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Chapter 7 Chemical Formulas & Chemical Compounds 7.1 Chemical Names & Formulas.

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Presentation on theme: "Chapter 7 Chemical Formulas & Chemical Compounds 7.1 Chemical Names & Formulas."— Presentation transcript:

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2 Chapter 7 Chemical Formulas & Chemical Compounds 7.1 Chemical Names & Formulas

3 Ions Cation: A positive ion Mg 2+, NH 4 + Anion: A negative ion Cl , SO 4 2  Ionic Bonding: Force of attraction between oppositely charged ions.

4 Predicting Ionic Charges Groups 3 - 12: Many transition elements have more than one possible oxidation state. Iron (II) = Fe 2+ Iron (III) = Fe 3+

5 Predicting Ionic Charges Groups 3 - 12: Some transition elements have only one possible oxidation state. Zinc = Zn 2+ Silver = Ag 1+

6 Formula Writing for Binary Ionic Compounds Magnesium Bromide Mg + 2 Br – 1 Mg 1 Br 2 MgBr 2 Calcium Sulfide Ca + 2 S – 2 Ca 2 S 2 CaS criss-cross the oxidation numbers to balance out the charge.

7 Writing Ionic Compound Formulas Example: Iron (III) chloride 1. Write the formulas for the cation and anion, including CHARGES! Fe 3+ Cl - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Not balanced! 3

8 Naming Ionic Compounds 1. Cation first, then anion 2. Monatomic cation = name of the element Ca 2+ = calcium ion 3. Monatomic anion = root + -ide Cl  = chloride CaCl 2 = calcium chloride

9 Naming Ionic Compounds some metal forms more than one cation use Roman numeral in name PbCl 2 Pb 2+ is cation PbCl 2 = lead (II) chloride Metals with multiple oxidation states

10 Elements with Multiple Oxidation Numbers Copper I Cu +1 Copper II Cu +2 Iron II Fe +2 Iron III Fe +3 Mercury I Hg +1 Mercury II Hg +2 Lead II Pb +2 Lead IV Pb +4 Tin II Sn +2 Tin IV Sn +4 Chromium II Cr +2 Chromium III Cr +3 Chromium VI Cr +6 Manganese IIMn +2 Manganese IIIMn +3 Manganese VIIMn +7 Cobalt IICo +2 Cobalt IIICo +3 Gold IAu +1 Gold IIIAu +3 Nickel IINi +2 Nickel IIINi +3 Nickel IVNi +4 **Silver Ag +1 **Zinc Zn +2 **Cadmium Cd +2

11 ♥ Poly Atomic Ions to Know and Love ♥ Name Formula Name Formula Acetate C 2 H 3 O 2 -1 (CH 3 COO -1 ) Hypochlorite ClO -1 Dichromate Cr 2 O 7 -2 Chlorite ClO 2 -1 Ammonium NH 4 +1 Chlorate ClO 3 -1 Nitrate NO 3 -1 Perchlorate ClO 4 -1 Nitrite NO 2 -1 Cyanide CN -1 Hydroxide OH -1 Carbonate CO 3 -2 Phosphate PO 4 -3 Chromate CrO 4 -2

12 Sulfite SO 3 -2 Hydrogen Carbonate HCO 3 -1 Sulfate SO 4 -2 Hydrogen Phosphate HPO 4 -2 Hydrogen Sulfite HSO 3 -1 Hydrogen Sulfate HSO 4 -1 Permanganate MnO 4 -1 Oxalate C 2 O 4 -2 Hydronium H3O+H3O+H3O+H3O+Silicate SiO 3 -2 Peroxide O 2 -2 Phosphite PO 3 -3 Bromate BrO 3 -1 Arsenate AsO 4 -2 ♥ More Poly Atomic Ions to Know and Love ♥ Name Formula Name Formula

13 Naming Compounds with Polyatomic Ions FormulaName (NH 4 ) 2 SO 4 ammonium sulfate ZnCO 3 zinc carbonate NH 4 Brammonium bromide Li 2 CO 3 lithium carbonate * Polyatomic & monatomic cation names remain the same, monatomic anions change their ending to –ide.

14 Writing Ionic Compound Formulas Example: Barium nitrate 1. Write the formulas for the cation and anion, including CHARGES! Ba 2+ NO 3 - 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. Not balanced! ( ) 2

15 Writing Ionic Compound Formulas Example: Ammonium sulfate 1. Write the formulas for the cation and anion, including CHARGES! NH 4 + SO 4 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. Not balanced! ( ) 2

16 Writing Ionic Compound Formulas Example: Aluminum sulfide 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ S 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. Not balanced! 23

17 Writing Ionic Compound Formulas Example: Magnesium carbonate 1. Write the formulas for the cation and anion, including CHARGES! Mg 2+ CO 3 2- 2. Check to see if charges are balanced. They are balanced! 3. Simplify to a formula unit.

18 Writing Ionic Compound Formulas Example: Zinc hydroxide 1. Write the formulas for the cation and anion, including CHARGES! Zn 2+ OH - 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. Not balanced! ( ) 2

19 Writing Ionic Compound Formulas Example: Aluminum phosphate 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ PO 4 3- 2. Check to see if charges are balanced. They ARE balanced!

20 More Examples… 1.Cr 2 O 3 2.Cr 2 O 3.CuSO 4 4.Ni(OH) 2 5.Cr 2 (C 2 O 4 ) 3 6.Cu 2 S 7.CuS 1.chromium (III) oxide 2.chromium (I) oxide 3.copper (II) sulfate 4.nickel (II) hydroxide 5.chromium (III) oxalate 6.copper (I) sulfide 7.copper (II) sulfide Chemical FormulaChemical Name

21 Hydrates Hydrate – when a water molecule (s) are chemically bonded to the ionic compound. Normal ionic naming protocol are used, then followed by the word “hydrate.” Prefixes are added to indicate the number of water molecules when naming hydrates.

22 Hydrate Prefixes # of water molecules prefix# of water molecules prefix 1mono-6hexa- 2di-7hepta- 3tri-8octa- 4tetra-9nona- 5penta-10deca-

23 Hydrates Example: MgBr 2 ∙ 6H 2 O Magnesium bromide hexahydrate The “ ∙ ” means “loosely bonded” Hygroscopic - easily absorb water molecules from the air. Deliquescent- very hygroscopic; takes out water from the air to dissolve completely to form a liquid solution. Anhydrous – when all of the water has been removed.

24 Naming Binary Covalent 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 P 2 O 5 = CO 2 = CO = N 2 O = diphosphorus pentoxide carbon dioxide carbon monoxide dinitrogen monoxide

25 Acids Acids always begin with Hydrogen AnionFormulaName Cl -1 HClHydrochloric Acid Br -1 HBrHydrobromic Acid SO 4 -2 H 2 SO 4 Sulfuric Acid SO 3 -2 H 2 SO 3 Sulfurous Acid NO 3 -1 HNO 3 Nitric Acid CN -1 HCNHydrocyanic Acid PO 4 -3 H 3 PO 4 Phosphoric Acid

26 Bases CationFormulaName Na +1 NaOHSodium Hydroxide K +1 KOHPotassium Hydroxide NH 4 +1 NH 3 Ammonia

27 Organic Compounds Organic compounds are named using a different set of rules. The simplest group is the hydrocarbons. These compounds are composed solely of the elements carbon and hydrogen. Carbon atoms can link to each other in chains and in rings.

28 Naming Hydrocarbons The stem of the compound name is then chosen from the following table: # of carbon atoms prefix# of carbon atoms prefix 1meth-6hexa- 2eth-7hepta- 3prop-8octa- 4but-9nona- 5penta-10deca-

29 Hydrocarbons: Alkanes These molecules have the generic formula: C n H 2n+2 They contain all single bonds. CH 4 methane C2H6C2H6 ethane C3H8C3H8 propane C 4 H 10 butane C 5 H 12 pentane C 6 H 14 hexane

30 Hydrocarbons: Alkenes These molecules have the generic formula: C n H 2n They contain double bonds between carbon atoms. C2H4C2H4 ethene C3H6C3H6 propene C4H8C4H8 butene C 5 H 10 pentene C 6 H 12 hexene

31 Hydrocarbons: Alkynes These molecules have the generic formula: C n H n They contain triple bonds between carbon atoms. C2H2C2H2 ethyne C3H3C3H3 propyne C4H4C4H4 butyne C 5 H 5 pentyne C6H6C6H6 hexyne

32 Chapter 7 Chemical Formulas & Chemical Compounds 7.2 Oxidation Numbers

33 Oxidation Numbers Oxidation Number – numbers assigned to atoms composing a compound or ion that indicate the general distribution of electrons among bonded atoms

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35 Chapter 7 Chemical Formulas & Chemical Compounds 7.3 Using Chemical Formulas

36 Molar Mass The mass of 1 mole of a pure substance is called its Molar Mass. Ex: Molar mass of Iron is 55.847 g/mol What is the molar mass of Platinum? 195.08 g/mol

37 Molar Mass The molar mass depends on the particles that compose the compound. If your element exists as a molecule, i.e. BrINClHOF, one mole of these particles contains 2 moles of the element as an atom. Determine the molar mass of oxygen molecules (O 2 ) (16.00 g/mol) x (2 atoms) = 32.00 g/mol The molar mass of oxygen molecules (O 2 ) is twice the molar mass of oxygen atoms!

38 Formula Mass The molar mass of a compound is the mass of the atomic mass units of one molecule. This takes into consideration the number of atoms of each element in a compound. Formula Mass is calculated the same way as molar mass except it is measured in amu, instead of g/mol.

39 Calculating Formula Mass Calculate the formula mass of magnesium carbonate, MgCO 3. 24.31 + 12.01 + 3(16.00) = 84.32 amu

40 Steps for Calculating Molar Mass for Compounds 1.List the elements 2.Determine how many atoms of each 3.Identify the atomic masses from the periodic table 4.Multiply how many atoms by the respective atomic mass 5.Add up the totals for the Molar Mass

41 Practice H 2 O H2 x 1.008 = 2.016 O1 x 15.99 = 15.99 18.006 C 6 H 12 O 6 C6 x 12.01 = 72.06 H 12 x 1.008 = 12.096 O6 x 15.99 = 95.94 180.096 NaCl Na1 x 22.9 = 22.9 Cl1 x 35.45 = 35.45 58.35 K 2 O K2 x 39.1 = 78.2 O1 x 15.99 = 15.99 94.19 g/mol

42 Calculating Percentage Composition Calculate the percentage composition of magnesium carbonate, MgCO 3. 24.31 + 12.01 + 3(16.00) = 84.32 amu 100.00

43 So…. In one mole of H 2 O, how many grams of Hydrogen are there? 2 mol H x 1.008g H = 2.016 g H in 1 mol H 2 O 1 mol H What % of Hydrogen, by mass, is in H 2 O? 2.016 g H x 100 = 11.2 % H 18 g H 2 0 *Must also find molar mass of H 2 O What % of Oxygen, by mass is in H 2 O? Mass Percent

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45 Formulas  molecular formula = (empirical formula) n [n = integer]  molecular formula = C 6 H 6 = (CH) 6  empirical formula = CH Empirical formula: the lowest whole number ratio of atoms in a compound. Molecular formula: the true number of atoms of each element in the formula of a compound.

46 Formulas Formulas for ionic compounds are ALWAYS empirical (lowest whole number ratio). Often, these are called formula units. Examples: NaClMgCl 2 Al 2 (SO 4 ) 3 K 2 CO 3

47 Formulas Formulas for molecular compounds MIGHT be empirical (lowest whole number ratio). Molecular: H2OH2O C 6 H 12 O 6 C 12 H 22 O 11 Empirical: H2OH2O CH 2 O C 12 H 22 O 11

48 Chapter 7 Chemical Formulas & Chemical Compounds 7.4 Determining Chemical Formulas

49 Empirical Formula Determination 1.Base calculation on assumption of 100 grams of compound. 2.Determine moles of each element in 100 grams of compound. 3.Divide each value of moles by the smallest of the values. 4.Multiply each number by an integer to obtain all whole numbers.

50 Empirical Formula Determination Adipic acid contains 49.32% C, 43.84% O, and 6.85% H by mass. What is the empirical formula of adipic acid?

51 Empirical Formula Determination (part 2) Divide each value of moles by the smallest of the values. Carbon: Hydrogen: Oxygen:

52 Empirical Formula Determination (part 3) Multiply each number by an integer to obtain all whole numbers. Carbon: 1.50 Hydrogen: 2.50 Oxygen: 1.00 x 2 352 Empirical formula: C3H5O2C3H5O2

53 Finding the Molecular Formula The empirical formula for adipic acid is C 3 H 5 O 2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid? 1. Find the formula mass of C 3 H 5 O 2 3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g

54 Finding the Molecular Formula The empirical formula for adipic acid is C 3 H 5 O 2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid? 3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g 2. Divide the molecular mass by the mass given by the emipirical formula.

55 Finding the Molecular Formula The empirical formula for adipic acid is C 3 H 5 O 2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid? 3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g 3. Multiply the empirical formula by this number to get the molecular formula. (C 3 H 5 O 2 ) x 2 = C 6 H 10 O 4

56 Determining Chemical Formulas from Mass Percents A sample has been analyzed, here are the results: 18.8 % Na 29.0 % Cl 52.2 % O How can you determine the chemical formula?

57 Step 1: Assume a 100 g sample. Then, your percent quantities become gram (mass) quantities. 18.8 g Na, 29.0 g Cl & 52.2 g O Step 2: Convert those masses to moles. 18.8 g Na x 1 mol Na = 0.817 mol Na 23 g Na 29.0 g Cl x 1 mol Cl = 0.817 mol Cl 35.5 g Cl 52.2 g O x 1 mol O = 3.26 mol O 16 g O

58 Step 3: Since your empirical formula is in small, whole number ratios, divide your mole amounts by the smallest mole quantity. 0.817 mol Na / 0.817 = 1.00 mol Na 0.817 mol Cl / 0.817 = 1.00 mol Cl 3.26 mol O / 0.817 = 3.99 ≈ 4.00 mol O Step 4: Use these values as subscripts in your formula Na 1 Cl 1 O 4 ≈ NaClO 4

59 Step 5: In the event the chemical formula is not the same as the empirical formula, you need the molar mass of the desired compound and you must compare it to the molar mass of the empirical formula. Step 6: Divide the given molar mass by the empirical molar mass to get the multiple quantity. Step 7: Multiply each subscript in the formula by that multiple quantity.

60 Ex: MM of molecular formula = 180 g/mol Using steps 1-4, you found that the empirical formula is CH 2 O. Find the molar mass of the empirical formula: MM EF = 30 g/mol Divide MM MF / MM EF to get a whole number. Ex: 180 / 30 = 6 C 1x6 H 2x6 O 1x6 C 6 H 12 O 6

61 Practice Problems: A sample has been analyzed to be 10.04 % C, 0.84 % H & 89.12% Cl. Find the Empirical Formula. A compound’s empirical formula has been determined to be HF. The compound’s molar mass is 40 g/mol. What is its chemical formula? A compound’s empirical formula has been determined to be CH 2. The compound’s molar mass is 42 g/mol. What is its chemical formula?


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