Chap 3-Lesson 3 Chemical Compounds-discussion PDF

Chem. 11- Chemistry for Engineers Chapter 3 – Lesson 3 Chemical Compounds  Elements combine to form compounds Compounds have different properties from the elements that make them Atoms of different elements are held together by chemical bonds Bonds help to determine...

Chem. 11- Chemistry for Engineers Chapter 3 – Lesson 3 Chemical Compounds  Elements combine to form compounds Compounds have different properties from the elements that make them Atoms of different elements are held together by chemical bonds Bonds help to determine the properties of a compound  CHEMICAL BONDS ▪ Atoms or ions are held together in molecules or compounds by chemical bonds. ▪ The type and number of electrons in the outer electronic shells of atoms or ions are instrumental in how atoms react with each other to form stable chemical bonds.  CHEMICAL BONDS Two of the most common substance on our dining table are salt and granulated sugar NaCl C12H22O11 The properties of substances are determined in large part by the chemical bonds that hold their atoms together  Chemical Bonds All chemical reactions involve breaking of some bonds and formation of new ones which yield new products with different properties.  Properties of Compounds Depend on atoms in the compound Depend on how the atoms are arranged in the compound Example: – C and H combine to form natural gas, auto gas, waxes in candle, plastics…etc.  Properties of Compounds are different than the elements that make them H2O (water) – H and O are colorless gases at room temperature – Water is a liquid at room temperature NaCl (salt) – Na is a metallic solid – Cl is a greenish-yellow gas that is poisonous – Table salt (NaCl) is used to flavor and preserve foods  Atoms combine in predictable numbers A given compound always contains atoms of elements in a specific ratio Ammonia NH3 always has a 1:3 ratio of Nitrogen to Hydrogen www.uh.edu  Chemical Formula Chemical Formula: uses the chemical symbols to represent the atoms of the elements and their ratios in the chemical compound. H2O 2:1 ratio of H to O  Chemical Bonds hold Compounds Together Chemical bonds are the “glue” that holds the atoms of elements together Chemical bonds form when the electrons in the electron clouds interact  Lewis Bonding Theory Atoms ONLY come together to produce a more stable electron configuration. Atoms bond together by either transferring or sharing electrons. Many of atoms like to have 8 electrons in their outer shell. – Octet rule. – There are some exceptions to this rule—the key to remember is to try to get an electron configuration like a noble gas. Li and Be try to achieve the He electron arrangement.  Lewis Symbols of Atoms Uses symbol of element to represent nucleus and inner electrons. Uses dots around the symbol to represent valence electrons. – Puts one electron on each side first, then pair. Remember that elements in the same group have the same number of valence electrons; therefore, their Lewis dot symbols will look alike. Li Be B C N O: :F: :Ne:  Lewis (Electron) Dot Structures Lewis Structure: atoms represented using the element symbol and dots for valence electrons – Two shared e- equals one bond – e- not used to form a bond are called “lone pairs” or “nonbonding pairs” Cations (positive ions) have Lewis structures without valence electrons Anions (negative ions) have Lewis structures with 8 valence electrons Valence electrons  Practice to write the Lewis symbol for Arsenic As is in group 15 (5A), therefore it has 5 valence electrons. As  Using Lewis Theory to Predict Chemical Formulas Compounds Predict the formula of the compound that forms between calcium and chlorine. ∙ Cl ∙∙ ∙∙ ∙∙ Draw the Lewis dot symbols ∙ Ca∙ of the elements. ∙∙ ∙∙ ∙ Cl ∙∙ ∙∙ ∙ Cl ∙∙ ∙∙ ∙∙ Transfer all the valance electrons Ca from the metal to the nonmetal, adding more of each atom as you go, until all electrons are lost Ca2+ from the metal atoms and all nonmetal atoms have 8 electrons. CaCl2 Examples for Lewis representation of some chemical bonds F F H O H O O O H O H O F F F F O O TYPES OF CHEMICAL BONDS Ionic bonds Covalent bonds Metallic bonds The three possible types of bonds. Ionic compounds consist of a cation and an anion the formula is always the same as the empirical formula the sum of the charges on the cation and anion in each formula unit must equal zero. Lewis bonding theory is able to explain ionic bonds very well. The ionic compound NaCl  Ionic bonding Ionic substances are formed when an atom that loses electrons relatively easily react with an atom that has a high affinity for electrons. ex. metal-nonmetal compound  Ionic Bonding Ionic Bond: a total transfer of one or more electrons from one atom to another Ionic compounds form into crystals of repeating formula units Ionic bonds are extremely strong Reactions between metals and nonmetals tend to form ionic bonds A positively charged ion (CATION) is attracted to a negatively charged ion (ANION)—attraction between ions is due to their opposite charges Example: Na+ + Cl- → NaCl  Ionic Bonds Ionic bonds are formed between metals and non-metals. Ionic bonds are formed between oppositely charged atoms (ions). Ionic bonds are formed by the transfer of electrons. – One atom loses (gives away) electrons. – One atom gains (receives) electrons.  ionic compound: a compound composed of a metal + nonmetal(s) - are held together by an ionic bond (electrostatic attraction) - a network of ions, with each cation surroundedby anions, and vice versa ionic compounds have the highest melting points compared to other substances – to melt the crystal solid, all the bonds between ions have to be broken  Ionic compounds Ionic bonds form between all nearby ions of opposite charge. Ionic bonds form between non-metal and metal atoms Ionic compounds are very stable and their crystals are very strong. The shape of the crystals formed depends on the ratio of positive to negative ions and the sizes of the ions  Ionic Compounds (Salts) Composed entirely of ions Electrically neutral (Ex: Al3+ and N3- combine to form AlN) Properties: Crystal lattice (geometric pattern) High melting points (solids at room temp.) Brittle and hard Liquid (molten) state conducts electricity; solid state does not Solutions are good conductors  Chemical Bonds Ionic bonds are formed by the attraction of oppositely charged ions.  Ionic Bonds Metal to nonmetal. Metal loses electrons to form cation. Nonmetal gains electrons to form anion. The electronegativity between the metal and the nonmetal must be > than 2. Ionic bond results from + to − attraction. – Larger charge = stronger attraction. – Smaller ion = stronger attraction. Lewis theory allows us to predict the correct formulas of ionic compounds. Ions that pack as spheres in a very regular pattern form crystalline substances.  More Gains and Losses Can elements lose or gain more than one electron? The element magnesium, Mg, in Group 2 can lose two electron and element oxygen in Group 6 can gain two electrons to form stable Nobel gas configurations. The ions can come together to form a crystal structure.  Ionic Bonds Ionic bonds usually are formed by bonding between metals and nonmetals.  Ionic Bonds One cation can bond to multiple (more than one) anion  Ionic Bonds When atoms form an ionic compound, their electrons are shifted to the other atoms, but the overall number of protons and electrons of the combined atoms remains equal and unchanged. Therefore, the compound is neutral. Al3+ and N3- ---→ Al N (+3) (-3) = 0 Mg 2+ and Cl- ---→ Mg Cl2 (+2) + (-1X2) =0 Relative sizes of some ions and their parent atoms.  Structure of ionic crystals Different types of crystals are formed depending on the ionic radii and the charge of the ions involved.  Property IONIC Bond Formation Electron transferred form M to NM Type of Structure Crystal latiice Physical Stae solid Melting Point high Solubility in Water yes Electrical Conductivity Yes ( solution or liquid) Other Properties  Covalent Bonds—Sharing Some atoms are unlikely to lose or gain electrons because the number of electrons in their outer levels makes this difficult. Consider the Lewis dot structure of carbon.. C.. C+4 + 4e- The alternative is sharing electrons.  Covalent Bonds A pair of shared electrons between atoms.(prefix co- means partner) Forms between non-metal atoms Neither atom gains or loses an electron The shared electrons are attracted to both positively charged nuclei. (nucleus has a positive charge because of protons) A covalent bond is represented by a line between the two atoms  Covalent Bonds Often found between two nonmetals. Typical of molecular species. Atoms bonded together to form molecules. – Strong attraction. Atoms share pairs of electrons to attain octets. Molecules generally weakly attracted to each other. – Observed physical properties of molecular substance due to these attractions.  Covalent Bonding Electron are shared by nuclei  The Covalent Bond Shared electrons are attracted to the nuclei of both atoms. They move back and forth between the outer energy levels of each atom in the covalent bond. So, each atom has a stable outer energy level some of the time. The formation of a bond between two atoms.  Examples of Convalent Bond The neutral particle is formed when atoms share electrons is called a molecule  Covalent Bond The number of covalent bonds that an atom can form depends on the number of electrons that it has available for sharing. – Atoms of Group 16 (O,S…)can form two covalent bonds. – Atoms of group 15 (N,P…) can form three bonds – Atoms of group 14 (C, Si…)can form four bonds  Single Covalent Bonds Two atoms share one pair of electrons. – 2 electrons. One atom may have more than one single bond. F F H O H F F H O H F F  Double Covalent Bond Two atoms sharing two pairs of electrons. – 4 electrons. Shorter and stronger than single bond. O O O O O O Chemical Bonds  Chemical Bonds Covalent bonds form when atoms share 2 or more valence electrons. Covalent bond strength depends on the number of electron pairs shared by the atoms. single double < triple < bond bond bond  Chemical Bonds Give all Materials their Structure Ionic Compounds (losing/gaining e-) – Most have a crystal structure – Solid at room temperature – High melting and boiling points (takes a lot of energy to break the bond) – Hard, brittle, good conductors of electricity once the ions are separated – Dissolve easily in water  Chemical Bonds Give all Materials their Structure Covalent Compounds (sharing valence e-) – Exist as individual molecules – Chemical bonds give each molecule a specific three-dimensional shape – Molecular shape can affect properties of the compounds – Melt and boil at lower temperatures (takes less energy to break up because atoms are organized as individual molecules)  IONIC COVALENT Bond e- are transferred from e- are shared between Formation metal to nonmetal two nonmetals Type of crystal lattice true molecules Structure Physical State solid liquid or gas Melting high low Point Solubility in yes usually not Water Electrical yes Conductivity (solution or liquid) no Other Properties  Properties of Ionic and Covalent Property Compounds Ionic Covalent Compounds Compounds State at Room Crystalline solid Liquid, gas, and Temperature solid Melting Points High Low Electrical Yes No Conductivity as a liquid Solubility in water Most have high Most have low solubility solubility Conducts electricity yes Not usually when dissolved in  Bond Polarity Bonding between unlike atoms results in unequal sharing of the electrons. – One atom pulls the electrons in the bond closer to its side. – One end of the bond has larger electron density than the other. The result is bond polarity. – The end with the larger electron density gets a partial negative charge and the end that is electron deficient gets a partial positive charge. d+ H Cl d-  Bond Polarity Most bonds are a blend of ionic and covalent characteristics.  Bond Polarity Nonpolar Covalent Bond – e- are shared equally – symmetrical e- density – usually identical atoms An electron density plot for the H2 molecule shows that the shared electrons occupy a volume equally distributed over BOTH Hydrogen atoms. Electron Density for the H2 molecule  Bond Polarity Polar Covalent Bond – e- are shared unequally – asymmetrical e- density – results in partial charges (dipole) d + d -  Bond Polarity ⚫ Nonpolar ⚫ Polar ⚫ Ionic Nonpolar and polar covalent bonds  Bond Polarity Electronegativity – Attraction an atom has for a shared pair of electrons. – higher e-neg atom  d- – lower e-neg atom d+ Probability representations of the electron sharing in HF. Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms electron rich electron poor region region e- poor e- rich H F H F d+ d- Trends in electronegativity across a period and down a group  Bond Polarity Electronegativity Trend – Increases up and to the right. Nature of bonds and electronegativity Electronegativity Bond difference (∆) ∆>2 Ionic 0.4 < ∆ < 2 Polar covalent ∆ < 0.4 Covalent  Which one of these bonds would be least polar? B-C C-N B-Si Boron: 2.04 Carbon: 2.55 Nitrogen:3.04 Silicon: 1.90  Bond Polarity 3.0-3.0 4.0-2.1 3.0-0.9 = 0.0 = 1.9 = 2.1 Covalent Ionic Pure Polar 0 0.4 2.0 4.0 Electronegativity difference Polar Molecules and Electric Field COMPOUND is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds H2 H2O NH3 CH4 A diatomic molecule contains only two atoms H2, N2, O2, Br2, HCl, CO A poly molecule contains more than two atoms O3, H2O, NH3, CH4  Metallic Bonding The model of metallic bonding can be used to explain the properties of metals. The luster, malleability, ductility, and electrical and thermal conductivity are all related to the mobility of the electrons in the solid. The strength of the metallic bond varies, depending on the charge and size of the cations  Metallic Bonds Metallic bonds are metal to metal bonds formed by the attraction between positively charged metal ions and the electrons around them. – Atoms are packed tightly together to the point where outermost energy levels overlap. This allows electrons to move freely from one atom to the next making them great conductors of electricity.  Metallic bonds Metallic bond – Closely packed metal atoms - Outer energy levels of metal atoms overlap - Valence electrons can move and flow freely throughout the substance - Allows flexibility of material to bend - Allows easy flow of electricity through material IONIC COMPOUNDS vs METALS  Metals have unique bonds Metallic bond: the equal sharing of electrons in all directions so electrons move easily among the atoms of the metal Atoms can slide past one another in metallic bonds which allows for easy shaping Properties of metals depend on bonds – Good conductors of electric current – High melting point – Solid at room temperature (except Hg) – Easily shaped and pounded  METALLIC Bond e- are delocalized Formation among metal atoms Type of Structure “electron sea” Physical solid State Melting very high Point Solubility in Water no Electrical yes Conductivity (any form) Other malleable, ductile, Properties lustrous  Molecules and Molecular Compounds Molecule - combination of at least two atoms in a specific arrangement held together by chemical bonds – May be an element or a compound – H2, hydrogen gas, is an element – H2O, water, is a compound 83  Diatomic molecules: – Homonuclear (2 of the same atoms) Examples: H2, N2, O2, F2, Cl2, Br2, and I2 – Heteronuclear (2 different atoms) Examples: CO and HCl Copyright MGraw-Hill 2009 84  Polyatomic molecules: – Contain more than 2 atoms – Most molecules – May contain more than one element – Examples: ozone, O3; white phosphorus, P4; water, H2O, and methane (CH4) 85  Molecular formula - shows exact number of atoms of each element in a molecule – Subscripts indicate number of atoms of each element present in the formula. – Example: C12H22O11 86  Allotrope - one of two or more distinct forms of an element – Examples: oxygen, O2 and ozone, O3; diamond and graphite (allotropic forms of carbon) Structural formula - shows the general arrangement of atoms within the molecule. 87 Naming of Compounds Things to Consider in Naming Compounds 1. Identify if compound is ionic or covalent. 2. Ionic compound do not uses prefixes like mono, di, tri to name the subscript. 3. Different rules in naming ionic and covalent compound. Naming molecular (covalent) compounds 1. Naming Binary Molecular compounds Composed of two nonmetals Name the first element Name the second element changing ending to “-ide” Use prefixes to indicate number of atoms of each element 90 GREEK PREFEXIS 91 92 Name the following: NO2 nitrogen dioxide N2O4 dinitrogen tetraoxide Write formulas for the following: Diphosphorus pentoxide P2O5 Sulfur hexafluoride SF6 93  Common Names B2H6 diborane SiH4 silane NH3 ammonia PH3 phosphine H2O water (vs. “dihydrogen monoxide) H2S hydrogen sulfide (vs. “dihydrogen sulfide”) 94  Acid - a substance that produces hydrogen ions (H+) when dissolved in water 2. Naming Binary Acids – Many have 2 names Pure substance Aqueous solution – Example: HCl, hydrogen chloride, when dissolved in water it is called hydrochloric acid 95  Naming binary acids – Remove the “–gen” ending from hydrogen (leaving hydro–) – Change the “–ide” ending on the second element to “–ic” – Combine the two words and add the word “acid.” 96 Name the following: HBr hydrogen bromide HBr (aq) hydrobromic acid Write formulas for the following: Hydrochloric acid HCl(aq) Hydrofluoric acid HF(aq) 97  Ions and Ionic Compounds Ion - an atom or group of atoms that has a net positive or negative charge Monatomic ion - one atom with a positive or negative charge Cation - ion with a net positive charge due to the loss of one or more electrons Anion - ion with a net negative charge due the gain of one or more electrons 98 Common Monatomic Ions 99  – Anions Name the element and modify the ending to “-ide” Example: Cl-, chloride Polyatomic ions - ions that are a combination of two or more atoms – Notice similarities - number of oxygen atoms and endings for oxoanions Nitrate, NO3- and nitrite, NO2- 100. Groups IA, IIA, IIIA elements, silver (Ag), zinc (Zn) and cadmium (Cd) form only one type of ion each: – Group IA elements , +1 charge always (e.g. Li+=lithium ion) – Group IIA elements ,+2 charge always (e.g. Mg+2=magnesium ion) – – Group IIIA elements, +3 charge always (e.g. Al+3=aluminum ion) – silver ion = Ag+; zinc ion = Zn2+; cadmium ion = Cd2+  Transition elements that have more than one charge Chromium(II) Cr2+ Chromous Iron(II) Fe2+ Ferrous Chromium(III) Cr3+ Chromic Iron(III) Fe3+ Ferric 6+ Chromium(VI) Cr Chromyl Lead(II) Pb2+ Plumbous Cobalt(II) Co2+ Cobaltous Cobalt(III) Co3+ Cobaltic Lead(IV) Pb4+ Plumbic Copper(I) Cu+ Cuprous Copper(II) Cu2+ Cupric Manganese(II) Mn2+ Manganous Manganese(III) Mn3+ Manganic Manganese(IV) Mn4+ Manganyl  Nickel(II) Ni2+ Nickelous Nickel(III) Ni3+ Nickelic Tin(II) Sn2+ Stannous Tin(IV) Sn4+ Stannic Mercury(II) Hg2+ Mercuric Mercury(II) Hg22+ Mercurous PolyatomicPolyatomic cations cations Ammonium Ammonium NH4+ NH4+ Hydronium Hydronium H3O+ H3O+ Nitronium + NO2+ Nitronium NO2 Uranyl 2+ UO22+ Uranyl UO2  Naming ions – Cations from A group metals Name the element and add the word “ion” Example: Na+, sodium ion – Cations from transition metals with some exceptions Name element Indicate charge of metal with Roman numeral Add word “ion” Example: Cu2+ ,copper(II) ion 104 ANIONS: formed only by nonmetals When a nonmetal forms an ion, it is named: element stem name + -ide suffix + ion O = oxygen atom , O2– = oxide ion N = nitrogen atom , N3– = nitride ion Formal Name Formula Simple Anions Arsenide As3− Azide N3− Bromide Br− Chloride Cl− Fluoride F− Hydride H− Iodide I− Nitride N3− Oxide O2− Phosphide P3− Sulfide S2− Peroxide O22− Oxoanions – oxygen containing polyatomic anions Dichromate Cr2O72− Arsenate AsO43− Iodate IO3− Arsenite AsO33− Nitrate NO3− Borate BO33− Nitrite NO2− Bromate BrO3− Phosphate PO43− Hypobromite BrO− Hydrogen phosphate HPO42− Carbonate CO32− Dihydrogen phosphate H2PO4− Hydrogen HCO3− Bicarbonate carbonate Permanganate MnO4− Hydroxide OH− Phosphite PO33− Chlorate ClO3− Sulfate SO42− Perchlorate ClO4− Thiosulfate S2O32− Chlorite ClO2− Hydrogen sulfate HSO4− Bisulfate Hypochlorite ClO− Sulfite SO32− Chromate CrO42− Hydrogen sulfite HSO3− Bisulfite Anions from Organic Acids Acetate CH3COO− Formate HCO2− Oxalate C2O42− Hydrogen oxalate HC2O4− Bioxalate Other Anions hydrosulfide HS− Bisulfide Amide NH2− Cyanate OCN− Thiocyanate SCN− Cyanide CN− Symbol Name Symbol Name H+ Hydrogen ion H- Hydride ion Li+ Lithium ion F- Fluoride ion Na+ Sodium ion Cl- Chloride ion K+ Potassium ion Br- Bromide ion Be2+ Beryllium ion I- Iodide ion Magnesium Mg2+ ion O2- Oxide ion Ca2+ calcium ion S2- Sulfide ion Ba2+ barium ion N3- Nitride ion Zn2+ zinc ion P3- Phosphide ion Formula Name Formula Name NO3- nitrate CO32- carbonate NO2- nitrite SO42- sulfate CN- cyanide SO32- sulfite MnO4- permanganate PO43- phosphate OH- hydroxide PO33- phosphite O22- peroxide ClO4- perchlorate HCO3- hydrogen carbonate ClO3- chlorate HSO4- hydrogen sulfate ClO2- chlorite HSO3- hydrogen sulfite ClO- hypochlorite HPO42- hydrogen phosphate CrO42- chromate H2PO4- dihydrogen phosphate C2H3O- 2 acetate  (Stock Symbol (Stock system) Common Symbol system) Common Cu+ copper(I) cuprous Hg22+ mercury(I) mercurous Cu2+ copper(II) cupric Hg2+ mercury(II) mercuric Fe2+ iron(II) ferrous Pb2+ lead(II) plumbous Fe3+ iron(III) ferric Pb4+ lead(IV) plumbic Sn2+ tin(II) stannous Co2+ cobalt(II) cobaltous Sn4+ tin(IV) stannic Co3+ cobalt(III) cobaltic Cr2+ chromium(II) chromous Ni2+ nickel(II) nickelous Cr3+ chromium(III) chromic Ni4+ nickel(IV) nickelic Mn2+ manganese(II) manganous Au+ gold(I) aurous Mn3+ manganese(III) manganic Au3+ gold(III) auric  Polyatomic ions Table Just have to memorize NH4+ ammonium ion CO32- carbonate ion CN- cyanide ion HCO3- hydrogen (or bi) carbonate ion OH- hydroxide  NO3- nitrate ion NO2- nitrite ion PO43- phosphate ion SO42- sulfate ion HSO4- hydrogen sulfate ion SO32- sulfite ion CH3COO- (C2H3O2-) acetate ion  CHEMICAL FORMULA of IONIC COMPOUNDS Ionic compounds - represented by empirical formulas – Compound formed is electrically neutral – Sum of the charges on the cation(s) and anion(s) in each formula unit must be zero – Examples: Al3+ and O2- Al2O3 Ca2+ and PO43- Ca3(PO4)2 113 Cross –over Rule is used to determine the number of atoms ( as subscript) of each element in the chemical formula 114 Formation of an Ionic Compound 115 Write empirical formulas for Aluminum and bromide AlBr3 barium and phosphate Ba3(PO4)2 Magnesium and nitrate Mg(NO3)2 Ammonium and sulfate (NH4)2SO4 116  Ionic Bonding Nomenclature To name Binary Ionic Compounds: 2 elements—one METAL and one NON-METAL Cation is always written first [Metal] Cation name stays the same Anion is written second [Non-metal] Change the non-metal’s ending to “-ide”. NO PREFIXES ARE USED FOR IONIC COMPOUND NAMING  Naming ionic compounds – Name the cation – Name the anion – Check the name of cation NOTE : If it is a A group metal you are finished If it is a transition metal, with some exceptions, add the appropriate Roman numeral to indicate the positive ionic charge 118 Rules in Naming and Writing Chemical Formula of Ionic Compounds 1. If both ions have charges that are exactly opposite (+1 & -1, 2+ & -2, etc.),formula contains one of each (This also applies for polyatomic ions.) A. Representative metal + Non-metal Monoatomic cation + Monoatomic anion Naming : name the metal first followed by the non- metal ending in ide K + Cl → K+ + Cl- → KCl potassium chloride Mg + O → Mg 2+ + O2- → MgO magnesium oxide  B. Representative Metal + Non-metal Monoatomic cation + polyatomic anion Naming: name the metal first followed by the polyatomic anion ending in ate or ite Al 3+ + PO43- → AlPO4 aluminum phosphate K+ + NO2- → KNO2 potassium nitrate C. Polyatomic cation + Polyatomic anion Naming: name the polyatomic anion first followed by the polyatomic anion ending in ate or ite NH4 + + NO3- → NH4 NO3 ammonium nitrate  Practice: Polyatomic Ions Compound Name Oxidation #s Chemical Formula Calcium phosphate Ca2+ PO43- Ca3(PO4)2 Sodium hydroxide Na1+ OH1- NaOH Ammonium sulfate NH41+ SO42- (NH4)2SO4 D. transition metal + non-metal Naming ionic compounds with more than charge ( transition element) Many metals form more than one compound with some anions. For these, roman numerals are used in the name to indicate the charge on the metal. Cu1+ + O2- = Cu2O copper(I) oxide copper(I) oxide cuprous oxide Cu2+ + O2- = CuO copper(II) oxide copper(II) oxide 2. For monatomic ions with different charges, use the crossover rule to have a neutral chemical formula Make the negative charge the subscript of cation, and make positive charge the subscript of anion Na+x O-2y = Na2O sodium oxide Ba2+ + Cl- -→ BaCl2 barium chloride Fe 3+ + Cl- -→ FeCl3 iron (III) chloride “Swap & Drop ( Cross-over) ” Method Given the name of an Ionic Compound, you can determine the chemical formula using the “swap and drop” method: 1. Write the symbols for each ion. 2. Determine the oxidation number of each ion. 3. Swap and Drop 4. Reduce (if necessary). 5. Rewrite Transition Metals--Ionic Compounds Transition metals are cations that have variable charges that makes them hard to name. – We use Roman numerals to indicate the charge of a transition metal. Example: – copper (II) oxide – charge of copper for this compound is +2 – titanium (IV) sulfide – charge of titanium for this compound is +4  Transition Metal Ionic Compounds To go from formula to name you need to determine the Roman numeral for your transition metal. 1. If there are no subscripts, simply give the transition metal the equal and opposite charge to the nonmetal. 2. Now use normal ionic bonding rules putting your new number in Roman numerals to the right of your transition metal ONLY.  Transition Metal Ionic Compounds To go from formula to name you need to determine the Roman numeral for your transition metal: 1. If there are subscripts present use the reverse “Swap and Drop.” 2. Now use normal ionic bonding rules putting your new number in Roman numerals to the right of your transition metal ONLY. 4. For polyatomic ions, where ions have different charges, also use the crossover rule —Express polyatomic ion with subscripts and parentheses Ba 2+ + NO3- → Ba (NO3)2 barium nitrate Fe 2+ + PO43- → Fe3(PO4)2 iron (II) phosphate  Polyatomic Ions To go from name to formula: 1. Write the symbols for each ion. 2. Determine the oxidation number of O 2- each ion. 3. Swap and Drop 4. Reduce (if necessary). 5. Put parentheses around the polyatomic (NH4)2O ion if receives a ** Remember charges CANCEL subscript greater than out each other!! one. 6. Rewrite EXCEPTION FOR CROSSOVER RULE: Ions with +4 and 2– charges Ionic compound formulas must reflect lowest ratio of elements Instead of Pb2O4, it should be PbO2 PbO2 lead (IV) oxide note : name still reflects the oxidation number of the transition metal Write names for the following: KMnO4 Sr3(PO4)2 Co(NO3)2 FeSO4 131 Write names for the following: KMnO4 potassium permanganate Sr3(PO4)2 strontium phosphate Co(NO3)2 cobalt(II) nitrate FeSO4 iron(II) sulfate 132 OXOACIDS ( Covalent Compounds) – When writing formulas, add the number of H+ ions necessary to balance the corresponding oxoanion’s negative charge – Naming formulas If the anion ends in “-ite” the acid ends with “-ous” acid If the anion ends in “-ate” the acid ends in “-ic” acid 133 Oxoacids with 4 forms are exhibited by group 7A ( halogens) oxoanions change ate to ic add acid change ite to ous add acid per – ic acid hypo- ous acid Oxoanion oxoacid Perchlorate ClO4− HClO4 Perchloric acid Chlorate ClO3− HClO3 Chloric acid Chlorite ClO2− HClO2 Chlorous acid Hypochlorite ClO− HClO Hypochlorous acid Oxoanion oxoacid Perbromate BrO4− HBrO4 Perbromic acid Iodate IO3− HIO3 Iodic acid Hypochlorite ClO− HClO Hypochlorous acid Name the following: H2SO3 HClO H3PO4 Write formulas for the following: Perchloric acid Nitric acid 136 Name the following: H2SO3 sulfurous acid HClO hypochlorous acid H3PO4 phosphoric acid Write formulas for the following: Perchloric acid HClO4 Nitric acid HNO3 137 Hydrates - compounds that have a specific number of water molecules within their solid structure – Hydrated compounds may be heated to remove the water forming an anhydrous compound – Name the compound and add the word hydrate. Indicate the number of water molecules with a prefix on hydrate. Example: CuSO4 · 5 H2O –Copper (II) sulfate pentahydrate 138 139 END of PRESENTATION