1. Noadswood Science, 2012

2. Mass, Moles & Formula  To know how to calculate formula mass and work out the percentage of an element within a compound Saturday, September 12, 2015

3. Mass Number  Each atom consists of a nucleus containing protons and neutrons, with electrons arranged around it.  Protons and neutrons have a relative mass unit of 1  Electrons have a very small mass compared to protons and neutrons with a mass unit of 1/2000 th  Most of the mass of an atom is found in its nucleus: -  The mass number of an atom is the total number of protons and neutrons it contains  The atomic number (also called the proton number) is the number of protons it contains

4. Chemical Symbol  The full chemical symbol for an element shows its mass number at the top, and atomic number at the bottom, e.g. carbon  It tells us that a carbon atom has six protons (it will also have six electrons, because the number of protons and electrons in an atom is the same)  The symbol also tells us that the total number of protons and neutrons in a carbon atom is 12  *You can work out the number of neutrons from the mass number and atomic number – in this example, it is 12 - 6 = 6 neutrons

5. Relative Atomic Mass (A r )  The atoms of each element have a different mass – the relative atomic mass is a way of saying how heavy atoms are in comparison to others  Carbon is given a relative atomic mass (RAM) of 12  The RAM of other atoms compares them with carbon ElementSymbolx Heavier Than Carbon (C)RAM HeliumHe1/34 BerylliumBe3/49 MolybdenumMo896 KryptonKr784 OxygenO1 & 1/316 SilverAg9108 CalciumCa3 & 1/340

6. Formula  Formulas tell you how many atoms are used but not the mass of each of the elements  For example in the reaction: - Mg + 2HCl → MgCl 2 +H 2 O  This shows us we need twice as many hydrogen and chlorine atoms, but not the masses needed (because different atoms have different masses)  Even though we are using more hydrogen atoms the mass of 2 atoms will be less than the mass of 1 magnesium

7. Relative Formula Mass (M r )  Relative formula mass is the sum of all the relative atomic masses added together  For example, MgCl 2 has a relative formula mass of 95: - MgCl 2  *Relative atomic mass for any element can be found using the periodic table – relative formula mass is also ‘molar mass’… Mg = 24Cl = 35.5 x 2

8. Relative Formula Mass  Work out the following relative formula masses using the atomic masses given (N = 14; H = 1; Na = 23; O = 16; Mg = 24; Ca = 40): -  Ammonia NH 3  Sodium oxide Na 2 O  Magnesium hydroxide Mg(OH) 2  Calcium nitrate Ca(NO 3 ) 2  Ammonia = 14 + (3 x 1) = 17  Sodium oxide = (2 x 23) + 16 = 62  Magnesium hydroxide = 24 + 2 x (16 + 1) = 58  Calcium nitrate = 40 + 2x (14 + (3 x 16)) = 164

9. % Mass  Percentage composition is just a way to describe what proportions of the different elements there are in a compound  If you have the formula of a compound, you are able to work out the percentage by mass of an element in it using relative atomic mass and relative formula mass… % mass of an element in a compound = A r x number of atoms (of element) x 100 M r (of whole compound)

10. % Mass Example – CO 2  % of oxygen in carbon dioxide  First find the atomic masses – C = 12 and O = 16  Formula = CO 2  Number oxygen atoms = 2  Atomic mass of O = 16  Formula mass CO 2 = 12 + (16 x 2) = 44  % oxygen is therefore… 16 x 2 x 100 44 = 72.7%

11. % Mass Example – Na 2 CO 3  % of sodium in sodium carbonate  First find the atomic masses – Na = 23, C = 12 and O = 16  Formula = Na 2 CO 3  Number sodium atoms = 2  Atomic mass of Na = 23  Formula mass Na 2 CO 3 = (23 x 2) + 12 + (16 x 3) = 106  % sodium is therefore… 23 x 2 x 100 106 = 43.4%

12. Moles  The mole is a name given to a Avogadro's number: 6.023 x 10 23 or 602,300,000,000,000,000,000,000  When you have precisely this number of atoms or molecules (of any element or compound) they weigh exactly the same number of grams as the relative atomic mass of the element or compound  1 mole of atoms of molecules of any substance will have a mass in grams equal to the relative formula mass (A r or M r ) for that substance

13. Moles  1 mole of atoms of molecules of any substance will have a mass in grams equal to the relative formula mass (A r or M r ) for that substance  For example: -  Carbon has an A r of 12 – 1 mole of carbon weighs 12g  Iron has an A r of 56 – 1 mole of carbon weighs 56g  Carbon dioxide has a M r of 44 – 1 mole of carbon dioxide weighs 44g  This means that 12g of carbon, 56g of iron, 44g of carbon dioxide etc… all contain the same number of atoms – 1 mole!

14. Moles Within A Mass  The number of moles within a given mass is worked out using the following equation: - Number of moles = Mass (g) of element or compound M r of element or compound  For example, to calculate the number of moles in 42g of carbon… Number of moles = 42g 12 = 3.5 moles

15. Molar Solutions  A 1 molar solution contains 1 mole per litre Number of moles = volume (litres) x molarity solution  For example – how many moles are in 185cm 3 of a 2M solution? 0.185 x 2 = 0.37 moles