Unit 1:
Measuring




Volume
Temperature
Mass
Distance
Volume
Reading the Meniscus
 Always read volume from the bottom of the meniscus. The meniscus is the curved surface of a liquid
in a narrow cylindrical container.
Try to avoid parallax errors.
 Parallax errors arise when a meniscus or needle is viewed from an angle rather than from straight-on at
eye level.
Graduated Cylinders
 The glass cylinder has etched marks to indicate volumes, a pouring lip, and quite often, a plastic
bumper to prevent breakage.
Measuring Volume
 Determine the volume contained in a graduated cylinder by reading the bottom of the meniscus at eye
level.
 Read the volume using all certain digits and one uncertain digit.
 Certain digits are determined from the calibration marks on the cylinder.
 The uncertain digit (the last digit of the reading) is estimated.
Use the graduations to find all certain digits
 There are two unlabeled graduations below the meniscus, and each graduation represents 1 mL, so the
certain digits of the reading are______
Estimate the uncertain digit and take a reading
 The meniscus is about eight tenths of the way to the next graduation, so the final digit in the reading is
_____
 The volume in the graduated cylinder is _______
10 mL Graduate
 What is the volume of liquid in the graduate? ______
25mL graduated cylinder
 What is the volume of liquid in the graduate? ______
100mL graduated cylinder
 What is the volume of liquid in the graduate? ______
Self Test
 Examine the meniscus below and determine the volume of liquid contained in the graduated cylinder.
 The cylinder contains: ______
Temperature
The Thermometer
 Determine the temperature by reading the scale on the thermometer at eye level.
 Read the temperature by using all certain digits and one uncertain digit.
 Certain digits are determined from the calibration marks on the thermometer.
 The uncertain digit (the last digit of the reading) is estimated.
 On most thermometers encountered in a general chemistry lab, the tenths place is the uncertain digit.
Do not allow the tip to touch the walls or the bottom of the flask.

If the thermometer bulb touches the flask, the temperature of the glass will be measured instead of the
temperature of the solution. Readings may be incorrect, particularly if the flask is on a hotplate or in an
ice bath.
Reading the Thermometer
 Determine the readings as shown on Celsius thermometers: 1. ___________
2. __________
Mass
Measuring Mass – Electronic balance
 Our balances read to the hundredths place, the uncertain digit is the hundredths place
_ _ _. _ X
Balance Rules
 In order to protect the balances and ensure accurate results, a number of rules should be followed:
 Always check that the machine is zeroed before adding any substance on it. This is done by
pushing the ZERO button before weighing any substance.
 Never weigh directly on the balance pan. Always use a piece of weighing paper, weigh boat, or
beaker to protect it.
 Do not weigh hot or cold objects.
 Clean up any spills around the balance immediately.
Determining Mass
1.
2.
3.
4.
Place container (weigh boat or beaker) on the pan
Zero the balance by hitting the zero bottom
Place object on pan
Read the digits to the last digit
Metric Ruler
•
•
•
•
Rulers are used to measure distance
Can be used to measure the volume of a cube (L x W x H) length x width x height
Determine the temperature by reading the scale on the thermometer at eye level.
Read the ruler by using all certain digits and one uncertain digit.
Reading a Ruler
• Determine the readings as shown below on centimeter ruler:
• Ruler 1 _________ Ruler 2__________
The SI System of Measurement
The Nature of Measurement
A Measurement is a quantitative observation consisting of TWO parts
Part 1 - number
Part 2 - scale (unit)
Examples:
20 grams
6.63 x 10-34 Joule·seconds
The Fundamental SI Units (le Système International, SI)
Base Quantity
Name of unit
Symbol
Length
Meter
m
Mass
Kilogram
kg
Time
Second
s
Temperature
Kelvin
K
Amount of Substance
Mole
mol
Prefix
Symbol
Multiply the base by
kilo
k
1000
hecto-
h
100
deca-
da
10
unit
m, L, s, g
1
deci-
d
0.1
centi
c
0.01
milli-
m
0.001
micro-
u
0.000 001
SI Prefixes Common to Chemistry
Example 1: Convert 18 liters to milliliters
Example 2: Convert 450 milligrams to grams
Example 3: Convert 20 kilograms to milligrams
Metric Dimensional Analysis Practice
1. Convert 400 mL to Liters
2. Convert 10 meters to mm
3. Convert 73 grams to kg
4. Convert 0.02 kilometers to m
5. Convert 20 centimeters to m
6. Convert 450 milliliters to dL
7. Convert 10 kilograms to grams
8. Convert 935 mg to cg
9. Convert 5.2 kg to mg
10. Convert 175 mL to cm3
Uncertainty and Significant Figures
Uncertainty in Measurement
•
A digit that must be estimated is called uncertain. A measurement always has some degree of
uncertainty.
Why Is there Uncertainty?
 Measurements are performed with instruments
 No instrument can read to an infinite number of decimal places
Precision and Accuracy
• Accuracy refers to the agreement of a particular value with the true value.
• Precision refers to the degree of agreement among several measurements made in the same manner.
Types of Error
• Random Error (Indeterminate Error) - measurement has an equal probability of being high or low.
• Systematic Error (Determinate Error) - Occurs in the same direction each time (high or low), often
resulting from poor technique or incorrect calibration.
Rules for Counting Significant Figures - Details
Nonzero integers always count as significant figures.
• 3456 has 4 significant figures
Zeros
• Leading zeros do not count as significant figures.
0.0486 has 3 significant figures
•
Captive zeros always count as significant figures.
16.07 has 4 significant figures
•
Trailing zeros are significant only if the number contains a decimal point.
9.300 has 4 significant figures
Exact numbers
• have an infinite number of significant figures.
1 inch = 2.54 cm, exactly
Sig Fig Practice #1
How many significant figures in each of the following?
1. 1.0070 m
2. 3.29 x 103 s
3. 17.10 kg
4. 0.0054 cm
5. 100,890 L
6. 3,200,000
Rules for Significant Figures in Mathematical Operations
Multiplication and Division: # sig figs in the result equals the number in the least precise measurement used in
the calculation.
• 6.38 x 2.0 = 12.76  13 (2 sig figs)
Sig Fig Practice #2
Calculation
Calculator says:
Answer
3.24 m x 7.0 m
100.0 g ÷ 23.7 cm3
0.02 cm x 2.371 cm
710 m ÷ 3.0 s
1818.2 lb x 3.23 ft
1.030 g ÷ 2.87 mL
Addition and Subtraction: The number of decimal places in the result equals the number of decimal places in
the least precise measurement.
• 6.8 + 11.934 = 18.734  18.7 (3 sig figs)
Sig Fig Practice #3
Calculation
3.24 m + 7.0 m
100.0 g - 23.73 g
0.02 cm + 2.371 cm
713.1 L - 3.872 L
1818.2 lb + 3.37 lb
2.030 mL - 1.870 mL
Calculator says:
Answer
Scientific Notation
Scientific Notation
• In science, we deal with some very LARGE numbers:
1 mole = 602000000000000000000000
•
In science, we deal with some very SMALL numbers:
Mass of an electron = 0.000000000000000000000000000000091 kg
•
Imagine the difficulty of calculating the mass of 1 mole of electrons!
0.000000000000000000000000000000091 kg
x 602000000000000000000000
???????????????????????????????????
•
A method of representing very large or very small numbers in the form:
M x 10n
• M is a number between 1 and 9.999
•
n is an integer
Example:
2 500 000 000
0.0000579
PERFORMING CALCULATIONS IN SCIENTIFIC NOTATION
ADDITION AND SUBTRACTION
4 x 106
+ 3 x 106
IF the exponents are the same, we simply add or subtract the
numbers in front and bring the exponent down unchanged.
4 x 106
- 3 x 106
The same holds true for subtraction in scientific notation.
4 x 106
+ 3 x 105
If the exponents are NOT the same, we must move a decimal to
make them the same.
A Problem for you…
2.37 x 10-6
+ 3.48 x 10-4
•
•
•
Density
Density is an important intensive property, which can be used to help determine the identity of an
unknown substance.
While the mass or the volume of a substance will vary from sample to sample, the density will remain
the same at a given temperature.
As you know, the density of a substance is a measure of how much mass is present in a given unit of
volume.
Density is the measure of the “compactness” of a material
•
•
•
How close the atoms or molecules are to each other
More than “heaviness” - density includes how much space an object takes up!!
All substances have density including liquids, solids, and gases
𝐷=
•
•
•
•
•
•
𝑚
𝑣
m = mass (g)
Which is the amount of matter in a substance
DO NOT confuse mass with weight
Weight – is the force of gravity exerted on an object
v = volume (ml) or (cm3)
Is measured by water displacement or using a ruler
Question
Which weighs more?
50 kilograms of iron
Or
50 kilograms of feathers
Which has a greater density?
iron
Or
feathers
Water
• Waters density = 1.000 g/ml at 4oC
• Ice floats on water
Densities of Common Substances
If water has a density of 1.00 g/cm3 , which of the following will
float or sink?
Factors that Affect Density
1. Temperature
 For most substances, as temperature increases the volume increases and as a result the density
decreases.
2. Pressure
3. Dissolved solids in Liquids
2 Ways to Measure the Volume of a Solid
1. Objects with regular sides calculate L x W x H using a ruler
2. Water displacement. Place irregularly shaped solid into a volumetric cylinder half filled with water and
place the solid into the cylinder and subtract the volumes
Useful Conversion Factors
1 cm3 = 1 ml
1 dm3 = 1 L
Manipulating the Density Formula
𝐷=
𝑚
𝑣
Sample Problem 1
A student determines that a piece of an unknown material has a mass of 5.854 g and a volume of 7.57 cm3.
What is the density of the material, rounded to the correct number of significant digits?
Sample Problem 2
Iron has a known density of 7.87 g/cm3. What would be the mass of a 2.5 cm3 piece of iron?
Sample Problem 3
Mercury has a density of 13.5 g/cm3. How much space would 50.0 g of mercury occupy?
Percent Error
Students often assume that each measurement that they make in the laboratory is true and accurate. Likewise,
they often assume that the values that they derive through experimentation are very accurate. However,
sources of error often prevent students from being as accurate as they would like. Percent error calculations
are used to determine how close to the true values, or how accurate, their experimental values really are.
The value that the student comes up with is usually called the observed value, or the experimental value. A
value that can be found in reference tables is usually called the true value, or the accepted value. The percent
error can be determined when the true value is compared to the observed value according to the equation
below:
% 𝐸𝑟𝑟𝑜𝑟 =
|𝑦𝑜𝑢𝑟 𝑟𝑒𝑠𝑢𝑙𝑡𝑠 − 𝑎𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑣𝑎𝑙𝑢𝑒|
𝑥 100
𝑎𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑣𝑎𝑙𝑢𝑒
Sample Problem 1
A student measures the mass and volume of a piece of copper in the laboratory and uses his data to calculate
the density o the metal. According to his results, the copper has a density of 8.37 g/cm3. Curious about the
accuracy of his results, the student consults a reference table and finds that the accepted value for the density
of copper is 8.92 g/cm3. What would be the student's percent error?
Sample Problem 2
A student experimentally determines the specific heat of water to be 4.29 J/g x Co. He then looks up the
specific heat of water on a reference table and finds that is 4.18 J/g x Co. What is his percent error?
Sample Problem 3
A student takes an object with an accepted mass of 200.00 grams and masses it on his own balance. He
records the mass of the object as 196.5 g. What is his percent error?
Classification of Matter
Classifying Matter by Composition
Elements- simplest kind of matter, made of one type of atom
An atom is the smallest unit of an element that maintains the properties of that element.
Cannot be broken down into simpler substances by ordinary chemical means
Ex. gold, copper, oxygen (on the periodic table)
Element
Draw Picture
Compounds – matter composed of the atoms of two or more elements chemically bonded
Compounds can be broken down by chemical methods
When they are broken down, the components have completely different properties than the compound.
Ex. Sugar, salt, water, carbon dioxide
Compound
Draw Picture
A mixture is a blend of two or more kinds of matter, each of which retains its own identity and properties.
A mixture is mixed together physically.
Variable composition, often expressed by a percent composition by mass or volume (Ex. 5% salt and 95%
water)
Mixture
Draw Picture
Homogeneous Mixture– matter with a uniform composition
Homogeneous mixtures are also called solutions.
Ex. Salt water and Kool –aid
A heterogeneous mixture is not the same throughout (not uniform).
Ex: M & M’s, Chocolate chip cookie, gravel, soil, rocks such as granite, blood, milk, salad, ocean water, etc.
Classify It
Copper wire
Magnesium
Aluminum foil
Pizza
Table salt (NaCl)
Calcium chloride
Granite
Orange juice
Apple Juice
Club soda
States of matter
Solid- matter that cannot flow and has definite volume and shape
Liquid- definite volume but no definite shape and can flow
Gas- a substance without definite volume or shape and can flow.
Plasma- a substance that is similar to a gas, but loses electrons due to its high temperature
Definite
Volume?
Definite
Shape?
Particle position and
movement
Solid
YES
YES
Packed tightly,
vibrate about fixed
point
Liquid
YES
NO
Close together, can
move past each other
- can flow
Gas
NO
NO
Far apart, move
rapidly - can flow
Drawings
Separating Mixtures
Mixtures are separated by their physical properties.
Primary methods of separating mixtures are:
 Filtration is a method used to separate the components of mixtures that contain an insoluble solid and a
liquid. Example: sand and water

Distillation is a method of separating substances in a mixture by evaporation of a liquid and
subsequent condensation of its vapor. Example: desalination of salt water

Centrifuge
Used to separate solid-liquid mixtures such as those in blood. The centrifuge spins rapidly and
causes the solid to settle to the bottom.
Ex. Separating blood

Chromatography is a method of separating mixtures that uses a stationary phase and a mobile phase.
Paper chromatography can be used to separate pigments because they move at different rates on the
paper.
Properties & Changes of Matter
Properties of Matter
Physical Property- a property that can be observed and measured without changing the identity of the
substance.
Examples? Mass, Density, Melting and Boiling Points
Chemical Property-relates to a substance’s ability to undergo changes that transform it into different
substances.
Examples? Reactivity, Toxicity, and Chemical stability
Chemists use properties to identify and separate matter. More than one property must be used for
identification.
Intensive Properties – do not depend on the amount of matter present
Ex. Melting pt., boiling pt, density, conduct electricity
Extensive Properties – depend on the amount of matter present
Ex. Volume, mass
Changes in Matter
A physical change does not change the composition or identity of the substance.
• Examples?
• Boiled water is still water.
• All phase changes are physical changes
Sublimation is a process in which a solid changes directly to a gas without going through the liquid phase.
Ex: dry ice  CO2
Deposition is a process in which a gas changes directly to a liquid without going through the liquid phase
Ex: liquid vapor to ice (frost on windshields)
A chemical change occurs when one or more substances are changed into new substances.
Reactants- substances that react
Products- substances that form
Products have NEW PROPERTIES
2H2 + O2 → 2H2O
reactants
product
Law of Conservation of mass – Also known as Conservation of Matter. Matter can be neither created nor
destroyed, though it can be rearranged. Mass remains constant in an ordinary chemical change.
Indications of chemical change
1.
2.
3.
4.
5.
Production of energy in the forms of heat, light, sound, or electricity
Production of a gas
Formation of a precipitate
A change in color
A change in odor
What is Energy?
Energy - is the ability to or capacity to do work or to produce change
Conservation of energy (1st Law of Thermodynamics) - Energy can be neither created nor destroyed; the
energy of the universe is constant.
Two types:
Kinetic energy – energy of motion
Potential energy – is the energy due to position of object
Chemical Energy
Chemical energy - is a special kind of potential energy
• Is the energy involved in chemical reactions
Energy Changes
•
•
Some changes in matter release energy.
• For example, the explosion that occurs when hydrogen and oxygen react to form water is a
release of energy.
Heat energy and light energy are released as the reaction takes place.
•
A change in matter in which energy is absorbed from the surroundings is an endothermic process
(heat enters).
• EXAMPLES: melting ice & boiling water
• When barium hydroxide reacts ammonium nitrate are mixed the test-tube feels cold to touch
because energy has been absorbed
•
A change in matter in which energy is released is an exothermic process (heat exits).
• Examples: freezing water & condensation
• Burning of paper gives off heat to the surroundings.
Phase Changes Graphs
What do you notice about the temperature during a phase change?
Phase Diagram




Represents phases as a function of temperature and pressure.
Critical temperature: temperature above which the vapor can not be liquefied.
Critical pressure: pressure required to liquefy AT the critical temperature.
Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).
Phase changes by Name
Phase Change chart for Water
Phase Diagram for Carbon dioxide
Intro to Thermochemistry bHeat and Temperature
Thermochemistry
Thermochemistry – is the study of the transfer of energy as heat that accompany chemical reactions and
physical changes
Heat and Temperature
Heat - is energy transferred between a system and its surroundings
• Heat absorbed or released is measured by a calorimeter
Heat
• Energy that flows from something warm to something cooler
• A hotter substance gives KE to a cooler one
• When heat is transferred (lost or gained), there is a change in the energy within the substance
Questions
A. When you touch ice, heat is transferred from
1) your hand to the ice
2) the ice to your hand
B. When you drink a hot cup of coffee, heat is transferred from
1) your mouth to the coffee
2) the coffee to your mouth
C. When you heat 200 g of water for 1 minute, the water temperature rises from 10°C to 18°C.
If you heat 400 g of water at 10°C in the same pan with the same amount of heat for 1 minute, what would you
expect the final temperature to be?
1) 10 °C
2) 14°C
3) 18°C
Temperature
Temperature – is a measure of the average kinetic energy of particles
• Measured in Kelvin
• Common Temperatures you need to KNOW
Celsius Scale (0C)
– 00C water freezes, 1000C water boils
0
Fahrenheit Scale ( F)
– 320F water freezes, 2120F water boils
Kelvin Scale (K)
– 273 K water freezes, 373 K water boils
– ABSOLUTE Zero 0 K
– Measured in Kelvin in chemistry
K = 273 + oC
Convert the following
K
o
C
50
357
32
298
Some Equalities for Heat
Heat is measured in calories or joules
•
1 kcal = 1000 cal
•
1 calorie = 4.18J
1 kJ = 1000 J
Specific Heat
• Why do some foods stay hot longer than others?
• Why is the beach sand hot, but the water is cool on the same hot day?
Different substances have different capacities for storing energy
It may take 20 minutes to heat water to 75°C. However, the same mass of aluminum might require 5
minutes and the same amount of copper may take only 2 minutes to reach the same temperature.
Specific heat - is the amount of heat needed to raise the temperature of 1 g of a substance by 1°C
water
aluminum
copper
silver
gold
cal/g°C
1.00
0.22
0.093
0.057
0.031
J/g°C
4.18
0.90
0.39
0.24
0.13
Questions
A. A substance with a large specific heat
1) heats up quickly
2) heats up slowly
B. When ocean water cools, the surrounding air
1) cools 2) warms
3) stays the same
C. Sand in the desert is hot in the day, and cool at night. Sand must have a
1) high specific heat 2) low specific heat
Measuring Heat
Requires
•
Grams of substance
•
Temperature change T

T = Tf - Ti
Specific heat of the substance
Calculating Heat
Energy Lost or Gained = mass x temp. change x specific heat
q = m x ∆ T x Cp
A few key ideas:
• If a substance receives heat and experiences an increase in temperature then Q is a positive number
and ∆T is a positive number.
• If a substance loses heat and experiences a decrease in temperature then Q is a negative number and
∆T is a negative number.
• Q, heat energy, can be measured in either Joules or calories. Just make sure that your units for c are
consistent with your units for Q.
• ∆T, change in temperature, can be measured in K, °C, or °F. Just make sure that your units for c are
consistent with your units for ∆T.
•
•
•
Always start a problem by listing the given information (with units) and writing down the specific heat
capacity equation.
You must ALWAYS show all work and make sure you have consistent units on your final answer.
The First Law of Thermodynamics states that if two substances exchange heat, the quantity of heat
gained by one substance is exactly equal and opposite to the quantity of heat lost by the other
substance.
Problem
The element hydrogen has the highest specific heat of all elements. At a temperature of 25°C, hydrogen’s
specific heat capacity is 14300J/(kg ∙ K). If the temperature of a 0.34 kg sample of hydrogen is to be raised by
25 K, how much heat will have to be transferred to the hydrogen?
A 0.59 kg brass candlestick has an initial temperature of 98.0°C. If 21,100J of heat is removed from the
candlestick to lower its temperature to 6.8°C, what is the specific heat capacity of brass?
A 0.38kg drinking glass is filled with a hot liquid. The liquid transfers 7032J of heat to the glass. If the
temperature of the glass increases by 22K, what is the specific heat capacity of the glass?