Chapter 6.3 “Ionic Bonding”

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 atoms’ electronegativities

More Subtract the electronegativities to determine the nature of the bond If the difference is greater than 1.7, the bond is ionic If the difference is from 0 to 0.3, the bond is non-polar covalent If the difference is from 0.3 to 1.7, the bond is polar covalent

Summary 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

Practice: What kind of bond? Na and Cl Sr and O C and O Ni and Fe N and O H and B Ti and Cr ionic polar covalent metallic non-polar covalent

Valence Electrons are…? The electrons responsible for the chemical properties of atoms, and are those in the outer energy level. Valence electrons - The s and p electrons in the outer energy level the highest occupied energy level Core electrons – are those in the energy levels below.

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 groups 1 and 2 Group 1: H, Li, Na, K, etc. have 1 valence e- Group 2: Be, Mg, Ca, etc. have 2 valence e-

More about Keeping Track For elements in groups 13-17: Subtract 10 from the group number This is the number of valence electrons

Electron Dot diagrams are… A way of showing & keeping track of valence electrons. How to write them? 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

The Electron Dot diagram 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. We have now written the electron dot diagram for nitrogen.

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 configuration; 8 in the outer level is stable Each noble gas (except He, which has 2) has 8 electrons in the outer level

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

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

Electron Dots For Cations Metals will have few valence electrons Metals will lose the valence electrons Ca

Electron Dots For Cations Metals will have few valence electrons Metals will lose the valence electrons Forming positive ions Ca2+ This is named the “calcium ion”. No dots are now shown for the cation.

Electron Dots For Cations Let’s do scandium, #21 The electron configuration is: 1s22s22p63s23p64s23d1 Thus, it can lose 2e- (making it 2+), or lose 3e- (making 3+) Sc = Sc2+ Point out that the dots are only shown to show you the different oxidation states, do not show dots with cations for the final electron dot structure Sc = Sc3+ scandium (II) ion scandium (III) ion

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

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)

Stable Electron Configurations All atoms react to try and achieve a noble gas configuration. Noble gases have 2 s and 6 p electrons. 8 valence electrons = already stable! This is the octet rule (8 in the outer level is particularly stable). Ar

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

Ionic Compounds Also called salts Made from: a cation with an anion (or literally from a metal combining with a nonmetal)

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.

Ionic Bonding Na+ Cl 1- Note: Remember that NO DOTS are now shown for the cation!

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.

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

Ionic Bonding Let’s do an example by combining calcium and phosphorus: Ca P All the electrons must be accounted for, and each atom will have a noble gas configuration (which is stable).

Ionic Bonding Ca P

Ionic Bonding Ca2+ P

Ionic Bonding Ca2+ P Ca

Ionic Bonding Ca2+ P 3- Ca

Ionic Bonding Ca2+ P 3- Ca P

Ionic Bonding Ca2+ P 3- Ca2+ P

Ionic Bonding Ca Ca2+ P 3- Ca2+ P

Ionic Bonding Ca Ca2+ P 3- Ca2+ P

Ionic Bonding Ca2+ Ca2+ P 3- Ca2+ P 3-

= Ca3P2 Ionic Bonding Formula Unit This is a chemical formula, which shows the kinds and numbers of atoms in the smallest representative particle of the substance. For an ionic compound, the simplest ratio of the ions is called a formula unit

Properties of Ionic Compounds Crystalline solids - a regular repeating arrangement of ions in the solid: Ions are strongly bonded together. Structure is a rigid crystal lattice. High melting points Coordination number- number of ions of opposite charge surrounding it

NaCl CsCl TiO2 - Page 198 Coordination Numbers: Both the sodium and chlorine have 6 NaCl Both the cesium and chlorine have 8 CsCl Each titanium has 6, and each oxygen has 3 TiO2

Do they Conduct? Conducting electricity means allowing charges to move. In a solid, the ions are locked in place. Ionic solids are insulators. When melted, the ions can move around. Melted ionic compounds conduct. NaCl: must get to about 800 ºC. Dissolved in water, they also conduct (free to move in aqueous solutions)

- 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.

Atoms and Ions Atoms are electrically neutral. Because there is the same number of protons (+) and electrons (-). Ions are atoms, or groups of atoms, with a charge (positive or negative) They have different numbers of protons and electrons. Only electrons can move, and ions are made by gaining or losing electrons.

F1- O2- An Anion is… A negative ion. Has gained electrons. Nonmetals can gain electrons. Charge is written as a superscript on the right. Has gained one electron (-ide is new ending = fluoride) F1- O2- Gained two electrons (oxide)

K+ Ca2+ A Cation is… A positive ion. Formed by losing electrons. More protons than electrons. Metals can lose electrons K+ Has lost one electron (no name change for positive ions) Ca2+ Has lost two electrons

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

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

Predicting Ionic Charges Group 13: Loses 3 electrons to form 3+ ions B3+ Al3+ Ga3+

Predicting Ionic Charges Do they lose 4 electrons or gain 4 electrons? Neither! Group 14 elements rarely form ions (they tend to share) Group 14:

Predicting Ionic Charges nitride Group 15: Gains 3 electrons to form 3- ions P3- phosphide As3- arsenide

Predicting Ionic Charges oxide Gains 2 electrons to form 2- ions Group 16: S2- sulfide Se2- selenide

Predicting Ionic Charges Gains 1 electron to form 1- ions Group 17: F- fluoride Br- bromide Cl- chloride I- iodide

Predicting Ionic Charges Stable noble gases do not form ions! Group 18:

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+

Naming cations Two methods can clarify when more than one charge is possible: Stock system – uses roman numerals in parenthesis to indicate the numerical value Classical method – uses root word with suffixes (-ous, -ic) Does not give true value

Naming cations We will use the Stock system. Cation - if the charge is always the same (like in the main group of metals) just write the name of the metal. Transition metals can have more than one type of charge. Indicate their charge as a roman numeral in parenthesis after the name of the metal

Predicting Ionic Charges Some of the post-transition elements also have more than one possible oxidation state. tin (II) = Sn2+ lead (II) = Pb2+ tin (IV) = Sn4+ lead (IV) = Pb 4+

Predicting Ionic Charges Some transition elements have only one possible oxidation state, such as these three: (memorize these) silver = Ag+ zinc = Zn2+ cadmium = Cd2+

Exceptions: Some of the transition metals have only one ionic charge: Do not need to use roman numerals for these: silver is always 1+ (Ag+) cadmium and zinc are always 2+ (Cd2+ and Zn2+)

Practice by naming these: Ca2+ Al3+ Fe3+ Fe2+ Pb2+ Li+

Write symbols for these: potassium ion magnesium ion copper (II) ion chromium (VI) ion barium ion mercury (II) ion

Anions are always the same charge Naming Anions Anions are always the same charge Change the monatomic element ending to – ide F- a fluorine atom will become a fluoride ion.

Practice by naming these: Cl- N3- Br- O2- Ga3+

Write symbols for these: sulfide ion iodide ion phosphide ion strontium ion

Polyatomic ions are… Groups of atoms that stay together and have an overall charge, and one name. Usually end in –ate, -ite, or -ide acetate: C2H3O2- nitrate: NO3- nitrite: NO2- hydroxide: OH- cyanide: CN-

Common Polyatomic Ions phosphate: PO43 ammonium: NH4+ sulfate: SO42- carbonate: CO32- chromate: CrO42- chlorate: ClO3- chlorite: ClO2- (One of the few positive polyatomic ions) If the polyatomic ion begins with H, then combine the word hydrogen with the other polyatomic ion present: H+ + CO32- → HCO3- hydrogen + carbonate → hydrogen carbonate ion

Writing Ionic Compound Formulas Example: iron (III) chloride (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Fe3+ Cl- 3 2. Check to see if charges are balanced. Not balanced! Now balanced. 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts. = FeCl3

Writing Ionic Compound Formulas Example: aluminum sulfide (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Al3+ S2- 2 3 2. Check to see if charges are balanced. Not balanced! Now balanced. 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts. = Al2S3

Writing Ionic Compound Formulas Example: magnesium carbonate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Mg2+ CO32- 2. Check to see if charges are balanced. They are balanced! = MgCO3

Writing Ionic Compound Formulas Example: barium nitrate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Ba2+ ( ) NO3- 2 2. Check to see if charges are balanced. Now balanced. Not balanced! 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance subscripts. = Ba(NO3)2

Writing Ionic Compound Formulas Example: ammonium sulfate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! ( ) NH4+ SO42- 2 Now balanced. 2. Check to see if charges are balanced. Not balanced! = (NH4)2SO4 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts.

Writing Ionic Compound Formulas Example: zinc hydroxide (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! ( ) Zn2+ OH- 2 2. Check to see if charges are balanced. Not balanced! Now balanced. 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts. = Zn(OH)2

Writing Ionic Compound Formulas Example: aluminum phosphate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Al3+ PO43- 2. Check to see if charges are balanced. They ARE balanced! = AlPO4

Naming Ionic Compounds 1. Name the cation first, then anion 2. Monoatomic cation = name of the element Ca2+ = calcium ion 3. Monoatomic anion = root + -ide Cl- = chloride CaCl2 = calcium chloride

Naming Ionic Compounds (Metals with multiple oxidation states) some metals can form more than one charge (usually the transition metals) use a Roman numeral in their name: PbCl2 – use the anion to find the charge on the cation (chloride is always 1-) Pb2+ is the lead (II) cation PbCl2 = lead (II) chloride

Things to look for: If cations have ( ), the number in parenthesis is their charge. If anions end in -ide they are probably off the periodic table (Monoatomic) If anion ends in -ate or –ite, then it is polyatomic

Practice by writing the formula or name as required… iron (II) phosphate potassium sulfide ammonium chromate MgSO4 FeCl3

Practice by writing the formula for the following: magnesium hydroxide iron (III) hydroxide zinc hydroxide

Hydrates Some compounds contain H2O in their struc-ture. These compounds are called hydrates. The H2O can usually be removed if heated. A dot separates water: e.g. CuSO4•5H2O is copper(II) sulfate pentahydrate. A Greek prefix indicates the # of H2O groups. sodium sulfate decahydrate nickel(II) sulfate hexahydrate Na2CO3•H2O BaCl2•2H2O Na2SO4•10H2O NiSO4•6H2O sodium carbonate monohydrate barium chloride dihydrate

Prefixes 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 deca

Hydrates Examples: I. Give the name of the following: 1. CuSO4  5H2O 2. MgCl2  6H2O 3. Na2SO4  10H2O II. Write the formula for: 1. zinc chloride hexahydrate 2. calcium phosphate dihydrate 3. copper (I) chloride pentahydrate

Helpful to remember... 1. In an ionic compound, the net ionic charge is zero (criss-cross method) 2. An -ide ending generally indicates a binary compound 3. An -ite or -ate ending means there is a polyatomic ion that has oxygen 4. Prefixes generally mean molecular; they show the number of each atom

Helpful to remember... 5. A Roman numeral after the name of a cation is the ionic charge of the cation