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Chemical Formulas and Equations Do Now • What is a mole in chemistry? • What was Avogadro's number? What did it represent? • What is molar mass? Do Now • What is a mole in chemistry? • Remember: Mole (mol) = the SI unit for amount • What was Avogadro's number? What did it represent? • Remember: Avogadro's number = the number of particles in a mole = 6.02 x 1023 • What is molar mass? • Remember: Molar mass = mass in grams of one mole of an element or compound Remember • You can convert from moles to particles, or moles to mass, and vice versa. • Ex: How many particles do you have if you have 2.5 mol of sulfur? • Ex: How many grams of carbon do you have if you have 2.44 x 1022 atoms? * You must use conversion factors to convert from number of atoms to moles and then to grams. Formula Mass • The molar mass is numerically equal to the atomic mass of monatomic elements and the formula mass of compounds and diatomic elements. • Ex: Find the formula mass of KBr • Ex: Find the formula mass of H2O Chemical Formula • Chemical formulas indicate the relative number of atoms of each kind in a chemical compound. • Identify the number of atoms in each of the following compounds: 1. KCl= _____________________ 2. C6H12O6=__________________ 3. NH3=__________________ Chemical Formulas • Formulas for covalent compounds show both the elements and the number of atoms of each element in a molecule. • Formulas for ionic compounds do not show numbers of atoms, but show the simplest ratio of cations and anions. Chemical Formulas • The meaning of formulas do not change when polyatomic ions are involved. • Polyatomic ions = a group of covalently bond elements that behave as a single ion Finding the number of atoms in a compound with polyatomic ions Compound 1. MgCO3 2. Al(ClO4)3 What does it look like expanded? Numbers of atoms Compound 3. Zn3(PO4)2 4. (NH4)2S What does it look like expanded? Numbers of atoms Molar Mass • Formulas can be used to calculate molar mass • Ex: ZnCl2 • Ex: ZnSO4 Do Now • Do you see a pattern in the following formulas? NH4NO2 NH2O Law of Definite Proportions • Law of Definite Proportions states that every pure substance always contains the same elements combined in the same proportions by weight. • For example: H2O, will always have the same percent by weight (11.2% H and 88.8% O) Percentage Composition • Percentage composition = the percentage by mass of each element in a compound • Percentage composition helps verify a substance’s identity and can be used to compare the ratio of masses contributed by the elements in two substances • Ex: Fe2O3 FeO 69.9% Fe 77.7% Fe 30/1% O 22.3% O Review Formulas • Molecular Formula = represents the number and kind of atoms in a molecule (not necessarily the smallest whole number ratio). • Structural Formula = indicates twodimensional arrangement of the bonds and lone pairs of electrons in a molecule. Empirical Formulas • Empirical formula = a chemical formula that shows the composition of a compound in terms of the relative numbers and kinds of atoms in the simplest ratio. Empirical Formula for Molecules (have covalent bonds) • The chemical formula that represents the simplest (lowest) atomic ratio in which elements can combine. • Formulas for molecular compounds are NOT NECESSARILY the empirical formula Molecular Formula Empirical Formula C2H6 C6H12O6 C4H8 C3H8 C6H10 C4H6 C2H4 C5H10 Empirical Formula for Ionic Compounds (Always written in empirical form) • The chemical formula that represents the simplest (lowest) atomic ratio in which elements can combine. • Formulas for ionic compounds are the empirical formula because they are ALWAYS written in the expressed as the lowest possible ratios. Empirical Form Ca2O2 Fe2O2 Pb2O4 Mg2O2 Empirical Formula for Ionic Compounds with Polyatomic Ions Empirical Form Fe2(CO3)2 Fe2(HSO4)2 Zn2(SO4)2 Pb2(SO3)2 Empirical or Molecular Formula C6H12O6 LiNO3 H2O2 C11H22O11 If written in molecular formula what is the empirical formula? Type of Bond Finding Empirical Formulas • You can find the empirical formula from percentage composition! – Step 1: Convert to mass in grams (assume you have 100g of the given substance) – Step 2: Convert from grams to moles (using the molar mass conversion factor) – Step 3: Reduce the molar ratio to the simplest whole-number ratio by dividing by the smaller amount – Step 4: Round to whole numbers and insert subscripts Finding Empirical Formulas • Ex: A given liquid has 60.0% C, 13.4% H, and 26.6% O by mass. Calculate the empirical formula. – Step 1: Convert to mass in grams (assume you have 100g of the given substance) 60.0% C x 100g = 60.0g C 13.4% H x 100g = 13.4g H 26.6% O x 100g = 26.6g O Finding Empirical Formulas – Step 2: Convert from grams to moles (using the molar mass conversion factor) 60.0g C x 1mol = 5.00mol C 12.01g C 13.4g H x 1mol = 13.3mol H 1.01g H 26.6g O x 1mol = 1.66mol O 16.00g O C5H13.3O1.66 ? Finding Empirical Formulas – Step 3: Reduce the molar ratio to the simplest whole-number ratio by dividing by the smaller amount C5H13.3O1.66 ? 5.00mol C / 1.66 = 3.01mol C 13.3mol H / 1.66 = 8.01mol H 1.66mol O / 1.66 = 1.00mol O – Step 4: Round to whole numbers and insert subscripts C3H8O Practice • What is the empirical formula of a compound that is 78.6% B and 21.4% H? Practice • What is the empirical formula of a compound containing 32.38% Na, 22.65% S, and 44.99% O? Practice • What is the empirical formula of a compound containing 26.56% potassium, 35.41% chromium, and the remainder is oxygen? Practice • In a 10.150g sample, 4.433g are phosphorous and the rest is oxygen. What is the empirical formula for this compound? Practice • A compound contains 0.606g nitrogen and 1.390g oxygen. What is the empirical formula of the compound? Do Now • What was the difference between the empirical formula and molecular formula? Do Now • What was the difference between the empirical formula and molecular formula? • Empirical formula = smallest possible whole number ratio of elements • Molecular formula = the actual formula of a molecule Molecular Formulas • For ionic compounds, the molecular formula is the same as the empirical formula. • For molecular compounds, the molecular formula is a whole number multiple of the empirical formula. *Both formulas are just different ways of representing the composition of the same molecule. Molar Mass • In order to determine the molecular formula, you must know the molecular mass! • The molar mass of a compound is equal to the molar mass of the empirical formula time a whole number, n. n(empirical formula) = molecular formula Molecular Formula Examples • Formaldehyde, acetic acid, and glucose each have the same empirical formula, CH2O – For formaldehyde, n = 1 – For acetic acid, n = 2 – For glucose, n = 6 Determining Molecular Formula • Step 1: Find the molar mass of the empirical formula using the molar masses of the elements from the periodic table • Step 2: Solve for n, the factor multiplying the empirical formula to get the molecular formula n = experimental molar mass of compound molar mass of the empirical formula • Step 3: Multiply the empirical formula by this factor to get the molecular formula Determining Molecular Formula • The empirical formula of a compound is P2O5. The experimental molar mass is 284 g/mol. Determine the molecular formula. – Step 1: Find the molar mass of the empirical formula using the molar masses of the elements from the periodic table 2 x molar mass of P = 2(30.97) = 61.94g/mol 5 x molar mass of O = 5(16.00) = 80.00g/mol molar mass of P2O5 = 141.94g/mol Determining Molecular Formula • The empirical formula of a compound is P2O5. The experimental molar mass is 284 g/mol. Determine the molecular formula. – Step 2: Solve for n, the factor multiplying the empirical formula to get the molecular formula n = experimental molar mass molar mass of empirical formula n = 284 g/mol 141.94 g/mol n=2 Determining Molecular Formula • The empirical formula of a compound is P2O5. The experimental molar mass is 284 g/mol. Determine the molecular formula. – Step 3: Multiply the empirical formula by this factor to get the molecular formula n(empirical formula) = 2(P2O5) = P4O10 Determining Molecular Formula • You can verify your answer by finding the molar mass of the molecular formula and compare it to the experimental molar mass 4 x molar mass of P = 4(30.97g/mol) = 123.88g/mol 10 x molar mass of O = 10(16.00g/mol) = 160.0g/mol molar mass of P4O10 = 283.88g/mol Practice • What is the molecular formula of a compound with the empirical formula BH3 and molecular mass of 28g? Practice • What is the molecular formula of a compound with a molecular mass of 34g that consists of 0.44g H and 6.92g O? Practice • What is the empirical formula of a compound that contains 65.5% carbon, 5.5% hydrogen, and 29.0% oxygen? What is the molecular formula if the molecular mass is 110g? Do Now • What does percentage mean? • What does a percentage represent? Review • Chemical formulas allow scientists to calculate a number of characteristics values for a given compound. • Chemical formulas represent the number and kind of atoms in a molecule. • If you know the chemical formula, then you can calculate the percentage composition. Review • Formula mass = the sum of the average atomic masses of all the atoms represented in the formula • Formula mass is numerically equal to the molar mass or gram formula mass • Ex: Find the formula mass of KClO3 • Ex: What is the molar mass of Ba(NO3)2? Percent Composition • Percentage composition = the percentage by mass of each element in a compound Finding Percent Composition • From the subscripts, you can determine the mass contributed by each element and add these to get the molar mass. • Divide the mass of each element by the molar mass. • Multiply by 100 to find the percentage composition of that element Finding Percent Composition • CO2 1mol x 12.01 g/mol = 12.01 g C +2mol x 16.00 g/mol = 32.00 g O mass of 1mol CO2 = 44.01 g Finding Percent Composition % C = 12.01 g C x 100 = 27.29% 44.01 g CO2 % O = 32.00 g O x 100 = 72.71% 44.01 g CO2 Finding Percent Composition • CO 1mol x 12.01 g/mol = 12.01 g C +1mol x 16.00 g/mol = 16.00 g O mass of 1mol CO = 28.01 g Finding Percent Composition % C = 12.01 g C x 100 = 42.88% 28.01 g CO % O = 16.00 g O x 100 = 57.71% 28.01 g CO Percent Composition • Percentage composition helps verify a substance’s identity and can be used to compare the ratio of masses contributed by the elements in two substances • Ex: Fe2O3 FeO 69.9% Fe 77.7% Fe 30/1% O 22.3% O Practice • Calculate the percentage composition of Cu2S, a copper ore called chalcocite. Practice • Calculate the percent of both elements in sulfur dioxide Practice • Calculate the percentage composition of ammonium nitrate, NH4NO3 Hydrates • Hydrates = salts that have crystallized from water solution • In the process of crystallization, the water molecules bind to the salt to form hydrates. • Hydrates are represented as follows: Na2CO3 ● 10H2O (sodium carbonate decahydrate) Practice • Calculate the molar mass of Na2CO3 ● 10H2O Practice • Calculate the percent composition of water in Na2CO3 ● 10H2O Do Now • Take out Reference Tables and open to Table E – QUIZ ON POLYATOMIC IONS ON WEDNESDAY • Turn to Table S – Where is the oxidation number noted? • How many aluminum atoms combine with how many sulfur atoms? How do we figure this out? OXIDATION NUMBER • Oxidation number = the charge an atom would acquire if all its bonds were treated as ionic bonds. • To determine how many atoms combine with one another in a compound we must determine each element’s OXIDATION NUMBER. Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions H+ Li+ Na+ K+ 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 Group 14: Lose of 4 electrons or gain of 4 electrons? Neither! Group 14 elements rarely form ions. 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 F1- Fluoride Br1- Bromide Cl1-Chloride I1- Iodide Group 17: Gains 1 electron to form 1- ions Predicting Ionic Charges Group 18: Stable Noble gases do not form ions! Predicting Ionic Charges Groups 3 - 11: Many transition elements have multiple oxidation states. Iron(II) = Fe2+ Iron(III) = Fe3+ Predicting Ionic Charges Groups 3 - 11: Some metals have only one possible oxidation state. Silver = Ag+ Zinc = Zn2+ Rules for assigning Oxidation Numbers 1. The atoms in a pure element have an oxidation number of zero. 2. Alkali metals always have an oxidation number of +1; alkaline earth metals always have an oxidation number of +2. 3. Fluorine always has an oxidation number of -1. 4. Oxygen has an oxidation number of -2 in almost all compounds. Exceptions are in compounds with a halogen, when it has an oxidation number of +2, and in peroxides (H2O2), when it has an oxidation number of -1. 5. Hydrogen has an oxidation number of +1 in almost all compounds except when combined with a metal when it has an oxidation number of -1. 6. The sum of all the oxidation numbers in a neutral compound is zero. 7. The sum of all the oxidation numbers in a polyatomic ion is equal to the charge of the ion. ASSIGNING OXIDATION NUMBERS • A compound has a total charge of ZERO so set your equation equal to ZERO. • Assign the variable X to your unknown oxidation number. • Given a compound, find the oxidation number of every element you know for certain. Then solve for others using algebra. • KMnO4 • CaCO3 Do you notice a pattern with the elements in these compounds? H2+1 O1-2 Mg1+2F2-1 Al2+3S3-2 Cu3+1P1-3 Ca1+2Cl2-1 Fe2+3O3-2 CRISS-CROSS METHOD to determine the chemical formula 1. Write the symbols for the elements side by side. 2. Write the oxidation states of each element to the top right of the symbol. When the nonmetal is combines with a metal, the oxidation state will always be the first number (the negative one) in the list of oxidation states. 3. Criss cross the charges DOWN and use the absolute values (-2 becomes 2). 4. Check to make sure the subscripts are the lowest ratio. Practice Criss-Cross Method 1. Na and S 2. K and P 3. Al and S 4. Mg and Br 5. Al and O What about Polyatomic Ions? • Write down the only cations on Table E: (positively charges ions) • Write down the only polyatomic ions that end in –ide. Polyatomic Ions Chart (Table E) Find the charge of the polyatomic ion using Table E Put parenthesis around the polyatomic ion. 1. PO4 7. HSO4 2. CO3 8. ClO4 3. SO3 9. CN 4. NH4 10. OH 5. ClO 11. S2O3 6. ClO2 12. SCN Polyatomic Ions • If you see a group of atoms together with a charge it is a polyatomic ion from Table E. • Put the polyatomic ion within Parenthesis. • Find the charge of the polyatomic ion. • Use the criss-cross method to determine subscripts. Na +1 CO3-2 How many atoms are present in the compound? NH4 +1 S-2 How many atoms are present in the compound? Rules for writing the formula of the compound composed of ions 1. Place all Polyatomic Ions in Parenthesis (Table E) 2. Determine all oxidation numbers of elements and polyatomic ions 3. Use Criss-Cross Method 4. Reduce to Empirical Form Practice 1. NH4 S 2. Na NO3 3. Cu Br 4. Al SO4 5. Fe CO3 6. Pb PO4 7. Ag ClO 8. Ca F 9. NH4 SO3 10. Cu OH 11. Ni I 12. Zn SO4 13. Pb ClO2 14. H I 15. Fe HSO4 16. Cu CO3 17. NH4 O 18. Ag S 19. Al ClO4 Do Now • What is the name of H2O? • What is the name of NaCl? • What is the name of NH4? Naming Compounds • Many chemical compounds have common names. • H2O = water • NaCl = table salt • NH4 = ammonia Review • Chemical formula = the relative number of atoms of each kind in a chemical compound • Molecular formula = the number and kind of atoms in a molecule (not necessarily the smallest whole number ratio) NAMING IONIC COMPOUNDS • An Ionic compound can quickly be determined if a METAL is bonded to a NONMETAL. • Naming Ionic compounds with metals that have only ONE oxidation state is fairly simple. Of the following metals listed below, check off all the elements that have more than one oxidation state: a. Mn___ b. Zn ___ c. K ____ d. Pb___ e. Mg___ f. Au ___ g. Ag ___ h. Ga__ i. j. Li ____ k. U ____ l. Cr ___ Sn __ Naming Ionic Compounds that contain metals with 1 oxidation state • Binary Ionic Compounds = ionic compounds with only 2 different elements. • Name the metal and end the nonmetal in –ide. • For example: CaBr2 = _______________ Naming Binary Ionic Compounds • Naming salts is very easy, because they are binary ionic compounds (made up of two elements). – The cation is named by borrowing the name of the element. – The anion named by combining the name of the element with an –ide ending. • The name of compound is made up of both the cation and anion name – Ex: NaCl = sodium chloride – Ex: ZnS = zinc sulfide – Ex: K2O = potassium oxide – Ex: Mg3N2 = magnesium nitride – Ex: Al2S3 = aluminum sulfide Name the following Binary compounds: 1. MgO= ________________________ 2. CaCl2 = _______________________ 3. AlBr3 = _______________________ 4. Ag3N = _______________________ 5. Al2O3 = ______________________ 6. LiI = _________________________ 7. BaF2 = _______________________ 8. Zn2C = _______________________ 9. Ba3N2 = ______________________ 10.CdO = _______________________ 11. Ga2S3 = _______________________ 12. K3N = ________________________ 13. SrO = ________________________ What about Polyatomic ions? (Table E) • These are ions consisting of more than one atom. The names of polyatomic ions end in –ate or –ite, except for two (ammonium and cyanide) Polyatomic ion chart (Table E) Ternary Ionic Compounds • • • • Ionic Compounds with 3 different elements. They usually contain Polyatomic Ions (Table E) Name the following Polyatomic Ions: a. NO3- _________ d. SO42- ___________ • b. ClO2- _________ e. SO32- __________ • c. CO32- _________ f. SCN- ____________ Naming Ternary Compounds • Name the metal and then name the polyatomic ion (if it has a negative oxidation number). • For example: KNO3 = ________________ Naming Ternary Compounds • Name the polyatomic ion (if it has a positive oxidation number) and then name the nonmetal. • For example: NH4Cl = ________________ Naming Ternary Compounds • If there are two polyatomic ions, name the positive polyatomic ion first and then name the negative polyatomic ion. • For example: NH4NO3 = ______________ • DO NOT CHANGE THE ENDINGS OF POLYATOMIC IONS! Name the following Ternary compounds: 1. NaC2H3O2 _____________________ 2. AgHCO3 _______________________ 3. LiNO2 _________________________ 4. Ga2(S2O3)3 ____________________ 5. Ca3(PO4)2 _____________________ 6. ZnSO3 __________________________ 7. KClO3 ___________________________ 8. Al(OH)3 _________________________ 9. RbSCN __________________________ 10. SrCO3 _________________________ Naming Ionic Compounds with Metals with Multiple oxidation states (multiple charges): Using the STOCK SYSTEM 1. Determine the oxidation state of the metal in the compound. 2. Name the metal, put the oxidation state in ROMAN NUMERALS in parenthesis and end the nonmetal in –ide. Review: Find the Formula Criss-Cross Pb+4 O-2 Cu+2 (SO4) -2 Sn+2(CO3) -2 Find Empirical formula Work Backwards: Start with the Empirical Formula to determine the Oxidation state of a Metal with Multiple Oxidation States Empirical Formula Fe1O1 Write in the oxidation number for the nonmetal or polyatomic ion you are sure of and criss-cross. Fe O Fe1(SO4)1 Fe (SO4) Cu1(SO4)1 Cu (SO4) Non-reduced form with Oxidation States Work Backwards: Start with the Empirical Formula to determine the Oxidation state of a Metal with Multiple Oxidation States Empirical Formula Write in the oxidation number for the nonmetal or polyatomic ion you are sure of and criss-cross. Sn1(SO3)1 Sn (SO3) Mn1(SO4)2 Mn (SO4) Cr1(PO4)2 Cr (PO4) Non-reduced form with Oxidation States Name the following compound using the Stock System: 1. Fe O ____________________________ 2. Fe Cl2 ___________________________ 3. Cu SO4 __________________________ 4. Pb Cl2 ___________________________ 5. Pb O2 ___________________________ 6. Cu3(PO4)2 ________________________ 7. Cu2 S ___________________________ 8. Fe2(CrO4)3 ______________________ 19. Sn CO3 _________________________ 10. Sn F4 ___________________________ Name each of the following compounds, use Roman Numerals only when necessary. Put a check next to every compound that begins with a metal with more than 1 oxidation state. Put parenthesis around all the polyatomic ions. 1. NH4 Cl _________________________ 2. Pb SO4 _________________________ 3. Co Cl3 __________________________ 4. Ba (NO3)2 ______________________ 5. Co2 (SO3)3 _____________________ 6. KH ____________________________ 7. NH4 F _________________________ 8. K2Cr2O7 _______________________ 9. Cu S __________________________ 10. Cu ClO2 ______________________ 11. Ag NO3 _______________________ 12. Fe Cl3 ________________________ 13. Cr F2 _________________________ 14. Na Cl _________________________ 15. Fe PO4 ______________________ 16. Li F _________________________ 17. Fe F3 _______________________ 18. Al (OH)3 _____________________ 19. Mg I2 ________________________ 20. Fe Cl3 _______________________ Do Now • What types of bonds are in molecular compounds? _____________________ • How can we tell if a formula has a covalent bond? ___________________________ Review Molecular Compounds • Molecule = a neutral compound held together by covalent bonds • Molecules may consist of identical atoms bonded together (O2) or different atoms bonded together (H2O) Naming Covalent Compounds • Covalent compounds are named in a similar way to ionic compounds – The first element in the formula is usually written first in the name – The second element has an –ide ending • Ex: SO2 = sulfur oxide – However, this is not completely correct… Naming Covalent Compounds • Since multiple covalent compounds can be made from the same elements, the name must distinguish them as different. – Prefixes are used to indicate the number of atoms of each element in the molecule. • Ex: SO2 = sulfur dioxide • Ex: SO3 = sulfur trioxide Naming Covalent Compounds • Most nonmetals have more than 1 oxidation state, therefore you can use the Stock System (Roman Numerals) or the Prefix System. • The Prefix System includes: MonoDiTriTetraPenta- HexaHectaOctaNonaDeca- Naming Molecular Compounds • According to the number of atoms of each element, state the prefix for the number of each atom before the name of the element and end the nonmetal in –ide. • For example: P2O5 = ______________ CO2 = ______________ Naming Molecular Compounds • Name the first element. Use a prefix ONLY if there is more than one. • Name the second element. ALWAYS use a prefix. Change the ending to –ide. • The prefix mono only needs to be used for the second half of the compound NOT the first element. Name the following covalent compounds: 1. CI4 ________________________________ 2. PCl5 ______________________________ 3. SI6 _______________________________ 4. P2S6 ______________________________ 5. N3O4 ______________________________ 6. SO2 ____________________________ 7. N2O4 ___________________________ 8. CO _____________________________ 9. NF3 ____________________________ 10. ICl5 ___________________________ 11. H2S3 __________________________ 12. N2O3 _________________________ 13. ClF7 __________________________ 14. SO3 ___________________________ 15. NI5 ___________________________ 16. BN2 __________________________ 17. P2O5 _________________________ 18. IF7 ___________________________ Do Now • What is the chemical formula for dihydrogen oxide? • What is the chemical formula for acetate? Writing Formulas • The chemical formula of a compound can be determined from the chemical name! General Steps to Determining Chemical Formulas: • 1. Determine what elements and/or polyatomic ions are in the compound. • 2. Write the symbols for each substance with oxidation states of each substance to the top right of the symbol. – If there is a metal with more than one oxidation state, it will be indicated as roman numerals in parenthesis after the element. – If there is a nonmetal with more than one oxidation state, use the negative number. General Steps to Determining Chemical Formulas (continued): • 3. All oxidation states in a neutral compound add up to ZERO, so figure out how many of each substance you need to make all oxidation states add up to ZERO. • 4. Put these numbers as subscripts and write the chemical formula. • 5. If you have more than one polyatomic ion, make sure you put the symbols in parenthesis. • 6. Check to make sure the subscripts are the lowest ratio. Polyatomic Ions (Table E) Writing Formulas for Ionic Compounds • All ionic compounds must be in empirical form. (Reduced Form) • If the compound ends in –ide, most likely it is a binary compound. Except for cyanide and hydroxide. • If the substance ends in –ate or –ite it contains a polyatomic ion. Put the polyatomic ion in parenthesis with oxidation numbers indicated before you criss-cross. Put the oxidation states on top and criss-cross the numbers. • For example: Aluminum Sulfide = Write the formulas for each of the following compounds: (Make sure you reduce it to empirical form) 1. Aluminum Chloride ___________________ 2. Silver Phosphate _____________________ 3. Lithium Hydride ______________________ 4. Magnesium Acetate __________________ 5. Potassium Sulfite _____________________ 6. Zinc Thiosulfate ____________________ 7. Strontium Nitride ____________________ 8. Calcium Oxide _______________________ 9. Gallium Oxalate ______________________ 10. Ammonium Hydroxide ________________ Writing formulas for compounds with metals with more than one oxidations state: (Roman Numerals will be given) 1. If there are Roman Numerals with the name, the Roman Numeral is the charge of the metals. For example: Iron (II) Oxide= Fe+2O-2=Fe2O2=FeO 2. If the compound ends in –ate or –ite, most likely you should look on the polyatomic ion chart. 3. Write the metal, then look up the polyatomic ion and place it in parenthesis. Put the charges on top and criss-cross. Reduce if necessary. For example: Zinc Carbonate= Gold (III) Thiocyanate = For each compound listed below, write the correct formula using the stock system. 1. Iron (II) Chloride _____________________ 2. Lead (IV) Phosphide __________________ 3. Tin (II) Oxide ________________________ 4. Copper (I) Iodide _____________________ 5. Nickel (III) Sulfide ____________________ 6. Cobalt (II) Thiocyanate ______________ 7. Manganese (IV) Oxide _________________ 8. Titanium (IV) Chromate _________________ 9. Iron (III) Sulfate _______________________ 10. Lead (II) Nitrate ______________________ 11. Tin (IV) Carbonate ______________ 12. Copper (II) Acetate _____________ Writing Formulas for Covalent Compounds • Sometimes the name will have prefixes (only when the two elements in the compound are nonmetals). Simply use the prefixes to figure out the formula! For each compound listed below, write the correct formula using the prefix system. 1. Diphosphorous pentoxide _____________________ 2. Silicon tetrafluoride ________________________ 3. Dihydrogen monoxide ________________________ 4. Tetraphosphorous trisulfide __________________ Naming Ionic Compounds using the Formula Determining subscripts of elements or polyatomic in ionic compounds Use Criss-cross method Example: Naming ionic compounds with metals with only 1 oxidation state Binary Ionic Compounds Ionic compounds with only 2 different elements. For example: CaBr2 = Calcium Bromide Name the metal and end the nonmetal in –ide. Ternary Ionic Compounds Ionic Compounds with 3 different elements. •Name the metal and then name the polyatomic ion. Naming ionic compounds with metals with multiple oxidation state 1. Determine the oxidation state of the metal in the compound. 2. Name the metal, put the oxidation state in ROMAN NUMERALS in parenthesis and end the nonmetal in –ide. For example: KNO3 = Potassium Nitrate Example: FeO= Iron (II) Oxide Writing the Formula of Ionic Formulas using the name of the Ionic Compound Formula Writing for Ionic Compounds Writing formulas for compounds with metals with more than one oxidations state: (Roman Numerals will be given) All ionic compounds must be in empirical form. (Reduced Form) If the compound ends in –ide, most likely it is a binary compound. Except for cyanide and hydroxide. If the substance ends in –ate or –ite it contains a polyatomic ion. Put the polyatomic ion in parenthesis before you criss-cross. Put the oxidation states on top and crisscross the numbers. If there are Roman Numerals with the name, the Roman Numeral is the charge of the metals. 2. If the compound ends in –ate or –ite, most likely you should look on the polyatomic ion chart. 3. Write the metal, then look up the polyatomic ion and place it in parenthesis. Put the charges on top and criss cross. Reduce if necessary. For example: Zinc Carbonate= Zn+2(CO3) -2 = Zn2(CO3)2 =Zn(CO3) Gold (III) Thiocyanate = Au+3 (SCN) Au(SCN)3 -1 = For example: Iron (II) Oxide= Fe+2O-2=Fe2O2=FeO THE END Naming Compounds • The POLYATOMIC ION CHART (TABLE E) • These are ions consisting of more than one atom. These names of polyatomic ions ens in –ate or –ite, except for two ( ammonium and cyanide). Naming Ionic Compounds • 1. Cation first, then anion • 2. Monatomic cation = name of the element • Ca2+ = calcium ion • 3. Monatomic anion = root + -ide • Cl = chloride Naming Ionic Compounds (continued) Metals with multiple oxidation states • - some metal forms more than one cation • - use Roman numeral in name • PbCl2 • Pb2+ is cation How to name compounds: • 1. Determine what the elements are in the compound. – A. If the compound is made of a metal and a nonmetal, name as follows: • 1. Name the metal first. – IF THE METAL IS S TRANSITION METAL WITH MORE THAN ONE OXIDATION STATE, TEHN YOU MUST ADD A ROMAN NUMBERAL INDICATING THE OXIDATION STATE AFTER THE METAL NAME. – 2. Name the NONMETAL second but change the ending to – ide. – 3. Do not worry about how many of each element there are. This does not matter in the name, just give the names of the elements. • 1. NaCl ____________________ • 2. K2O _____________________ • 3. MgBr2 ____________________ • • 4. Al2S3 _____________________ • 5. Fe2O3 ____________________ • 6. FeO ______________________ 7. AuCl3 _____________________ 8. LiI ________________________ b. If the compound is made up of more than 2 different elements, check to see if there is a POLYATOMIC ION in the compound. Refer to the polyatomic chart. You should familiarize yourself with these ions so that you can easily recognize them. 1. If the compound is composed of a metal and a negative polyatomic ion: Metal + negative polyatomic ion - name the metal - then name the polyatomic ion. MgCO3 _____________________________ If the metal is a transition metal with more than one possible oxidation state, you must indicate the oxidation state using a Roman Numeral in parenthesis after the name of the metal. Ni(NO3)3 _______________________________________ 2. If the compound is between a positive polyatomic ion and a nonmetal: polyatomic ion + nonmetal - name the polyatomic ion first - then name the nonmetal and change the ending to –ide. NH4Cl _____________________________ 3. If the compound is between two polyatomic ions: polyatomic ion + polyatomic ion - name the positive ion first - name the negative ion second. • NEVER CHANGE THE ENDINGS OF POLYATOMIC IONS!!!!!!! Practice: 1. NaNO3 ________________________ ________________________ • 2. Al2(SO4)3 _______________________ _____________________________________ • 3. Ca(NO3)2 _______________________ • _______________________ 4. NH4Br ___________________________ ___________________________ 5. CuSO4 ___________________________ ___________________________ c. If the compound is made of 2 nonmetals, there are two ways to name it. You can name the compounds the same way you name the metals and nonmetals but you must use an oxidation number roman numeral in parenthesis to indicate the oxidation number of the first element. Naming Binary Compounds Compounds between two nonmetals - First element in the formula is named first. - Second element is named as if it were an anion. - Use prefixes - Only use mono on second element - • - • • • • P2O5 = diphosphorus pentoxide CO2 = carbon dioxide • 1. CO2 _____________________ • 2. CO ______________________ • 3. P2O5 ____________________ • 4. SiO2 _____________________ • 5. N2O4 ____________________ Naming Acids 1. Elements that begin with hydrogen are normally acids. 2. Naming Binary Acids a. Begin with the word HYDRO- and end with the non metal in –IC b. Then add the word ACID For example: HCl= Hydrochloric acid 3. Naming Ternary Acids a. Cover the hydrogen and look up the polyatomic ion being used. b. If the polyatomic ion ends in –ITE change the ending to –OUS and add the word ACID. For example: HClO2 = chlorite = Chlorous Acid c. If the polyatomic ion ends in –ATE change the ending to –IC and add the word acid. For example: H SO = sulfate = Sulfuric Acid Name the following acids: 1. HCl _______________________________ 2. HBr _______________________________ 3. HF _______________________________ 4. H2S ______________________________ 5. H3P ______________________________ 11. H3PO4 _________________________ 12. H2SO3 _________________________ 13. H2S2O3 ________________________ 6. H2Se ___________________________ 7. HI ______________________________ 8. H2CO3 __________________________ 9. H2C2O7 _________________________ 10. HNO2 __________________________ For each of the following acids below, write the correct formula. 1. Sulfuric Acid ________________________ 2. Hydrochloric Acid ____________________ 3. Clorous Acid ________________________ 4. Cloric Acid __________________________ 5. Thiocyanic Acid _______________________ 6. Acetic Acid _______________________ 7. Nitric Acid ___________________________ 8. Dichromic Acid _______________________ 9. HydroIodic Acid ______________________ 10. Nitrous Acid ________________________ How can you tell the difference between a binary acid and a ternary acid from its name?