Honors Chemistry 2012-13

Unit #1:

Matter and MeasurementsIntroduction: Chemistry is concerned

with matter and energy and how the two interact with each other

Unit topics:

• Matter• Measurements• Properties of substances

Types of Matter

Matter has mass and takes up space:

Phases of matter:

Solids• Fixed volume and shape

Liquids• Fixed volume, indefinite shape

Gases• Indefinite shape and volume

Matter

Pure substances• Fixed composition• Unique set of properties

Mixtures• Two or more substances in some combination

Figure 1.1 - Classification of Matter

A. Elements:

1.Elements cannot be broken down into two or more pure substances• 115 elements; 91 occur naturally

2. Common elements• Carbon (found in charcoal)• Copper (found in pipes, jewelry, etc.)

3. Rare elements• Gold• Uranium

Atomic Symbols

Elements are given symbols:

1. Chemical identifier

2. Elements known to ancient times often have symbols based on Latin names• Copper, Cu (cuprum)• Mercury, Hg (hydrargyrum)• Potassium, K (kalium)

3. One element has a symbol based on a German name• Tungsten, W (wolfram)

Table 1.1 - Elements and Abundances

• Some elements are common, some are rare

Compounds

1. Compounds are combinations of two or more elements chemically bonded to one another.

C. Compounds

2. Compounds always contain the same elements in the same composition by mass.

Water by mass:• 11.19% hydrogen• 88.81% oxygen

3. Properties of compounds are often very different from the properties of elements from which the compounds form.

D. Mixtures

1. Two or more substances in such a combination that each substance retains a separate chemical identity.

A. Copper sulfate and sand • Identity of each is retained

B. Contrast with the formation of a compound• Sodium and chlorine form sodium chloride

D. Mixtures

2.Homogeneous mixtures• Uniform• Composition is the same throughout• Example: seawater

3.Heterogeneous mixtures• Not uniform• Composition varies throughout• Example: rocks

Figure 1.3 – Sodium, Chlorine and Sodium Chloride

Figure 1.2 – Cinnabar and Mercury

Figure 1.4 – Copper Sulfate and Sand

Figure 1.5 – Two Mixtures

4. Separation of mixtures (methods)

Filtration -• Separate a heterogeneous solid-liquid mixture• Barrier holds back solid and lets liquid pass

through• Filter paper will hold back sand but allow water to

pass through

Distillation -• Separates homogeneous mixtures• Salt water can be distilled, allowing water to be

separated from the solid salt

Figure 1.6 – Distillation Apparatus

4. Separation of mixtures (methods)

Chromatography –separation of mixtures based on the size of the particles in the mixture

* commonly used in research and industry

1.2 Measurements

Metric System:

1. Based on the decimal and Powers of ten

2. Four major units:• Length• Volume• Mass• Temperature

Table 1.2 - Powers of Ten

A. Instruments and Units

1. Length• Unit of length is meter; measure of distance• A meter is slightly longer than a yard• Precise definition is the distance light travels in

1/299,272,248 of one second

2. Volume• Unit of volume is liter; measure of the amount of

space matter occupies• Other common units of volume:• Cubic centimeters• Milliliters

• 1 mL = 1 cm3

Table 1.3 – Units and Unit Relations

Measuring volume:

• Graduated cylinder• Pipet or buret• Used when greater accuracy is required

Figure 1.8 – Measuring Volume

3. Mass

• Unit of mass is grams; measure of the amount of matter an object contains

• Other common units of mass:• Kilogram• milligram

Figure 1.9 – Weighing a Solid

4. Temperature

• Factor that determines the direction of heat flow• Temperature is measured indirectly:• Observing its effect on the properties of a

substance• Mercury in glass thermometer• Mercury expands and contracts in response to

temperature

• Digital thermometer• Uses a device called a thermistor

Figure 1.10 – Fahrenheit and Celsius Scales

Temperature Units

• Degrees Celsius• Until 1948, degrees centigrade

• On the Celsius scale• Water freezes at 0 °C• Water boils at 100 °C

The Fahrenheit Scale

• On the Fahrenheit scale vs. Celcius• Water freezes at 32 °• Water boils at 212 °F

The Kelvin Scale:

The lowest possible temperature, in theory 0 K or

-273.16 °C

Unlike the other two scales, no degree sign is used to express temperature in K

Relationships Between Temperature Scales:

• Fahrenheit and Celsius

• Celsius and Kelvin

328.1 CF tt

15.273 CK tT

Example 1.1

Uncertainties in Measurements

1. Estimation• Every measurement carries uncertainty• All measurements must include estimates of

uncertainty with them• There is an uncertainty of at least one unit in the

last digit

Figure 1.11 – Uncertainty in Measuring Volume

Example 2

Rounding Rules:

1. If the first digit to be discarded is 5 or greater, round up

2. If the first digit to be discarded is 4 or smaller, drop off

Addition and Subtraction Rule:

1. Perform the addition(s) and/or subtraction(s)

2. Count the number of decimal places in each number

3. Round off so that the resulting number has the same number of decimal places as the measurement with the least number of decimal places.

Multiplication and Division Rule:

1. When multiplying or dividing numbers, the result is rounded to match the number of significant digits in number with the least number of total significant digits

Example: 2.40 X 2 =

Example: 3.66 / 1.275 =

Example (not in notes)

Exact Numbers:

1.Some numbers carry an infinite number of significant figures

2.These are exact numbers

3. Exact numbers do not change the number of significant figures in a calculation

4. Numbers like π= 3.14 should be considered exact

More on Exact Numbers

5. In some problems in the text, numbers will be spelled out in words

6. “Calculate the heat evolved when one kilogram of coal burns”

7. Consider these numbers to be exact, as well as, conversion factors

C. Conversion of Units

1. Conversion factors are used to convert one set of units to another

Choosing a conversion factor

2. Choose a conversion factor that puts the initial units in the denominator

A. The initial units will cancel

B. The final units will appear in the numerator

Table 1.3 – Length, Volume and Mass Units

Example 4

Properties of Substances

Chemical properties• Require chemical change resulting in new substances

Physical properties• No chemical change

Gold Metal

Examples of Chemical Properties:

• Digestion, burning, rusting

Examples of Physical Properties:

• Color, phase changes, state of matter

• Intensive property – a property that is independent of the sample size

• Extensive property – a property that is dependent on the sample size

A. Density

mass divided by its volume

v

md

Figure – Density of Wood and Water

Example 1.5

B. Solubility

• The process by which one substance dissolves in another is a physical change

• The resulting mixture is a solution/homogeneous mixture

• Solutions may be classified by the relative amount of solute and solvent

1.Saturated: maximum amount of solute

2.Unsaturated: less than maximum amount of solute

3.Supersaturated: more than maximum amount of solute

Figure 1.13 – Sugar Crystals

Figure 1.12 – Solubility and Temperature

Example 1.6

C. Color

• Some substances can be identified by color• Color arises from the absorption and transmission of

specific wavelengths of light• Copper sulfate is blue• Potassium permanganate is deep violet

D. Visible Light

• Visible light ranges from 400 to 700 nm• Below 400 nm is the ultraviolet• Ultraviolet light leads to sunburn

• Above 700 is the infrared• Heat• Absorption of infrared light leads to warming up• Global warming and carbon dioxide

Table 1.4 – Color and Wavelength

Figure 1.14-1.15

Supplementary Information:

Precision and Accuracy in Measurements• A. Precision vs. accuracy (% error calculations)

• Precision – the consistency of a measurement; statistically reported as the range• Accuracy – the closeness to the correct or accepted

value for a measurement; statistically reported as the % error

% Error

• Formula: correct – lab value x 100 = % E

correct

B. Scientific/Exponential Notation

• A method for placing numbers with a large number of zeros in a usable form based on powers of ten

• Example: Express the following numbers in scientific notation:• A) 0.00005607• B) 560700000000

C. Significant Figures

• Rules for significant figures:• A) nonzero digits – are always significant• B) initial zeros – are not significant if they are in

front of the nonzero digits in a decimal number• C. in-between zeros – are always significant• D. final zeros – are significant if they are to the

right of the nonzero digits in a decimal number

Example:

• State the number of significant figures in the following set of measurements:

• A) 30.0 g B) 29.9801 g

• C) 0.03 kg D) 31,000 mg

• E) 3102. cg

• What statements can you make about the accuracy and precision of measurements A-E?

Example:

• Solve the following problems and state the answers with the proper number of significant figures.

• A) Calculate the area of an object with a length of 1.345 m and a width of 0.057 m.

• B) Calculate the density of a substance with a mass of 12.03 g and a volume of 7.0 mL.

Example:

• Express the answers to the following computations in the correct number of significant figures:

• A) 129.0 g + 53.21 g + 1.4365 g

• B) 10.00 m – 0.0448 m

C. Area and Volume Conversions

• To convert a squared or cubed unit, square or cube the conversion factors. If the conversion factor is 1mL=1cm3, do not cube either.

• Ex: Express the area of a 27.0 square yards of carpet in square meters.

• Ex: Convert 17.5 quarts to cubic meters.

(1L = 1.057 qt, 1 ft3 = 28.32 L)

Key Concepts

1. Convert between Fahrenheit, Celsius and Kelvin.

2. Determine the number of significant figures in a measured quantity.

3. Determine the number of significant figures in a calculated quantity.

4. Use conversion factors to change from one quantity to another.

5. Use density to relate mass and volume.

6. Given the solubility, relate mass to volume for a substance.