chapter 7 chemical formulas and … 7.1 and 7.2 day1 chemical names and formulas ... list the rules...

CHAPTER 7 CHEMICAL FORMULAS AND CHEMICAL COMPOUNDS

SECTION 7.1 AND 7.2 DAY1 CHEMICAL NAMES AND FORMULAS

Objectives: ¢ Explain the significance of a chemical formula. ¢ Determine the formula of an ionic compound

formed between 2 given ions. ¢ Name an ionic compound given its formula. ¢ Using prefixes, name a binary molecular

compound from its formula. ¢ Write the formula of a binary molecular

compound given its name.

OBJECTIVES CONTINUED

¢ List the rules for assigning oxidation numbers. ¢ Give the oxidation number for each element in

the formula of a chemical compound. ¢ Name binary molecular compounds using

oxidation numbers and the Stock system.

OXIDATION NUMBERS

¢ What is the charge on the bromide ion in NaBr? ¢ Charges are physically real; oxidation numbers

are just for bookkeeping to help keep track of electrons.

OXIDATION NUMBERS (IN YOUR TEXTBOOK)

¢ Number precedes the sign

¢ Example: 2+

¢ Sign precedes the number

¢ Example: +2

Ionic charge Oxidation number

OXIDATION NUMBERS Assigning oxidation numbers: 1.  Atoms in a pure element have an oxidation number of zero. 2.  The more-electronegative element in a binary molecular

compound is assigned the number equal to the negative charge it would have as an anion. The less-electronegative atom is assigned the number equal to the positive charge it would have as a cation.

3.  Fluorine has an oxidation number of -1. 4.  Oxygen has an oxidation number of -2. (except in peroxides

in which it is -1 --- H2O2 and in compounds with halogens in which it is +2 --- OF2)

5.  Hydrogen has an oxidation number of +1. (except in compounds with metals in which it is-1)(CH4)

6.  In a compound, the sum of the oxidation numbers will equal zero.

7.  The sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion.

OXIDATION NUMBERS

Assign oxidation numbers to each of the following compounds:

a.  UF6

b.  H2SO4

c.  MgCl2

OXIDATION NUMBERS

Assign oxidation numbers to each atom in the following compounds or ions:

a.  HCl b.  CF4

c.  PCl3

d.  SO2

e.  HNO3

f.  KH g.  P4O10

h.  HClO3

i.  N2O5

j.  GeCl2

SECTION 7.1 CHEMICAL NAMES AND FORMULAS

¢ Formula

¢ Compound

Word Definition

DIFFERENTIATE THE 2 COLUMNS:

I II MgCl2 Ca(OH)2

KF KClO3 SO2 NH4OH

N2O5 FeCrO4

GeCl2 NaCH3COO

Al2S3 Ca(NO3)2

CuBr2 NaMnO4


I II MgCl2 SO2

Al2S3 N2O5

CuBr2 CO2

NaCl NH4

HOW MANY IONS OF EACH ELEMENT ARE PRESENT IS THE FOLLOWING COMPOUNDS?

1.  MgCl2

2.  KBr 3.  Al2S3

4.  CaI2 5.  NaCl What is the charge on each of the above atoms? Total charge on a compound is ZERO. These compounds are composed of monatomic ions. Monatomic ions = ions formed from a single type of

atom The above compounds are called binary ionic

compounds. Why?

NAMING BINARY IONIC COMPOUNDS

Rules: ¢ Name the cation with the elements name. ¢ Use the root of the anion and change the ending

to –ide. ¢ Example: NaCl = sodium chloride

NAMING BINARY COMPOUNDS

1.  MgCl2

2.  KBr 3.  Al2S3

4.  CaI2 5.  NaCl

Common anions: Fluoride Nitride Chloride Oxide Bromide Sulfide Iodide Phosphide

COMPLETE THE PRACTICE ON P. 207 #1-2. 1.  Write formulas for the binary ionic compounds formed

between the following elements: a. K and I b. Mg and Cl c. Na and S d. Al and S e. Al and N

2.  Name the binary ionic compounds indicated by the following formulas: a. AgCl b. ZnO c. CaBr2 d. SrF2 e. BaO f. CaCl2

ADDITIONAL PRACTICE:

1.  Write the formula for the following binary ionic compounds: a. Mg and I b. K and S c. Al and Cl d. Zn and Br e. Cs and S f. Sr and O g. Ca and N

2.  Name the following binary ionic compounds: a. BaF2

b. CaO c. AgF d. CdO e. K3N f. NaI g. AlBr3

STOCK SYSTEM OF NOMENCLATURE ¢  Some elements have 2 or more cations. ¢  Example: Fe+2 and Fe+3

¢  Mark these elements on your periodic table: Cu+1 and Cu+2

Cr+2 and Cr+3

Fe+2 and Fe+3

Pb+2 and Pb+4

Sn+2 and Sn+4

V+2 , V+3, and V Ag+1

¢  Roman numerals are placed after the name( never in a formula ) to indicate the charge or oxidation number

¢  Example: copper (I) chloride = CuCl iron (III) bromide = FeBr3

PRACTICE:

1. Write the formula and give the name for the compounds formed between the following ions: a. Cu+2 and Br-

b. Fe+2 and O-2

c. Pb+2 and Cl-

d. Hg+2 and S-2

e. Sn+2 and F-

f. Fe+3 and O-2

2. Give the names for the following compounds: a. CuO c. SnI4

b. CoF3 d. FeS

DAY 1 HOMEWORK: NAME THE FOLLOWING IONIC COMPOUNDS

1.  MgCl2

2.  KF

3.  GeCl2

4.  Al2S3

5.  CuBr 6.  CuBr2

7.  FeO 8.  Fe2O3

9.  MgS 10.  CaI2

11.  K2S 12.  CrCl2

13.  Ag2O 14.  CaO 15.  Ba3P4

16.  NaF 17.  Na2O 18.  BeS 19.  MnO 20.  BaCl2

21.  FeBr2

22.  CaCl2

23.  AgBr 24.  Na3P 25.  AlI3

26.  CdBr 27.  SnO 28.  Ba3N2

29.  VO 30.  NaCl

DAY 1 HOMEWORK: NAME THE FOLLOWING IONIC COMPOUNDS

DAY 1 HOMEWORK: WRITE THE FORMULA FOR THE FOLLOWING.

1.  tin (II) fluoride 2.  potassium oxide 3.  aluminum sulfide 4.  silver iodide 5.  copper (II) chloride 6.  magnesium oxide 7.  zinc sulfide 8.  lead (IV) nitride 9.  barium oxide 10.  lithium chloride

11.  sodium phosphide 12.  vanadium (II) chloride 13.  strontium (III) chloride 14.  silver fluoride 15.  cesium oxide 16.  aluminum bromide 17.  gold oxide 18.  potassium iodide 19.  titanium phosphide 20.  iron (II) oxide

DAY 1 HOMEWORK: WRITE THE FORMULA FOR THE FOLLOWING.

DAY 2 CHEMISTRY I POWER POINT PRESENTATION

Objectives: ¢ Explain the significance of a chemical formula. ¢ Determine the formula of an ionic compound. ¢ Design a presentation.

CHEMICAL COMPOUNDS POWERPOINT PRESENTATION

¢ This will be 2 TEST GRADES!!!!!!! ¢  1 grade for content ¢  1 grade for presentation


¢ Each student will select a metal and a non-metal to form their compound. No two students in a class are allowed to have the same compound.

¢ Take a few minutes to decide, then sign up your two elements beside your name on the list on my desk.

¢ This presentation will count as 2 TEST GRADES!


Items that you must have in your project to receive a grade of 100 on the content grade.

1. Chemical compound a. Chemical formula b. Name of chemical compound c. Common name if applicable d. Interesting facts e. Usages f. Physical description g. State of matter


2. For each element present in compound a. Symbol b. Number of protons, neutrons, and electrons c. Most common charge(s) d. Interesting facts e. Usages f. Most common state of matter g. Physical description


3. PowerPoint slides a. Minimum of 4 slides b. Minimum of 4 pictures c. 15 lines or less per slide d. Logical flow to slides e. Correct grammar and spelling f. Last slide must be work cited

LAST SLIDE

Works cited:

http://www.greensboro-nc.gov/NR/rdonlyres/3DE716E2-060E-43E2-ACC2-D4FDEE5B6825/0/formulas.jpg

http://www.findsounds.com/ISAPI/search.dll www.homewood.k12.al.us/~kreaves/oxidation.JPG


4. Day of presentation a. Paper copy must be turned in with

presentation b. Must be on time ---- NO EXCEPTIONS c. Must be on a flash drive to present to class and must be checked in the library before used on my computer


Items that you must have in your project to receive a grade of 100 on the presentation grade.

These are covered on the Presentation Rubric handout.


I II MgCl2 Ca(OH)2

KF KClO3 SO2 NH4OH

N2O5 FeCrO4

GeCl2 NaCH3COO

Al2S3 Ca(NO3)2

CuBr2 NaMnO4

POLYATOMIC IONS 1- 2- 3-

Acetate Carbonate Phosphate

Bromate Chromate Arsenate

Chlorate Dichromate

Chlorite Hydrogen phosphate

cyanide Oxalate

Dihydrogen phosphate

peroxide

Hydrogen carbonate (Bicarbonate)

Sulfate

Hydrogen sulfate Sulfite

Hydroxide

Hypochlorite

Nitrate

Nitrite

Perchlorate

Permanganate

POLYATOMIC IONS

¢  Oxyanions = polyatomic ions that contain oxygen ¢  Sometimes oxyanions are formed by the same 2

elements. ¢  Example: NO3

- and NO2- (NO3

- is the most common)

¢  The most common ion is given the ending –ate. ¢  The ion with one less oxygen is given the ending –

ite. ¢  Sometimes there are more ions. ¢  Example: ClO- ClO2

- ClO3- ClO4

-

¢  The anion with one less oxygen than the –ite ending has the hypo- prefix. The anion with one more oxygen than the –ate ending has the per- prefix.

¢  ClO3- is the most common.

¢  Name the anions above.

POLYATOMIC IONS

¢ Note: There are only 2 cations. ¢ Note: NH4

+

¢ What is the name of this cation? ¢ Do not confuse this with NH3. ¢ Remember an ion has a charge.

WRITE THE FORMULA FOR TIN (IV) SULFATE.

PRACTICE

Give the names for the following compounds: a. Ag2O b. Ca(OH)2

c. KClO3

d. NH4OH e. FeCrO4

f. KClO

PRACTICE:

Write the formulas for the following ionic compounds: a. Sodium iodide b. Calcium chloride c. Potassium sulfide d. Lithium nitrate e. Copper (II) sulfate f. Sodium carbonate g. Calcium nitrate h. Potassium perchlorate

EXTRA PRACTICE

Write the formulas for the following ionic compounds:

a.  Copper(II) nitrate b.  Potassium iodide c.  Sodium hydroxide d.  Ammonium acetate e.  Calcium carbonate f.  Potassium permanganate g.  Sodium sulfate h.  Iron(III) nitrate

EXTRA PRACTICE

Give the names of the following compounds: a.  Ag2S b.  NaMnO4

c.  Ba(OH)2

d.  NH4NO3

e.  Fe(ClO)2

f.  Ca(NO3)2

g.  K2SO3

h.  NaCH3COO


I II MgCl2 SO2

Al2S3 N2O5

CuBr2 CO2

NaCl NH3

NAMING BINARY MOLECULAR COMPOUNDS

¢ Unlike ionic compounds, molecular compounds are composed of individual covalently bonded units, or molecules!!!!!!

¢ There are 2 nomenclature systems to name binary molecules.

¢ The older system uses prefixes. ¢ The newer system is the Stock system.


Rules for prefix system: 1.  Less-electronegative element is given first. It is given a

prefix only if there is more than one atom. 2.  Second element is named by combining (a) a prefix

indicating the number of atoms, (b) the root of the second element, and (c) the ending –ide.

3.  The o or a at the end of a prefix is usually dropped when the word following the prefix begins with a vowel (example: monoxide or pentoxide).

Unlike ionic formulas, molecular formulas are not reduced to the simplest ratio of atoms in the compound. The actual number of atoms is shown.

Example: benzene C6H6


Number Prefix

1 mono-

2 di-

3 tri-

4 tetra-

5 penta-

6 hexa-

7 hepta-

8 octa-

9 nona-

10 deca-


Formula Prefix-system name

N2O

NO

NO2

N2O3

N2O4

N2O5


Give the name of As2O5. Write the formula for oxygen difluoride.


Name the following binary molecular compounds: a.  SO3

b.  ICl3

c.  PBr5

Name these using the Stock system. Write the formula for the following compounds: a.  Carbon tetriodide b.  Phosphorus trichloride c.  Dinitrogen trioxide How would these be named using the Stock

system?


Name the following binary molecular compounds: a.  PF5

b.  XeF4

c.  CCl4

Write formulas for the following compounds: a.  Carbon dioxide b.  Dinitrogen pentoxide c.  Sulfur hexafluoride

SECTION REVIEW P. 215

2.  Write the formulas for the compounds formed between the following:

a.  Aluminum and bromine b.  Sodium and oxygen c.  Magnesium and iodine d.  Pb2

+ and O2-

e.  Sn2+ and I-

f.  Fe3+ and S2

-

g.  Cu2+ and NO3

-

h.  NH4+ and SO4

2-

SECTION REVIEW

3.  Name the following compounds using the Stock system:

a.  NaI b.  MgS c.  CaO d.  K2S e.  CuBr f.  FeCl2

SECTION REVIEW

4.  Write formulas for each of the following compounds:

a.  Barium sulfide b.  Sodium hydroxide c.  Lead(II) nitrate d.  Potassoum permanganate e.  Iron(II) sulfate f.  Diphosphorus trioxide g.  Disulfur dichloride h.  Carbon diselenide

DAY 2 HOMEWORK: NAMING NON-BINARY COMPOUNDS 1.  NaNO3

2.  Ca(OH)2

3.  K2CO3

4.  NH4Cl 5.  MgSO4

6.  AlPO4

7.  (NH4)2SO4

8.  Na3PO4

9.  CuSO4

10.  NH4OH

11.  Li2SO3

12.  Mg(NO3)2

13.  Al(OH)3

14.  (NH4)3PO4

15.  KOH 16.  Ca(NO3)2

17.  K2SO4

18.  Pb(OH)2

19.  Na2O2

20.  CuCO3

Day 2 Homework: formulas with polyatomic ions – complete the table

OH -1 NO3-1 CO3

-2 SO4-2 PO4

-3

H HOH HNO3 H2CO3

Na

Mg

NH4+1

Ca

K

Al

Pb+4

DAY 2 HOMEWORK: NAMING BINARY COVALENT COMPOUNDS

1.  CO 2.  CO2

3.  SO2

4.  NO2

5.  N2O 6.  SO3

7.  CCl4

8.  NO 9.  N2O5

10.  P2O5

11.  N2O4

12.  CS2

13.  OF2

14.  PCl3

15.  PBr5

Day 3 Section 7.3 Using Chemical Formulas

Objectives: ¢ Calculate the formula mass or molar mass of any

given compound. ¢ Use molar mass to convert between mass in

grams and amount in moles of a chemical compound.

¢ Calculate the number of molecules, formula units, or ions in a given molar amount of a chemical compound.

¢ Calculate the percentage composition of a given chemical compound.

FORMULA MASSES

¢ Formula mass = sum of the average masses of all atoms represented in its formula

¢ Formula mass = molecular mass ¢ Atomic masses are rounded to 2 decimal places

for all calculations in this book. ¢ Formula mass and molar mass are numerically

equal, but different in units. ¢ Formula mass = amu ¢ Molar mass = g/mol

FIND THE FORMULA MASS FOR POTASSIUM CHLORATE.

What is the formula?

FIND THE MOLAR MASS OF EACH BELOW:

a.  Hydrogen sulfate b.  Calcium nitrate c.  Phosphate ion d.  Magnesium chloride

TEST NEXT CLASS ON NOMENCLATURE

DAY 3 HOMEWORK: NAME THE FOLLOWING

1.  BaCl2

2.  N2O 3.  Ag2O 4.  CuBr 5.  CuBr2

6.  NH4OH 7.  Fe2O3

8.  PBr5 9.  Al2O3

10.  Al2(SO4)3

11.  K2S 12.  CrCl2

13.  CrCl3

14.  NO3 15.  Ba3P2

16.  Hg2I2

17.  Na2O 18.  CO 19.  Pb(OH)3

20.  Mn2O3

DAY 3 HOMEWORK: NAME EACH AND CALCULATE THE MOLAR MASS

1.  Na2CO3

2.  NaOH 3.  MgBr2

4.  KCl 5.  FeCl2

6.  FeCl3

7.  Zn(OH)2

8.  Be2SO4

9.  CrF2

10.  Al2S3

Name Molar Mass

DAY 3 HOMEWORK: NAME EACH AND CALCULATE THE MOLAR MASS

11.  PbO 12.  Li3PO4

13.  TiI4

14.  Co3N2

15.  Mg3P2

16.  Ga(NO2)3

17.  Ag2SO3

18.  NH4OH 19.  Al(CN)3 20.  Be(CH3COO)2

Name Molar Mass

Day 3 Homework: write the formula and calculate the molar mass

31.  Iron (III) oxide 32.  Gallium nitride 33.  Iron (II) bromide 34.  Vanadium (V)

phosphate 35.  Calcium oxide 36.  Magnesium acetate 37.  Aluminum sulfate 38.  Copper (I) carbonate 39.  Barium oxide 40.  Ammonium sulfite

Formula Molar Mass

DAY 3 STUDY FOR TEST TEST NEXT CLASS ON NOMENCLATURE

DAY 5 WHAT IS THE MASS IN GRAMS OF 2.50 MOL OF OXYGEN GAS?

Oxygen gas = O2

WHAT IS THE MASS IN GRAMS OF 6.25 MOL OF COPPER (II) NITRATE?

HOW MANY MOLES OF COMPOUND ARE THERE IN THE FOLLOWING:

a.  6.60 g (NH4)2SO4

b.  4.5 g Ca(OH)2

HOW MANY MOLECULES ARE THERE IN THE FOLLOWING:

a.  25.0 g H2SO4

b.  125 g sugar, C12H22O11

PERCENTAGE COMPOSITION

Find the percentage composition of Cu and S in copper (I) sulfide.

Do your results equal 100%?

PRACTICE

Find the % compositions of the following: a.  PbCl2

b.  Ba(NO3)2


1.  Determine both the formula mass and the molar mass of ammonium carbonate.

2.  How many moles of atoms of each element are there in one mole of ammonium carbonate?

3.  What is the mass in grams of 3.25 mol iron(III) sulfate?

4.  How many moles of molecules are there in 250 g of hydrogen nitrate?

5.  How many molecules of aspirin, C9H8O4 are there in a 100.0 mg tablet of aspirin?

6.  Calculate the % composition of each element found in ammonium carbonate.

DETERMINE BOTH THE FORMULA MASS AND THE MOLAR MASS OF AMMONIUM CARBONATE.

HOW MANY MOLES OF ATOMS OF EACH ELEMENT ARE THERE IN ONE MOLE OF AMMONIUM CARBONATE?

WHAT IS THE MASS IN GRAMS OF 3.25 MOL IRON(III) SULFATE?

HOW MANY MOLES OF MOLECULES ARE THERE IN 250 G OF HYDROGEN NITRATE?

HOW MANY MOLECULES OF ASPIRIN, C9H8O4 ARE THERE IN A 100.0 MG TABLET OF ASPIRIN?

CALCULATE THE % COMPOSITION OF EACH ELEMENT FOUND IN AMMONIUM CARBONATE.

DAY 5 HOMEWORK: P. 236 #32-34

DAY 6 SECTION 7.4 DETERMINING CHEMICAL FORMULAS

Objectives: ¢ Define empirical formula, and explain how the

term applies to ionic and molecular compounds. ¢ Determine an empirical formula from either a

percentage or a mass composition. ¢ Explain the relationship between the empirical

formula and the molecular formula of a given compound.

¢ Determine a molecular formula from an empirical formula.

SECTION 7.4 DETERMINING CHEMICAL FORMULAS

¢ Empirical formulas = formula showing the smallest whole-number mole ratio of atoms in the compound

¢ Problem: Find the empirical formula of a compound that contains 32.38% Na, 22.65% S, and 44.99% O. First change % to grams then calculate moles of each. Divide each by the smallest of these numbers. Round to the nearest whole number.

PROBLEM

Analysis of a 10.150 g sample of a compound known to contain only P and O indicates a P content of 4.433 g. What is the empirical formula of this compound?

PROBLEM

Using the information form the last problem, determine the molecular formula of this compound if the molar mass is 283.89 g/mol.


1.  A compound is found to contain 36.48% Na, 25.41% S, and 38.11% O. Find its empirical formula.

2.  Find the empirical formula of a compound that contains 53.70% Fe and 46.30% S.

3.  Analysis of a compound indicates that it contains 1.04 g K, 0.70 g Cr, 0.86 g O. Find its empirical formula.

4.  If 4.04 g of N combine with 11.46 g of O to produce a compound with a formula mass of 108.0 amu, what is the molecular formula of this compound?

5.  The molar mass of a compound is 92 g/mol. Analysis of a sample of the compound indicates that it contains 0.606 g N and 1.390 g O. Find its molecular formula.

A COMPOUND IS FOUND TO CONTAIN 36.48% NA, 25.41% S, AND 38.11% O. FIND ITS EMPIRICAL FORMULA.

FIND THE EMPIRICAL FORMULA OF A COMPOUND THAT CONTAINS 53.70% FE AND 46.30% S.

ANALYSIS OF A COMPOUND INDICATES THAT IT CONTAINS 1.04 G K, 0.70 G CR, 0.86 G O. FIND ITS EMPIRICAL FORMULA.

IF 4.04 G OF N COMBINE WITH 11.46 G OF O TO PRODUCE A COMPOUND WITH A FORMULA MASS OF 108.0 AMU, WHAT IS THE MOLECULAR FORMULA OF THIS COMPOUND?

THE MOLAR MASS OF A COMPOUND IS 92 G/MOL. ANALYSIS OF A SAMPLE OF THE COMPOUND INDICATES THAT IT CONTAINS 0.606 G N AND 1.390 G O. FIND ITS MOLECULAR FORMULA.

CLASS WORK: P. 236 #26, 30, 36, 37, 47, 50

chapter 7 chemical formulas and … 7.1 and 7.2 day1 chemical names and formulas ... list the rules...

Documents