International System Of Units (Metric System)

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# International System Of Units (Metric System) - PowerPoint PPT Presentation

International System Of Units (Metric System). Types of Measurements. 1- QUALITATIVE MEASUREMENTS: observations of reactions — changes in color and physical state. 2- QUANTITATIVE MEASUREMENTS : which involve numbers . Use SI units — based on the metric system.

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Presentation Transcript
Types of Measurements

1- QUALITATIVE MEASUREMENTS:observations of reactions — changes in color and physical state.

2- QUANTITATIVE MEASUREMENTS:which involve numbers.

• UseSI units— based on the metric system
What is Scientific Notation?
• Scientific notation is a way of expressing really big numbers or really small numbers.
• For very large and very small numbers, scientific notation is more concise.

Writing Numbers in Scientific Notation

• Scientific notation is a method of expressing a quantity as a number multiplied by 10 to the appropriate power.
• For example, the measurement 300,000,000 m/s can be written as 3.0  108 m/s in scientific notation.
• The same is true of small measurements. For example, the quantity 0.0015 kg can be written as 1.5  10-3 in scientific notation.
Move decimal point

# of spaces the decimal moves is the power of 10

If exponent is positive, move decimal to the right

If exponent is negative, move decimal to the left

4.285 x 102  428.5(move decimal 2 spots right)

4.285 x 10-4  0.0004285(decimal moves 4 spots left)

Converting From Scientific to Standard Notation

Learning Check
• Express these numbers in Scientific Notation:
• 405789
• 0.003872
• 3000000000
• 2
• 0.478260

The International System

• To avoid confusion, scientists established the International System of Units, or SI, in 1960 as the accepted system for measurement.
• There is Seven SI base units
Metric Prefixes
• Kilo- means 1000 of that unit
• 1 kilometer (km) = 1000 meters (m)
• Centi- means 1/100 of that unit
• 1 meter (m) = 100 centimeters (cm)
• 1 dollar = 100 cents
• Milli- means 1/1000 of that unit
• 1 Liter (L) = 1000 milliliters (mL)
To convert a larger units to smaller units : multiply
• Ex: 8Kg = 8 * 1000 = 8000g
• To convert a smaller units to larger units : divide
• Ex: 7g = 7/1000 = 0.007 Kg

Length

• Length is defined as the distance between two points.
• The meter (m) is the SI unit of length. Smaller objects can be measured in centimeters (cm) or millimeters (mm). The length of your textbook or pencil would be measured in centimeters.
• To measure long distances, you use kilometers.
• Kilometers might be most familiar to you as the distance traveled in a car or the measure of a long-distance race.

O—H distance =

9.4 x 10-11 m

9.4 x 10-9 cm

0.094 nm

Units of Length
• ? kilometer (km) = 500 meters (m)
• 2.5 meter (m) = ? centimeters (cm)
• 1 centimeter (cm) = ? millimeter (mm)
• 1 nanometer (nm) = 1.0 x 10-9 meter

Volume

• Volume is the amount of space that something occupies. The volume of liquids are usually given in liters (L) or milliliters (mL). The volume of solids can be given in cubic meters (m3), cubic centimeters (cm3), or cubic millimeters (mm3).
• Units of Volume

The SI unit of volume is the amount of space occupied by a cube that is 1 m along each edge. This volume is the cubic meter (m)3. A more convenient unit of volume for everyday use is the liter, a non-SI unit.

A liter (L) is the volume of a cube that is 10 centimeters (10 cm) along each edge (10 cm  10 cm  10 cm = 1000 cm3 = 1 L).

Common metric units of volume include the liter, milliliter, cubic centimeter, and microliter.

The volume of 20 drops of liquid from a medicine dropper is approximately 1 mL.

Mass

• The mass of an object measures the amount of matter in the object.
• The kilogram (kg) is the SI unit for mass.
• You can determine mass with a triple-beam balance.
• The balance compares an object to a known mass. Weight and mass are not the same. Mass depends only on the amount of matter in an object.

Weight

• Weight is a force that measures the pull on a given mass by gravity
• The SI unit for weight is the Newton (N).
• Weight depends on gravity, which can change depending on where the object is located.
• If you were to travel to other planets, your weight would change, even though you would still be the same size and have the same mass.
• This is because gravitational force is different on each planet.

Platinum

Mercury

Aluminum

DENSITY - an important and useful physical property

13.6 g/cm3

21.5 g/cm3

2.7 g/cm3

Densityis the amount of matter in a given volume. Density can be expressed in grams per milliliter (g/mL) or grams per cubic centimeter (g/cm3).

PROBLEM: Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg?

• Specific Gravity:

Sp.Gr. =Density of substance (g/ml) / Density of water (g/ml)

Learning Check

If blood has a density of 1.05 g/mL, how many liters of blood are donated if 575 g of blood are given?

1) 0.548 L

2) 1.25 L

3) 1.83 L

Temperature

• The physical property of temperature is related to how hot or cold an object is.
• Thermometers are used to measure temperature.
• Temperature is measured in SI with the Kelvin (K) scale.
• There is three common scales used to determines temperature
• 1- Fahrenheit
• 2- Kelvin
• 3- Celcius
On the Celsius scale, the freezing point of water is 0°C and the boiling point is 100°C.
• On the Kelvin scale, the freezing point of water is 273.15 kelvins (K), and the boiling point is 373.15 K.
• The zero point on the Kelvin scale, 0 K, or absolute zero, is equal to 273.15 °C.
• The Kelvin scale starts at 0 K. In theory, 0 K is the coldest temperature possible in nature.
• Because one degree on the Celsius scale is equivalent to one kelvin on the Kelvin scale, converting from one temperature to another is easy. You simply add or subtract 273, as shown in the following equations.
Fahrenheit FormulaḞ = 9/5 ċ + 32
• Celsius Formulaċ = 5/9 * ( Ḟ - 32)
Learning Check

The normal temperature of a chickadee is 105.8°F. What is that temperature in °C?

1) 73.8 °C

2) 58.8 °C

3) 41.0 °C

Precision and Accuracy

• Precisionis a description of how close measurements are to each other.
• Suppose you measure the distance between your home and your school five times and determine the distance to be 2.7 km.
• Suppose a friend measured 2.7 km on two days, 2.8 km on two days, and 2.6 km on the fifth day.
• Because your measurements were closer to each other than your friend’s measurements, yours were more precise.

Accuracy

• Accuracy - a measure of how close a measurement is to the true value of the quantity being measured.
Example: Evaluate whether the following are precise, accurate or both.

Accurate

Not Precise

Not Accurate

Precise

Accurate

Precise

What is a mole?
• The mole, whose abbreviation is “mol”, is the SI base unit for measuringamount of a pure substance.
• A counting unit
• Similar to a dozen, except instead of 12, it’s 602 billion trillion 602,000,000,000,000,000,000,000
• 6.02 X 1023 (in scientific notation)
• 1 dozen Al atoms = 12 Al atoms
• 1 mole of Al atoms = 6.02 X 1023 atoms
• A mole is Avogadro’s number of particles, that is 6.02 × 1023 particles.

1 mol = Avogadro’s Number = 6.02 × 1023 units

A Mole of Particles Contains 6.02 x 1023 particles

1 mole C

1 mole H2O

1 mole NaCl

= 6.02 x 1023 C atom

= 6.02 x 1023H2O molecules

= 6.02 x 1023NaCl “molecules”

(technically, ionics are compounds not molecules so they are called formula units)

6.02 x 1023 Na+ ions and

6.02 x 1023Cl– ions

6.02 x 1023 particles

1 mole

or

1 mole

6.02 x 1023 particles

Note that a particle could be an atom OR a molecule!

6.02 × 1023 atoms Na

= 7.22 × 1022 atoms Na

0.120 mol Na ×

1 mol Na

Mole Calculations I
• How many sodium atoms are in 0.120 mol Na?
• Step 1: we want atoms of Na
• Step 2: we have 0.120 mol Na
• Step 3: 1 mole Na = 6.02 × 1023 atoms Na

1 mol K

1.25 × 1021 atoms K ×

= 2.08 × 10-3 mol K

6.02 × 1023 atoms K

Mole Calculations I
• How many moles of potassium are in 1.25 × 1021 atoms K?
• Step 1: we want moles K
• Step 2: we have 1.25 × 1021 atoms K
• Step 3: 1 mole K = 6.02 × 1023 atoms K
Molar Mass
• The atomic mass of any substance expressed in grams is the molar mass (MM) of that substance.
• Equal to the numerical value of the average atomic mass (get from periodic table)

1 mole of C atoms = 12.0 g

1 mole of Mg atoms = 24.3 g

1 mole of Cu atoms = 63.5 g

• The atomic mass of iron is 55.85 amu.
• Therefore, the molar mass of iron is 55.85 g/mol.
Molar Mass of Compounds
• The molar mass (MM) of a compound is determined the same way, except now you add up all the atomic masses for the molecule (or compound)
• Ex. Molar mass of CaCl2
• Avg. Atomic mass of Calcium = 40.08g
• Avg. Atomic mass of Chlorine = 35.45g
• Molar Mass of calcium chloride = 40.08 g/mol Ca + (2 X 35.45) g/mol Cl 110.98 g/mol CaCl2

20

Ca40.08

17Cl 35.45

Mole Calculations II
• Now we will use the molar mass of a compound to convert between grams of a substance and moles or particles of a substance.

6.02 × 1023 particles = 1 mol = molar mass

• If we want to convert particles to mass, we must first convert particles to moles and than we can convert moles to mass.

g

mol

g/mol

Formula

g/mol

g

mol (n)

Equation

HCl

0.25

g= g/mol x mol

H2SO4

53.15

NaCl

3.55

Cu

1.27

Converting between grams and moles
• If we are given the # of grams of a compound we can determine the # of moles, & vise-versa
• In order to convert from one to the other you must first calculate molar mass

g = mol x g/mol

mol = g  g/mol

• Thiscanberepresentedinan“equationtriangle”

Atoms or Molecules

Flowchart

Divide by 6.02 X 1023

Multiply by 6.02 X 1023

Moles

Multiply by atomic/molar mass from periodic table

Divide by atomic/molar mass from periodic table

Mass (grams)

47.88 g Ti

1.33 mole Ti ×

= 63.7 g Ti

1 mole Ti

Mass-Mole Calculations
• What is the mass of 1.33 moles of titanium, Ti?
• We want grams, we have 1.33 moles of titanium.
• Use the molar mass of Ti: 1 mol Ti = 47.88 g Ti

1 mol Pb

207.2 g Pb

2.55 × 1023 atoms Pb ×

×

1 mole Pb

6.02×1023 atoms Pb

Mole Calculations II
• What is the mass of 2.55 × 1023 atoms of lead?
• We want grams, we have atoms of lead.
• Use Avogadro’s number and the molar mass of Pb

= 87.8 g Pb

1 mol O2

6.02×1023 molecules O2

0.470 g O2 ×

×

1 mole O2

32.00 g O2

Mole Calculations II
• How many O2 molecules are present in 0.470 g of oxygen gas?
• We want molecules O2, we have grams O2.
• Use Avogadro’s number and the molar mass of O2

8.84 × 1021 molecules O2