Pre-AP Chemistry NOTES 9-1 VALENCE ELECTRONS ......When the electron is transferred, oppositely-charged ions are formed, which are then attracted to each other by electrostatic forces:

Pre-AP Chemistry NOTES 9-1

VALENCE ELECTRONS AND THE FORMATION OF IONS

There is a driving force among atoms to gain or lose electrons in order to reach a low-energy, stable state. This stable state is obtained when an atom has an electron configuration of s2p6 in its outermost energy level. The atom is then said to have a full outer shell of electrons. For example: Argon – Krypton – Both of these elements belong to what group on the periodic table?__________________________________ Noble gases (with the exception of helium) have a stable electron configuration (s2p6). This type of configuration, therefore, is called a noble gas configuration. Why is helium stable, even though it doesn’t have an s2p6 configuration? Why is it virtually impossible to get a noble gas to react or bond with anything? VALENCE ELECTRONS The total number of electrons present in the final s-p configuration is very important. These are the electrons which are involved when atoms bond with each other. They are known as valence electrons – the electrons in the outermost energy level (s and p sublevels) involved in bonding. The number of valence electrons (for the representative elements) can be determined by noting the number of the group the element is in: Barium has __________________ valence electrons Carbon has __________________ valence electrons Iodine has ___________________ valence electrons

A Lewis Dot Structure (also called an “electron dot structure”) is a method used to show the number of valence electrons in an atom:

EXAMPLES: sodium beryllium nitrogen oxygen argon helium FORMATION OF IONS Octet Rule – atoms react by changing the number of their electrons so as to acquire the stable electron structure of a noble gas (s2p6 – 8 valence electrons) In doing this the atoms form ions (positively or negatively charged atoms) EXAMPLES: sodium 1s22s22p63s1 lose 1 electron [Na+] sodium ion aluminum fluorine oxygen

FORMATION OF IONS

symbol

number of valence

electrons

Lewis dot structure

number of electrons lost or

gained

charge

ion symbol

Ba

S

N

K

C

Ne

He

ELECTRON CONFIGURATION OF IONS (Extra Credit) Metals lose electrons to achieve noble gas configuration and become cations (positively-charged ions). Nonmetals usually gain electrons to become stable, forming anions (negatively-charged ions). S2- Na+ Fe3+ Zn2+


CHEMICAL PERIODICITY

FOCUS ON FLUORINE

(Fluorine is the “Rudy” of the periodic table: Fluorine is little but strong)

Atomic Radius - the radius of the atom

Atomic radius decreases as you move across the periodic table and increases as you move down

Which has the larger atomic radius? chlorine sulfur

Which has the larger atomic radius? chlorine bromine

Ionization Energy - the amount of energy required to remove an electron from a gaseous atom

In general, ionization energy increases as you move across the periodic table and decreases as you move down

Which has the larger ionization energy? oxygen nitrogen

Which has the larger ionization energy? oxygen sulfur

Electronegativity-the attraction an atom has for electrons when it is chemically bonded to another atom

In general, electronegativity increases as you move across the periodic table and decreases as you move down

Which has the larger electronegativity? magnesium sodium

Which has the larger electronegativity? potassium sodium


CHEMICAL BONDING: METALLIC AND IONIC BONDS

Metallic Bond – consists of closely-packed cations surrounded by free-floating valence electrons This type of bond is formed between metals only (either pure or alloys):

Why are metals such good conductors of electricity? Why are metals ductile (capable of being drawn into a thin wire) and malleable (capable of being hammered into thin sheets)? Ionic Bond – a bond formed when electrons are transferred between atoms with largely different electronegativity values

Sodium wants to lose and electron, while chlorine wants to gain an electron:

When the electron is transferred, oppositely-charged ions are formed, which are then attracted to each other by electrostatic forces:

The ions are then held together in lattice structures by electromagnetic attractive forces:

Ionic compounds do not have molecules

CHARACTERISTICS OF IONIC COMPOUNDS:

*Most are crystalline solids with high melting points

*Many are soluble in polar solvents (such as water) and insoluble in nonpolar solvents

*The molten form of the compound conducts electricity (it contains mobile charged particles) –

however solids cannot conduct electricity (the charged particles cannot move)

*The aqueous solutions of ionic compounds also conduct electricity well (mobile charged particles)

*They tend to be brittle and shatter when hammered


CHEMICAL BONDING: COVALENT BONDS Covalent Bond – a bond formed when two atoms of similar electronegativity values (usually nonmetals) share one or more pairs of electrons.

Neither atom will give up an electron, so they share:

The two shared electrons are attracted to orbitals in both fluorine atoms and this holds the atoms together. The atoms do this to attain noble gas configuration:

Covalent compounds are made up of discrete molecules, rather than lattice structures.

Lewis Dot Structures can be used to represent the covalent compound formed: *The “dash” refers to a shared pair of electrons Double Covalent Bond (Double Bond) – made up of 2 shared pairs of electrons Triple Covalent Bond (Triple Bond) – made up of 3 shared pairs of electrons LEWIS STRUCTURES OF MOLECULES:

The structure of a molecule can be determined using the Valence Shell Electron Pair Repulsion Theory

(VSEPR). The electron geometry of the molecule can be determined using the following guidelines:

1. Add the total number of valence electrons for all atoms in the molecular formula to determine the number of pairs of electrons. (If the species is charged, add one electron for each negative charge and subtract one electron for each positive charge.) 2. Arrange the peripheral atoms around the central atom, placing one electron pair (represented by a “dash”) between each. Subtract the total number of electrons used from the original total. 3. Use the remaining electrons to fulfill the octet rule (eight electrons in the outer shell) on each peripheral atom. (Don’t fill octets for Groups IA, IIA, or IIIA – these don’t get octets.) Subtract the electron pairs used. 4. If any electron pairs still remain, place them on the central atom.

5. Check to be sure that all atoms fulfill their octet. If some atoms do not, then take some of the electrons from the peripheral atoms and form multiple bonds with the central atom. (NOTE: Only C,N,O,P,S form multiple bonds.) 6. Elements with atomic numbers lower than carbon are full with few electrons (ie. beryllium is full with 2 pairs of electrons, boron is full with 3 pairs of electrons, etc.) 7. Nonmetals from VA on up (beginning with phosphorus) can have expanded octets (more than 4 pairs of electrons around the central atom. The “d” orbitals in these atoms are available to form these types of structures. Without the “d” orbitals, atoms are too small to fit more than 4 electron pairs around them.

EXAMPLES: CCl4 PH3 H2O ClO3

- NH4

+ NaCN CO2 NO3

-

Watch for exceptions to the rule: BF3 PCl5 SF6 ICl2

- XeF4


POLARITY OF BONDS AND MOLECULES

BOND POLARITY

A polar bond forms when two atoms of different electronegativity values bond together covalently. This

results in pair of electrons that is shared “un-equally” between the two atoms. The diagram below represents

a hydrogen atom (on the left) bonded with a fluorine atom (on the right):

hydrogen fluorine

Since fluorine in the above example has a higher electronegativity value (3.98) than hydrogen (2.20), their

shared electron pair is displaced more toward fluorine, resulting in a partial negative “pole” on the fluorine

end and a partial positive pole on the hydrogen end – hence, a polar bond.

Molecules may be polar or non-polar, depending upon their molecular structure. Water, for instance, has an

asymmetrical structure that consists of two hydrogen atoms and two “lone pairs” of electrons in a tetrahedral

orientation around a single oxygen atom. This gives the water molecule an uneven distribution of electric

charge, resulting in a molecule that has a partial positive pole and a partial negative pole – a polar molecule:

Molecules like carbon tetrachloride consist of polar bonds. However, the symmetrical structure of the

molecule results in an even distribution of charge – the molecule as a whole is nonpolar in nature:

In general, molecules are nonpolar if *there are no lone pairs on the central atom and *all peripheral atoms are the same ELECTRONEGATIVITY AND TYPES OF CHEMICAL BONDING The following scale shows that the type of chemical bonding is determined by electronegativity differences:

Electronegativity Difference

Type of Bond

0 – 0.39

nonpolar covalent

0.40 – 0.99

moderately polar covalent

1.00 – 1.99

highly polar covalent

> 2.00

ionic

EXAMPLES:

molecule

electronegativity difference

type of bond

HF

NaCl

Br2

NO


MOLECULAR GEOMETRY

TYPES OF COVALENT BONDS *Single Bond – consists of 1 pair of shared electrons forming a sigma bond (σ) (2 “p” orbitals overlap “head-to-head”) *Double Bond – consists of 2 pairs of shared electrons *1 sigma bond *1 pi bond (π) – formed when 2 “p” orbitals overlap “side-to-side” *Triple Bond – consists of 3 pairs of shared electrons *1 sigma bond 2 pi bonds

RELATED FACTS CONCERNING COVALENT BOND TYPES 1. A single sigma bond is stronger than a pi bond 2. The triple bond is the shortest of the bonds and is stronger than the rest 3. The single bond is the longest of the bonds and is weaker than the rest *A single bond has a bond order of “1” *A double bond has a bond order of “2” *A triple bond has a bond order of “3” 4. Bond Dissociation Energy – the energy required to break a chemical bond (also referred to as “bond energy” for short); increases with bond order

The molecular geometry is determined by the number of electron pairs on the central atom, and whether these electron pairs are bonding pairs or lone pairs (non-bonding pairs).

The electron pair geometry can be determined using the table below:

TOTAL ELECTRON PAIRS ON CENTRAL ATOM

ELECTRON PAIR OR STRUCTURAL GEOMETRY

HYBRIDIZATION

2 linear sp

3 trigonal planar sp2

4 tetrahedral sp3

5 trigonal bipyramidal sp3d

6 octahedral sp3d2

MOLECULAR GEOMETRIES

Linear

Trigonal Planar

Tetrahedral

Pyramidal

Bent

Trigonal

Bipyramidal

See-Saw

T-Shaped

Linear

Octahedral

Square Pyramidal

Square Planar

Websites to view molecular geometries:

http://intro.chem.okstate.edu/1314F97/Chapter9/VSEPR.html

http://www.chemmybear.com/shapes.html

EXAMPLES:

Molecule

Lewis Structure Electron Pair Geometry

Molecular Geometry Polar or Nonpolar?

BeCl2

BF3

CH4

NH3

H2O

http://intro.chem.okstate.edu/1314F97/Chapter9/VSEPR.html

http://www.chemmybear.com/shapes.html

Molecule

Lewis Structure Electron Pair Geometry

Molecular Geometry Polar or Nonpolar?

PCl5

SeF4

BrI3

ICl2-

SCl6

ClF5

XeF4


INTERMOLECULAR ATTRACTIVE FORCES

In addition to the covalent bonds that exist between atoms in a molecule (H2O for instance), there are also weak attractions between the molecules themselves. They are responsible (among other things) for whether a compound is a liquid, solid, or gas at room temperature. 1. London Dispersion Forces (AKA Induced dipole-Induced dipole, van der Waal’s Forces) Caused by the attraction of the positively charged nucleus of one atom for the electron cloud of an atom in a nearby molecule, inducing a temporary dipole in neighboring atoms/molecules.

*The weakest of the intermolecular attractive forces *All covalently bonded substaces are attracted by London dispersion forces *Bigger atoms / molecules have larger London dispersion forces and are therefore attracted more strongly (their large electron clouds are not attracted as strongly to their own nuclei, so they are more easily polarizable For example: The bigger the molecule, the higher the boiling point (note the halogens) 2. Dipole-Dipole Interactions - Caused by the attraction of polar molecules to each other

*Similar to ionic bonds, but much weaker since only partial charges are involved

*Occur only in polar molecules

3. Hydrogen Bonding Occurs when a hydrogen atom that is covalently bonded to a very electronegative element (F, O, N) is attracted to a lone electron pair on another atom in a nearby molecule *The strongest of the intermolecular attractive forces (~5% the strength of a covalent bond) *Explains many of the “odd” characteristics of water PROPERTIES OF COVALENT (MOLECULAR) COMPOUNDS 1. Composed of nonmetals 2. Solid, liquid, or gas at room temperature – based, in general, on molecular weight 3. Low melting point . . . Why? 4. Poor or non-conductors of electricity in any state (solid, molten, or aqueous) . . . Why?

EXAMPLE: Determine the predominant IMF present in the following. Place an “X” through any substance that does not have intermolecular attractive forces. LiBr PF3 CF4 CH3OH RULES:

1. Determine if it is ionic or molecular. (Ionic substances do not have IMF’s)

2. Determine if it has hydrogen bonding (H + F,O,N)

3. Draw the rest of the structures: Polar = “Dipole-Dipole Interactions” Non-Polar = “London Dispersion Forces”

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Pre-AP Chemistry NOTES 9-1 VALENCE ELECTRONS ......When the electron is transferred, oppositely-charged ions are formed, which are then attracted to each other by electrostatic forces: