Transcript Chapter 1 notes

Chemistry
Chapter 1
The Science of Matter
1.1 Objectives
• Students will be able to…
– Understand matter, its phases, and
composition
– Know the differences quantitative and
qualitative expressions
– Explain the if something is a substance or
a mixture
– Know the differences between
homogeneous and heterogeneous
mixtures
I. What is Chemistry?
The study and investigation of the
structure and properties of matter
II. A Picture of Matter
A. composition, structure, and behavior
1) matter— a substance that takes up
space and has mass
2) mass— the amount of matter in an
object
3) properties of matter— characteristics
and behavior; can be physical or
chemical
III. Using Models in Chemistry
Scientific model— tools to help you
understand the relationship between the
macroscopic and submicroscopic views of
matter
II. A Picture of Matter
B. examining matter
1) macroscopic view— matter large
enough to be seen
2) submicroscopic view— dealing with
atoms
1.5 x10-10 m
IV. Classifying matter
A. classification by composition:
1) qualitative —verbal description, no exact
figures
2) quantitative —description by numbers
IV. Classifying matter
B. Pure substance vs. Mixture
1) pure substance—matter with the same
definite composition and properties
examples of elements and compounds:
C H2O NaCl NH3
2) mixture—physical blend of two or more
substances
IV. Classifying matter
C. Mixed matter
1) Mixtures
heterogeneous— not uniform
-has different “phases”
homogeneous—uniform
-has one “phase”
-a called a solution
What Kind of Mixture is it?
Review Questions
• What three characteristics of matter
does chemistry deal with?
• The population of India is….
Is this a quantitative or qualitative
expression? Why?
• How does a mixture differ from a
substance?
• Select a common item and describe
chemically using the proper terms.
IV. Classifying matter
2) Phase
-solid, liquid, gas, plasma
IV. Classifying matter
States of matter
solid
a) matter with a definite, fixed shape and
volume
IV. Classifying matter
States of matter
liquid
a) matter with variable shape and fixed
volume
b) exhibits flow
c) takes the shape of its
IV. Classifying matter
States of matter
gas
a) matter with variable shape and volume
b) exhibits flow
c) takes the shape and volume of its container
d) normally stable as a gas at room temperature
e) vapor is a gas even though the normal state is not
f) volatile—changing to a gas easily at room
temperature
IV. Classifying matter
States of matter
Plasma- ???
- low-density ionized gas
IV. Classifying matter
3) Solutions
solute— the substance being dissolved
solvent— the substance doing the dissolving
aqueous solutions (aq)—water containing
dissolved materials
Can be separated by physical means:
evaporation, filtration, distillation, etc.
Evaporation
Filtration
Distillation
IV. Classifying matter
D. Separation of matter into pure substances
physical change
a) alterations that do not change the substance’s
identity and composition
b) boil, freeze, melt, condense, dissolve, crush,
break, cut…
physical properties
conductivity
density
solubility
melting point
ductility
boiling point
malleability
odor
refractive index
IV. Classifying matter
Intensive property- does not depend on the
amount of matter
Extensive property- depends on the
amount of matter
V. Substances: Pure Matter
V. Substances: Pure Matter
A. elements
simplest form of matter retaining the properties of that
matter
examples : Ag Pb O W
COMMON ELEMENTS TO KNOW:
Ag, Al, Ar, As, Au, B, Ba, Be, Bi, Br, C, Ca, Cl, Co, Cr, Cs, Cu, F, Fe, Fr, H,
He, Hg, I, K, Kr, Li, Mg, Mn, N, Na, Ne, Ni, O, P, Pb, Ra, Rb, S, Sb, Si, Sr, Sn,
U, W, Zn
B. chemical symbols:
-each element has a different symbol
-capitalize the first letter only
-word roots from English and other languages (Latin)
Periodic Table of Elements
V. Substances: Pure Matter
C. compounds
-more than one element in a type of matter
-can only be separated by chemical methods
examples: NaHCO3 CO H2CO CaCO3
D. formulas of compounds
-formula—correct combination of chemical
symbols
H2O
H2SO4
Jimmy took a drink but he will drink NO MORE,
because what he thought was H2O was H2SO4!!!
Review Questions
• What three characteristics of matter
does chemistry deal with?
• How does a mixture differ from a
substance?
• How does a compound differ from a
mixture?
• Select a common item and describe
chemically using the proper terms.
Homework
• Read section 1.1 and 1.2
• Copy Vocabulary into your notebooks
1.1 Objectives
• Students will be able to…
– Understand the law of the
conservation of matter
Review Questions
• Is milk a homogeneous or
heterogeneous mixture? Why?
• Forms of energy like sunlight and heat
are not matter. Why?
Ch. 1 The Science of Matter
Section 1.2
1.2 Objectives
• Students will be able to…
– Distinguish between physical and
chemical properties
– Contrast physical and chemical
changes
– Apply the law of conservation of
matter
– Complete density and dimensional
analysis calculations
Warm Up Questions
• What is the difference between and
element and a compound?
• Describe the picture. Are these
quantitative or qualitative descriptions?
• What is the chemical formula of the
chemical model?
– Black= Carbon (C)
– White= Hydrogen (H)
– Red= Oxygen (O)
VII. Chemical Properties and
Changes
A. Atoms and chemical change
1) chemical properties—ability to form
new substances as a result of chemical
reactions
2) chemical changes—alterations that
changes a substance’s identity and
composition to something new,
through a chemical reaction
VII. Chemical Properties and
Changes
3) chemical reactions—the changing of
substance(s) into new ones
a) reactants—starting substances in a rxn.
b) products—new substances formed in a rxn.
c) clues that a chemical rxn. has occurred
energy is given off (gets hotter)
color change
production of a gas
energy is absorbed (gets colder)
odor change
production of a solid (precipitate; ppt.)
A Chemical Reaction
H2SO4 + 2NaOH -> Na2SO4 + 2H2O
Reactants
Products
VII. Chemical Properties and
Changes
4) Law of Conservation of Mass—
matter cannot be created nor
destroyed, it merely changes
form
(burn a candle)
a) reactant mass = product mass
b) exceptions are nuclear rxns.
VII. Chemical Properties and
Changes
B. chemical reactions and energy
1) energy—the capacity to do
work
2) some types of energy
a) potential—energy at rest;
energy
of position
b) kinetic—energy of
motion
c) thermal—heat energy
d) radiant—light energy
e) chemical—energy in
chemical
bonds
VII. Chemical Properties and
Changes
3) Law of Conservation of Energy:
in a physical or chemical
change, energy cannot be
created nor destroyed, it merely
changes form
4) energy changes in reactions
a) exothermic - giving off
heat
b) endothermic - absorbing
heat
SI Units
• SI stands for “le Systeme International
d’Unites”
• France, 1789.
• There are two types of units:
– base units;
– derived units.
• There are 7 base units in the SI system.
• Scientific (Exponential) Notation is used for
convenience with smaller or larger units in
the SI system.
SI Base Units
SI Prefixes
VI. Identifying Matter by its
Properties
B. Density
1) density = mass (g) / volume (ml)
2) D = M / V
3) density usually decreases as temp.
increases (due to increased volume)
Derived Units
• Derived units are obtained from the 7 base
SI units.
• The are combined by multiplication or
division
• Examples: km/hr
g/mL
m/s
cal/g/oC
Density Problems:
• A metal bar has a mass of 35.50 g and a
volume of 262 cm3. What is its density
in g/cm3?
D=M/V
D= 35.50 g / 262 cm3
D= 0.135g/ cm3
Density Problems:
• 500.0 mL of a liquid has a density of
0.447 g/mL. What is its mass?
D=M/V
M= DV
M= (0.447g/ml )(500ml)
M= 224g
Density Problems:
• 4.2 g of a substance has a density of
0.89 g/m3. How much space, in m3, does
it occupy?
D=M/V
V=M/D
V=4.2g / (0.89g/m3)
V=4.7m3
Volume
• The units for
volume are given
by (units of
length)3.
SI unit for volume is 1 m3.
• We usually use
1 mL = 1 cm3.
• Other volume units:
1 L = 1 dm3 = 1000
cm3 = 1000 mL.
Volume
Uncertainty in Measurement
• All scientific measures are subject to error.
• These errors are reflected in the number of
figures reported for the measurement.
• These errors are also reflected in the observation
that two successive measures of the same
Precision and Accuracy (often confused)
• Measurements that are close to the “correct”
value are accurate.
• Measurements that are close to each other are
precise.
Precision and Accuracy
Significant Figures (sig-figs)
• The number of digits reported in a
measurement reflect the accuracy of the
measurement and the precision of the
measuring device.
• Report the fewest significant figures
• Fewest number for multiplication and
division
• Fewest decimal places for addition and
subtraction
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Significant Figures (sig-figs)
Non-zero numbers (e.g. 1, 2, 3…9) are
significant.
Zeros between non-zero numbers are always
significant. (e.g. 204 ml) (the sandwich rule)
Zeros before the first non-zero digit are not
significant. (e.g. 0.0003 has one.)
Zeros at the end of the number after a
decimal place are significant. (e.g. 123.00 g)
Zeros at the end of a number before a
decimal place are ambiguous (e.g. 10,300 g).
Examples:
How many sig-figs are in each of the following?
4
10 m
= 3; 5.23 x
2. 0.000487 kg = 6; 4.87 x 10-4kg
3. 29.0400 s
= 7; 2.90400 x 10s
4. 507 people = 3; 5.07 x 102 people
5. 230,050 cm = 5; 2.30050 x 105 cm
= 5; 4.5600 x 10 L
6. 45.600 L
1. 52300 m
Scientific Notation
• Numbers written in scientific notation
include a numeral with one digit before the
decimal point, multiplied by some power of
ten (6.022 x 1023)
• All digits are significant.
1. Convert to scientific notation:
a. 450 000 000
b. 0.000 000 047
c. 46.04
2. Convert to non-scientific notation:
a. 7.09 x 10-6
b. 3.39 x 105
c. 8.00 x 106
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Dimensional Analysis
Method of calculation using a knowledge of
units and some basic algebra.
Given units can be multiplied or divided to
give the desired units.
Conversion factors are used to manipulate
units
Treat units just like you would treat numbers
Given unit  (conversion factor) = Desired
unit
The conversion factors are ratios equal to
“1”
Using Two or More Conversion Factors
• To convert length in meters to length in
inches:
• Given: 1 m = 100 cm and 1 inch = 2.54 cm
• What are the conversion factors for these
two?
• How many inches are in 1.5 m?
Using Two or More Conversion Factors
• In dimensional analysis ask three
questions:
• What data are we given? (where are we?)
• What quantity do we need? (where do we
want to be?)
• What conversion factors are available to
take us from what we are given to what
we need? (how do we get there?)
Examples: Convert the following:
1. 240 000 000 cm to km.
2. 0.00657 m3 to cm3.
3. 2.57 years to s.
4. 4.98 feet to mm.
Homework
• Read Chapter 2
• Copy Vocabulary into your notebooks
• Worksheet 1.1
Stuff to add