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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 ionic polar covalent metallic polar covalent non-polar covalent metallic 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. l First we write the symbol. lThen add 1 electron at a time to each side. lNow they are forced to pair up. lWe have now written the electron dot diagram for nitrogen. l N The Octet Rule l l l 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 l 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 2+ Ca 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 2+ = Sc scandium (II) ion Sc = 3+ Sc 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. P 3(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 1) Also called salts 2) 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 2+ Ca P Ionic Bonding 2+ Ca Ca P Ionic Bonding 2+ Ca Ca P 3- Ionic Bonding 2+ Ca P Ca P 3- Ionic Bonding 2+ Ca P 2+ Ca P 3- Ionic Bonding Ca 2+ Ca P 2+ Ca P 3- Ionic Bonding Ca 2+ Ca P 2+ Ca P 3- Ionic Bonding 2+ Ca 2+ Ca 2+ Ca P P 33- Ionic Bonding = Ca3P2 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 1. Crystalline solids - a regular repeating arrangement of ions in the solid: – – Ions are strongly bonded together. Structure is a rigid crystal lattice. 2. High melting points • Coordination number- number of ions of opposite charge surrounding it - Page 198 Coordination Numbers: NaCl Both the sodium and chlorine have 6 CsCl Both the cesium and chlorine have 8 TiO2 Each titanium has 6, and each oxygen has 3 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. 3. 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. An Anion is… • A negative ion. • Has gained electrons. • Nonmetals can gain electrons. • Charge is written as a superscript on the right. 1F Has gained one electron (-ide is new ending = fluoride) 2O Gained two electrons (oxide) A Cation is… A positive ion. Formed by losing electrons. More protons than electrons. Metals can lose electrons + K 2+ Ca Has lost one electron (no name change for positive ions) 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 B3+ Al3+ Ga3+ Group 13: Loses 3 electrons to form 3+ ions Predicting Ionic Charges Neither! Group 14 elements rarely form ions (they tend to share) Group 14: Do they lose 4 electrons or gain 4 electrons? Predicting Ionic Charges N3- nitride P3- phosphide As3- arsenide Group 15: Gains 3 electrons to form 3- ions Predicting Ionic Charges O2- oxide S2- sulfide Se2- selenide Group 16: Gains 2 electrons to form 2- ions Predicting Ionic Charges F- fluoride Cl- chloride Group 17: Gains Br- bromide 1 electron to form 1- ions I- iodide Predicting Ionic Charges Group 18: Stable noble gases do not form ions! Predicting Ionic Charges Many transition elements have more than one possible oxidation state. Note the use of Roman iron (II) = Fe2+ numerals to show charges iron (III) = Fe3+ Naming cations • Two methods can clarify when more than one charge is possible: 1) Stock system – uses roman numerals in parenthesis to indicate the numerical value 2) 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: • • • • • • • Na+ 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 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 2•O • 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 • • • • • 2- sulfate: SO4 carbonate: CO32chromate: CrO42chlorate: ClO3chlorite: ClO2- • phosphate: PO43 + • ammonium: NH4 (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- → HCO3hydrogen + 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- 2. Check to see if charges are 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. 3 Not balanced! Now balanced. = FeCl3 Writing Ionic Compound Formulas Example: aluminum sulfide (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are 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. 3+ Al 2 2S 3 Not balanced! Now balanced. = Al2S3 Writing Ionic Compound Formulas Example: magnesium carbonate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. Mg2+ CO32They 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! 2+ Ba ( NO3 )2 2. Check to see if charges are 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. Now balanced. Not balanced! = 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! 2. Check to see if charges are balanced. ( NH4+ ) SO42- 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. 2 Now balanced. Not balanced! = (NH4)2SO4 Writing Ionic Compound Formulas Example: zinc hydroxide (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! 2+ Zn 2. Check to see if charges are 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. ( OH- )2 Not balanced! Now balanced. = 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! 2. Check to see if charges are balanced. 3+ Al PO4 3- 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: 1) If cations have ( ), the number in parenthesis is their charge. 2) If anions end in -ide they are probably off the periodic table (Monoatomic) 3) 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 structure. 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. Na2SO4•10H2O sodium sulfate decahydrate nickel(II) sulfate hexahydrate NiSO4•6H2O sodium carbonate monohydrate Na2CO3•H2O BaCl2•2H2O barium chloride dihydrate Prefixes • • • • • • • • • • 1 2 3 4 5 6 7 8 9 10 mono di tri tetra penta hexa hepta octa nona 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