CHEMISTRY 110 LECTURE EEXXAAMM IIII MMaatteerriiaall

I. The Atom - history: A. Democritus (approx. 400 BC) - Matter is made of minute indivisible, invisible particles called atoms

B. Dalton’s atomic theory (approx. 1800 AD)- Experimentally based: Atoms are the building blocks of matter.

a. Each element is made of tiny individual particles called atoms

b. Atoms are indivisible and invisible. They cannot be created or destroyed.

c. All atoms of each element are identical in every respect.

d. Atoms of one element are different from atoms of any other element.

e. Atoms of one element may combine with atoms of another element in the ratio of small whole numbers to form chemical compounds.

II. Subatomic Particles: -

electron proton neutron Location in atom

charge

mass

1 amu= 1.99 x 10-23 g

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III. Atomic number, atomic mass, Nuclear symbol and Isotopes Periodic table 14 atomic number=number of protons Si 28.09 atomic mass=weighted average of all isotopes A X Z

A. 1. Atomic number, ,, equals the number of protons

2. Mass number, , equals the sum of protons and neutron (nucleons). That is, the number of subatomic particles that contribute substantially to the mass of the atom

3. X is the symbol of the element

The number of electrons, , equals the number of protons in a neutral atom 4. Isotopes are atoms with the same number of protons but different number of neutrons Isotopes of Silicon:

Nuclear symbols

28

Si 14

29

Si 14

30

Si 14

Si-28 Si-29 Si-30 Number of protons Number of electrons Number of neutrons Isotopic atomic mas 27.9769 amu 28.9765 amu 29.9737 amu abundance 92.21% 4.69% 3.10 %

C. IONS-Ions are atoms ( or a group of atoms) with a charge Na atom Cl atom 23 protons 35 protons 23 electrons 35 electrons

Ion-Problem 1. a. Calculate the number of protons, electrons and neutrons of an O-16 atom that has a -2 charge.

b. How many subatomic particles does this ion have?

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IV. Principal levels and sublevels A. Model of an atom:

Electrons in the atoms are found in principal energy levels (n) also called shells. The shells are represented here by concentric spheres around the nucleus. The shell closest to the nucleus has the lowest potential energy. Each principal level is subdivided into sublevels designated as: s-,p-,d-,and f –sublevels. Each sublevel contains orbitals

Maximum number of electrons that can occupy a principal energy level = 2n2 n=∞ potential energy n=4 _________32 e- increases n=3___________18e- n=2 ___________8 e- n=1 ___________2 e- Sublevels: The s-, p-, d-, and f-sublevels have one- , three-, five-, and seven -orbitals respectively. . Each orbital may be occupied by no more than two electrons. Therefore, a filled s-, p-, d-, and f- sublevels may be designated as s2, p6, d10, and f14 respectively. ___________ f _____ ______d_____ ____p __ _ s

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The Electronic Configuration (Distribution of Electrons among Sublevels) Electrons will fill the sublevels of lower potential energy before proceeding to those of higher energy a) Using TheTriangular Array: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 8s Exercise: Write the electron configuration for : i) Co (Z=27) ii) S (Z= 16) iii) Ca (Z= 20) iv) S2- (Z =16) b) Using The Periodic Chart: 1 s 1 s 2 s

2 p 3 s

3 p 4 s 3 d

4 p 5 s 4 d

5 p 6 s 5 d

6 p 7 s 6 d

7 p 4 f 5 f

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Why do elements in the same group (A-groups) have similar properties?

Group I A ( All have one valence electron) H (Z=1) 1s1 Li (Z=3) 1s2 2s1 Na (Z= 11) 1s2 2s2 2p6 3s1 K (Z= 19) 1s2 2s2 2p6 3s2 3p6 4s1 Rb (Z=37) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 Cs (Z=55) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s1 Group II A ( All have 2 valence electrons) Be (Z=4) 1s2 2s2 Mg (Z= 12) 1s2 2s2 2p6 3s2 Ca (Z= 20) 1s2 2s2 2p6 3s2 3p6 4s2 Sr (Z=38) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 Ba (Z=56) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 Group VII A ( All have 7 valence electrons) F (Z=9) 1s2 2s22p5 Cl (Z= 17) 1s2 2s2 2p6 3s23p5 Br (Z= 35) 1s2 2s2 2p6 3s2 3p6 4s24p5 I (Z=53) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 5p5 At (Z=85) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s26p5 Valence shell= Outermost occupied shell of a particular atom. The Rare Gases ( Inert Gases, Noble Gases) Group VII IA ( All have 8 valence electrons) He (Z=2) 1s2 Ne (Z=10) 1s2 2s22p6 Ar (Z= 18) 1s2 2s2 2p6 3s23p6 Kr (Z= 36) 1s2 2s2 2p6 3s2 3p6 4s24p6 Xe (Z=54) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 5p6 Rn (Z=86) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s26p6

Valence electrons These are the electrons in the Valence Shell (outermost shell). Valence electrons are involved in reactions.

Note: # valence e- = the group number for the "A" elements. Exercise: Write the electron configuration for each of the following : 1. O (Z= 8) (Representative element, A-Group)

2. Ni (Z= 28)(Transitional element, B-Group)

3. Na+ (Z= 11)

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4. S2- (Z= 16)

5. Co (Z= 27) (Transitional element, B-Group)

Lewis Electron- dot Structure for the Representative Elements( A-Groups) The valence electrons are represented by dots N O S Al Cl Xe

VI) THE VARIATION OF PROPERTIES A) ATOMIC SIZE: Atoms are very tiny, they have diameters of about 10 -10 meter. 1) As we proceed down within a group, the size of atoms, generally, increases and that as we proceed from left to right across a period, a gradual decrease in size is observed. 2) The factors that determine the size of the atoms are: i) The order of the outer shell. ii) The amount of nuclear charge that the outer electrons feel. B) IONIC SIZE: In general positive ions are smaller than the neutral atoms from which they are formed, while negative ions are larger than neutral atoms. The decrease in size that accompanies the creation of a positive ion is often a result of the removal of all electrons from the outer shell of the atom. Example: When negative ions are produced from neutral atoms, electrons are added to the outer shell without any change to the nuclear charge. The effective nuclear charge felt by any one electron in the outer shell decreases. B) IONIZATION POTENTIAL (ENERGY): It is defined as the energy required to remove an electron from an isolated gaseous atom. First ionization energy: Second ionization energy:

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The variation of the first ionization energy across periods and down groups, parallels the trends in atomic size. Thus as we proceed down within a group, the increase in size that occurs is accompanied by a decrease in ionization energy. As we move across a period, from left to right, the increased effective nuclear charge experienced by the outer shell electrons causes the shell to shrink in size and also makes it more difficult to remove an electron, hence an increase in the ionization potential is observed. C) ELECTRON AFFINITY: It is the energy that is released or absorbed when an electron is added to a neutral gaseous atom . Example: As with the ionization energy, the variations in electron affinity generally parallel the variations in atomic size. Therefore atoms that are very small and have outer shells that experience a high effective nuclear charge have very large electron affinities (elements in the upper right of the periodic table). On the other hand, atoms that are large and whose outer shells feel the effect of a small effective nuclear charge have small electron affinity (elements in the bottom left of the periodic table).

The Octet Rule: Atoms react to obtain an "octet" ( 8 valence e-) Elements achieve an octet by: 1)sharing valence electrons with other atoms forming covalent bond. 2) losing or gaining valence electrons forming ionic bond.

A. Covalent Bond 1. A covalent bond is a pair of electrons shared between two nonmetal atoms 2. Each of the bonding atoms will complete its valence shell and achieve an octet by sharing electrons. 3. Each atom will have eight electrons around it, except if it is: a. hydrogen ( will have only 2 electrons, forming a duet) b. an element in the third or lower period on the periodic table that occupies the center of a molecule (may have more than 8 electrons).

Lewis Electron-dot Structure for covalent molecules: a. Diatomic elements H2

Cl2

O2

N2

*There are never more than 3 covalent bonds between two atoms

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b. Binary covalent compounds (Nonmetal-nonmetal compounds)

Lewis electron-dot struc Lewis electron-dot strucH2O

CCl4

CO2

NH3

CH4

EXCEPTIONS TO THE OCTET RULE: 1) ELECTRON DEFICIENT MOLECULES: This is typical of molecules where the central atom belongs to group IIIA. Example: 2) EXPANDED VALENCE SHELL: This is typical of molecules where the central atom belongs to the third, fourth, fifth, sixth or seventh period. Example: 3) ODD MOLECULES: This is typical of molecules where the total number of valence electrons is an odd number. B) RULES FOR DRAWING LEWIS STRUCTURE: 1) Count all the valence electrons for the atoms. ( If the species is an ion, add an additional electron for each negative charge or subtract an electron for each positive charge.) 2) Place one pair of electrons for each bond. 3) Complete the octets of the atoms bonded to the central atom. (Remember that the valence shell of any hydrogen atom is complete with only two electrons.) 4) Place any additional electrons on the central atom in pairs. 5) If the central atom still has less than an octet, you must form multiple bonds so that each atom has an octet.

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C) PRACTICE: Draw a Lewis structure (Electron-dot structure) for each of the following: 1) SCl4

2) SO2

3) NH4+ 4) ICl4 -

5) ClF3 6) NO3 -

7) CO32- 8) CO2

9) HCN 10) SO2

11) SF4 12) BrF5

13) NH3 14) H2O

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B. Ionic Bond Ionic bond is the electrostatic attraction between positive and negative ions. Examples of ionic compounds (metal-nonmetal compounds): Write the chemical formula for the following ionic compounds 1) sodium chloride 2) calcium chloride 3) aluminum oxide

Periodic Table: Simple Cations and Anions

IA

H+

IIA

IIIA

IVA

VA

VIA VIIA

Li+

Be2+ N3- O2- F -

Na+ Mg2+

IB

IIB Al3+ P3- S2- Cl -

K+ Ca2+

Cr2+

Cr3+ Mn2+

Mn3+ Fe2+

Fe3+ Co2+

Co3+ Ni2+

Ni3+ Cu+

Cu2+ Zn2+

As3+

As5+ Se2- Br-

Rb+ Sr2+

Ag+

Cd2+

Sn2+

Sn4+ Sb3+

Sb5+ Te2- F -

Cs+ Ba2+

Au+

Au3+ Hg2+

Hg22+

Pb2+

Pb4+ BI3+

BI5+

Monoatomic IONS 1. Non-metal Ions H+ hydrogen Group VA Group VIA Group VIIA N3- nitride O2- oxide F- fluoride P3- nitride S2- sulfide Cl- chloride Br - bromide I – iodide

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2. Metal Ions that exhibit a fixed charge Group IA Group IIA Group IIIA B-Group Li+ lithium Mg2+ magnesium Al3+ aluminum Ag+ silver Na+ sodium Ca2+ calcium Zn2+ zinc K+ potassium Sr2+ strontium Cd2+ cadmium Ba2+ barium Metal Ions that exhibit variable charges Use the Roman Numeral or the Classical common name Group IVA Group VA Pb2+ lead (II) Pb4+ lead(IV) Sb3+ antimony(III) Sb3+ antimony(III) (plumbous) (plumbic) Sn2+ tin(II) Sn4+ tin(IV) As3+ arsenic(III) As5+ arsenic(V) (stannous) (stannic) Group IB Cu+ copper(I) Cu2+ copper(II) (cuprous) (cupric) Au+ gold(I) Au3+ gold(III) (aurous) (auric) Group IIB Hg2

2+ mercury(I) Hg2+ mercury(II) (mercurous) (mercuric) Other B- Groups Cr2+ chromium(II) Cr3+ chromium(III) (chromous) (chromic) Mn2+ manganese(II) Mn3+ manganese(III) (manganous) (manganic) Fe2+ iron(II) Fe3+ iron(III) (ferrous) (ferric) Co2+ cobalt(II) Co3+ cobalt(III) (cobaltous) (cobaltic) Ni2+ nickel(II) Ni3+ nickel(III) (nickelous) (nickelic) Writing Chemical Formulas and Nomenclature of Ionic Compounds Containing Monoatomic Ions Practice:

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Polyatomic Ions: A group of atoms bonded together to form an ion *(Memorize the list of polyatomic ions given below) NH4

+ ammonium OH- hydroxide CN- cyanide OCN- cyanate SCN- thiocyanate C2H3O2

- acetate

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MnO4

- permanganate C2O4

2- oxalate --------------------------------------------------------------------------------------------------------------------- CrO4

2- chromate Cr2O72- dichromate

------------------------------------------------------------------------------------------------------------------- CO3

2- carbonate HCO3- hydrogen carbonate, or bicarbonate

-------------------------------------------------------------------------------------------------------------------- NO2

- nitrite NO3- nitrate

---------------------------------------------------------------------------------------------------------------------- SO3

2- sulfite SO42- sulfate

HSO3- hydrogensulfite, or bisulfite HSO4

- hydrogensulfate, bisulfate

S2O32- thiosulfate

---------------------------------------------------------------------------------------------------------------------- PO3

3- phosphite PO43- phosphate

HPO42- monohydrogenphosphate

H2PO4- dihydrogenphosphate

---------------------------------------------------------------------------------------------------------------------- ClO- hypochlorite BrO- hypobromite IO- hypobromite ClO2

- chlorite BrO2- bromite IO2

- bromite ClO3

- chlorate BrO3- bromate IO3

- iodate ClO4

- perchlorate BrO4- perbromate IO4

- periodate --------------------------------------------------------------------------------------------------------------- Prefixes and suffixes -ite = one less oxygen atom than “ate” per- =one more oxygen atom than “ate” hypo- = one less oxygen atom than “ite” thio- = one oxygen atom is replaced by S bi- = one H+ added to divalent anion di- = two

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Writing Chemical Formulas and Nomenclature of Ionic Compounds Containing Polyatomic Ions Practice:

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Practice Chemical Formulas Write chemical formulas for the ionic compounds made from each set of ions: Name of cation Name of anion formula of

cation Formula of anion

Formula of compound

number of ions

1

sodium chloride

Na + Cl -

2 calcium chloride

Ca 2+ Cl -

3 aluminum chloride

Al 3+ Cl -

4 calcium nitrite

Ca 2+ NO2 -

5 calcium sulfate

Ca 2+ SO4 2-

6 sodium phosphate

Na + PO4 3-

7 potassium bromide

K + Br -

8 calcium bicarbonate

Ca2+ HCO3 -

9 magnesium bisulfate

Mg 2+ HSO4 -

10 barium nitrate

Ba 2+ NO3 -

11 ammonium bromide

NH4 + Br -

12 aluminum phosphate

Al 3+ PO4 3-

13 potassium phosphate

K+ PO43-

14 potassium sulfate

K + SO4 2-

15 aluminum bicarbonate

Al 3+ HCO3 -

16 aluminum bisulfate

Al 3+ HSO4 -

17 calcium hydroxide

Ca 2+ OH -

18 potassium chlorate

K + ClO3-

19 magnesium chlorate

Mg 2+ ClO3 -

20 sodium nitride

Na + N 3-

21 sodium nitrite

Na + NO2 -

22 sodium nitrate

Na + NO3 -

23 lithium sulfide

Li + S 2-

24 lithium sulfite

Li + SO3 2-

25 lithium sulfate

Li + SO4 2-

26 lithium hydroxide

Li + OH -

27 aluminum hydroxide Al 3+ OH -

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Name of cation Name of anion Formula of

cation formula of anion

formula of compound

number of ions

28 aluminum carbonate

Al3+ CO32-

29 strontium carbonate

Sr 2+ CO3 2-

30 strontium nitrate

Sr 2+ NO3 -

31 strontium nitrite

Sr 2+ NO2 -

32 strontium nitride

Sr 2+ N 3-

33 sodium acetate

Na + C2H3O2 -

34 sodium oxalate

Na + C2O4 2-

35 barium oxalate

Ba 2+ C2O4 2-

36 aluminum phosphide

Al 3+ P 3-

37 aluminum phosphate

Al 3+ PO4 3-

38 ammonium iodide

NH4 + I -

39 ammonium phosphate

NH4 + PO4 3-

40 ammonium acetate

NH4 + C2H3O2 -

41 ammonium oxalate

NH4 + C2O4 2-

42 ammonium sulfite

NH4 + SO3 2-

43 ammonium bicarbonate

NH4 + HCO3 -

44 strontium bisulfate

Sr 2+ HSO4 -

45 silver chloride Ag + Cl -

46 silver sulfate Ag + SO4 2-

47 zinc bromide Zn 2+ Br -

48 zinc phosphate Zn 2+ PO4 3-

49 cobalt (II) chloride Co 2+ Cl -

50 cobalt (III) chloride Co 3+ Cl -

51 nickel (II) bromide Ni 2+ Br -

52 nickel (III) perchlorate Ni 3+ ClO4 -

53 iron (II) nitrate Fe 2+ NO3 -

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54 copper (I) carbonate Cu + CO3 2-

Name of cation Name of anion Formula of cation

Formula of anion

Formula of compound

number of ions

55 copper (II) carbonate Cu 2+ CO3 2-

56 tin (II) hydroxide Sn 2+ OH -

57 tin (IV) hydroxide Sn 4+ OH -

58 gold (I) bromide Au +

Br -

59 gold (III) bromide Au 3+ Br -

60 lead (II) sulfite Pb 2+ SO3 2-

61 lead (IV) sulfite Pb 4+ SO3 2-

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Practice: Give the name for each of the following ionic compounds: Formula of

compound Name of compound

1) NaCl

2) CaCl2

3) AlCl3

4) Ca(NO2)2

5) CaSO4

6) Na3PO4

7) KBr

8) Ca(HCO3)2

9) Mg(HSO4)2

10) Ba(NO3)2

11) NH4Br

12) AlPO4

13) K3PO4

14) K2SO4

15) Al(HCO3)3

16) Al(HSO4)3

17) Ca(OH)2

18) KClO3

19) Mg(ClO3)2

20) Na3N

21) NaNO2

22) NaNO3

23) Li2S

24) Li2SO3

25) Li2SO4

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26) LiOH

27) Al(OH)3

Chemical formula Name of compound 28) Al2(CO3)3

29) SrCO3

30) Sr(NO3)2

31) Sr(NO2)2

32) Sr3N2

33) NaC2H3O2

34) Na2C2O4

35) BaC2O4

36) AlP

37) AlPO4

38) NH4I

39) (NH4)3PO4

40) NH4C2H3O2

41) (NH4)2C2O4

42) (NH4)2SO3

43) NH4HCO3

44) Sr(HSO4)2

45) AgCl

46) Ag2SO4

47) ZnBr2

48) Zn3(PO4)2

49) CoCl2

50) CoCl3

51) NiBr2

52) Ni(ClO4)3

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53) Fe(NO3)2

54) Cu2CO3

55) CuCO3

Formula of compound

Name of compound

56) Sn(OH)2

57) Sn(OH)4

58) AuBr

59) AuBr3

60) PbSO3

61) Pb(SO3)2

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Exercise: Write a chemical formula for each of the following Ionic compounds: 1. Sodium sulfide 1. ____________________ 2. Aluminum oxide 2. ____________________ 3. Nickel (II) chloride 3. ____________________ 4. Nickel (III) chloride 4. ____________________ 5. Cobalt (II) nitride 5. ____________________ 6. Cobalt (III) nitride 6. ____________________ 7. Potassium phosphate 7. ___________________ 8. Iron (III) sulfate 8. ____________________ 9. Calcium carbonate 9. ___________________ 10. Ammonium acetate 10. __________________ Exercise: Give the Name for each of the following Ionic compounds 1. Cu2S 1. ____________________ 2. Mg(ClO2)2 2. ____________________ 3. Zn3N2 3. ___________________ 4. Al(OH)3 4. ___________________ 5. Mn3(PO4)2 5. ___________________ 6. Cu(H2PO4)2 6. ___________________ 7. Co(CN)2 7. __________________ 8. Sn(SO3)2 8. ___________________ 9. (NH4)2CrO4 9. ___________________ 10. Cr(BrO3)2 10. ____________________

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Nomenclature of Binary Covalent Compounds (Nonmetal-Nonmetal) The chemical formula of any covalent compound always starts with a nonmetal a. Keep the name of the first element the same. b. Change the ending of the second element by adding the suffix, ide. c. Indicate the number of atoms of each element in the compound using the following Greek prefixes.

(Memorize the following prefixes)

hexa- = 6 Apply the above Greek prefixes only to binary covalent compounds. Exercise Name the following binary covalent compounds: 1. SO3 1. _______________________ 2. BrF5 2. _______________________ 3. S2Cl7 3. ______________________ 4. NO 4. _______________________ 5. N2S5 5. ______________________ 6. H2O 6. ______________________ (use common name) 7. NH3 7. ______________________

(use common name) Exercise Write formulas for the following covalent compounds: 1. diphosphorus pentoxide 1. _____________________ 2. dinitrogen trioxide 2. _____________________ 3. carbon tetrachloride 3. ______________________ 4. tetraphosphorus decoxide 4. ______________________

Prefix Number

Mono- 1 Di- 2 Tri- 3 Tetra- 4 Penta- 5 Hexa- 6 Hepta- 7 Octa- 8 Ennea-/Nona- 9 Deca- 10

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Nomenclature of Acids Formula starts with a "H". [H2O is excluded]

All acids must be dissolved in water to exhibit their properties

a. Binary Acids

1. HF 1. ____________________

2. HCl 2. ____________________

3. HBr 3. ____________________

4. Hl 4. ___________________

5. H2S 5. __________________

6.* HCN (not binary) 6. hydrocyanic acid_____ *exception

b. Oxy- Acids (Ternary Acids)

Polyatomic ion

Name of ion Formula of acid Name of acid

NO2- nitrite HNO2 nitrous acid

NO3

- nitrate HNO3 nitric acid

SO32-

SO4

2-

PO33-

PO3

3-

ClO -

ClO2 -

ClO3

-

ClO4 -

CO3

2-

C2H3O2-

Lewis Structure of Oxy Acids (Ternary Acids) a) HNO3

b) H2SO4

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Summary: Naming Compounds

Ionic Compounds Molecular Compounds Acids Chemical formula starts with a metal ion Need to use prefixes: mono,di,tri..etc Chemical formula starts with H or NH4

+ ion. Fixed charge metal variable charge metal Binary Acids Oxyacids Example: Na+,Ca2+ Example: Fe2+and Fe3+ Example: HCl Example: H2SO4 Molecular compounds versus ionic compounds: Molecular compounds are made of molecules. Ionic compounds are made of ions

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I. Atomic mass/ Molecular mass/Formula mass: THE MOLE

Avogadro's number

13 Al 26.98 Atomic mass Molar mass 26.98 amu 26.98 g {1 atom} = 1mole of Al atoms = 6.02 x 1023 Al atoms II. MOLAR MASS a. Consider Cl2O7 :

2 atoms Cl 2 moles Cl 7 atoms O 7 moles O = 1 molecule Cl2O7 = 1 mole of Cl2O7 Molecular mass of Cl2O7 =mass of one molecule = _________ amu Molar mass of Cl2O7 = mass of one mole of molecules = mass of 6.02 x 1023 molecules Cl2O7 = __________g b. Consider Na2SO4 2 atoms Na 2 moles Na 1 atoms S 1 moles S 4 atoms O 4 moles O = 1 formula unit Na2SO4 = 1 mole of Na2SO4 Formula mass of Na2SO4 = mass of one unit of Na2SO4 = _________ amu Molar mass of Na2SO4 = mass of one mole of Na2SO4 =mass of 6.02 x 1023 units of Na2SO4 = _________ g

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Exercises 1. How many Na atoms in 3.0 mole Na?

2. How many moles of H are in 44.0 moles of H3PO4? Plan: moles H3PO4 moles H 3. How many grams of O are in 25.45 moles of HClO3? Plan: moles HClO3 moles O mass O 3. How many P atoms are in 73.2 g P2O4? Plan: grams P2O4 moles P2O4 moles P atoms P gram 4. If 2.74x 1024 atoms of an element weigh 277.5 g, find its molar mass. Plan: Need to find g/mole 5. Find the mass of carbon in 25.00 g C4H10O2 Plan: grams C4H10O2 mole C4H10O2 mole C mass C

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III. Percentage Composition Percent by mass = mass of element x 100 Mass of compound 1. Find the percentage composition of Na3PO4

2. Find the percentage composition for Al(NO3)3

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IV. EMPIRICAL FORMULA Empirical formula shows the smallest ratio of atoms in a compound. Examples: Molecular formula = C6H12O6 Empirical formula= CH2O Calculation of Empirical Formula and Molecular Formula Exercise: 1. Determine the empirical formula for a compound that has the following percentage composition: 43.64 % P and 56.36 % O

Step 1. Express the percent as grams Assume 100 g of material. Step 2. Change the grams into moles

Step 3. Divide by the smallest number to obtain ratios as whole numbers. Step 4. If the simplest mole ratio is not a whole number, multiply by a factor

Answer: ____________

b. If the molar mass is determined to be about 288 g/mole, find the molecular formula.

Answer: ____________

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2. Determine the empirical formula for a compound that has the following percentage composition: 29.08 % Na, 40.56 % S, and 30.36% O 3. A 0.8640 g sample of a compound made of C, H, and O is burned (reacted with O2) . The products are 1.727 g CO2 and 0.7068 g H2O only. Find the empirical formula of the compound.

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CHEMISTRY 110

100 POINTS SHOW ALL YOUR WORK. YOUR ANSWERS MUST HAVE THE CORRECT NUMBER OF SIGNIFICANT FIGURES AND UNITS. CORRECT SPELLING MUST BE USED. __________________________________________________________ 1. COMPLETE THE FOLLOWING TABLE: ASym Number of Number of Number of Mass number Z protons neutrons electrons ------------------------------------------------------------------- ______ 16 _______ 18 32

56Fe ________ _______ _______ _______ 2. Give the isotope symbol for the following: a. An anion of nitrogen with the same number of neutrons as oxygen-15 b. An iron cation with the same charge and number of subatomic particles as 58Co2+ 3. Calculate the molar mass of Na3PO4

4. How many grams of H are there in 3.0 x 1025 molecules of H2SO4? 5. Name or give the chemical formula for the following:.

oxalic acid magnesium hydrogen carbonate _____________________________________ _____________________________________ mercurous nitride ammonium carbonate _____________________________________ _____________________________________ silver nitrate aurous iodide _____________________________________ _____________________________________ plumbic acetate iodine tribromide _____________________________________ _____________________________________ calcium peroxide hydrobromic acid _____________________________________ _____________________________________ potassium phosphide sulfurous acid _____________________________________ _____________________________________ nickelous permangante cobaltous sulfide _____________________________________ _____________________________________

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CS2 Co2O3 _____________________________________ _____________________________________ Ni(NO2)2 Bi(NO3)3 _____________________________________ _____________________________________ Ba3N2 HClO3(aq) _____________________________________ _____________________________________ Ca(OH)2 N2O5 _____________________________________ _____________________________________ Sr(HSO3)2 Hg(HCO3)2 _____________________________________ _____________________________________ H2CO3(aq) PbO2 _____________________________________ _____________________________________ SO3 _____________________________________ _____________________________________ HF HBrO2(aq) _____________________________________ _____________________________________ HC2H3O2(aq) Au3PO4 _____________________________________ _____________________________________ N2O3 Cu(Cl0)2 _____________________________________ _____________________________________ HCN(aq) Al(OH)3 _____________________________________ _____________________________________ KH _____________________________________ _____________________________________

6. The percentage composition of a compound is 63.133% C, 8.831% H, and 28.04% O.

The Molar mass = 171.21 g/mol

a. What is its empirical formula?

b. What is its molecular formula?

7. The chemical formula of DDT is C14H9Cl5. In a 0.750 gram sample:

a. How many moles of C14H9Cl5 are present?

b. How many grams of carbon are present?

c. What is the total number of atoms present?

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d. What is the percent hydrogen in C14H9Cl5?

8. How many grams of Na has the same number of atoms as 13.0 g N?

9. 1.450 moles of element Y weighs 0.30044 kg. a. What is the molar mass of Y?

b. What element is this?

10. For the following questions , identify the element whose atoms fit the following descriptions (Use the periodic table)

a. ____ Which has d electrons?

a) hydrogen b) Copper c) nitrogen

b. ____ A metalloid in period 5

c. ____ The element in period 4, group IIIA

d. ____ The element with a total of 3 electrons in the 2nd main energy shell

e. ____ The smallest alkali metal

11. a. Write the electron configuration of a iodine atom. b. Write the electron configuration of tin 12. Write the electron dot structure for the following compounds

H2SO3(aq) NO3-

REMEMBER TO DO THE PROBLEMS IN THE TEXTBOOK