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## Topic 6- Relative Formula Mass

• Relative Atomic Mass, Ar.
• Ar is same as mass number
• Bigger number for each element= Relative Atomic Mass (Ar)
• Relative Formula Mass, Mr.
• Relative formula mass= relative atomic masses added together
• For MgCl2= Mg- (24 +) Cl2 = (35.5 x 2) = 95
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## Topic 6- Two Formula Mass Calculations

• Pecentage mass of an element in a compound = Ar x No. of atoms (of that element)/Mr (of whole compound) x 100
• Percentage mass of sodium carbonate: Na2Co3
• Ar of sodium = 23  Ar of carbon = 12  Ar of oxygen = 16
• Mr of Na2CO3= (2x23) + 12 + (3 x 16) = 106
• ((23x2)/2) x 100 = 43.4%
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## Topic 6- Finding the Empirical Formula (from masse

• List all the elements in the compound
• Underneath them write their experimental masses
• Divide each mass by the Ar for that particular element
• Turn the numbers you get into a nice somple ratio by multiplying or dicind them by well chosen numbers
• When the ratio is in its simplest form, this tells you the empirical formula of the compound.
• Example: Find the empirical formula of the magnesuim oxide produced when 9.6g of Magnesium react with 6.4g of Oxygen. (Ar for Magnesium= 24, Ar for Oxygen = 16.)
• List two elements: Mg- 0
• Write in experimental masses: 9.6                   6.4
• Divide by the Ar: 9.6/24=0.4  6.4/16=0.4
• Multiply by 10:     4:4
• Then divide by 4:     1:1
• Simplest formula: 1 atom of Mg to 1 atom of O
• MgO
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## Topic 6- Calculating Masses In Reactions

• Three Important steps
• Write out balanced equation
• Work out Mr- just for the bits you want.
• Apply the rule: Divide to get one, then multiply to get all.
• Example: What mass of magnesium oxide is produced when 60g of magnesium is burned in air?
• 2Mg + 02 >> 2mg0
• 2 x 24 >> 2 x (24 +16)
• 48 >> 40
• 48 and 80 tell us that 48g of Mg react to give 80g of MgO.
• 1g of Mg reacts to give 1.67g of MgO
• 60g of Mg to give us 100g of Mgo
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## Topic 6- Percentage Yield

• The yield of a reaction is the mass of product it produces
• The more reactants you start with, the higher the actual yield will be- that's pretty obvious.
• You can work out the theoratical (predicted) yield of a reaction from the balanced reaction equation.
• The percentage yield compares the actual yield with the theoratical yield. It doesn't depend on the amount of reactants you started with.
• Percentage yield is always somewhere between 0% and 100%.
• 100% - means that you got all the product you expected to get.
• 0%- means no reactants were converted into product. i.e. no product was made at all
• Percentage yield = actual yield (grams)/ theoratical yield (grams) x 100
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## Topic 6- Yields are always less than 100%

1) Incomplete reactions- if not all the reactants are converted to product, the reaction is incomplete and the yield will be lower than expected.

2) Practical losses during preparation- you always lose a bit when you transfer chemicals from one container to another. Think about transferring a liquid to a new container- some of it always get left behind.

3) Unwanted reactions- things don't always go exactly to plan. Sometimes you get unexpected reactions happening, so the yield of the intended product goes down. There an be caused by imurities in the reactants, but sometimes just changing the reaction conditions affects what products you make.

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## Topic 6- Waste In Reactions Cost Money

• Reactions that make lots of waste aren't usually profitable. This is becasue waste products aren't commercially useful.
• If a waste product is harmful, it can pose a threat to people and the enviroment. Disposing of harmful waste products safely can be very expensive.
• Chemists in industry are always looking for ways to produce products in a way that safely makes the most profit. They work to find reactions with there characterisitics:
• They give high percentage yield- so lots of product is made from the expensive raw materials.
• All of the products are comerically useful so there isn't any waste.
• They are a suitable speed- so the products are made quickly and safely.
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