So, what is Chemistry?
What does the word “chemistry”
make you think?
This?
Or this?
Or maybe this?
And possibly this..
But, Chemistry really is…
Baby stuff!
And…
• Cool cars!
And…
• Things you need
And…
• Things that you really need
And..
• Things that some of us really, really need..
And…
• Fun stuff…
But, seriously..
• How
would we
survive
without it?
This is a part of chemistry
So is this..
And this..
Chapter 1 Notes
• Math requirements for general
chemistry:
1. Significant figures and usage
2. Scientific notation and usage
3. Basic algebra
4. Dimensional analysis
5. Graphing and interpretation
LO’s for SigFigs and SciNot
• Explain the meaning of and identify
significant figures in experimental data.
• Correctly perform calculations using
significant figures
• Convert any number to or from scientific
notation
• Correctly perform calculations using
scientific notation with a calculator.
Significant Figures
Significant Figures are affectionately
referred to as SigFigs by chemists.
SigFigs are used as a method to convey
how precisely a measured quantity is
known. SigFigs can tell you how carefully
a quantity was measured.
All experimental data has some uncertainty
associated with it, usually due to the
limitations of the equipment used to obtain
the data.
The mass of a beaker was
measured on 3 different balances.
Which value is correct?
Which digit is uncertain?
2.3 g
2.31 g
2.3145 g
What would the 3 balances read if you
had a beaker with a mass of exactly
2500 g?
Rules for Finding SigFigs
1. All non-zero digits are always Sig.
2. All captive zeros are always Sig.
3. Leading zeros (to the left) are Never
Sig
4. Trailing zeros (to the right) are only
Sig when the number has a decimal
point.
5. Counting numbers have infinite
SigFigs.
Practice- How many SigFigs?
4509 -____
1.007 - ____
4510 -____
10.0 - ____
12,000 -_____
0.00501 - ____
120,020 -_____
0.02100 - _____
Doing math with SigFigs
Multiplication and division:
- Count the number of SigFigs in each
factor. Your answer will have the smallest
number of SigFigs in the factors.
- Examples: 2.000 x 5.00 = 10.0
- 3.14 x 2.2 = 6.908 = 6.9
- 1500 / 6.24 = 240.3846 = 240
- 100000 / 75.202 = 1329.752=1000
Addition and subtraction:
Add or subtract the numbers as normal.
To determine the number of sigfigs, start
from the left end of your answer and
check, place by place (hundreds, tens,
ones, tenths) if there are sigfigs in that
place for all factors. When 1 factor runs
out of sigfigs, all sigfigs to the right in your
answer are dropped. (Ignore empty
spaces to the left of the decimal point).
Examples: 2010
41.75
+2.1
2053.85 = 2050
1.204
0.0072
+5.25
6.4612 = 6.46
200,000
-10.0
199,990.0=200,000
Scientific notation
Serves two purposes:
1. To simplify writing very large (6.02 x 1023) and
very small (1.6 x 10-26) numbers.
2. To clarify confusion about the number of SigFigs
in a value.
2100 versus 2100.
1,000,000 vs 1,000,000.
For numbers larger than 1: Re-write the number dropping
all zeros that are not sigfigs. Put in a decimal place so
that there is one and only one digit to the left of the
decimal place. Write “x10” after the number.
•
Ex: 107,000 = 1.07 x 10
Count the number of places that the decimal place must be
moved to get from the original number to the re-written
number. Write the number of places moved as the
exponent for the 10.
•
Ex: 107,000 moves 5 places so we write 1.07 x 105
For numbers less than 1: The rules are exactly the same
except that the exponent is written as a negative
number.
•
Ex: 0.00107 moves three places so we write1.07 x 10-3
Practice:
12,000 =
6.1 x 104 =
1,060,000 =
2.07 x 107=
0.0021 =
4.04 x 10-3=
0.0003050 =
2.90 x 10-4=
• Do this calculation with your calculator!!!
6.25x108 / 1.70x10-4
SciNot Using Calculators
Do not use the “carrot key” Ʌ
Enter the number (ex: for 6.02 x1023)
6.02
Press 2nd EE (on the TI30)–do not enter x 10
Enter the exponent
23
Your calculator should read
6.02E23
The “E” means “x10 to the”
Practice
6.02x1023 / 4.5x1012 =
1.75x10-5 x 2.17x108 =
(4.05x106) (3.1x10-7) / 8.82x1015 =
Basic Algebra
In chemistry, you will often have problems
where you must solve an algebraic equation
for an unknown. You must be familiar with
how to do this.
Brief review:
M
D=
V
Solve for M
Solve for V
To solve for an unknown, you must get the
unknown by itself, on top, on one side of
the equal sign.
Move either the unknown or the other
variables until solved.
When moving variables (multiplication and
division) a variable on top (numerator) on
one side of the equals goes to the bottom
(denominator) on the other side. A variable
on the bottom goes to the top when moved
to the other side.
Practice:
ax =
b
y
t
2.5z =
15
PV=nRT
Finally, Chemistry!!!!
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•
•
•
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Chemistry Learning objectives Chapter 1-B
Differentiate between and discuss characteristics of 3
common states of matter
Model structure of 3 states of matter
Identify products and reactants in a chemical reaction
Differentiate between and give examples of chemical
and physical changes
Differentiate between and give examples of endothermic
and exothermic reactions
Discuss energy changes in products and reactants for
exo and endothermic reactions
• Chemistry is defined as the study of
matter and the changes matter undergoes.
• Matter – Anything that has mass and
volume (takes up space)
• Essentially, chemistry is the study of
everything in the universe and what
happens to it
Physical States of matter
Solid
Particles closely
packed
Liquid
Gas
Closely packed Very loose
Strong attraction Medium
between particles attraction
Weak to no
attraction
Fixed shape
No fixed shape
No fixed shape
Fixed volume
Fixed volume
No fixed volume
Physical change – A change where the chemical
identity of the material does not change.
Ex:
Chemical change – A change where new
substances are formed. The chemical identity
does change.
Ex:
Reversible and not reversible are not good
indicators of physical and chemical
changes
Chemical Reactions
In any chemical reaction, atoms are rearranged to form different substances!
No new atoms are formed – no atoms are
destroyed.
All reactions are written as:
Reactants → Products
Reactants – left side of arrow – what you are
starting with.
Products – Substances that are produced in
the reaction. Right side of the arrow.
→ - The action!!! Can be read as “yields”
“reacts to form” “produces” “gives”
“decomposes to”
Examples:
Exothermic reactions: Energy is released as
a products of the reaction. Energy stored
in the reactants is released to the
surroundings (beaker gets HOT).
Example:
Endothermic reactions: Energy is absorbed
from the surroundings and added to the
reactants. (beaker gets cold)
Example:
Graphs:
Signs of a chemical reaction
1.
2.
3.
4.
5.
Giving off a gas.
Precipitation – Formation of a solid.
Change in heat. Absorbed or given off.
Production of light.
Change in color.
Units and the SI system
In science, values or numbers do not occur
by themselves. A number must have a unit
with it to give the number meaning.
Think of 3!!!!
Seven SI base units
•
•
•
•
•
The ampere (A) - unit of measurement of electric current
The kilogram (kg) - unit of measurement of mass
The metre (m) - unit of measurement of length
The second (s) - unit of measurement of time
The kelvin (K) - unit of measurement of thermodynamic
temperature
• The mole (mol) - unit of measurement of amount of
substance
• The candela (cd) - unit of measurement of luminous
intensity
Units that you should be familiar with
Milliliter (mL)– volume – about _____ drops
Liter (L)– volume – a little more than __ quart
Gram (g)– mass – a penny has a mass of __ g
Kilogram (kg)– a little more than ___ pounds
Meter (m)– distance - a little more than a yard
Centimeter (cm)– distance - a little less than
____inch
Dimensional Analysis
A method of calculation that uses the UNITS
of the values and conversion factors to
set up the calculation.
A conversion factor is any mathematical
relationship between two units. Used to
convert from one unit to another.
Examples:
How to do Dimensional analysis:
1. In the problem, find the UNITS that you are
looking for and the UNITS that you are given.
2. Find an appropriate conversion factor (or
series of factors) that will allow you to relate
given units to looked for units.
3. Draw a grid and place the given unit in the
upper left. Enter conversion factor into the grid
so that units cancel (top and bottom) until only
the looked for unit is left.
4. Put in the numbers in appropriate places and
run them thru your calculator for the answer.
Example 1: An elephant weighs 12 tons.
How many pounds is this?
Looked for unit –
Given unit –
Conversion factor –
Example 2:
The fence in a baseball field in 330 ft from
home plate. How far is this in yards?
Looked for units –
Given units –
Conversion factor -
Example 3:
A month is 30 days long. How long is this in
hours? In minutes? In seconds?
Looked for units –
Given units –
Conversion factor -
How old are you in seconds?
SI unit conversions – The prefix tells you the
conversion factor. All prefixes are used the
same way.
Ex:
Milli means 1/1000 (10-3). This means that:
There are 1000 milligrams in 1 gram
Or
1 milligram is 1x10-3 grams
1000 mg
1g
1g
1 mg
1000 mg
10-3 g
Practice
Properties of matter
Properties are characteristics or descriptions
of matter.
1. Chemical properties – Can only be
observed when a substance changes.
Tells how a substance reacts with others.
2. Physical properties – Can be measured
or observed without changing the
substance.
Chemical property examples:
Physical property examples:
• Chemical change – Produces new
substances. The chemical identity of
substances change.
• Examples:
• Physical Change – Only affects the
physical properties. No new substances
formed.
• Examples:
Density-A physical property
• Density is defined as the amount of mass per
unit of volume.
• Can be thought of as the amount of matter
packed into a given volume.
• Most common units are g/mL
• May often see g/L or g/cm3
D=
M
V
Notice that density relates the mass of a
material to the volume. It can be used as a
conversion factor in dimensional analysis
to convert between mass and volume.
Practice 1: 50.0 mL of gold has a mass of
965g. What is the density of gold?
Practice 2: Helium has a density of 0.179
g/L. A balloon that holds 2.5 L of He would
have what mass of He?
Practice 3: A chunk of gold that weighs 2 lbs
would take up how much space?
(1 lb = 454 g)
The Nature of Matter
Atom – The basic building block of matter.
Think Legos!!!! There are 92 different
kinds of atoms that occur naturally on
earth. (We will have much, much more to say about this!!!!)
Element – A substance that is composed of
only 1 type of atom.
Cannot be broken down into a simpler
type of matter by any chemical means.
• Molecules – Two or more atoms that are bonded
together. (Snapping Legos together). If the same
type of atoms are bonded then the molecule
makes up an element.
• Compounds – Two or more different types of
atoms bonded together.
Can be broken down into simpler substances by
chemical change.
Properties are different than those of component
elements.
Always a fixed ratio of elements (must have a
formula).
• Mixtures – Two or more compounds
and/or elements physically mixed but not
chemically bonded.
Can be separated by physical means.
Keeps the properties of the component
substances.
Homogeneous mixtures – Same throughout
Heterogeneous mixtures – Different
components are apparent.
Matter
Pure
substance
Element
Compound
Mixture
Homogeneous
Heterogeneous
• Some elements are diatomic – They occur
in nature as molecules composed of 2
atoms.
You should remember these 7 Diatomics!!!
Bromine
Oxygen
Fluorine
Iodine
The BOFINCH elements
Nitrogen
Chlorine
Hydrogen
• Allotrope – Different molecular or
crystalline form of the same element.
Examples: