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Chapter 1 - Notes
1-0 Introduction
A) Main idea of section:
B) Chemistry is:
C) Examples of branches of chemistry:
D) Basic questions of chemistry:
1-1 Matter and Energy
A) Definitions:
1. Matter
2. Mass
3. Energy
4. Kinetic energy
5. Potential energy
6. Exothermic
7. Endothermic
B) Law of Conservation of Matter
C) Law of Conservation of Energy
D) Why have a “Law of Conservation of Matter and Energy”?
APChem,6/25/2017
1-2 States of Matter
A) Complete the following drawing: (label phase transitions)
B) For each phase, sketch a kinetic molecular model (particle model) (see figure 1-3) (also
see http://phet.colorado.edu/en/simulation/states-of-matter and
http://www.chemteam.info/Matter/StatesofMatter.html )
Phase of matter
Solid
Liquid
Gas
Type of Motion
Thermal expansion
Definite volume
Definite shape
Compressibility
Density
C) Other states?
p. 2
1-3 Chemical and Physical Properties
A) Main idea of section:
B) Definitions:
1. Chemical properties:
 examples?
2. Physical properties:
 examples?
3. Extensive properties
4. Intensive properties
1-4 Chemical and Physical Changes
A) Chemical change:
 examples?
B) Physical change:
 examples?
1-5 Mixtures, Substances, Compounds, and Elements
A) Main idea of section:
B) Definitions:
1. mixture
2. substance:
3. heterogeneous mixture:
 examples?
p. 3
4. homogeneous mixture
 examples?
5. compund
 examples?
6. element
 examples?
C) Graphic Organizer: Matter (Figure 1-7)
D) Separation methods (NIB) – see
http://www.chemteam.info/Matter/SeparationOfMixtures.html
E) Law of Constant Composition:
 Ex:
p. 4
1-6 Measurements in Chemistry
A) Metric prefixes to know:
Prefix
Abbreviation
Meaning
Example
megakilodecicentimillimicronanopico1-7 Units of Measurement
A) Mass & Weight
1. Difference?
2. Unit (mass)
B) Volume:
1 L = ______ mL = ______ dm3
1 mL = _____ cm3
1-8 Use of Numbers – Significant figures (try the links for practice problems, then complete the
notes) This is a very useful website for this – go to ALL the sections and be sure you
understand them – this entire section is here on-line
http://www.chemteam.info/SigFigs/SigFigs.html
A) Significant Figures in Measurements (Fig 1-15)
1. In measurements, you will need ______ estimated digit(s)
2. So the reading will be based on the smallest increment + one estimated digit (smallest
increment = smallest measurement between two adjacent lines/marks)
3. Example:
1
2
 Smallest increment: 0.1 cm
 So: reading for the above arrow 1.52 cm (or 1.53 cm) the 2 is the estimated digit
 Note: if the measurement device has a “readout” (for instance, the balances), record
the measurement AS GIVEN by the device
p. 5
cm
B) Counting sig figs
1. Two types of numbers:
 exact: either counting numbers (3 dogs, 75 eggs) or (many) conversion factors
(100 cm = 1 m) – these have ______________ significant figures
2. Measurements:
 All nonzero digits ___________ significant
189.7 (___ sf)
1.276535 (__ sf)
 All zeros between nonzero digits __________ significant
1000.07 (___sf)
104.2 (___sf)
 Zeros before the first nonzero digit ___________ significant
0.0027 (___sf) 0.75 (___sf)
 Zeros after the last nonzero digit ________ significant if a decimal is present
 Zeros after the last nonzero digit ________ significant if a decimal is not present
122.70 (__sf) 1240 (___sf)
 Note: If the number is in scientific notation – use the same rules as ___________
C) Rounding (note: not the way you learn in elementary school!) Not In Book (NIB)
1. Use the digit just after the final decimal place you want to keep
2. If that digit is
 0-4 – the last digit __________________
 6-9 – the last digit __________________
 5 - the last digit ____________________
3. Examples: Round to the hundredths place
0.324 = ________
0.325 = ________
0.399 = _________
0.326 = ________
0.335 = ________
0.395 = _________
D) Calculations:
1. General: Do all calculations, keeping track of the significant figures as you go, but do
not round until the end
2. Note: the basis for this method is that you cannot have more than one estimated digit.
3. Addition/Subtraction
 rule:
 Examples:
p. 6
(1) 0.301 + 1.44 + 0.2 = 1.941 = _______________
(2) 144.60 + 0.3334 + 10.7865 = 155.7199 = ________________________
4. Multiplication/Division
 rule:
 Examples:
(1) 0.301 x 1.44 x 0.2 = 0.086688 = ________________________
(2) 144.60 x 0.3334 x 10.7865 = 520.013282 = ______________________
5. Mixed example: If you do an experiment, and the theoretical value is 5.000, and your
experimental result is 5.020, what is your percent error?
1-9 The Unit Factor (Dimensional Analysis) – AKA Factor Label
A) What is a “unit factor”?
B) Method:
1. CAUTION!!! You must have units on all numbers, even intermediates!
2. Know the unit you have for your initial value
3. Use unit factors to replace the units you don’t want with units you do want
4. Check that the units cancel!!!
5. Sig figs on the answer are generally what you start with!
C) Examples
1. convert 3.0 ft to inches
2. convert 2.55 miles to inches (1 mile = 5280 ft.)
3. convert 0.75 L to cm3
4. convert 15 cm3 to in3
p. 7
1-10 Percentage
A) Percentage can be used as a conversion factor
B) Percentage can be “interpreted” several different ways:
C) Percentage should be specified (by mass, by volume, etc)
1-11 Density and Specific Gravity
A) Density =
B) Density is
 intensive, and frequently used to help identify a substance
 a function of temperature
 a conversion between _____________ and _____________
C) Specific gravity is
D) Note: density of water at 3.98 ºC = 1.000 g/mL
1-12 Heat and Temperature
A) Temperature is: A measure of the kinetic energy of a substance
B) Conversion formulas:
1. Between T (in K) and T (in °C): _______________________________
2. Between T (in °C) and T (in °F): _________________________________
p. 8
C) Temperature scales:
ºF
ºC
K
Normal bp of water
“Normal” body temperature
“Room temperature”
98.6
70
Normal mp of water
1-13 Heat Transfer and the Measurement of Heat we will do this section together – skip for now
A) Units:
SI: Joules (J), also use calorie
B) Calorie definition:
the amount of energy required to raise the temperature of 1 g of water by 1 Co
1 calorie =
4.184 Joules
EXACT!
note: food calories: Calories = 1000 calories = 1 kcal
1 g protein = 4 Cal, 1 g carbohydrate = 4 Cal, 1 g fat = 9 Cal
C) Calorimetry - measurement of the heat released or absorbed in a process (physical or
chemical change)
 Equation: q = m Cp T, where
q = heat
Cp = specific heat
..
For water:
Cp = 1.000 cal/g°C
m = mass
T = change in temperature
or 4.184 J/g°C for water
D) Example: do ex. 1-19
p. 9