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Chemistry 11 – Mole Concept Study Guide 1

Avogadro’s Number:

• There is a convenient unit when working with small particles such as atoms and molecules.

• This number is known as _________________, which is called __________________________, to honour the Italian scientist whose hypothesis led to its determination.

• If written out, this number looks like this:

602 000 000 000 000 000 000 000

Why do we need to know such a big number like this in Chemistry?

• Elements have different atomic masses (eg. Mg = 24.3g/mol, Ca = 40.1g/mol)

• Atoms and molecules are _____________ – we will work with _______ quantities of them.

• Need a convenient _____ to count atoms and molecules

• Example: A dozen ! instead of counting 1, 2, 3, we count by 12’s ! a known quantity that is understood by all. ! A convenient unit

• Chemists use a special term that is similar to “a dozen”.

• This term is called the “_________”

• 1 Mole is the same as having 6.02 x 1023 “things”

• Examples:

" 1 mole of Carbon atoms = 6.02 x 1023 Carbon atoms " 1 mole of Oranges = 6.02 x 1023 Oranges " 1 mole of Basketballs = 6.02 x 1023 Basketballs " 1 mole of Teachers = 6.02 x 1023 Teachers

• Notice in your comparisons that even though the ______ of a teacher is significantly

greater than the _______ of an orange, you will still get 6.02 x 1023 of them in a mole!

• Molar Mass: ___________________________________________________________.

" For an element, the molar mass is shown on the __________________________ " For a compound, it is the _____ of the atomic masses for each of the elements that

comprise it.


Molar Mass Calculation Notes

Step 1. Begin by separating out the different atoms found in the compound. H2O:

Step 2. Beside each atomic symbol, write how many are found in the compound. H2O: Step 3. Find the atomic mass of each atom and write the mass next to number corresponding to its element. H2O: Step 4. Multiply the atomic mass with the number of atoms of that element in the compound formula. H2O: Step 5. Add all totals to get the molar mass of the compound. H2O: Try these harder ones: (NH4)2SO4: Cu(NO3)2 • 6H2O:


Diatomic Elements:

There are certain elements that exist as diatomic gas rather than single atom in nature. They are: Note:You must know these! To find molar mass of these gases, don’t forget to multiply the elemental molar mass by ___. Eg. Find the molar mass of hydrogen gas Molar Mass = Eg. Find the molar mass of chlorine gas Molar Mass =


Gram ! Mole & Mole ! Gram Conversions:

• The use of the MOLAR MASS allows us to calculate the mass of a given number of moles of a substance, and the calculation of the number of moles in a given mass of a substance.

• The unit for moles is “_____”

• In order to calculate from grams to moles, you need to use a ______________________!

• Conversion factors are used to relate the number of moles to the mass of material present.

• Since 1 mole of a given substance “X” has a mass of (molar mass of X) g…

or ____________________s Examples: What is the mass of 3.25 mol of CO2? [molar mass of CO2 = 44.0g/mol] What you want = (What you have) X (Conversion Factor) Mass of CO2 = 3.25 mol CO2 X molar mass of CO2 1 mol of CO2 = 3.25 mol CO2 X 44.0 g CO2

1 mol CO2 = 143 g CO2 How many moles of N2 are there in 50.0 g of N2? [molar mass of N2 = 28.0g/mol] What you want = (What you have) X (Conversion Factor) Moles of N2 = 50.0 g N2 X 1 mol N2 s Molar mass N2 = 50.0 g N2 X 1 mol N2 28.0 g N2 = 1.79 mol N2 • When you are calculating molar mass, the units for molar mass is ____________!


Moles ! Molecules & Molecules ! Moles Conversions:

• There are times when chemists would like to know the number of particles or molecules involved in a reaction.

• Thus, we need another set of Conversion Factors!

_____________________ or ______________________s

Examples: How many molecules are there in 0.125 mol of molecules? What you want = (What you have) X (Conversion Factor) # of molecules = 0.125 mol X 6.02 x 1023 molecules s 1 mol = 7.53 x 1022 molecules How many moles are in 1.00 x 1022 molecules of H2O? What you want = (What you have) X (Conversion Factor) Moles of H2O = 1.00 x 1022 molecules X 1 mol s 6.02 x 1023 molecules = 0.0166 moles

Try these problems in Hebden: pg 82 Q’s #8-10; pg 84 Q’s #15 bcdeg, 17 bd


Mole Ratios (using the three step method) Mole ! Mole

Magnesium metal reacts with hydrochloric acid to produce magnesium chloride and hydrogen gas. The chemical equation is given as the following:

Mg (s) + 2HCl (aq) ! MgCl2 + H2 (g)

If 4 moles of HCl is used, how many moles of MgCl2 will be formed in the reaction? Step (1): What is the mole ratio for HCl to MgCl2? Step (2): Write two conversion factors for this mole ratio. Step (3): Calculate the number of moles of MgCl2 formed.

_____________mol MgCl2 Now try these example questions on a separate sheet of paper using the 3-step method! 1. 3CaSi2 + 2SbCl3 ! 6Si + 2Sb + 3CaCl2

a) If 0.65 moles of CaSi2 is used, how many moles of Sb will be formed in the reaction?

b) If 1.47 moles of Si is formed in the reaction, how many moles of SbCl3 was used at the start of the reaction?

2. In a Chemistry 11 experiment, John mixes 0.010 moles of calcium chloride together with silver nitrate in a beaker. This results in the formation of silver chloride and calcium nitrate. How many moles of silver chloride are formed in the reaction?


• In chemistry, we often find ourselves asking the question: How many grams of chemical #1 must we use in order to produce 25 g of chemical #2?

• To do this, we must use stoichiometry! Defn: Stoichiometry: __________________________________________________________ __________________________________________________________ __________________________________________________________

• Previously we learned how to calculate gram to mole conversions and mole to mole conversions (mole ratio). We can now put them together and calculate gram to gram conversions.

Let’s look at an example:

4NH3 + 5O2 ! 6H2O + 4NO

If 10.0 g of ammonia is used, how many grams of water will be produced?

Step 1: Make a Road Map! We need to get from ______________ ! _________________ Road Map: ____________ ! _____________ ! ____________ ! ______________ Step 2: Carry out the conversions one step at a time! First do Grams NH3 ! Moles NH3: 10.0 g NH3 x 1 mol NH3 = 0.588 mol NH3 17. 0 g NH3

Next do Moles NH3 ! Moles H2O: 0.588 mol NH3 x 6 mol H2O = 0.882 mol H2O 4 mol NH3

Finally do Moles H2O ! Grams H2O: 0.882 mol H2O x 18.0 g H2O = 15.9 g H2O 1 mol H2O _________ g of water are produced from 10.0 g of ammonia


• So far we’ve learned how to calculate g! mol, mol ! mol, and g ! molecules calculations.

• We can also learn how to calculate molecules ! atoms as well! From Mole Map:

Example 1: How many atoms of oxygen are found in 1.00 x 1023 molecules of ozone (O3)? Road Map: molecules O3 ! atoms O Note: There are ____ atoms of oxygen for every molecule of ozone. Calculate:

1.00 x 1023 molecules O3 x 3 atoms O = 3.00 x 1023 atoms O 1 molecule O3 Example 2: How many atoms of hydrogen are there in 1.00 g of H2O? Road Map: g H2O ! mol H2O ! molecules H2O ! atoms H Note: There are ____ atoms of hydrogen for every molecule of H2O. Calculate: 1.00g H2O x 3 ato ms O = 3.00

Molecules (Particles)

Atoms

X # of atoms 1 molecule

X 1 molecule # of atoms


Ideal Gas: 1. John Dalton: (On the right)

Focused on ___________ of gas. E.g., 11.1g of H2 combine with 88.9g of O2 46.7g of N2 reacts with 53.3g of O2 42.9g of C reacts with 57.1g of O2 → No real pattern is observed :(

2. Joseph Gay-Lussac: Focused on __________ of gas. E.g., 1L of H2 combine with 1L of Cl2 to produce 2L of HCl 1L of N2 reacts with 3L of H2 to make 2L of NH3 2L of CO reacts with 1L of O2 to make 2L of CO2 → Gives nice relationship, but he could not explain why…

3. Avogadro: (On the left)

Came up with a hypothesis to explain Joseph’s experimental data: Avogadro’s hypothesis:

1. What does the hypothesis mean?→ Notice how the hypothesis is focused on the __________, not ________. This hypothesis tells us that 2L of gas AB will be made by combining 1L of reactant A and 1L of reactant B. 2. What does the hypothesis mean? → If two gases contains the same number of ___________ and are found under the same ___________ and ___________, both gases occupy the same ___________.

Standard Temperature and Pressure (STP)

Standard Temperature = ___________oC Standard Pressure = ___________kPa

Molar Volume (L/mol) (The ___________ occupied by ___________ mole of the gas.)

The volume of one mole of ANY gas at STP is___________L

Conversion factors: (22.4L/mol is true for only gases!, not liquid or solid)

!!!!!!!!!!!!!!!! or


Eg. How many moles of H2 are there in 10.0L of H2(g) at STP? Road Map: ______ of hydrogen → _______ of hydrogen

Try the harder one: Eg. How many atoms of chlorine are there in 10.0L of chlorine gas at STP? (Hint: Halogens are diatomic elements.) Road Map : ________ of chlorine → ________ of chlorine → ________ of chlorine → ________ of chlorine

When density is mentioned at any point in a problem, you should immediately remember that:

density (g/L) = mass (g) volume (L)

And find a way to get ________ and ________ through ________. (Use the heart of mole to help you get to the mass and volume.) Eg. How many moles are contained in 7.50L of C2H5OH(l)? (density 0.789g/mL) Road Map: Use ________ → ________ → ________

Another useful formula for a problem dealing with density is:

density (g/L) = molar mass (g/mol) molar volume (L/mol)

Note: ________ can be substituted for molar volume if gas is at STP.


Chemistry 11 – Percentage Composition:

Mission:

You have synthesized a sample thought to be caffeine, the stimulant found in coffee, tea and cola beverages in a lab. By elemental analysis,

determine if the sample you made is caffeine or not.

Every time when a chemist prepares a “new compound”, its chemical formula must be established beyond reasonable doubt. Usually, the molecule is analyzed by several methods, and each technique reveals information about the compound’s identity. When a compound is analyzed in the lab, the analysis involves finding the mass of each element in the compound. It is easier to compare the composition of compounds having the same elements (such as FeO and Fe2O3) if the composition is described by percentage composition. Percentage composition: _______________________________________________________

Let’s look at an example on how to calculate the percentage composition of H2O. You may think H2O is simply made of 67% hydrogen and 33% oxygen because of the molar ratio (H:O =2:1). But is this the case?

1. Assume there is _______ of the compound. Find the __________ of H2O.

2. Find the total mass of each element in the compound. (This is the amount in the brackets!)

Total mass of H = Total mass of O =

3. Divide the mass of each element by the molar mass and multiply by 100%.

% of H in compound = % of O in compound =

This means _______% of the mass of H2O comes from its hydrogen atoms. Every 100g of H2O contains _____g of hydrogen. If the calculation have been done right, the sum of the % composition of the individual elements will be 100%.

________ % H + ________% O = 100.%

We can use this to help us complete our mission!


Example: The sample you synthesized gave the following analysis:

A 50 g sample contains 24.8 g C, 2.6 g H, 14.4 g N, and 8.25 g O.

Does this analysis agree with the chemical formula of caffeine, which is C8H10N4O2 (MM= 194g/mol)? Look at the % composition of carbon in both samples! Your sample: Caffeine: The percentages indicate that your sample is _____________________!


Empirical and Molecular Formula

When we are analyzing an unknown compound in a lab, we can determine the percentage composition of the compound from the data results. Using the percentage composition, we can figure out the simplest ratio of atoms in this compound, also known as the empirical formula.

• Empirical Formula: _______________________________________________________ _______________________________________________________

For example, what do CH2, C2H4, C3H6, C4H8, and C5H10 have in common? They all contain ______ as many hydrogens as carbons. The empirical formula for all of these compounds is _______. Let’s take a look at how we can determine the empirical formula from the percentage composition data. What is the empirical formula of a compound consisting of 80.0 % C and 20.0 % H?

1. Assume you have _____ of the compound. Find the mass of each element in the sample.

2. Find the number of moles of each element by dividing the mass of each element by its

molar mass.

3. Find the ratio of atoms of each element by dividing the number of moles of each element by the smallest number of moles.

Use the following visual organizer

C H % 80.0 20.0 g 80.0 20.0

mol 6.67 20.0 /small 1 2.99 ~ 3

Empirical formula CH3

Therefore, the simplest ratio of atoms is ______________. The empirical formula is ________. There are times when the ratio may not be easily rounded to a whole number. If this is the case, you need to multiply this fraction or decimal number by a number (start with x2, x3, x4…. until decimals become __.99, __.98, or __.97 ) so that it becomes a whole number again.


Example: What is the empirical formula of a compound containing 39.0 % Si and 61.0 % O?

Si O % 39.0 61.0 g 39.0 61.0

mol 1.3879 3.8125 /small 1 2.746~2.75

Multiplier 4 10.987 ~ 11 Empirical formula Si4O11

Finding the Molecular Formula The empirical formula only gives us the simplest ratio of the atoms for each element in the compound. Thus, CH2, C2H4, C3H6, C4H8, and C5H10 all have the same empirical formula CH2. Say we were given one of these 5 organic compounds. How can we figure out which one we have for sure? We need to find the molecular formula for this compound!

• Molecular Formula: ______________________________________________________ ______________________________________________________

To find the multiple that converts the empirical formula to the molecular formula, you need to be given the molar mass of the compound. Multiple = N = Molar mass = Molar mass of compound 111221 Empirical mass Molar mass of empirical formula Once we have “N”, multiply it to the empirical formula to obtain the molecular formula!

Molecular Formula = N x (Empirical Formula Example:

A molecule has an empirical formula of HO and a molar mass of 34.0 g. What is the molecular formula? Use the following visual organizer

Empirical Molecular mass (g/mol) 17.0 34.0 formula HO ?

H2O2


Preview of Solution Chemistry:

What is solution? Solutions are made by dissolving a _________ into a ___________. Now, imagine making a cup of tea where the _______ added is the solute and the ___________ is the solvent. Depending on the __________ of sugar you decide to add, the tea will either taste a little bit sweet, or super sweet! There are two terms given to describe this:

• Concentrated: A ________ amount of solute is dissolved in the solution • Dilute: _______________ solute is dissolved in the solution

Thus, the super sweet tea is termed ______________, while the little bit sweet tea is termed ________. In chemistry, you often see bottles of solutions such as ______ with the terms _____ or ______ attached to the label. These units describe how concentrated or dilute the solution is. The most common expression of solution concentration involves the use of the “_______”.

• Molarity (Molar Concentration): ____________________________________________

____________________________________________

The units of molarity are ___________ or ______.

The _________ the number, the _______ concentrated a solution is. Thus, 1 M HCl is considered _________, while 12 M HCl is considered quite ____________________. Example: 10.0 g of NaCl is dissolved in water to make 100. mL of solution. What is the molar concentration of the solution? [Molar mass NaCl = 58.5 g / mol]

Step 1: Convert the mass to moles NaCl. Step 2: Convert the volume to litres. Step 3: Calculate the molarity of the solution.


Note: You may also be asked to solve for moles or Litres from the Molarity equation.

•

• s Example: What mass of NaNO3 is needed to make 1.50 L of a 0.200 M solution? [Molar mass = 85.0 g / mol] Road Map: ! ! Step 1: Step 2: Dilutions

Q: What happens to the concentration of a solution (eg. apple juice) when more water is added?

A: The concentration of the solution would ____________! (ie. apple juice doesn’t taste as ______) The process of making a solution ___________________ is called ___________. In general, to dilute a solution, you make the __________ of the solution __________. This is usually done by _________________ to the solution. To calculate the concentration of a diluted solution, the following dilution equation is used:

where: Mi = Initial concentration of the solution Vi = Initial volume of the solution *** Mf = Final concentration of the solution Vf = Final volume of the solution *** *** The units must be the same, not necessarily Litres!


Example: Ash adds 250. mL of water to 250. mL of a 0.200 M solution of NaCl. What is the concentration of the diluted solution?

Mi = ________, Vi = ________, Vf = ___________________, Mf = ________

Using the dilution equation:

Example: How much water was added to 200. mL of a 0.100 M solution of KCl to make a 0.0500 M solution?

Mi = _________, Vi = _________, Mf = __________, Vf = __________

Using the dilution equation:

However, Vf represents the total volume of the solution. Therefore, to determine the volume of water added:


Ion Concentrations in Solution:

So far, we know that in order to make a solution, a solute has to be dissolved in a solvent. How does an _______ compound (such as NaCl) ___________ when placed in a solution of ________? In order to dissolve, the ionic compound is broken down into its ______. One of the ions will be ______________ charged, while the other ion will be _______________ charged. A _______________________ is written to show the ions that are formed when a compound dissolves. Examples:

• LiF ! • K2O ! • Al2(SO4)3 !

In the last few classes, we have seen how to calculate the molarity of a substance, and how to calculate the molarity of a substance after dilution has occurred. What we need to do now is to investigate how to calculate the concentrations of ions in a solution after an ionic compound dissolves. Example: What are the concentrations of the ions in a 1.5 M Na3PO4 solution?

Step 1: Write the balanced dissociation reaction for the compound.

Step 2: Compare the mole ratios in the dissociation reaction.

Step 3: Calculate the molarity of the ions in the solution using the mole ratios.

In other words: Na3PO4 ! 3Na+ + PO43–

How would you calculate the concentration of the ions when you mix two different solutions together that contain the same _____________________?


Example: What is the concentration of each ion when 200. mL of 0.20 M NaCl is mixed with 200. mL of 0.50 M CaCl2?

Step 1: Calculate the concentration of the solutions after dilution (mixing) has occurred.

[NaCl]: Mi = 0.20 M Vi = 200. mL Mf = ? Vf = (200. + 200.) mL = 400. mL

(0.20 M) (200. mL) = Mf (400. mL) Mf = (0.20 M) (200. mL) = 0.10 M (400. mL)

[CaCl2]: Mi = 0.50 M Vi = 200. mL Mf = ? Vf = (200. + 200.) mL = 400. mL

(0.50 M) (200. mL) = Mf (400. mL) Mf = (0.50 M) (200. mL) = 0.25 M (400. mL)

Step 2: Write the balanced dissociation equations for each compound and calculate the ion concentrations.

Step 3: Add up the ion concentrations for the common ions in each compound.


Density Every solution (and object) has a unique _____________ that distinguishes it from one another. This property is known as the _____________.

• Density: The density of a solution is given by the following equation: Density = Mass or d = m Volume V The units for density are ___________ or ____________. For example, ______ has a density of _________. This means that _____ of water has a mass of _____. In a sense, the density of a solution can be used as a _____________________ in solving problems! 1 g H2O or 1 mL H2O 1 mL H2O 1 g H2O Example: Concentrated sulphuric acid has a density of 1.86 g / mL. What is the mass of the acid

present in 15.0 mL of the acid?

Mass = Density x Volume g H2SO4 = 1.86 g x 15.0 mL = 27.9 g

mL One other important fact is that the density of a substance can help us determine whether it will ________ or _________ when placed in another substance. A _________ dense liquid or object will float on liquids having a _______________ density! Have you ever wondered… Why does a person tend to float more easily while swimming in the ocean when compared to swimming at Watermania?

The water at Watermania has a density of about ____________. Since the human body is made up of around 60 – 75% water, we have a density just a bit ________ than 1.0 g / mL. Thus, we just ___________________ in water! The density of ocean water is about ______________. Thus, a human body would ________ more readily in the ocean as it is considerably _________________ than ocean water.

mole concept compiled sg - · pdf file! instead of counting 1, 2, 3, ... chemistry 11...

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