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Ionic Bonding Section 4.1.

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Presentation on theme: "Ionic Bonding Section 4.1."— Presentation transcript:

1 Ionic Bonding Section 4.1

2 Introduction to Bonding
Chemical bond: an interaction between atoms or ions that results in a reduction of the potential energy of the system, thereby becoming more stable Three types of bonds: ionic, metallic, and covalent The bond type depends on the atom's electronegativitites

3 More If the atoms have very different electronegativities, then ionic bonding occurs If they both have high electronegativities, then covalent bonding occurs If they both have low electronegativities, then metallic bonding occurs

4 Practice: What Kind of Bond?
Na and Cl Sr and O C and O Ni and Fe N and O Li and N Ti and Cr Ionic Covalent Metallic metallic

5 Valence Electrons Valence electrons are the electrons in the outermost energy level, which is the highest occupied energy level They are the electrons responsible for the chemical properties of atoms Core electrons – are those in the energy levels below.

6 Keeping Track of Electrons
Atoms in the same group have the same outer electronic structure and therefore the same number of valence electrons. The number of valence electrons is easily determined. It is the group number for a short group element Group 2: Be, Mg, Ca, etc. Each has 2 valence electrons

7 Electron Dot (Lewis Dot)diagrams
A way of showing & keeping track of valence electrons. Write the symbol - it represents the nucleus and inner (core) electrons Put one dot for each valence electron (8 maximum) They don’t pair up until they have to (Hund’s rule) X

8 The Electron Dot Diagram (Lewis Structure) for Nitrogen
Nitrogen has 5 valence electrons to show. First we write the symbol. N Then add 1 electron at a time to each side. Now they are forced to pair up as one side has two electrons

9 The Octet Rule The noble gases are unreactive in chemical reactions
In 1916, Gilbert Lewis used this fact to explain why atoms form certain kinds of ions and molecules The Octet Rule: in forming compounds, atoms tend to achieve a noble gas structure; 8 in the outer level is stable Each noble gas (except He, which has 2) has 8 electrons in the outer level

10 Formation of Cations Metals lose electrons (are oxidized) to attain a noble gas structure. They make positive ions (cations) If we look at the electronic structure, it makes sense to lose electrons: Na: 1s22s22p63s1 shows 1 valence electron Na1+: 1s22s22p6 This is a noble gas structure with 8 electrons in the outer level.

11 Electron Dots For Cations
Metals have few valence electrons (usually 3 or less); calcium has only 2 valence electrons Ca

12 Electron Dots For Cations
Metals will lose the valence electrons Ca

13 Electron Dots For Cations
Form positive ions Ca2+ This is named the “calcium ion”. No dots are now shown for the cation.

14 Electron Configurations: Anions
Nonmetals gain electrons to attain noble gas electronic structures. They make negative ions (anions) S = 1s22s22p63s23p4 = 6 valence electrons S2- = 1s22s22p63s23p6 = noble gas structure. Halide ions are ions from chlorine or other halogens that gain electrons

15 Electron Dots For Anions
Nonmetals will have many valence electrons (usually 5 or more) They will gain electrons to fill outer shell. 3- P (This is called the “phosphide ion”, and should show dots)

16 Stable Electron Configurations
All atoms react to try and achieve a noble gas structure. Noble gases have 8 valence electrons and so are already stable This is the octet rule (8 in the outer level is particularly stable). Ar

17 Ionic Bonding Anions and cations are held together by opposite charges (+ and -) Simplest ratio of elements in an ionic compound is called the formula unit (also called the empirical formula). The bond is formed through the transfer of electrons (lose and gain) Electrons are transferred to achieve noble gas structure.

18 Ionic Bonding Na Cl The metal (sodium) tends to lose its one electron from the outer level. The nonmetal (chlorine) needs to gain one more to fill its outer level, and will accept the one electron that sodium is going to lose.

19 Ionic Bonding Na+ Cl - Note: Remember that no dots are now shown for the cation

20 Ionic Bond Negative charges are attracted to positive charges.
Negative anions are attracted to positive cations. The result is an ionic bond. A three-dimensional crystal lattice of anions and cations is formed.

21 Ionic Compounds The ionic substance is held together by strong electrostatic attractions between all ions in all three dimensions No molecules present An ionic lattice is formed This gives them distinct physical properties

22 Preserve Electroneutrality
When ions combine, electroneutrality must be preserved. In the formation of magnesium chloride, 2 Cl- ions must balance a Mg2+ ion: Mg Cl- → MgCl2

23 NaCl CsCl TiO2

24 Please go to the “Naming ions, compounds and molecules presentation.

25 Properites of Ionic Compounds
Hard, brittle crystalline solids Relatively high melting and boiling points Do not conduct electricity when solid, but do when molten or in aqueous solution Are more soluble in water than other solvents

26 - Page 198 The ions are free to move when they are molten (or in aqueous solution), and thus they are able to conduct the electric current.

27 Predicting Ionic Charges
Group 1A: Lose 1 electron to form 1+ ions H+ Li+ Na+ K+ Rb+

28 Predicting Ionic Charges
Group 2: Loses 2 electrons to form 2+ ions Be2+ Mg2+ Ca2+ Sr2+ Ba2+

29 Predicting Ionic Charges
Group 3: Loses 3 electrons to form 3+ ions B3+ Al3+ Ga3+

30 Predicting Ionic Charges
Nitride Group 5: Gains 3 electrons to form 3- ions P3- Phosphide As3- Arsenide

31 Predicting Ionic Charges
Oxide Gains 2 electrons to form 2- ions Group 6: S2- Sulfide Se2- Selenide

32 Predicting Ionic Charges
Gains 1 electron to form 1- ions Group 7: F- Fluoride Br- Bromide Cl- Chloride I- Iodide

33 Predicting Ionic Charges
Stable noble gases do not form ions! Group 0:

34 Predicting Ionic Charges
Many transition elements have more than one possible oxidation state. Note the use of Roman numerals to show charges Iron (II) = Fe2+ Iron (III) = Fe3+

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