Chem Test 3 A

Front 1

Saturation

Back 1

Maximum amount of solute that can be dissolved in a solution at equalibrium. This amount is called solubility

Front 2

Unsaturated

Back 2

Contains less than the maximum amount which can be dissolved.

Front 3

Supersaturated

Back 3

Contains more than the maximum amount which can be dissolved at equalibrium.

Front 4

Electrolytes

Back 4

Are substances whose water solutions will conduct an electric current because they dissociate into ions in solution.

Can produce electricity and conduct electricity.

Front 5

Strong electrolytes

Back 5

Substances which are essentially completely dissociated in aqueous solution.

Completely separate in solution

Front 6

Weak electrolytes

Back 6

Dissociated to a relativley small extent in solution. And produce a few ions the solution is thus much smaller conductor than a strong electrolyte.

Front 7

Nonelectrolytes

Back 7

are substances which produce nonconducting solutions because they do not dissociate into ions in water. Examples: Sugars, alcohols

Front 8

Arrhenius definition of acid and base

Back 8

Acid- substance produces hydrogen ions in solution

(Acid anhydride

Generally non Metallic oxides which reacts with water to give acidic solutions)

Base- substance that produces hydroxide ions in solution

(Base anhydrides generally metal oxides which react with water gives basic solutions)

Front 9

Amphoteric substances

Back 9

Can function as either an acid or a base. In water they do not disolve and behaves as neither acids nor bases.

This indicates that not all metal oxides or hydroxide are basic in water.

Front 10

Bronsted Lowry definition

Back 10

Acid: proton donor

Base: proton acceptor

Proton = H+ ion

Conjugate acid base theory

Acid That gives an H+ will yield conjugate base.

Base that will accept an H+ will yield conjugate acid.

Front 11

Lewis definition

Back 11

Acid: electron pair acceptor

Base: electron pair donor

Front 12

Electro negativity of O

Back 12

Each lone O will pull an electron from the center element allowing the H to leave more readily

Acidity increases

Front 13

Pg 216

Back 13

A

Front 14

Bases in water

OH

NH

Back 14

1 bases are strong (100% ionized)

2 bases are weak (few % ionized)

Front 15

Salts

Back 15

Ionic compounds containing cations and anions

Strong acid and base = neutral _____

Strong acid weak base = acidic _____

Weak acid strong base= basic _____

Front 16

Oxidation

Back 16

Loss of electron

Increased oxidation state

Front 17

Reduction

Back 17

Gains electron

Decreasing in oxidation state.

Front 18

Oxidation states

Back 18

0 Free elements (Cu Na O2 Cl2)

1+ Hydrogen with nonmetals

(1- if with metals hydrides)

2- Oxygen in most compounds

(1- in peroxides H2O2)

1- Halides (simple salts NaCl cacl2)

1+ Group IA alkali metals

2+ Group IIA alkaline earth metals

Front 19

Oxidation states of metals

Back 19

A moving down a representative group of metal the lower oxidation state is more stable

E.G. Al prefers 3+

while TI prefers 1+

Front 20

Transition (B) groups

Back 20

(1) Lose s electrons first

(2) lose any number of d electrons next

Front 21

Balancing oxidation and reduction reaction step 1

Back 21

Assign oxidation states to all reactants and products

Front 22

Balancing oxidation and reduction reaction step 2

Back 22

Decide what is oxidized and what is reduced

Front 23

Balancing oxidation and reduction reaction step 3

Back 23

Mass balance

Pull out the species involved in each half reaction and balance each half reaction with respect to atoms (mass)

First balance the atoms oxidized and reduced then:

If acidic

Add water to the side needing an O for balancing

Ad H+ to the opposite side to balance the Hs

If basic

Add OH- to the side needing O (normally double)

Add water to the opposite side to balance Hs

Front 24

Balancing oxidation and reduction reaction step 4

Back 24

Charge balance

Balance with respect to charge using electrons.

Check see that the oxidation loses the appropriate number of electrons and the reduction gains the appropriate number of electrons.

Front 25

Equivalents

Back 25

Number of H+,OH-, or e- gained or lost, per mole.

Front 26

Eq wt

For wt

Back 26

g/eq

g/mole

Front 27

Normality

Back 27

EQ solute/L Soln

N=EQ/V

EQ=NV

Front 28

Molarity

Back 28

Moles solute/L soln

M=MOLES/V

MOLES=MV

Front 29

Titration equations

Back 29

Eq acids=Eq base

Eq oxidized = EQ reduced

Eq=NV (eq/L ×L= Eq)

EQ= wt/eq wt (g/(g/eq) = Eq)

Front 30

Balancing oxidation and reduction reaction step 5

Back 30

Balance so that when we add half reaction the electrons will cancel from opposite sides of the equation

Front 31

Ionic (electrovalent)bond

Back 31

Occurs when elements on the left electro positive metals meet those on the right electronegative nonmetallic if electronegative difference on the Pauling scale is greater than 1.7 this will form

Face centered cubic

Front 32

Covalent bond

Back 32

Sharing of electrons formed when both atoms are on the right (electro negative nonmetals) and has a small difference in electronegativities
Unequal sharing

Front 33

Ionic bond

Back 33

C

Front 34

Metal bonds to nometals

Back 34

Forms ionic salt
Metals conduct electricity
High melting point

Front 35

Electrostatic attraction

Back 35

Example NaCl
Na atom Losses of an electron (+)
Cl gains the electron (–)

Front 36

Water and ionic compounds

Back 36

Water breaks bonds and electrolyte are free and can now conduct electricity.

Front 37

Vesper theory

Back 37

like charges repell
Valance shell like charges 180 degrees apart.

Front 38

Lewis structures

Back 38

A

Front 39

Poly atomic ions cation(s)

Back 39

Ammonium
NH4+ only positive

Front 40

Polyatomic ions anion(s)

Back 40

Nitrate NO3^–
Sulfate SO4^(2–)
Phosphate PO4^(3–)

Front 41

Allotropes

Back 41

Different forms of the same element as a result of the differences in molecular structure such as O2 and O3 or different packing of atoms in a solid.

Front 42

Sp Hybridization

Back 42

2 paired electrons
Orbital geometry
Linear

Front 43

Sp^2 hybridization

Back 43

3 pairs of electronsOrbital geometry Triangular planar

Front 44

Sp^3 hybridization

Back 44

4 sets of paired electrons

Orbital geometry Tetrahedral

Front 45

Dsp^3 hybridization

Back 45

5 paired electronsOrbital geometry Triangular bipyramidal

Front 46

D^2sp^3 hybridization

Back 46

6 electron pairs

Orbital geometry
Octahedral

Front 47

Factors influencing formation of bonds

Back 47

Energy must be put into system (endothermic) more energy in than out

Front 48

Coordinate covalent bond

Back 48

When both electrons shared come from the same atom.

Front 49

Polar covalent bond

Back 49

Any difference in electronegativity will result in an unequal sharing of electrons

Front 50

Polar and nonpolar shape

Back 50

Dipole cancel in p usually linear
Dipoles do not cancel in np will not be symmetrical in all positions

Front 51

Hydrogen bonding

Back 51

An extreme polarization when the smallest most electronegative element (N,O,F) attach to H.
Strong bond responsible for water being liquid at room temp.
And for water to be polar enough to disolve ionic salts.

Front 52

unpolarized

Back 52

Ionic potential equals q/r (low)
More ionic in nature
Downward trend of periodic table
More soluble I'm H2O
More basic
Less color
Higher melting point
Less cation hydrolysis
(S orbitals)

Front 53

Polarized

Back 53

Ionic potential equals q/r (high)
More covalent in nature
Trends down periodic table
Less soluble in water
More acidic
More color
Lower melting point
More cation hydrolysis

Front 54

Pg 182

Back 54

Examples

Front 55

Effect of anion on the character of compounds
Polarization (distortion) increases with

Back 55

1 increase in size
2 increase in charge
Rows right to left
Groups top to bottom

Front 56

Valence bond theory

Back 56

Explains how a bond forms between 2 neighboring atoms using ovelapping one atomic orbital from each yielding a new wave form which holds two electrons of opposite spin.
The degree of the overlap is due to the strength of the covalent bond.

Front 57

Molecular Orbital (MO)

Back 57

Theory considers all of the atomic orbitals from both atoms and shows how these interact to form orbitals on the resulting molecule called molecular orbitals Even orbitals which are not involved in the bond are described. Two types of covalent bonds are formed:
sigma head on head overlap
and pi off internuclear axis

Front 58

Hybridization of d orbitals

Back 58

From the third period down on down we get hybridization ro include an aditional available (d) orbital. These structures will exceed the octet rule.

Front 59

Multiple bonds n2

Back 59

Triple bond
1 sigma
009999652627351861618

Front 60

Resonance

Back 60

Is the mathematical combination of two or more valence bond structures that are equivalent

Front 61

Electron pair repulsion theory

Back 61

Principle
Pairs of electrons in the valence shell(whether bonded or nonbonded) are always arranged to allow the greatest spacial separation. The arrangement depends only on number of electron pairs.