CHEM Chapter 1 PDF

Preview

Summary

basic notes teacher made us print out...

Description

1

Chapter 1. Chemical Foundations (Review) Chemistry is….. the science that describes matter – its properties, the changes it undergoes and the energy changes that accompany those processes.

1.1 Chemistry: An Overview Why is Chemistry important? [Insert your own belief here]

1.2 The Scientific Method Importance of the Scientific Method The scientific method attempts to minimize the influence of bias or prejudice in the experimenter. Even the best-intentioned scientists can't escape bias. It results from personal beliefs, as well as cultural beliefs, which means any human filters information based on his or her own experience. Unfortunately, this filtering process can cause a scientist to prefer one outcome over another. For someone trying to solve a problem around the house, succumbing to these kinds of biases is not such a big deal. But in the scientific community, where results have to be reviewed and duplicated, bias must be avoided at all costs. That's the job of the scientific method. It provides an objective, standardized approach to conducting experiments and, in doing so, improves their results. By using a standardized approach in their investigations, scientists can feel confident that they will stick to the facts and limit the influence of personal, preconceived notions. Even with such a rigorous methodology in place, some scientists still make mistakes. For example, they can mistake a hypothesis for an explanation of a phenomenon without performing experiments. Or they can fail to accurately account for errors, such as measurement errors. Or they can ignore data that does not support the hypothesis. --- George Harris, “How stuff Works” To put it simply: the scientific method is a series of techniques for investigating phenomena, acquiring new knowledge, or correcting and integrating previous knowledge.

2

Scientific Models Terms to know: Theory: a set of tested hypothesis that gives an overall explanation of some natural phenomena. – A theory is an interpretation – a possible explanation of why nature behaves in a particular way. Natural law: a statement that expresses general behavior – for example, the observation that the total mass of materials is not affected by chemical changes in those materials is called the law of conservation of mass. Difference between theory and law: A law summarizes what happens; a theory is an attempt to explain why it happens

1.3 Units of Measurements In science, there are fundamental SI (Systeme Internationale) units you should know and memorize. Physical Property Length Mass Time Electrical current Temperature Amount of substance

Name of Unit meter kilogram seconds ampere Kelvin mole

Symbol m kg s A K mol

IMPORTANT: The SI unit for Volume is not Liter (L), it’s a mixed or combined unit.

3

You should also know these number prefixes and what they mean. Prefix Peta Giga Mega kilo hecto

milli micro nano pico

Abbreviation Meaning T 1012 109 G 106 M 103 k h 102 Base unit m, g, s, L m 10-3 µ 10-6 n 10-9 p 10-12

Example 1 Tg = 1*1012 g 1 Gg = 1*109 g 1 g = 1*106 g 1 kg = 1*103 g 1 hg = 1*102 g

1 mg = 1*10-3 g 1 µg = 1*10-6 g 1 ng = 1*10-6 g 1 pg = 1*10-6 g

1.4 Uncertainty in Measurements In chemistry, we’re interested in quantification. We want to know how much? For example:

2 H2(g) + O2(g)  2 H2O(l)

If we had 2.0 g of hydrogen: Exactly how much oxygen is needed to produce water and exactly how much water can be produced? Example: A student measured a certain volume of liquid, what is the volume of liquid the student should report?

Remember: In any measurement, some digits are exact and some are inexact.

4

Some conventions to remember! •

Some numbers are exact a) Counted numbers (ex. The number of pencil in a box containing a gross of pencils: 1 gross of pencils is 144 pencils b) Some conversion factors are exact (ex. 1 in. = 2.54 cm or 1 cm3 = 1 mL)

Important: Exact Numbers do not limit the number of sig figs in a calculation. •

Some numbers are inexact Important: Any measured quantity always contains some uncertainty.

There are two important terms related to the uncertainty of measured values: Accuracy: How closely a measurement agrees with the “true” or accepted value Precision: How closely individual measurements agree with each other

Since all measurements possess some degree of uncertainty, we indicate how much uncertainty using significant and insignificant uncertain numbers.

5

Rules for Determining Which Digits are Significant 1) all non-zero digits are significant

5.32 2332

2) a zero between other sig figs is significant

5.02 5.002

3) leading zeros are not significant

0.000532 0.0005002

4) final zeros in a number with a decimal point are significant 5.00 50.000 0.05020 150.

5) final zeros in a number with no decimal point are ambiguous 150 2000 6) Scientific notation is often required: Express 1500 to three sig figs: Express 9000 to two sig figs:

6

Significant Figures in Calculations Addition and Subtraction Ask:

The answer has the same number of decimal places as the least precise measurement used in the calculation. Example: Perform the following calculation and answer using the correct number of significant figures 29.52 + 3.001 + 7219.5 =

• Important: When you add and subtract, it is possible to end up with more significant figures than you started with, or with fewer significant figures than you started with. Example: The mass of a beaker containing a volatile liquid was determined to be 25. 57 g. After 15 minutes, the mass of the beaker and its contents was determined again, and this time it was found to have a mass of 25.49 g. What mass of liquid evaporated?

Important: When you add and subtract, it is possible to end up with more significant figures than you started with, or with fewer significant figures than you started with. Example: You obtained 5.68 g of a compound for an experiment. A few minutes later, you obtained an additional 8.32 g of the same compound. What is the total mass of the substance?

7

Multiplication and division Ask:

2.99 ∗ 7.3 = 5982 ∗ .00201 = 8.672∗10−6 1.376∗104

=

Calculations involving Multiple Steps: Perform the calculation one step at a time, applying the appropriate rule at each step. Important: DO NOT ROUND UNTIL THE END! KEEP TRACK OF SIGNFICANT DIGITS

12.3 + 295.703 8.75 12.3 + 897.21 67.982

Rounding. If left most digit to be removed > 5, round up If left most digit to be removed < 5, stays the same. Examples: Round the following numbers to three significant digits. 3.55711 = 3.55209 = 9.22500001 = 9.235000 =

8

Significant Figures and Scientific Notation:

9.2678 ∗ 10−4 + 8.0 ∗ 10−5 =

9.972𝐸 − 9 + 5.80𝐸 − 11

1.6 Learning to Solve Problems Systematically IMPORTANT: EXACT NUMBERS DO NOT LIMIT THE NUMBER OF SIGNIFICANT FIGURES IN A CALCULATION

SOLVING PROBLEMS BY FOCUSING ON THE UNITS 1. Write down what you want to know. 2. Start with the given measured quantity. 3. Apply conversion factors to cancel unwanted units.

1.7 Dimensional Analysis Example: How many inches are in 62. cm?

Example: The length of the marathon race is approximately 26.2 miles. What is the distance in kilometers?

Example: Express 65.0 miles/hr in m/s. (given: 1 mile = 1.609 km)

Example: Convert 0.00275 dL to µL.

9

1.8 Temperature Conversions Three Temperature Scales

To convert between Kelvin and Celsius

𝑻𝒄(𝑪𝒆𝒍𝒔𝒊𝒖𝒔) + 𝟐𝟕𝟑. = 𝑻𝑲(𝑲𝒆𝒍𝒗𝒊𝒏) Example: The boiling point of water at the top of Mt. Everest is 70.˚C. Covert this temperature to the Kelvin Scale.

Example: On a summer day the temperature in the laboratory, as measured on a lab thermometer is 28.˚C. Express this temperature in C.

10

To convert between Fahrenheit and Celsius

𝑻𝑭(𝑭𝒂𝒓𝒆𝒏𝒉𝒆𝒊𝒕) = 𝟏. 𝟖𝟎𝑻𝑪(𝑪𝒆𝒍𝒔𝒊𝒖𝒔) + 𝟑𝟐 𝑻𝑪(𝑪𝒆𝒍𝒔𝒊𝒖𝒔) =

𝑻𝑭(𝑭𝒂𝒓𝒆𝒏𝒉𝒆𝒊𝒕) − 𝟑𝟐 𝟏. 𝟖𝟎

Example: One of the body’s responses to an infection or injurty is to elevate its’ temperature. A certain flu victim has a body temperature of 101. ˚F. What is this temperature on the Celsius scale?

Example: On a summer day the temperature in the laboratory, as measured on a lab thermometer is 28.˚C. Express this temperature in ˚F.

1.9 Density What is the definition of density of solution?

density of solution (dsoln ) =

mass of solution(g) volume of solution(mL)

Mass density can be defined as the amount of mass of an object, in units of grams, per unit volume, in units of mL. In the first part of CHEM 1411 we will be using the simple definition of density?

𝑑𝑒𝑛𝑠𝑖𝑡𝑦(𝑑) =

𝑚𝑎𝑠𝑠(𝑔) 𝑣𝑜𝑙𝑢𝑚𝑒(𝑚𝐿)

11

Example: Mercury has a density of 13.6 g/mL. What volume of mercury must be taken to obtain 225 g. of the metal.

Example: The density of liquid mercury is 13.6 g/mL. Express this in units of lbs/ft3.

Example: What is the mass of 1.5 m3 of ethanol whose density is 0.789 g/mL , expressed in pounds?

1.9 Matter Term to know: What is matter? Not what’s the matter ☺ According to your book matter, the “stuff” of which the universe is composed, has two characteristics, mass and it occupies space. There are several states of matter, three of which you should know:

12

Physical and Chemical Changes Important: You should be able to distinguish between chemical and physical properties of matter as well the differences between physical and chemical changes. A chemical property of matter is ability of matter to combine with or change into another substance resulting in a change in chemical identity. An example of a chemical property is iron is that it can react and combine with air, in this case oxygen, to form iron oxide (rust).:

Water Air (O 2 ) Iron (Fe)

Iron Oxide (Fe 2 O3 ) Rust

Important: This transformation to go from is also known as a chemical change . A physical property of matter is a characters of a substance that can change without the substance becoming something chemically different. An example of a physical property of water is that water can undergo phase changes and in all phase’s water is still water.

13

Important: These transformations are often known as physical changes. Example: Classify each of the following as a physical or a chemical property. 1. The boiling point of a certain alcohol is 78 C. 2. Diamond is very hard. 3. Sugar ferments to form alcohol. 4. A metal wire conducts an electric current.

14

Intensive and Extensive Properties Properties of matter can be further classified according to whether or not they depend on the amount of substance present. Intensive properties do not depend on the amount of substance present. Examples of intensive properties:

Extensive properties does depend on the amount of substance present. Examples of extensive properties:

Elements and compounds Important terms to know Element: This is a substance of matter that cannot be broken down into simpler substance by chemical or physical means Example: [Look at the periodic table of ELEMENTS]

Compounds: This is a substance which contains constant composition of elements that can be broken down into its fundamental elements through a chemical process. Water can not only undergo a physical change, but it also can undergo a chemical change.

15

The Law of Constant Composition Pure substances, such as water, no matter where you find it, whether it be the ocean, river, always has the same fixed ratio. 11% Hydrogen and 89% Oxygen by mass. The Law of Constant Composition:

16

Mixtures and Pure Substances Important terms to know Mixture: This is a material that contains two or more pure substances. There are two types of mixtures: Heterogeneous mixture

Homogenous mixtures:

Example: Identify each of the following as a pure substance, a homogeneous mixture, or a heterogeneous mixture.

1. gasoline 2. a stream with gravel at the bottom 3. air 4. brass 5. copper metal

17

Separation of Mixtures A mixture can be separated by physical means, not chemical, into pure substances. Each pure substance will retain its own composition and properties.

Sea Water Homogeneous or Heterogenous Mixture?...