bonding unit

chemical bond

• involves the electrons only

• nuclei of the different atoms are not affected

Quantum Theory

• electrically neutral (# of protons = # of electrons)

• acquire the more stable electron structure of the nearest Noble Gas

orbital

region of space where electrons are most likely to be found

non-valance electrons

in the inner energy levels

valance electrons

• in the outer shell

• paired or single

lone pair

• two electrons (pair)

• repelling effect on electrons in any nearby orbitals

• reduce the bond angle by 2.5°

• closer to central atom - more repulsion than bonded pairs

bonding electron

• single electron

• share that electron with another atom

Octet Rule

• obeyed by Main Group atoms

• max. eight electrons can occupy orbitals in the valence energy level to become like noble gasses

Pauli Exclusion Principle

• two electrons may share the same region of space (never more)

• occupy empty valence orbitals before forming electron pairs

Chemical Formulae

• structure of atoms

• identify element

.

Molecular Formula - do not reduce

.

empirical formula - reduce

.

Lewis Structure

.

Structural Formula (Kekulé Structure)

.

Perspective Drawing (Stereochemistry)

.

Ball and Stick Model

rules for placing electrons

• One electron dot is placed in each of the four valence orbitals before electron pairing

• If there are five to eight valence electrons, a second electron with a bonding electron

ionic compound

• transfer of electrons

• electronegative atom - anion

• less electronegative atom - cation

naming inorganic compounds

• ionic or molecular

• ionic - how many ions does the metal form (eg. Cu (ii))

• monatomic or polyatomic

NH^₄+

make ionic compounds due to positive charge

name → formula

right to left

formula → name

left to right

ionic compound shape

• solid at room temperature

• form crystal (ionic) lattice

lewis structure for ionic compounds

• draw lewis structure of both atoms

• draw half-headed arrow to show transfer

• add more atoms to balance the charges

• draw square brackets around ions and add charges (lost/gained)

lattice

• maximize forces of attraction

• shape of ionic compounds

• unit cell defined by the lattice points

• array of alternating positive and negative charges

properties of ionic compounds

• lattice

• high MP and BP

• non-conductive as solid

• conductive as molten or in aqueous solution

high MP and BP

• strong attraction between cation and anion

• hard to seperate

non-conductive as a solid

cation and anion are fixed in the solid state - no movment

conductive as molten/in solution

cation and anions are free to move

brittle

• when knocked out of place, positive aligns with positive, negative aligns with negative

• repulsion occurs, causing a fracture

intermolecular forces

• the forces that exist between molecules

• London dispersion

• dipole-dipole

• hydrogen bond

intramolecular forces

• forces that hold atoms together within molecules

• covalent bonds

• ionic bonds

london dispersion

• temporary

• non-polar

• determined by the number of electrons in the molecule

• low boiling point - weak attraction

dipole-dipole

• permanent

• polar

• stronger dipole = stronger force

• attracts positive end with negative end

• higher boilng point

hydrogen bond

• permanent

• N, O, F is bonded to H

• strongest of all IMFs

• high boiling point - strong attraction, hard to break

covalent bonds

• atoms share valence electrons

• non-metal + non-metal

• non-polar/polar

• ΔEN less than 1.7

ionic bonds

• transfer of electrons

• ΔEN more than 1.7

• metal (cation) + non-metal (anion)

bond strength order (strongest to weakest)

• metallic

• ionic

• molecular

• hydrogen bonding

• dipole-dipole

• london dispersion

Van der Waals’ Forces

• define attraction of intermolecular forces

• London dispersion

• Dipole-dipole

LDF causes

• LDFs are in all atom, ion, and molecule

• temporary shift, attraction between electron and proton, off when shifts back

boiling point increase

• size increase

• more electrons, more surface area

strength of LDF

• number of electrons

• shape and size

more electrons

stronger LDF

shape and size

• affect how close the molecules approach each other in solid and liquid states

• closer molecule = stronger attraction/repulsion

graph

• top - hydrogen bonding - highest boiling point due to strong attraction

• drop - only LDFs - lowest boiling point due to weak attraction

• increasing - LDFs increase attraction with an increase in size (more electrons)

electrostatic attraction

negative side of the dipole bonds with the unshielded proton

hydrogen bonding requirements

• first molecule has hydrogen attatched to a highly electronegative atom (N, O, F) - hydrogen bond donor

• second molecule has a lone pair on a small highly electronegative atom (N, O, F) - hydrogen bond acceptor

properties of hydrogen bonded substances

• increased melting/boiling point

• increase solubility - like disolves like

• shape and stability

key points of IMF

• boiling points are a measure of IMF

• IMF increase with polarization (seperation of positive/negative charges) of bonds

• boiling point increases with molar mass/surface area

• strength of IMF - LDF < DDF < hydrogen bonding < covalent < ionic < metallic

molecular

related to/has molecules

polarity

• contrasted properties/direction

• having poles

• opposite powers in opposite parts/directions

• being opposite

dipole moments

• separation of charge

• large difference in electronegativity = large dipole moment

• ionic/covalent bond

• measure polarity

• electrons unequally shared

• one atom is more electronegative than another

shape and polarity of bonds

• determines overall polarity

• asymmetrical = polar

• symmetrical = non-polar

determinations for polar molecules

• charges cannot cancel out

• unequal distribution of charge

• may include lone pairs

• may have different types of bonds

• asymmetrical

• different peripheral atoms

Electron Domain

region around the central atom of a molecule where electrons are concentrated

Bonding Domains

• single, double, or triple bonds connected to the central atom

• every bond counted as only one electron domain

Non-Bonding Domains

lone pairs connected to the central atom

.

Triganol planar

.

linear

.

tetrahedral

.

Triganol pyramidal

.

bent

VSEPR Theory

• shape of an atom

• due to the repulsions between electron pairs in the valence shell

• depends on number of bonding groups and non-bonding electrons

number of electron pairs

determined by writing the Lewis structure

strength of repulsion (highest to lowest)

1. lone pair - lone pair

2. lone pair - bonded pair

3. bonded pair - bonded pair

AX_mE_n

• A = central atom

• X = bonded/peripheral atom

• E = lone pairs

• m and n = number of lone pairs or bonded atoms

Electronegativity

ability of an atom to attract electrons in a chemical bond

down a group

• Atomic radius increases (bigger)

• Increase in nuclear charge is too small to be important

• More shielding - more shells

• electronegativity decreases

• Less attraction between nucleus and bonding electrons

shielding effect

• block

• inner electrons shields outer electrons from attractive force of the nucleus (protons)

going across a period

• atom becomes smaller

• nuclear charge increases

• atomic radius decreases

• greater attraction between nucleus and bonding electron

• electronegativity increases

nuclear charge

greater than repulsion charge

smaller atom

• more electrons being put in the same energy level

• less shielding

sharing

not equal

non polar covalent electronegativity

• no pull

• equal

• ΔEN - ≤ 0.4

polar covalent electronegativity

• higher electronegativity

• pulls electron towards it

• 0.4 < ΔEN > 1.7

ionic bond electronegativity

ΔEN - 1.7 <

polar covalent

• between two extremes

• electron sharing unequal

• electron density greater around more electronegative atom

• different non-metals - different electronegativity

nonpolar covalent

• same or different non-metals that have same/similar electronegativity

• electron shared equally

• purely covalent

Electrical Charge

• deficiency/excess of electrons in a chemical species

• result is a net positive charge or net negative charge

Molecular Polarity

• electrical charge not symmetrically distributed among the atoms in a molecule

• partial positive and partial negative charges on opposite ends of the molecule

𝛿

• lower case delta

• partial charge

Metallic Bond

sharing of free electrons among a structure of positively charged ions

No Bond

• Noble Gases have a full outer valence shell

• non-reactive

theory of covalent bonding

• fill their valence orbitals (except H, He, Be, and B)

• bonding is attraction of a pair of shared electrons

• Unpaired electrons can bond

• Two atoms with bonding electrons form covalent bond

• The number of covalent bonds that each atom forms is limited (fill valence shell)

limitation for octet rule

• first period - only two electrons

• B and Be - do not need full valance shell

• odd number of valence electrons

• third period < exceed rule

Noble Gases

• similar properties

• low reactivity

• don’t form bonds

halogens

• highly reactive

• need one more electron

Nomenclature

• naming chemical compounds

• easily identified as separate chemicals

covalent nomenclature

• least electronegative element first

• second element - suffix to –ide

• USE prefixes

one

mono

two

di

three

tri

four

tetra

five

penta

six

hexa

seven

hepta

eight

octa

nine

nona

ten

deca