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Topic 1: An Introduction to Chemistry Matter & Change (Chapter 1 in Modern Chemistry)

Chemistry is a physical science. It is the study of the composition, structure, and properties of matter, the processes that matter undergoes, and the energy changes that accompany these processes.

Branches of Chemistry

(6 main areas of study)

x Organic chemistry-the study of most carbon-containing compounds (except carbon & carbonate)

x Inorganic chemistry-the study of non-organic substances x Physical chemistry-the study of the properties and changes of matter and their relation to

energy x Analytical chemistry-the identification of the components and composition of materials x Biochemistry-the study of substances and processes occurring in living things x Theoretical chemistry-the use of mathematics and computers to understand the principles

behind observed chemical behavior and to design and predict the properties of new compounds

A chemical is any substance that has a definite composition.

Mass is a measure of the amount of matter.

Matter is anything that has mass and takes up space. Matter is stuff.

Matter has characteristic properties. These properties can be used to distinguish among substances and to separate them. Properties can be classified as physical or chemical.

Physical properties are characteristic that can be observed or measured without changing the identity of the substance. It may look different but it has the same makeup. For example: boiling, melting, and freezing points are physical properties. Water has a formula of H2O when it boils and becomes steam which also has a formula of H2O. It has this same formula when liquid water freezes and changes to ice. Physical changes are changes in a substance that does not involve a change in the identity of the substance. All phase changes (changes of state) are physical changes.

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Matter in the solid state has definite volume and definite shape or according to the Kinetic Molecular Theory (KMT), solid particles vibrate around fixed points. Matter in the liquid state has a definite volume but does not have a definite shape. It assumes the shape of its container. According to the (KMT), liquid particles vibrate around moving points. Matter in the gas state has neither definite volume nor definite shape or according to the (KMT), gas particles move in random straight line motion until they hit something. A fourth state of matter is plasma. Plasma is a high-temperature physical state of matter in which atoms lose most of their electrons.

http://www.dlt.ncssm.edu/tiger/Flash/phase/KineticEnergy-Solid.html http://www.dlt.ncssm.edu/tiger/Flash/phase/KineticEnergy-Liquid.html http://www.dlt.ncssm.edu/tiger/Flash/phase/KineticEnergy-Gas.html Physical properties can be intensive or extensive. Extensive properties depend on the amount of matter that is present. For example: volume, mass, and the amount of energy in a substance Intensive properties do not depend on the amount of matter present. For example: melting point, boiling point, density, and ability to conduct electricity and to transfer energy as heat

Chemical properties relate  to  a  substance’s  ability  to  undergo  changes that transform it into different substances. The substances react to form new substances. A chemical change or chemical reaction is a change in which one or more substances are converted into different substances. The substances that react in a chemical change are called the reactants. The substances that are formed by the chemical change are called the products.

carbon plus oxygen yield carbon dioxide.

carbon + oxygen Æ carbon dioxide

C + O2 Æ CO2

Task 1a

1. Label the following as chemical or physical change/property. a. Melting b. Digestion c. Rusting of iron d. Tearing paper

2. Label the following as an intensive or extensive physical property. a. Length b. Mass c. Density d. Color

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Classification of Matter

Matter is broken down into mixtures and pure substances.

Mixtures are blends of two or more kinds of matter, each of which retains its own identity and properties. They are simply physically mixed together and have variable composition. In other words, they do not have a formula. They can usually be separated relatively easily. There are heterogeneous mixtures and homogeneous mixtures. A heterogeneous mixture is not mixed uniformly, for example: granite, clay & water. Homogeneous mixtures are uniform in composition. They still do not have a formula, but the particles are physically mixed evenly all the way through. Homogeneous mixtures are also called solutions. Examples of solutions are salt water, air, and kool-aid. Some mixtures can be separated by filtration, evaporation, distillation, or chromatography.

Pure Substances are all homogeneous but are not mixtures because they are chemically combined instead of physically combined. A pure substance has a fixed composition (has a formula). Every sample of a given pure substance has exactly the same characteristic properties and has exactly the same composition. Pure substances can be broken down into elements and compounds.

Elements are pure substances that cannot be broken down into simpler, stable substances and is made of one type of atom. Elements are found on the periodic table. They are made of only one type of atom. An atom is the smallest unit of an element that maintains the chemical identity of that element.

A compound is a substance that can be broken down into simple stable substances. Each compound is made from the atoms of two or more elements that are chemically bonded. They have definite formulas.

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Task 1b

1. Label the following as heterogeneous mixture, solution, compound or element. a. Salt, NaCl b. Tea c. Iced tea d. Iron e. Concrete f. Copper g. Sugar

The Periodic Table

The periodic table is one of the most important references for chemists. It is imperative that you learn the symbols and names of common elements. Be sure to study the list supplied in the memory work section of edline.

Periodic table key: blue main group metals pink transition metals teal metalloids purple nonmetals green noble gases

Symbol color key (at rt): Black = solid White = gas Yellow = liquid

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The vertical columns on the periodic table are called groups or families. We refer to the groups by number and the families by name. For example, column 2 is called group 2 or the alkaline earth metal family. Elements in the same group have similar properties. There are a few common family names that you need to know. Other families are known by the first element in that group.

Group 1 is the alkali metals family Group 2 is the alkaline earth metal family Group 17 is the halogen family Group 18 is the noble gas family The horizontal rows of elements in the periodic table are called periods. Physical and chemical properties change somewhat regularly across a period. Periods are designated by the number of row (1-7). Elements can also be classified as metals, nonmetals, and metalloids. Metals are to the left of a stair-step line that begins on the left of Boron (except for Hydrogen which is a nonmetal). Nonmetals are to the right of the stair-step line. Elements that touch the stair-step line on a top, bottom or side (not corner) except for aluminum are metalloids. Metals are good conductors of heat and electricity. They are malleable and ductile. They tend to lose valence electrons when they bond. Nonmetals are poor conductors of heat and electricity. They are gaseous or brittle. They tend to gain electrons when they bond. Metalloids sometimes act like metals and sometimes act like nonmetals. (See color codes on periodic table) Task 1c

1. Write the symbol for the following elements. a. Copper b. Sodium c. Argon d. Oxygen e. Zinc

2. Write the name of the following element symbols. a. K b. Li c. S d. He e. Fe

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3. List the group and period of the following elements. a. Xe b. Cr c. I d. U

4. Label the following as a metal, nonmetal, or metalloid. a. Aluminum b. Nickel c. Germanium d. Carbon

Measurements & Calculations (Chapter 2 in Modern Chemistry)

The Scientific Method is a logical approach to solving problems by observing and collecting data, formulating hypotheses, testing hypotheses, and formulating theories that are supported by data. Observing uses the senses. Observations can be qualitative or quantitative. Qualitative observations are non-numerical, for example: the room is cold. Quantitative observations are numerical, for example: the room is 58oF.

Chemists study systems. A system is a specific portion of matter in a given region of space that has been selected for study during an experiment or observation. An example of a system is a test tube and its contents.

An hypothesis is a testable statement (educated guess). Many times hypotheses are stated as an “if-then”  statement.

Hypotheses are tested using an experiment. For an experiment to be valid, the experiment will have controls or conditions that are constant throughout the testing. The experiment will also have variables, or conditions that change. Experiments will usually have a dependent variable and an independent variable. The independent variable (usually on the x axis of a graph) is the variable that is typically being manipulated by the experimenter while the dependent variable (usually on the y axis of a graph) is the observed result of the independent variable being manipulated. An example would be the amount of fertilizer vs. plant growth. The amount of fertilizer would be the independent variable and the growth would be the dependent variable.

After the hypothesis has been tested and the data has been recorded and analyzed, then conclusions can be drawn from the experiment. Theorizing about the experiment could start by constructing a model. A model in science is more than a physical object; it is often an explanation of how phenomena occur and how data or events are related (For example: the

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atomic model). If a model successfully explains many phenomena, it becomes part of a theory. A theory is a broad generalization that explains a body of facts or phenomena. A scientific theory summarizes a hypothesis or group of hypotheses that have been supported with repeated testing. A theory is valid as long as there is no evidence to dispute it. Therefore, theories can be disproven. This is different from a law. A law generalizes a body of observations. At the time it is made, no exceptions have been found to a law. Scientific laws explain things, but they do not describe them. One way to tell a law and a theory apart is to ask if the description gives you a means to explain 'why'.

Task 1d

1. Label the following as qualitative or quantitative observations. a. The liquid floats on water. b. The metal is malleable. c. The liquid has a temperature of 55.6oC.

SI Measurements & Units

Quantity is something that has magnitude, size, or amount. A measurement is a quantity with a unit of measurement. For example: 1.5 g. Scientists us the SI system of measurement (Le Systeme  International  d’Unites). We will also refer to this as the metric system. Here is a table of the base (fundamental) units used in the SI system. We will only be using the first five.

Important fundamental measurements

Mass- a measure of the quantity of matter, measured with a balance, units g, kg, etc.

Length- the distance between two points, measure with a ruler, units m, km, cm, mm, etc.

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Derived SI Units

Units not listed on the base unit table are derived units. Derived units are combinations of SI base units.

Important derived measurements

Weight- a measure of the gravitational pull on matter, measured with a scale, unit N (newtons).

Volume- the amount of space occupied by an object, measures with a ruler or graduated cylinder, units m3, cm3, mL, L, etc. For regular shape objects, you can use a mathematical formula for volume such as V = l x w x h IMPORTANT: 1 cm3 = 1 mL and 1 dm3 = 1 L.

Density- the ratio of mass to volume, or mass divided by volume. D = m/V, the units for density are g/cm3, g/mL, kg/L, kg/dm3, etc.

D = m V

Task 1e 1. Label each of the following measurements by the quantity each represents. For instance,

a measurement of 10.6 kg/m3 represents density. a. 5.0 g/mL b. 37 s c. 39.56 g d. 47 J e. 25.3 cm3 f. 500 m2 g. 30.23 mL

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Task 1f 1. What is the density of a block of marble that occupies 310. mL and has a mass of 853 g? 2. Diamond has a density of 3.26 g/cm3. What is the mass of a diamond that has a volujme

of 0.351 cm3? 3. What is the volume of a sample of liquid mercury that has a mass of 76.2 g, given that the

density of mercury is 13.6 g/mL? 4. A block of sodium that has the measurements 3.00 cm x 5.00 cm x 5.00 cm has a mass of

75.5 g. Calculate the density of sodium.

Dimensional Analysis (Factor Label Method)

Dimensional Analysis is a mathematical technique that allows you to use units to solve problems involving measurements. THIS IS AN IMPORTANT TECHNIQUE!

Use conversion factors to go from the quantity given to the quantity sought. Factors are numbers, labels are units.

When using the factor-label method, problems consist of three parts: 1. a known beginning – GIVEN 2. a desired end – WANTED 3. a connecting path – CONVERSION FACTORS

A conversion factor is a ratio derived from the equality between two different units that can be used to convert from one unit to the other. For example:

4 quarters = 1 dollar

This can be written as two conversion factors:

4 quarters or 1 dollar 1 dollar 4 quarters

Notice that each conversion factor equals 1. They equal each other. Notice conversion factors can be flipped.

Here is an example using conversion factors & dimensional analysis: How many seconds are in 2.5 days?

2.5 days 24 hours 60 minutes 60 seconds 1 day 1 hour 1 minute

Multiply the top, multiply the bottom, and divide the answers. All units cancel out except seconds.

This equals 216 000 seconds.

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Metric Conversions

Learn this chart! (Here’s  an  easy  way:    My  kangaroo  has  dance  until  dawn  ‘cause  music  makes  noise).

106 103 102 101 100 10-1 10-2 10-3 10-6 10-9

M k h da u d c m P n

Mega- kilo- hecta- deka- deci- centi- milli- micro- nano-

You can use dimensional analysis with this information or move the decimal the appropriate number of places. For example:

Convert 35 m to km. 35 m 1 km = 0.035 km 1000 m Or  “m”  has  no  prefix  so  it  is  in  the  units  column,  “km”    has  kilo  as  its  prefix,  so  it  is  in  the  “k”  column. Move the decimal 3 places to the left, just like on the chart. If you use this method, watch out for prefixes that change by more than one change of 10 (mega, micro, nano).

35 m becomes 0.035 km

Task 1g

1. Complete the following conversions a. 10.5 g = _____ kg b. 1.57 km = _____ m c. 3.54 Pg = _____ g d. 3.5 mol = _____�Pmol e. 1.2 L = _____ mL f. 37.8 mL = _____ cm3

2. Use dimensional analysis to determine the following.

a. If 1 mag = 13 bops and 1 bop = 4.6 skuts, how many mags are in 583 skuts? b. How many centimeters are in 2.5 yards? (1 inch = 2.54 cm)

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Temperature Conversions

Use the following formulas for temperature conversions.

K = oC + 273 oC = K – 273

oC = 5/9(oF-32) oF = 9/5 oC +32

Task 1h

1. Make the following temperature conversions. a. 58.7oC = _____ K b. 323 K = _____ oC c. 75oF = _____ oC d. 368 K = _____ oF

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Using Scientific Measurements

Accuracy refers to the closeness of measurements to the correct or accepted value of the quantity measured. Precision refers to the closeness of a set of measurements of the same quantity made in the same way.

Percentage Error is calculated by subtracting the accepted value from the experimental value, dividing the difference by the accepted value, and then multiplying by 100.

Percentage error = experimental - accepted x 100 accepted

Task 1i

1. What is the percentage error for a mass measurement of 17.7 g, given that the correct value is 21.1 g?

2. A volume is measured experimentally as 4.26 mL. What is the percent error, given that the correct value is 4.15 mL?

3. During an experiment, a student obtains the following density data: 9.12 g/mL, 9.11 g/mL, and 9.13 g/mL. The literature value for the density of this substance is 8.78 g/mL. Is this student accurate? Is this student precise? What is the % error of this students information?

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Significant figures (Significant Digits) in a measurement consist of all the digits known with certainty plus one final digit, which is somewhat uncertain or is estimated.

Rules for Significant Figures Read from the left and start counting sig figs when you encounter the first non-zero digit

1. All non zero numbers are significant (meaning they count as sig figs) 613 has three sig figs 123456 has six sig figs

2. Zeros located between non-zero digits are significant (they count) 5004 has four sig figs 602 has three sig figs 6000000000000002 has 16 sig figs!

3. Trailing zeros (those at the end) are significant only if the number contains a decimal point; otherwise they are insignificant (they don’t  count) 5.640 has four sig figs 120000. has six sig figs 120000 has two sig figs – unless  you’re given additional information in the

problem

4. Zeros to left of the first nonzero digit are insignificant (they don’t  count); they are only placeholders! 0.000456 has three sig figs 0.052 has two sig figs 0.000000000000000000000000000000000052 also has two sig figs!

Rules for addition/subtraction problems Your calculated value cannot be more precise than the least precise quantity used in the calculation. The least precise quantity has the fewest digits to the right of the decimal point. Your calculated value will have the same number of digits to the right of the decimal point as that of the least precise quantity. In practice, find the quantity with the fewest digits to the right of the decimal point. In the example below, this would be 11.1 (this is the least precise quantity).

7.939 + 6.26 + 11.1 = 25.299 (this is what your calculator spits out)

In this case, your final answer is limited to one sig fig to the right of the decimal or 25.3 (rounded up).

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Rules for multiplication/division problems The number of sig figs in the final calculated value will be the same as that of the quantity with the fewest number of sig figs used in the calculation. In practice, find the quantity with the fewest number of sig figs. In the example below, the quantity with the fewest number of sig figs is 27.2 (three sig figs). Your final answer is therefore limited to three sig figs.

(27.2 x 15.63) 1.846 = 230.3011918 (this is what you calculator spits out)

In this case, since your final answer it limited to three sig figs, the answer is 230. (rounded down) Rules for combined addition/subtraction and multiplication/division problems First apply the rules for addition/subtraction (determine the number of sig figs for that step), then apply the rules for multiplication/division. Task 1j

1. Provide the number of sig figs in each of the following numbers: (a) 0.0000055 g _____ (b) 3.40 x 103 mL ______ (c) 1.6402 g _____ (d) 1.020 L _____ (e) 16402 g ______ (f) 1020 L _______

2. Perform the operation and report the answer with the correct number of sig figs.

(a) (10.3 m) x (0.01345 m) = ___________________ (b) (10.3) + (0.01345) = ______________________ (c) [(10.3) + (0.01345)] = ____________________________

[(10.3) x (0.01345)]

3. Polycarbonate plastic has a density of 1.2 g/cm3. A photo frame is constructed from two 3.0 mm sheets of polycarbonate. Each sheet measures 28 cm by 22 cm. What is the mass of the photo frame?

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Scientific Notation (You should already know this)

Scientific Notation was developed in order to easily represent numbers that are either very large or very small. Scientific Notation is based on powers of the base number 10.

Examples: The number 200,000,000,000 stars in scientific notation is written as 2 x 1011 stars

The number 0. 000,006,645 kilograms in scientific notation is written as 6.645 x 10-6 stars

The first number 6.645 is called the coefficient. The coefficient must be greater than or equal to 1 and less than 10.

The coefficient contains only significant digits. The second number is called the base. The base must always be 10 in scientific notation. The number -6 is referred to as the exponent or power of ten. The exponent must show the number of places that the decimal needs to be moved to change the number to standard notation. A negative exponent means that the number written in standard notation is less than one. To Change from Standard Form to Scientific Notation:

1. Place decimal point such that there is one non-zero digit to the left of the decimal point. 2. Count number of decimal places the decimal has "moved" from the original number.

This will be the exponent of the 10. 3. If the original number was less than 1, the exponent is negative; if the original number

was greater than 1, the exponent is positive.

To Change from Scientific Notation to Standard Form:

1. Determine the number of places the decimal must be moved from the exponent. 2. Decide if the standard form will be a number greater than one or less than one. 3. Move the decimal in the coefficient adding place holders if necessary.

Task 1k

1. Write the following numbers in scientific notation. a. 96 400 b. 0.361 c. 0.0057300 d. 6 587 234 000

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2. Write the following numbers in standard notation. a. 3.97 x 103 b. 8.862 x 10-1 c. 6.251 x 109 d. 512 x 10-8

Data Manipulation

All measurements taken in lab must be in the correct significant digits. Data tables should be prepared ahead of time if possible. You will be expected to understand what is occurring, not only following directions. This is not just cookbook chemistry.

You will have to graph data that you collect in lab. Remember that the independent variable goes on the x-axis. The dependent variable goes on the y-axis.

Two quantities are directly proportional to each other if dividing one by the other gives a constant value. y/x = k or y = kx (in the form of y = mx + b) This is a straight line graph. An example of this is density.

Two quantities are inversely proportional to each other if their product is constant. xy = k This graph produces a curve called a hyperbola. Pressure & volume of gases give this type of inverse proportion.

States of Matter (Chapter 10 Modern Chemistry)

Watch this!! This will review the six phase changes that matter undergoes.

http://www.kentchemistry.com/links/Matter/PhaseChanges.htm

Phase Diagrams

A phase diagram is a graph of pressure versus temperature that shows the conditions under which the phases of a substance exist.

The triple point of a substance indicates the temperature and pressure conditions at which the solid, liquid, and vapor of the substance can coexist at equilibrium.

The critical point of a substance indicates the critical temperature and critical pressure. This is the point above which the substance cannot exist in the liquid state.

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The normal melting point or boiling point is the melting or boiling point at standard pressure. In order to find the normal points you must read the graph across from 1 atm, 760 mmHg, or 101.3 kPa.

General Phase Diagram

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Task 1L

1. Using phase Diagram for water (a), what is the critical point of water? 2. What  is  water’s  triple  point? 3. Using the phase diagram for carbon dioxide (b), what is the normal melting point? The

normal boiling point?

Heating & Cooling Curves

Heating & Cooling curves can help you tell at what temperature a substance melts/freezes or boils/condenses at the current pressure. Note that as a phase is change the temperature doesn’t  change.    All  the  energy  is  going  toward  breaking  the  particles intermolecular bonds so the change can occur. After the phase change occurs the energy can now be used to increase the temperature (or vice versa).

http://www.dlt.ncssm.edu/tiger/Flash/phase/HeatingCurve.html

This link allows you to perform three melting/boiling experiments while simultaneously graphing the data. You should be able to determine the melting point and boiling point of each substance.

http://www.kentchemistry.com/links/Matter/PhaseChanges.htm

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Solutions (Chapter 12 Modern Chemistry)

Solubility & Solubility Curves

Solubility of a substance is the amount of that substance required to form a saturated solution with a specific amount of solvent at a specified temperature.

A saturated solution contains the maximum amount of dissolved solute. An unsaturated solution contains less solute than a saturated solution under the existing conditions. A supersaturated solution contains more dissolved solute than a saturated solution contains under the same conditions.

As you can see in the diagram below, the solubility of solids generally increase with temperature, while the solubility of gases decrease with a temperature increase.

Task 1m

4. Using the solubility curves, what is the solubility of KClO3 in 100 g of H2O at 50 oC? 5. How many grams of KCl is needed to make a saturated solution at 50 oC in 100 g of

H2O? What about in 200 g of H2O at this same temperature? 6. How many grams of NaCl will dissolve in 100 g of H2O at 60 oC? How much salt would

sink to the bottom of the beaker if 30 g of NaCl is added to 100 g of H2O? What is 60 g of NaCl is added to 100 g of H2O?

7. What type of solution would 40 g of NH3 in 100 g H2O at 20 oC?