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Unformatted text preview: Chapter 5 Types of Bons to form Compounds
4.1 Valence e’
e’ in outermost shell are called valence electrons
Ca (20 e’) 1s22s22p63s23p64s2 2 valence e’ S (16 e’) 1s22s22p63s23p4 6 valence e’ Ar (20 e’) 1s22s22p63s23p6 8 valence e’ The group number refers to the valence e’ Electron Dot StructurePut the number of valence e’ around the atomic symbol. i
L N Be O B F C e
N 1 4.2 Octect Rule
Noble gasses are extra stable because they have eight valence e’. The p orbitals are filled.
All other elements want to achieve the noble gas e’ configuration. Positive ions- 2e Mg 12 p 12 e' 8e 2e 2e 1s22s22p63s2 1s22s22p6
Same as Ne Na loses 1 e’ to become Na+ 1s22s22p63s1 --------- 1s22s22p6 + 1e’ 2 Negative Ions- 2 e' 2e 8e 6e 2e S 16 p 16 e'
Same as Ne N gains 3e’ to become N3- 1s22s22p3 + 3e’ -------- 1s22s22p6 Group number will determine the number of electrons given up or accepted Group electrons
Group 1 2 3 5 6 7 Charge +1 +2 +3 -3 -2 -1 Metals take on positive charges, Non-metals negative charges
Examples H+ Be2+ Al3+ N3- S2- Cl- 3 Ionic Compounds
Positive and negative ions come together to form a compound. Ionic bond. Ex.
K. + :Br: Æ K+ + Br- Æ KBr Æ Mg2+ + Cl- Æ MgCl2 Æ Al3+ + Cl- Æ AlCl3 . ..
Ca: + :F:
. Write subscripts when doing ionic compounds they do not change. Two wheels per bike.
Nomenclature1. Monoatomic IonsA. Positive ions (cations)
Name of element followed by ion
Na+ sodium ion
Ca2+ Calcium ion B. Transition elements (cations)
4 Name of element followed by a Roman numeral that corresponds to the charge.
Cu+ copper(I) ion Fe2+ Iron(II) ion Cu2+ copper(II) ion Fe3+ copper(III) ion Some transition elements just have one charge: Zn2+, Ag+
C. Negative ions (anions)
Name of element + ide followed by ion
Cl- chloride ion
S2- Sulfide ion 2. Ionic Compounds
metal + non-metal –ide
KCl potassium chloride MgBr2 Magnesium bromide Al2O3 aluminum oxide Fe2O3 Iron (III) oxide copper (I) nitride Cu3N sodium phosphide silver (I) maybe sulfide Ag2S Iron (II) nitride Fe3N2 Na3P 5 Balance out the charges so everything is neutral. 3. Covalent Compounds- sharing of e’
Look at Cl each atom want one more e’, neither wants to form a positive charge. l
C + l
C Diatomic- X2, other diatomic molecules Group 7 (F2,Cl2, Br2)and N2, O2 H2 Look at NH3 H
N + H H N H H
H Covalent bonds- Non-metal, non-metal 6 Group Number Valence e’ Covalent bonds
(neutral molecule) 1 1 1 4 4 4 5 5 3 6 6 2 7 7 1 Must full fill octet
O H + H O H H C + Cl
C C l
C Can have multiple bonds Steps for writing electron-dot structures 7 1. Determine the central atom. Place other atoms around it.
2. Determine the total valence e’.
3. Place a single bond (2 e’) between all atoms.
4. Subtract the bonding e’ from total.
5. Full fill octects of all atoms with remaining e’. SO3
O 1. S 3.
O 2. 4 x 6 = 24 4. S O 24 - 6 = 18
O S O 8 For Charged species add an e’ or take one away from total. O
N 2- O 1. N 3.
N 2. 2 x 6 = 12
18 e' 4. O 18 - 4 = 14 5. -1 O N O Exceptions to octet1. Free radicals. Ex NO, NO2, ClO2, unpaired e' ,Cl.
2. Too few e'.
a. Boron can have 6e'. BF3
b. Be has 4e' Ex BeF2.
c. Too many e', P= 10e', S = 12e', EX. PCl5, SF6 Nomenclature of Covalent compounds.
Need tom know the difference between NO, NO2, N2O
Prefix- element Prefix- element -ide 9 Atoms Prefix 1 Mono- 2 Di- 3 Tri- 4 Tetra- 5 Penta- 6 Hexa- 7 Hepta- 8 Octa- Exceptions1. Often the mono is emitted when it is the first element.
CO carbon monoxide CO2 carbon dioxide 2. When o and o or a and o appear together the first vowel is omitted.
N2O4 dinitrogen tetroxide Ex.
SiO2 Silicon dioxide diphosphorus pentoxide P2O5 SO2 sulfur dioxide dinitrogen oxide 4. Bond Polarity 10 e’ can be shared equally – non-polar bond
or unequally- polar bond
Depends on a the electronegativity of the elements envolved.
Electronegativity is the ability of an atom to attract shared e’.
F has the highest at 4.0 then O, and N. Look at periodic table trends.
Have a difference in electronegativity in a bond the bond is a polar covalent bond. Non-symmetrical e’ cloud of the bond or the entire molecule. d+ d- d+ d- C-N H-S Polar Polar d+ dF-F Si-O C-H K-Cl Non-polar Polar non-polar Ionic
difference bigger than
1.6 Predicting bond typeNonmetal- Nonmetal -Covalent- Usually different atoms polar covalent bond.
Metal- Nonmetal- Ionic Don’t use polar non-polar 6. Polyatomic Ions 11 Many atoms coming together to form a charge (usually with oxygen, and
Look at NO2Main ones (on sheet)
SO42- - sulfate PO43- - phosphate NO3- - nitrate CO32- - carbonate OH- -hydroxide NH4+ - ammonium A. Nomenclatue of Polyatomic Ions 1. Contain one less oxygen end in –ite
SO32- - sulfite PO33- - phosphite NO2- - nitrite CO22- - carbonite 2. Add a hydrogen (H+) add: (still an ion)
1 hydrogen- add hydrogen
HSO4- - hydrogen sulfate HPO42- - hydogen phosphate 2 hydrogens only phosphate put dihydrogen
H2PO4- - dihydogen phosphate 12 B. Ionic Compounds from Polyatomic Ions
metal + Polyatomic anion
Ca2+ + CO32- ---- CaCO3 Calcium carbonate Na+ + NO2- ---- NaNO2 Sodium nitrite K+ + HSO4- ---- KHSO4 Potassium hydrogensulphate Calcium Nitrate Ca(NO3)2 Use parenthesis Magnesium Phosphite Mg3(PO3)2 Ammonium hydroxide NH4OH 5. VSPER theory.
7. Three kinds of intermolecular forces can be traced to electrostatic attractions;
1. Dipole forces: attraction between polar molecules
Example: I-Cl is a covalent; both non-metals. Attraction is the positive pole of
one molecule and the negative pole another.
2. Induced Dipole forces: attraction between substances with non-polar molecules. 13 Dependent upon size of e’ cloud; the greater the cloud, the more it can distort it
Result of shifting electron clouds with the molecule.
Example: C-H bonds are relatively non-polar.
Also called dispersion forces, London forces & London dispersion forces.
3. Hydrogen bonds: H-atom bonded to an atom that is small and highly
electronegative and has at least one unshared pair of electrons. Nitrogen, Oxygen
and Flourine are the only elements that satisfy this rule.
H-bond is 1/10th that of a covalent bond. 14 ...
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