Bohr-Rutherford Diagram

1. Bohr-Rutherford Diagram

2. Bohr-Rutherford Diagram • Way to show the arrangement of electrons around the nucleus for the first 20 elements • Nucleus holds the protons and neutrons, which we know from the atomic number • The electrons are drawn on the orbits (rings) around the nucleus

3. Look at the Element Hydrogen

4. There is a maximum number of electrons each orbit can hold • Excess electrons move to the next orbit around the nucleus • The first orbit holds 2 • The second and third orbit holds 8 • The fourth orbit holds 18

5. Examples Lets draw the Bohr Rutherford diagram of Potassium K Lets draw the Bohr-Rutherford diagram of fluorine F

6. Do you notice anything about the electrons on the outer-most orbit?

7. The electrons on the outer-most orbit are responsible for the reactivity of an atom • Atoms want their outer orbit full • The first orbit wants two electrons • The second and third want eight electrons • The fourth wants eighteen electrons

8. They fill their outer orbit by reacting with other elements • This explains why elements form compounds • Elements that have a full outer orbit are more stable

9. Activity

10. Ionic Compounds/Covalent Compounds • Ionic compounds - occurs between ions, which are negatively charged or positively charged ions • Atoms that don’t have a full outer orbit either give up electrons or take electrons to fill their shell giving them a positive or negative charge • Ionic compounds occur between metals and non-metals • Covalent compounds/molecular compounds - want to fill their outer shell but they share the electrons between atoms instead of losing or gaining electrons • There is no charge • Molecular compounds occur between non-metals and non-metals

11. Ions and Ionic Compounds • Look at Sodium and Chloride

12. You can look at the periodic table to determine whether an ion gives electrons or takes electrons • Ions have similar reactivity according to their group, i.e. Alkali Metals are likely to give up one electron, etc. • This tells us how many electrons the ion needs to give up or gain to have a full outer-shell; this is important when writing chemical formulas for ionic compounds

13. Chemical Formulas for Ionic Compounds • When writing the chemical formula of an ionic compound you don’t just write the ions; for example Na+ Cl-, it is written as NaCl • If we look at the ionic compound that involves Mg2+ and Cl1- Magnesium needs to give up 2 electrons to have a full outer shell, while Chlorine needs only 1 electron to have a full outer shell • Therefore, we need 2 chlorine atoms per magnesium ion to satisfy both ions

14. Look at Mg2+ and Cl-

15. Therefore, the chemical formula for Mg2+ and Cl1- is MgCl2 • You can also figure out the chemical formula for an ionic compound by using the criss-cross method, and drop the charges • Mg2+ Cl1- = MgCl2 • Example write the chemical formula for Al3+ and O2-

16. Al2O3 • Some ions have multiple charges i.e. Iron (Fe) can either give up 3 electrons or 2 electrons depending on the ion it is bonding with • You can identify which ion is used in an ionic compound by looking at the chemical formula • For example if we look at Fe2O2

17. For example if we look at Fe2O2 • Using the criss-cross method you can identify which Fe ion was used • Fe2 O2 which becomes Fe2+ and O2- • You always reduce ionic compounds when writing their chemical formula, so Fe2O2 becomes FeO • Using the chemical formula Ni2S2 which Nickel ion is involved in this ionic compound.

18. Naming Ionic Compounds • Naming ionic compounds involves writing the metal first than the non-metal and changing the non-metals ending to “ide” • Ex. MgCl2 Magnesium Chloride • Name this ionic bond, Al2O3

19. Activity Sheet