round to 1 Decimal

8.3 is the final answer

Adding and Subtracting

10.0 + 123 = _____

23.456 – 23.0 = _____

100.12 + 56.45 = _____

1,000 + 12.234 = _____

Calculating w/sig. figs (cont.)

Multiplication and Division- use same # sig. fig. as the measurement w/the least sig. fig.ex. 2.38 3 sig. fig

x 9.0 2 sig. fig 21.42 round to 2 sig. Fig

21 is the final answer

Multiplying and Dividing

100.0 x 10 = _____

34.56 x 23.45 = _____

12.045 x 34.008 = _____

50.04 x 23 = _____


1. Distinguish between accuracy and precision.

2. Determine the number of significant figures in measurements.

3. Perform mathematical operations involving significant figures.

Assignment: WS 2-6 and sig fig WS

Quiz

How many significant figures are in the following numbers?

1. 8,000 _____2. 100.01 _____3. 0.00056_____4. 4500.10 _____5. What is precision?

Chapter 2 - Sec.3 Day 2

Objectives:

1. Perform mathematical operations involving percent error.

Percent Error Observed value – based on lab

measurements

True value – based on generally accepted references

Error exists in any measurement d/o measurer, instrument, conditions

Percent Error

% error = true value – obs. value x 100

true value

Exampleatomic mass of Al = 28.9 gmeasured mass = 27.0 g

What is the % error?28.9 g – 27.0 g x 100 = 7.00 %

28.9 g


1. Perform mathematical operations involving percent error.

Assignment: WS 2-5 and % error WS

Quiz

How many significant figures are in the following numbers?

1. 8,104 _____2. 100.01 _____3. What does % error tell us?4. What is accuracy?5. What is precision?

CHEMISTRY – Chapter 2 Sec.3 Day 3

Objectives:1. Use dimensional analysis to convert

measurements.2. Convert measurements into scientific

notation.3. Perform mathematical operations using

exponents.

Problem Solving Rules

Write down what is known.- mass = 346 g volume = 34.6 cm3

2. Write down unknown.- density = ?

3. Write the equation to use.D = m/V

4. Fill in knowns.D = 346 g/34.6 cm3

5. Solve for unknown and label.D = 200 g/cm3

6. Check your work.

Dimensional Analysis

- use with conversion factors to change from one unit to another

Steps:convert 2550 m to km

1. Determine conversion factor- 1000 m to 1 km

2. Set up T-bars3. Write given # in first box

Dimensional Analysis (cont.)

4. Write conversion factor in 2nd box- unit on bottom matches unit of given #

5. Matching labels cancel- if 1 from conversion factor is on top divide

- if 1 from conversion factor is on bottom multiple

Scientific Notation

- Used to represent very large or very small numbers

- There are two parts- Basic form is M x 10n

- M is a number- n is a number representing how many

places to move the decimal

Scientific Notation (cont.) If n is negative, your number is a decimal If n is positive, your number is a large

number

Examples:60,000,000 = 6 x 107

0.000005 = 5 x 10-6

125,000 = 1.25 x 105

Scientific Notation (cont.)

Write the following in scientific notation. 1,000,000,000 23,456 0.0005678 0.034 14,239.1

Scientific Notation (cont.)

Write the following in long hand.1. 1 x 10-9

2. 3.5 x 105

3. 7.123 x 10-3

4. 5 x 102

5. 4.56 x 10-2

Multiplication w/exponents

Step 1 Multiply coefficients

Step 2 Add exponents

ex. (2 x 102) (2.5 x 105) = 5 x 107

Division w/exponents

Step 1 Divide coefficients

Step 2 Subtract exponents

ex. (5 x 10-2) (1.0 x 107) = 5 x 10-9

Addition & Subtraction w/exponents

All numbers must be written in the same power of 10

ex. 5.8 x 103 + 2.16 x 104

- change to 0.58 x 104 + 2.16 x 104 = 2.74 x 104

Scientific Notation & sig figs

All numbers in front of the x 10 are significant

ex. 2.00 x 102 = 3 sig fig2 x 102 = 1 sig fig

Scientific Notation & calculators

5.44 x 107/8.1 x 104

5.44 (EE or exp) 7 / 8.1 (EE or exp) 4= 6.7 x 102


1. Use dimensional analysis to convert measurements.

2. Convert measurements into scientific notation.

3. Perform mathematical operations using exponents.

Assignment: p. 57 ~ 1-7 and WS

chemistry – chapter 1 & 2

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