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Chapter 10. Chemical Quantities. Before We Begin…. I can write numbers in scientific notation. I can write numbers in standard notation. I can multiply numbers written in scientific notation. I can divide numbers written in scientific notation. Before We Begin….

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Chapter 10

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## Chapter 10

Chemical Quantities

### Before We Begin…

• I can write numbers in scientific notation.

• I can write numbers in standard notation.

• I can multiply numbers written in scientific notation.

• I can divide numbers written in scientific notation.

### Before We Begin…

• We need to review some scientific notation.

• Scientific notation is a way of writing very large and very small numbers.

### How to Write Numbers in Scientific Notation

• Always written as a coefficient multiplied by 10 raised to a power.

3.5 x 1034

coefficient

power

### Examples:

• Write the following in scientific notation:

• 234560000

• 0.00056974

• 8524000000

• 0.000000044258

### How to Multiply in Scientific Notation

• To multiply numbers written in scientific notation you multiply the coefficients and add the powers.

(2.35x1014) x (3.25x10-23)

Multiply

(2.35x3.25) x 1014+-23

### How to Multiply in Scientific Notation

• To multiply numbers written in scientific notation you multiply the coefficients and add the powers.

(2.35x1014) x (3.25x10-23)

Multiply

### Examples:

• Multiply the following numbers:

• (1.23x104) x (4.56x107)

• (7.89x10-1) x (1.23x1010)

• (4.56x107) x (7.89x10-10)

• (1.23x10-11) x (4.56x10-23)

### How to Divide in Scientific Notation

• To divide numbers written in scientific notation you divide the coefficients and subtract the powers.

(2.35x1014) ÷ (3.25x10-23)

Divide

Subtract

### How to Divide in Scientific Notation

• To divide numbers written in scientific notation you divide the coefficients and subtract the powers.

(2.35x1014) ÷ (3.25x10-23)

Divide

Subtract

### Examples:

• Divide the following numbers:

• (1.23x104) ÷ (4.56x107)

• (7.89x10-1) ÷ (1.23x1010)

• (4.56x107) ÷ (7.89x10-10)

• (1.23x10-11) ÷ (4.56x10-23)

### Section 1

The Mole: A Measurement of Matter

### Section 1 Learning Targets

10.1.1 – I can describe methods of measuring the amount of something.

10.1.2 – I can define Avogadro’s number as it relates to a mole of a substance.

10.1.3 – I can distinguish between the atomic mass of an element and its molar mass.

10.1.4 – I can describe how the mass of a mole of a compound is calculated.

### Measuring Matter

• You often measure the amount of something by one of three different methods – by count, by mass, and by volume.

### Example:

• If 0.20 bushel is 1 dozen apples and a dozen apples has a mass of 2.0kg, what is the mass of 0.50 bushel of apples?

### What Is a Mole?

• Mole (mol) – 6.02x1023 representative particles of that substance (SI unit for measuring the amount of something).

• A mole of any substance contains Avogadro’s number of representative particles, or 6.02x1023 representative particles.

### Converting Number of Particles to Moles

• You can use Avogadro’s number as a conversion factor.

### Example:

• How many moles is 2.80x1024 atoms of silicon?

### Converting Moles to Number of Particles

• The reverse also works.

### Example:

• How many molecules are in 5.6 moles of NO2?

### The Mass of a Mole of an Element

• The atomic mass of an element expressed in grams is the mass of a mole of the element.

• Molar mass – the mass of a mole of an element.

• Find the element on the periodic table and the mass that’s listed is the mass of one mole.

### The Mass of a Mole of a Compound

• To calculate the molar mass of a compound, find the number of grams of each element in one mole of the compound.

• Then add the masses of the elements in the compound.

### Example:

• What is the mass of 1.00 mol of sodium hydrogen carbonate?

### Section 2

Mole-Mass and Mole-Volume Relationships

### Section 2 – Learning Targets

10.2.1 – I can describe how to convert the mass of a substance to the number of moles of a substance, and moles to mass.

10.2.2 – I can identify the volume of a quantity of gas at STP.

### The Mole-Mass Relationship

• Use the molar mass of an element or compound to convert between the mass of a substance and the moles of a substance.

### Example:

• Find the mass, in grams, of 4.52x10-3mol of C20H42.

• The reverse is also true.

### Example:

• Calculate the number of moles in 75.0g of dinitrogen trioxide.

### The Mole-Volume Relationship

• Avogadro’s hypothesis – states that equal volumes of gases at the same temperature and pressure contain equal numbers of particles.

• Standard temperature and pressure (STP) – means a temperature of 0°C and a pressure of 101.3kPa or 1 atmosphere (atm).

• At STP, 1 mole or 6.02x1023 representative particles, of any gas occupies a volume of 22.4L

• Molar volume – the 22.4L of a gas.

### Calculating Volume at STP

• 22.4L = 1 mol at STP provides a nice conversion factor.

### Example:

• What is the volume of 3.70 mole N2 at STP?

### Example

• How many moles are in 102 L of carbon dioxide, CO2?

### Calculating Molar Mass from Density

• Different gases have different densities and is usually measured in g/L so we can calculate different things using density as a conversion factor.

### Example:

• A gaseous compound composed of sulfur and oxygen, which is linked to the formation of acid rain, has a density of 3.58 g/L at STP. What is the molar mass of this gas?