CHEMISTRY – Chapter 1 & 2
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Transcript of 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
Chapter 1 – Section 2
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
Review and Assignment
1. Distinguish between a mixture and a pure substance.
2. Define what matter is.
Assignment: WS
Chapter 1 – Section 2
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
Review and Assignment
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 -
Review and Assignment
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
Review and Assignment
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
Review and Assignment
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 = _____
Review and Assignment
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
Review and Assignment
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
Review and Assignment
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