Honors Chemistry 2013-14

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

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.

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

Figure 1.4 – Copper Sulfate and Sand

Figure 1.3 – Sodium, Chlorine and Sodium Chloride

Mixtures

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

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

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

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

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

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

• Water freezes at 32 °F• Water boils at 212 °F

The Kelvin Scale:

The lowest possible temperature, in theory 0 K or

-273.15 °C

Unlike the other two scales, no degree sign is used to express temperature in K and there are no negative values

Relationships Between Temperature Scales:

• Fahrenheit and Celsius

• Celsius and Kelvin

328.1 CF tt

15.273 CK tT

Precision and Accuracy in Measurements

Precision and Accuracy in Measurements• 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

Percent Error

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

correct

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

Significant Figures

• Rules for significant figures:• 1. nonzero digits – are always significant• 2. initial zeros – are not significant • 3. in-between zeros – are always significant• 4. final zeros – are significant if the decimal point

is written in the number• 5. SN -If the number is in SN, the entire coefficient

is significant

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 F) 2.40x102 g

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 digit in final

place of a number

Figure 1.11 – Uncertainty in Measuring Volume

Example 2

Rounding Rules:

1. If the first digit to the right of the place you are rounding to is 5 or greater, round up

2. If the first digit to the right of the place you are rounding to 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.

Example: 3.55 + 2.1 =

Multiplication and Division Rule:

1. Preform the multiplication or division.

2. Count the number of total significant digits in each starting number.

3. The result is rounded to match the total number of significant digits in number with the least number of total significant digits

Example: 2.40 X 2 =

Example: 3.66 / 1.275 =

Exact Numbers:

* Some numbers carry an infinite number of significant figures

* These are exact numbers like conversion factors, known values or numbers spelled out in words.

Examples: 12 inches = 1 foot

π= 3.14

fifteen minutes

• 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

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

Example:

• A gasoline station in Manila, Philippines, charges 37.57 pesos per liter of super unleaded gasoline at a time when one US dollar (USD) buys 47.15 Philippine pesos (PHP). The car you are driving has a capacity of 14.00 US gallons, and gets 24 miles per gallon.

• A. What is the cost in USD/gallon?

• B. How much would a tank full cost is USD?

C. If you are empty and have 1255 pesos, how many km can you drive?

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 = 1.06 qt, 1 cm3 = 1 mL)

Properties of Substances

Chemical properties• Require chemical change resulting in new substances• Digestion, burning, rusting

Physical properties• No chemical change (appearance chages)• 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 volume for an object

v

md

Figure – Density of Wood and Water

Example:

Two students were given identical cylindrical bars with the following data:

Mass = 96.03g, length = 10.7 cm, diameter = 9.82 mm

A. Student X was asked to determine whether her bar was made of pure palladium. (Dpalladium = 12.02 g/mL)

B. Student Y was asked to determine how many grams of ethyl alcohol (d = 0.789 g/mL) his bar would displace

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.

• C) Mercury thermometers are being phased out because of the toxicity of mercury vapor. A common replacement for mercury is the organic liquid isoamyl benzoate, which boils at 262°C. What is the boiling point in °F? and K?