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Chemistry - Unit 1

The Language of Chemistry

Mark Holcomb

on 19 August 2014

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Transcript of Chemistry - Unit 1

Unit 1:
The Language of

Scientific Method
Process used to solve problems and answers questions.
1. Questioning and observation
2. Hypothesis - educated prediction
3. Experiment and gather data
4. Interpret data to form a conclusion
More about the Scientific Method
Variable – what is being tested
Controls – aspects of the experiment kept the same
Experimental group – group that contains the variable
Control group – group that does not contain the variable; used for comparison to the experimental group
Only ONE variable can be tested at a time
Common Terms
Energy – the capacity to do work or produce heat. All chemical reactions either require energy or give off energy.
Law of Conservation of Energy – energy is neither created nor destroyed but it can be transferred.
Matter – anything that has mass and volume
Law of Conservation of Matter – matter is neither created nor destroyed but it can be transformed.
Chemistry – study of matter and the changes it undergoes
Volume – amount of space an object takes up
Mass – amount of matter in an object
Weight – force of gravity on an object
We consider mass and weight to be synonymous on the earth but weight changes depending on the amount of gravity, but mass remains constant.
Element – simplest form of matter (cannot be broken down any further)‏
Atom – smallest particle of an element with all the properties of that element
Compound – two or more elements chemically bonded together
Molecule – smallest particle of a compound
Example: There are many atoms of copper in a penny and there are many molecules of water in a drop of water.
Properties of Matter
Chemical properties – properties that cannot be observed without altering the substance (ex. flammability, reactivity with other substances)‏
Physical properties – properties that can be observed without altering the substance (ex. color, texture, density, melting or boiling point, shape, etc.)‏
Changes in Matter
Chemical changes – changes that alter the identity of a substance (ex. burning, reacting with another substance, rusting, cooking)‏
Indications of a chemical change: new gas produced, new solid produced, or color change
Physical changes – changes that do not alter the identity of a substance (ex. tearing, cutting, bending, crushing, dissolving, phase changes)‏
States (Phases) of Matter
Solid – definite volume, definite shape, vibrational motion of particles only
(ex. ice, sand, salt, wood)‏
Liquid – definite volume, no definite shape (takes shape of container), flow of particles around each other
(ex. water, alcohol, syrup, gasoline)
Gas – no definite volume (fills up container), no definite shape (takes shape of container), rapid,random motion of particles
(ex. oxygen, steam)
More about States of Matter
Changes in state occur when a certain temperature/pressure is reached (melting point, boiling point, etc.)‏
Discuss water changes...
Discuss butane in a lighter...
Matter Classification
All matter is either a pure substance or a mixture of pure substances.
Pure substances can be elements or compounds.
Elements – simplest form of matter (ex. Cu, Fe, C, Na...all found on periodic table)‏
Compound – two or more element chemically bonded together (ex. NaCl, H O, K SO )‏
There are two types of Mixtures:
Homogeneous Mixture – mixture of two or more pure substances that looks the same throughout (ex. Lotion, air, pudding, salt water)‏
Heterogeneous Mixture – mixture of two or more pure substances that contains visibly different parts (ex. Human, chocolate chip cookie, salad)‏
What weighs more?
One kg of cotton One kg of lead
1 kg = 1 kg
Regardless what the material is!
BUT which is more dense?
1kg of cotton
1kg of lead
We will need the volume also to figure this out...
Density is a property of matter that describes its “compactness”
It is a ratio of mass per unit volume
Density = Mass/Volume
D = ♥
Units are usually g/mL or g/cm
Density of water is 1 g/mL, which means that 1 g of water has a volume of 1 mL.
What is the total volume this flask can hold if it was filled to the top?
HINT: The volume written on the flask is NOT the answer.
We can use the density formula!
Let's say we filled the flask with water and found its mass.
= Mass of water
1547g - 500g = 1047g
Question: Why do we need the mass of the water?
Water is a liquid which takes the volume of its container.
By knowing its mass and density we can solve for the volume of the container!
(Density of water @ 25 C = 0.9970 g/mL)
You have the mass of water = 1047g
You have the density of water = 0.9970 g/mL

Formula for Density is:

Rearranged formula to solve for volume:

Plug in your values:
v = (1047g / 0.9970g/mL)
v = ~1050mL

NOTE: The volume is not equal to the volume of the written on the flask.

Question: Could you solve for mass of the water if you knew the volume instead?

Rearranged formula to solve for mass:
Usually looked up...
Usually found...
Dealing with Numbers
Units are very important in science.
Units tell us two things:
TYPE of measurement
MAGNITUDE of measurement
Example: 32 seconds
Type: time Magnitude: small
Example: 32 tons
Type: weight Magnitude: large
Example: 32 meters
Type: length Magnitude: medium
In chemistry we use the metric system!!!
Metric vs. English System of Units
Metric system is based on powers of ten and is used by most other countries and scientists.
Example: Meters, grams, liters
English system is mainly only used by Americans.
Example: Feet, pounds, fluid ounces
Why 10?
Example: You ate 3 eggs for breakfast out of a dozen. What percent of the eggs did you eat? What percent remains in the carton?
Example: There are 45 people employed at Walgreens. 35.6% of them work part time. How many employees are full time?
Accurate AND Precise
Accurate but not Precise
Precise but not Accurate
Not Accurate nor Precise
Precision vs. Accuracy
Precision – consistency in data
(getting the same or close to the same result over and over)

Accuracy – getting close to the accepted value
Example: Three trials of an experiment showed that the density of aluminum was 4.45 g/mL, 4.48 g/mL, and 4.47 g/mL. The correct value is 2.70 g/mL. These results are precise but not accurate!
Significant Figures
Significant figures are only used for measured values, not for exact numbers.
The more sig figs a measurement has the more accurate the measurement and vice versa.
For example a mass of 80 g (one sig fig) could range anywhere from 75 g to 84 g. However your balance is probably more accurate than that! A mass of 80.0 g (3 sig figs) could range anywhere from 79.95 g to 80.04 g. This is a big difference!
The number of sig figs in a given measurement depends on the accuracy of the instrument used. Sig figs are all the exactly known numbers and the first uncertain or estimated number.
Measure the following using both rulers.
1. All non-zero digits are significant.
23.8 has 3 sig figs
2. Zeros between other non-zero digits are significant
5007 has 4 sig figs
35.00009 has 7 sig figs
3. Zeros in front of non-zero digits are not significant.
0.0892 has 3 sig figs
0.00005 has 1 sig fig
4. Zeros at the end of a number are significant ONLY if a decimal is present.
5.3000 has 5 sig figs
5000.00 has 6 sig figs
639000 has 3 sig figs
Significant Figure Rules
Note: When using scientific notation, ignore the “ x 10?” part and only count the sig figs in the base number.
3.20 x 10 has 3 sig figs
7.000 x 10 has 4 sig figs

The Problem: You measure the length of an object to be exactly 200 cm. This value only has one sig fig, which means your measurement has a great deal of error…anywhere from 150 cm to 249 cm.
The Solutions: (1) Use a decimal 200. cm
(2) Use scientific notation 2.00 x 10 cm
Now the measurement is anywhere from 199.5 to 200.4, which is much more accurate!
Determine the number of significant figures in each of the following values.
0.0845 kg _________
37.00 h _________
8630000.000 mi _________
0.0000000217 g _________
5.750 in _________
0.5003 s _________
500,000 lbs __________
0.0003005400 mL _________
Practice Counting Sig Figs
Rules of rounding with Addition and Subtraction
Round to the fewest digits after the decimal point.
Example: 28.0 + 23.538 + 25.68 = 77.218
Round to the fewest number of significant figures.
Example: 2.24 x 3.4 = 7.616
Rules of rounding with Multiplication and Division
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