Ions

1. Ions Unit 7 Chapter 6

2. Getting Started • When the difference in electronegativity of two atoms is greater than 1.6, • electrons are transferred from one atom to the other. • Since electrons are negatively charged, when an atom loses an electron, • it becomes positively charged.

3. Q is the Loneliest Number… • When an atom gains an electron, • it becomes negatively charged. • You can calculate the charge, Q, • by subtracting the number of electrons from the number of protons: • Q = p+ – e- • Aluminum has 13 protons. If it loses 3 electrons, Q = 13 – 10 = +3

4. Positively Shocking • A cation is a positively charged ion. • They form when atoms lose electrons • There are more protons than electrons. • All metals form cations. • Cations have no special name: • Na+ is the sodium ion.

5. Don’t Be So Negative! • An anion is a negatively charged ion. • They form when atoms gain electrons • There are more electrons than protons. • Most nonmetals will form anions.

6. The “-ides” of Chemistry • Anions are named by dropping the final syllable(s) of the element • and adding “-ide”. • Chlorine becomes chloride, Cl- • Oxygen becomes oxide, O-2

7. Mono-a-Mono • Monatomic ions are ions made up of single atoms. • For example, magnesium forms the Mg+2 cation • as one single positively charged atom. • Bromine forms the Br-1 anion • as one single negatively charged atom.

8. Poly Wanna Covalent Ion!?! • Polyatomic ions are ions made up of more than one atom. • In polyatomic ions, covalent bonds form between atoms within the ion. • The covalent bonds do not dissociate in water. • Electrons have been gained or lost during formation.

9. The Cyanide Ion (CN-) • Carbon has 4 valence electrons and N has 5. • A covalent bond forms between C and N sharing 3 electrons each. • Carbon steals an extra electron • The whole molecule acquires a negative charge. N C -1

10. Polyelectric…is that like GE? • Polyatomic ions form ionic bonds with other ions. • These ionic bonds will dissociate, • and because of this, polyatomic ions can conduct electricity • just like monatomic ions.

11. Making ‘em Big! • The sulfate ion is SO4-2. • It is made up of one sulfur atom • covalently bonded to four oxygen atoms. • When the sulfate ion was formed, • it consumed two extra electrons • (probably from some metal).

12. A Little of this… • The magnesium ion, Mg+2, • can combine with the sulfate ion, SO4-2, • to form magnesium sulfate, MgSO4 • also known as Epsom Salts.

13. It’s ‘Cause They’re So Cool • When we talked about the Noble Gases, • we learned that they are nonreactive because they have full s and p orbitals (d & f, too). • When the outer s and p orbitals are full, the atom is very stable. • Ions form when atoms attempt to fill their outer s and p orbitals to be • Isoelectronic with the Noble Gases.

14. The Yin & Yang of Stabilization • Orbitals are stabilized when they are completely full, completely empty, or half-full (d and f orbitals only). • Since an s-orbital can hold 2 electrons and a p-orbital can hold 6, • atoms are stable when they have 8 electrons in their valence shell • (we usually ignore d’s and f’s).

15. Eight is Enough • The Octet Rule: • Atoms will attempt to gain or lose electrons • to have 8 electrons • in their outer s & p orbitals.

16. Hangin’ on the p-Block • s- & p-block elements almost always obey this rule, • but d- & f-block elements do not always. • d and f-block elements have lots of orbitals with nearly overlapping energies. • Because of this, many elements in the d and f-blocks can form ions with more than one charge.

17. Predicting Charges • We can look at the electron configuration of a neutral atom to predict the charge of its ion. • Na = [Ne]3s1 [Ne]3s0 = Na+1 • S = [Ne]3s23p4  [Ne]3s23p6 = S-2 • Be = [He]2s2  [He]2s0 = Be+2 • N = [He]2s22p3 [He]2s22p6 = N-3

18. Ions