Chapter 1:
Chemistry & Measurement
Classification of Matter
Chemical Composition
pure substances:
only one type of matter
definite composition
distinct properties
two categories of pure substances:
elements - cannot be broken down into simpler substances
by chemical change
118 known; 112 named
compounds - 2 or more elements chemically combined
(bonded together); have fixed composition
Law of Definite Proportions (Joseph Proust):
A specific compound must always be composed of the same
proportion of its elements by mass.
mixtures:
more than one type of matter
more than one substance physically combined
have variable composition
homogeneous mixtures
heterogeneous mixtures
solids: rigid
ordered arrangement or particles
fixed volume and shape
not compressible
particles very close together
liquids: fluid (flow)
fixed volume but variable shape
very slightly compressible
short range order; short range motion
gases: fluid
least dense form of matter
shape and volume are variable
highly compressible (lots of empty space)
constant, random, chaotic motion
physical property or change:
describes or involves only a change of phase
ex. sublimation of dry ice CO2 (s) ! CO2 (g)
chemical property or change:
describes or involves change of chemical identity
ex. hydrogen combines with oxygen to form water
intensive property:
independent of the amount of sample
ex. mp of water is 0oC (at 1 atm)
density
molar mass
extensive property:
depends on the size (extent) of the sample
ex. mass; volume; heat of combustion
Uncertainty in Measurements
Uncertainty in Measurements
Precision
How close are values
to one another?
standard deviation
vs.
Accuracy
How close is your
experimental value to
the known value?
percent error
Significant Figures - all certain digits plus the first
uncertain digit
How many significant figures do you record when
making a measurement?
How many significant figures are in a number?
How many significant figures do you record in the
answer to a calculation?
How many significant figures are in a given number?
1. Non-zero integers are always significant.
2a. Leading zeros are not significant;
2b. Captive zeros are significant;
2c. Trailing zeros may or may not be significant;
(trailing zeros are always significant if the
number contains a decimal point)
2d. When in doubt, use scientific notation.
3. Exact numbers have an infinite number of
significant figures.
examples: counting, definitions, integers
Dimensional Analysis and Unit Conversions
dimensional analysis - carry numbers and units
through algebraic manipulations; treat the unit
itself as an algebraic entity
ex. (2x)2 = 4x2;
(4cm)2 = 16cm2
unit conversions - convert quantities from one
unit scale to a different unit scale using one or
more conversion factors
ex. 125.0 in = ??? cm
conversion factor - statement of equality between
unit scales
ex. 1 in = 2.54 cm
How many significant figures do you keep as an
answer to a calculation?
1. Multiplication and Division:
number of significant figures in the answer should
be the same as the least number of significant
figures in the data
2. Addition and Subtraction:
number of decimal places in the answer should
be the same as the least number of decimal
places in the data
Unit Conversions: Intrasystem Conversions
metric system conversions
English system conversions
examples:
54.5 ng = ________ pg
25.0 mi = ________ ft
(1 mi = 5280 ft)
Unit Conversions: Intersystem Conversions
metric " English system conversions
example:
115 mm = _______ ft
(1 in = 2.54 cm)
Unit Conversions: Combined Unit Conversions
examples:
55 mi/h = ________ m/s
(1 mi = 1.6093 km)
2580 cm2 = ________ m2
2580 cm2 = ________ in2
12.4 g/cm3 = ________ kg/m3
Density as a Conversion Factor:
density = mass/volume
examples:
Calculate the density of a liquid if a 43.7 g
sample occupies a volume of 55.7 mL.
The density of an alloy is 6.286 g/cm3.
Determine the mass of a spherical sample of
this alloy if the sphere’s radius is 7.84 mm.
Temperature Conversions:
Fahrenheit, Celsius, and Kelvin Temperature Scales