Chemical Bonding!

Chemical bonding only involves the valence e-

Formed when e- attracted by nucleus of another

Electrostatic force- opposite charge attract; like charge repel

- greater distance, smaller attraction

- greater charge, greater force.

- the force operates equally in all directions


Intermolecular force=force between molecules

Intramolecular force=force within a molecule

Only weak bonds break in melting process



Chemical bonds exhibit various degrees of sharing, depending on the amount of attraction.



Polarity=molecule's electrical balance, but imbalance with electrical charge of elements



Ionic=electrons are transferred

Non polar covalent=shared equally

- Electronegativity difference <0.5

- Attain full electron shells

- Simultaneously attracted

- high melting points(except CH4, O2, F2)

Polar covalent= shared unequally

- higher energy=partial negative charge=δ-

- lower energy=partial positive charge=δ+

- arrow indicates the migration of electrons

Example

Organic Chemistry May 26, 2011

The chemistry of carbon compounds
.

Importance
-organic compounds are seen anywhere in our life.
-Examples of organic compounds: sugar, chair.......

Properties of organic compounds.
-low melting point
-weak or non-electrolytes
-can forms chains of carbon atoms that are linked in a
①straight-line

 

②circular pattern

③branched pattern

    -can link with other atoms in:

    single bonds

  double bonds

  triple bonds

*Versatility organic compounds makes it such an important branch of chemistry.

Alkanes (unbranched/straight chain)

-A hydrogen : only contains O,H

There are different types of hydrogen

                             ways to present them

-no-polar molecules → immiscible with water

-geometry: tetrahedron

-alkanes are saturated alkanes( all C atoms are bonded by single bonds)

*Saturated: not possible for another atom to bond to the structure.

-Naming of alkanes: the names of all hydrocarbons end in "-ane"

Ex. The structure of C2H6?

①          

     full structure                     

②CH3CH2CH3  

Condensed structure

③C3H8

molecular formula

④

 

Ball-and-Stick model

Names of Alkanes.

methane   CH4

ethane      C2H6

propane   C3H8

butane      C4H10

pentane    C5H12

hexane     C6H14

heptane    C7H16

octane      C8H18

nonane     C9H20

decane     C10H22

They are homologous series: a series of organic compounds with similar general formula, possess similar properties.

CnH(2n+2) -----only for alkanes.

Branched hydrocarbons

hydrocarbons have "side branches" which are also hydrocarbon chains.

(substituted carbon/ branched carbon)

Ex

* Alkyl group: an alkane which has lost one hydrogen atom.

The bottom part are alkyl group(CH3)

 

The name of this branced hydrocarbon is 2-methylpropane

*"-pronpane" is the parent hydrocarbon(longest chain)

Carbon has 4 bonds.

  10 hydrogens

Naming: the names of all alkyl groups end in "-yl", because they are alkyl.

Ex. C2H6      ethyl  C2H5

      C3H8      propyl C3H7

Electron Dot & Lewis Diagrams

- nucleus is represented by the atomic symbol

- for individual elements, # of valence electrons

- electrons are represented by dots around the symbol

- a orbit hold a max of 2 electrons

- each orbital gains 1 electron before pair up

Each bond represents 2 electrons

All valence electron must be used

Each element must have a full valence orbital except H

Ex. CH4

Period Table Trend

Metallic Properties:
- the change from metallic to non-metallic going from left to right across the table
- elements become more metallic going down a family

Atomic Radius:
- Decreases across a row from left to right, increase down a group

Reactivity:
- metal&non-metal show different trend
- most metal: Francium
- most non-metal: Fluorine


Ion charge:
- elements ion charges depend on their group


Melting&Boiling point:
- centre of the table-->highest melting/boiling point
- noble gases have lowest melting point

Ionization Energy
- increases up and right
- hellium highest
- Francium lowest

Electronegativity:
- fluorine has highest electronegativity

- Francium has lowest electronegativity

Predicting the # of valence electron

Valence  electron:  electron which can take part chemical reactions also called the " reactable electron" in the outermost(energy level) open electron shell of atom.
-open shell: shell contains less than max. # of electron
-closed shell: contains exactly max. # of electron

valence electron: not in the core, not it d- & f- subshells

ex.
Al=[Ne]3s2 3p1  has 3 valence electron
Pb=[Xe]6s2  4f14  5d10  6p2  has 4 valence electron
Xe=[kr]5s2  4d10  5p6  or[Xe] has 0 valence electron (noble gases)

The periodic law summmarizes the period table
-The periodic law: properties of the chemical elements recur periodically when the elements are arranged from lowest to highest number.
            

ELECTRON CONFIGURATION!!

-the elctronic configuration of an atom is notation that describes the orbitals in which the electrons occupy and the total number of electrons in eaqch orbital.

-when an electron absorbs or emits a specific amount of energy it instantaneously moves from one orbital to another.

-an energy level is the amount of energy, which an electron in an atom can possess.

-energy difference between two energy level is quantom of energy.

Ground State: when all the electrons of an atom are in their lowest possible energy level.

Excited State: when one or more of an atom's electrons are in higher energy level than the lowest electron.

Orbital: an orbital is the actual region of space occupied by an electron in a particular energy level.

Shell: a shell is the set of all orbitals having the same n-value.

Subshell: a subshell is the set of orbitals of the same type.

The letters s, p, d, f refer to the four different types of orbitals, each circle represents an orbital.

For a given value of "n", different types of orbitals are possible for:
n=1 only the s- type
n=2 s- and p- types
n=3 s- , p- and d- types
n=4 s- , p- , d- and f- types.

A maximum of 2 electrons can be placed in each orbit. (Pauli Exclusion Principle)

Writing Electronic Configuration for Neutral Atoms
-always start with the lowest energy level.
-figure out how many electrons you have, then start at the lowest (1s) and keeping adding.
-each electron has an opposite spin designated by upward and downward arrows.

ex. C has 6 electrons, 2 in the 1s, 2 in the 2s and 2 in the 2p.
notice the 2 electrons in the 2p pccupy separate suborbitals and are not paired up.
the electron configuration is 1s2 2s2 2p2.

Writing Electronic Configuration for Ions
for negative ion:
-add electron to the unfilled subshell, starting where the neutral atom left off.

for positive ion:
-start with neutral atom, remove electron from the outermost shell first.
-in both s- and p- , remove from p-, first.

Core Notation
the set of electrons for an atom can be divided into two subsets: the core electrons and the outer electron.
-the core of an atom is the set of electron with the configuration of the nearest noble gas that comes before it.
-the outer electron consist of all electron outside the core electron normally take part in chemcal reactions.

locate the atom and note the noble gas at the end of the row ABOVE the element.

ex. S: 1s2 2s2 2p6 3s2 3p4
          [NE] 3s2 3p4

 

Atomic Structure~

proton: has a relative mass of 1 and electric charge of +1; located in the nucleus.
neutron: has a relative mass of 1 and no electric charge; located in the nucleus.
electron: has a relative mass of almost 0, has a electric charge of -1; located in the cloud surrounding the nucleus.


Atomic number=the proton number


Atoms have no overall electric charge. Therefore, # of p+ = # of e-
anion: when e- is added to an atom
cation: when e- is removed from a atom.


Mass number= # of proton + # of neutron. (therefore, # of neutron = mass # - # of proton)
Atomic mass= the average mass of a element's isotopes


Isotope: have same atomic number but different atomic mass (in another word, different # of neutron)