Chapter 6 Manipulating Polyatomic Ions and Chemical Bonding.

Chapter 6

Manipulating Polyatomic Ions and Chemical Bonding

Basic Polyatomics

Name Formula Name Formula

Hydroxide OH-1 Nitrate NO3-1

Ammonium NH4+1 Phosphate PO4

-3

Acetate C2H3O2-1 Chromate CrO4

-2

Carbonate CO3-2 Dichromate Cr2O7

-2

Sulfate SO4-2 Chlorate ClO3

-1

Ways to expand your polyatomics

Polyatomic ions vary in their charges, number of oxygen

atoms, and number of hydrogen atoms.

1. To change the number of oxygens:One more oxygen ClO4

-1 perchlorate

Memorized ClO3-1 chlorate

One less oxygen ClO2-1 chlorite

Two less oxygens ClO-1 hypochlorite


2. Family Members Whatever is true for chlorine, is also

true for fluorine, bromine, and iodine. Memorized ClO3

-1 chlorate

F substitution FO3-1 fluorate

Br substitution BrO3-1 bromate

I substitution IO3-1 iodate


3. If you add a hydrogen, you have to make the ion more positive (one less negative) and call the ion “bi________”

Memorized CO3-2 - carbonate

HCO3-1 – bicarbonate

Memorized SO4-2 - sulfate

HSO4-1 - bisulfate


4. Combinations of #1, #2, and #3 are possible:HSO3

-1 is called bisulfite

FO2-1 is called fluorite

Rules for expanding your list of polyatomic ions

Rule #1

To Change the number of oxygens: Remove one oxygen = change ending of name to –

ite

Remove two oxygens = change ending of name to –ite and beginning of name to Hypo-

Add one oxygen = change beginning of name to Per-


Name Formula Name Formula Name Formula Name Formula

ClO4-1 SO5

-2 NO4-1 PO5

-3

ClO3-1 SO4

-2 NO3-1 PO4

-3

ClO2-1 SO3

-2 NO2-1 PO3

-3

ClO-1 SO2-2 NO-1 PO2

-3

Examples:


Rule #2 Other Family Members Elements near each other in the same column tend to

form similar polyatomic ions.

Name Formula Name Formula Name Formula

Chlorate ClO3-1 Sulfate SO4

-2 Phosphate PO4-3

Fluorate FO3-1 Selenate SeO4

-2 Arsenate AsO4-3

Iodate IO3-1

Bromate BrO3-1


Rule #3 Add a hydrogen Add only one H = change the beginning of the name to

bi- and make the charge one less negative (due to hydrogen’s positive one charge)

Name Formula Name Formula

Carbonate Sulfate

Bicarbonate Bisulfate


Rule #4 Combinations of 1, 2, & 3 Combinations of #1, #2, and #3 are possible:

HSO3-1

Memorized SO4-2 Sulfate

Lose an “O” SO3-2 Sulfite

Add an “H” HSO3-1 Bisulfite

HFO2 Memorized ClO3

-1 Chlorate Substitute an F FO3

-1 Fluorate Lose an “O” FO2

-1 Fluorite Add an “H” HFO2

Bifluorite


Combos Cont’dEx1: What is the formula for hypoiodite?

Find I on the P-table (near Cl). Chlorine forms chlorate (ClO3

-1). Thus, Iodine forms iodate (IO3

-1). The –ite and hypo- in hypoiodite mean that iodate lost two oxygens.

Hypoiodite = IO-1


Combos Cont’d Ex2: What is the formula for Biperselenate?

Find Se on the periodic table. It is near S. Sulfur forms sulfate (SO4

-2). Therefore, selenium forms selenate (SeO4

-2). The per- in biperselenate means that selenate has gained one oxygen. Also, the bi- means that it has gained a hydrogen (don’t forget to change the charge!).

Biperselenate = HSeO5-1

Monatomic Ions

For nonmetals, almost all single names that end with –ide indicates a single charged atom.

Simply write the symbol and the charge. The periodic table column indirectly indicates the element’s charge. Remember, elements want to have 8 electrons in their outer shell (Octet Rule).

Monatomic Ions

Column #1 elements have a +1 charge Column #2 elements have a +2 charge Column #3 = +3 Column #15 = -3 Column #16 = -2 Column #17 = -1

Monatomic Ions

Ex1: What is the formula for chloride? Cl-1

Ex2: What is the formula for an aluminum ion? Al+3

Ex3: What is the name of the S-2 anion? Sulfide

Ex4: What is the name of the Mg+2 cation? Magnesium Ion

6.1 Introduction to chemical bonding

Most elements are not found alone in nature. They

are “stuck” to other atoms. Chemical Bond - Link between atoms that results

from the mutual attraction of their nuclei for their electrons.

Types of chemical bonds: Ionic - transfer of electrons (metal + nonmetal) Covalent - sharing of electrons (2 nonmetals) Metallic - happens in metals when there is only one

type of element

Introduction to Chemical Bonding

Covalent bonds may be polar or nonpolar Polar - unequal sharing of electrons (HCl)

Nonpolar - equal sharing of electrons (H2)

There are two ways to predict polar vs. nonpolar ( and covalent vs. ionic)

Introduction to Chemical Bonding

#1 Use electronegativity difference 0 = nonpolar covalent 0.4 - 1.7 = polar covalent greater than 1.7 = ionic

Examples NaCl Cl = 3.16 HCl Cl = 3.16 Na= - 0.93 H = - 2.20 2.23 .96 Ionic Polar Covalent

Cl2 Cl = 3.16 Cl = - 3.16 0 Nonpolar Covalent

Introduction to chemical bonding

#2 - There is an easier way to predict Ionic = metal + nonmetal

Polar Covalent = 2 different nonmetals

Nonpolar Covalent = 2 of the same nonmetals

Ionic Bonds

Ionic compound - a substance composed of positive and neg. ions so that the charges are equal. It involves a transfer of electrons.

Ca+2 with Cl–1 will form the compound CaCl2. It takes two chlorine ions to cancel out the the +2

charge on the calcium ion. Formula unit - lowest whole # ratio of ions Ionic Bond = a METAL + a NONMETAL

Metals - lose e- - why? low IE NM - gain electrons - why? high electronegativity

Ionic Bonds

Metals lose electrons until they become like a noble gas . (8 valence e-)

Nonmetals gain e- until they do the same.

Both go to s2p6 - 8 valence e- - called a stable octet The tendency to arrange e- so each atom has 8 is

called the octet rule or rule of 8

Ionic Bonds

The formation of an ionic bond:Na to Cl = [Na]+1[Cl]-1

Na 1s 2s 2p 3s

Cl

1s 2s 2p 3s 3p

Ionic Bonds

Ionic bonding picture:Ex1: Na to Cl= [Na]+1[ Cl ]-1

Na Cl

Ex2: Ba to Cl = [Ba]+2 2[ Cl ]-1

Ba Cl

Cl

Ionic Bonds

The easy way:Find the charge of each atom “criss cross” the charges – charge cancels

out and you are left with a neutral compound

EX3: Al NEX4: Na SEX5: Al S

Ionic Bonds

A few more examples not in your note packet using Polyatomic Ions

USE PARENTHESIS

Li and NO3-1

Ca and C2H3O2-1

Magnesium PhosphiteAluminum hyponitrite

Ionic Bonds

Energy is involved in all chemical reactions. Na + Cl yields NaCl + 769 kJ

Lattice energy - energy released when an ionic compound

forms.

NaCl = - 769 kJ/mole NaF = - 922 kJ/mole

KCl = -718 kJ/mole

smaller ions have higher lattice energies

Ionic Bonds

Properties of ionic compounds:HardShatter Conduct electricityHigh melting pointOdorless

6.2 Covalent Bonding

In covalent bonding atoms share electrons. In the H2 molecule, each H atom says, "I only need one more e- to be like a noble gas (helium)." Since each hydrogen has only one electron, when two hydrogens bond they can share their electrons.

Covalent Bonding

Molecule - smallest quantity of matter that exists by itself and retains the properties of that substance. Describes a covalently bonded substance.

monatomic molecules - He, Ne, Ar, (noble gases are always monatomic)

diatomic molecules – H2 O2 N2 Cl2 Br2 I2 F2 (you must memorize these!!)

polyatomic molecules - P4, S8, C6H12O6

Covalent Bonding

The formation of a covalent bond

Bond Length vs. Bond Energy Bond length = Bond Energy

=

Covalent Bonding

Diatomic Molecules and Orbital Notation (Orbital overlap or notation diagrams):

H2 O2

H 1s O 1s 2s 2p

H 1s O 1s 2s 2p

N2

N 1s 2s 2p

N 1s 2s 2p

Covalent Bonding

Why are these atoms forming bonds? Octet Rule- Atoms lose, gain, or share electrons to

have 8 electrons in their outer shell.

HF – orbital notation

H 1s

F 1s 2s 2p

Lewis Dot Diagrams of molecules (covalent

compounds) and polyatomic ions

Basic rules Each atom wants 8 electrons (except H wants 2). Each atom goes for close to the right # of bonds. The least electronegative atoms goes in the middle

OR The atom that makes the most bonds goes in the

middle. (H always on the outside.) OR The “single guy” (the atom that does not have a

subscript after it) goes in the middle. Symmetry is key!!! Place the atoms in order (left, right, bottom, and

top) around a central atom

Lewis Dot Diagrams

To determine the number of bonds: S = N – A S = shared electrons (bonds)

2 N = needed e- (all elements need 8 except for H which needs 2)

A = how many e- an atom actually has (# of valence e-)

If using S=N-A, add the charge into the Actual amount of electrons.

Put [ ] around the molecule and include the charge

Lewis Dot Diagrams

Examples: Draw the following Lewis structures EX1: CH4

Ex2: H2O

Ex3: PCl3

Lewis Dot Diagrams

Ex4: SiH2F2

Ex5: CS2

Ex6: C2H6

Lewis Dot Diagrams

Ex7: C2H4

Ex8: C2H2

Ex9: CH2O

Lewis Dot Diagrams

Ex10: HCN

Ex11: FON

Drawing polyatomic ions

count electrons: if the charge is - 3, add 3 electrons to A EX: PO4

-3

less bonds than atoms want = negative charge more bonds than atoms want = positive charge

P wants 3 bonds, has 4: + 1 charge Each O wants 2, has 1 so each O = -1 Total = - 3 Ex11: PO4

-3

Coordinate covalent bond

Coordinate covalent bond- 2 shared electrons in a bond are donated by 1 atom Examples:

NH4+

OH-1

sulfate nitrate nitrite carbonate bicarbonate H2SO4

H3PO4

6.4 Metallic Bonding - “Sea of electrons theory”

The nuclei are arranged in a systematic lattice.

The bond strength relies on the nuclear charge and the number of valence e- Ex. Mg is stronger than Na

The valence electrons form a sea of free moving electrons that are attracted to multiple positive nuclei.

Metallic Bonding

Conducts Electricity as a result of free electrons.

Malleability and ductility results from the nuclei's ability to move passed each other

Metallic Bonding

Remember: in ionic bonds some atoms want e- and some don’t in covalent bonds, all atoms share – in metals, no one atom wants the e-

6.5 The Properties of Molecular Compounds

Valence shell electron pair repulsion theory (VESPER) – e- pairs get as far away from each other as possible Because of this we can predict the shape of molecules

based on how many bonds and lone pairs are on the central atom

Shapes

Example SharedPairs

LonePairs

Shape Example Angle(s) and drawing

AB 1 0 Linear HCl 180o

AB2 2 0 Linear CO2180o

Shapes

Example SharedPairs

LonePairs

Shape Example Drawing

AB2E 2 1 Bent SO2 119.5o

AB2E2 2 2 Bent H2O

H2S104.5o

AB3 3 0 TriangularPlanar

CH2O 120o

Shapes

Example SharedPairs

LonePairs


AB3E 3 1 Triangularpyramidal

NH3 107.5o

AB4 4 0 Tetrahedron CH4 109.5o

Shapes

Example SharedPairs

LonePairs


AB5 5 0 TriangularBipyramidal

PCl5 90o

1200

1800

AB6 6 0 Octahedron SF6 900

1800

Shapes

Examples: Predict the shapes of the following (show all work):

Ex1: CCl4 Ex2: HBr Ex3: SO3

Ex4: SO2

Ex5: H2S Ex6: NH3

Ex7: ClO4 -1

Ex8: PF5

Intermolecular Forces

Intermolecular forces (IMF)- forces that hold molecules together happens between covalent compounds intermolecular forces - can be weak or strong

Intramolecular forces – chemical bonds (ionic, covalent, metallic) happens within a molecule or compound always strong

H -------------Cl H -------------Cl


Types of intermolecular forces dipole-dipole

dipole - when electrons are unevenly distributed Ex1: predict the IMF that occurs with HCl

Ex2: predict the IMF that occurs with H2O (***one of the most impt. Ever!)


hydrogen bonding - H-bonding is a “super-duper” dipole-dipole

H-bonding happens any time H is bonded to F, O, or N Hydrogen bond is FON!!! Why does this happen?

A large difference in electronegativity between F, O, or N and H results in one

end of the molecule being very negative, while the other end is very positive.


Effect of H-bonds on physical properties:H-bonding tends to cause the following in

substances: Boiling Point Heat of Vaporization Vapor Pressure Melting Point H-bonds causes water to expand when it freezes. H-bonding is also responsible for the shapes of

proteins.


Recap: Polar molecules have dipole-dipole IMF holding

them together. H-bonding is “super” dipole-dipole IMF These two types of IMF usually result in substances

being solids or liquids at room temp. Most nonpolar covalent substances are gases at

room temp. as the forces holding them together are not strong enough to keep the molecules attracted - hence they are gases

O2, H2, N2 - straight nonpolar substances CO2 - have dipoles, but nonpolar due to its

molecular geometry


Van der Waals Forces (London Forces) Temporarily induced dipoles caused by the motion of

electrons. more electrons = more attraction so, bigger atoms have

stronger Van der Waals forces Occur with noble gases

Bonding

Bond Energy -what is the strength of chemical bonds? bond energy - energy needed to break a bond -

measured in kJ/mole bond strength and stability:

stronger bond - more stable -needs more energy to break the bond

weaker bond - takes little energy to break the bond so the chemical is unstable

chemical changes favor lower energy states - exothermic reactions are favored

Bonding

Bond Strength which is stronger? - single, double, or triple

bond? Triplewhich is shortest bond length? s, d, or t?

triple which is stronger, short or long

bonds? short

Chapter 6 Manipulating Polyatomic Ions and Chemical Bonding.

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Transcript of Chapter 6 Manipulating Polyatomic Ions and Chemical Bonding.